Dynamics of surfaces and reaction kinetics in heterogeneous catalysis: proceedings of the international symposium, Antwerp, Belgium, September 15-17, 1997

Studies in Surface Science and Catalysis 109 DYNAMICS OF SURFACES AND REACTION KINETICS IN HETEROGENEOUS CATALYSIS Stu...

Author: G.F. Froment | K.C. Waugh

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Studies in Surface Science and Catalysis 109 DYNAMICS OF SURFACES AND REACTION KINETICS IN HETEROGENEOUS CATALYSIS

Studies in Surface Science and Catalysis 109 DYNAMICS OF SURFACES AND REACTION KINETICS IN HETEROGENEOUS CATALYSIS

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

Vol. 109

DYNAMICS OF SURFACES AND REACTION KINETICS IN HETEROGENEOUS CATALYSIS Proceedings of the International Symposium, Antwerp, Belgium, September 15-17, 1997 Editors

G.E Froment

Laboratorium voor Petrochemische Techniek Universiteit Gent, Gent, Belgium

K,C.Waugh

Department of Chemistry, UMIST Manchester, UK

1997 ELSEVIER Amsterdam

m Lausanne

m New

York m Oxford m Shannon--

Singapore--

Tokyo

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

ISBN 0-444-82609-2 91997 Elsevier Science B.V. All rights reserved. No part of this publication may be reproduced, stored in a retrieval system or transmitted in any form or by any means, electronic, mechanical, photocopying, recording or otherwise, without the prior written permission of the publisher, Elsevier Science B.V., Copyright & Permissions Department, P.O. Box 521, 1000 AM Amsterdam, The Netherlands. Special regulations for readers in the U.S.A.- This publication has been registered with the Copyright Clearance Center Inc. (CCC), 222 Rosewood Drive, Danvers, MA 01923. Information can be obtained from the CCC about conditions under which photocopies of parts of this publication may be made in the U.S.A. All other copyright questions, including photocopying outside of the U.S.A., should be referred to the copyright owner, Elsevier Science B.V., unless otherwise specified. No responsibility is assumed by the publisher for any injury and/or damage to persons or property as a matter of products liability, negligence or otherwise, or from any use or operation of any methods, products, instructions or ideas contained in the material herein. This book is printed on acid-free paper. Printed in The Netherlands

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Preface

xi

Catalytic surface reaction pathways and energetics from first principles M. Neurock Molecular studies of the mobility of surface metal atoms and adsorbates during catalytic reactions G.A. Somorjai and G. Rupprechter

35

Molecular kinetics of heterogeneous catalytic reactions R.A. van Santen, A. v.d. Runstraat and R.J. Gelten

61

From supersonic beams and single crystal microcalorimetry to the control of catalytic reactions D.A. King

79

Dynamic phenomena at the oxide/water interface : the interplay of surface charge formation, metal complex adsorption, and dissolution/reprecipitation J.-F. Lambert and M. Che

91

The influence of oxygen poisoning on a multiply promoted iron catalyst used for ammonia synthesis : a temperature-programmed desorption and reaction study F. Rosowski and M. Muhler

111

Importance of dynamics in real catalyst systems H. Tops~e, C.V. Ovesen, B.S. Ciausen, N.-Y. TopsCe, P.E. Hejlund Nielsen, E. T6mqvist and J.K. N~rskov

121

Reaction kinetics as a basis for optimal transient operation of catalytic reactors Y. Sh. Matros, G.A. Bunimovich and V.O. Strots

141

vi

THEORETICAL AND EXPERIMENTAL STUDIES ON THE DYNAMICS OF SURFACES

A probabilistic model for the deactivation of a dual function catalyst by coke formation accounting for reaction and surface migration S. Sh~gh and G.F. Froment

159

Self-sustained isothermal oscillations in N20 decomposition on Cu overexchanged ZSM-5 P. Ciambelli, A. Di Benedetto, E. Garufi, R. Pirone and G. Russo

175

Adsorption characteristics of pyridine bases on zeolite (010) examined by atomic force microscopy (AFM) M. Komiyama

185

Transient and steady-state studies of the effect of water on cobalt Fischer-Tropsch catalysts K.F. Hanssen, E.A. Blekkan, D. Schanke and A. Holmen

193

Modelling of the dynamics of complex catalytic phenomena based on surface mobility processes and the remote control mechanism P. Ruiz, Y.-W. LL E. Gaigneaux and B. Delmon

203

Adsorption and reactions of methane on ferric molybdate using DRIFTS technique S. Fuangfoo, A. S. Chellappa and D. S. Viswanath

217

Dynamics of multi-component adsorption with interactions : a mean - field approach M. Dubel and S.D. Prasad

227

Models of adsorption kinetics on rough surfaces M. Giona and A. Adrover

241

Investigation of the structure sensitivity of nitrogen adsorption on single crystal ruthenium clusters using density functional theory D.J. Dooling and L.J. Broadbelt

251

CA TAL YSIS AND KINETICS

Oxydehydrogenation of propane on NiMoO, catalyst under transient and steady-state conditions S. Pietrzyk, M.L. Ould Mohamed Mahmoud, T. Rembeczky, R. Bechara, M. Czernicki and N. Fatah

263

vii

Catalytic ignition during methane oxidation on platinum : experiments and modelling G. Veser, J. Frauhammer, L.D. Schmidt and G. Eigenberger

273

Modelling catalytic cracking kinetics using estimated adsorption equilibrium constants B. Sowerby and S.J. Becket

285

Model discrimination for reactions with stop-effect S. Golay, O. Wolfrath, R. Doepper and A. Renken

295

Methanol oxidation over supported vanadium oxide catalysts : new fundamental insights about oxidation reactions over metal oxide catalysts from transient and steady state kinetics I.E. Wachs, G. Deo, M.V. Juskelis and. B.M. Weckhuysen

305

The effects of alkali promoters on the dynamics of hydrogen chemisorption and syngas reaction kinetics on Ru/SiO2 surfaces D.O. Uner, N. Savargoankar, M. Pruski and T.S. King

315

Interparticle migration of hydrogen on zeolite and their participation in the hydrogenation of adsorbed species and catalytic reaction I. Nakamura, A. Zhang, Y. Fan and K. Fujimoto

325

"State-defining" TAP pulse response experiments J.T. Gleaves, G.S. Yablonskii, P. Phanawadee and Y. Schuurman

333

Transient and steady-state microkinetic models of catalytic reactions on nonuniform surfaces L.J. Broadbelt and J.E. Rekoske

341

Transient kinetics of methane dehydrogenation and aromatisation : experiments and modelling Y. Schuurman, T. Decamp, A. Pantazidis, Y.-D. Xu and C. Mirodatos

351

The desorption of C O 2 from the surface as a kinetically relevant step in the CO oxidation reaction over platinum T.A. Nijhuis, M. Makkee, A.D. van Langeveld and J.A. Moulijn

361

Non-linear steady-state kinetics of complex catalytic reactions : theory and applications G.S. Yablonskii and M.Z. Lazman

371

Neural network based model of the kinetics of catalytic hydrogenation reactions E.J. Molga and K.R. Westerterp

379

Microkinetic analysis of temperature-programmed experiments in a microreactor flow system O. Hinrichsen, F. RosowskL M. Muhler and G.Ertl

389

viii A microkinetic analysis of the reverse water gas shift reaction E. Tserpe and K. C. Waugh

401

TRANSIENT OPERA TION OF REACTIONS AND REACTORS

Selectivity enhancement in consecutive reactions using the pressure swing reactor A.J. Kodde and A. Bliek

419

Experimental studies of transient thermal effects during catalytic oxidation in a packed-bed reactor S. Marengo, P. Comotti, S. Scappatura and M. Vasconi

429

Dynamic operation of trickle bed reactors H.W. Piepers and A.A.H. Drinkenburg

439

Simulation of a catalytic converter of automotive exhaust gas under dynamic conditions A.J.L. Nievergeld, E.R. v. Selow, J.H.B.J. Hoebink and G.B. Matin

449

Effect of variables on the periodic operation of a trickle bed reactor L. Gabarain, J. Cechini and P. Haure

459

Oxidative coupling of toluene under periodic conditions on Pb/Li/MgO : a selective pathway to 1,2-diphenylethane S. Dubuis, M. LorenzL R. Doepper and A. Renken

469

Reduction-oxidation-cycling in a fixed bed reactor with periodic flow reversal H. Seiler and G. Emig

479

Ether decomposition activity of CaNi5 hydrogen storage alloy H. ImaL T. Hosoya and S. Futami

491

Chemical kinetics of a two component phase segregated system. A simple rate model A.A. AI-Haddad and J. Mathew

495

Influence of some phenomena occurring on the surface and in the active phase of the vanadium catalyst on the reactor dynamics K. Gosiewski

511

A detailed kinetic model for the hydrogenolysis, isomerization and dehydrogenation of n-butane I. Mach,~, 7-. Romero and M.M. Rarnl~rez de Agudelo

517

ix Limitation of metal particle size to carbon chain growth in FischerTropsch synthesis Y. Yang, K. Xie and X. Li

523

Network simulation of catalytic cracking reactions C.I.C. Pinheiro, F. Lemos and F. Ram6a Ribeiro

529

Non-catalytic carbon gasification modelling I. Santos Silva, C. Palma, F. Lemos, F. Ram6a Ribeiro and J. Sousa Lobo

535

Effect of solubility parameter on the MTBE synthesis kinetics C. Fit#, J. Tejero, M. Iborra, F. Cunill and J.F. Izquierdo

541

Hydrogen spillover effect over the oxide surfaces in supported nickel catalysts V. Aimasan, T. Gaeumann, M. Lazar, P. Marginean and N. Aldea

547

Transient investigation of the catalytic activity of copper in NO decomposition over Cu-ZSM5 R. Pirone, P. Ciambelli, E. Garufi and G. Russo

553

Thermodynamic transition-state theory and extrathermodynamic correlations for the liquid-phase kinetics of ethanol derived ethers R. Datta, K. Jensen, P. Kitchaiya and T. Zhang

559

Hydrodemetallation kinetics of residual petroleum fractions M.T. Martfnez, M.A. Callejas, E. Carbdand A. Hern~ndez

565

Development of a computational tool for the transient kinetics of complex chemical heterogeneous reaction systems G.A. Carrillo Le Roux, I. Bergault, H. Delmas and X. Joulia

571

Methods of elimination and the problem of nonuniqueness of inverse problem solutions in models of non-stationary chemical kinetics S.Io Spivak and R.M. Asadullin

577

Authors' index

587

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xi

!!i!i!iii!iJi !ii !!iii i!iii!iiiii ii!iiii!iiiiiiiiii iiiii Many processes of the chemical industry are based upon heterogeneous catalysis. Two important items of these processes are the development of the catalyst itself and the design and optimization of the reactor. Both aspects would benefit from rigorous and accurate kinetic modeling, based upon information on the working catalyst gained from classical steady state experimentation, but also from studies using surface science techniques, from quantum chemical calculations providing more insight into possible reaction pathways and from transient experimentation dealing with reactions and reactors. This information is seldom combined into a kinetic model and into a quantitative description of the process. Generally the catalytic aspects are dealt with by chemists and by physicists, while the chemical engineers are callud upon for mechanical aspects of the reactor design and its control. The symposium "Dynamics of Surfaces and Reaction Kinetics in Heterogeneous Catalysis" aims at illustrating a more global and concerted approach through a number of prestigeous Keynote Lectures and severely screened oral and poster presentations. G.F. Froment K.C. Waugh

The International Symposium "Dynamics of Surfaces and Reaction Kinetics in Heterogeneous Catalysis" 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, congresses 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]

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91997 Elsevier Science B. V. All rights reserved. Dynamics of Surfaces and Reaction Kinetics in Heterogeneous Catalysis G.F. Froment and K.C. Waugh, editors

Catalytic Surface Reaction Pathways and Energetics from First Principles Matthew Neurock Department of Chemical Engineering, School of Engineering and Applied Science, University of Virginia, Charlottesville, VA, 22903-2442 Abstract First-principle quantum chemical methods have advanced to the stage where they can now offer qualitative, as well as, quantitative predictions of structure and energetics for adsorbates on surfaces. Cluster and periodic density functional quantum chemical methods are used to analyze chemisorption and catalytic surface reactivity for a series of relevant commercial chemistries. DFT-predicted adsorption and overall reaction energies were found to be within 5 kcal/mol of the experimentally known values for all systems studied. Activation barriers were over-predicted but still within 10 kcal/mol. More specifically we examined the mechanisms and reaction pathways for hydrocarbon C-H bond activation, vinyl acetate synthesis, and ammonia oxidation. Extrinsic phenomena such as substituent effects, bimetallic promotion, and transient surface precursors, are found to alter adsorbate-surface bonding and surface reactivity. I. I n t r o d u c t i o n Heterogeneous catalysis is responsible for the commercial production of nearly 90% of all chemical intermediates. Despite its great importance, our understanding of the elementary mechanisms, the nature of the active sites, and the dynamics of the surface is still rather poor. The ability to "design" optimal catalysts is, therefore, still very much an experimental art. The complexity of the active surface, coupled with the myriad of competing physicochemical steps obscure our ability to elucidate the elementary features that control catalytic properties and reaction mechanisms. Advances in state-of-the-art analytical techniques, fundamental catalytic studies, organometallic cluster chemistry, surface science and theory, however, are beginning to help unravel the nature of the active site and surface. As our understanding of the adsorbatesurface interaction, surface reactivity and catalytic properties matures, our ability to rationally design new materials will improve. Over the past 30 years, surface science has played a leading role in understanding and quantifying adsorbate-surface interactions and surface reactivity on model surfaces. While surface science has yet to lead to the direct discovery of a new catalyst, the fundame~.tal information gleaned from these studies has certainly advanced our understanding of the elementary surface reaction processes. This implicitly contributes to the design of new systems by providing the synthetic-organic and catalytic chemists with an important database regarding the fundamental reactivity of different metal surfaces. This information is subsequently used in the selective identification of possible catalyst precursors. Theory, for the most part, has been used to provide a conceptual understanding of adsorption and reactivity. Much of the previous literature has focused on complementing surface science and organometallic chemistry in the analysis of adsorbates on model clusters and surfaces. The results from these studies have helped to establish the fundamental electronic factors that control surface chemisorption and reactivity. While theory has helped provide a wealth of information on the basic adsorbate-adsorbate and adsorbate-surface interactions, little however, has been achieved in terms of quantitative analyses. For a more indepth analysis of the quantum chemical

applications to heterogeneous catalysis and surface reactivity, the interested reader is referred to a series of reviews by Ruette [1], van Santen and Neurock [2], Whitten [3] and Pacchioni [4]. Advances in the first principles quantum chemistry, novel algorithms, and computer hardware, however, are now beginning to make it possible to compute properties and reactivity on surfaces with much more reliable accuracy. As accuracy and CPU performance continue to improve, the size of the systems that can be examined will increase. Theory will likely play an importam role in future catalysis research in estimating intrinsic energetic properties, elucidating reaction mechanisms and developing structure reactivity relationships. Computation can be combined with experiments to offer a valuable tool for assisting design efforts. In this paper, I review the recent advances and developments of first-principle quantum chemical methods and discuss their application to modelling chemisorption, surface reactivity of reactants/intermediates, and the catalytic behavior for a series of relevant commercial chemistries. We focus primarily on the static representation of the surface.

II. Background A) Formal Chemisorption Theory Much of our current theoretical understanding of surface chemisorption and reactivity can be tied to early pioneering developments by Newns-Anderson [5], Lang and Williams [6], Nr and Luinqvist [7, 8], Hoffmann [9, 10], and van Santen [2, 11-14] in what can broadly be defined as Formal Chemisorption Theory. Chemisorption Theory refers to the formal mathematical analysis of the fragment orbital interactions between the adsorbate and the metal adsorption site. It was born out developments from both the solid-state physics and the chemistry communities. Chemisorption theory elegantly describes the electronic features and molecular orbital interactions that control surface bonding and reactivity for model systems. The classic picture used to illustrate the possible adsorbate-surface interactions that govern chemisorption on a metal surface is shown in Fig. 1 [2, 10, 13]. These include the donation of electrons from the adsorbate to unoccupied metal surface orbitals, the backdonation of electrons from the metal into the lowest energy antibonding orbital of the adsorbate, and Pauli repulsion due to the interaction between tidied metal surface orbitals at the top of the valence band and filled molecular orbitals of the adsorbate. In the adsorption of CO on a transition metal surface [14-17] such as Pt, for example, there is electron donation from the CO 5c~ orbital to a dz2 orbital of the metal surface and backdonation from the dxz, and dyz surface orbitals into the antibonding n* CO orbital [15]. This leads to a strong metal-carbon interaction and a weakened C=O bond. Pauli repulsion weakens the adsorbate-surface interaction. Each of these features: electron donation, electron backdonation and Pauli repulsion, are readily described in terms of Frontier Molecular Orbital (FMO) theory. A number of valuable conceptual ideas were born out of Formal Chemisorption Theory. Most of which are now well understood, and have been used to explain the results from quantum chemical analyses of chemisorption on a number of transition metal surfaces. Formal theory, for example, has been used to identify a series of general trends and concepts relevant to the understanding adsorption phenomena on metal surfaces [2]. These include the following ideas: 1) As you move from left to right across the periodic table the adsorption energy for a stable molecular species, such as ammonia, generally decreases due to the increase in the Fermi level, the f'dling of the metal d-band, and Pauli repulsion. 2) Stable molecular species prefer to sit atop to minimize Pauli repulsion effects. 3) Reaction intermediates with incomplete valence shells prefer to sit at higher coordination sites to increase donative effects.

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.."'-'..Donation

.-" :.

..

.

.

-..

"-44. tl

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i~

Pauli Repulsion

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Fig. 1. The fundamental interactions between the frontier orbitals of an adsorbate and a metal surface band structure that occur upon chemisorption. Formal chemisorption theory has also been used to described a number of other important chemisorption phenomena, such as the stabilizing effects of neighboring electropositive adsorbates (K, Na), the destabilizing effects of electronegative adsorbates (C1, F), surface relaxation and surface reconstruction [18]. More recently, Hammer and Ncrskov [17] applied formal theory to elegantly explain the results from a series of large-scale periodic density functional quantum chemical calculations for adsorption on transition metal and bimetallic surfaces. Chemisorption theory will undoubtedly continue to play an important role in describing relevant concepts in chemisorption and surface reactivity well into the future. More quantitative results from theory, however, will require more sophisticated quantum mechanical methods.

B) Ab Initio Based Methods The ability to quantitatively model transition metal catalyzed surface chemistry relies on two major assumptions. The first is that the quantum chemical method employed is capable of accurately predicting the electronic structure of the system being modeled. The second refers to the choice of chemical model that is chosen to represent this system. Is the model appropriate for simulating the bulk properties of the actual system? Most of previous studies aimed at quantum chemical modelling of chemisorption and reactivity have used finite metal clusters as models of the surface and/or the bulk metal lattice. Quantitative results established from cluster models can be quite sensitive to the model choice (size as well as structural conformation). Special care, therefore, is required in choosing appropriate cluster models. The accuracy of quantum chemical methods has improved considerably over the past few years and is no longer the limiting factor in predicting reliable information. The choice of cluster versus slab, however, is still an open topic of debate and is discussed in detail herein. Most of the initial attempts at the application of ab initio quantum chemistry to predict chemisorption and reactivity on transition metals were met with little success due to the

limitations in the methods that were used. Hartree-Fock molecular orbital theory (HF) was found to yield poor energetic predictions for systems where electron-correlation is important. HF calculations of transition metals, therefore, offer little beyond the qualitative results already determined by simple Extended Htickel theory. More advanced methods for treating electron correlation via many-body perturbation methods, couple-cluster theory, and/or multi-reference wave functions now provide a much higher level of accuracy for transition metal systems. These approaches, however, are computationally intensive and scale anywhere between N5-N 10, where N is the number of basis functions. They have, therefore, only been applied to very small systems. In addition to the limitations in the basic theory, a number of early studies implemented a series of limiting approximations which compromised the quantitative accuracy of the results. Some of these approximations include: 1) small basis sets, 2) one-electron effective core potentials 3) frozen cluster approximation, 4) non-optimized spin-states, and 5) fixed adsorbate-cluster geometries. Most of these early efforts were, therefore, limited to qualitative descriptions only. The tremendous advances that have occurred over the past decade, however, are now making it possible to provide much more reliable fundamental data from f'trst-principles. The development of density functional theory, in particular, has had a major impact on the field. Density functional theory (DFT) is an alternative formalism for solving the energy of an N-electron system, where the electronic interactions, and therefore electron correlation, are explicitly accounted for via a fundamental quantity, the density [19]. Density functional theory, is tied back to the Thomas-Fermi model [20, 21] of the homogeneous electron gas and the Xot theory developed by Slater [22, 23]. It was not officially considered an ab initio approach until the landmark paper by Hohenberg and Kohn [24], which formally proved that the energy of the ground state of an N-electron system is a unique functional of the density. Kohn and Sham [25] later demonstrated that the electron density, and thus the energetics, could be described in terms of single electron operators, and written in a formalism which is quite similar to that for the single electron wave functions adapted in MO theory. The explicit accounting of electron correlation in the DFT approach enables a much higher level of accuracy than simple HF approaches. In many cases DFT provides information that is as accurate as some of the highest level CI methods and costs significantly less in terms of CPU resources. DFT methods scale as N 3, whereas the higher-level CI methods scale between N5-N 10, where N is the number of basis functions. The exchange-correlation functional is, at this stage, one of the limiting features in the theory toward predicting more accurate energetics. One can expect that a great deal of future research in the field of DFT methods will be devoted to developing more accurate exchangecorrelation potentials.

C) Density Functional Quantum Chemistry Many of the initial applications of theory applied toward modelling heterogeneous catalytic systems are reviewed in detail by Ruette [1], van Santen and Neurock [2], Whitten [3] and Pacchioni [4]. The application of DFT to modelling chemisorption began in earnest in the late 1980's. Much of this initial work demonstrated that the approach could be used to predict electronic structure, energetics and properties for adsorbates on small clusters. Most of this work, however, was performed via single-point calculations with no optimization of the geometric structure. Analytic gradients in DFT methods had just begun to emerge in the late 1980's. The reported energies were over-predicted due to a poor approximation of the local spin density as a perfectly smooth electron gas [26]. It wasn't until the early 1990's that nonlocal gradient corrections, spin unrestricted computations, relativistic scalar corrections, and geometry optimization schemes were incorporated into DFT-codes [27-29]. Today, density functional methods are one of the most popular techniques for predicting structural, energetic and spectral properties for organometallic systems. Ziegler presented an excellent review of the status of DFT methods for modelling inorganic systems back in 1992 [26]. Structural predictions for

transition metal-containing species were found to be within 0.05 ~ for bond lengths, 3-4 ~ for bond angles, whereas energetic predictions where within 4-5 kcal/mol of the experimental values. Spectral properties such as infrared and Raman frequencies, XPS and UPS were found to be within about 5% of the measured values. While these benchmarks have improved in recent years due to the introduction of mixed DFT/HF methods and more accurate algorithms, more work still needs to be done toward providing more accurate exchange-correlation potentials. The expectation values reported by Ziegler hold only for well-defined organometallic structures. The application of DFT calculations to metal clusters and surfaces, however, is still an open area. We describe our results from cluster calculations below. Here the issue of accuracy is more closely tied to the cluster approximation rather than on the application of the DFT methods. The development of DFT band structure calculations in the condensed matter physics community have followed a similar course as that described for the molecular codes in chemistry. DFT was an extension of the already established Slater X~ methods. By the 1990's DFT periodic calculations were well-established in providing the electronic structure for the bulk systems. In the periodic-DFT calculations, a unit cell is defined and repeated in either 1-, 2-, or 3-dimensions to represent linear chains, slabs and bulk systems. Wave functions are represented by wavevectors in k-space. Generalized gradient approximations are used to determine the exchangecorrelation potentials. This method is now emerging as a powerful tool for examining adsorption, surface relaxation, and even surface reactivity on transition metal surfaces.

D) Cluster versus Slab Calculations As discussed, two general approaches currently exist for modelling surface chemistry with quantum mechanics: the cluster approach [2, 30-37] and the extended-band surface approach [3841]. Both methods have clear advantages and disadvantages.

1) The Cluster Approach In the cluster approach, the local molecular fragment orbitals are explicit, thus making the local chemical interactions, chemical bonding, and charge transfer mechanisms between the adsorbate and the metal surface orbitals very easy to elucidate [42]. This detailed level of focus, however, makes it difficult to treat the bulk electronic structure. The cluster approach has been criticized for its inability to describe the extended-band electronic structure [42, 43]. Instead of a continuous conduction and valence bands, the cluster approach is based on discrete orbitals which have specific energy gaps. It is well known that there are periodic oscillations in the adsorbate binding energy with metal cluster size for small metal particles [2, 32, 33, 36, 37, 44-46]. This was found in both experiments as well as computations. The computed ionization potential, for example, is typically greater than the measured work function due to the discrete gap between the highest occupied and lowest unoccupied molecular orbitals in the cluster calculations. Previous MO-based ab initio efforts indicate that cluster size oscillations are large, and that one needs to "prepare" the cluster for bonding by exciting an electron in the base cluster prior to binding the adsorbate in order to predict more reliable energetic information [42]. Through firstprinciple DFT calculations, however, we have shown that while these oscillations in binding energy exist, they are substantially lower than predicted by the ab initio MO calculations [2, 35, 36, 47]. Reliably accurate predictions of adsorbate-cluster interactions can be gained through a careful analysis of: 1) the lowest energy cluster conformation and geometry, 2) the determination of the lowest energy spin state, and 3) the optimization of adsorbate-induced structural relaxation [35].

2) The Band Approach The advantage of periodic-DFT methods is that they treat the extended surface and even the 3D bulk, thus providing a more accurate representation of the material's electronic structure. Specific challenges that exist for the extended-band method relate to its difficulty in isolating

local bonding and specific orbital interactions. In addition local adsorbate-surface relaxation, adsorbate optimization and frequency predictions are also still rather difficult to treat at this level. Structural optimization and activation barriers, which are now possible, require substantial CPU resources. Despite these drawbacks, DFT extended band approaches are successfully being used to predict adsorption and surface reactivity for small adsorbates: 02 [48], H2 [49, 50] and CH4 [41] on metal surfaces. We have used the band approach to determine the cohesive energy as well as the adsorbate binding energies on various transition metal systems (Mo, Cu, and Ni). The results have demonstrated good agreement with experimental results. Bulk cohesive energies are within 10 kcal/mol of the experimental results. Both the cluster, as well as periodic approaches, will likely play invaluable roles in the future toward the quantitative prediction of transition metal surface chemistry. Herein, we discuss some of the recent developments on the application of DFT-cluster calculations to chemisorption and reactivity of adsorbates on metal surfaces. We demonstrate how these results can subsequently be used to begin to model overall catalytic cycles and interpret different selective oxidation chemistries.

II. Chemisorption on Metals The quantitative analysis of adsorption on metal clusters, as discussed earlier, can be sensitive to the cluster chosen to model the chemisorption site. In a systematic series of studies, we examined the adsorption of various atomic and molecular adsorbates on different metal surfaces. We found that it was important to optimize: 1) size, 2) structural configuration, 3) spin state, and 4) the relaxation that occurs upon chemisorption of the cluster. For all transition metal elements beyond the 3d row, relativistic corrections become very important as well [35]. Many of the previous studies in the literature, however, have not carefully considered each of these features, and therefore, have lead to incomplete results. The palladium dimer, for example, was found to be quite sensitive to relativistic effects and non-local gradient corrections. The correct predictions of bond length, energetics and spin state required the specific accounting of relativistic corrections. Nakao et al. [51] have shown that relativistic corrections are responsible for lowering the energy of the 5s-5s bonding orbital. This orbital drops enough in energy whereby it can now accept an electron from the dz2-dz2 antibonding orbital, thus resulting in a change from a singlet to a triplet ground state. Relativistic corrections sTengthen .ed the degree of orbital overlap, thus leading to a decrease in the bond length from 2.77 A to 2.55 A. The binding energy, therefore, increased by 7 kcal/mol.

A) Cluster Size Effects The lowest energy spin states for a series of DFT-optirnized Pdx clusters were found to be triplet states. All calculations included nonlocal gradient and relativistic corrections internal to the SCF convergence scheme. The lowest energy configurations were that of the closed-packed arrangement, where metal-metal bonding was maximized. The lowest energy configurations and spin states for each of these clusters were then used to analyze the chemisorptive properties [35]. Figure 2 depicts a plot of the binding energy for atomic oxygen against Pd cluster size. The oscillatory behavior is consistent with experimental evidence for similar cluster systems [44, 45, 52, 53]. Other theoretical studies have also demonstrated oscillations with cluster size [42, 46, 54]. The interesting feature here is that beyond 5 to 6 metal atoms, the binding energy oscillations are close to the error associated with the DFT approximation itself. Oscillations are within 3 kcal/mol of the experimental value for atomic oxygen on palladium (-365 kJ/mol). This is in contrast with previous MO-ab initio results on fixed clusters which indicate that these oscillations are much more dramatic and require "cluster preparation" prior to calculation [42]. This cluster preparation technique is a scheme that was devised to minimize the energy required to excite an electron in the metal cluster from an occupied to unoccupied orbital. While this technique has been used to predict values in better agreement with the experiment, we hold reservations because the bond-prepared cluster is not a true ground state.

-250

~

-3OO

-3so

.4oo

-4so 0 1 2 3 4 5 6 7 8 9 1011121314151617181920 C l u s t e r Size

Fig. 2. The effect of cluster-size on the prediction of binding energies. A comparison of DFT-predicted (data points) binding energies for oxygen on Pd(111) cluster models with the experimental value of 365 kJ/mol (horizontal line). Returning to DFT methods, Rochefort et al. [55] also found that reliable energetic information could be established from the DFT-cluster approach by allowing for strict accounting for cluster relaxation and spin-state optimization rather than adopting the bond preparation scheme. One possible explanation for the differences displayed between the early HF cluster calculations and current DFT methods, is that HF methods provide a much poorer representation of the virtual orbitals than density functional theory. This would lead to substantial overestimates for the energy required to excite an electron. Indeed, more robust MCPF calculations which include core-valence correlation effects now indicate that the initial MO methods were substantially off in their predictions of binding/adsorption energies. The newer MO ab initio results are in good agreement with both DFT and experirnental binding energies [56]. In addition to the results presented in Figure 2 for the oscillations of O/Pd(111), we have also examined the binding of hydrogen (atomic), CO, and acetate on palladium. These systems are also well-behaved in terms of binding energy oscillations with cluster size. Realistic binding energies were found on all clusters greater than 6-7 palladium atoms. B) Adsorption Energy Predictions

A general feature that has emerged from the cluster calculations is that all atoms directly involved in the adsorption complex should maintain the coordination number that they would have in an actual surface. A minimum cluster size for an atop-bound adsorbate in an fcc surface, for example, would be 6 atoms in the top layer and 3 atoms in the second layer, resulting in a coordination number of 9. We found that a 19 atom metal cluster (M 19) with 12 atoms in the top layer and 7 just beneath, provides a reasonable compromise toward balancing CPU resources and computational accuracy [47] for a number of systems. The results from our analysis of adsorbates on different transition metal surfaces are summarized in Table 1. In general, the cluster calculations provide energies to within the 4-5 kcal/mol. This agrees quite well with the established organometallic literature reviewed by Ziegler [26]. Beyond this level of accuracy, we begin to bump into the accuracy of the DFT method itself rather than the cluster approach.

10 Table 1

Comparison of DFT-predicted and experimentally measured adsorption energies. All energies are reported in kJ/mol.

A T O M I C ADSORBATES Oxygen on Cu(111 ) Oxygen on Ni(111) Oxygen on Pd(111) Oxygen on Rh(100) Hydrogen on Pd(111)

,

ADSORBATES 02 on Pd(111) 02 on Pt(111) CO on Cu(111) CO on Ni(111) CO on Pd(111) NH3 on Cu(111) Ethylene on Pd(111) Water on Pd(111) Formaldehyde on Pd(111) Maleic Anhydride

DFT (kJ/mol) ' 449 44 1 349 456 251

Experimental (kJ/mol) 431 469 364 450 259

MOLECULAR

30 31 63 142 130 40 61 40 47 80

25-50 32 66 155 146 --50 59 42 50 93

Molecular adsorbates and surface intermediates present an interesting challenge in that they can typically bind at different sites. Carbon monoxide, for example, can sit at atop, bridge and hollow positions on a number of fcc metal surfaces such as Pt and Pd [16]. Each of these modes are actually local minima on the potential energy surface. Establishing the lowest energy mode, therefore, requires the structural optimization and energetic predictions for all three of these adsorption geometries. On Pd(111), we f'md that the three-fold hollow is the lowest energy site (145 kJ/mol), followed by the two-fold bridge site, and then the atop site. The results are consistent with UHV single crystal studies. At lower coverages, CO occupies the 3-fold sites in a -~/3x-v~ R 30 ~ structure. At coverages greater than 0.32, CO occupies bridges sites. At coverages greater than 0.5, CO is found at the atop adsorption sites. A second example involves the adsorption of ethylene on transition metal surfaces and offers an interesting challenge, in that it can bind via rc or di-t~ adsorption modes. Complete structural optimizations were performed for ethylene in both coordination geometries (Fig. 3). In the nmode, the ~ orbital on ethylene interacts with the dz 2 orbital on the metal center. There is a backdonation of electron density into the antibonding re* orbital of ethylene which leads to a small weakening of the C-C bond length. This is noted by the slight increase (0.05 A) in the C-C bond from the gas phase value 1.34A. There is considerably more backdonation of electron density from the surface into the n* orbital when ethylene is bound di-o. This is evident from the substantial increase of the C-C bond (1.45 A) when bound di-t~. This geometry is in excellent agreement with the experimental value 1.44 A [57, 58]. The predicted binding energy for the di-o mode is -61 kJ/mol which is slightly more stable than the r~-bound mode (-31 kJ/mol), and closer to the experimental measured value of-59 kJ/mol. Both ~ and di-a modes, however, are thought to be present in the experiment as indicated from HREELS data [59-61].

11

Di-c~ ~C ...~. ~:.:~,. .,. . . . . . . . . .

1.: ~ ::~

AEads= -30 kJ/mot

2.1

"i:

AEads=-61 kJ/mol Experimental -59 kJ/mol

Fig. 3. Adsorptionmodes for ethylene on Pd(111). DFT computed structures and binding energies. Both examples presented above highlight the multifunctional nature of the catalytic surface. Different sites can ultimately lead to different reaction paths. In addition to the complex functionality of the surface, many of the industrially-relevant reactants and intermediates contain more than one functional group that can interact with the surface. Multifunctional adsorbates are currently an active area of research [62], [63], Shekhar, 1995 #129; Davis, 1992 #127; Davis, 1991 #89; Houtman, 1994 #150; Delbecq, 1995 #117]. In Fig. 4, we present our DFT results for the adsorption of acrolein on Pd(111). Acrolein contains both a C=C and a C=O functional groups as well as two lone pairs of electrons on the oxygen. This gives rise to various different modes of adsorption. These include an 114 (both C=C and C=O surface interactions), 1]3 (di-~ (C=C) and lone pair surface interaction), rl 2 {di-c (C=C)}, 1"12{di-c (C=O)}, as well as all of the interactions through the rc (C=C) and ~ (C=O) interactions that can exist. The UHV HREELS results by Davis and Barteau [64] indicate that acrolein binds through the C-C bond and through some form of interaction with oxygen. The C-C stretch mode is substantially shifted from 1724 to 1670 cm -1 upon adsorption. This is indicative of di-a bonding of the ethylene moiety. There is also a shift in the C=O stretch. Based on the appearance of the 8(CCC) mode, Davis and Barteau suggested that the shift is due to a direct interaction of the C=O bond with the surface. Our computational results indicate that the rl4(C=C and C=O) interaction is hindered, and that the rl 3 (di-~ C=C and the lone pair oxygen) interaction shown in Fig. 4 is more energetically favored. Despite the differences in the suggested mode of adsorption, the predicted adsorption energy is in good agreement with the TPD data taken by Davis and Barteau [64]. The behavior of acrolein is similar to that found for acetone on palladium, where the CH3 substituents on acetone hinder the rl2(C=O) adsorption mode and make the tilted rl 1 mode more favorable.

12

~::i

....

1.490 e~

)C

.....

Acrolein on Pd(111)

: " 2.76

Mode 1]2 (C=C) a n d 111 ( C - O ) Adsorption Energy ..~

.f

AEADs = --47 k J / m o l AII~

Exp. "~.~,DS = --50 k J / m o l

Fig. 4. DFT-predictedadsorption of acrolein on Pd 19 model of Pd(111).

C) Adsorbate Induced Relaxation Effects In most of the results presented above, we have taken special care to optimize both the adsorbate, as well as the metal surface, in the chemisorbed state. The relaxation of the metal surface can, in some instances, lower the binding energy by up to 30 kJ/mol and distort the cluster geometry as is seen in Fig. 5 and 6. Fig. 5 depicts the structural changes in the small Pd6 metal cluster upon the addition of atomic oxygen, and atomic hydrogen. The metal-oxygen and metal-hydrogen bonds are substantially stronger than the metal-metal bonds. Atomic oxygen and atomic hydrogen therefore form strong bonds with the three metal atoms of the cluster that they coordinate with. To compensate, all other bonds to these atoms become weaker. This is apparent by the longer metal-metal bonds involving atoms Pdl, Pd3, and Pd6 in the adsorbate cluster complexes in Fig. 5. Next nearest-neighbor metal bonds become somewhat stronger as noted by their shorter bond lengths. This is consistent with Bond Order Conservation [65, 66] principles. Similar results are found for atomic adsorbates on larger metal clusters which offer a better description of the surface. This is shown in Fig. 6. The three metal atoms in the surface adsorbate complex are pulled up out of the surface by 0.1/~ and elongate by about 0.1 A. Surface relaxation increased the binding energy for oxygen on Pdl9 by 20 kJ./mol. While the relaxation on the larger Pdl9 atom system is more constrained than that on the smaller Pd6 cluster, it is still present. The classic example of ethylidene on Pt was examined and found to relax the metal surface atoms involved in the complex by 0.1 A,. This is consistent with LEED measurements by van Hove and Somorjal [67-69]. Atomic adsorbates are more strongly bound to metal surfaces than molecular intermediates and therefore, lead to much larger relaxation effects.

13 A)

B)

c)

~d3

1

1

~4

1-3 2.756I 1-6 2.631 3-6 2-4 2-6 4-5

2.756 2.756 2.632 2.757

Pd-X

1-3 1-6 3-6 2-4 2-6 4-5

2.918 2.898 2.918 2.667 2.756 2.661

1-7 1.768 3-7 1.768 6-7 1.767

1-3 1-6 3-6 2-4 2-6 4-5

3.131 3.074 2.918 2.651 2.757 2.628

1-7 2.068 3-7 2.068 6-7 2.068

["'] Primary effects (local)

Fig. 5.

Adsorbate induced cluster relaxation. The optimized structures for A) Pd6, b) I-l~d6 and c) O/P6.

Atomic Oxygen on Pd(111) A)

........

....

Pd 6 P d - P d R e l a x a t i o n - 0.3 A

Fig. 6

B)

2.

o

o,

Pdl 8 P d - P d R e l a x a t i o n -- 0.1 A

The effects of cluster size on adsorbate-induced surface (cluster) relaxation. A) Oxygen chemisorption on Pd6, and B) oxygen on Pdl9.

14

D) Higher Surface Coverages: The Extended Band Approach As we begin to explore higher surface coverages, the cluster approach becomes limited due to the size of the cluster required to accurately represent the system. The extended-band method provides an excellent alternative. Higher surface coverages are actually easier to treat because the size of the unit cell is substantial reduced, thus saving considerably in terms of computational efforts. To demonstrate, we present recent work on modelling high coverage saturation of carbon monoxide on transition metal surfaces. At high saturation coverages, CO was found to preferentially bind at the atop position on Cu(100) to help reduce Pauli repulsion. Results reported here are consistent with the LDA-DFT calculations by te Velde and Baerends [70]. The copper surface is represented as a slab which is infinite in both x and y directions and three atomic layers deep. The results for full coverage are presented in Fig. 7 A and B. The predicted adsorption energy, 17 kcal/mol, is in excellent agreement with the experimental value of 17.6 kcal/mol. Crystal orbital overlap population (COOP) analysis is used to distinguish the relative amount of orbital overlap between adsorbate-surface orbitals and specific location of the overlap bands with respect to the Fermi level. In addition, the sign on the overlap characterizes whether or not the given bands are bonding or antibonding. The results for CO on Cu(111) slab are presented in Fig. 7B. The antibonding features just below the Fermi level correspond to the 2n* antibonding orbital of oxygen. This clearly indicates the backdonation of electrons from the surface into antibonding 2~* orbital to help stabilize the adsorbate on the surface.

A)

B)

CO chemisorbed on Cu(lO0) 3 layer slab

CrystalOrbitalOverlapPopulation(COOP)diagram

2

Bonding States

ii

0

ji

s i'

co'2~* i il v iiiii~ i~' -6

!

Antibonding States ~t "

n

n

li

II

g

-

~

Binding Energy(DFT) = 17.6 kcal/mol Experiment = 17 kcal/moi

i:EFermi=-12.483 eV ,

-14 -20.41

!i

-13.60

-6.80

0

6.80

13.60

20.41

E - Efermi

Fig. 7. ExtendedDFT-band results for the adsorption of carbon monoxide on Cu(100). A) The mode and energetics of adsorption. B) Crystal orbital overlap population predictions [70b].

E) Structure-Chemisorption Properties The real value in computation lies in its ability to easily change substituents, surface structure, surface composition in an effort to explore structure-property relationships. While we have spent a considerable effort in modelling the adsorption of ethylene on palladium, the extension to substituted-ethylene structures is straightforward and used to provide simple structure-property relationships. The results are depicted in Fig. 8, where we replace one of the terminal hydrogens of ethylene with different electron-withdrawing groups. The more electron-withdrawing

15 substituents increase the charge on the carbon, thus making it less negative. This slightly decreases electron donation to the surface and ultimately weakens the binding energy. The computed values for ethylene, vinyl alcohol, 1-fluoro-ethylene are -61, -50, and -45 kJ/mol respectively [71]. While a rigorous experimental study has not been performed, we can compare adsorption energy of other substituted ethylene species. Acrolein, for example, has an adsorption energy which is nearly 20 kJ/mol less favorable than ethylene, despite the fact that both the ethylene and the oxygen functional groups interact with the surface. The ethylene moiety in acrolein, clearly has a lower overall binding energy than ethylene itself. We return to structureactivity issues in our discussions on structure reactivity

A)

CH2=CH2 ,~...1.452/:fli~ ..................... ................... .:..:. ~~

A

,~

i

B)

CH2=CHOH

-0.41 .

.

~

AEADS = -61 kJ/mol

~ .

.

.

1.45 .

,!:;;;~" .

~:............. ,i,i:..........................~

.......: ....

AEADs"- --54 kJ/mol

C)

-0.068

CH2=CHF

~

+0.02

..:+::..:.!.i.~:ii:~:~.~.:~............. ::.. .:-..:..... . 2.12 i

~.2.13

AEaos = --44 kJ/mol

Experimental =-59 kJ/mol Fig.8. The effects of substituents on the adsorption of substituted ethylenespecies (CH2--CHX) on Pd(111). X refers to: A) -H, B) -F, and C) -OH. F) Spectral Predictions for Adsorbates on Surfaces

As was discussed, infrared and raman spectra for organometallic systems can typically be computed to within 5% of the experiment. Unlike adsorption energy predictions, structure and vibrational frequencies are fairly insensitive to differences in the DFT methods (local vs. nonlocal spin density). Even some of the earliest reported local-spin-density approximation (LDA) DFT calculations which ignored adsorbate and surface relaxation predicted frequencies to within 10 percent of the measured values. For example, Ushio et al. have shown that LDA calculations for formate on small Ni4 clusters (frozen at its bulk atomic positions) provide very good agreement with experimental HREELS studies on Ni(111) [72]. Unlike adsorption energy predictions, structure and vibrational frequencies are fairly insensitive to gradient-corrections. We have used frequency predictions to corroborate experimental HREELS results and to help resolve active surface intermediates. We analyze the results for two specific systems: the binding of acetate and the adsorption of maleic anhydride. In the first system, Barteau and Davis speculated that the surface acetate intermediate is bound di-~ [73]. By examining the energy of adsorption for the monodentate, bidentate, and di-c acetate intermediates, we were able to help confirm that the di-~ was the lowest energy state. To further support this, frequency calculations

16 were performed on the optimized di-o bound species and compared with the HREELS results from Davis and Barteau. The results for the predominant modes are given in Table 2. DFT predictions are within 5% of the HREELS data. Table 2

Comparison between DFT-predicted and experimental frequencies determined by HREELS for acetate on Pd(111) and maleic anhydride on Pd(111). DFT predictions were computed by in our laboratories [36]. Experimental HREELS data reported in [73].

Vibrational Mode

Frequency HREELS

DFT

Vs (Pd-O)

320

295

8 (O-C-O)

685

665

Vs (C-C) d (CH3) Vs (OCO)

925 1215 1415

950 1380 1469

Table 3.

Comparison of DFT-predicted and HREELS measured frequencies for maleic anhydride gas phase, n-bound on Pd(111) monolayer, and di-(~ bound on Pd(111). DFI" predictions were computed by in our laboratories [71]. Experimental HREELS data reported in [74].

Multilayer

rc-Pd(lll)

Mode

Exp.

DFT

Exp.

C=O in plane bend In plane ring bend C-H Bend (wag) C-C stretch C-H Scissor C-O-C (symm. stretch) C=C stretch C=O (assym. stretch) C=O (s~nn. stretch)

408 695 863 863 1060 1250 1595 1785 1880

379 680 860 860 1048 1235 1617 1787 1841

400 397 680 706 835 828 835 841 1175 1247 (1300) 1466 1610 (1580) .........

DFT ,

di-o Pd(lll) Exp.

DFT

400 690 800 800 1171 NS 1827

379 691 788 879 1078 1200 1434/1460 1780

In the second system, we examined the lowest energy modes for the adsorption of maleic anhydride on Pd(111) [71]. The optimized geometries for 1"11, re, and di-o bound maleic anhydride are shown in Fig. 9. Second derivative calculations for MA on smaller Pdl0 surface clusters were performed to determine frequencies. The di-c~ mode, where maleic anhydride lays flat to the surface was found to be the lowest energy state. This is in good agreement with the experimental HREELS data of Xu and Goodman [74]. The computed vibrational frequencies are compared with the results from HREELS data for the gas phase species, MA r~-bound, and MA di-o bound on Pd in Table 3.

III. Surface Reactivity The reliability of DFT calculations to the prediction of adsorption energies on metal, metal oxide, and metal sulfide surfaces is now fairly well-established. The application of DFT methods to predict activation barriers, however, is still an emerging area. The available literature is,

17 therefore, limited. For the most part, previous efforts have focused on the activation of small molecules such as H2, 02, CO, NO, N2, and CI-I4 on metals and in zeolites.

A)

1"11 ~'~...........Y .. 1.44 , "i;.

B)

rt

9

C)

di-G ..

~ ~

1.459 A:~'1.476 .

2.18

z ~ s = -28 kJ/mol

A E ~ s = -34 kJ/mol

AE~)s = -83 kJ/mol AEe~p = -90 kJ/mot

Fig. 9. Adsorptionmodes and corresponding DFT-predicted adsorption energies for maleic Pdl9 model of Pd(lll).

anhydride on the

While DFT has been used to successfully model activation barriers for some gas phase systems and well-defined organometaUic complexes, there is a concern that DFT methods under-predict the barriers for some free radical abstraction systems. The under-estimation is primarily due to the over-accounting of the self-interaction of electrons in the SCF procedure [75]. Surface-bound free radical intermediates demonstrate less localized unpaired spin due to their strong interaction with the surface. This is likely to reduce some of the problems related to self interaction effects. Below we summarize results for DFT computed barriers for the activation of simple adsorbates over different transition metals. A more thorough and systematic investigation is required to better understand what controls the accuracy in activation barrier predictions on metal surfaces.

A) Small Molecule Activation: Dimers Through a rigorous set of single-point calculation we mapped out the potential energy surface for the dissociation of NO over a model Cu(111) cluster [76]. Fig. 10 summarizes the overall reaction coordinate for NO dissociation. NO adsorbs from the gas-phase at a three-fold site and subsequently dissociates to form N* and O* which sit in three-fold hollow sites. The reaction path involves an N-O stretch along with a metal atom insertion [76]. The transition state is late with respect to the N-O stretch. The strong metal-oxygen and metal-nitrogen interactions help stabilize this state on the surface. The transition complex looks much more like the products than the reactants. The Cu(8,3) model cluster used in this study was somewhat small for modelling the dissociated state. Edge effects for N and O on this cluster can lead to an overestimation of their binding energies, van Daelen et. al [77, 78] found similar results for the activation of a series of small molecules, CO, O2, N2, and NO over Cu. More recently, fullyperiodic DFT calculations have been performed on the chemisorption of H2 [50, 79-81], 02 [48]

18 and CI-I4 [41]. The results from these band analyses are in excellent agreement with known experimental values.

Surface Reactants 1.15 AEAD S = -83 kJ/mol

~I

2"1~r~2.177

Transition State

AE~'= +88 kJ/mol

2 175 2.710 ~1~:" 1 " 9 ~

.... ~ . _ . . . ~ .

1-967

L,2,,7

Surface Products

1.975~ .

0

...........~#~1.967AErxn

= - 4 6 kJ/mol

Fig. 10 Reactionpath analysisfor NO dissociationover a model Cu(111)cluster.

B) C.H Bond Activation C-H bond activation is essential in a number of commercially relevant processes including methane activation and selective oxidation of alkanes and alkenes. It is also important in that it is the microscopic reverse step for hydrogenation. Our interest in C-H activation is driven by attempts to model two specific industrial processes, the acetoxylation of ethylene for synthesis of vinyl acetate, and the selective hydrogenation of maleic anhydride (MA) to tetrahydrofuran (THF). In VAM synthesis, one of the speculated rate-determining steps is the activation of the C-H bond of ethylene to form a surface vinyl species [82]. The first step in the hydrogenation of MA to THF involves the addition of hydrogen to the ethylene moiety. This step is the microscopic reverse for the C-H bond activation of the surface-bound maleic anhydryl. The transition states for both ethyl and maleic anhydryl C-H bond activation are very similar. This transition state structure is likely common for other C-H bond activation paths.

19 We have analyzed the 13-hydride elimination mechanism and predicted the activation barriers for a homologous series of substituted adsorbed ethyl species, CH2-CH2X. o~ and [3 refer to the positions of the hydrogen along the carbon backbone in the adsorption complex, a corresponds to the carbon atom which is attached to the surface, whereas [3 is its nearest heavy-atom neighbor. Various electron-withdrawing substituents were substituted for X to probe their effect on the transition state and the computed activation barriers. More specifically, we analyzed ethyl (CH3CH2o), 1-fluoro-ethyl (CFH2-CH2~ and 1-propyl aldehyde (CH2(CHO)-CH2~ surface intermediates. The activation of the C-H bond of the adsorbed ethyl intermediate involves a much more complex potential energy surface than those for the simple dimers discussed above which only had a few degrees of freedom. The predominant mode for the ~-hydride elimination involves a C-H bond stretch. To isolate the region of the transition state on the potential energy surface, a series of optimizations were performed along the chosen reaction coordinate. The reaction path was frozen at various positions along the coordinate while all other modes were optimized. A plot of this 2D coordinate provides a rough location of the activation barrier. Frequencies were then computed for this approximate transition state to determine the mode of the true reaction coordinate. A more rigorous mode-following technique was then used to isolate the true transition state in this system. The corresponding structure for the transition state for ethyl activation on Pd7 is shown in Fig. 11A along with the reaction coordinate vector and its frequency, -228. The activation barrier on the small 7 atom cluster is +59 kJ/mol. This is extended to the larger Pd 19 cluster by performing sequential reaction coordinate optimizations on the larger system. The results for the adsorbed-, transition-, and product-states on the Pd19 cluster are shown in Fig. 12. The reactant and product states are fully optimized. Only the handoptimized transition state can be found for the large 19 atom systems. Frequency calculations for any transition metal lower than the 3d series in the periodic table require numerical second derivative calculations. These become computationally very intensive as you move beyond more than 10 metal atoms. Frequency calculations for ethyl on Pd7, for example, took nearly 50 hours on a 4-processor Cray C90. The Pd 19 cluster was, therefore, much too large to perform the same true transition-state isolation search.

A)

"".

1.478 ~.~.

k'

~

Ethyl

~ ~ 1.582

v = 228.2 cm-1

Acetate

v = 264.7 cm-1

Fig. 11. The isolatedtransition state for the 13-Heliminationof A) surface-bound ethyl species and B) acetate on Pd7 model of Pd(111). The negative eigen modesare shown beneath, 264.7and 228.2 cm-1, respectively.

20 The reaction coordinate identified in Fig. 12 involves a C-H stretch along with a C-C-surface bending motion toward the surface. The transition state is late in the C-H stretch with a C-H bond length of 1.68/~. M-C and M-H interactions form and help to stabilize the transition state. The M-C and M-H bond lengths here are 2.26 A and 1.58/~ respectively. The transition-state is quite similar to that found for methane activation on organometallic clusters and metal surfaces [3, 30, 41, 83-87]. On Ni, for example; the C-H bond stretch for methane is 1.80 A, while the M-C and M-H bonds are 2.21 A and 1.5 A respectively [30, 84]. The main difference between methane and the ethyl group described here is that the initial CH2 moiety of the ethyl group anchors the ethylene backbone to the surface and provides for enhanced stability. This helps to lower the activation barrier from that of methane or ethane. The predicted barrier, here is +69 kJ/mol which is only slightly higher than the experimentally measured barrier of 47-57 kJ/mol for ethyl on Pd(100) that was determined by Kovacs and Solymosi [88]. CH2-CH 3. -> CH2=CH2. + H*

A)

,~ ~:

B)

C) ~

1.479 ~ '

:............!:!i.......................

. . . . . . 2.16 ,*l ............................... ~ 2.25 1.71

AE = 0 kJ/mol

AE* = + 69 kJ/mol

z~Erx n =

- 3 kJ/mol

Experimental 4 0 - 5 7 kJ/mol

Fig. 12. The I3-Helimination reaction path analysis for ethyl on Pd(111). DFT predictedA) adsorptaon-, t~) transition-, and C) product-states. The overall reaction products are adsorbed ethylene and atomic hydrogen. Hydrogen can bind to either the two-fold bridge position or the three-fold hollow site. The three-fold fcc site is energetically favored over the hcp three-fold and the two-fold bridge sites for H/Pd(111). The overall reaction is exothermic at -6 kJ/mol. These results are in good agreement with the overall energy for [3-hydride elimination from ethyl on Pt(111) which predict a value -28 kJ/mol [89]

C) Substituent Effects By understanding the nature of the transition state, we can begin to explore the effects of substituents, promoters, and inhibitors on the adsorbate-surface bond. The effect of adding electron-withdrawing groups at the [3-carbon on the proposed transition state and activation barrier were examined here by substituting one of the [3-hydrogen atoms with a fluorine atom and then with a -CHO group. The transition state isolation process described above was repeated to locate the structure and to determine the barrier for the [3-hydride elimination of the 1-fluoroethyl species. While the structure was nearly identical to that for ethyl, the barrier for the fluoroethyl group is over 20 kJ/mol higher (E* --- +92 kJ/mol). The results are consistent with those of Forbes and Gellman [90-92] who found that fluorine substitution at the 7-carbon of the surface propyl group substantially increased the barrier for [3 hydride elimination on Cu. They proposed that the CF3- group is electron-withdrawing, leading to greater charge transfer from the carbon

21 and therefore destabilizing H & transfer. Our results closely follow their speculated analysis with the minor exception that they propose the activation over a single metal atom center (/) rather than the two metal atom complex (I/) proposed above. ~5"

~5+

C

~5-

C

I.

I,

I

C-.

I,

I

M'95 +. . . . .

~5+

C

I,

I

I

M ' .....

H'8

- 7_i 9

M' "2 ". . . .

-

9

M

I II Our predicted transition state (Fig. 12) more closely follows that reported by Bent [89] who used

stereoselective groups to confirm that the mechanism involves a syn [3-elimination step [931. The CH2(CHO)-CH2~ group contains a -CHO substituent which is slightly less electronwithdrawing than the F- substituent on the CH2F-CH2~ group. The DFT computed barrier for the CH2(CHO)-CH2o species (E* = +75 kJ/mol)was found to lie in between that for ethyl and the fluoro-ethyl systems as might be expected from a simple correlation between the electronwithdrawing ability and the barriers for [3-hydride elimination. In Fig. 13, we depict the [3-hydride elimination surface reaction for the CH2(CHO)-CH2-o surface species to give acrolein. This microscopic reverse of this step involves the selective hydrogenation of acrolein, a valuable selective oxidation intermediate. Acrolein is a structural moiety for maleic anhydride, and therefore an ideal model for the hydrogenation of maleic anhydride to succinic anhydride. The predicted transition state is shown in the center of Fig. 13. The corresponding barrier for addition of hydrogen to adsorbed acrolein ( the reverse reaction) is +82 kJ/mol.

(OHC)CH2-CH2 * -> (OHC)CH=CH2 * + H*

A)

~

B)

if: 1 522 1.525 . ~ .....

~ ~ 2.101 ~ii! ,~ ~ . .~. . .,. . , . . ~ . ~. , ...... . L ~

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

C) ~ ....................... i ~"~ 1.495

1.. ~2.015~ ,............................ ,~"~

:

17o

. ~ ~ .:. . . . .1. ...~. . %

2.12 t ~ " ~ 2 " 3 ~}: :1.6r:' ~' 2.06 ~

~

N,~:.. 1.464 ~r 14~ 2 . 1... ~ ............................ 1 7 " :":il, , 9 ".......................

""

:i~..........................~ " .. 1 . 7

~

AErx n = + 75 kJ/mol

AErx n = - 7 kJ/mol

Fig. 13. The 13-Heliminationreaction path analysis for CH2-CH2(CHO) on Pd(111). DFT predicted A) adsorption, B) transition-, and C) product-states.

22 D) Acetate Activation The chemistry of acetate on transition metal surfaces is important for a variety of selective oxidation processes. Methanol and vinyl acetate syntheses are two such important oxidation chemistries where acetate intermediates have been postulated. In VAM synthesis, acetate is a critical intermediate in both VAM formation, as well as in its decomposition to CO2. The latter unselective decarboxylation path becomes important at higher operating temperatures. Understanding the mechanism for decarboxylation and VAM synthesis may ultimately aid in the design of new catalyst formulations on new operating conditions. Acetate decomposition has been examined experimentally on Pd(111) [73, 94], Rh(111) [95], Rh(100) [96], and Rh(ll0) [96] at ultrahigh vacuum conditions. In a combined TPD and HREELS investigation on Pd(111), Barteau and Davis [73] found that acetic acid weakly adsorbs and forms a series of catemeric structures at low temperatures. As the temperature is increased to 200 K, acetic acid decomposes to form acetate surface intermediates which are stable up to 300 K. HREELS data indicate the acetate intermediate is bound di-~ to the surface. Our DFT cluster results, described earlier, help to confirm the di-a adsorption mode (Fig. 14A). At 300 K Davis and Barteau found that CO2 desorbs from the surface. They speculated that the activation of the C-H bond on the terminal methyl group is the rate-controlling step for decarboxylation. The CH2=CO2" surface intermediate (Fig. 14C) that forms is thought to undergo rapid C-C bond scission to release CO2.

A)

..............~

AEad s

-212 kJ/mol

~1.517

2.14

....

[

...

::,. :.. 2.0

~:

~

o

AErx n = +69kJ/mol

C) ~~i~.:~.::.

[2.19

1.70 ..~~

Fig. 14. The C-H bond activation of surface-boundacetate on Pd19 cluster model of Pd(111). The A) adsorbed- B) transition-, and C) the product-states.

23 The resolution of the reaction coordinate for the C-H bond activation of adsorbed acetate is complex in that there are a number of degrees of freedom which need to be considered. The acetate surface species can approach the surface through various different ways from its initial perpendicularly adsorbed state. In a series of detailed first-principle calculations, we found that it was energetically most favorable to bend at the angle created by the C-C bond and the surface normal [36, 71]. The activated complex is depicted in Fig. 13B is very similar to that shown for the activation of theoethyl group (See Fig. 11A and B). There is a considerable stretch in the CH bond (from 1.1 A to 1.78 A ). Metal-hydrogen and metal-carbon bonding are evident, as indicated by the M-H (1.70 A) and M-C (2.45 A) bond lengths, and help to stabilize the transition state. The predicted barrier for C-H activation of the surface acetate (+115 kJ/mol) is somewhat higher than the experimental value reported by Davis and Barteau of 85 kJ/mol [73]. The value reported by Davis and Barteau, however, is an overall barrier for acetate decarboxylation and may be lower than the intrinsic C-H activation step which is computed here. The overall energy for acetate activation was found to endothermic by +69 kJ/mol. The CH2=CO2" surface intermediate that forms lies parallel to the surface. There is a considerable backdonation of electrons into the antibonding x*C=C orbital, thus weakening the C-C bond. Carbon-carbon bond scission releases CO2 into the gas phase and produces a CH2* surface species.

E. The Effect of Alloying: Pd/Au Bimetallics It is well established that alloying an active metal with a second can lead to marked increases in the activity and selectivity of a heterogeneous catalyst [97]. While considerable work has been done to understand the nature of charge transfer, there is still a ongoing debate as to whether the measured increases in activity and selectivity are due to geometric or electronic effects [97]. Recent surface science experiments indicate that shifts in the core level XPS binding energies can be used to deduce the electronic interactions of a monolayer deposited over a second welldefined metal surface [98-100]. Rodriguez [101,102] and Hammer and N0rskov [17, 103, 104] have performed elegant analyses of CO bonding on bimetallic pseudomorphic overlayers. Very little, however, has been done to distinguish how the composition and spatial properties affect the fundamentals of adsorption and surface reactivity on the primary metal. First principle computations provide a nice framework by which we can readily substitute atoms of a second metal in for the primary metal lattice to examine their effects. We demonstrate this here by replacing Pd atoms in the surface with gold to establish the effects of gold on the barrier for acetate decomposition. The results on the bare palladium surface indicate that the barrier on Pd7 cluster is +150 kJ/mo!. This is depicted in the bottom path shown in Fig. 15. By exchanging a palladium atom that is not directly involved in the transition state complex, but is a first nearest-neighbor with gold, the barrier is increased by + 17 kJ/mol. This is shown in middle path of Fig. 15 (The atom which is substituted is colored black). This suggests that the electronic effect, while relevant, is weak. If we substitute a palladium atom that is directly involved in the initial adsorption complex, the barrier is increased by +100 kJ/mol to a value of +250 kJ/mol. This suggests that this route is unfavorable. This is shown in the upper path in Fig. 15. Finally, if we substitute the central palladium atom in the complex with gold, the barrier for C-H activation goes up to +550 kJ/mol. This is clearly an unlikely step. The effect of gold in this system, therefore, is primarily a geometric one. Gold atoms in the surface are inactive and act to break up the palladium ensembles which are required to carry out C-H bond activation.

24

300

I1 '

'

'

I

.'

'

'

I

'

'

'

I

'

'

'

I

'

'

'

I

f

'

'

'

-i

200

150 =

100

0

1

1.2

1.4

1.6

1.8

2

2.2

Reaction Coordinate C-H Bond Length (Ang~.rom) Fig. 15. The spatial effects of replacing Pd atoms with Au in the Pd(111) surface on the activation barrier for C-H bond activation of surface-bound acetate. The lowest energy curve corresponds to pure Pd(111) surface. The middle curve is that/or a gold atom substituted at a neighboring palladium site, but not directly involved with the adsorbate-surface complex. The top curve corresponds to substituting Au at one of the sites where acetate adsorbs. The final plot, which is not shown here, has a barrier of 550 kJ/mol and corresponds to the system where Au is substituted at the central Pd cite which is responsible for activating the C-H bond. These results are consistent with a recent study that was published by NOrskov et al. [41] on the activation of methane over Ni and Ni/Au. Through rigorous first-principle periodic D F T - b a n d calculations, they demonstrated that gold raises the barrier for methane activation by + 16 kJ/mol for a single atom substituted into the adsorption complex. The barrier was found to depend upon the location, as well as the amount of gold added. In addition, they h a v e s h o w n that H2 activation over a gold surface is inactive [79] Preliminary results indicate that the binding energy of acetate and CO increases if gold is substituted for palladium at sites which are located one layer beneath the surface. The binding energy, however, decreases if the gold is actually substituted into the surface. M o r e work on the spatial and compositional effects, however, is required to better understand Pd/Au and other bimetallic systems. F) The effect of Surface Oxygen: From Transient Surface Precursors to Site Poisons The presence of oxygen can open up a number of additional reaction pat hways that can control the actual surface chemistry. Madix has demonstrated that adsorbed atomic o x y g e n can behave as a nucleophillic center and attack surface bound hydrocarbon intermediates or as a Br0nsted base for hydrogen transfer reactions [63]. C h e m i s o r b ed atomic oxygen can also act as a poison on different transition metal surfaces.

25

Both atomic and molecular oxygen have also been identified on various metal surfaces as transient precursors [105-116]. These are energetically mobile intermediates which can readily move across the surface and react with other surface species. "Hot" atomic oxygen species have been identified by Ertl [105], and Harrision [106]. These species move abnormally long distances across the surface from their site of dissociation, indicating that their kinetic energy is slow to dissipate to the surface in the formation of chemisorbed 0 2- *. Molecular oxygen, which weakly binds to transition metal surfaces, has also been identified as a transient oxygen precursor in hydrocarbon oxidation chemistry. Roberts et al. [108-116] have used adsorbed ammonia on metal surfaces at UHV conditions to isolate and identify transient oxygen intermediates. Their results indicate that both atomic and molecular oxygen intermediates can be present and that the predominant species for carrying out surface chemistry is a function of the metal as well as the operating conditions. To illustrate the multifunctional role of oxygen, we present two specific examples. The first system examines ammonia oxidation. In the second, we re-visit the acetate decomposition system on Pd(111) that was discussed earlier.

1. Ammonia Oxidation Roberts et. al. [115, 116] have demonstrated that a transient oxygen intermediate is responsible for opening up a non-activated route for the oxidation of ammonia over copper. The activation of ammonia in the gas phase is a highly unlikely, requiring over +400 kJ/mol. On Cu(111) the barrier is lowered to +344 kJ/mol but is still insurmountable [76]. Copper has a nearly filled dband which makes it difficult for it to accept electrons or carry out N-H activation. In the presence of atomic oxygen, the barrier is lowered substantially to + 132 kJ/mol [76]. The overall energy is endothermic by +92 kJ/mol. While the barrier is still somewhat high, it can now be overcome at higher temperatures. Analogous to the C-H activation mechanism discussed above, there is an agostic N-H stretch which is required to surmount the barrier. The N-H bond is stretched a considerable distance before the N-H antibonding orbital is low enough in energy to be populated by electrons from the surface. An intermediate which could effectively stabilize the transferring hydrogen would likely help to lower the barrier. The activation of ammonia in the presence of molecular oxygen was found to require a barrier which was only half that for atomic oxygen. The weakly bound oxygen intermediate flits and stretches toward the approaching hydrogen to form a hydroperoxo surface intermediate that can stabilize N-H bond activation. The established reaction paths for both the atomic and the molecular paths are illustrated in Fig. 16 A and B, respectively. The intrinsic barrier for N-H activation assisted by the molecular oxygen precursor, however, is only +60 kJ/mol. If we factor in the heat of adsorption of ammonia and oxygen, however, there is no apparent activation barrier [76]. These results are consistent with the experimental evidence from Roberts et al [115, 116]. Our results for oxygen-assisted acetate decomposition [36, 47] are consistent with the experimental work of Davis and Barteau [94].

2) Oxygen-Assisted Acetate Decomposition The presence of preadsorbed atomic oxygen stabilizes the di-tJ acetate surface intermediate by 15 kJ/mol (see Fig. 17) [36, 47]. Since acetate is already strongly bound to palladium, little is gained via attractive lateral surface interactions. These results are consistent with the experimental UHV results from Davis and Barteau [94] who demonstrate a weak increase in binding of acetate in the presence of atomic oxygen. Surface-bound atomic oxygen can, however, act as a Br~nsted base and change the kinetics and thermodynamics for C-H bond activation. The overall reaction energy (+25 kJ/mol) is less endothermic than that for acetate decomposition on the clean Pd(111) surface (+69 kJ/mol). A careful examination of the reaction coordinate and the transition state indicates that there is a repulsive interaction between the approaching -CH3 group on the acetate

26 and the surface oxygen. This causes the surface oxygen to move from a next-nearest neighbor 3fold hollow site to a bridge site and then on to a next-next nearest neighbor 3-fold hollow site as you traverse the reaction coordinate. The reactive surface ensemble, therefore, requires up to eight palladium atoms. At very low surface coverages oxygen raises the barrier slightly from 115 to 130 kJ/mol. At moderate to higher surface coverages, however, the decomposition route is essentially shut down due to a dramatic decrease in the number of active sites. There will be few ensemble-sizes which are greater than 8 atoms and, therefore, acetate remains stable on the surface. As the temperature increases, the coverage decreases. Decomposition occurs rapidly over a small range in temperature, exhibiting "domino" behavior and autocatalytic kinetics. Autocatalytic kinetics has been observed on Rh [96] as well as Ni [117].

A)

Surface Reactants -*~i' .~"008" x2.45

Transition Complex N 1~ "~i::~,1.86

Surface Products V

~ ~ .

AE = 0 kJ/mol

B)

, 1.66.. 4[jO (~' .::..z::~*""" ;~i.07 11.36 2.17[ ~(,)

AE = +132 kJ/mol ;.,1.00 ~##~[}O "~;~ ''~8:~5 ~; ~ TM i~

AE = 0 kJ/mol

2.0!a ~:11

AE = +67 kJ/mol

..

AE = +48 kJ/mol

~, 1.9

d~

~..~ o 0.98 I 97 ~.5N82 1.58 1.

.00

AE = -25 kJ/mol

Fig. 16. The activation of ammonia by a) atomic oxygen, and B) molecularoxygen transient precursors on a Cu(111) model cluster.

IV. Catalytic Cycles By collecting the energetics of adsorption for a large number of surface intermediates, their energies of formation, and coadsorption interaction energies, we can begin to fashion together overall catalytic cycles. The energies for each of the proposed steps for an elementary overall cycle can be examined and used to elucidate which steps may likely control. By comparing a series of proposed cycles, the most probable systems can then be distinguished. To demonstrate we present the results for the oxidation of ammonia over copper [76] and synthesis of vinyl acetate [36, 47]. In the ammonia oxidation system shown in Fig. 18, ammonia is adsorbed in the presence of oxygen on Cu(111). At lower surface coverages, atomic oxygen increases the adsorption energy of ammonia by 30 kJ/mol. The adsorption energy for ammonia in the presence of oxygen is -48

27 kJ/mol. In addition to stabilizing ammonia, oxygen also lowers the barrier for N-H bond activation. The overall energy of reaction for NH3* + O* --> NH2* + OH* is +48 kJ/mol. While the products are NH2* and OH*, we show only the species involved in the next reaction step, NH2* and O*. The OH* intermediates are not shown until a later step where they recombine and desorb as water. While ammonia favors the atop adsorption mode, the more electron deficient NH2* fragment prefers a bridging Cu site. Subsequent N-H activation then leads to NH* and N* sm'face intermediates. The reaction of NH2* to NH* is thermoneutral while the final reactions of NH* +O* --> N* + OH* is exothermic. The overall reaction energies are directly related to the stability of the NHx* product that is formed. The relative ordering of binding energies for NHx species are: NH3* < NH2* < NH* < N*. Therefore the overall thermodynamics become more favorable for each subsequent step. The product N* species recombine to form N2 (g). The presence of surface oxygen provides an additional pathway for the removal of N* as NO (g). While the reaction of N* + N* --> N2(g) is energetically favored over that of N* + O* --> NO(g), at higher oxygen surface coverages NO formation becomes the more predominant route. The OH* species can subsequently recombine to form water which desorbs, and surface oxygen. By examining the energetics for each of these steps we can see that the surface reaction step and the hydroxyl recombination steps are likely rate limiting candidates. The actual barrier for the NH3* + O* reaction was already reported above. Reaction coordinate calculations for the hydroxyl elimination reaction should also be performed to distinguish which is more likely to be rate controlling. This was just one speculated scheme. A series of others were also examined and tested for their potential likelihood [76].. A)

•

1.529~

~

•AEads =-227 kJ/mol

:,~:1.28

2.18 ~

I

....~ii"~E!~

2 01_,~ 2.05

~~AE*=+13OkJ/mol B)

1.49

~

2.141 1.31 , ~ ~ ~

'!:~ .32

C)

..... ............I!T M

1.30

9

/ AErxn= 25 kJ/mol

Fig. 17. Oxygen-assisted C-H bond activation of acetate on Pd. The A) adsorbed-, B) transition-, and C) product states for acetate activation on O/Pd(111).

28

~.,,...,~,-:.#

-107 Uh.,,I ~ ~ . . \2~fl\ ,,

~

~

~

~

"1' "1-

+48 kJ/tool

~

\

a

"~............, ~

.........

~

]A

.,~

+o.

~ ,

~

-,~!!~:~C,~r......

~.,.,.,~,~,~.~ .......... 9

~~-

Fig. 18. The overall catalytic route for the oxygen assisted decompositionof ammoniato N2, 02, and NO.

A) Vinyl Acetate Synthesis The foregoing analysis can be extended from the chemistry of ammonia to a more complex catalytic system such as vinyl acetate synthesis. Vinyl acetate is produced by the acetoxylation of ethylene in the presence of oxygen over supported Pd/Au particles. While this is a wellestablished commercial route, the mechanism is still poorly understood. It was postulated that the chemistry could occur in the liquid layer via homogeneous solublized Pd-acetate complexes. Recent evidence, however, indicate that the chemistry occurs on the Pd metal surface rather than on Pd(2+) particles. While we have explored both homogeneous as well as heterogeneous [36, 47, 118] mechanisms, we discuss only the heterogeneous results here. The overall cycle depicted in Fig. 19 involves the formation of VAM on a model Pd(111) surface carried out in the absence of oxygen. The fist step involves the adsorption of acetic acid. Acetic acid binds weakly to Pd via a van der Waais interaction with the lone pair of electrons on hydroxyl group along with a very weak interaction with the carbonyl oxygen. The computed adsorption energy, -5 kJ/mol, is slightly weaker than that determined by Davis and Barteau at UHV conditions [73]. The experiment results contain extra stabilization due to the formation of hydrogen bonded catemers. Hydrogen bonding can easily be responsible for the 30 kJ/mol difference between the experimental and DFT (monomer) predictions. Acetic acid dissociates at low temperature to form surface acetate species. In the absence of oxygen, ethylene adsorbs and reacts to form surface vinyl and hydride groups. The overall energy for this step is +67 kJ/mol (endothermic). The surface vinyl and acetate groups can subsequently couple to form vinyl acetate. This step is also endothermic (+67 kJ/mol). The VAM product desorbs at a cost of +20 kJ/mol. Hydrogen, which is produced in the acetic acid and ethylene activation steps, associatively recombines and desorbs as H2(g). The two surface reaction steps, ethylene activation and vinyl and acetate coupling are the two most endotherrnic steps, and are therefore ideal rate-limiting candidates [36, 47]. The barriers are still being worked out.

29 Heterogeneous Mechanism (H2 Formation) .

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

.

Acetic Acid Adsorption

,.

.

.

.

.

.

.

LL

~:

Dissociation of Acetic Acid

~~

IIIII

II II

I

Ii~..

\..7 Ae:-so

-~"2 +

Ethylene A

2CE----6I

~

....~ ....... ~: ..............................,.

~ H2 Associatil

<

.~,,.............?.!~,<.,~:~:s

:

Ethyl .... D

i

s

~

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

~

esolption

AE = + 2 9 . t

~. AE = +230 ] Vinyl and Acetate [! ~sutfa . . . . . . tion to VAM VAM desorption AE = + 20

~ ~ . . .:.......~............. :

:..,~:.:.:~,.

,:~e:~.~.....................................[."

100 Therm ochemical Prediction s Associati ve H2 D esorp lqon

_~ 50 E 0 Dissociation of AcOH

~4D t_

VAM formation

-50

-100

Ethylene Adsorption

Ethylene to Vinyl Surface Reaction

-150 Reaction Path Fig. 19. The overall catalytic route for the synthesis for vinyl acetate formation from ethylene, acetic acid in the absence of oxygen on a model Pd(111) surface.

While the reaction path presented in Fig. 19 provides some interesting ideas, it is only likely at very low surface coverages and in the absence of oxygen. The presence of oxygen will likely influence both the kinetics as well as the thermodynamics. The effect of oxygen on the overall energetics of the elementary steps of the VAM process is depicted in Fig. 20. In addition, oxygen decreases the endothermicity of ethylene to vinyl surface reaction by 20 kJ/mol. The

30 vinyl/acetate coupling and VAM desorption steps do not appear to be strongly influenced by the presence of oxygen. The hydroxyl groups which are formed associatively recombine to form water and surface oxygen, thus completing the overall cycle. The ethylene C-H bond activation and vinyl/acetate coupling reactions appear as if they may be likely rate-limiting steps. 50

Heterogeneous Mechanism (H2 0 Formation)

z ..= .,~

VAM Desorpt ion

;~

/

H.vdmx yl Recombination

+20

Acetic Acid Associative Adsorptiotz

I

-58

-50 ~,~

+67.5

VAM formation acetate and vinyl su~. ace reaction

I

+30~Water [ Desorption. I .

I - 100

- 150

Ethylene Ad .... rption

61

/ ~Ethvlene +47 1 ......~ace reacti~176 vinyl

t Reaction Path Energetics

Fig. 20. The overall catalytic route for oxygen-assistedvinyl acetate synthesis on a model Pd(l I 1) surface.

V. Summary and Conclusions First principle quantum chemical methods have reached the stage where they can now begin to provide reliable information on the structure, spectral and energetic properties of adsorbates on surfaces. Both the cluster, as well as the periodic band methods, were found to be quite successful. In contrast to early MO-based cluster studies, the DFT cluster methodology provides more accurate predictions, provided that special care is taken to optimize the structural configurations, the geometries, and spin states for all dusters. These results are subsequently used to elucidate the most favored adsorption sites, the most favored adsorption modes, and overall reaction energies for specific surface reaction steps. From this data, we can begin to construct overall catalytic cycles and examine their likelihood in carrying out proposed process chemistry. More detailed reaction coordinate searches are required to analyze the actual mechanism, where transition states and corresponding activation barriers are rigorously computed. While the transition states found herein appear to be quite reasonable, the predicted activation barriers are slightly high. This may be the result of cluster size effects rather than the DFT accuracy. By establishing the structure of the relevant adsorption and transition states, it becomes significantly easier to begin to develop structure-property relationships. The effects of bimetallics, adsorbate substituents, transient surface precursors were all examined herein on a series of commercial relevant model catalytic systems. The results have helped to elucidate the effects of gold on acetate decomposition, electron-withdrawing substituents on ]3-hydride elimination, and transient oxygen precursors for ammonia and acetate activation. More information, however will be required before we can establish more general structure-reactivity relationships.

31

VI. Future Directions Much of what has been discussed has targeted intrinsic adsorbate-surface bonding and reactivity. These efforts have been of great value in determining binding energies, activation barriers, and substituent effects on ideal surfaces. The chemistry that occurs under industrial conditions, however, can be substantially different due to competing extrinsic phenomena. The dynamics of the surface under actual processing conditions can involve a complex competing physicochemical phenomena. Defects sites, metal-support interactions, surface reconstruction, local coverages and surface compositions, solvent interactions, and deactivation can all affect the intrinsic adsorbate-surface chemistry. The ability to elucidate and quantitatively model each of these effects along with adsorbate-surface bonding holds a great challenge for computational chemical modelling. Increased CPU resources, faster processors, increased RAM, code parallelization, however, will dramatically decrease computational time and efforts. This will greatly increase the size of the models that can be examined and push us closer to more realistic situations which begin to embody these extrinsic effects. Three distinct algorithms will likely play a role in modelling these features. They are classified as first-principle, empirical and hybrid methods. While first-principles models offer the most accurate and rigorous solutions, they are often limited in terms of the size of the systems that can be examined. When coupled either directly or indirectly with empirical models they can be used to solve much larger problems. For example, ab initio molecular dynamic simulations enable the dynamic simulation of solvent interactions in homogeneous catalytic systems. DFT/MD approaches offer hope toward modelling surface dynamics for some systems. Car-Parinello algorithms are already being used to treat dynamic relaxation of transition metal surfaces. However, even with the advanced resources available in the future, direct dynamic modelling of all the physicochemical surface phenomena will be likely beyond reach. The time scales for chemical reaction phenomena preclude direct dynamic simulations. Hybrid quantum chemical/molecular mechanics and quantum chemical/dynamic Monte Carlo methods, however, which can decouple the dynamics and the detailed electronics stand to offer a much better chance for modelling a series of complex competing surface phenomena and will likely lead the efforts toward more realistic chemical modelling of surface chemistry. The greatest contributions for modelling will most likely come from combined experimental and modelling efforts which complement one anothers' strengths.

Acknowledgments I would like to thank Dr. George W. Coulston (DuPont), Mr. P. S. Venkataraman (Virginia), Dr. David A. Dixon (Pacific Northwest Laboratories), Dr. William D. Provine (DuPont), Dr. Carlonda Reilly (DuPont), Dr. Jan J. Lerou (DuPont), Dr. Kerwin Dobbs (DuPont), Mr. David Kragten (Eindhoven University of Technology)., and Professor Rutger A. van Santen (Eindhoven University of Technology) for their helpful technical discussions. I would also like to thank DuPont Chemical Company and the Donors of the Petroleum Research Fund of the American Chemical Society for their financial support of this work.

32

VII. Literature Cited

.

3. 4. 5. 6. 7. 8. .

10.

11. 12. 13. 14. 15. 16. 17. 18. 19. 20. 21. 22. 23. 24. 25. 26. 27. 28. 29. 30. 31. 32. 33. 34. 35. 36.

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91997 Elsevier Science B. V. All rights reserved. Dynamics of Surfaces and Reaction Kinetics in Heterogeneous Catalysis G.F. Froment and K.C. Waugh, editors

Molecular Studies of the Mobility of Surface Adsorbates During Catalytic Reactions

35

Metal

Atoms

and

G. A. Somorjai and G. Rupprechter Deparmaent of Chemistry and Center for Advanced Materials, Materials Sciences Division, E.O. Lawrence Berkeley National Laboratory, University of California at Berkeley, Berkeley, CA 94720 1. I N T R O D U C T I O N According to the classical view of catalytic surface reactions the reacting molecules adsorb on a surface where each atomposition is well-defined. After bond rearrangements in the adsorbed surface layer that was thought to behave as a two-dimensional gas the products desorb and the surface is available again for the next turnover. Modern surface science techniques that permit atomic scale monitoring of the structures of both the adsorbates and the substrate metal atoms during chemisorption and during the catalytic reaction have shattered the myth of the rigid catalyst surface waiting for the reaction to occur. Low energy electron diffraction (LEED) surface crystallography and scanning tunneling microscopy (STM) revealed the restructuring of the metal surface atoms on terraces and at defects (steps and kinks) during chemisorption and catalytic reactions. STM detected the mobility of atoms and molecules in the adsorbed monolayer. Sum frequency generation (SFG) - surface vibrational spectroscopy showed that the chemisorption of certain reactants disrupts the substrate surface structure (corrosive chemisorption) and that some of the strongly chemisorbed species are mere spectators during the catalytic turnover. The molecules that react during catalytic hydrogenation for example are weakly adsorbed and appear only at higher reactant pressures after all the strongly adsorbing sites on the metal surface are occupied. In this paper we show experimental evidence for the dynamic character of the catalytic surface on which the reaction occurs and for the mobility of both adsorbate and substrate atoms or molecules during the catalytic process. First surface science techniques (STM and SFG) that permit monitoring of the catalyst surface during the reaction are introduced. Then we discuss the restructuring of the metal substrate, the surface defects and the adsorbate layer. Finally we describe molecular level studies how platinum catalyzes the hydrogenation of light olefins (C 2(24) and the oxidation of CO and suggest mechanisms for both types of reactions in light of the, newly aquired experimental data. 2 . STM AND SFG STUDIES OF PLATINUM SURFACES AT H I G H R E A C T A N T PRESSURES During the last two decades a number of surface science techniques were developed that provided atomic level information on atomic or electronic surface structure as well as surface composition with ever improving spatial and time resolution [1]. LEED, XPS, AES, and SIMS have been routinely applied for surface characterization of single-crystal or thin film model catalysts used in the course for catalytic reaction studies. However, their drawback has been the need for a high vacuum working environment. To overcome this restriction, UHV compatible high pressure cells were developed [2,3] that combined capabilities to perform catalytic

36 reactions at high pressures using small surface area (--1 cm 2) model catalysts with surface characterization in ultrahigh vacuum (UHV) before and after the reaction, along with surface cleaning by ion bombardment and the deposition of additives by vapor phase condensation, respectively. On the other hand, this approach does not allow in situ studies. During the past five years, particularly two new surface science techniques proved to be capable to bridge the pressure gap; that is, they can be used to obtain molecular level surface information during chemical change at both low and high ambient pressures: Scanning tunneling microscopy (STM) [4] and infrared-visible sum frequency generation (SFG) surface vibrational spectroscopy [5]. Both of these techniques can operate within a 14 order of magnitude pressure range (lO"~-lO'Torr)without significant change in signal quality in terms of spatial or energy resolution. STM readily detects changes in atomic surface structure, while SFG yields the vibrational spectra of adsorbates often with about 1% of a monolayer sensitivity and with a 5 cm ~ energy resolution even at atmospheric pressures. Using these two techniques, we can monitor both substrate and adsorbate structures during reactions at high pressures. 2.1.

Scanning Tunneling Microscopy

(STM) at High Pressures

The first high pressure STM developed in our group was enclosed in a reaction cell, and was designed using piezo-electric tubes to control the motion of the scanning tip. The model catalyst Pt(111) surface had to be cleaned and characterized in a separate UHV chamber equipped with ion bombardment, AES and LEED. In order to transfer the clean Pt sample to the high pressure reaction chamber, it was was then placed in a small vacuum "suitcase" which was pumped with an ion pump at 10 -6 Torl" pressure. The surface was additionally protected by a passivating monolayer of sulfur after exposure to H:S. The protective sulfur layer was subsequently removed in the reaction cell by heating in oxygen at 1 atm pressure before annealing the surface while in atmospheric pressures of gases. STM Vessel i

I!

1

UHV Chamber

i

B

.............[,,1..........

F

~ii~~

l

Figure 1. Schematic diagram of the STM operating inside a high pressure chemical reactor which is attached to a UHV surface characterization chamber.

37 Our new design is a system which directly combines a UHV sample preparation chamber with an STM vessel (--2 L) that can be operated at controlled pressures in the range of 1 0 9 m l 0 3 Tort. This apparatus is shown in Fig. 1. A turbopumped airlock makes possible in vacuo sample transfer to and from a separate UHV chamber using a transfer rod. After cleaning in UHV, the single-crystal sample is transferred into the variable STM reaction chamber and imaged under UHV conditions. An infrared spot lamp located outside the chamber was used to heat the sample (through a viewport) up to 1400 K. Controlled pressures (tap to 1 atm) of gases are then introduced and the sample is re-imaged. The system was equipped with a quadrupole mass spectrometer which can be differentially pumped for gas analysis while the chamber is pressurized. The STM is of the walker type Besocke design that provides improved thermal stability as compared to our previous STM design. With the development of this specialized version of an STM, we can study the atomic-level structure of surfaces in situ, under controlled environments, with pressures ranging from UHV to atmospheric and temperatures from 300 to 425 K [6,7].

2.2. STM Studies of Adsorbate-lnduced Restructuring of Pt STM at high pressures clearly reveals the ability of the metal surface to rearrange in different ways depending on the type of molecule that is chemisorbed. In an atmosphere of hydrogen the Pt(110) face has a missing-row reconstruction surface structure as shown in Fig. 2a. The images consist of clearly separated parallel rows (hills) running for several hundreds of nanometers along the [110] direction. When oxygen is substituted for H z, again at one atmosphere and at 300 K, the surface appears to be very different form the surface in hydrogen. Instead of the nested, small scale reconstruction, we find the surface dominated by larger (100-300 A) structures (Fig. 2b). The larger hill and valley structure is due to enlarged (111) microfacets. The surface is actually comprised of enlarged (111) microfacets separated by steps, which are not resolved because of tip convolution effects and the angle of the steps relative to the tip motion. It is clear that we can obtain good quality images of flat surfaces and yet have difficulty deducing the structures of moderately rough surfaces. The surface in equilibrium with carbon monoxide is again different from that in either hydrogen or oxygen. On the scale of 1 to 30 A, the surface appears to be smooth relative to the hydrogen case, showing no hints of missing-row reconstructions (Fig. 2c). On a larger scale, we see features, but of a much different nature than the surface in oxygen, appearing as though the surface has formed multiple height steps. This is in agreement with UHV studies, which show that carbon monoxide rifts missing-row reconstructions. In this case, we can expect to see flat (1 x 1) terraces separated by steps due to the displacement of platinum atoms to or from nearby terraces. It should be noted that the reconstructions were stable before and during heating to 425 K for 5 h. As well, the reconstructions were present after evacuating the reactor cell, just prior to introducing carbon monoxide. Consequently, as the adsorbate is altered, so does the arrangement of surface metal atoms. Surface defects like steps and kinks also change their structure upon chemisorption. For instance, when sulfur is adsorbed on a stepped P t ( l l 1) crystal surface it forms a p(2 x 2) ordered structure, and induces doubling of the terrace widths and step heights [8]. Upon deposition of carbon the steps become highly kinked [9].

38

"

0i ..........ff If

.......

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

h yd. :I::os n : }1,7 a:t:;:m..: ~:~:"~ ~o' 0 A' x: '~'0~ )...... A.

oxyg~

i ~.t:m,

"T~ested" mi,.~sil~g-l'ow ~-eco~str.uclior~s

9 0 0 jii, x "~S0 ..?i.....

(111) mic:,::>facets,

k [Ixo]

[xxo]

[oo~]

c c~,rbon m o n o x i d e " 1. a:tm~, 7 7 0 A x ~4:(~A~

u t ~ t ' e c : o ~ s t ~'tJcted ( l •

[er'l'acos

separale,::t by mult. iple h e i g h t st:ePS

Figure 2. a) Topographic image of the Pt(110) surface in 1.7 atm of hydrogen after heating to 425 K for 5 hours, showing (n x 1) missing-row reconstruction randomly nested. Image size: 730 A x 700 A. Vertical range: Az = 10 A; b) Topographic image of the Pt(ll0) surface in 1 atm of oxygen after heating to 425 K for 5 hours. Image size: 900 A x 780 A. A~ = 25 A; c) Topographic image of the Pt_(110) su.rface in 1 atm of carbon monoxide after heating to 425 K for 4 hours. Image size: 770 A x 740 A. A~ = 42 A.

2.3. STM Studies of Adsorbate Monolayer Restructuring by Coadsorption The adsorption of coadsorbates may cause the rearrangement of the adsorbate overlayer as well as of the underlying substrate atoms in order to accommodate all surface species present. A reconstruction of adsorbate overlayers by coadsorption has been demonstrated in a STM and I.F:.ED study of sulfur chemisorption on Pt(111) surfaces [10]. Surface structures imaged by STM are consistent with electron diffraction patterns obtained in complementary LEED studies (Fig. 3a). At a sulfur coverage of 0.25, a p(2x2) sulfur structure can be observed. The coadsorption of carbon monoxide molecules induced a reordering of the sulfur structure (Fig. 3b). The sulfur overlayer is compressed into an ordered structure of higher local coverage, creating empty space on the surface for carbon monoxide adsorption, i.e. (~/3 x ~3)R30 ~ structure. New terraces form which contain exclusively CO and are separated by monoatomic steps from the terraces containing the compressed sulfur overlayer. The CO molecules have a high surface mobility and are not visible in STM experiments. The change of the sulfur overlayer structure is reversible and the original p(2x2) structure can be restored after desorbing CO at high temperature.

39 Competitive adsorption and mobility of adsorbates on the surface attribute to the coadsorption-induced reconstruction of adsorbate overlayers. If surface species is immobile because of a high activation energy for surface diffusion, coadsorption cannot take place. On the other hand, the adsorption energy of one adsorbate must be sufficient to compress the other adsorbate into a more compact layer. If coadsorbates have a very low heat of adsorption, the thermodynamic driving force for adsorbate overlayer reconstruction is absent.

Figure 3. STM images (a) before and (b) after the reordering of sulfur overlayer on Pt(111) induced by CO exposure.

2 . 4 . Sum Frequency Generation (SFG) Infrared-visible SFG, an interface specific vibrational spectroscopy technique, was utilized to monitor the chemisorbed molecules over a ten orders of magnitude pressure range, up to one atmosphere. SFG is a three wave mixing process which involves the addition of infrared (r and visible light (r to produce light at the sum of these two frequencies (c%=r [11,12]. The visible beam is held at fixed frequency while the infrared beam is tuned over the vibrational range of interest. The electromagnetic radiation source for the experiment is generated by a passive, active mode locked Nd:YAG laser which outputs a 20 ps pulse at 1064 nm when using #5 dye from Kodak. The repetition rate of the laser is 20 Hz and the total energy per pulse is approximately 50 rnJ. The energy is divided into three portions. The fn'st of which is frequency doubled from the 1064 nm Nd:YAG fundamental using KD*P to produce green light at 532 nm. The remaining light is employed to generate infrared light by one of two angel tunable optical parametric generator/amplifier (OPG/OPA) stages capable of producing infrared radiation over most of the range from 1100--4000 c m x. The first OPG/OPA generates light from 2600 cm 1 to 4000 cm 1 with FWHM of 12 cm 1 at 2880 cm -1 using LiNbO 3, while the second OPG/OPA is angle-tunable from 1100 cm 1 to 2400 c m 1 with a FWHM of 4 c m 1 by difference frequency generation in a AgGaS 2 crystal [13]. The infrared and visible beams are temporally aligned and focused concentrically on the late transition metal catalyst inside the batch reactor. The spot size of the green beam is approximately 1.55 mm FWHM at the sample and the beam makes an angel of 50 ~ with respect to the surface normal. The input energy is approximately 400

40 gJ. These values yield a total power of approximately 0.8 gigawatts/cm 2. This regime has been shown experimentally to be at least a factor of 4 below powers that cause either photochemistry or damage the platinum sample even over extended periods of time. The infrared radiation is focused to approximately 1.2 mm inside the green spot at an angle of 56 ~ with respect to the surface normal and the maximum energy out of the OPG/OPA is always used. The infrared power varies widely with frequency. The maximum infrared power is a 2850 cm -1. At this frequency the IR energy at the sample is approximately 275 I.d/pulse when using Kodak #5 dye (approximately 20 ps pulses). A vibrational spectrum of surface species is obtained by monitoring the sum frequency signal as a function of the incident infrared photon energy, but it should be noted that for a vibrational mode to be observable by SFG, it must be both infrared and Raman active. Only those modes which lack centrosymmetry can in the dipole approximation simultaneously obey both rules. Therefore, in the experiments described in this paper the (isotropic) gas phase and the fcc lattice of the bulk platinum sample possess inversion symmetry and give nearly zero contribution to the signal. Henceforth, the dominant contribution is generated by the modes of the adsorbed monolayer at the platinum surface, where inversion symmetry is always broken [ 14]. The Pt(111) single crystal was cleaned by conventional surface science procedure in a UHV chamber. The sum frequency spectra were taken at room temperature from 10 1~ Torr to 700 Torr CO pressure by successively increasing the gas phase pressure. A batch reactor coupled via a gate valve to an UHV system was fitted with CaF 2 windows to allow infrared and visible light to pass into it and to allow sum frequency light to pass out to a photomultiplier tube (Fig. 4). The reactor could be charged to 1.5 atmospheres of total pressure and was equipped with a re.circulation pump in the reaction loop. The loop contained a septum for gas abstraction and analysis by a gas chromatograph. The mixing time of the chamber was determined to be approximately 3 minutes by following the dilution of a plug of methane gas mix into a background of pure Ar. The reactor volume was 62.3 liters and could be pumped out to a base pressure of 5 • 10 11 Torr by opening the gate valve to an ion and turbomolecular pump. The sample could be resistively heated to 1000 K, and cooled to 120 K under vacuum using liquid nitrogen (250 K under ambient pressure). When the reactor was under vacuum, typical surface analysis techniques such as retarding field analyzer for Auger and LEED were available. An A r + gun could be used to sputter the crystal clean.

Nd:YAG I 1064nmJhonlinea- "" ~100-1100cm-1\ Laser 50mJIpps ulse~120 r

I-Detect~ Computer

/

/'~

!I _- ( ~

~111)/ Z~-~

Septum ..... -

~

UHVIBatch Reactor

Figure 4. The UHV/batch reactor coupled to an Nd:YAG laser for in situ SFG studies.

41

2.5. Corrosive CO Chemisorption on P t ( l l l ) at High Pressures Studied by SFG and STM During infrared-visible SFG studies of carbon monoxide on Pt(111) in the 101%700 Torr pressure range, dramatic changes in the vibrational spectrum of surface species were observed at high CO pressures. The vibrational signatures of the adsorbed species implicate the formation of carbonyl-platinum cluster analogs coadsorbed with an incommensurate CO overlayer. Figure 5 shows the SFG vibrational spectra of carbon monoxide obtained at 107-700 Torr of CO and at 295 K. When the clean P t ( l l l ) surface was exposed to 10 L (1 L=10 -6 Torr sec) of CO in UHV, two peaks at 1845 cm 1 and 2095 cm 1 were observed which are characteristic of CO adsorbed at bridge and atop sites. I~EED revealed that a c(4 x 2) structure was formed in which an equal number of carbon monoxide molecules occupied atop and bridge sites [15]. Such results are in agreement with previous HREELS [16] and reflection-absorption infrared spectroscopy (RAIRS) [17] studies. The much higher relative intensity of atop bonded CO to bridge bonded CO in the SFG spectra is due to the specific selection rule for the SFG process [ 18]. As mentioned earlier, SFG is a second order, nonlinear optical technique and requires the vibrational mode under investigation to be both IR and Raman active, so that the SFG intensity includes contributions from the Raman polarizability as well as the IR selection rule for the normal mode. Raising the pressure to 1 Torr, the bridge bonded CO disappeared and the frequency of atop species shifted to a higher value (2105 cml). This is accompanied by an increase in the atop concentration with a corresponding decrease in bridge bonded CO coverage- as found on the same system at high CO exposure in UHV [19]. The frequency shift is the result of the enhancement of the dipole coupling between CO molecules on the surface because of their closer packing at high pressure, thereby weakening their bonds to the metal [20]. Surprisingly, by further increasing the pressure to 11.5 Torr, the intensity of the atop CO was attenuated without further frequency shift or broadening; at the same time, two new features of 1980 cm ~ and 2045 cm 1 became visible. At 150 Torr of CO and higher (a pressure region that is not readily accessible to HREELS and IR spectroscopy), a broad peak coveting 1980 cm 1 and 2045 c m 1 features appeared and extended to 1800 cm 1 with a monotonic decrease of the intensity. At 700 Torr of CO, the spectrum was dominated by the new features and a strong background extended to 1700 cm 1 again with a monotonic decrease in intensity. The spectra corresponded to the changes in the nature of bonding and concentration of the adsorbed CO and were completely reversible and reproducible with variation of the gas pressure. Raising the carbon monoxide pressure caused the adsorption of more CO on the Pt(111) surface and, hence, a compressed overlayer was formed. This increased the repulsive interaction between CO adsorption as judged by the rapidly declining heat of adsorption with coverage [21]. Under pressure, the interactions between CO molecules and platinum atoms are strong enough to overcome the CO-CO repulsion. The carbon monoxide molecules on the surface lost their registry of sites and formed an incommensurate overlayer as lower symmetry locations were occupied. Indeed, the strong background seen in the 1800 cmX-1980 cm 1 frequency range was probably the result of the formation of an incommensurate overlayer due to the ambiguity of site occupation. The presence of the 2045 cm 1 adsorption peak is intriguing. The frequency is characteristic of terminally bonded CO (N-CO) molecules; therefore one reasonable explanation could be that atop CO has been squeezed out of its normal site and caused the frequency to shift to the lower value. The intensity of atop CO decreased with increasing pressure, while the 2045 cm ~ feature intensity went up. Another interpretation of the 2045 cm ~ peak is the multiple bonding species formation. This frequency is comparable to the vibrational frequency of terminally-bonded CO in platinum carbonyl binary complexes Pt-(CO)n (n=1,2,3,4) synthesized by Ozin et al. by co-condensation of platinum atoms with carbon monoxide molecules in an argon matrix at very low temperature [22]. We prefer the assignment of this peak to the multiple bonded platinum carbon monoxide

42

cluster analogs on the surface. But, in our case, these cluster analogs were affected by the platinum substrate underneath, so the CO vibrational frequency is slightly red-shifted from Ozin's binary complexes.

700 torr

ai v c

150 torr

~,~,

9

9

o

o

~I~!

U. CO

o~

N

0 0

E 0 Z

rr

O0

0

O

O @

1.0 torr

0

0

1600 1700 1800 1900 2000 2100 2200

100 A x 100 A

Frequency (cm-1)

Figure 5. left: Pressure dependence of SFG spectra for CO adsorbed on Pt(111) at 295 K. The spectra change indicates that new surface species appeared at high pressures. right: Scanning tunneling microscope pictures of carbon monoxide over the Pt(111) crystal face at three different pressures: 1.0 Torr, 100 Torr, and 500 Torr.

43 The spectra for elevated CO pressure were consistent with CO chemisorption accompanied by displacive reconstruction (i.e., distortion or relaxation of the transition metal surface) [23]. In this model some of the Pt atoms could be pulled out from the crystal lattice, perhaps at a step or kink site, and become more exposed to the adsorbing CO molecules; thereby forming multiple bonded CO platinum complexes (Ptm(CO)n with n/m >1) or carbonyl analogs. However, in our case, because of the repulsive interaction of CO ligands, these complexes are not stable. They can only exist on the surface in the presence of high CO pressure. Most catalytic reactions require high reactant pressures, in the range of stabilization for the cluster-like species. As will be described below, our SFG studies during CO oxidation on the Pt(111) crystal face indicated the CO vibrational spectrum to be characteristic of the cluster-like species. High pressure STM studies confirm the dramatic changes on the Pt(111) surface due to CO chemisorption at high pressures. Figure 5 shows the surface structure observed as the CO pressure is increased. It appears that the steps on the metal surface become diffuse and the surface species that are imaged can be readily moved by the STM tip. Thus the platinum carbonyl clusters that form at high coverages are highly mobile. There are also structures that exhibit a moire pattern as shown in Fig. 5. These are likely caused by the incommensurate CO overlayer. 3. STM AND SFG STUDIES OF OLEFIN H Y D R O G E N A T I O N

3.1. Ethylene, Propylene and Isobutylene Hydrogenation on Pt(lll) Studied by SFG In this section we show results of SFG studies of light olefm hydrogenation which reveal the n-bonded species of ethylene, propylene and isobutene as the dominant intermediates which subsequently hydrogenate to alkanes at 300 K and atmospheric pressures. The other species present during this high turnover catalytic reaction, the strongly chemisorbed ethylidyne, propylidyne and isobutylidyne and the corresponding di-cy bonded species, are spectators that hydrogenate much too slowly to be important sources of the hydrogenated products. The nbonded reaction intermediates are generally present in concentration of a few percent of a monolayer and are desorbed as the reactant pressure is decreased, indicating relatively weak surface bonding. Ethylene hydrogenation on platinum surfaces is a very simple catalytic reaction and takes place at room temperature. A reaction mechanism proposed by Horiuti and Polanyi [24] suggests that ethylene is hydrogenated stepwise by H atoms on the surface. In the absence of hydrogen, several distinct species have been identified in UHV conditions for an ethylene monolayer absorbed on Pt(111). At temperatures below 50 K, monolayer molecules interact weakly with the surface via ~-coordination. Ethylene is physisorbed with its C-C bond parallel to the surface, and interatomic distance between two carbon atoms is almost unchanged with respect to the gas-phase molecule. At temperatures between 60 K and 240 K, the carbon atoms attain nearly s p 3 hybridization and bond directly to the substrate. This is referred to as di-~bonded coordination. A further annealing of the surface induces dehydrogenation of ethylene molecules. Ethylidyne is formed by losing one hydrogen and transferring a second hydrogen atom to the other carbon, and is the prominent hydrocarbon species on the surface at temperatures of up to 450 K. Several groups have studied the reaction intermediates in the hydrogenation reaction using infrared spectroscopy [25,26]. However, they were all carried out in the absence of ethylene in the gas phase. By contrast, SFG allows us to monitor the surface in situ under catalytic conditions [27]. The conventional turnover rate (TOR) of ethylene hydrogenation on Pt(111) at 295 K with 100 Torr H 2 and 35 Torr C~H4 per platinum atom as determined by gas chromatography is 11 +1 ethylene molecules converted to ethane per second. Three features were present on the surface under these conditions, at 2880 cm -1, 2910 cm -1 and a small peak

44 around 3000 1 (Fig. 6a). Compared with fingerprint spectra of surface species prepared in vacuum, the largest peak in the vibrational spectrum at 2880 cm 1 is the vs(CH 3) of ethylidyne (M=CCH3), a decomposition species which resides in the fcc three fold hollow site and is present on the surface during reaction. The feature at 2910 cm x originates from the vs(CH 2) of chemisorbed ethylene. The species is usually referred to as di-cy bonded ethylene because the carbon-carbon double bond of the gas phase molecule is broken and two sigma bonds are formed with the underlying platinum surface atoms. The small peak just below 3000 cm 1 is the vs(CH z) of weakly bonded (physisorbed) ethylene. The molecule's n orbital bonds directly with the underlying platinum surface and is often called n-bonded ethylene. The weak vs(CH z) signal for rt-bonded molecules can be attributed to the surface-dipole selection rule for metal surfaces, that dynamic dipoles parallel to the surface plane are canceled by image dipoles inside the metal. The spectrum was stable for hours while the reaction rate remained constant. After the gas mixture was evacuated from the reactor, the surface was saturated with ethylidyne, while dio and n-bonded ethylene molecules disappeared (Fig. 6b). After the reaction cell was recharged with an ethylene/H 2 mixture, the reaction rate recovered (as shown in Fig. 6c) and n-bonded ethylene was restored on the surface, while the intensity from the di-cy species did not return. This demonstrated a direct competition for adsorption sites between the ethylidyne and the di-o ethylene. Since a high turnover rate was again observed in the second run in absence of di-obonded species, the hydrogenation through the di-G ethylene is at most a minor channel in the reaction scheme. It was also observed that the ethylene hydrogenation reaction occurred at the same rate regardless of the presence of ethylidyne on the surface. In addition, preadsorbed ethylidyne groups do not effect the reaction rate. This indicates that ethylidyne is not directly involved in the hydrogenation reaction. The same conclusion was reached from the results of a transmission infrared spectroscopy study of 14C ethylidyne hydrogenation in hydrogen [28]. In contrast to the behavior of di-cy-bonded ethylene and ethylidyne, the appearance of rtbonded species is directly correlated with the reaction rate. It is most likely the key intermediate in ethylene hydrogenation. It should be emphasized that the other surface species present, di-cy bonded ethylene and ethylidyne, are spectators and do not contribute to the turnover rate in any significant way [29]. Interestingly, only these strongly chemisorbed species are normally detectable by studies in UHV, while the weakly bond n-bonded ethylene is more readily found at atmospheric reaction conditions. Surface ethyl groups are visible at very high hydrogen pressures. Under 727 Torr H z and 60 Torr C/H4 at 295 K, two additional peaks are present on the SFG spectrum (Fig. 7). These can be assigned to an ethyl species on the surface. The low surface concentration of ethyl groups at medium hydrogen pressure suggests a high degree of reversibility with the addition of the first hydrogen into adsorbed ethylene. UHV spectroscopic measurements were conducted to calibrate the concentration of ~bonded ethylene on Pt(111) under reaction conditions. The calibration was achieved by exposing the clean Pt(111) surface to a near-saturation coverage of oxygen at room temperature, followed by exposure to ethylene at 120 K [30]. This results in a mixture of rt-bonded and diG-bonded ethylene on the surface at a concentration of 6% of a monolayer and 10% of a monolayer respectively [31]. Using this spectrum as a reference, the intensity of the 3000 cm 1 peak correspond to approximately 4% (error bars from 2-8%) of a monolayer of reactive nbonded ethylene (the error bars for this result are a factor of 2 at the 90% confidence level using worst case assumptions about C-H bond reorientation). This means that the turnover rate for ethylene hydrogenation is actually 25 times faster per reactive intermediate species than when estimated per exposed platinum atom. Therefore, the absolute turnover rate of physisorbed ethylene is approximately 275 ethane molecules formed per surface intermediate per second under the above conditions.

45

a

p. ,,,,~,,, O; ~ i

TOF = 11 •

~,,~

9

c;' ~

o

_z,

0 o :

p

r

~ , ~ ~ ~

r H\

6

-------

do 9

oo

I

2850

n c~c:.a

.

s

oo .o~b

%00000000000 o

I

1

I

2900

2950

3000

%00

301

Frequency (in cm "1)

,~. ,,~,,,, TOF =12• ~,,, ,,,,

o .~

:'o 9

~

ethytidyne

b~ 9

~

O' ~

~

.

~

~

c;

o~ i o

:

oo

lk,

:

q

: " 0

iS

I

2750

2800

I

2850

!

2900

Frequency (in cm 1)

I"

2950

'

3000

!

2850

~

lk

?~ I

2900

_...u

o~176 %0o0~ o

I

1

2950

3000

Frequency (in cnf I)

Figure 6. a) SFG spectrum of the P t ( l l l ) surface during ethylene hydrogenation with 100 Torr I-Iz, 35 Torr CzH 4, 615 Torr He at 295 K; b) The vibrational spectrum of the same system after the evacuation of the reaction cell; c) SFG spectrum under the same conditions as a), but on a surface which was pre-covered in UHV with 0.52 monolayers of ethylidyne.

3050

46 The concentration of surface intermediates is too high for this reaction to be dominated by defects. In fact, it has been demonstrated that the rate of ethylene hydrogenation is independent of the platinum surface structure over which the reaction takes place (a so-called structure insensitive reaction) [32]. Evidence from inorganic cluster analogs and homogeneous phase catalysis suggests that the reaction probably occurs on atop sites [29]. Hydrogenation at such sites would lead to a structure insensitive reaction, because they are available on all crystal faces. The concentration of surface intermediates is much lower than one intermediate per surface platinum atom and remains constant at approximately 4% of a monolayer over a wide range of ethylene pressures. Therefore, this represents saturation coverage of =-bonded ethylene under reaction conditions, consistent with the kinetics of ethylene hydrogenation which are nearly zero order in ethylene partial pressure above approximately 25 Torr of C2H4 [33]. It has been observed that ethylene hydrogenation is unaffected by decomposition species present on the surface during reaction [34,35]. Because the concentration of reaction intermediates is only 4% of a monolayer and the saturation coverage for the decomposition species, ethylidyne, is only 25% of a monolayer, it appears that there should always be sufficient sites available on which the reaction may occur.

TOF

- 61

O0 r

ethyl

r

U

0

og

%~

9

~

=- ethylene

b

o%

oo%0

I

2850

I

2900

I

2950

oo POo

bOOo o

I

3000

3050

Frequency (cm "1) Figure 7. Hydrogenation of ethylene on Pt(111) with 727 Torr H 2 and 60 Torr C2H 4 at 295 K. The two peaks marked with arrows are features characteristic of an ethyl species.

47

ethylidyne

ethyl

x - ethylene

H~--H

H,~

/H

/C

%

1

1

(1)

(2)

(3)

. . . . e - . ~ . ".~....................... N~

I .

.

.

.

/

I x

C:

Figure 8a. Proposed mechanism for ethylene hydrogenation on Pt(111).

c~

48 Clean P t ( l l l ) dcq dl2 = 229 +0.01A I

~ 1

do~

tb,

t ~////A//i,'////J'///A

d23 = 227 + 0.03A

d12

Ib~

d ~ = 2.23 + 0.04,~

d23

d b = 2.26~, Top View d c c = 1.49 + 0.05

do1 - 1.21 + 0.03A d 12 " 230 + 0.03A q = 0.1 + 0.08~. r~ = 0.0 +0.09&

d23 = 2.26 + 0.04A, b~ = 0.11 + 0.05A b 2 = 0.08 + 0.09

Figure 8b. The surface structure of ethylidyne adsorbed on Pt(111) as obtained by low-energy electron diffraction surface crystallography. ' We therefore propose a model in which n-bonded ethylene adsorbs on atop sites where it can be hydrogenated to ethane in the presence of mechanistically unimportant decomposition species. Fig. 8a shows a schematic representation of the hydrogenation process. LEED surface crystallography studies determined that both ethylidyne and di-o-ethylene species occupy threefold, face centered cubic adsorption sites with structures shown as in Figs. 8b and 8c. After these strongly adsorbing sites are occupied, n-bonded ethylene adsorbs at atop sites. In the process, the ethylidyne species may change location by moving from fcc to hcp three-fold sites to leave more space for the n-bonded species. The rapid motion of ethylidyne at 300 K is indicated because the molecule cannot be imaged by STM, and calculations indicate a low activation energy for ethylidyne diffusion between the three-fold sites. After sequential hydrogenation to ethyl and then to ethane the catalytic hydrogenation process repeats. Propylene [36] and isobutene [37] hydrogenation have also been studied by SFG on Pt(111) at room temperature. Similar to ethylene hydrogenation, propylene and isobutene are hydrogenated stepwise on platinum by atomic hydrogens formed by H 2 dissociation on the surface. The reactions proceed via alkyl intermediates. An issue of interest is the regioselectivity of hydrogen addition to propylene and isobutene, i.e. the position selectivity of hydrogen atom addition to two nonequivalent carbon atoms in the C=C double bond. Hydrogen addition to the terminal carbon yields 2-propyl, while hydrogen addition to the internal carbon forms 1-propyl. Similarly 2-isobutyl and 1-isobutyl groups result from hydrogen addition to terminal and internal carbons of isobutene. Under reaction conditions of propylene hydrogenation, 2-propyl and r~-bonded propylene species were observed on the surface. Vacuum studies indicated that 2-propyl and 1-propyl groups had similar hydrogenation rates. This demonstrates that the 2-propyl group is the reaction intermediate and not just a spectator residing on the surface. 2-propyl groups are formed in a preferential addition of a hydrogen atom to the terminal carbon in propylene molecules (Fig. 9). Isobutene hydrogenation has presumably a different reaction mechanism. Although 2isobutyl groups and n-bonded isobutene were the dominant surface species under reaction conditions, the hydrogenation rate of 2-isobutyl was much slower than that of 1-isobutyl. As a

49 result, isobutene hydrogenation may be forced to proceed from u-bonded isobutene through the slow kinetic step of 1-isobutyl formation. It provides an explanation of a slow isobutene hydrogenation rate, which is at least one order of magnitude lower than for those of 1-butene and cis-2-butene (Fig. 9) [38]. The ability to monitor the surface concentration of active intermediates with chemical specificity during reaction should greatly contribute toward understanding the molecular details for a wide range of heterogeneous catalytic reactions. It will be invaluable to exploit SFG vibrational spectroscopy to explore the nature of the surface chemical bonds and the concentration of intermediate species for a wide range of catalytic reactions. Significantly, it has been shown for ethylene hydrogenation, that the surface species present under UHV conditions are different than under ambient reaction conditions. Indeed, this is probably the first of many reactions that will show significantly different surface chemistry at high pressures now that the pressure gap can be bridged.

top view

side view bc-c

,oo

v

fcc site

~_ ~,.~__J ~ a ~ _ ~

,

dl 2

dp

, do0

...bcc

site

fcc site

hcp site

o

bl

23~

1.92 A

+2'0 +0.15 A -11

22~

1.g6 A

b~

bCC

bu

dp

2.07 A

1.56 A

0.03 A

+0.30 ]t, +0.50 A +0.07 A

2.05 A

1.53 A

0.02 A

+21o +0.15A +0.30A ~0.40A • -11 o

doo,

0.83 A

0.61 A

-1-0.25 A

+0.07 A

0.74 A

0.57 h

do1 1.32 A

-t-0.12 A

1.35 A

dn 2.27 A +0.05 A

2.27 A

A +0.25A +O.OTA +0.12A •

Figure 8c. The surface structure of di-o ethylene on Pt(111) as obtained by low-energy electron diffraction surface crystallography.

50 H

I-I, H-C~

...4-I C--.4-I

+H )

\\\\\\\\\\\\\\\

\\x\\\\\\\~\X\\

n-bonded

9

+H )

I

---H

~\\\~\\\\\\\\\\

ethyl

ethylene

H

H H

HH

H \\\\\\\\\\\\\\\

+H )

H

H

H--C~ \\\\\\\\\\\\\\\

n-bonded propylene

+H )

9

H

\\\\~\\\\\\\\\\

2-propyl H

,

,,",,~,,~

n-bonded isobutylene

+H )

A~H H

~\~\\\~\\~\\\~\

)

H

/

\\\\\\\\\\\\\\\

isobutyl

Figure 9. Suggested reaction pathways for ethylene, propylene and isobutene hydrogenation on Pt(111).

51

3 . 2 . STM Studies of Propylene Hydrogenation and Thermal Decomposition (Coking) on P t ( l l l ) at High Pressures In this study the high pressure STM was applied to minitor surface structural changes of a Pt(111) model catalyst during propylene hydrogenation. As described in section 2.1., the metal surface was cleaned and characterized in a separate UHV chamber, and subsequently sulfurpassivated before transfer to the high pressure STM. After removing the protective S layer by heating in I arm of 02, the reactor cell was pumped down to 10.5 Torr and a 1 arm mixture of propylene (10%) and H 2 (90%) was admitted to the chamber with the sample at 300 K. It should be noted that, under these conditions, the Pt surface is active in catalyzing the hydrogenation of propylene to propane [39]. Propylene readily adsorbs on the surface of Pt at 300 K and forms ordered structures of propylidine (-=C-CH2-CH3) (which have been studied in detail) [40,41,42,43]. The STM Images show the surface consisting of flat terraces and monatomic height steps that are characteristic of the initially clean Pt(111) surface. No atomic resolution was obtained; again, probably as a result of the rapid mobility of the propylidyne and other possible small fragments. It is important to note that no change of Pt step morphology takes place during reaction conditions at 300 K. We proceeded to heat the surface in the hydrocarbon/hydrogen atmosphere using a infrared spot heater. We found that carbonaceous clusters did not form until after heating to 700 K. This is -150 K higher temperature than necessary to form carbonaceous clusters in vacuum [9]. This indicates that the dehydrogenation/polymerization reactions are inhibited by the presence of atmospheric pressures of hydrogen. This is likely due to the efficient rehydrogenation of any fragmentation product to more saturated and mobile hydrocarbons, such as propylene and ethylene or to fully hydrogenated propane, ethane and methane, that desorb to the gas phase. Methane desorption has indeed been detected by IR under these conditions [44]. A very different behavior is observed in pure H z. Clusters do not form in this environment until at least >900 K due to rapid hydrogenation of any C residues, including those produced by heating first in pure CO environments at 800 K [9]. STM images taken after heating the surface to 800 K in the presence of a 90% ydrogen~ 10% propylene mixture show 20-30 A particles forming elongated aggregates (200 x 75 A). The elongated nature of the aggregates is not understood, but could be due to coadsorbate (hydrogen and/or propylene) interactions, which force the hydrocarbons to the edge of the adsorbate islands. The experiments described have produced a number of significant new observations about the surface chemistry of hydrocarbons as they decompose when heated to increasingly higher temperatures under various ambient conditions on Pt(111). First, we noticed that the step structure of the platinum surface remains essentially unchanged when covered by the propylidyne monolayer formed by the room temperature adsorption of propylene. This is true both in UHV and under catalytic reaction conditions of atmospheric pressures of Hz/propylene mixtures. The stability of the step structure is maintained even when moderate decomposition, leaving small fragments, takes place (after heating t o - 5 5 0 K in UHV). Severe dehydrogenation, however, results in graphite formation and causes step pinning, a phenomenon observed previously by Land et al. using ethylene- although in that case the temperature used was 1200 K [45]. During the dehydrogenation studies [9], it was further discovered that a clean platinum STM tip, "activated" by pulses of several tenths of a volt, can be used for local catalytic hydrogenation or oxidation of carbonaceous deposits at 300 K on the nanometer scale ("tipcatalysis" [46]).

52

4. SFG STUDIES OF CO OXIDATION CO oxidation experiments were carried out at a total pressure of a few hundred Torr with various carbon monoxide to oxygen ratios. The gases were mixed in the reactor and the total pressure was brought up to one atmosphere using helium. The order in which CO and 02 were introduced did not influence the results. CO pressure was varied from tens to hundreds of Tort. The partial pressure of CO at a given total pressure was also an important factor. Therefore the chemistry of the reaction was explored by varying the CO/O 2 ratios at a given total pressure. The SFG spectra obtained during carbon monoxide oxidation at different platinum crystal temperatures (with an initial condition of 100 Torr of 0 2, 40 Torr of CO and 600 Torr of He) are shown in Fig. 10a.

=

9

v

%oR=2238

m

oo, O

r,~ IJ.

= m

642 K

c~o

o

"~

.0.'

o

TOR=1692

o

r,,O i J,.

o

590 K

r3'j It.

TOR=705

vm

540 K

r.~

b

%.===,

I,,l.. r4)

1800

TOR=28

[

I

[

]

1

1900

2000

2100

2200

2300

2400

Frequency (cm1) Figure 10a. Temperature dependence of SFG spectra observed during high pressure CO oxidation over Pt(111) at 100 Ton" 02 and 40 Torr CO. The temperature and turnover rate are also shown.

53

The reaction rate was very low when the sample temperature was low, and the SFG spectrum was dominated by adsorbed CO surface species at atop sites. With increasing temperature, the atop CO intensity decreased due to the thermal desorption of carbon monoxide and the reaction accelerated. At 540 K, a reaction rate of 28 molecules of COJplatinum surface site/second was obtained. When the platinum sample temperature increased to above 600 K, the reaction became self-sustained, proceeding to a constant high temperature without the need of heating because of the high exothermicity of this oxidation. The temperature at which the reaction becomes self-sustaining is defined as the ignition temperature. The ignition temperature is a function of CO partial pressure - this dependence is shown in Fig. 11.

O ,,.->

=

v m

t~ (.-

LI.. CO

11

R=4480

=

I

v i

C

750K ~

O3 IL CO

p

O

,

~, TOR=2685 1

,,-:..

=

c. r

~0K

O3 I,J. O0

,===~=====~r

6

'

6 (~'s

TOR=660 - r . t / ~ r t-I-, 9t - s

13t~ r x 3

o ..

=. v m

e-

03

590 K

6

TOR=243

IJ.. CO

I

1700

1800

1900

2000

1

l

2100

2200

2300

Frequency (cm1) Figure 10b. Temperature dependence of SFG spectra of high pressure CO oxidation over Pt(111) at 100 Torr CO and 40 Torr 02. The temperature and turnover rate are also shown.

54 The SFG spectra obtained at temperatures above ignition are dramatically different from those obtained in the low-temperature regime. Atop CO, which is the dominant feature in the low-temperature regime, disappeared completely above ignition temperature, while there were three new features at 2050 cm", 2130 cm" and 2240 cm1 that showed up as the turnover rate shot up to 2238 molecules/platinum site/second. The 2130 cm 1 peak was assigned to the stretch mode of carbon monoxide adsorbed on the oxidized platinum sites, which has been shown not to be important in CO oxidation [47]. The broad peak centering at 2050 cm 1 was assigned to a new CO species, which is the only CO species present on the surface in the high-temperature regime (>600 K). The assignment of the feature at-2240 cm ~ is difficult to make because the frequency is too high to be readily assigned to any of the surface CO species, since the frequency of CO stretch mode in gas phase is only 2143 cm -1. However, such a high frequency CO stretch has been reported when the molecule is bound to positively charged platinum. Consequently, because of the frequency regime, and the presence and the formation of carbon dioxide species on the surface, this feature could also be assigned to a surface CO2 related species. In the presence of excess CO (at 100 Torr CO and 40 Ton" O z, for example) the vibrational spectrum became considerably simpler. The spectral features we assign to CO adsorption at an oxidized Pt site and to CO 2 both disappeared. Only the peaks associated with the presence of incommensurate and terminal CO species are detectable in this circumstance. This is shown in Fig. 10b. The kinetics of CO oxidation changed with the change of the surface vibrational spectrum. Different apparent activation energies were observed for the two different temperature regimes, below and above ignition temperature. An activation energy of 42 Kcal/mol was obtained for the low-temperature regime below ignition temperature, which is similar to the CO desorption energy. Above ignition temperature, where the reaction is self-sustained, a value of 14 Kcal/mol was observed in the high-temperature regime (>600 K), which is close to the activation energy of 11.7 Kcal/mol obtained in molecular beam studies [48].

800 High Reactivity R

G)

"

.-"""

7oo~

E l'--

Unstable regime

e-

....."

._m_o 600 i ..,., e-

500 0.0

w Reactivity Regime

I

1

1

I

0.2

0.4

0.6

0.8

1.0

CO Partial Pressure

Figure 11. The CO partial pressure-dependence of the ignition temperature of CO oxidation over Pt(111).

55

A correlation was also shown between the reaction turnover rate and the coverage of the surface CO species. In Fig. 12a, the relative coverage of atop CO is plotted as a function of the turnover rate at a constant temperature of 590 K. The coverage was altered by changing the CO/O 2 ratio and was normalized to the low-temperature SFG intensity of the same species. With the decrease of atop CO coverage, the reaction rate increased greatly, indicating that atop CO is not a key intermediate during CO oxidation. Its presence on the surface inhibits the CO oxidation reaction. However, the reaction rate was proportional to the concentration of the new surface CO species centered at 2050 cm -1 (shown in Fig. 12b), which is strong evidence that this species is active for oxidation to CO 2. Approximately half order dependence on both oxygen and carbon monoxide was found in the high-temperature regime above ignition temperature. It is well known that this reaction has a negative order in CO at low temperature. There were two different reaction regimes for CO oxidation over P t ( l l l ) in the investigated temperature range. In the low-temperature regime (below ignition temperature), atop CO dominated the surface and the reaction has an activation energy of 42 Kcal/mol. The reaction rate was inversely proportional to the surface concentration of atop CO. Above ignition temperature, new CO species (incommensurate overlayer and CO terminally bonded at distorted platinum sites) were dominant, and in excess oxygen CO adsorbed on oxidized Pt. A CO 2 complex was also formed. The reaction rate increased linearly with the surface concentration of the new CO species in the high-temperature regime (>600K), indicating the active role of these CO species in CO oxidation.

300

,,..,, 2 5 0 -i

i i

-'-

"-

"5

200

o E ,,..4

n,. E

"

15o-

"0-...

e~

"'0

50 0.7

0.8

0.9

1.0

Relative Coverage of Atop CO

Figure 12a.

A plot of the reaction rate as a function of atop CO coverage at 590 K.

56 3000

2500 -

2000

1500 0.6

~

J

=

0.7

0.8

0.9

1.0

Relative Coverage (a.u.)

Figure 12b. The reaction rate as a function of the new CO species surface concentration at 720 K.

5. C O N C L U S I O N We have used STM and IR-visible SFG capable of providing chemically specific information at both high and low pressures (i.e., UHV and atmospheres of gases) for the in situ study of heterogeneous catalytic surfaces. The importance of these techniques is that they can provide information on the molecular structure of the substrate and the adsorbed reaction intermediates as well as provide a measurement of their surface coverage under realistic reaction conditions. With a specialized version of a scanning tunneling microscope which allows us to study in situ the atomic structure of surfaces under variable pressures (UHV-atmospheric) and temperatures (30(0-425 K), we have investigated the structures of the initially clean (110) surface of single crystal platinum while in environments of hydrogen, oxygen, and carbon monoxide. The surface in 1.7 atm of hydrogen appears to be dominated by various sizes of nested missing-row reconstruction. The surface in 1 atm of carbon monoxide, however, does not have the small scale missing-row reconstruction, but does appear to have flat terraces separated by multiple height steps. The surface in 1 atm of oxygen appears to have enlarged (111) microfacets. An extreme case of chemisorption-induced restructuring of metal surfaces is corrosive chemisorption as observed by SFG. In this circumstance, metal atoms break away from step or kink surface sites and form bonds with several adsorbate molecules. Carbon monoxide can form several carbonyl ligand bonds with platinum atoms leading to the creation of metalcarbonyl species. Thus, metal-metal bonds are broken in favor of forming metal-carbonyl clusters that are more stable at high CO pressures. The SFG vibrational spectra detect the reversible formation of new adsorbed carbon monoxide species above 100 Torr on Pt(111), that appear to be platinum-carbonyl clusters Ptn(CO) m, with (m/n) > 1 and a CO commensurate overlayer.

57 During ethylene hydrogenation over Pt(111) the reaction intermediate appears to be weakly bound ~-bonded ethylene which produces most of the ethane, while ethylidyne and di-~ bonded ethylene are spectators during the catalytic process. The surface concentration of nbonded ethylene is 4% of a monolayer during the turnover, which yields an absolute turnover rate 25 times higher than the turnover rate per platinum atom. The high pressure STM was also used to study the structure of a Pt(111) model catalyst surface during the thermal decomposition of propylene. Carbonaceous clusters were produced by partial dehydrogenation and polymerization of the hydrocarbon. Finally, SFG revealed new CO species in CO oxidation and provided strong evidence for their active role for oxidation to CO2. All our studies of chemisorption and catalytic reactions at high pressures indicate that the metal surface is flexible; it changes its surface structure as the chemisorbing atoms or molecules have changed. There is mobility in the adsorbed overlayer, and the surface defects also restructure. The dynamics of adsorbate-induced restructuring explains how transition metals such as platinum can be high activity and selectivity catalysts under both oxidizing and reducing reaction conditions. Each adsorbate causes restructuring in different ways so as to optimize the adsorbate/substrate binding energy. Thus, CO or hydrocarbon oxidation occurs on one type of platinum surface structure, while the "reforming" of naphtha which produces high octane gasoline and is carried out in excess hydrogen is carried out on platinum surfaces that have very different surface structures from those present under oxidizing conditions. The experimental evidence from the application of surface science techniques clearly indicates that chemisorption restructures metal surfaces. The restructuring may be subtle, involving the motion of atoms around the chemisorption bond perpendicular to the surface plane and/or rotation parallel to the surface. Often massive adsorbate-induced restructuring occurs, especially at high reactant pressures, that must involve atom transport over large distances (as compared to interatomic spacing) to produce new surface planes and rearrange surface defects (steps an kinks). In the presence of the adsorbate, the metal surface assumes a different and thermodynamically more stable configuration as compared to the surface structure in the absence of the adsorbate. This, then, is the catalytically active surface that is stabilized by the reactant mixture. ACKNOWLEDGEMENT This work was supported by the Director, Office of Energy Research, Office of Basic Energy Sciences, Materials Sciences Division, of the U.S. Depamnent of Energy under Contract No. DE-AC03-76SF00098.

REFERENCES 1. 2. 3. 4.

G.A. Somorjai, Introduction to Surface Chemistry and Catalysis, John Wiley & Sons, Inc., New York, 1994. D.W. Blakely, E. Kozak, B.A. Sexton and G.A. Somorjai, J. Vac. Sci. & Techn., 13 (1976) 1091. A.L. Cabrera, N.D. Spencer, E. Kozak, P.W. Davies and G.A. Somorjai, Rev. Sci. Instrum., 53 (1982) 1888. B.J. McIntyre, M. Salmeron and G.A. Somorjai, J. Vac. Sci. Technol. A, 11 (1993) 1964.

58

5. 6. 7. 8. 9. 10. 11. 12. 13. 14. 15. 16. 17. 18. 19. 20. 21. 22. 23. 24. 25. 26. 27. 28. 29. 30. 31. 32. 33. 34. 35. 36. 37. 38. 39.

P.S. Cremer, B.J. Mclntyre, M. Salmeron, Y.R. Shen and G.A. Somorjai, Catal. Lett., 34 (1995) 11. B.J. Mclntyre, M.B. Salmeron and G.A. Somorjai, Catal. Lett., 14 (1992) 263. B.J. Mclntyre, M.B. Salmeron and G. A. Somorjai, Rev. Sci. Intrum., 64 (1993) 687. J.D. Batteas, J.C. Dunphy, G.A. Somorjai and M. Salmeron, Phys. Rev. Lett., 77 (1996) 534. B.J. Mclntyre, M. Salmeron and G. A. Somorjai, J. Catal., 164 (1996) 184. J.C. Dunphy, B.J. Mclntyre, J. Gomez, D.F. Ogletree, G.A. Somorjai and M. B. Salmeron, J. Chem. Phys., 100 (1994) 6092. Y.R. Shen, Nature, 337 (1989) 519. Y.R. Shen, Surf. Sci., 299/300 (1994) 551. J. Zhang, J. Huang, Y. Shen and C. Chen, J. Opt. Soc. Am., 10 (1993) 1758. X. Zhu, H. Suhr and Y.R. Shen, Phys. Rev. B, 35 3047 (1987). H. Steininger, S. Lehwald and H. Ibach, Surf. Sci., 117 (1982) 342. N.R. Avery, J. Chem. Phys., 74 (1981) 4202. B.E. Hayden and A.M. Bradshaw, Surf. Sci., 125 (1983) 787. R. Superfine, J.Y. Huang and Y.R. Shen, Opt. Lett., 15 (1987) 1276. G.S. Blackman, M.-L. Xu, D.F. Ogletree, M.A. Van Hove and G.A. Somorjai, Phys. Rev. Lett., 61 (1988) 2352. A. Crossley and D.A. King, Surf. Sci., 68 (1977) 528. C. E. Wamaby, A. Stuch, Y. Y. Yeo and D. A. King, J. chem. Phys., 102 (1995) 22. H. Huber, E.P. KiJndig, M. Moxkovits and G.A. Ozin, J. Amer. Chem. Soc., 95 (1973) 332. A. Wander, M.A. Van Hove and G.A. Somorjai, Phys. Rev. Lett., 67 (1991) 626. I. Horiuti and M. Polanyi, Trans. Faraday Soc., 30 (1934) 1164. S. Mohsin, M. Trenary and H. Robota, J. Phys. Chem., 92 (1988) 5229. J.E. Rekoske, R.D. Cortright, S.A. Goddard, S.B. Sharma and J.A. Dumesic, J. Phys. Chem., 96 (1992) 1880. P.S. Cremer and G.A. Somorjai, J. Chem. Soc. Faraday Trans., 91 (1995) 3671. T. Beebe and J. Yates, J. Am. Chem. Soc., 108 (1986) 663. P. S. Cremer, X. Su, Y. R. Shen, and G. A. Somorjai, J. Am. Chem. Soc., 118 (1996) 2942. H. Steiniger, H. Ibach and S. Lehwald, Surf. Sci., 117 (1992) 685. P.S. Cremer, X. Su, Y.R. Shen and G.A. Somorjai, Catal. Lett., 40 (1996) 143. J. Schlatter and M. Boudart, J. Catal., 24 (1972) 482. R. Cortright, S. Goddard, J. Rekoske, and J. Dumesic, J. Catal., 127 (1991) 342. S. M. Davis, F. Zaera, B. E. Gordon, and G. A. Somorjai, J. Catal., 92 (1985) 240. T. Beebe and J. Yates, J. Am. Chem. Soc., 108 (1986) 663. P.S. Cremer, X. Su, Y.R. Shen and G.A. Somorjai, J. Phys. Chem., 100(1996) 16302. P.S. Cremer, X. Su, Y.R. Shen and G.A. Somorjai, J. Chem. Soc. Faraday Trans., 92 (1996) 4717. C. Yoon, M.X. Yang and G.A. Somorjai, J. Catal., (1996), submitted. P.H. Otero-Schipper, W.A.Wachter, J.B. Butt, R.L. Burwell Jr. and J.B. Cohen, J. Catal., 50 (1977) 494.

59

40. R.J. Koestner, J.C. Frost, P.C. Stair, M.A.V. Hove and G.A. Somorjai, Surf. Sci., 116 (1982) 85. 41. R.J. Koestner, M.A.V. Hove and G. A. Somorjai, J. Phys. Chem., 87 (1983) 203. 42. K.M. Ogle, J.R. Creighton, S. Akter, and J.M. White, Surf. Sci., 169 (1986) 246. 43 N.R. Avery and N. Sheppard, Proc. R. Soc. London Ser. A, 405 (1986) 1. 44. G. Shahid and N. Sheppard, Spectrochimica Acta, 46A (1990) 999. 45. T.A. Land, T. Michely, R.J. Behm, J.C. Hemminger and G. Comsa, J. Chem. Phys., 97 (1992) 6774. 46. B.J. Mclntyre, M. Salmeron, and G.A. Somorjai, Catal. Lett., 39 (1996) 5. 47. J.A. Anderson, J. Chem. Soc. Far. Trans., 88 (1992) 1197. 48. C.T. Campbell, G. Ertl, H. Kuipers and J. Segner, J. Chem. Phys., 73 (1980) 5862.

This Page Intentionally Left Blank

91997Elsevier Science B. V. All rights reserved. Dynamics of Surfaces and Reaction Kinetics in Heterogeneous Catalysis G.F. Froment and K.C. Waugh, editors

61

Molecular kinetics of heterogeneous catalytic reactions R.A. van Santen, A. v.d. Runstraat, R.J. Gelten Schuit Institute of Catalysis, Faculty of Chemical Engineering, Eindhoven University of Technology, The Netherlands Abstract To relate knowledge of elementary reaction rate constants with the intrinsic overall kinetics of a heterogeneous catalyst one has to explicitly deal with non-ideal mixing effects of adsorbed reactants. This is illustrated for hydroisomerisation catalysis catalyzed by zeolites and carbonmonoxide oxidation catalyzed by the Pt(100) surface.

1.

Introduction

There is a substantial increase in our understanding of elementary reaction steps on a variety of catalytically relevant materials. Especially surface science has greatly contributed to our knowledge of the reactivity of transition metal surfaces 1'2'3. The other area of spectacular advance is acid catalysis by zeolitic protons 4'5'6. This creates an opportunity to attempt to modeling of overall reaction rates based on elementary reaction steps according to a particular mechanism. Especially in zeolite catalysis theory appears to have been advanced to a stage where

computed

reaction

energy

schemes

enable

discrimination

between

different

mechanisms. Nonetheless a major issue concerns the relative importance of diffusion or adsorption. One is interested to know which aspects of the reaction relate to elementary reaction rate constants and which aspects relate to physical process parameters. This issue will be addressed here by taking as an example the hydroisomerisation reaction of hexane, catalyzed by noble metal acidic zeolites. As a second example we will discuss the microscopic basis of noble metal catalyzed reactions. Essential here is recognition of the importance of surface-composition and diffusion to the overall reaction. Since surface adatoms and adsorbed molecules and their fragments show a strong non-ideality in their mixing behavior, classical kinetics approaches do often not apply. This holds especially when strong lateral interaction lead to surface island formation. As an

62 example the oscillatory kinetics of the CO oxidation reaction will be taken. Monte Carlo studies can be used to take care of the statistical mechanical aspects of the overall kinetics problem.

2.

Hydroisomerisation kinetics

Interesting kinetic information of the importance of adsorption effects to overall kinetics of a zeolitic reaction is obtained by studying the reaction at conditions where pore filling varies 8. Experimental data indicate that especially the orders of alkane and hydrogen are sensitively dependent on this. Whereas at high pressures pores are completely filled and hence in a bifunctional mechanism alkane and alkene are equilibrated within the micropore 9, this is not necessarily the case at low pressures. A microkinetics model has been constructed that explicitly describes the dependence of overall reaction rate on micropore flling. The model can be illustrated with figure 1. Molecules from the gas phase adsorb to the zeolitic micropore sites. Transport steps are introduced between molecules adsorbed in the micropore with the metal sites and with the acidic protons, also located in the micropores. Communication of molecules between catalytically active sites again is only possible via the micropore sites. Clearly molecular exchange between gas phase and active sites becomes inhibited with increasing micropore filling. As an overall result 8 one finds that the orders in hexane and hydrogen differ only in absolute value at low micropore filling to respectively 0 and -1 at high micropore filling. Differences between zeolites appear not to be controlled by differences in intrinsic acidity 1~ but by differences in heats of adsorption 11'~2. At intermediate pressures the order in hexane is low and hence differences in apparent activation energy, that arise from adsorption effects, are minor g. Diffusion effects play a role at high pore filling. This can be considered a consequence of strong adsorption energies. When properly optimized, the metal function of the catalysts essentially maintains hexanehexene equilibrium. The reaction energy scheme deduced 1~ for hexene protonation and isomerisation is shown in figure 2. Protonation of hexene consists of two energy terms: adsorption of hexene in the micropore and subsequent protonation. Only the first term depends significantly on micropore dimensions. The overall effect of course is a difference in protonation energy of gas phase olefine for different zeolites. With respect to the protonated state the true activation energy for

63 isomerisation is found to be of the order of ~ 130 kJ/mole. This is substantially higher than the corresponding value o f - 3 0 kJ/mole found in superacidic systems. The concept of micropore blocking at high filling suggests an answer to a classical question in hydroisomerisation catalysis. According to Spivey and Bryant 13, Pt-H-Mordenite and Pd-H-faujasite behave dramatically different for the hydroisomerisation of mixtures of pentane and hexane. Whereas in mixtures of pentane and hexane the rate of hexane isomerisation increases and the rate of pentane is relatively high when catalyzed by Pt-H-Mordenite, with Pd-H-Faujasite the rate of hexane isomerisation is found to decrease and hexane is found to suppress isomerisation of pentane. Comparison of the rate changes of pentane conversion, shows that dilution with hexane decreases the rate of pentane conversion in both cases, but most in the catalytic reaction with Faujasite. The difference in heats of adsorption between Faujasite and Mordenite causes the micropores of Mordenite to be nearly completely occupied, where in Faujasite this is much less the case. The Mordenite behavior may correspond to the case where micropore alkane/alkene equilibrium is not equilibrated with gas phase, whereas the Faujasite behaves as if equilibrium between gas phase and micropore us established. E.g. one deduces for high pore filling and no equilibrium with gas phase.

r6 =k~-06 =k~.Keq- 06. PH2 (1) ~k~-

(

6

h) { 1+

6

}

--u 1+ Kads 9Ptot

---;--6--

The corresponding expression for the low pore filling and equilibrium between gas phase and micropore becomes:

64 6 Kads "Keq r6 = k ~ . 0 6 = k ~ .

6 1 + Kad s .Keq

P6

P6 5 - - + Kads 9Keq PH 2

PS PH2 (2)

6 "Keq "(Ptht-P5) Kads

(P.2 + K ds Ken Ptht)

(K5ds - K6ds) P5 PH 2 Keq

+ K6ds "Ptot

k6i is the elementary constant for isomerisation of hexene in the adsorbed state. 0 is the surface coverage. K ~ is the equilibrium constant for adsorption of hexane in the micropore: K ~ > K]~ because of the higher heat of adsorption of hexane 12 Keq is the equilibrium constant of the alkene/alkane equilibrium. This is assumed to be the same for C5 and C6. Expressions (1) and (2) both show a decreasing rate of the rate of hexane isomerization when hexane is deluted with pentane. However, when there is no equilibration with the gas phase, the decrease is significantly less than when micropore composition and gas phase composition are equilibrated (PHz / Keq >> 1). These simple model calculations illustrate the dramatic effect of micropore filling on catalytic performance. Since dilution by pentane does not increase the rate ofhexane in Mordenite. The experiments by Spivey and Bryant indicate that diffusion has to be explicitly considered. It is likely that not only the higher heat of adsorption of alkanes in Mordenite compared to Faujasite, but ultimately the one dimensional structure and resulting single file diffusion TM are responsible for the unique behavior of the Mordenite catalyst. Also the much smaller reduction of pentane conversion when mixed with hexane in Mordenite compared to Faujasite is a strong indication of the non-equilibrium situation in the medium pore zeolite.

3.

The oxidation of CO catalyzed by Pt(lO0) surface

Experimentally, on well defined single crystal surfaces and under particular conditions, the rate of the CO oxidation reaction has been found to develop complex oscillatory behaviour 15.

65 Additionally, reactant concentrations can form surface pattems in a variety of shapes such as concentric circles or rotating spirals 16. Simulation of such spatio-temporal behavior based on atomistic elementary reaction rate parameter provides a sensitive test for macroscopic rate simulations. Here we report dynamic Monte Carlo simulations 17'~s of spatio-temporal patterns and kinetic oscillations. Surface reconstructions are responsible for oscillatory behavior of a wide range of catalytic reactions ~9. Also in the CO oxidation under consideration here, the feedback mechanism responsible for oscillatory behavior is a reconstruction of the Pt(100) surface, that is driven by the surface concentration of adsorbate molecules 2~

The most stable structure of the clean

Pt(100) surface is the hexagonal phase. Upon adsorption of CO, the surface transforms into the square (lxl) form. Adsorption rates of CO on the two different phases are hardly different. In contrast, oxygen adsorbs a thousand times slower on the hexagonal phase than on the square phase, while its adsorption probability on the square phase is of the same order of magnitude as that of CO. Crucial to the simulations presented here is the inclusion of surface reconstruction, together with correct time-dependence of the reactions. As such, the method provides an extension of earlier important computer simulations of CO oxidation on Pt surfaces 23'24'25. A dynamic Monte Carlo method is used 26, based on the solution of the master equation of the reaction system. The reaction system consists of a regular grid with periodic boundary conditions. The largest grid used in our simulations contained ca. eight million reaction sites. A short description of the model is presented in Fig. 3 and in Table I, that shows the parameters of the rate constants considered. As shown in figure 4 at low temperature, below 400 K, the reaction rate is zero because of surface poisoning by CO or oxygen, depending on the partial gas pressures and the sticking coefficients chosen in this temperature regime. In the range between 400 K and 440 K, molecules start to desorb and reaction is initiated. A steady-state reaction rate of CO2 production is found, which increases with temperature. Increasing the temperature from 440 K, our simulations show a decreasing reaction rate, going to zero at 530 K. The rate of reaction then becomes limited in the adsorbate surface concentration. Above 530 K, reactant concentrations on the surface are so low, that no CO2 production occurs. Between 470 and 520 K, oscillations in the CO2 production rate are found. The amplitudes of these oscillations are strongly temperature dependent. The most regular oscillations and the largest amplitudes

Table 1

reaction

psi,,, (Pa)

pexp(Pa)

S~im

exp so

CO ads.

1.3.10 .3

1-4" 10 .3

0.8

,~0.8

[12,13,36]

02 ads.

5_15.10 .2

2.7.10 .2

0.1

~0.1

[13,14]

V sim (S -l )

vexp (s -! )

E sim act (kJ / mol)

CO des.

1.1015

1-3.105

175

117-159

[13,32-36]

CO2 prod.

2.10 l~

10 I~

84

50-100

[37]

l x l --+ hex

1.10 9

105

105

[lo,35]

Nucleation

0.03

0

~0

[33-36]

Trapping

0.03

0

~0

[35,36]

ref

ref

exp act (kJ / mol)

67 hexanes(g), hexenes(g)

/

adsorption

\

j n-a oxyl\

Figure 1: Schematic model of hydroisomerisation reaction scheme including transport steps.

Table 1: Comparison between experimental parameters and simulation parameters, p stands fpr pressure, v for prefactors, Eac, for activation energy and So for the initial sticking coefficient. In our model, CO adsorption occurs on both the hexagonal and the 1 x 1 phase, with equal rates. Parameters for O2 ads. are for dissociative oxygen adsorption on the 1 x 1 surface only. Adsorption of oxygen on the hexagonal phase is neglected, since it is a thousand times slower than on the 1 x 1 phase. Oxygen desorption does not occur, because of the strong bonding of the fragments to the surface. The surface reconstructions are described in the text.

Pt-H-Mordenite

Plot

C~/CB

H2/Hc

rc5

0.79

pure C5

10/1

0,241

0.79

pure C6

10/1

0.79

15/85

8/1

50/50 75/25

(MPa)

Pd-H-Faujasite

514K

Ptot

CJC6

H2/Hc

rc5

0.79

pure C5

10/1

0.187

0.337

0.79

pure C6

10/1

0.101

0.374

0.79

15/85

10/1

0.042

0.293

10/1

0.137

0.519

50/50

10/1

0.095

0.208

10/1

0.143

0.507

75/25

10/1

0.081

0.183

FC6

rc~

0.431

Mixing decreases rate for C5 conversion less with Pt/H/Mordenite than with Pd/H-Faujasite

Mixing increases rate for C 6 conversion with Pt/H/Mordenite, but it decreases with Pd/H/Faujasite

J. Spivey, P.A. Bryant, Ind. Eng. Chem. Proc. Dev. 21,750 (1982)

69

Mixtures High pore filling" No equilibrium between gas phase and micropore (Mordenite) ga6s[C6]

Keq

0 hex

1+ Ka6ds[C6] q.- Kads[Cs]

K eq PH~(1 + Ka6d~ 9P~oht)

i

f

:s(eo

_

- Kads)as 1 l+ (KAY, - -- ' 7~6-- - -- (i~ 1 + K~d , "P,o,

Low pore filling" Equilibrium micropore (Faujasite)

between

Ka6s Keq

gas

phase

p~

9

r6 - k~ . o L

- k~.

1 + Ka6s . K eq

P6 + K~, 9K eq

Ka6~ "Keq "(Ptoht-- Ps)

(~ + Kads "Keq 9~oht) 9 1+ 6

~eq +

K6ds" etht

and

70

I

9~ 8 6+ --d--~--

/

R-C=C\

SjO,,,~f O

-A~,~=c6~8oI R_cl =c'"%""iit~/5~ ," Si~i

'.'.*

I

Z~Ipr~ 80 adsI"!sCjt~\A(O R-CH- / I

gactis~ - 130

C 120

C

I

R'-C=C

,,' c '~t;R'-C=C ~o, /I -80 I R'-C-C I['Sl S i ~ 83 is iso-alkoxy species

n-alkoxy species

Figure 2" Reaction energy scheme of alkene protonation and isomerisation.

71 are found at 490 K. Here the simulates oscillation period is around 200 seconds. Excellent agreement with experiment is found between 460 and 540 K. On large grids, many fronts are generated in parallel, all forming local oscillations in CO2 production. If these local oscillations are not in phase, they result in a vanishing amplitude of the oscillations in the overall CO2 production rate. synchronization occurs when crossing reaction fronts extinguish one another periodically. On large grids, another form of self-organization, yielding-temporal pattern formation becomes possible. Wave fronts in the form of pulsing concentrical circles as well as rotating spiral structures can occur. Such patterns have been observed experimentally on Pt(110). The experimental observation of concentrical wave fronts has been ascribed to structural defects in the surface 2v. On these defects, oxygen dissociates faster, and they therefore act as a periodic pulse generator for reaction fronts. With each pulse, a new front is generated, which then grows continuously. This behavior is indeed reproduced by our simulations when such a defect is included. Nonetheless, concentrical circles can also be simulated on a perfectly homogeneous surface. A circular reaction wave front is initiates spontaneously, e.g. by oxygen adsorption on a site that is vacated by CO desorption, at a position we will call the primary center. Inside this front, at a position different from the primary center, a new front is initiated. Structural defects form a stabilizing, rather than a necessary factor for spatiotemporal pattern formation in the form of concentrical circles. Monte Carlo simulations show that, during a simulation, due to stochastic fluctuations, synchronized kinetic oscillations spontaneously change into spatio-temporal patterns with steady state kinetic behavior, and vice versa. Also, rotating spirals can change into pulse generators and vice versa. In

summary,

surface

reconstructions

provide

a

microscopic

mechanism

for

the

synchronization of local oscillators. Global synchronization is necessary to obtain oscillations in macroscopic dimensions.

72

2 - D grid of unit cells, two sites per cell.

y

1 x 1 phase" use only site 1 unit cell"

hexagonal phase" use both sites unit cell"

neighbor relation with 4-fold symmetry"

neighbor relation with 3-fold symmetry"

,, i

~ %% %

9 '~%

% %%

%

%

%

%1

% %%

|

%%

%%

%%

'~%

% %%

%%

%% %

'~

% %

% %%

.7

73

Figure 3: The simulation grid consists of unit cells containing tow sites, in order to allow differentiation between the 1 x 1 and the hexagonal phase, which have different symmetries. For graphical representations, we can choose unit vectors arbitrarily; for the simulation program only relative coordinates (x, y, site) ate specified. When site 2 of a unit cell is marked with a label, this site becomes unavailable for adsorbates, hence we have effectively one site per unit cell. This results in a fourfold neighbor relation as on the 1 x 1 phase. In the absence of a label on site 2, this site is indeed available for adsorbates, hence we have two equivalent sites per unit cell. This results in a threefold neighbor relation as on the hollow sites of the hexagonal phase. Reactions

are specified as changes of the contents of one or more

sites.

Surface

reconstructions are specified as reactions that turn the label on site 2 off or on. In a similar way, the top-sites on both lattices can be incorporated, without altering our results significantly. The algorithm of the simulations p r o g a m works schematically as follows. (1) It reads the initial configuration and the reactions that are specified. (2) It then calculates a time at which the first reaction will occur, using a probabilistic procedure based on a stochastic model of the reaction kinetics. (3) One of the reactions that is possible on the grid is selected with a probability proportional to its rate and executed. (4) Goto (2).

0.12 0.1 0.08 (1)

I~"

!

!

average rate 4imits o

(>

,

o~

9

9 ', ", , ~"',

,,,,

"-..~

...oe~176

"

'~

0.06 0.04 0.02 0

400 420 440 460 480 500 520 540

75

Figure 4: CO2 production rate versus temperature, simulated on a grid of 256 x 256 unit cells. Similar data obtained on larger grids do not differ significantly from the ones shown here. The reaction rate is given in molecules CO, produced per platinum atom. The average rates are drawn as well as the upper and lower limits of the rate in the oscillatory region. The amplitudes of the oscillations are the difference between the upper and lower limits.

76

References: 1.

J.K. Norskov, Rep. Prop. Phys., 53 (1990) 1253; B. Hammer, J.K. Norskov, in 'Chemisorption and reactivity on supported clusters and thin films: NATO-ASI Series 331, R.M. Lambert, G. Pacchiani (Ed.), Kluwer, 1996, to appear.

2.

J. Whitten, H. Yang, Surf. Sci. Rep., 44 (1993).

3.

R.A. van Santen, M. Neurock, Catal. Rev. Sci. Eng., 37 (1995) 557.

4.

J. Sauer, Chem. Rev., 89 (1989) 199.

5.

C.R.A. Catlow, J.M. Thomas, Phil. Trans. Roy. Soc. London A, 341 (1992) 255.

6.

R.A. van Santen, G.J. Kramer, Chem. Rev., 95 (1995) 637.

7.

S.R. Blaszkowski, R.A. van Santen, J. Am. Chem. Soc., 118 (1996) 5152. S.R. Blaszkowski, R.A. van Santen, Topics in Catal., in press.

8.

A. van de Runstraat, J. van Grondelle, R.A. van Santen, Ind. Eng. Chem. Rev., to appear.

9.

G.F.~Froment, Catal. Today, 1 (1987)455.

10. J. J~inchen, J.H.M.C. van Wolput, L.J.M. van de Ven, J.W. de Haan, R.A. van Santen, Catal. Lett., 39 (1996) 147. 11. A. van de Runstraat, P.J. Stobbelaar, J. van Grondelle, L.J. van IJzendoorn, R.A. van Santen, Stud. Surf. Sci. Catal., 105 (1997) 1253. 12. S.P. Bates, W.J.M. van Well, R.A. van Santen, B. Smit, J. Am. Chem. Sot., 118 (1996) 6753. 13. J. Spivey, P.A. Bryant, Ind. Eng. Chem. Proc. Dev., 21 (1982) 750. 14. H. Lin, G.D. Lei, W.M.H. Sachtler, Appl. Catal. A, 137 (1996) 167. 15. R. Imbihl, G. Ertl, Chem. Rev., 95 (1995) 697. 16. S. Jakubith, H.H. Rotermund, W. Engel, A. von Oertzen, G. Ertl, Phys. Rev. Lett., 65 (1990) 3013. 17. R.J. Gelten, J.P.L. Seegers, J.J. Lukien, A.P.J. Jansen, P. Hilbers, R.A. van Santen, manuscript in preparation. 18. R.J. Gelten, J.P.L. Seegers, J.J. Lukien, A.P.J. Jansen, P. Hilbers, R.A. van Santen, manuscript in preparation 19. M. Kiskinova, Chem. Rev., 96 (1996) 1431. 20. S. Hagstr6m, H.B. Lynn, G.A. Somorjai, Phys. Rev. Lett., 15 (1965) 491.

77 21. P. Heilman, R. Heinz, R. Mtiller, Surf. Sci., 83 (1979) 487. 22. A. Hopkinson, D.A. King, Chem. Phys., 177 (1993) 433. 23. R.M. Ziff, E. Gulari, Y. Barshad, Phys. Rev. Lett., 24 (1986) 2553. 24. P. M611er, K. Wertzl, M. Eiswirth, G. Ertl, J. Chem. Phys., 85 (1986) 5328. 25. M. Eiswirth, P. M611er, K. Wetzl, R. Imbihl, G. Ertl, J. Chem. Phys., 90 (1989) 510. 26. A.P.J. Jansen, Comp. Phys. Comm., 86 (1995) 1.

This Page Intentionally Left Blank

91997 Elsevier Science B. V. All rights reserved. Dynamics of Surfaces and Reaction Kinetics in Heterogeneous Catalysis G.F. Froment and K.C. Waugh, editors

79

From supersonic beams and single crystal microcalorimetry to the control of catalytic reactions D. A. King Department of Chemistry, University of Cambridge, Lensfield Road, Cambridge CB2 1EW, UK

Abstract Modem surface science techniques provide a means for a detailed unravelling of the mechanisms of catalytic reactions and the surface processes which accompany them. The focus of the present review is on one surface, Pt{ 100}, and several important catalytic reactions: CO oxidation, a critical reaction in car exhaust catalysis, NO reduction by H2, and NH3 oxidation, the industrial process used to manufacture NO and hence nitric acid. The { 100} surface of Pt is of particular interest, because it can be prepared in two forms, a metastable (lxl) bulk termination structure and the stable hexagonal top layer structure, denoted hex. Using a novel single crystal adsorption calorimeter, we have determined the energy differences between the clean (lxl) and hex surfaces, and also between the two surfaces with adsorbates which lift the hex reconstruction. Molecular beam studies with CO and D2 revealed the mechanism for this adsorbate restructuring process: it proved to be strongly non-linear. This non-linearity was subsequently shown to be critical in the widespread observation of regimes of sustained oscillations in many catalytic processes, such as CO + 02, CO + NO, and NO + H2, observed on Pt{ 100}. Detailed modelling using only experimentally determined kinetic parameters gives remarkably good agreement with experimental measurements of oscillatory existence regimes and periods. Finally, a combination of calorimetric and molecular beam techniques have produced a new mechanism for the industrially important ammonia oxidation reactions (the Ostwald process) over Pt catalysts. This includes a tested recipe for a very substantial improvement in the operating conditions, with high selectivity to the desired product (NO) at low temperatures, and at high rates.

1. Introduction The greatest potential for i m p r o v e m e n t s in catalytic processes in the long term lies in a reductionist approach, which has been outlined elsewhere 1. For a given reactive system of interest, the techniques developed by surface scientists, both experimental and theoretical, can be deployed to tackle each of the following questions.

80 (1)

Where are the atoms comprising the surface at various stages of the process? Why are they there? What structural changes accompany surface processes? What are the bond energies, and interaction energies between adsorbed species, and what is the nature of the surface chemical bonds? What are the factors controlling the bond making and breaking processes at the surface? How can these processes be controlled to yield the desired products and improved rates?

(2) (3) (4)

The focus of the present brief review will be on one surface, Pt{100}, and several important catalytic reactions, including CO oxidation and NO reduction, of importance in car exhaust catalysis, and ammonia oxidation, the Ostwald reaction, which is still an important industrial process for the production of NO. The importance of these processes can be simply demonstrated: of the world's consumption of 4.71 million ounces of Pt in 1996, 1.87 million ounces was used in the production of car exhaust catalysts alone. The {100} surface of Pt has itself been subjected to fairly exhaustive structural analyses 2-4 ; as shown in figure 1, the top layer adopts a pseudo-hexagonal structure, while the second layer retains the square symmetry of the ideal {100} surface, with a corrugation across the large two dimensional unit cell of about 0.5,~ arising from the misregistry of the two layers. The structure can be understood within the framework of effective m e d i u m t h e o r y s : the co-ordination around surface atoms is reduced by the loss of neighbouring atoms compared with bulk atoms, and hence the valence charge

2o - ~

eooe

e.oe

9c : ~ : o : o : o : o o o . o o ~ o

oomoo 9 ~o:o 9 9

9 9.....-

-..,o,., - , o o o o o oo o oo.o o ~ o oos o o Oo5ToOoOo+ ,~ o D~. .~ ~O,q),O,O,O:O,O~O~O~O.O.O,OOO O ".

O_-QO~-~.

9 9 9 9o_o 9 9 9 9 9o~a_:a_:e~o :o :o ~o

O'O.

Figure 1. A model of the top two layers of the stable hex phase of Pt{ 100}, with a pseudo hexagonal top layer and bulk-like square second layer.

~ ~,O~-O~O--O--.O0-,O000,

density is lowered below the optimal value. The surface layer atoms respond in two ways: by relaxing towards the second layer, and by reconstructing so as to increase the density of atoms in the top layer. Both effects combine to increase the local embedding charge density to the optimal value characteristic of bulk atoms. When an adsorbate, such as CO, NO or H, is introduced to the hexagonally reconstructed surface, referred to henceforth as Pt{100}-hex, the reconstruction is lifted and the top layer Pt atoms adopt the ideal (lxl) structure 4 . The formation of a chemical bond between top layer Pt atoms and the adsorbate increases the local e m b e d d i n g charge density experienced by the surface Pt atoms, the hexagonal top layer now experiences too large a charge density, and the (lxl)

81 surface becomes more stable. A metastable clean (lxl) surface can be generated stable up to - 380 K, by adsorbing NO, desorbing N2 and titrating off the remaining O with hydrogen. 2. Energetics: Calorimetric Measurements on Pt{lO0}

An instrument has recently been developed to measure adsorption heats and reactions at single crystal surfaces calorimetrically6, 7 . It is based on the use of thin film single crystals mounted on a metal ring, and pulsed molecular beam dosing with remote temperature sensing of the rise in crystal temperature. This is achieved by detecting the increase in black body radiation of the unsupported central 2 mm diameter section of the single crystal. This section has a heat capacity of only - 1 ~tJ K -1, and the technique is sensitive to gas doses of 1% of a monolayer of, for example, CO on Ni{100}, with good signal-to-noise. The amount adsorbed per gas pulse is determined by measuring the flux using a stagnation detector and the sticking probability using the King and Wells 8 reflection-detector technique. The technique has been deployed to measure the energetics for irreversible and reversible processes at surfaces, providing, inter alia, a critical benchmark for theory and a surface bond dissociation energy database which may be useful in the prediction of catalytic activity 9 . CO/Pt{100} system:

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

0.90

i

0.77

9

Figure 2. Schematic energy diagram illustrating the clean hex (lxl) phase transition, and the relative stabilities of the two phases with adsorbed CO. All energies are quoted in units of eV per (lxl) unit cell areal 0.

1

P~{l~}(lxl) - 1r CO(~

The only direct measurement of the energy difference between two surface phases has been made by single crystal adsorption calorimetry (SCAC), on Pt{100}. This was achieved by comparing the integral heats of adsorption of CO and of C2H4 on the two different initial states of the surface, the metastable (lxl) and the stable hex, at a total coverage of 0.5 monolayers. At this coverage the hex reconstruction is completely lifted, and, whatever the initial state, the final state is 0.5 ML CO (or C2H4) on Pt{100} (lxl), with a c(2x2)-CO structure. In this way the energy difference between the two clean surface initial states, U l x l - Uhex, was found to be 13 kJ (mol Pts) -1. As shown schematically in figure 2, adsorption of

82 CO on the hex phase without lifting the reconstruction occurs with a heat of adsorption (120 kJ (mol CO) ~ which is 60 kJ (mol CO) -1 lower than that on the metastable (lxl) surface. This much higher adsorption heat on the (lxl) phase is the factor that reverses the stability, so that with 0.5 ML CO the (lxl) phase is 17 kJ (mol Pts) -1 more stable than the hex phase. 3. Energetics: the Switch from Dissociative to N o n - D i s s o c i a t i v e N O Adsorption

Calorimetric data for dissociative a d s o r p t i o n have highlighted the importance of lateral interactions between adatoms, particularly for catalysis. For example, the differential adsorption heat for 02 dissociative adsorption on Ni{100} falls from 570 kJ mo1-1 at zero coverage to 170 kJ mo1-1 at 0.5 MLll, 9 , which indicates a next-nearest-neighbour O-O pairwise repulsive interaction energy as high as 40 kJ mo1-1. These strong a d a t o m - a d a t o m repulsive

350~+0,

400 _

Figure 3. Differential heat of adsorption for NO on Ni{ 100 } at

o 300 E

300 K 12.

200 "T" 150 100 ~-

50 ~ 0

k ,

!

,

9

!

-

'

t

0.2 0.4 0.6 Apparent Coverage, ML

'

0.8

interactions play a critical role in NO chemisorption, as clearly demonstrated on Ni{100} 12,13 . As shown in figure 3, the N-O, O-O and O-N repulsive interactions drive the adsorption heat down from 380 kJ mo1-1 at zero coverage to 135 kJ mo1-1 at 0.6 ML; further adsorption then occurs into the non-dissociatively chemisorbed NO state, for which the adsorption heat is higher than it would be for the dissociated state. Due to repulsive adatom-adatom interactions, at 0 = 0.2 ML NO, the differential heat of adsorption into the dissociated state w o u l d be - 60 kJ mol 1, significantly less than that for the molecular state at 135 kJ mo1-1. In Section 6 below we shall see that this has critically important implications for selectivity in NOx removal and in NO synthetic catalysis. We note that on Pt at room temperature NO adsorption is non-dissociative, with an adsorption heat of 140 kJ mo1-1, even at zero coverage, but at higher temperatures the same considerations apply as on Ni{100}. 4. M o l e c u l a r Beam Studies: the M e c h a n i s m of Adsorbate-Induced Reconstruction

From earlier studies14, is it is known that the lifting of the hex phase reconstruction of Pt{100} by CO adsorption proceeded through the nucleation and g r o w t h of islands of ( l x l ) phase in which the local CO coverage was

83 approximately 0.5 ML, while the CO coverage on the hex phase was very low. It was a s s u m e d that a CO molecule landing on a hex region of the surface could adsorb and then diffuse across the surface; on arrival at a (lxl) island b o u n d a r y , it w o u l d be transferred to the (lxl) phase, producing island growth. H o w e v e r , this m e c h a n i s m for island growth was not borne out in m o r e recent d e t a i l e d molecular b e a m studies16,17 . At temperatures between 285 and 340 K, the CO sticking probability s on the initial hex phase surface w a s f o u n d to fall monotonically from 0.8 at zero coverage to - 0.02 at 0.5 ML CO, where a c(2x2) CO structure covers the whole Pt(lxl) surface. In this t e m p e r a t u r e r a n g e s is 0.8

~, .= 0.7 .o r 0.6

i . . . .

..O o

0.5

r

0.4

-~

o

!

. . . .

i . . . .

!

"'"'--..

Z

!

5.4x10

0.3

.3

K

[

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""-.

"...

..... ...

.......

o.~

0.0

. . . .

""'"....

0.2

"~

. . . .

.... ...,...,,(,5.0x10 3 ML/s at TS= 315 K "'... " 1"9x10"2 ML/s i T - 400 ""',.. 1.0x10 -2 ML/s I S -

0.0

0.1

0.2

0.3

Total CO coverage

0.4 (ML)

0.5

Figure 4. Sticking probability measurements for CO on Pt{ 100}-hex at a temperature of 400 K, at variable incident molecular beam fluxes from 8.5x10 -4 ML s -1 to 1.9x10 -2 ML s-1. The dashed line is the measured sticking probability at temperatures below 340 K 17.

i n d e p e n d e n t of both t e m p e r a t u r e a n d b e a m flux. H o w e v e r , at h i g h e r t e m p e r a t u r e s a curious phenomenon was discovered" the sticking probability is very strongly flux (or pressure) dependent. Results obtained at 400 K are shown in figure 4. At a coverage of just 0.04 ML, for example, the sticking probability varies between 0.03 at a flux of 3.5 x 10.4 MLs -1 to 0.66 at 1.9 x 10 -2 ML/~ (This is equivalent to a pressure increase from about 4x10 -10 mbar to 2x10 -8 mbar). These results have a relatively straightforward explanation. At temperatures above 380 K the lifetime of CO on the hex phase is < 1 s, whereas the ( l x l ) phase, at t e m p e r a t u r e s below about 430 K, the lifetime is effectively infinite. This is a consequence of the differences between adsorption heats on the two phases, as discussed in Section 2 above. In this temperature range there is competition for CO m o l e c u l e s on hex either to transfer to (lxl) islands or to desorb at a rate rd, given by Z (Sa - Sn), where Sa is the absolute sticking probability, Sn the measured sticking probability and Z the collision rate. As shown in figure 3, Sa is given by the sticking probability curve at Ts < 340 K. This means that the rate law governing island g r o w t h can be determined, since the CO coverage on the hex phase is governed by a first order desorption expression at very low CO coverages:

ohex

CO = r d / k d

where kd is a measured rate constant 17, and the island growth rate is simply

rg = Z s

n.

84 r is strongly non-linear16,17; the rate expression governing the A plot of rg vs. "CO adsorbate-induced reconstruction is:

_/,-,hex)4.1

rg = c~tTco

(1) Figure 5. Schematic illustration of the CO-induced Pt{ 100 }-hex to (lxl) surface phase transition. Light spheres: Pt atoms in (lxl) phase; darkened spheres: Pt atoms in hex phase; hatched spheres: region of hex surface to be switched to (lxl) by density fluctuation leading to 4 CO molecules (black dots) on the hatched patch.

We have s u g g e s t e d that this rate law is indicative of a co-operative phenomenon. In order for an area of hex phase to be transformed to (lxl), a m i n i m u m number of surface Pt atoms are involved; as suggested in the sketch in figure 5, this may be 7. For one CO molecule the adsorption heat advantage on the (lxl) surface is 60 kJ (mol CO) ~ which spread amongst 7 surface Pt atoms reduces to -- 9 kJ (mol Pts) -1, insufficient to overcome the inherent energy advantage of the clean hex phase (13 kJ (mol Pts)-l). Apparently around 4 to 5 CO molecules need to be simultaneously present on a minimum hex area adjacent to a (lxl) island to cause growth; this is produced by a local density fluctuation in the CO adlayer on the hex phase. Perhaps surprisingly, the same rate law, equation (1), was also found for the lifting of the hex reconstruction on Pt{100} by hydrogen (deuterium) although in this case the f l u x - d e p e n d e n t sticking probability occurs at a m u c h lower temperature, of a r o u n d 240 K 18 . A similar process occurs with 02 adsorption, except that in this case the adsorbate on the hex phase is immobile; a Monte Carlo analysis of sticking probability data showed that the restructuring is only initiated w h e n the local island size reaches 4 0 adatoms 19 .

5. The Role of Non-Linear Restructuring in Oscillatory Reactions 2~ The catalytic activity for CO oxidation on the unreconstructed (lxl) phase is considerably higher than on the hex phase, which has been traced back to the very low sticking probability for dissociative 02 adsorption on the hex phase (~ 10 -4) compared with the (lxl) phase (> 0.2) 21,19,22 . Imbihl and Ert123,24 w e r e the first to associate this phase transition directly with the existence, u n d e r specified

85 conditions, of an oscillatory regime in the CO + 02 reaction over Pt{100}, although they did not incorporate the non-linear growth law, equation (1), into the mechanism. In our own recent modelling, we have d e m o n s t r a t e d that a common mechanism based on the non-linear p o w e r law together with an autocatalytic sequence applies to oscillatory regimes observed in the CO + NO 2s , C O + 0226,27 and NO + H228,29 reactions. For the CO + 02 reaction, the critical importance of the power factor n in the growth law (where n = 4.1 in equation (1)) was demonstrated in a set of computations, based on a mean-field approximation, in which n was varied, but all other parameters such as rate constants were taken from i n d e p e n d e n t experiments 27. With the experimental value, n = 4.1, the existence regime and the period of the oscillations experimentally observed were reproduced well, but as the reaction order n was reduced the period of oscillations was found to decrease, and completely vanished for n < 3. Experimental results for the oscillatory behaviour of the N O / H 2 reaction on Pt{100} provide a particularly severe test for kinetic modelling. In this case the oscillatory period was found to be a very sensitive function of substrate temperature with given partial pressure of reactants, varying within a range of less than 10 K from about 400 s to only a few seconds with increasing t e m p e r a t u r e 3~ . The essential features of this are well reproduced by the nonlinear model 28, which also provides an explanation for this remarkable temperature sensitivity. Decreasing the temperature in the region of 450 to 430 K causes a significant decrease in the rate of h y d r o g e n desorption, and hence increase in the hydrogen coverage. Surface hydrogen reacts with dissociated NO, and therefore decreases the coverage of molecular NO on the (lxl) areas of the surface. The critical NO coverage required to switch from the upper to the lower rate branch is then approached more slowly, which leads to an increase in the oscillatory period. Finally, below a certain temperature the critical NO coverage is no longer reached, and the reaction stays on the upper rate branch, with oscillations terminated. In this work we did not adjust any parameters, but used values quoted by others from experimental data. Subsequently we have improved the fit to the data, by re-examining the experimental parameters and also by introducing non-zero rate terms for the production of N2 O29 . In this way one further detail from the experimental results 31 is accurately reproduced in the model: as shown in figure 6, the oscillatory production of N20 is found to be out of phase with the production of NH3, H2 and H20. This ability to describe the major features of the experimentally observed behaviour, using parameters for rate constants based on independent experiments, raises the confidence level in the non-linear model for this system.

86 1 o.o

_-1_

50

0~; / M L

. . .. . ,

3

o

z

. i a

L

. .. . .

I

,

,

,

I

I

'-

a

I

. . . .

~ "~I : : : :, , : : : ,.1

0.25

'

' - '

i

1. . . .

I

. . . .

!

I

. . . .

I

. . . .

I

7

: -,, ,, -, 1- ,,- -: i* ,, , :~ : ~: : I~-

0,}~.o / M L 0.00 50

Oo~ / M L ' o

x 0.25

0~ '~ / M L 0.00 100

Figure 6. Modelled oscillations for the NO + H2 reaction on Pt{ 100} at T = 434 K, PNO = 1.1x106 mbar, PH2 = 2x10 -5 mbar. Local adsorbate coverages, reaction rates and the fraction of the surface in the (lxl) phase are shown 29.

r,. / MLs-~ ~• o

0.25

rH. o / M L s -~ 0.00 30

I

I

I

I

0

50

I

I

i

,

A

,

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,

,

,

,

_

I

,

,

,

,

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,

,

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,

_

r,.., / MLs-' x 5.0

r.,.o

A

,

,

,

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,

,

,

~

/ MLs-' 2.5

1 O0

150 time

200

250

S

6. The Ostwald Process: Catalytic Oxidation of NH3, Over Pt{100} A m m o n i a is oxidised over platinum to form NO as the first step in the industrial synthesis of nitric acid, a process that has been in large-scale use for over 70 years. The reaction is strongly exothermic, enabling reactors to be run adiabatically, and it results in high conversions of NH3 to NO (94-98%) 32 . Despite this efficiency, problems are encountered with Pt loss, which is believed to be due to the formation of PtO2(g) at the high o p e r a t i o n t e m p e r a t u r e s (> 1000 K) 33 . These high temperatures are necessary both because the reaction exhibits a temperature dependent product selectivity (below ~ 800 K the oxidation reaction mainly produces N2) and also for the activation of new Pt gauzes. A new gauze is inactive until it is at temperature in the presence of the reactants; the wires restructure causing their surfaces to become highly facetted. This facetting is believed to arise from the surface energy and activity differences between the different crystal planes 34 . Academic interest in the reaction lies in the challenge of elucidating details of the reaction mechanism over which, despite n u m e r o u s studies, there is still much uncertainty. Industrially, the challenge is to bring the r e a c t i o n t e m p e r a t u r e d o w n while m a i n t a i n i n g a h i g h efficiency of NO production and avoiding N2 production. R e c e n t l y Bradley et a135 investigated this reaction on Pt{100} over the t e m p e r a t u r e range 300 to 800 K using molecular beams, arriving at a new m e c h a n i s m for the reaction. Firstly, using mixed NH3 + 02 b e a m s they d e m o n s t r a t e d that the single crystal results mimicked those from polycrystalline gauze catalysts: at crystal temperatures between 450 and 550 K, the dominant

87 0.02

Figure 7. Reaction products NO, N2 and H20 monitored as a function of time as an NH3 beam of 0.019 ML s -1 is incident on an O-saturated Pt { 100} surface, at a surface temperature of 450 K35.

Ts =450K

NO 0.01-

0.00

-10

N2

'

I

0

'

I

'

i

'

l

'

i

'

l

'

l

'

i

'

l

'

I

'

10 20 30 40 50 60 70 80 90 100

Time / s products were N2 and H20, while above 600 K NO was the dominant product, with N2 production progressively suppressed with increasing temperature. These results were obtained for all NH3/O2 mixtures between 30 and 90% 02. In a second series of experiments, using temperature p r o g r a m m e d reaction spectroscopy, it was found that after adsorption of 0.3 ML NH3 with varying amounts of adsorbed O, with coverages below 0.15 ML O the dominant product was gaseous N2, with a peak temperature of ~ 380 K, but with larger amounts of adsorbed O, NO desorption became dominant, peaking at 470 K, with eventually complete suppression of the 370 K N2 peak. Crucial results were obtained in a series of isothermal experiments, with O adatoms preadsorbed to varying coverages prior to beaming NH 3 at the surface, at temperatures between 400 and 550 K. A typical result, obtained at 450 K with 0.63 ML preadsorbed O adatoms and an NH3 beam flux of 0.019 ML s -1, is shown in figure 7. Initially, only NO is produced in the gas phase, plateauing at the high rate of 0.013 ML s -1. This is consistent with a sticking probability for NH3 on the O-saturated surface of N 0.70, demonstrating that sites for NH3 adsorption, designated below as 's', are different from sites for O adsorption, designated '*'. However, once a significant proportion of the surface O adlayer has been consumed in the reaction, the product N2 is observed, and becomes dominant. By back-integration it was determined that N2 is only formed when the O coverage falls below 0.16 ML. It was concluded 35 that N2 production arises from the dissociation of NO p r o d u c e d by oxidation of adsorbed NH3. The observed oxygen coverage dependence of the product formation is explained by a sharp fall in the heat of adsorption for dissociation of NO to Na + Oa with increasing Oa coverage, as already illustrated for NO on Ni{100} at 300 K, in figure 2; at O coverages above a critical value, estimated at around 0.16 ML, NOa formed in the reaction on the surface is more stable than Na + Oa, and therefore desorbs with high efficiency. At coverages less than 0.16 ML O adatoms, however, dissociation is energetically favourable, and the recombinative desorption of N2 is fast: gaseous N2 becomes

88 the dominant product. The key steps in the overall reaction, consistent with the data, are as follows. s = top site

NH3(g) + s --+ NH3(a) O2(g) + 2* --+ 20(a) NH3(a) + O(a) + * -+ NH(a) + OH(a) + H(a) NH(a) + 20(a) --+ NO(a) + OH(a) NO(a) +* -+ N(a) + O(a) 2N(a) --+ N2(g) NO(a) --+ NO(g)

* = hollow site

"~ 0o > 0.16 ML

J 0o > 0.16 ML

This mechanism provides the basis for improving the operating conditions for the Ostwald process, to avoid the costly loss of Pt catalysts as PtO2 at the high temperatures employed. The catalyst can be run at temperatures as low as 500 K, with very high efficiency to NO production, provided that the O 2 / N H 3 ratio is kept high enough to maintain a steady state oxygen coverage > 0.16 ML, to prevent NO dissociation. Bradley et al demonstrated that 100% efficiency could be achieved by operating at temperatures as low as 500 K with an 02:NH3 ratio of > 46. It remains to be seen whether this can be used to industrial advantage. References o

2. 3. 4. ~

6.

o

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

D. A. King, Surf. Sci., 299/300, 678, (1994). P. Heilmann, K. Heinz and K. M611er, Surf. Sci., 83,487, (1979). A. Borg, A. M. Hilman and E. Bergene, Surf. Sci., ~ 10, (1994). Reviewed in: S. Titmuss, A. Wander and D. A. King, Chem. Rev., 96, 1291 (1996). J. K. Norskov and N. D. Lang, Phys. Rev., B21, 2131, (1980). C. E. Borroni-Bird, N. A1-Sarraf, S. Andersson and D. A. King, Chem. Phys. Letters, 183,516 (1991). A. Stuck, C. E. Wartnaby, Y. Y. Yeo, J. T. Stuckless, N. A1-Sarraf and D. A. King, Surf. Sci., 349, 229, (1996). M. G. Wells and D. A. King, Surf. Sci., 2_9.,9454, (1972). Reviewed in: W. Brown and D. A. King, Chem. Rev., submitted. Y. Y. Yeo, C. E. Wartnaby and D. A. King, Science, 268, 1731, (1995). J. T. Stuckless, C. E. Wartnaby, N. A1-Sarraf, St. J. B. Dixon-Warren, M. Kovar and D. A. King, J. Chem. Phys., 106, 2012, (1997). L. Vattuone, Y. Y. Yeo and D. A. King, J. Chem. Phys., 104, 8096, (1996). L. Vattuone, Y. Y. Yeo and D. A. King, Cat. Letters, 4_!,1119 (1996). P. Thiel, J. Behm, P. R. Norton and G. Ertl, J. Chem. Phys., ~ 7437; 7448, (1983) Jackman, K. Griffiths, Davies and P. R. Norton, J. Chem. Phys., 79_, 3529 (1983).

89

16. 17. 18. 19. 20. 21. 22. 23. 24. 25. 26. 27. 28. 29. 30. 31. 32. 33. 34. 35.

A. Hopkinson, J. M. Bradley, X.-C. Guo and D. A. King, Phys. Rev. Lett., 7_.G1 1597, (1993). A. Hopkinson, X.-C. Guo, J. M. Bradley and D. A. King, J. Chem. Phys., 9_.G9 8262, (1993). A. T. Pastuer, St. J. Dixon-Warren and D. A. King, J. Chem. Phys., 103, 2251, (1995). X.-C. Guo, J. M. Bradley, A. Hopkinson and D. A. King, Surf. Sci., 310, 163, (1994). M. Gruyters and D. A. King, J. Chem. Soc. Faraday Trans., submitted, (1997). P. R. Norton, K. Griffiths, P. E. Bindner, Surf. Sci., 138, 125, (1984). J. M. Bradley, X.-Cu. Guo, A. Hopkinson and D. A. King, J. Phys. Chem., 104, 4283, (1996). M. P. Cox, G. Ertl and R. Imbihl, Phys. Rev. Letters, 5__4,1725, (1985). M. Eiswirth, P. M611er, K. Wetzl, R. Imbihl and G. Ertl, J. Chem. Phys., 9_.G0 510, (1989). A. Hopkinson and D. A. King, Chem. Phys., 177, 433, (1993). M. Gruyters, T. Ali and D. A. King, Chem. Phys. Lett., 232, 1-6, (1995). M. Gruyters, T. Ali and D. A. King, J. Phys. Chem., 100, 14417, (1996). M. Gruyters, A. T. Pasteur and D. A. King, J. C. S. Faraday Trans., 92, 2941, (1996). A. V. Walker, M. Gruyters and D. A. King, Surf. Sci., submitted (1997). M. Slinko, T. Fink, T. L6her, H. H. Madden, S. J. Lombardo, R. Imbihl and G. Ertl, Surf. Sci., 264, 157, (1992). P. D. Cobden, J. Siera and B. E. Niewenhuys, J. Vac. Sci. Technol., A10, 2487, (1992). M. H. Stacey, Catalysis (London), 3_, 98, (1980). G. Chinchen, P. Davies and R. J. Sampson, in Catalysis: Science and Technology, J. R. Anderson and M. Boudart, eds., Springer-Verlag, N.Y., Vol. 8, p. 1, (1987). L. D. Schmidt, J. Vac. Sci. Technol., 1_2,2341, (1975). J. M. Bradley, A. Hopkinson and D. A. King, J. Phys. Chem., ~ 17032, (1995).

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91997 Elsevier Science B. V. All rights reserved. Dynamics of Surfaces and Reaction Kinetics in Heterogeneous Catalysis G.F. Froment and K.C. Waugh, editors

91

Dynamic Phenomena at the Oxide/Water Interface: the interplay of surface charge formation, metal complex adsorption, and dissolution/reprecipitation. by J-F. LAMBERT and M. CHE 1 L a b o r a t o i r e de R d a c t i v i t ~ de S u r f a c e ( U R A 1 1 0 6 C N R S ) , U n i v e r s i t d P i e r r e M a r i e Curie, Tour 54 - 4, Place J u s s i e u 7 5 2 5 2 P A R I S C E D E X 05 - F r a n c e

et

Keywords:

Heterogeneous catalysts preparation, interface (oxide/water), interfacial coordination chemistry, adsorption mechanisms, kinetics of interfacial phenomena.

Abstract:

The present review is concerned with phenomena occurring during the earliest stages of the preparation of supported transition metal catalysts, i.e. at the interface between aqueous solutions and solid oxides. In an order of increasing time scales, we consider successively the formation of surface charge (protonation/deprotonation) and its structuring effect on the neighboring solution, the various mechanisms of transition metal complex (TMC) adsorption- electrostatic, outer sphere complexes and inner sphere complexes - and dissolution/reprecipitation reactions. It is shown that a global theory of interface phenomena relevant in heterogeneous catalysts preparation should combine concepts and results derived from various disciplines: colloid science, electrochemistry, supramolecular chemistry, coordination chemistry and geochemistry.

1 Institut e-mails:

Universitaire de France [email protected] -

[email protected]

92

Introduction The present communication is concerned with dynamical phenomena occurring during supported metal catalysts preparation, i.e., during all steps prior to the establishment of reaction conditions. This includes all pretreatment steps such as drying, calcination, reduction, but we will mostly focus on the very first steps during which transition metals are deposited on the surface, most often from aqueous solutions. This field of research has been comparatively neglected until recently, because it was felt that any phenomena occurring at such an early stage would be blurred out during the remainder of catalyst history. It seems, however, that the details of the inception of metal/support interaction do matter, and that their fine control is essential to achieve reproducible catalyst preparation. Even though "molecular design" of supported transition metal catalysts has not yet reached the sophistication attained in the field of homogeneous catalysis, this is not an impossible goal and some recent reviews have contributed to build bridges between these two disciplines I-4. We will chiefly be concerned with "wet" interfaces, between solid supports and an aqueous phase. Unfortunately, most of the methods that allow precise interface structural studies are not applicable in these conditions; however, several spectroscopic techniques may be used, preferably in combination with macroscopic adsorption measurements, to arrive at a reasonable degree of local characterisation of the solid/water interface.

I. Catalyst supports and their surfaces" structural problems Supports currently used for metals dispersion are mostly high-surface oxides (excluding carbon, but including days and zeolites). In some cases, the support oxide is well-crystallised and exposes a small number of different faces with low Miller indices. This situation is illustrated in Figure la by the case of the (110) plane of TiO2 (rutile). The upper part of the figure shows a fresh cleavage plane in which atoms have been removed so as to insure electroneutrality: note that there is a high surface concentration of coordinatively unsaturated (cus) Ti4+- in coordinence 5 instead of 6. The lower part indicates a possible transformation of this surface in the presence of water: it is to be understood that a molecule of water has adsorbed on each cus Ti4+, and then protonated a neighbouring oxide in an acid-base reaction, resulting in two different types of surface OH- (see w I.A.). This deceptively simple picture was obtained from results obtained for single crystals under controlled atmosphere, in the presence of water vapour, and may not be completely transposable to rutile in water suspension. Another immediate objection is that very small oxide particles will expose a large proportion of atoms at edges and comers: the.se atoms will be undercoordinated with respect to those found on low-Miller index surfaces. The problem is of

93

Ti4+ 0

0H

(terminal)

o,.@ 9

@

~ @ (!~2)

02" @ @ (ter0miHal~ 0 AI3+ 16-c1 OH O 0 AI3+ (4-c) (l~2)

b

a S

S:

C

- ~) Si4+ (~

(terminal) (~) 0X (geminal)O

Figure !: Models of oxide/ aqueous solution interfaces la. TiO2 (rutile) - above: fresh cleavage plane; below: after water adsorption lb. above: structure of a cleavage plane of y A1203; below: assumed structure for the hydrated surface of y A1203 (model constructed from the data of ref. 8) lc. amorphous SiO2; the figures in the cycles indicate the numbers of (SiO4) tetrahedra. Adapted from ref. I 0.

94 course also relevant to metal particles modelling s. An instance for oxide surfaces is the CoO-MgO system with rocksalt structure - Co 2+ ions on edges and comers, with coordinence 4 and 3, have different adsorptive properties from those on (100) faces (coordinence 5) 6,7. The case of alumina is more complicated than that of TiO2, since most forms contain 6-coordinated as well as 4-coordinated A13+. Figure lb represents the most commonly accepted model for the surface of ~/A12038 (an improved version has been published recently9). This oxide has a spinellerelated structure and is known to chiefly expose its (110), (100) and (111) planes. We illustrate the (110) plane, or rather one of its two termination modes (plane C). Note that the termination indicated does not insure electroneutrality; however, the condition is satisfied by adding OH- groups to each cus A13+ (lower part of the figure). In the present case, no relevant data can be provided by monocrystals studies since the latter are concerned with a different form, (xAI203 (corundum). Finally, highly dispersed silicas are generally amorphous. There is no general consensus on the best representation of a silica surface; most authors use either a crystalline form of SiO2 (especially cristobalite) as an approximation of the real structure, or randomly connected networks of (SiO4) tetrahedra, which are terminated in the surface plane by Si-OH groups. Figure I c, adapted from ref. 1o, illustrates a model obtained by the simulated annealing technique that evidences the coexistence of cycles made of 3 to 8 (SiO4) tetrahedra, 6-cydes being most frequent.

II. Surface charge formation II.A.

Origin

The existence of surface hydroxyls is evidenced in the above schemes. These surface groups may be protonated (reverse of reaction(I)) or deprotonated (reaction (2)), according to the solution pH:

+ H+ t

S OH

(1)

Ka2 S-O-

+

H+

(2)

Therefore, a pH-dependent surface charge will be developped. In particular, there will be a value of pH for which the total electric charge of the surface is zero: the point of zero (net protonic) charge, or PZC (PZNPC). At lower pH values, the surface bears a global positive charge, and at higher pH, a global negative charge. Typical PZCs are 2-2.5 for silica, 8-8.5 for alumina and 6.5 for titania.

95 The most commonly used model Cl-site, 2-pK") assumes that only one type of surface groups is contributing to the formation of protonic charge. In that case, PZC= 1/2(pKal+pKa2). Such an assumption is clearly an oversimplification for surfaces such as those of alumina and titania, for which several different surface OHs are expected. For instance, it is obvious from Figure 1 a that two different OH types should be present on the surface of futile (110): bridging OHs are expected to be more acidic than terminal OHs, and therefore to have lower values for pKal and pKa2. The same conclusion holds true for yA1203 (110) (Figure lb), where the number of OHs is increased by the presence of A13+ with coordinence 4 and 6. Taking into account the other exposed faces, it is generally assumed that 5 or 6 types of OH are present on the surface of yA12038. As for silica (Figure lc), one can distinguish between terminal and geminal Si-OH, in approximate proportions 3:1. Recently, a more sophisticated mathematical treatment was developped to extract pKa values from the proton titration curve of a chemically heterogeneous surface ("multisite, 1-pK") 11, allowing to estimate pK values for each distinct OH type. This approach seems to have interesting prospects 12. Proton adsorption/desorption kinetics may be studied by pressurejump type techniques. Protonation is usually very fast; deprotonation may be slower but time scales of a few tens of seconds are not exceeded 13. For practical purposes, the oxide surface charge can be considered as being instantaneously established on contact with the metal-containing solution. II.B. S t r u c t u r i n g

effect on the

solution

Even in dilute electrolyte solutions, many properties of water are strongly modified in the immediate vicinity of a charged interface, especially its density. This effect has chiefly been studied by electrochemists interested in charged metallic surfaces: water close to a silver electrode bearing a surface charge of 0.25 Coulomb / m 2 reaches a density of 2.0 14,1 s, i.e., dose to that of ice VII which is formed at a pressure of 2.2 GPa (22000 atmospheres). This phenomenon is called electrostriction. In the first layer of water molecules, the interaction of water dipoles with surface electric field can be an order of magnitude stronger than hydrogen bonds (which are themselves an order of magnitude higher than kT at room temperature) and therefore the electric dipoles of the water molecules are strongly oriented, in the direction of the surface or away from it depending on the sign of the surface charge. As a consequence, the electric permittivity ~ falls down from Ebulk = 78, to values as low as 2.516. It is expected that ion pairing should be favoured in the interface region with respect to the bulk solution. Furthermore, transition metal ions located in this region will have a strongly decreased mobility as though they were imbedded in a glassy matrix 17.

96 The density and other properties fluctuate wildly as a function of the distance from the surface. One may consider that important perturbations from bulk water properties will be encountered in a region 10 to 15/~ thick 18. Furthermore, in a non-zero electric potential ~, such as that imposed by a charged surface, the concentrations of charged species are strongly modified. This effect will be developped in w III.A, as it is the key to electrostatic adsorption of metal ions; however, one must realise that it also has an influence on local pH - in the vicinity of a negatively charged surface, for instance, (OH)- ions will be repelled while (H30 +) ions will be concentrated, resulting in a pH drop with respect to the bulk solution. The opposite is expected close to a positively charged surface. This may well have effects on the speciation of transition metal precursors. It is sometimes hard to evidence because one needs a sensitive spectroscopic probe able to distinguish a surface species, possibly present in small amounts, against the background of unmodified bulk species. For the molybdate/ y alumina system, 95Mo NMR has shown that the alumina surface could promote the basic hydrolysis of heptamolybdate to monomolybdate in its vicinity 19_21. Until recently, the extent of pH modification could only be inferred indirectly. Measurements have become available using the second harmonic generation technique. For the aqueous solution/silica interface, they yielded pH variations between bulk and surface of-0.6 (at pHbulk = 6) or -2.3 (at pHbulk = 12)22; the effect is expected to be more important for the aqueous solution/alumina interface, since the extent of surface charging is higher for alumina.

III. Metal complex adsorption mechanisms: successive degrees of intimacy in the metal/surface interaction III.A. E l e c t r o s t a t i c

adsorption:

the

double

layer

The existence of a surface charge cr will result in an electric potential ~, dependent on the distance to the surface, and related to o through Poisson's equation. 9 is not easy to measure directly, but the problem of evaluating it as a function of o and the distance from the surface (x) has long been solved 23, at least for simple surface geometries: 9 decreases very quickly as a function of x (it becomes negligible within a few tens of A in realistic situations). With respect to the bulk, a species with charge Z has electrostatic energy Zq~ (with q=elementary charge), and its concentration is multiplied -z_~ by the Boltzmann factor, e k T 9 For example, the concentration in H30 + at the surface, [H30+]s, is related to the bulk concentration [H30+]b and to the

97 _~o

surface potential ~~ by: [H 30+] s=[H 30+] b e leT , and the consequences of this dependence have been treated in w II.C. More generally, all ions in the solution will have concentrations affected by the Boltzmann factor. The result is illustrated in Figure 2 for a positively charged, planar surface immersed in a 1:1 electrolyte: in the region of nonnegligible ~, there is a substantial excess of anions over cations. This region is called the diffuse layer (or diffuse swarm) and will contain a net global charge (negative in the example) of equal value, but opposite sign to that of the surface layer (positive here). The "double layer" (surface + diffuse layer) is analogous to a condenser, with capacitance values of several tens or h u n d r e d s of ~ C / c m 2.

I

. .

-

~'~'~-'~~

~O

Figure 2: A very simple model of electrostatic adsorption on a negatively charged oxide surface with formation of a "double layer" (surface + diffuse layer). Small dosed circles are cations, larger open circles are anions. Oo-:surface charge density; x: distance from the surface into the solution; ~:: thickness of double layer; ~: electric potential; c+(x) and c-(x): local concentrations in cations and anions, respectively. The shaded area represents the excess of cations over anions in the diffuse layer, and therefore the amount of cations that are electrostatically adsorbed. Thus, if the solution contains an appreciable a m o u n t of metal ions with a charge opposed to that of the surface, the diffuse layer will be enriched in them with respect to the bulk. They will be "strongly held" in the sense that they cannot be washed away with distilled water, because the charge of the surface layer must remain compensated; however, they can be easily replaced if the solid is washed with a solution containing ions of the same sign, because the electrostatic interaction is non-specific. We m a y

98 contrast this situation with inner sphere complex formation (w II.C) which implies chemical bonding and is therefore highly specific. We know that support oxides may have positive as well as negative charge as a function of pH. It follows that, as one raises the solution pH and crosses the p H p z o the support will change from an anion exchanger to a cation exchanger. This importance of pH control and pHpzc determination has been realised for a long time 24, even though other notions from colloid chemistry (such as specific adsorption and the triple layer) have not yet been fully integrated. One of the terminological problems in the literature on catalyst preparation is that many authors speak of "strong interaction with the support" without bothering to determine if the interaction is specific or not. A purely electrostatic interaction is non-specific, so that any adsorbed ion can be replaced by an ion of equivalent charge; in opposition, the formation of inner sphere complexes (w III.C) implies a chemical reaction and is often highly specific for a given complex. A second problem concerns the use of the term "ion exchange" to refer to electrostatic adsorption in the double layer dose to a charged surface. This term is certainly appropriate for adsorption on solid materials which have substitutional, pH-independent lattice charge (or rather adsorption inside them, since these materials, chiefly clays and zeolites, generally have important microporous volumes). However, it may be confusing for adsorption on surfaces bearing OH groups, because some models in colloid chemistry use "ion exchange" to refer to a process of specific adsorption (see w III.C), probably involving inner sphere complex formation. "Electrostatic adsorption" should therefore be preferred. III.B. H y d r o g e n

Bonding

and

Outer Sphere

Complexes

An intermediate situation between the non-specific electrostatic adsorption and the highly specific inner sphere complex formation is the formation of an outer sphere complex where the original ligands remain coordinated to the metal ion, but form weak specific bonds with the surface, for instance hydrogen bonds. An obvious analogy exists with the field of Supramolecutar Chemistry, where second sphere coordination has been evidenced in the interaction of transition metal complexes and macrocydes. Macrocycles are cyclic molecules containing a fixed number of oxygen atoms which are Br/3nsted and Lewis bases, such as the crown ether 18C6 (6 oxygens in a 18member ring). Complexes with ammine ligands form stable adducts through a regular pattern of hydrogen-bonding with cycle oxygens 25. The siloxane oxygen atoms in 6-cycles of (SiO4) tetrahedra on the surface of silica have an arrangement rather similar to 18C6 (figure 3). Precise structural identification of outer-sphere metal complexes in adsorption studies is infrequent, but a few examples can be found in the literature 26. Recently, we have proposed that specific hydrogen bonding between the ammine ligands and the silica surface played an important role in [Ni(chxen)2]2+ adsorption on SiO227 (where chxen is the cyclohexanediammine ligand). Further progress will depend on the

99

I

A9

ct

)NH3

Oo

Figure 3: (left) structure of the adduct [Pt(NH3)2CI212:18C6:(DMA)2(from2 5), outlining the pattern of hydrogen-bonding between complex ion and macrocyde; (right): a hypothetical outer-sphere complex of the same complex ion with surface groups of silica. application of spectroscopic techniques providing information bonding, such as IR in the OH stretching region, or 1H NMR. III.C.

Inner

Sphere

on

H-

Complexes

III.C.1 Solution vs. Interracial Coordination

Chemistry: Thermodynamics and Kinetics It is difficult to evidence pure cases of inner sphere complex formation between surface groups and transition metal ions (also called "grafting") since other phenomena are usually occurring in parallel. The clearest instances are observed when "spectator" ligands are inert to substitution, either bec_aj_~seof chelate effects (c/s- [Ni(en)2(H20)2] 2+ on various supports 25) or because of high crystal field activation energies ([Co(NH3)5(RO)] x+ on rlA12032 9; RO= OH, H20 or alcohol). The Ni II complex ds-[Ni(en)2(H20)2] 2+ can be grafted on the surface of SiO2, yA1203 and Y zeolite to give the cis-disubstituted [Ni(en)2(SO)2]x+, where (SO) is a surface oxide or hydroxide group 2 8. This grafting starts upon contact with the solution and is completed upon drying in mild conditions. It is evidenced most directly by shifts to low-energy of the UV bands corresponding to the d-d electronic transitions of the Ni II ion. These spectroscopic results can easily be interpreted since the theory of d-d transitions in metal complexes has long been established in the field of homogeneous coordination chemistry. The various ligands encountered in homogeneous chemistry have been classified according to the intensity of the crystal field (z~) they will impose on a transition metal ion: this is called

100

the spectrochemical series of ligands. For example, one has: ... ACI_ < AH20 < ANH3 ... Our results showed how to fit different surface groups of oxide support in this series: Acl_ < AAIO < AZO < ASiO < AH20 < ANH3 ... (A10 = surface group of yA1203, etc.). Thus, common oxide supports will behave as weak-

field ligands, a fact that has some theoretical justificatio 30. As a consequence, the crystal field stabilisation energy

(CFSE) for grafted metal complexes will be less negative than for the original, homogeneous-phase complex, if one starts with aqua complexes (which is most usual in catalyst preparation) or with ammine complexes. Since the change in CFSE is a major component of the adsorption enthalpy AHads, transition metal adsorption is not expected to be strongly exothermic. There is unfortunately a dearth of thermodynamical data concerning transition metal complexes adsorption on oxide surfaces. We may, however , draw some conclusions from related systems such as Cd 2+ / oxides which were studied for environmental reasons. In a very thorough study of cadmium adsorption on TiO2 (rutile) and Fe203 (hematite) 31, Fokkink et al. could resolve the enthalpic and entropic contribution to inner sphere cadmium adsorption. The (intrinsic) AHads had negligible values, b e t w e e n - 7 and +5 kJ/mole; still adsorption occurred because it was entropically favoured, with ASads in the +110 to +120 J/(mole.K) range. The authors' interpretation was that the strongly negative hydration entropy of cadmium ions is reduced when they are "stuck" to the surface and thus deprived of about one half of their hydration shell. This phenomenon of solvation loss should be most important when the complex electric charge is reduced upon ISC formation. The idea that entropic effects can be predominent in metal complex adsorption will have to be considered in forthcoming studies, especially in the light of our newly gained knowledge on interface organization (w II.B). The kinetics of adsorption by inner sphere complex formation is even less studied than its thermodynamics, but electrochemists have devised good operating models for it32 where electrostatic factors again play an important role. Characteristic times for ISC adsorption range from 10s to 105s at RT (for Zn II and Cr III on yA1203, respectively). Hachiya et al. 33,34 have evidenced a good correlation between the rate constants for adsorption of various metals on yA1203, and the rate constants for hydrolysis of the same metals in aqueous solutions (both transitional and non-transitional metals were considered: Cu 2+, Mn 2+, Co 2+, and Cd 2+, Pt~+). This is not surprising since both ISC adsorption and hydrolysis are ligand exchange reactions: the same ranking of metal ions in order of lability (vs. inertia) can be observed in both cases, and is justifed once again in terms of crystal field theory.

101 III.C.2 The triple layer and "specific adsorption" The triple layer model attempts to take into account inner sphere complex formation and electrostatic adsorption simultaneously by considering "specifically adsorbed" ions 3s which are supposed to be maintained very close to the surface, whether it be through the formation of covalent bonds with some surface groups, or of some outer sphere complex. No specific interpretation of the bonding is required, provided one can define a plane of specific adsorption, located a few A from the surface and containing those ions: this is called the Stern layer. The theory distinguishes then between three successive parallel layers: the surface plane proper, the Stern layer, and the diffuse layer. .

A

r

/

y,,~

f

D/,,~

Diffuse layer

l= diffuse swarm

O

9 AI3+

O2"

@ O H (terminal) @ O H (bridging)

~

OH2 + (protonated OH)

~(Mo7024)6. ~ in Stern layer layer ~,~/,~ ~f alumina (NH4)+ ~/~ in diffuse (Gouy) Surface ~

layer

Su ce plane

_~

--Gouylayer

Outer Helmholtz plane, front of diffuse layer Inner "d" Helmholtz plane or Stern layer

Figure 4: Semi-realistic (A) and schematized (B) representations of the triple layer in the case of ammonium molybdate adsorption on alumina at low pH (modified, from refs. 38-39). The surface plane is positively charged, the Stern layer contains an excess of negatively charged heptamolybdates, and the diffuse layer mostly contains NH4+.

It may occur that specifically adsorbed ions overcompensate surface charge: for example, the surface is positive because pH < pHpzc, hut a~ excess of anions lie in the plane of specific adsorption so that the global charge, of the oxide particle is negative Csuperequivalent adsorption ''36 with "charge reversal"37). This is exactly the situation described by Spanos et al.38 for

102 heptamolybdate adsorption on y alumina (figure 4): in the pH range 5-8, the surface charge measured by acid-base titration is positive while the global particle charge measured by electrophoresis is negative, so that coadsorption of ammonium ions occurs. We have witnessed a similar situation for paratungstate adsorption on ~/alumina 39. It may be remarked that the low values of the dielectric constant in the vicinity of the surface should favour ion pairing, and thus coadsorption. The triple layer theory has also been applied recently by Kn6zinger e t al. for [Pd(NH3)4] 2+ and [PtC16]2- adsorption on ~/A120340. III.C.3 M o l e c u l a r d e f i n i t i o n o f a d s o r p t i o n sites In order to fully reconcile the coordination chemistry and colloid chemistry approaches to inner sphere adsorption, one would like to compare the ligand properties of surface groups of various supports with those of molecules in homogeneous solution. Some elements of such a comparison have already been outlined in w but a systematic parallel study would necessitate a precise characterization of the adsorbed complex at the molecular level. There are indeed examples of this kind of characterization, but they constitute the exceptions rather than the rule.

A

o2

(~~ ~ H20

erminal) ~

B

(terminal)

Figure 5: A comparison of inner sphere adsorption sites of [Cu(H20)6]2+ on TiO241 (A) and on the Al13 pillars of a pillared clay 4 2 (B). [Cu(H20)6] 2+ adsorption on TiO2 (anatase) has been studied by Ludwig et al. who showed that the support acts as a bidentate ligand 41. The nature of

the inner sphere adsorption sites is illustrated in Figure 5A, and can be formulated as Ti-OH-Ti-OH(2), i.e. both a bridging OH (between two TiTM)

103 and a terminal OH (possibly protonated, hence the OH(2) notation) penetrate the CuII coordination sphere. An independent study of [Cu(H20)6] 2+ adsorption into pillared clays 42 revealed a very similar mechanism, with inner sphere adsorption on AI-OH-A1-OH(2), also implying a bridging and a terminal OH borne by the same cation (Fig. 5B). Another instance of precise site identification may be found in the Ni/silica system. In contrast with other examples quoted in the present review, it is concerned with the solid/gas interface. Phyllosilicate-containing Ni/SiO2 samples (cf. w IV.C) were thermally treated under flowing 02 at 773K, then under vacuum at 973K. Coordinatively unsaturated Ni(II) ions were formed in the process; studies by W-visible and EXAFS spectroscopy, and EPR after photoreduction to Ni(I)43, 44 revealed that the ions were coordinated to 3 oxide ligands from the support. However, their exact disposition was not dear until an extended H(ickel treatment of several competing models was carried out 4s. The most probable model for the Ni(II)/silica surface interaction is represented in Figure 6.

) Bridging 0 ) OH (terminal)

3-coordinated coordinatively unsaturated (cus) Ni(II) at the surface of silica after high-temperature treatment in vacuum. Ni(II) is coordinated with 3 terminal silanol groups of a flexible 6-cycle of (SiO4) tetrahedra. This example is mentioned here, although not directly concerned with the solution/oxide interface, to emphasize the large amount of work that is sometimes needed to unravel the structure of adsorption sites at the molecular level. It should not be surprising therefore that in most cases, molecular information on adsorption sites is scanty or even absent.

IV.

Dissolution/reprecipitation IV.A. W e a t h e r i n g "

the

lessons

of geochemistry

It has become dear in recent years that the formation of mixed phases, containing both the transition metal and an element from the support, did not necessitate high calcination temperatures: indeed it occurs very commonly during the deposition step itself. In particular, talc- and nepouite-

104 related Ni(II) phyllosilicates were formed in the [NiII(H20)6]/SiO2 system46, 47, and layered double hydroxides (LDHs or hydrotalcites) in several [MII(NH3)6]/A120348 and [NiII(H20)6]/A120349 systems. It turned out that this came as no surprise to the scientific community of geochemists. They have been studying for years the alteration of primary minerals on weathering, i.e,. those changes brought about by climatic factors, including exposure to the soil solution. Dissolution followed byneoformation of secondary phases usually results. Of course geochemistry is often concerned with "geological" time scales, but this is not necessarily so. Many relevant data are to be found in ref.50. To give orders of magnitude for support oxides, the half-life of a surface A1 is 669 days (5.8"107 s) for ctA1203, but only 44 days (3.8"106 s) for 6A1203s 1. As we will discuss in w IV.B, the presence of TMC ions may speed up dissolution phenomena. In the case of [MgII(H20)6]/SiO2, the time scale for dissolution/reprecipitation was typically a few tens of hours, or 105ss2, i.e., an order comparable to the duration of unit operations in catalyst preparation. Furthermore, catalyst preparation often includes a drying step at temperatures between 90 and 120~ In this temperature range, some water solution may remain in the porosity for a significant lapse of time before drying is completed, and of course dissolution reactions will be speeded up with respect to room temperature s3. We shall now mention a few geochemical phenomena that are likely to be relevant in the understanding of TMC deposition procedures, but this is far from exhausting the subject. IV.B. P r o t o n - p r o m o t e d dissolution

and

ligand-promoted

One of the experimental facts long established in the field of geochemistry is that the solubility of an oxide may be profoundly altered by the presence of various ions in solution, including TMCs. The effect of protons is to increase the dissolution rate through surface protonation. For instance, if a bridging surface OH group is protonated, this will weaken the bonds to both neighbouring A13+ and thus facilitate their detachment from the surface 50 as indicated in Figure 7, scheme A: As for TMCs, they will have an influence on both the thermodynamics and kinetics of reaction. The thermodynamic effect is easier to understand. It results from the fact that species normally formed by the dissolution of oxides (such as [Al(H20)6]3+ or HxSiO4) may react with transition metal complexes in solution, thus dragging the reaction of oxide dissolution to the right. Such a phenomenon seems to be present in the molybdate / alumina system s 4. The kinetic.effect is due to the fragilisation, or inversely the protection, of support surface gr0ups by adsorbed complexes. For instance, it is established that inner-sphere adsorption of dianions (such as MoO42- or W O42-) on a surface cation fragilises its remaining bonds to the solid, thus increasing the rate of its dissolution (Figure 7, scheme B).

105

!~iii!!!!L!.!i!!!!iii!~i/ii!iii!i!!!!!!!!!!!!!!i!!!!i!i!!iiil

+ H+

@ 20

@

+ OH"

I

2 OH" . . . . . . . . . . . . . . . . . . . . . . . ...............................................

@ ~20 + OH"

x~i~i~ii~i~iiii~i~i;~i~ii~i~i~iiiiii~i~i~i~

Fig~ure 7: Several models

:!,i}:{ ,i J ii~i i/i i~i/:ili31USil}:i!iiii:i;ii

for monotungstate ISC adsorption on an alumina surface, and their

effect on alumina solubility (partly from ref. 5 tl, modified). See discussion in text.

106 This is supposed to be due to the transfer of electronic density from the anion acting as ligand to the A13+ cation, resulting in a weakening of its bonds with framework oxide ions. However, one might imagine a different mode of adsorption, in which tungstate adsorption indeed slows down A13+ dissolution by forming a bridging group between previously unconnected surface ions, as illustrated in scheme C of Figure 7. Finally, scheme D is probably more realistic than either of the previous ones. It is hard to determine without further study which effect will predominate in this case, but in any case tungstate adsorption may be expected to exert an influence on alumina dissolution rates.

IV.C. P r e c i p i t a t i o n

and

coprecipitation

Cases are known in soil science where the adsorption of transition metal complexes on a surface is followed, at higher loadings, by the formation of metal-oxygen-metal linkages, and finally by the nucleation of a new metal-containing phase which is often a metal hydroxide. A typical example may be found in the work of Chisholm-Brause et al. on the CoII/yA1203 systemS s, 56. It must be underlined that the metal hydroxide phase forms on the surface under conditions where there is no bulk precipitation, whether it be due to surface conditions different from the bulk, to heterogeneous nucleation on the grafted metal complex, or to a combination of both effects; whatever the mechanism, the term of surface precipitation is used for this phenomenon. We have previously underlined 57 the possible role of olation and oxolation reactions in the formation of metal hydroxide/metal oxide nuclei. In further steps of catalyst pretreatment, such nuclei may actually be reduced to metal(0) particles without significant size changes: thus, metal dispersion may be determined from the earliest stages of supported metal preparation s3, and particularly before the final reduction stage, a fact that is not often realized. An example of surface precipitation in the field of supported catalysts preparation can be found in the deposition-precipitation method, wherein slow metal precipitation is caused by the hydrolysis of urea at high temperatures. Burattin 55 has shown that deposition-precipitation of Ni II on low-surface area silica resulted in the formation of lamellar Ni hydroxide, Ni(OH)2. However, on high surface area silica, the same procedure gives rise to the formation of nickel phyllosilicates. It is not yet known with certainty if this occurs through surface dissolution of the silica support (eventually promoted by Ni II adsorption, cf. w IV.B) followed by coprecipitation of a mixed phase, or through an interface reaction between Ni complexes in the solution and the modified silica surface. The neoformed silicate layers are observed to be in narrow association with the silica particles, which would rather suggest an interface

107

reaction; but much work remains to be done to elucidate the mechanism of phyllosilicate formation.

V. Conclusion. In this review, we have presented a short overview of interface phenomena relevant to the first steps of supported metal catalysts preparation. However sketchy the presentation may be at times, it gives an idea of the complexity of the oxide/TMC solution interface, where many intertwined processes are occurring with time scales ranging from a few seco.nds to 105 s or more. Figure 8 attempts to specify the relevant time scales according to the type of phenomenon considered; we also included data on electron transfer or redox reactions, which were not discussed here but may play a role in the reguletion of inner sphere complex formation. Two m a i n conclusions may be drawn: first, none of the above phenomena may be neglected when considering the typical durations of elementary operations implied in catalyst preparation, be it in the laboratory or in industrial processes; second, within each class of phenomena, very large variations may be encountered so that it is not hopeless to try and isolate the effect of a single type of interaction by careful parameter setting.

Duration of

Interference

unit operations

with diffusional phenomena

/i :/:::i):!:~)j :/ii~i~Li~ii,':~:ii?!ii)~i~/i:~!i)i~:i ~i!~:i!~i!:~i:ii~,~ 84 : i:~:~~i!i~:,~~i~i!(i:i!?i~/i:i~:~(~::iii!~~i!:~),?i~,:~!~i~!Ei|ga i n diii!i ~~ie~hian~:e:ii!~ (ilil;~~i!~(:ii:i:~~!i!)z,~ii!i,7, ~, ~i:)~/i::i, ~?:i(:~i : ,7(,i:i ~ii:

:/

:

: : i : ,i!,: i :

Electron transfer

:Oi

~id !0~C i a ~

I

'i)!ii!ii /i 'J: : :.

~leprec'ip i t a t i O n

1 "0-10 1"0 -8 1"0.6 1"0.4 1"0 -2 1"00 1"02 1"04 1"06

1"08

Time scale (s)

Figure 8: Approximatetime scalesfor interfacephenomenarelevantto supportedcatalyst

preparation. OSC adsorption and proton transfer are discussed in w Ill.B, ligand exchange and ISC adsorption in w UI.C, dissolution-reprecipitation in w IV. The data for redox phenomena in solution were taken from ref. $ 9.

According to the class of phenomena to be considered, we drew heavily on results obtained in different scientific disciplines: colloid chemistry for interface charging and electrostatic adsorption, supramolecular chemistry for outer-sphere interactions, homogeneous-phase coordination chemistry for inner-sphere adsorption, and soil science or geochemistry for dissolution/reprecipitation. These different aspects are still far from being integrated into a coherent picture. The scientist specialising in heterogeneous catalyst synthesis works in an area of chemistry that is still in the

108 c o n s t i t u t i o n stage. H e / s h e m u s t be an o p p o r t u n i s t i c s c a v e n g e r , r e a d y to i m p o r t a n d a d a p t concepts from s e e m i n g l y u n r e l a t e d fields of science: b u t each one of these fields has n o w m a t u r e d to such a degree that their contact and s y n e r g y is likely to bring significant b r e a k t h r o u g h s in our u n d e r s t a n d i n g and utilisation of interface p h e n o m e n a .

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91997 Elsevier Science B. V. All rights reserved. Dynamics of Surfaces and Reaction Kinetics in Heterogeneous Catalysis G.F. Froment and K.C. Waugh, editors

111

The Influence of Oxygen Poisoning on a Multiply Promoted Iron Catalyst Used for Ammonia Synthesis: A Temperature-Programmed Desorption and Reaction Study E Rosowski* and M. Muhler** * Fritz-Haber-Institut der Max-Planck-Gesellschaft, Faradayweg 4-6, D- 14195 Berlin * Ruhr-Universit~it Bochum, D-44780 Bochum, Germany A multiply promoted iron catalyst used for ammonia synthesis was studied in a microreactor flow system equipped with a calibrated mass spectrometer. By feeding synthesis gas with about 5 ppm of oxygenic compounds at 603 K, the effluent ammonia concentration was decreased by a factor of 4. It was possible to regenerate the catalyst at 723 K by feeding purified synthesis gas indicating reversible poisoning. The temperature-programmed desorption of N2 (N2 TPD) and the temperature-programmed surface reaction (TPSR) of adsorbed atomic nitrogen with H2 were studied in the active and in the poisoned state. The activation energy of N2 desorption was found to increase from 146 kJ/mol to 174 kJ/mol due to oxygen poisoning in good agreement with values observed for potassium-promoted and potassium-free catalysts, respectively. The TPSR experiments in the poisoned state revealed that the onset temperature of NHa formation was shifted by 10 K to higher temperatures, and that the peak shape resembled potassium-free catalysts. These effects indicate that oxygen poisoning mainly affects potassiumpromoted sites thus transforming the catalyst from the promoted state with an oxygen-deficient K+O coadsorbate layer into an essentially unpromoted state with an oxygen-saturated K+O coadsorbate layer. 1. I n t r o d u c t i o n

Since the beginning of the industrial synthesis of NHa from N2 and H2 on multiply promoted iron catalysts, the phenomenon of oxygen poisoning has been recognized. A review on the deactivation of NHa synthesis catalysts was published recently [1]. Oxygen-containing compounds like 02, H20, CO or CO2 are designated as temporary poisons since they lower the activity while present in the synthesis gas whereas permanent poisons like sulfur-containing compounds accumulate on the catalyst surface and cause an irreversible loss of catalytic activity [2]. Especially alkali promoted iron catalysts are highly sensitive to oxygen poisoning which was reported as early as 1926 [3]. It has been demonstrated recently that already a few ppm of 02 in the synthesis gas lower the NHa activity of alkali-promoted iron catalysts considerably [4]. The poisoning efficiency of the oxygenic compounds was found to be equivalent per oxygen atom due to the conversion to H20 at the top layer of the catalyst bed [3,5]. Hence only H20 is present in the major part of the catalyst bed establishing a dynamic equilibrium between H20 and H2 in the gas-phase and adsorbed atomic oxygen (O-,) on the catalyst surface. The equilibrium coverage of O-, (| depends on the thermodynamic parameters such as gas c o m -

112 position, pressure and temperature. A procedure to evaluate the synthesis gas purity has been suggested based on the time necessary to achieve equilibrium after changes in these parameters [4]. Recently the effect of adding 2.9 % H20 to the feed stream was studied over iron-based ammonia synthesis catalysts leading to interesting transient phenomena [6]. NH3 synthesis in the absence of poisons is relatively well understood on the atomic level and a wealth of experimental information is available [7]. Experimental information on NH3 synthesis under partial poisoning by oxygenic compounds is much less abundant and the mechanism on the atomic level is still poorly understood. Stoltze and N0rskov [8] assumed as the simplest conceivable mechanism that O - . adsorbs competitively with the other intermediates. At constant temperature, gas-phase composition and pressure, the relative loss in activity is equal to( 1-(9o )2

[81. Recently, a new experimental approach to study the interaction of N2 and H2 with iron catalysts has been presented. It comprises the combined application of the temperature-programmed desorption (TPD) of N2 and of the temperature-programmed surface reaction (TPSR) of N - . with H2 in a microreactor flow system [9,10]. In addition to the quantitative results which allow on-line determination of the number of Fe surface atoms, the computer-aided simulation provides a way to derive kinetic parameters of the involved elementary steps which are part of the mechanism of NHa synthesis. N2 desorption studies revealed a significant shift of the N2 TPD peak to lower temperatures with the addition of potassium [ 11,12]. In the present publication, the combination of N2 TPD and N - . TPSR experiments will be applied to a multiply promoted iron catalyst before and after oxygen poisoning at 603 K in order to provide deeper insight into the mechanistic aspects of oxygen poisoning and the role of potassium.

2. Experimental The experiments were carried out in an all stainless steel microreactor system with four gas lines which was operated at pressures up to 100 bar. The gases were supplied by Linde with the following purities: He 99.9999 %, N2 99.9999 %, H2 99.9999 %, the mixture of 25% N2 in H2 used as synthesis feed gas 99.9996 %. The feed gas was further purified by means of a purification unit described elsewhere [4]. The flows were regulated by electronic mass flow controllers. The reactor consisted of a glass-lined U-tube similar to the one described in ref. [13]. It was not possible to detect the desorption of N2, H2 or NH3 from the empty tube within the limits of detection. The U-tube was placed in a copper block to ensure isothermal operation. Gas analysis was performed using a mass spectrometer (Balzers GAM 445) which was calibrated for He, H2, N2 and NHz by using a reference gas mixture. The calibration for H20 was carried out using a He stream saturated with H20 at room temperature. A multiply promoted iron-based catalyst (KM1) was used which was supplied by Haldor Tops0e A/S. The BET area determined on-line after reduction was 12.9 mZ/g. Usually 140 to 400 mg of the 250 ~ m - 8 0 0 ~ m sieve fraction were used thus preventing limitations by heat or mass transport [2]. The reduction was carried out in synthesis gas using 40 Nml/min with a heating ramp of 10 K ~ up to 723 K. After 48 h at 723 K, the concentration of H20 was below the detection limit of 1 ppm and the concentration of NH3 had reached steady state. The absence of poisoning by oxygen-containing compounds was tested as described in ref. [4]. It was possible to run NH3 synthesis at steady state at temperatures as low as 543 K. The number of Fe surface atoms was determined on-line by N - . TPSR yielding 94 ~mol/g.

113 When analysing the TPD and TPSR experiments the design equation of a continuous flow stirred tank reactor (CSTR) was used, e.g 9

PN2 po

S dON =

F 2dr

5/3 dON =

F 2dT

(1)

where PN2 is the partial pressure of N2, Po is the ambient pressure, S is the dimensionless total number of sites, ON is the coverage of N - , ranging from 0 to 1, F is the carrier flow rate in number of molecules/s, [3 is the heating rate in K/s. This equation is a good approximation for sufficiently small heating rates and a high amount of inert carrier gas. The heating rates used ranged from 1-10 K/rain, and the He or H2 flow was 50 Nml/min resulting in maximum concentrations of about 270 ppm.

3. Results and Discussion 3.1. Poisoning and regeneration The result of a typical poisoning experiment is displayed in fig. 1. NH3 synthesis was run at 603 K at steady state using 40 Nml/min synthesis gas. After 4 hours the gas purification was bypassed exposing the catalyst to the impurities present in the synthesis gas with a purity of 99,9996% (table 1).

Table 1 The impurities in ppm present in H2 and N2 as specified by the supplier. Hydrogen Nitrogen purity 99.9996% 99.9999% 99.9996% 99.9999% 02 2 0.3 0.5 0.3 H20 1 0.5 1 0.5 CO, C,,,Hm 0.1 0.1 0.1 0.1

Due to the conversion of 02 to H20, the total amount of H20 is about 5 ppm in good agreement with the amount determined on-line by mass spectrometry. When bypassing the purification unit, the NHa yield was found to decrease continuously over a period of about 30 hours until a new steady state was established. The effluent NHa mole fraction decreased by a factor of 4. This experiment demonstrates the high sensitivity of the multiply promoted iron catalyst towards oxygen poisoning. Similar deactivation experiments with time have been described in detail in refs. [1,4]. The amount of O - , deposited during the exposure of the catalyst to the unpurified synthesis gas was determined by a transient H2 step experiment. The flow was switched from the unpurified synthesis gas to He at 603 K and the temperature was increased to 723 K. During the flushing in He, the desorption of H2 and N2 was observed, whereas the coverage of O - , was not affected. The higher temperature had to be chosen in order to achieve the complete removal of O-, because of the high endothermicity of the reaction O - , + H2 = H20 [8]. Then the flow was switched from He to purified synthesis gas while monitoring the effluent mole fractions of H20 and NHa.

114 1800 E

1600

c

1400

c~ c~

0 0

~

$ guard bypassed

"~

active state

poisoning at 603 K ynthesis gas

1200

0

1000

E 7z

_

800

_

C

600

LU

poisoned

400 0

10

20

30

40

Time / h

Figure 1. Poisoning experiment at 603 K. After 7 h NH3 synthesis at steady state using a total flow of 40 Nml/min and a catalyst weight of 140 mg, the purification unit was bypassed exposing the catalyst to the impurities present in 99.9996% pure synthesis gas.

60

E c~

I

i

NH 3

I

r ~O

1

3000

He ~ synthesis gas

2000

40

(D O

regeneration at 723 K 40 Nml/min synthesis gas

E

O -1- 20

E C ~O

O

E 1000

I z

C

C

UJ

W

0

I

I

I

I

100

200

300

400

Time / min

Figure 2. Regeneration experiment at 723 K using a catalyst weight of 300 mg. After 30 min in He, the flow was changed to 40 Nml/min synthesis gas.

115

A typical result is shown in fig. 2. After 30 min in He at 723 K, purified synthesis gas was passed over the catalyst bed resulting in the instantaneous formation of H20 and NH3 (fig. 2). Integrating the amount of H20 formed yields an amount of 20/zmol/g H20 which is equivalent to 20/zmol/g O-,. After passing through a maximum, the concentration of H20 decreased continuously approaching the detection limit of 1 ppm after about 7 hours. During this period, the concentration of NH3 continued to increase reaching the previous unpoisoned steady-state value of 2400 ppm which is determined thermodynamically. After the regeneration, NHz synthesis was run at lower temperatures far from equilibrium yielding the same activity as prior to the poisoning. The experiments thus demonstrate that the effect of cofeeding oxygenic poisons at the ppm level is indeed reversible. 3.2. The influence of O - , on the desorption kinetics of N2 In a recent publication, the results of N2 TPD experiments with various heating rates and initial coverages of N - , have been presented [10]. A monolayer coverage of N-, has been achieved by switching from steady state NH3 synthesis at 603 K to N2 only and cooling in N2 to 300 K. Hence it should be possible to study the influence of O-, on N2 TPD experiments by either using purified synthesis gas at 603 K or unpurified synthesis gas followed by cooling in N2. Trace A in fig. 3 was obtained by using purified synthesis gas and a heating rate of 5 K/rain. The narrow peak with a maximum at 644 K is assigned to desorption occurring mainly without readsorption whereas the broad shoulder at about 770 K is assumed to be due to desorption with readsorption. From experiments with various heating rates, a rate constant of desorption k,~s = 2 x 109s-1 exp(-146 kJ/mol/RT) was derived. Modelling the experimental data yielded the best fit for a coverage-dependent net activation energy of adsorption E* = (-15 + 30 ON) kJ/mol assuming only one type of atomic nitrogen species. When carrying out the N2 TPD experiment in the poisoned state, the observed peak (trace B in fig.3) differs considerably from the unpoisoned result. A single broad peak at about 740 K is found which has roughly the same height as the narrow peak at 644 K observed for the unpoisoned state. One might have anticipated that the main difference between the poisoned and the unpoisoned peaks would be the amount of N2 missing due to site blocking by O-,. Instead, the whole TPD spectrum is found to be changed. Integrating the N2 TPD traces up to 850 K reveals that oxygen poisoning decreased the amount of N2 desorbed from 44/,tmol/g to 36.9/zmol/g. The difference of 7.1 ~mol/g N2 corresponds to 14.2/zmol/g N-,. The influence of oxygen on the N2 desorption kinetics was further investigated by varying the heating rates in the poisoned state. The results are shown in fig. 4 using/3 = 1, 5, and 10 K/rain (traces A, B, C, respectively). The peak height is found to scale with the heating rate and the peak maximum temperature is found to shift to higher temperature with increasing heating rate. The broad maxima, however, do not allow to determine the rate constant of desorption accurately. Assuming an unchanged preexponential factor Aaes = 2 x 109s-1, an activation energy of adsorption Eaes = 172 kJ/mol is obtained which reproduces the shifting peak maxima reasonably well. As derived in ref. [10], the determined rate constant actually corresponds to the rate constant k,,ec of the recombination of two N-, to form N2-* which is the rate-determing step of the desorption of N2.

2N-,

~

N2--,+,

(2)

N2-,

~

N2+,

(3)

116

N2 TPD

150 E Q_ c3 c"

.9 o

A active state

100

A

ID

E r

z

50

_

c (D

_.z UJ

o 500

:

550

600

,

650 700 750 Temperature / K

,

,

800

,

,

850

Figure 3. N 2 TPD data obtained with/3 = 5 K/min after N2 dosing at 603K and subsequent cooling in N2. The catalyst weight was 300 mg. Trace A was obtained in the active state, trace B in the poisoned state by using unpurified synthesis gas prior to dosing N2.

The subscript "rec" used in the following refers to the recombination of two N - , to form N2-*, "dis" refers to the dissociation of N2-*, and "des" and "ads" designate the desorption and adsorption of N2-*, respectively. In the unpoisoned state, the N2 TPD peak structure was found to be strongly influenced by readsorption of N2 within the catalyst bed [ 10]. The following equation for the rate of desorption was derived from equations (2) and (3):

--dON 2dr

=

A~.~cexp(-E~e~/RT)O2N Adis Aads

e z p ( - - ( E d i ~ + (Ead~ -- E d e s ) ) / R T ) p N 2 ( 1

-- ON) 2

(4)

Ades

Hence it is the net activation energy of adsorption E* = Edis + (Eads - Edes) which determines the influence of readsorption on the peak shape. The negligible asymmetry of the TPD traces in fig. 4 to higher temperature indicates that readsorption has only a minor effect indicating activated adsorption in the poisoned state. The results of the TPD experiments in the active and the poisoned state are summarized in table 2. The presence of O-, caused a significant increase of E,.~c by 26 kJ/mol. B. Fastrup [ 12] found E,.ec = 163 kJ/mol for potassium-free iron catalysts compared with 146 kJ/mol for potassium-promoted catalysts. Hence oxygen poisoning seems to affect mainly potassium-promoted sites, thus removing the potassium-induced decrease of E,.ec. When using E,.er = 172 kJ/mol to calculate the outcome of a TPD experiment with/3 = 10 K/s at saturation under UHV conditions, a N2 TPD peak at 880 K results in good agreement with about 850 K observed by Bozso et al. [14] on Fe(111).

117

E

r C"

o

760 K -,,-C ~ ~ ' ~

poisoned

.

.

.

.

N2 TPD

200

0

5

B ~ "

f/740K

,

E 7

100

.//

W

- - ...... ; 550

~ 600

705 K

i 650

!

'

.. A

!~

700 750 Temperature / K

800

850

Figure 4. N2 TPD data in the poisoned state obtained with/3 - 1 K/min (trace A), 5 K/min (trace B), and 10 K/min (trace C) after N2 dosing at 603K and subsequent cooling in N2. The catalyst weight was 300 rag. Poisoning was achieved by using unpurified synthesis gas prior to dosing N2.

Table 2 The rate constants of N2 adsorption and desorption in the active state (ref. [10]) and in the poisoned state. active state poisoned

A,,ec / s -1 2 x 109 2 x 10 ~

E,,ec / kJ/mol 146 172

A* / Torr-ls -1 7.6 x 10 -2 -

E* / kJ/mol -15+30 x ON -

3.3. The influence of O - , on N - , hydrogenation The temperature-programmed reaction between N-* and H2 has been shown to be a reliable tool to determine the active metal area [9]. The saturation coverage of N - , was achieved by switching from steady state NH3 synthesis at 673 K to pure N2 and maintaining the sample at 673 K in N2 for 1-2 hours before cooling in N2 to room temperature. A typical result obtained with the KM 1 catalyst by following this procedure is shown in the upper half of fig. 5 (trace A). As indicated by the arrow, the NH3 concentration has not yet reached the baseline at 700 K indicating an ongoing formation of NH3. In order to remove N - , completely in this case, a holding period at 700 K has to be included in the experimental procedure yielding the amount of 94 #mol/g NH3. When lowering the NH3 synthesis temperature to 603 K prior to switching to N2 at the same temperature, trace B was obtained displayed in the lower half of fig. 5. The onset temperature of 360 K, the peak maximum temperature of 420 K and the tailing to higher temperature of the TPSR trace obtained after dosing at 673 K are closely reproduced. It might be assumed

118

I TPSR 60

E {3.

A N2 dosing

40

c~ r

._o

20

(3 I

.............................................................., V

o

E

C'3

"1" Z

UJ

60

40

20

300

400

500

600

Temperature ! K

700

Figure 5. N-* + H2 TPSR data obtained in the active state after dosing at 673 K (trace A) and 603 K (trace B) using a flow of 50 Nml/min H2, a heating rate of 2 K/min, and a catalyst weight of 140 mg. The spectrum in the poisoned state was obtained by using unpurified synthesis gas prior to dosing N2 at 603 K (trace C).

that atomic nitrogen segregating from the bulk to the surface at higher temperatures gives rise to the observed tailing. However, a lower NH3 concentration after dosing at 603 K should have been observed at high temperature due to the lower bulk solubiltity of atomic nitrogen at 603 K compared with 673 K. This is not the case in agreement with the computer-aided modelling results presented in ref. [9] which identified the tail as a consequence of the assumed Langmuir-Hinshelwood (LH) kinetics. Poisoning the catalyst prior to switching to N2 resulted in trace C displayed in the lower half of fig.5. The onset of NH3 formation is found to be shifted to higher temperatures by about 10 K, and the peak maximum concentration is is reduced by about 20 ppm. The traces of the active and the poisoned state coincide from about 500 K on. The peak structure observed after poisoning resembles the TPSR traces obtained with potassium-free catalysts shown in ref. [9]. The latter were found to consist of a sharp onset of NH3 formation at about 400 K followed by a broad maximum and the tail to higher temperatures without any further structures. In

119 agreement with the N2 TPD results, oxygen poisoning affects the potassium-promoted sites to a large extent leading to an essentially unpromoted state of the catalyst. The difference in the amount of NH3 formed due to poisoning is obtained by subtracting trace C from B yielding 18.7 ~,mol/g in good agreement with the decrease in N-. of 14.2 t, mol/g obtained by N2 TPD and the amount of 20 ~mol/g H20 detected during regeneration. This quantitative agreement demonstrates the reproducibility of the poisoning experiment. It is noteworthy that the effluent NHa concentration was decreased by 75 %, whereas the active metal area was found to decrease by only 20 %. This comparison suggests that also next-nearest neighbour sites are affected by the presence of O-..

TPSR 120

"--------

"- 100 o 80

NH3?

poisonedstate 20%H2 in He

'~

60

"~ 40 20 0 ,_.

300

400

500

600

700

800

Temperature / K

Figure 6. TPSR data obtained in the poisoned state using 50 Nml/min 20% H2 in He and a catalyst weight of 400 mg. The heating rate was 2 K/rain.

The close interrelation between the two ways of removing of N - . either as NH3 in H2 or as desorbing N2 in He is illustrated in fig. 6 which reproduces the result of a TPSR experiment using only 20% H2 in He. The computer modelling of the TPSR experiment predicted that a lowered partial pressure of H2 should increase the fraction of N - , desorbing as N2 [9]. Indeed, the formation of N2 is observed starting at 600 K accompanied with the decreasing concentration of NH3 approaching 0 ppm at higher temperatures. These observations can be interpreted on the basis of the LH mechanism which predicts a lowered coverage of H - . due to the lowered partial pressure of H2 which in turn decreases the rate of NH3 formation. Hence when reaching the onset temperature of N2 desorption, residual N - , is still present on the surface being observed as desorbing N2. The N2 TPD peak shape agrees with the results of the coverage-dependent modelling presented in ref. [ 10]. It is dominated by readsorption of the desorbing N2 molecules within the catalyst bed due to the high number of empty sites from which N-, has already been removed as NH3.

120

4. Conclusions Oxygen poisoning at 603 K achieved by cofeeding of about 5 ppm of oxygenic compounds was found to be reversible. The decrease of the effluent NH3 concentration by 75 % compared with t9o ~ 0.2 indicates additional poisoning of the next-nearest neighbour sites by O-.. In the poisoned state, a broad symmetric N2 TPD peak shifted to higher temperatures was observed in agreement with N2 TPD data obtained with potassium-free catalysts. From TPD experiments with various heating rates, an activation energy of N2 desorption of 172 kJ/mol was derived compared with 146 kJ/mol for the active state. The TPSR data in the poisoned state resembled those of potassium-free catalysts. Both observations can be rationalized by assuming that oxygen poisoning affects mainly potassium-promoted sites. Thus, the influence of potassium on the rate of N2 desorption and on the TPSR kinetics is removed to a large extent transforming the catalyst into an essentially unpromoted state. On the atomic scale, the K+O coadsorbate layer on the active catalyst is obviously in a substoichiometric oxygen-deficient state which is becoming oxygen-saturated and consequently inactive by cofeeding oxygen-containing compounds. N-. TPSR experiments were performed subsequent to N2 adsorption at 673 K and 603 K temperatures in pure H2 and with 20 %H2 in He. The unchanged tailing at higher temperatures when dosing N2 at 603 K and the observation of desorbing N2 when using diluted H2 were identified as consequences of the assumed Langmuir-Hinshelwood kinetics.

Acknowledgement The authors benefited from discussions with B. Fastrup, O. Hinrichsen and G. Ertl, and are grateful to Haldor TopsCe for supplying the iron catalyst. REFERENCES 1. EE. HCjlund Nielsen, in: Catalytic Ammonia Synthesis, 1st ed., ed. J.R. Jennings (Plenum Press, New York, 1991) p. 285. 2. A. Nielsen, J. Kjaer and B. Hansen, J. Catal. 3 (1964) 68. 3. J.A. Almqvist and C.A. Black, J. Am. Chem. Soc. 48 (1926) 2814. 4. B. Fastrup and H.N. Nielsen, Catal. Lett. 14 (1992) 233. 5. A.T. Larson and R.S. Tour, Chem. Met. Eng. 26 (1922) 647. 6. K.C. Waugh, D. Butler and B.E. Hayden, Catal. Lett. 24 (1994) 197. 7. G. Ertl, in: Catalytic Ammonia Synthesis, 1st ed., ed. J.R. Jennings (Plenum Press, New York, 1991)p. 109. 8. P. Stoltze, Physica Scripta 36 (1987) 824. 9. B. Fastrup, M. Muhler, H.N. Nielsen and L.E Nielsen, J. Catal. 142 (1993) 135. 10. M. Muhler, E Rosowski and G. Ertl, Catal. Lett. 24 (1994) 317. 11. B. Fastrup, J. Catal. 150 (1994) 345. 12. B. Fastrup, Top. Catal. 1 (1994) 273. 13. T.Z. Srnak, J.A. Dumesic, B.S. Clausen, E. T6rnqvist and N.-Y. Tops~e, J. Catal. 135 (1992) 246. 14. E Bozso, G. Ertl, M. Grunze and M. Weiss, J. Catal. 49 (1977) 18. 15. Z. Paal, G. Ertl and S.B. Lee, Appl. Surf. Sci. 8 (1981) 231.

91997 Elsevier Science B. V. All rights reserved. Dynamics of Surfaces and Reaction Kinetics in Heterogeneous Catalysis G.F. Froment and K.C. Waugh, editors

121

Importance of Dynamics in Real Catalyst Systems H. Topsoe 1, C. V. Ovesen 1 B.S Clausen ~ N.-Y. Topsoe ~ P.E. Hojlund Nielsen ~, E. Tornqvist 1, J.K. Norskov 2 1 Haldor Topsoe Research Laboratories DK-2800 Lyngby, Denmark 2 Department of Physics, Technical University of Denmark DK-2800 Lyngby, Denmark

1. Introduction In the kinetic modelling of catalytic reactions, one typically takes into account the presence of many different surface species and many reaction steps. Their relative importance will depend on reaction conditions (conversion, temperature, pressure, etc.) and as a result, it is generally desirable to introduce "complete" kinetic fundamental descriptions using, for example, the microkinetic treatment [1]. In many cases, such models can be based on detailed molecular information about the elementary steps obtained from, for example, surface science or in situ studies. Such kinetic models may be used as an important tool in catalyst and process development. In recent years, this field has attracted much attention and, for example, we have in our laboratories found the microkinetic treatment very useful for modelling such reactions as ammonia synthesis [2-4], water gas shift and methanol synthesis [5,6,7,8], methane decomposition [9], CO methanation [10,11], and SCR deNOx [12,13]. It is important to realize that in the kinetic treatments of real catalyst systems, it has almost always been assumed that the total number of sites (or the total surface area) remains constant (see, e.g., [ 14]) and that the only changes taking place are in the relative surface coverages by the different surface species. The above assumptions may not always be valid since as the local conditions inside the catalyst bed change, the surface structure (and number of sites) may also change. In fact, many surface science experiments (see, e.g., [15-18]) and experiments on real catalyst systems [19] have shown that such effects can be very important. However, in view of the lack of quantitative in

122 situ information, such complications have until recently not been introduced into microkinetic models and have been assumed neglected. In the present paper, we will discuss in some detail the results of methanol synthesis catalyst since in this case, the dynamic changes occurring in the catalyst structure have been described in some detail and it has been possible to use this insight to formulate a dynamic microkinetic model. The article is to a large extent based on the results published in recent articles [20,21]. However, some previously unpublished infrared, transient and industrial studies will also be presented since these throw further light on the importance of the origin of the dynamic aspects.

2. I n situ E v i d e n c e f o r D y n a m i c R e s t r u c t u r i n g d u r i n g M e t h a n o l S y n t h e s i s

Below we will discuss in situ EXAFS and Fourier transformed infrared (FTIR) spectroscopy results for Cu/SiO2 and Cu/ZnO catalysts which show that the surface and bulk structures may change with reaction conditions. It is furthermore seen that these changes depend very much on the nature of the support and the changes are also to a large extent reversible.

2.1 EXAFS

Figures l a and l b show the variation in the average nearest-neighbor coordination number determined from the Fourier transformed EXAFS above the Cu K-edge of Cu/SiO2 and Cu/ZnO catalysts, respectively [20,22]. The coordination numbers are shown as a function of the changes in the gas composition. The observed initial increase in the overall coordination number is likely to be due to some sintering of the small metallic Cu particles during the different treatments. The changes in the gas composition have essentially no effect on the coordination number of the small Cu particles of the Cu/SiO2 catalyst (Figure l a). Within the scatter of the data, the numbers are identical indicating that the gas composition has no significant influence on the structure and dispersion of the catalyst.

In contrast to the behavior of the Cu/SiO2 catalyst, large changes in the magnitude of the

123 coordination numbers for the Cu/ZnO catalyst (Figure l b) are observed with changes in the gas composition. This suggests significant structural and morphological changes of the Cu metal particles. Initially, the Cu particles are very small and some sintering takes place. However, of more interest, it is seen that in the CO/CO2/H2 methanol synthesis gas mixtures with the highest reduction potential, the coordination numbers are always relatively low. Upon changing to more oxidizing atmospheres, like the H20/CO/CO2/H2 gas mixture, the average coordination number increases, but upon returning to the dry methanol synthesis gas, the coordination number drops again. As seen from Figure l b, these changes are reversible.

The EXAFS results for the Cu/ZnO catalysts which show an essentially reversible change in the Cu-Cu coordination number upon changing the oxidation potential of the synthesis gas may, as discussed below, be interpreted in terms of a change in the particle dispersion due to a wetting/non-wetting phenomenon of the small Cu particles on the ZnO support. An increase in the oxidation potential ("wet conditions") gives rise to non-wetting, i.e. the Cu particles become more spherical in shape giving relatively high coordination numbers. The subsequent lowering of the oxidizing potential of the reaction mixtures ("dry conditions") results in increased wetting of the support, i.e. the Cu particles obtain a more disc-like shape particle. This particle will have relatively many low coordinated surface atoms and the average coordination number will be lower.

2.2 FTIR Below we will present some infrared results on Cu/ZnO and Cu/SiO2 catalysts [23] which also demonstrate that reversible changes take place in the Cu/ZnO catalysts and which throw some light on the nature of the surface changes occurring. It should be pointed out that these experiments were carried out prior to the EXAFS results on catalysts with a lower Cu loading (1%). Thus, only the qualitative nature of the changes will be addressed here.

A CO absorption band at 2085 cm 1 is observed upon CO adsorption at room temperature on the Cu/ZnO sample following reduction at 220 ~ in a gas mixture consisting of 0.5% CO, 4% CO2, 4% H2 (balance Ar). This CO absorption band is seen to shilt upwards to 2102 cm-I alter

124

10.5 I'

I

'

i

'

t

'

1

'

'

I

i

'

I

a

[ syn.gas syn.gas+H~

9.5

10%002/H2 . ~ ..... ._-~,~.~

6c

i , D " -

c 0 c 0 0

syn.gas

........ ~ .... '~ syn.gas+H20

V

syn.gas

syn.gas

8.5

0

<

7.5 5%Cu/Si02 6.5

l

1

,

!

2

,

t

3

~

1

,

4

Exp. no.

I

5

,

!

6

,

I

7

125

10.5

I

'

I

'

1

'

'I

'

1

'

I

'

b syn.gas+H20

9.5

6 c-

syn.gas+H20

co

om...

~

co o

8.5

" 10%CO2/H2

o 6. c~ <

syn. gas s

"

n.gas

syn.

7.5

5%Cu/ZnO

9 syn.gas

6.5

l

1

,

1

2

,

I

3

,

1

,

4

1

5

,

1

6

,

I

7

Exp. no. Figure 1

Variation in the apparent Cu-Cu nearest neighbor coordination number with changes in the gaseous environment, a) 5%Cu/SiO2 catalyst [22], b) 5%Cu/ZnO catalyst [20].

126 the sample was treated at 220 ~ in a methanol synthesis gas mixture consisting of 4.6% CO, 4,7% CO2, 87.6% H2 (balance Ar). According to the literature [24], the CO band at 2085 cm1 is due to CO adsorbed on low index Cu(111) surfaces, whereas the higher frequency CO band at 2102

c m "1

is attributed to CO adsorbed on high index Cu(ll0) surfaces. Thus, these

observations also suggest that the surface morphology of Cu on ZnO changes after exposure to environments with varying reduction potential with respect to the occurrence of changes in the structure. However, from the FTIR results we get some indication of the surface changes taking place. The observation that the disc-like particles produced after exposure to the most reduced reaction conditions expose predominantly (110) planes could suggest that the Cu particles will also expose (110) planes of the Cu-ZnO interphase. This is also in accord with the findings of Campbell for surface science model system [25].

Significant changes in sample transparency in the IR have been observed when the Cu/ZnO sample was exposed to different conditions [23]. A large decrease in IR throughput (sample transparency) of ZnO is observed as the sample undergoes a reduction treatment. Similar phenomena were also observed by Taylor and Amberg [26] and later also by Boccuzzi et al. [27]. Without going into details, the loss of transmission upon reduction can be explained in terms of electronic effects in that the reduction leads to an increase in defect structure in ZnO due to the formation of oxygen vacancies. The IR results on the 1% Cu/SiO2 sample show that, in contrast to the Cu/ZnO catalyst, no change of IR throughput occur, regardless of the reducing potential of the gaseous surrounding.

The above results are very important since they clearly show that oxygen vacancies can be formed in the ZnO under methanol synthesis conditions. This supports the suggestion presented in reference [20,21] that the morphological changes are caused by changes in the vacancy concentration at the interface between Cu and ZnO. This will be discussed further in section 3.2.

127

3. Catalytic significance of dynamic morphology changes during methanol synthesis

In the following section, we will discuss the kinetic implications of the dynamical changes in catalyst morphology during methanol synthesis. First, we will present an analysis of steady state kinetic experiments using a static, microkinetic model where it is assumed that the number of sites are constant. Then, we will introduce the dynamic aspect into the microkinetic modeling and also discuss some recent transient experiments.

3.1 Analysis of steady state results using a static microkinetic model A microkinetic analysis of the methanol synthesis reaction has recently been performed [7,8]. The starting point for the analysis is that the active phase for methanol synthesis is metallic copper. The following mechanism based on primarily surface science studies over Cu single crystals was proposed:

H20(g)+*

~

H20*

(1)

*

+"+

OH* + H*

(2)

2OH*

~-~

H20* + O*

(3)

~

O*+H*

(4)

~-~

H2(g)+2*

(S)

CO(g)+*

~-~

CO*

(6)

CO*+O*

~

C02"+*

(7)

*-+

CO2(g)+*

(8)

H20*

+

OH*+* 2H*

CO2" CO2" + H * HCO0* +H*

*-'+ HCOO* + *

(9)

~-~

H2COO* + *

(10)

H2COO* + H*

~

HsCO* + O*

(11)

H3CO* + H*

~

CHgOH* + *

(12)

CHsOH*

~

CHsOH(g)+*

(13)

where * represents an empty site and X* an adsorbed specie. In this mechanism, methanol is synthesized from CO2 in accordance with isotope labeling experiments [28,29]. Formate is a key intermediate in the synthesis of methanol. The water gas shift reaction takes place through a surface redox mechanism involving reaction steps (1)-(8) [5,6].

8E-6

Ix3

O

OO

7E-6 O

6E-6

~

x2 ~Z 9

5E-6 4E-6 3E-6

O 9

A A

2E-6

9 zx

CO/CO2-5.7 CO/CO2=2.7

1E-6 r~

9

0E0 0E0

1E-6

2E-6

3E-6

4E-6

5E-6

6E-6

7E-6

8E-6

calculated rate of MeOH synthesis (mol/sec/g cat)

Figure 2 Comparison of the calculated rate with the measured rate of methanol synthesis over a Cu/ZnO/AI203 catalyst. The calculated rate is obtained from the static microkinetic model. Inlet gas compositions: 12% CO, 2.1% CO2, 85.9% H2 (solid circle), 17.9% CO, 6.7% CO2, 75.4% H2 (empty triangle) [2 l].

129 The Langmuir Hinshelwood kinetic model based on this reaction scheme is formulated assuming that all reactions are in equilibrium except for reaction steps (2), (4), (7) and (11). Reaction steps (2), (4) and (7) are all steps which may be slow during the shift reaction [5,6], whereas reaction step (11) represents the slow step for methanol synthesis [8]. The kinetic and thermodynamic parameters are taken from available Cu single crystal experiments. We call this type of model a static microkinetic model since the number of active sites are assumed constant (i.e., independent of reaction conditions) [21 ].

Figure 2 shows the agreement between this model and measured rates of methanol synthesis over a Cu/ZnO/A1203 at ambient pressure and temperature by Graaf et al. [30,31]. It is seen that overall the microkinetic model is able to explain the magnitude of the rate quite well, but the data are grouped into two families. A sensitivity analysis of the microkinetic model showed that the data and the model could not be brought in agreement unless the parameters were severely changed which would bring them in disagreement with the surface chemistry of the reaction.

3.2 Dynamic Microkinetic Model It is noted that the two families of points in Figure 2 are characterized by having different reduction potential. In section 2, we discussed the EXAPS and FTIR results which showed that such changes in gas composition lead to structural changes. In view of this, it is likely that the difficulty in obtaining agreement using the static model is due to the fact that the assumption of constant number of active sites (and constant bulk and surface structure) used in the static model is not valid. Below we will examine this by modifying the static model to include the dynamic aspect. However before presenting the dynamic model we will give some further catalytic evidence that exposure to different gas compositions changes the catalyst behavior.

Figure 3 shows experiments that have been carried out over a methanol catalyst in a flow reactor operating at industrial conditions [32]. The measurements are carried out at constant outlet conditions thus, in each measurement the flow is varied to obtain the same outlet

130 1.5

~

> ~ 0 0 t-t~

1.0

t"

E

"O N

E

0.5

O

z

~:1

0.0

i

0

~

i

1 O0

,

i

200

No 002

,

i

~

300

I

400

,

.

i

500

Run'time (h)

Figure 3 The measured rate of methanol synthesis over a Cu/ZnO/Al203 catalyst against time on stream. The data is obtained in a plug flow reactor operating under steady state conditions at 494 K and 63 atm total pressure and exposed to the feed gas: 5% CO, 5% CO2, 88% H2, 2% N2. After 168 run hours the CO2 is removed from the gas for 4 hours [32].

131 concentrations. In this way the methanol synthesis rate is measured directly. After the catalyst has been reduced it is exposed to the synthesis gas. Initially, the catalyst shows some sintering and the measured rate decreases. When the catalyst is stable the CO2 is removed from the gas for a while. No synthesis of methanol is observed showing that CO2 is necessary for the synthesis of methanol. When the catalyst is later exposed to the CO2 containing gas an increase in the rate of approximately 50% compared to before the CO2 was removed is observed. The rate slowly decreases to the same level as before the exposure to the CO2 free gas. This experiment shows that the exposure to more reducing gas has changed the catalytic properties of the catalyst. These changes are reversible since the catalyst returns to its initial stage.

The above findings also show that it is necessary to include the dynamic aspect in a microkinetic model to be able to account quantitatively for the rate of methanol synthesis. A dynamic microkinetic model of the methanol synthesis was recently presented [21]. To introduce the dynamic aspect into the microkinetic model, it is necessary to have a description of the changes in particle shape, i.e. surface area, with change in the gaseous environment. As the change in particle shape is induced by change in contact surface free energy between the Cu particle and the support, we will first present results for such calculations.

The description of the change in particle morphology is based on the Wulff constructed particles [20,21 ]. This gives the surface area and distribution of surface planes for a particle for changes in the contact surface free energy y/y0 between particle and support. This is shown in Figure 4a where the Cu particle is supported with the (110) facet against the substrate. For positive values of y/yo it is seen that the (111) facet is the dominating facet, whereas for large, negative values of y/y0, the (110) facet dominates. As mentioned in section 2, the change in contact surface free energy y/y0 is related to the changes in the number of oxygen vacancies at the interphase between the Cu particle and the ZnO support [21]. Thus, the following two equations are important in determining the number of oxygen vacancies at the interphase: H2(g) + Zn -O- Cu ~-~ H20(g) + Zn -r-I-Cu

(14)

CO(g) + Zn -O- Cu ~ CO2(g) + Zn -IZI- Cu

(15)

132

iim,,--===--~,~--,,,==,--,,mmm

mm,==~,==

.... . . . . .

....

100

....

100

110

.....

110

111

....

111

totol

totol

r'3 r

-~..

,,

9

-~.

..9.

~-.~

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

.~

-9"

...-.

. - - - - ~

_

, .....

. . . .

.,, /

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

l .0

I

i

I -0.6

i

I

i

I

i

I

I

-0.2

I 0.2

i

I

i

I

i

I

,

0.6

1.0

0

y/to

1000

2000

reduction

3000

4000

potential

5000

6000

7000

8000

9000

(PH2Pco/PH2oPc02)

Figure 4 The dimension-less surface area A/V ~

of the different facets as well as the total area

calculated from the Wulff construction assuming that

the (110) plane of the particle is

attached to the substrate. A: The surface area is shown as a function of the contact surface free energy ?/y0 between particle and substrate. B: The surface area is shown as a function of the reduction potential of the gas phase [21]. The relation between contact surface free energy and reduction potential is given by equation (16).

133 where-I"1- is an oxygen vacancy in the Zn-O-Cu interphase. These equations are the basis for relating the contact surface free energy to the gaseous environment. As shown in [21] the following relation between reduction potential and the relative surface contact free energy can be derived:

1-

?'/70

=

~/ PH:Pco K1 K2 PH:o Pco----~, I

1+

PH:Pc~

(16)

K~ K,_ PH:o Pco:

where K~K2, is the product of the equilibrium constant of reactions (15) and (16). K1K2 is expressed as a function of the free energy of reaction (14) and (15) is given by: K1K2 = e x p -

(17)

A procedure for estimating AG is given in [21]. Figure 4b shows the surface area of a Cu particle as a function of the reduction potential using the value of zXG = -2 kJ/mol reported in [21] for the (110) interface. It is seen that large changes in the surface area takes place and especially abundance of the (110) facet.

In the dynamic, microkinetic model, it is furthermore taken into account that the rate of methanol synthesis is different over the three low index facets [21]. The observed rate of the methanol synthesis for a given catalyst is therefore an average of the rates over the exposed facets and can be expressed as r.,bs = (rl rl00 + s rll0 + (1 - r 1 - s) rill) 9N

(18)

where rl is the ratio of the number of sites on the (111) plane relative to the total number of sites N; ~ is the ratio of the number of sites on the (110) plane relative to the total number of sites; and r~ is the specific rate over a surface site in plane i. The distribution of surface facets is determined from Figure 4.

The result of using this dynamic microkinetic model to the data by Graaf et al. [31, 32] is shown in Figure 5. It is seen that the separation of the data into two groups is no longer

8E-6 o e~t) r

O

7E-6 6E-6

~ ~,,,I

5E-6 ~Z 9 9

4E-6 3E-6

9

2E-6

9 zx

1E-6 0E0 0E0

1E-6

2E-6

3E-6

4E-6

5E-6

calculated rate of MeOH synthesis (mol/sec/g

Figure 5 Comparison of the calculated rate with the measured rate of methanol Cu/ZnO/Al203 catalyst. The calculated rate is obtained fiom the dynamic Inlet gas compositions: 12% CO, 2.1% CO2, 85.9% H2 (solid circle), 75.4% H2 (empty triangle) [21 ].

135 present and a much better description of the data is obtained. For the most reducing gas composition, the catalyst exposes a larger fraction of the (110) facet. This causes a large increase in the production of methanol since the rate over this surface plane is higher than over the two other low index surface planes. Thus, the structure sensitivity of the reaction is also taken into account.

3.3 Transient experiment Further indication of changes in the number of active sites with changes in gaseous environment are obtained from transient studies of the methanol synthesis reaction over a Cu/ZnO/Al203 catalyst and a CH/AI203 catalyst [33]. Figure 6 shows the measured methanol trace from a flow reactor operating at atmospheric pressure in which the catalysts first was exposed to a very reducing gas (5% CO in H2) and then to a less reducing gas (5% CO, 5% COa in H2). Results for both Cu/ZnO/Al203 (dashed line) and Cu/Al203 (solid line) are shown. For the Cu/ZnO/Al203 catalyst (dashed line), it is seen that after the change to the less reducing gas there is an initial increase in the rate of methanol synthesis which after a while slowly decays to the steady state value. Using the static microkinetic model it has not been possible to explain this maximum in production of methanol and the slow decay to steady state (for example by changes in the amount of adsorbed species). However, in view of the in situ observations (section 2) and the dynamic, microkinetic model, the catalyst will have a high surface area (and a high number of (110) active sites) after the exposure to the very reducing 5% CO in HE gas mixture. Therefore, there will initially be a high production rate of methanol when the catalyst is exposed to the CO2 containing gas. However, as the Cu particles equilibrates to the less reducing gas the surface area will decrease causing a decrease in the rate of methanol production. For the Cu/Al203 catalyst (solid line), a quite different transient behavior is observed. This catalyst is seen to adjust to the new conditions without going through a maximum. The transient behavior of this catalyst is possible to explain by the static microkinetic model. The above shows that the dynamic changes are unique for the ZnO supported systems.

136

Addition of CO 2 to the stream o~

v

E 03

c"'

1 0 -1

o..,, t""

1

,

O o_

",. . . . .

I

'air I

lilm i

4ml l

!

I

I

I

I

(am'I

Ilmimdill i

I

I

,aim i

ill, i

I . .

Q lim, imm m I

,me,t

m i

#

4ms,.

Cu/ZnO/AI203

m m

C

m s

0

Cu/Ai203

'

0 0

0 C C~

1 0 -2

,

0

,

t

,

10

I

~

20

Time

. . . . .

I

30

. . . . . .

,,

,

40

(rain)

Figure 6 Methanol concentration at the reactor outlet for two different catalysts, Cu/ZnO/Al203 and Cu/A1203 loaded in the reactor at 220 ~

First, the inlet concentration of the reactor is 5~ CO

and 95% 1-12 and then (at the arrow) CO2 is added to the stream given a concentration of 5% CO, 5% CO2 and 90% H2 at the inlet [33].

137 Conclusion and Outlook

The EXAFS and FTIR investigations of Cu/ZnO methanol synthesis catalyst show that depending on the reduction potential of the reaction gas the surface and bulk morphology of the Cu particle change. Microkinetic analysis of the steady state and transient kinetic experiments show that the morphological changes have a strong impact on the activity behavior of the catalysts. As a result of such changes, one cannot assume in the kinetic modelling of the methanol synthesis reaction that the structure and the total number of sites remain constant. It was seen that it is possible to modify the static microkinetic model based on surface science results and the measured structure sensitivity and adjust this model to take into account the structural changes. Such a model has been termed a dynamic microkinetic model. The dynamic model also explains nicely the trends observed in transient experiments. The ZnO support is seen to play an essential role and the Cu-ZnO interphase energy determines to what extent spreading and morphological changes occur. The structural and kinetic transient effects observed for the ZnO containing catalysts are very different from those observed for other supported Cu catalysts.

In recent years, phenomena such as surface reconstructions, surface segregations, etc. have been investigated extensively both experimentally and theoretically, and a much better quantitative understanding is emerging. Consequently, it is very likely that dynamic, microkinetic models can also be developed to take into account such phenomena and several model systems are presently being investigated.

We have also seen that the combined QEXAFS/XRD method [34] has some unique advantages to provide the in situ structural insight. The recent improvements in time resolutions by several orders of magnitude will allow the investigation of new types of dynamic phenomena.

138 References

1. Dumesic, J.A., Rudd, D.F., Aparicio, L.M., Rekoske, J.E., and Trevifio, A.A, "The microkinetics of Heterogeneous Catalysis", American Chemical Society, Washington DC, (1993) 2. Stoltze, P., and Norskov, J.K., Phys. Rev. Lett., 55, 2502 (1985) 3. Stoltze, P., and Norskov, J.K., J. Catal., 110, 1 (1988) 4. Stoltze, P., and Norskov, J.K., Topics in Catal., 1, 253 (1994) 5. Ovesen, C.V., Stoltze, P., and Norskov, J.K., and Campbell, C.T., J. Catal, 134, 445 (1992) Ovesen, C.V., Clausen, B.S., Hammershoi, B.S., Steffensen, G., Askgaard, T., Chorkendorff, I., Norskov, J.K., Rasmussen, P.B., Stoltze, P., and Taylor, P., J. Catal., 158, 170 (1996) Rasmussen, P.B., Holmblad, P.M., Askgaard, T.S., Ovesen, C.V., Stoltze, P., and N~rskov, J.K., and Chorkendorff, F., Catal. Lett., 26, 373 (1994) 8. Askgaard, T.S., Norskov, J.K., Ovesen, D.V., and Stoltze, P., J. Catal., 156, 229 (1995) 9. Alstrup, I. and Tavares, M.T., J. Catal., 135, 147 (1992) 10. Alstrup, I., J. Catal., 151, 216 (1995) 11. Tavares, M.T., Alstrup, I., Bernardo, C.A., Rostrup-Nielsen, J.R., J. Catal., 158 402 (1996) 12. Dumesic, J.A., Topsoe, N.-Y., Slabiak, T., Morsing, P., Clausen, B.S., T6rnqvist, E., and Topsoe, H., New Frontiers in Catalysis, "Proceedings of the 10th International Congress on Catalysis, Budapest" (L. Guczi, F. Solymosi and P. Teteny, Eds.), Akad6miai Kiado, Budapest, 1993, p. 1325 13. Dumesic, J.A., Topsoe, N.-Y., Topsoe, H., Chen, Y. and Slabiak, T., J. Catal., 163, 409 (1996) 14. Boudart, M. and Djega-Mariadasson, G., "Kinetics of Heterogeneous Catalytic Reactions", Princeton Univ. Press, Princeton, 1984. 15. Brill, R., Richter, E.L., Ruck, E., Angew. Chem., lnt. Ed. Engl., 6 882 (1967) 16. Ertl, G., Topics in Catal., !, 305 (1994) 17. Samorjai, G.A., "Introduction to Surface Chemistry and Catalysis", Wiley, New York, 1994

139 18. Besenhacher, F., Spmnger, P.T., Ruan, L., Olesen, L., Stensgaard, I., and La~gsgaard E., Topics in Catal., 1, 325 (1994) 19. Dumesic, J.A., Topsoe, H, Bourdart, M., J. Catal., 37 513 (1975) 20. Clausen, B.S., Schiotz, J., Gr~b~ek, L., Ovesen, C.V., Jacobsen, K.W., Norskov, J.K., and Topsoe H., Topics in Catal., 1, 367 (1994) 21. Ovesen, C.V., Clausen, B.S., Schiotz, J., Stoltze, P., Topsoe, H., and Norskov, J.K., J. Catal., in press 22. Clasen, B.S., unpublished results 23. Topsoe, N.-Y., to be published 24. Pritchard, J., Catterick, T., and Gupta, R.K., Surface Science, 53, 1 (1975) 25. Ludviksson, A., Zhang, R., Campbell, C.T., Griffiths, K., Surf. Sci., .313, 64 (1994 26. Taylor, J.H., and Amberg, C.H., Can. J. Chem. 39, 535 (1961) 27. Boccuzzi, F., Ghiotti, G., and Chiorino, A., Surface Science, 183, L285 (1987) 28. Rozovskii, A. Ya., Kin. Kat., 21, 97 (1980) 29. Chichen, G.C., Denny, P.J., Parker, D.G., Spencer, M.S., and Whan, D.A., Appl. Catal., 30, 333 (1987) 30. Graaf, G.H., Stamhuis, E.J. and Beenackers, A.A.C.M., Chem. Ing. Sci., 43, 3185 (1988) 31. Graaf, G.H., Ph.D. thesis, Rijksuniversiteit, Groningen, 1988 32. Hojlund Nielsen, P.E., to be published 33. Tornqvist, E., to be published 34. Clausen, B.S., Gr~b~ek, L., Steffensen, G., Hansen, P.L., and Tops~e H., Catal. Lett., 20, 23 (1993)

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91997 Elsevier Science B. V. All rights reserved. Dynamics of Surfaces and Reaction Kinetics in Heterogeneous Catalysis G.F. Froment and K.C. Waugh, editors

141

Reaction kinetics as a basis for optimal transient operation of catalytic reactors

Yurii Sh. Matros, Grigorii A. Bunimovich and Vadim O. Strots Matros Technologies, 14963 Green Circle Dr., St. Louis, Missouri 63017, USA

1. INTRODUCTION Optimal design and operation of chemical plants is ultimately dictated by economical criteria: the lowest cost of making a desirable product or of destroying harmful pollutants, given certain constrains related to process safety, waste minimization, etc. This optimization problem involves many parameters (cost of raw materials, equipment manufacturing and erection, labor, maintenance, utilities, waste disposal, etc.) that are interesting to analyze but will generally be left out of the scope of this paper. Instead, we will focus on optimal organization of processes taking place within a catalyst bed. Usually, the problem is to find a combination of temperature, pressure and composition that provides for maximum use of the catalyst ability to accelerate desirable reaction and is within given constrains. This combination defines a theoretical optimal regime, regardless whether it can be realized in a reactor. Such a regime is commonly sought using mathematical model of the process. Solution of model equations requires knowledge of the reaction kinetics expressed as rate equations for all stages involved. Generally, steady-state operation is assumed which must be adhered to during the process operation. This, traditional approach to optimization of continuously operated reactors usually treats transient processes (catalyst deactivation, runaway phenomena, etc.) as undesirable. If those cannot be avoided, processes are designed to or minimize their effects. For example, catalytic cracking or dehydrogenation of alkanes is designed as transient, including stages of reaction and regeneration, because of quick and unavoidable catalyst coking. However, there is substantial evidence that external forcing of process parameters can improve the reactor performance over the steady-state optimum. In this paper, we will focus on forced unsteady-state operation (FUSO) applied to continuous processes with nearly constant catalyst activity. In particular, we are interested in the effect of dynamic processes occurring on the catalyst surface on the FUSO, and in methods of FUSO optimization. Examples of successful practical application of FUSO will be considered. The term 'catalyst surface' is used broadly here. The dynamic processes of interest may include surface restructuring, diffusion of species into the bulk of solid catalyst pellets, reactions in liquid phase supported in a porous catalyst pellets, etc.

142 2. FORCED UNSTEADY-STATE OPERATION Optimum reactor performance obtained under optimal steady-state conditions does not determine an absolute limit for a heterogeneous catalytic system. Reason for this statement lies in a fact that a set of the steady states is comprised within a broader set of possible unsteady or cyclic states of the dynamic system. Therefore, the maximum value of an object or performance function ( J " ~ ) obtained under optimal unsteady-state conditions may not be lower than that for optimal steady-state (J,'~), or J " ~ > J,"~. Two major factors contribute to the potential of FUSO [1 ]: 1. The dynamic properties of catalyst. Unsteady conditions in the fluid phase induce a catalyst state (composition and structure) changes which, under certain conditions, result in a 'resonance' in the catalyst behavior, expressed as its selectivity and/or activity increase. 2. The dynamic characteristics of an entire reactor system. Interaction of external forcing with heat and mass transfer may result in temperature and composition patterns in the reactor that are close to optimum and unattainable in steady state. A number of reactor schemes operating under forced unsteady-state conditions is known. An unsteady-state process in a fixed bed reactor can be created by forcing of the inlet composition or temperature. As a rule, a stepwise periodic control is preferable because it is simpler to perform and affects a system stronger than other forcing modes, such as sine wave. Widely applied is periodical flow reversal in a fixed bed reactor, organized by changing the flow direction with switching valves or by continuous rotation of the bed of monolithic catalyst relative to stationary direction of flow. Both methods create continuous migration of a temperature or adsorption zone along the catalyst bed. The direction of this migration is changed periodically. Reactor configurations are known that allow the zone to move in one direction through a system of catalyst beds [2]. In fluidized beds, favorable unsteady state of the catalyst can be obtained by catalyst circulation inside the reactor or between the fluidized bed and riser reactors. Another group of FUSO combines chemical reaction with separation of products. These methods can employ a reactor with circulating bed of catalyst [3,4] or by periodic changes of feed and product ports in a reactor with several fixed beds, known as simulated moving bed reactor [5-7]. Reaction and separation can include periodic pressure changes using the known separation technique of pressure-swing adsorption. 2.1. The Problem of the Research Approach

History of investigation of FUSO aimed at radical improvement of the performance of chemical processes is nearly 30 years long. Many experimental studies have demonstrated the performance improvement of various processes of heterogeneous catalysis (see reviews by Renken [8, 9], Silveston [10,11 ] Yadav and Rinker [12]), including SO2 oxidation on a vanadium catalyst, NH3 synthesis on a promoted iron, H2S oxidation by SO2 on a bauxite, Fisher-Tropch synthesis on ruthenium and cobalt catalysts, CO oxidation on various catalysts, methanol synthesis on a Cu/ZnO catalyst, etc. Typically, an isothermal fixed bed reactor was used with periodically oscillating inlet gas composition. Most valuable experimental strategy includes measurement of steady-state reaction rates in the entire region of possible inlet composition and comparison with the cycle averaged reaction rate at the same cycle average composition of the gas mixture.

143 However, there has been limited success in the development and commercial application of unsteady-state process operation. In part this was caused by lack of dramatic improvements worthy of commercial applications. It is extremely difficult to achieve such improvements because of tremendous inherent complexity of forced unsteady-state catalytic systems. Not completely understood by this moment, interplay of reaction kinetics, heat and mass transport phenomena, nature and parameters of concentration forcing, etc. causes this complexity. An example below shows how this complexity makes the commonly used trial-and-error approach to the design of FUSO fail so often. Partial oxidation of hydrocarbons on oxide catalysts occurs via a redox mechanism, including reduction of an active site by hydrocarbons and its oxidation by gas phase oxygen [13,14]: Activation &oxygen: 02 ~-~ 2 O, ~ OL ~ Ov Partial oxidation: R + OL ~ RO Complete oxidation: R + O, ~ COx; R + OL --~ COx This scheme includes three forms of active oxygen: weakly-bonded adsorbed oxygen, Os, lattice oxygen at or near the surface, Or, and lattice oxygen in the bulk of the catalyst, Ov. While both Os and Or participate in complete oxidation, only Or leads to partial oxidation. This mechanism makes it attractive to separate oxidation and reduction steps of catalytic cycle in space and/or time. The FUSO may include the steps of a) oxidation of the catalyst, b) removal of weakly-bonded oxygen form the surface, for example, by flushing the surface with inert gas, c) contacting a hydrocarbon with the catalyst for a highly-selective partial oxidation to occur, and d) flushing the surface before the next cycle begins. A researcher has to find a periodic forcing strategy to uncover the FUSO potential. Using simplest bang-bang forcing function leaves at least 11 parameters for optimization: four temperatures, two compositions, duration of each stage, and four residence times. With only four values of each parameter tested, the number of experiments one has to run is 4~1, or over four million. Apparently, it is unlikely to find optimal periodic regime experimentally with an affordable effort. In the best case, a researcher with good insight may find some local extremum which demonstrates the advantage of FUSO but typically is not convincing for industry to initiate serious and expensive efforts aimed to develop commercial processes. Therefore, a different approach for the process optimization is needed, based on detailed quantitative description of dynamic processes taking part in the catalytic reaction. Such quantitative dynamic models are difficult and expensive to develop, this is why they are rarely available. A simpler approach is reasonable that uses quantitative information on observable rates of some hypothetical processes on the surface. A dynamic kinetic model can be presented as follows [2]:

a0 u

r);

f = f ( c , O, or, r);

w =

q~ = q0(c, 0, or, T)

r);

at

o,

r),

o,

r)]

Here, the apparent rate of chemical transformation, W, is comprises functions f and q~that describe the catalytic cycle and effect of the reaction mixture on the catalyst properties,

144 respectively, and depend on gas phase composition, c, and temperature, T, concentrations of surface intermediates, 0, and on the composition and structure of a catalyst as a whole, ct. Development of dynamic kinetics models is most effective when transient kinetic experiments combined with physico-chemical methods of investigation of catalyst surface. There are examples of models that describe well the dynamic processes on the catalyst, such as studies by Balzhinimaev et. a/. [ 15], Sadhankar and Lynch [ 16], Jobson et. al., [ 17]. Once a model of dynamic processes on catalyst surface is devised, it can further be used for numerical optimization of the periodically forced reactor. Invariance of such model, where all equations and parameters are independent on time at every space scale such as pellet, catalyst bed or reactor at any time, simplifies the further scale up. 2.2. Analysis and optimization of cyclic processes

The first problem to be solved is to establish whether the optimal steady-state operation can be improved using forced unsteady-state conditions. If the answer is positive, then the optimal operation should be found. The last stage obviously is the cost comparison between the optimal unsteady-state operated system and the steady-state one. For the class of systems described by ordinary differential equations the general optimum periodic control problem includes: - equations for variables of state: :~ = f ( x , u ) , where x=[xl(t),x2(O ..... x,(O] is a vector of state variables, continuous time dependent functions; - definition for forcing control variables, u(t)= [u~(t),u2(t) ..... um(t)] , where u(O is a vector of piecewise continuous functions determined over the time interval [O,t c ]; - periodicity constrains for control functions, u(O)-u(tc); periodicity constrains for state variables x(O)= x(tc) are usually assumed too; t

- definition for objective function, J = l l f o ( X , u ) d t ;

wherefo(X,u) is an instantaneous

tr 0

process performance index; - constrains on instantaneous values of control variables 9u~'~' ~ uj(t)~ u ~ , 1=1 ..... L; or on their averaged values. For steady states, solution of the system is determined from the equation 0= f ( x , , u s) and the objective function is determined as: J, = fo(x,,us). A number of techniques was developed to solve the optimization problem. Some deal with two limiting cases of periodic operation: relaxed steady states obtained at high frequency of the forcing function and quasi-steady states with forcing period much longer than the system response time [18]. For the intermediate range of frequencies and low amplitudes, the most widely used method is the H-criterion developed by Guardabassi et. al., [19]. This method have been used in Refs. 20-22 for analysis of chemical reaction systems. Shape of the optimal control function can be found numerically using an algorithm by Horn and Lin [23]. In Refs. 24 and 25 this technique was extended to the simultaneous optimization of a forcing function shape and cycle period. There is a number of examples where the effect of periodical forcing has been elucidated. Simple systems with two parallel or consecutive reactions (A ~ B , A ~ C; A ~ B , B C; A ~ B, B ~ A; A ~ B , A + B ~ C) were found in Ref. 22 to give rise to a selectivity

145 enhancement due to high-frequency temperature oscillations if activation energy of a desirable reaction (A --~ B) is higher than that for a side process. In more complex model systems, selectivity increase can be expected if: there exist a competition between two gaseous species for empty active sites, and a relaxation time of a 'useful' intermediate concentration is less than that for that leading to a side product. The following example was analyzed by Zolotarskii and Matros in Ref. 26: A + [Z] ~ [AZ], B+[AZ] <-~ P+[Z], B+[Z] +-->[BZ] andA+[BZ] ~ R+[Z].

(1)

If concentration of the intermediate [BZ] responds slower than the concentration of [AZ], periodic forcing of concentration of the component A increases selectivity compared to the optimal steady-state. An intermediate frequency periodic control is preferred. Even larger selectivity results from simultaneous variation of both inlet concentrations A and B (Fig. 1).

E

0"21

5 II

I..

= 0.14

1

~ -o.o

-o.1

-0.2 _

o.1

w ~,.i

~ ~ ' ~ 1

1

~ w ~,~.!

lO Cycle,

sec

~

lOO

~ 3~m

I

lOOO

Figure 1. Effect of concentration forcing on cycle average selectivity and conversion for reaction scheme (1). Computer simulation results adapted from [26]. 1: steady-state, 2 and 3: cycle-average rates of product C and product D formations; 4: selectivity regarding to product D formation at periodic variation of product A concentration; 5: selectivity at simultaneous periodic variation of products A and B.

Ozgalsen et. al. [24] and Chen et. el. [25] conducted numerical optimization of ethylene oxidation on silver in a CSTR. In both papers, substantial improvement in ethylene oxide yield was found, but at different parameters of cyclic regimes. This difference results from the fact that numerical techniques often identify some local maxima. Unfortunately, these papers do not contain interpretation of the results from the reaction mechanism point of view. Truffer and Renken [27] found a positive effect in the reaction of ethylene addition to acetic acid that involves substantial reaction inhibition by an adsorbed product. The effect is achieved at high frequency oscillation of acetic acid concentration because the cycle average concentration of inhibiting intermediate product is lower than in steady-state. This facilitates ethylene adsorption and a favorable distribution between the two adsorbed species. Thullie and Renken [28] found that periodical interruption in the feed of reactant causing formation of non-reactive surface species increases the rate of catalytic reaction of dehydration of alcohols or deamination of primary amines on acid base catalyst. The process can be generally described by one of the two simple schemes:

146 (a): A + [X] ~-~ lAX], A +[Z] ~-~ [,4Z]; lAX] +[.4 Z] ~ B + C +[Z]+ [X], and (b): A +[Z] ~ [A Z], [A ZI +A ~ [A ZA ] and [A Z] --~ B + C +[Z] .

Optimal operation corresponds to a short total cycle while the fraction of time for a reactant A feed interruption depends on the inlet concentration of the reactant. Processes exhibiting multiplicity and/or self-sustained oscillations can be improved by imposing a forced perturbation of the inlet composition, as has been shown by Graham and Lynch [29] for the example of CO oxidation over Pt supported catalyst. The reaction exhibits two steady-states at intermediate CO coverage. Simultaneous oscillations of CO and oxygen inlets brings the system to an average catalyst state corresponding to the upper regime of reaction performance. Strots et. al. [30] used a realistic dynamic model of practically important process, SO2 oxidation, to optimize periodic regimes in a CSTR using stepwise composition forcing. Analysis of various operating modes showed that it is beneficial to change SO2 and O2 concentrations in the same phase, and at low concentration of SO3 in the active melt of catalyst. This conclusion leads to an improved periodical regime where simultaneous forcing of feed concentrations and flow rate, organized so that to provide for periodical flushing of SO3, resulted in over-equilibrium SO2 conversion, unattainable in any steady-state regime. It is practically important to develop optimization techniques for dynamic processes in distributed systems. In Ref. 31, OzgOlsen and (~inar discussed an approach for numerical simulation of tubular wall-cooled reactor with periodical forcing of reactant concentration and flow rate. They simulated CO oxidation with single steady-state rate equation that did not include any dynamic processes on the catalyst surface. CO conversion exceeding the steadystate value was obtained, but even for this simple kinetic model computation time demand is too high to implement numerical optimization algorithms. 3. COMMERCIAL APPLICATION OF FUSO 3.1. Partial oxidation in fluidized bed and riser reactor.

A practically important example of dynamically operated partial oxidation system is the riser reactor (Fig. 2) developed by DuPont [32] for n-butane oxidation to maleic anhydride over a vanadyl pyrophosphate catalyst. In the riser zone, small (--- 100 ban) spherical catalyst particles contact with hydrocarbon feed during a short time, about 10-30 s, while they are transported by the gas stream with a velocity about 0.5 m/sec. The catalyst is regenerated in the fluidized bed regeneration zone. Separation of oxygen and butane feeds allows substantial increase in the feed n-butane concentration, exceeding the explosion limit. Pilot tests demonstrated 80-85 % reaction selectivity, which is by 5 to 10 % higher than that achieved in a tubular reactor. This selectivity gain is attributed to suppression, in the riser reactor, of surface oxygen species which accelerate reaction steps leading to carbon oxides formation. Steam stripping of the catalyst after the reoxidation zone also improved selectivity. Riser technology appears to be quite versatile. Patience and Mills [33] investigated propylene oxidation into acrolein and found that this technique has a potential for the commercial scale production of acrolein. Their kinetic model was based on a simplified single site redox mechanism involving consecutive-parallel reactions for the partial and complete oxidation of propylene. Its predictions of the performance of the reactor gave correct trends,

147 but it was understood that accurate simulation and optimization of the reactor needs a different dynamic model, valid for process conditions actually encountered in a riser reactor.

2

2 Air

/nlet

1

Inert

2

gas

2

t

Hydrocarbon

Fig. 2. Riser reactor flow diagram (from [32]). 1: riser, 2: regenerator, 3: stripper, 4: catalyst.

Fig. 3. Flow diagram of the reactor with periodic flow reversal. 1: fixed bed of catalyst, 2: switching valve.

Particles of catalyst circulating in a fluidized catalyst bed also are in unsteady state. If the characteristic time of catalyst state changes is longer than the time scale for particle dispersion, the catalyst state will adjust to averaged parameters of the gas mixture in the reactor volume rather than to the local properties of the current location in the reactor. This situation can lead to improved selectivity of partial oxidation process. For instance, higher selectivity o-xylene oxidation to phthalic anhydride on V2Os/TiO2 at high conversion can be achieved if the catalyst state corresponds to low o-xylene conversion [34]. Therefore, stirring of catalyst in fluidized bed is advantageous, especially in a reactor with "organized" fluidized bed equipped with packing promoting mass transfer between low and high density phases of fluidized bed.

3. 2. Fixed bed reverse-flow reactor (RFR).

This technique uses periodic reversal of the flow direction of the reaction mixture through the fixed bed of catalyst (Fig. 3) which thereby serves both as an accelerator of chemical reactions and as a heat regenerator and/or accumulator [2], often flanked by layers of heat regenerative inert, typically ceramic, packing. Flow reversals induce continuous back-andforth migration of the heat and reaction waves through the catalyst bed (Fig. 4). This allows for continuous autothermal operation without or with minimum external heat input. Theoretical basis of RFR operation, experimental results and commercial applications were described in a number of papers [2,35].

148 Three commercial processes, complete oxidation of volatile organic compounds (VOC) for purification of industrial exhaust gases, SO2 oxidation for sulfuric acid production, and NOx reduction by ammonia, have employed the periodic flow reversal concept. In all cases the reaction kinetics, including dynamic phenomena plays a crucial role. RFRs for VOC oxidation (also referred to as regenerative catalytic oxidizers, or RCOs) are often designed assuming single irreversible exothermic reaction. Mechanisms of complete oxidation of organics are complex, particularly for oxidation of multicomponent mixtures [36,37]. However, simple first-order rate equations are useful for analysis and accurate design of RCOs for two primary reasons: a) they describe well the main reaction features such as exponential reaction rate increase with temperature, and b) most rate equations are reduced to this form at high temperatures and low VOC concentrations.

6oo t

~1764oo 3

t.--

'Co

200

8

1'

60 "~

g

40 8 20 0

0.0

02

04

06

08

Depth of bed, dimensionless

1.0

Fig. 4. Behavior of flow reversal reactor. 1-5: temperature profiles, 1'-5' " conversion profiles during a period between flow reversals. Example of exothermic reversible reaction.

Optimization of an RCO involves a number of process parameters which often act in opposite directions. The catalyst temperature increase helps to reduce its amount but requires more packing and larger reactor vessel and pressure drop, or more auxiliary fuel to sustain the reactor operation. Increase in superficial velocity of the gas flow results in smaller reactor and less heat energy requirements but increases pressure drop. In Ref. 38, Matros et. al. showed that extensive computer simulation makes it possible to find an optimum combination of process parameters to minimize a cost criterion that combines annualized capital and operating costs.

A further development of RCO, reported by Zagoruiko et. al. [39] uses the observation that oxidation of a many VOCs proceeds via formation of partially oxidized, polymer film on the catalyst surface as follows: CdtmOk + [(9]

<-->

[P]

[P] + 02

~-~

C02 + H 2 0 + [(9].

149 A copper-chromite catalyst combines large VOC adsorption capacity, 3.5-10"~ mol/m2 [40] with large internal surface area (a property shared by many oxide systems used for VOC control), up to 100 m2/g [41]. This allows over 150 hours of the catalyst operation in adsorption mode. When the catalyst is saturated, periodical flow reversals and external heating are turned on. Once ignited, the catalyst then uses the heat of oxidation of adsorbed VOCs, until these are depleted. Then the catalyst operates as an adsorbent only, and the cycle is repeated. This method appears to be effective for treatment of very diluted VOC streams, using up to 15 times less heat energy than an 'usual' RCO for treatment of a gas flow containing 200 mg/m3 of styrene. Of course, application of this method requires investigation of multicomponent adsorption of various VOCs for each particular case and continuous monitoring of the output gas composition. Oxidation of SO2 into SO3 is a classic example of exothermic reversible reaction. Optimal temperature regime for such a reaction requires starting at as high temperatures as a catalyst can handle, then the temperature decrease along with progressing conversion of a reactant. It is traditionally performed in multi-bed adiabatic reactors with intermediate cooling. Temperature profile in an RFR (Fig. 4) has lower temperatures at both ends of the catalyst bed, suggesting that an RFR would perform close to the theoretical optimum. Theoretical analysis [2] suggests a general strategy to improve RFR performance by providing larger residence time along the decreasing temperature profile. This strategy works for other exothermic reversible processes, such as of synthesis of methanol and ammonia (see Refs. 2 and 42). This can be obtained using larger catalyst pellets, lower linear velocity, or lower adiabatic temperature rise. However, these changes in parameters do not favor stability of operation and require a longer catalyst bed. Several industrial catalytic reactors with periodical flow reversal are used in non-ferrous metallurgy for treatment of lean sulfurous gases with SO2 concentration varying from 1 to 4.5 %, providing for simplified design, lower metal weight, and lower pressure drop at the same performance and catalyst loading as in traditional, multi-bed units [2]. Pilot tests carried out in early 80s (see Refs. 2 and 43) showed that dynamic properties of vanadium catalysts strongly affect the performance of an SO2 oxidation RFR. In particular, large amount of SO3 is stored in the catalyst and exits the reactor with significant delay. When it does, its instantaneous concentration exceeded input concentration of SO2. To simulate this complex behavior, Bunimovich et. al. [43] used the dynamic kinetic model by Balzhinimaev et. al. [15] based on the reaction mechanism including three-step catalytic cycle and side processes of reduction of active V 5§ to V4§ and saturation of vanadium complexes in the melt of active component, turning them into inactive form. The RFR simulation showed that the catalyst state dynamics causes substantially lower than predicted with steady-state reaction kinetics SO2 conversion at low SO2 concentrations (<6 %), and especially at short cycles. At higher input SO2 concentration the difference decreased (Fig. 5). Higher input temperature of the gas flow also results in improved performance of the RFR and smaller difference between the results obtained with dynamic and steady-state kinetic models [43]. This work resulted in a number of practical conclusions: a) though steady-state kinetics can be used for the design of RFR for SO2 oxidation, it overestimates SO2 conversion at low input concentrations, and a correction must be made; b) negative effect of transient processes in the catalyst can be diminished by increasing input temperature of the gas flow, for example, by installing flanking beds of inert packing.

150

d

0. ._

I,.. > tO 0

100..3'

95

..,....,.., --"-

90

(/)

85

=>

80 75

o

70

0

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

40

. . . . . .

z , , , , , , , , , J , , , , , , , , , i

80

120 Cycle period, rain

160

Fig. 5. Effect of cycle period and inlet SO2 concentration on the performance of flow reversal reactor. (1-3): SO2 conversion predicted with dynamic kinetic model, (1'-3') SO2 conversion predicted with steady-state kinetic model. (1,1'): 9 % SO2 in the reactor inlet, (2,2') : 6 % SO2, and (3,3): 3 % S02.

Process of selective catalytic reduction of nitrogen oxides by ammonia (SCR) involves injection of ammonia into a gas stream containing nitrogen oxides, then reduction of NOx by ammonia on the surface of a catalyst typically containing Vanadium oxide on titania. The reactions involved are mildly exothermic (additional heat is required in most cases). Limits of the optimal process temperature, usually from 200 to 350~ are dictated by catalyst activity at low temperatures and by the reaction selectivity at high temperatures. The NOx-containing gas flows often have low temperature and variable flow rates and concentrations. This combination of factors makes application of an RFR to NOx reduction advantageous. One industrial unit for NOx selective catalytic reduction was reported to operate in Russia [44], with ammonia water injection between two catalyst beds. Advantages of utilizing high ammonia adsorption capacity of the vanadium catalyst for NOx reduction were recognized by Agar and Ruppel [45]. Later, Noskov et. al. [40] used the following reaction mechanism: NHs + [ ] ~ I-NHs] [NH3] + NOx + 02 ~ N2 + H20 + [ ] l ~ d + 02 -~ N2 + 1-120 + [],

and showed that the dynamics of ammonia interaction with the catalyst results in superior performance of RFR over traditional systems. When ammonia is added to the gas flow in between two halves of the catalyst bed, it is adsorbed on the downstream catalyst, and is partially used for NOx reduction. After the flow reversal, the incoming NOx react with this adsorbed ammonia and clear the catalyst surface. Thus, a wave of adsorbed ammonia, similar to heat wave, is formed in the catalyst bed, as shown by our simulation results in Fig. 6. Reversals of the flow do not permit ammonia to leave the bed, and this allows operation at increased NH3/NO~ ratio, with better SCR efficiency and lower ammonia slip. Besides, ammonia adsorption makes an RFR much less sensitive to short-term oscillations of incoming NOx loading. Distributed ammonia input, when a larger part of ammonia is added at the point within the catalyst bed closer to the gas inlet, allows further improvement of the process performance due to better utilization of the catalyst's adsorption capacity [44].

151

400 -

tO

Temperature rise due to fuel injection

Catalyst

o

~

~

t

bed

/

~/Catalysti ~

1::

2~ 3 0 0 CZ.

I---

200

~"

100

o

e--

o

'

0.0

'

i 0.4 [.,.

'

0.8

"

- z ~

(6")600-7

NH3 injection Length of beds, dimensionless

Catalyst =

j ~ Catalyst

= =

'

I,..

Adsorbent

(9 Q.

-.~

o~

E 4o0F-

m~

f o.

1 , I , I , I , 1-~ t 0.0 0.4 0.8 Length of beds, dimensionless

o 0o

0

\

1" 0

,

0

x

o zo

'

I

'

I

'

1 '

i

'

1

I

I

0.1

'

"

0.2

Length of beds, m

0.0 0.4 0.8 Length of beds, dimensionless

Fig. 6. Behavior of De-NOx flow reversal reactor. (Computer simulation based on model SCR reaction kinetics). (1,1 ', 1", 1"'), (2,2',2",2"') and (3,3',3",3"'): temperature, gas phase ammonia concentration, ammonia coverage, and NOx conversion profiles at the beginning, middle and end of flow reversal period.

Fig. 7. Temperature profiles in flow reversal reactor during gradual increase in inlet temperature. Modeling of automotive converter operation. 1-8: profiles at 0, 1, 2, 4, 6, 8 and 10 min after converter start-up. Cycle duration is 1 min.

3.3. Adsorption-catalytic system for CO and NOx control A technology utilizing dynamic processes of adsorption and surface reaction was developed by Goal Line (USA) for treatment of exhausts from gas turbine cogeneration facilities [46]. During first phase of the treatment cycle, the exhaust passes through a proprietary catalyst where CO is oxidized to C02, and NO to NO2. The NO2 is chemically reacted or adsorbed on a coating included in the Goal Line catalyst while CO2 is released to the atmosphere. When

152 the adsorbent-catalyst is saturated with NO2, it undergoes short periodical regeneration by an undisclosed 'regeneration gas' that reduces NO2 to N2. Unfortunately, no further technical details is available to us. 3.4. FUSO in automotive applications Catalysis is used in automobiles for treatment of engine exhausts. Operation of catalytic converters is inherently unsteady-state, where the external forcing is imposed by the engine operation. Typically, the composition of the spark-ignition gasoline engine fluctuates between slightly lean (oxygen excess) and slightly reach (oxygen deficiency) with a period of-- 1 sec. Modem automotive catalysts perform three reactions (therefore - three-way catalyst, or TWC): oxidation of carbon monoxide, oxidation of hydrocarbons and reduction of nitrogen oxides. Operation of TWC requires nearly stoichiometric ratio between fuel and oxygen in the engine combustion chamber thus making an unique example of a tail-pipe treatment unit determining how the main production line (an engine) works. Silveston [47] recently reviewed several studies dealing with effect of input composition forcing on the converter operation. The main conclusion is that oscillations are unlikely to positively affect the performance because of sub-optimal cycling period and high conversion levels achieved. However, more complex FUSO modes are expected to solve a number of impending problems of automotive exhaust control, as illustrated by the following examples. It is known that lean-bum combustion is attractive from the standpoint of fuel economy and CO emission reduction, but it hard to implement because state-of-the-art catalysts do not provide sufficient NOx removal because of oxygen excess. Extensive efforts aimed at development of catalysts for NO decomposition and for lean NO reduction by hydrocarbons were not successful to date. B6gner et. al. [48] investigated a method of overcoming this problem by adding a NOx adsorbent to the TWC catalyst. Engine operation was periodically switched from lean-bum to rich-bum. Nitrogen oxides were stored in the adsorbent during the lean phase: NO + 8902 +->NO2 (over Pt) MOx + NO2 +-~ IMNOj.,~,

and removed during rich phase, along with adsorbent regeneration: [MNOJ,d, ~ NO + MO~ + 8902 NO + CO ~ 89N2 + C02.

The adsorbent provided for optimum performance at 30 s lean and 30 s rich phases, with about 94 % of NOx removal. However, significant improvement in fuel economy requires longer cycles for this approach to be successful, that makes it necessary to have a higher capacity adsorbent. Another problem is related to the expanding use of alternative fuels, such as natural gas. For example, natural gas fueled diesel exhausts often are too cold for currently available oxidation catalysts that need -~ 450~ for methane destruction required in Japan and Europe. Our simulation showed that the reverse-flow operation makes it possible to solve this problem. Once ignited, the bed of monolith supported oxidation catalyst is able to maintain high temperatures over all spectrum of engine operation modes, from idle to full load, with adequate methane destruction. Other hydrocarbons and CO are also completely removed.

153 RFR was proposed in by Houdri and Thomas [49] but have not been applied in automobiles because conventional converter provided for adequate exhaust treatment. Now, an RFR appears to be useful both for wider application of natural gas and for startup emission treatment. Cold start of internal combustion engines results in the release of large amounts of hydrocarbons that cannot be treated by a converter because it has not yet been heated up to operating temperatures. Future regulations in USA, Europe and Japan will require abatement of these start-up emissions. Solutions discussed in the literature include installing a converter close to the engine exhaust manifold, electrical heating, and using adsorbents to retain hydrocarbons until the catalyst is heated up. We suggest to combine dynamic properties of a reverse-flow reactor with the ability of modem adsorbent to retain large amounts of hydrocarbons at low temperatures and release them as the temperature increases. In particular, a bed of zeolite or other appropriate adsorbent can be placed between two beds of catalyst in an RFR, so that the exhaust flow through all three sections is reversed periodically. Fig. 7 shows temperature profiles during start-up of such a converter. Initially, cold exhaust gas flow enters cold convener where hydrocarbons are adsorbed. When the exhaust temperature increases, catalyst becomes heated first. Reversal of the flow provides for preheating both catalyst beds surrounding the adsorbent. When the catalyst becomes hot, the adsorbent temperature begins to increase. Hydrocarbons are released from the adsorbent and pass to the already preheated catalyst where they are destroyed. When both catalyst and adsorbent are sufficiently preheated, the converter can further operate with or without flow reversals, depending on the exhaust temperature. Though only preliminary simulation studies have been conducted to date, development and optimization of such a converter will require information on the kinetics of reactions involved, as well as of adsorption on the reaction mixture components on various promising adsorbents. CONCLUSIONS Interaction of the dynamic properties of a catalyst, a micro-scale physic-chemical system, and the dynamic properties of the macro-scale reactor creates an opportunity to improve the performance of catalytic processes using forced unsteady-state operation. Forced dynamic operation makes it possible to generate spatio-temporal patterns of temperature, composition and catalyst states that cannot be attained under steady-state operation. Probably, a potential for increase of efficiency of partial oxidation processes due to FUSC is tremendous. Frequently observed small effects apparently result from current research philosophy: trial-and-error experimentation with different forced periodic regimes while searching for drastic improvement. We support different paradigm. Money should be spent for elucidation of the process mechanism, developing models appropriate for numerical simulation, search for optimums using computer simulation, then testing the results in the lab or pilot unit. This is how we have been working for a long time, and this is how success was achieved in the development of reverse-flow operation. The above examples demonstrate that incorporation of the catalyst surface dynamics opens new ways for further improvement of the reactor operation, beyond the results achieved using reverse-flow and riser reactors. Transient operation of chemical reactors created new class of optimization problems, more complicated than those being solved to optimize traditional, steady-state fixed-bed reactors.

154 Fortunately, fast increase in the availability of low-cost computing resources simplifies handling multidimensional systems of differential equations of chemical reactor dynamics. A fundamental basis for cyclic optimization of catalytic reactors has been developed. It is based on detailed knowledge of reaction kinetics and fundamental process of mass and energy transport. Power of mathematical modeling and computer simulation has been demonstrated for several reaction systems. It is recommended to invest in fundamental investigations of reacting systems and development of adequate reactor models that could further employ continuously decreasing cost of computer simulation to achieve optimal regimes of chemical reactor performance. We have not discussed many interesting results obtained in areas of optimal control of processes involving rapid deactivation and regeneration of catalyst, processes in inherently unsteady fluidized beds, biological systems.

REFERENCES

1. G.K. Boreskov and Yu.Sh. Matros, Catal. Rev. - Sci. Eng., 25 (1984) 551. 2. Yu.Sh. Matros, Catalytic Processes under Unsteady-State Conditions, Elsevier, Amsterdam, 1989. 3. S. Viswanathan and R. Aris, Proc. 3rd ISCRE, Adv. Chem. Ser., 133 (1974) 191. 4. T. Petroulas, R. Aris, and R.W. Carr, Chem. Eng. Sci., 40 (1985) 2233. 5. B.B. Fish, R.W. Carr and R. Aris, Chem. Eng. Sci., 43 (1988) 1867. 6. B.B. Fish and R.W. Carr, Chem. Eng. Sci., 44 (1989) 1173. 7. A. Ray, A.L. Tonkovich, R. Aris and R.W.Carr, Chem. Eng. Sci., 45 (1990) 2431. 8. A. Renken, Int. Chem. Eng., 24 (1984) 202. 9. A. Renken, Int. Chem. Eng., 33 (1993) 61. 10. P.L. Silveston, in Reactions and Reaction Engineering (R. A. Mashelkar and R. Kumar, Eds.), Indian Academy of Sciences: Bangalore, India, pp. 217-246. 11.P.L. Silveston, in Unsteady State Processes in Catalysis (Ed.: Yu. Sh. Matros), VSP, Utrecht, The Netherlands, Tokyo, Japan, p. 217-232. 1990. 12.R. Yadav and R. G. Rinker, Chem. Eng. Sci., 44 (1989) 2191. 13. V.D. Sokolovskii, Catal. Rev. - Sci. Eng., 32 (1989) 1. 14.E.A. Mamedov, Appl. Catal. A: General 116 (1994) 49. 15. B.S. Balzhinimaev, A. A. Ivanov, O. B. Lapina, V. M. Mastikhin and K. I. Zamaraev, Faraday Discuss. Chem. Soc., 87 (1989) 133. 16. R. R. Sandhankar and D. T. Lynch, J. Catal., 149 (1994) 278. 17. E. Jobson, O. Hjortsberg, S.L. Andersson and I. Gottberg, Reactions over a Double Layer Tri-Metal Three-Way Catalyst, SAE Paper 960801, 1996. 18. J. E. Bailey, Chem. Eng. Commun., 1 (1973) 111. 19. G. Guardabassi, A. Locatelly, and S. Rinaldi, Joumal of optimization Theory and Applications, 14 (1974) 1. 20.D. Sincic and J.E.Bailey, Chem. Eng. Sci., 35 (1980), 1153. 21. L.E. Sterman and B.E.Ydstie, Chem. Eng. Sci., 45 (1990) 737. 22. L.E. Sterman and B.E.Ydstie, AIChE J., 37 (1991) 986. 23. F.J.M. Horn and R.C. Lin, Ind. Eng. Proc. Design. Develop. 6 (1967) 21.

155 24 F. Ozg~lsen, R.A. Adomaitis and A. t~inar, 1992, Chem. Eng. Sci., 47 (1990) 605. 25 C.-C. Chen, C. Hwang and R.Y.K. Yang, Can. J. Chem. Eng. 72 (1994) 672. 26 I.A. Zolotarskii, Yu.Sh. Matros, React. Kinet. Catal. Lett., 20 (1982) 321. 27 M.A. Truffer and A. Renken, AIChE J., 32 (1986) 1612. 28 J. Thullie and A. Renken, Chem. Eng. Sci., 46 (1991) 1083. 29 W.R.C. Graham and D.T. Lynch, AIChE J., 36 (1990) 1796. 30. V. O. Strots, Yu.Sh. Matros and G.A. Bunimovich, Chem. Eng. Sci., 47 (1992) 2701. 31. F. 0zgialsen and A. t~inar, Chem. Eng. Sci., 49 (1994) 3409. 32. R. M. Contractor, H. E. Bergna, H. S. Horovitz, C.M. Blackstone, U.Chowdhry and A.W. Sleight, Catalysis 1987, J.W. Ward (eds.), Elsevier Science Publishers B.V., Amsterdam, p. 645, 1988. 33. G. S. Patience and P.L. Mills, in: New Developments in Selective Oxidation l-I, V. C. Corberan and S. V. Bellon (eds.), Elsevier Science B.V., Amsterdam, The Netherlands, p.1, 1994. 34. A. A. Ivanov and B. S. Balzhinimaev, in: Unsteady-State processes in Catalysis, Yu. Sh. Matros (eds.), VSP, Utrecht, The Netherlands, p.91, 1990. 35. Yu. Sh. Matros, and G.A. Bunimovich, Cat. Rev. - Sci. Eng, 38 (1996) 1. 36.J.J. Spivey, Catalysis. A Review of Recent Literature, 8 (1987) 157. 37. A. A. Barresi and G. Baldi, Ind. Eng. Chem. Res., 33 (1994) 2964. 38. Yu. Sh. Matros, G. A. Bunimovich, S. E. Patterson and S. F. Meyer, Catalysis Today, 27 (1996) 307. 39. A.N. Zagoruiko, O. V. Kostenko and A. S. Noskov, Chem. Eng. Sci., 51 (1996) 2989. 40. A.S. Noskov, L.M. Bobrova, G.A. Bunimovich, O.V. Goldman, A.N. Zagoruiko and Yu.Sh. Matros, Catal. Today, 27 (1995) 315. 41. Yu.Sh.Matros, A.S. Noskov, V.A.Chumachenko, Kataliticheskoe Obezvrejivanie Otkhodiaschikh Gasov Promyshlennykh Proizvodstv (Catalytic Processes in Air Pollution Control), V.N. Parmon (eds.), Nauka, Sibirskoe Otdelenie, Novosibirsk, 1991 (in Russian). 42. K. M. Vanden Bussche, S.G. Neophitydes, I. A. Zolotarskii, and G. F. Froment, Chem. Eng. Sci., 48 (1993) 3335. 43. G. A. Bunimovich, N. V. Vernikovskaya, V. O. Strots, B. S. Balzhinimaev and Yu. Sh. Matros, Chem. Eng. Sci., 50 (1995) 565. 44. A.S. Noskov, L.M.Bobrova and Yu.Sh.Matros, Catalysis Today, 17 (1993) 293. 45. D. Agar and W. Ruppel, Chem. Eng. Sci., 43 (1988) 2073. 46. Goal Line Environmental Technologies, SCONOx, Catalytic Adsorption System for Natural Gas Fired Power Plants to Reduce or Eliminate: Nitrogen Oxides and Carbon Monoxide, Technical Brochure. 47. P.L.Silveston, Catalysis Today, 25 (1995) 175. 48. W. BOgner, M. Kramer, B. Krutzsch, S. Pischinger, D. Voitlaander, G. Wenninger, F. Wirbeleit, M.S. Brogan, R.J. Brisley and D.E. Webster, Appl. Catal. B: Environ., 7 (1995) 153 49. E. J. Houdri and W. R. Thomas, Apparatus for Improving the Purification of Exhaust Gases from an Internal Combustion Engine, U.S. Patent No. 3,189,417 (1965).

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91997 Elsevier Science B.V. All rights reserved. Dynamics of Surfaces and Reaction Kinetics in Heterogeneous Catalysis G.F. Froment and K.C. Waugh, editors

159

A p r o b a b i l i s t i c model for the deactivation of a dual f u n c t i o n catalyst by coke formation a c c o u n t i n g for reaction and surface migration. Santokh Singh and Gilbert F. Froment L a b o r a t o r i u m voor Petrochemische Techniek U n i v e r s i t e i t Gent K r i j g s l a a n 281, B-9000 Gent, Belgium. ABSTRACT

The paper considers a sequence of reactions catalyzed by two types of sites, one of w h i c h is subject to deactivation through site coverage by coke. Since the reaction steps take place on d i f f e r e n t types of sites, surface migration of the r e a c t i n g c o m p o n e n t s is also a c c o u n t e d for. The distance over which the m o l e c u l e s have to travel before reaching an apropriate free site also depends on the coke content of the catalyst and evolves with time. A p r o b a l i s t i c model is developed for such a process. It is a p p l i e d to e x p e r i m e n t a l data on n.pentane isomerization on a P t / a l u m i n a catalyst. The effect of coking on the m i g r a t i o n d i s t a n c e s is shown to be significant. INTRODUCTION

Hydrocarbon conversion processes like reforming or h y d r o c r a c k i n g are b a s e d on a dual function catalyst in which one c o m p o n e n t of the catalyst, Pt e.g., catalyzes h y d r o g e n a t i o n and d e h y d r o g e n a t i o n and the other component, alumina or a zeolite with acidic character, the isomerization and cracking. A normal paraffin, e.g., is d e h y d r o g e n a t e d on a Pt-site into a n - o l e f i n that desorbs before moving, either through the gas phase or over the surface to an acid site, where the corresponding i-olefins are produced. These then desorb and move to a Pt-site, w h e r e h y d r o g e n a t i o n into an i-paraffin and desorption take place. These reactions are always a c c o m p a n i e d by undesirable side reactions leading to coke formation. As sites become covered by coke the rate of reaction decreases, but since the distance b e t w e e n sites r e m a i n i n g active increases an additional effect has to be a c c o u n t e d for in the rate of the global process c o n s i s t i n g of reaction and migration. The process c o n s i d e r e d here is shown in Fig. i.

AI.O

~

~

A2.a

A2.'~ -

.

.

.

.

.

.

.

--~

~

A3..~

.

A3.a .

c

.

.

.

.

.

.

~

~

A4.o

.

C

Fig. i. Schematic r e p r e s e n t a t i o n of a process consisting of various adsorptions, desorptions, reactions and migrations.

160 In this F i g u r e o r e p r e s e n t s sites on the P t - p h a s e and r a c i d sites on the a l u m i n a p h a s e . Coke f o r m a t i o n may o c c u r out of a d s o r b e d A a or a d s o r b e d A 3. It d e a c t i v a t e s the i-sites, leading s i m u l t a n e o u s l y to c h a n g e s in L z and L 2. A s s u m i n g s t e a d y s t a t e b u t no rate d e t e r m i n i n g step, L a n g m u i r a d s o r p t i o n and s i n g l e s i t e r e a c t i o n s leads to the f o l l o w i n g e q u a t i o n for the rate of d i s a p p e a r a n c e of A I (viz A p p e n d i x A) : CA. r A~ =

-

(i)

CA4/K

- DEN

w h e r e K is the o v e r a l l e q u i l i b r i u m c o n s t a n t and D E N c o n t a i n s the rate c o e f f i c i e n t s k z, k 2, k 3, .... , the a d s o r p t i o n e q u i l i b r i u m c o n s t a n t s Kz, K 2, ... a n d the m i g r a t i o n d i s t a n c e s L I and L 2. In w h a t follows e x p r e s s i o n s for L I and L 2 w i l l be derived. A PROBABILISTIC APPROACH FOR P R E S E N C E OF COKE F O R M A T I O N

THE

MIGRATION

PHENOMENON

IN

THE

The sites are s u b j e c t to a d s o r p t i o n , d e s o r p t i o n r e a c t i o n a n d coking, so that a site a p p r o a c h e d by a m i g r a t i n g m o l e c u l e c o u l d be e i t h e r free or o c c u p i e d .

t \1 Se~e.'~;al m a ~ o n

on a dual Nnc~on Ca,,alysl

Reaction

and m i g r a t i o n

on the

Fig. 2. surface

of a dual

function

catalyst

Fig. 2 shows a s q u a r e g r i d of o- and r - s i t e s w i t h only one site, o or i, at each g r i d point. A s s u m e that i) all free sites h a v e e q u a l a c t i v i t y for the c o r r e s p o n d i n g reaction. 2) no more than one m o l e c u l e can o c c u p y a site at a g i v e n time. 3) the r e a c t i o n s a l w a y s o c c u r on a s i n g l e site 4) the m i g r a t i n g m o l e c u l e c h o o s e s its d i r e c t i o n r a n d o m l y a f t e r e n c o u n t e r i n g a site w h i c h is c o v e r e d by a r e a c t i n g c o m p o n e n t or

161 by coke. 5) e a c h m i g r a t i o n step or s e g m e n t is l i m i t e d to the a d j a c e n t sites only. The p a t h f o l l o w e d by a m i g r a t i n g m o l e c u l e in its search for a free r- or o - s i t e (rf or or) c o n s i s t s of one or more segments. The n u m b e r of s e g m e n t s is a r a n d o m v a r i a b l e r e p r e s e n t e d by U. The m i g r a t i o n d i s t a n c e s L z and L 2 are r a n d o m variables. The c o n d i t i o n a l p r o b a b i l i t y that the n u m b e r of segments that the m o l e c u l e of the i n t e r m e d i a t e A 2 has to m i g r a t e to reach a free site w h i c h is of the r - t y p e is :

w h e r e u(_> 0) is an i n t e g e r and P~r/s is the p r o b a b i l i t y that site e n c o u n t e r e d by the m o l e c u l e of A m is a free r-site. index s r e f e r s to the e v e n t that the site is e n c o u n t e r e d by m o l e c u l e . The e x p e c t e d n u m b e r of segments, Nz, involved in m i g r a t i o n of A 2 is :

N, = ~

the The the the

u<~ - P : , / ~ ) u P f , / s

u=O

N~ : { ~ - P~'/~}

(2)

P~/s

If no d i s t i n c t i o n is m a d e b e t w e e n the side and d i a g o n a l of a unit of the g r i d the e x p e c t e d d i s t a n c e for the m i g r a t i o n of A 2 can then be w r i t t e n : Lz = { 1 - Pf,/S}d.

(3)

Pf~/s w h e r e d(A) is the l e n g t h of a segment. S i m i l a r l y the e x p e c t e d d i s t a n c e in the m i g r a t i o n by:

of A 3 is given

L2 = { 1 - Pf,,/S }d. Pfo/s

(4)

where P~o/s is the conditional probability e n c o u n t e r e d by the A 3 - m o l e c u l e is a free o-site. P R O B A B I L I T I E S OF T H E O C C U R E N C E OF C E R T A I N P A T H OF T H E M I G R A T I N G M O L E C U L E

that

a

site

T Y P E S OF SITES ON THE

D u r i n g its m i g r a t i o n the m o l e c u l e of the i n t e r m e d i a t e A 2 or A 3 m a y h a p p e n to e n c o u n t e r one of the types of sites w h i c h is e i t h e r n o t free or is not a p p r o p r i a t e for the next reaction in

162 the sequence. What it r e a l l y n e e d s is a free t-site, t+, if it is an A2-molecule and a free c-site, c+ if it is an A3-molecule. Let P+a, Poa, P+~, Poa, and Pd~, respectively, be the p r o b a b i l i t i e s that a site on the c a t a l y t i c s u r f a c e is a c+, Co, t+, t o, r d site. Pa and P~ are the p r o b a b i l i t i e s t h a t a site is a c- or a r-site, i r r e s p e c t i v e of w h e t h e r it is free, o c c u p i e d , or d e a c t i v a t e d by coke. T h e n f r o m Bayes' rule, the p r o b a b i l i t y t h a t the site e n c o u n t e r e d by the A 2 - m o l e c u l e is a f r e e site can be w r i t t e n : Pf, l s -

-

Pslf.pfa

Ps/t'Ps" + Psloapoa + P s l f ~ p f ,

+ Pslo~po~ + Psld~Pd~

(5)

w h e r e Psi+a, Psloa, Pslf~, Psla~, Psld~' are the p r o b a b i l i t i e s that a of, c o, r+, to, or a t d is e n c o u n t e r e d by the m i g r a t i n g m o l e c u l e . Since the m o l e c u l e w i l l o n l y m o v e to an a d j a c e n t site, w h a t e v e r type of site it m a y be, Psi fo = Psio~ = Psi f, = P s i =

the

= Psi<~

U s i n g the m u l t i p l i c a t i o n rule of p r o b a b i l i t y , a b o v e identity, (5) b e c o m e s Pf~Is-

together with

PfI'P~

pf/apa + polopa + pfl~p, + Pol~P~ + pal~p~

(6

)

w h e r e P+ o, Polo are the p r o b a b i l i t i e s that a a - s i t e is free or o c c u p i e g , and P+/~, Po/~, Pdl~ the p r o b a b i l i t i e s t h a t a t - s i t e is free, o c c u p i e d by At, A 2, A s , A 4 or by coke and, t h e r e f o r e , deactivated. Once the r e a c t i o n A 2 = A 3 has t a k e n p l a c e on a r-site, the m o l e c u l e of the i n t e r m e d i a t e A s s t a r t s its s e a r c h for a free csite. U s i n g the same a r g u m e n t as u s e d in d e r i v i n g (6), the p r o b a b i l i t y that the site e n c o u n t e r e d by the A - m o l e c u l e is a free c - s i t e is written: Pfa/s :

Pf/oPo Pf/<,Po + Po/aPa + Pf/~P~ + Po/,P~ + Pd/,P~

(7 )

The d u a l f u n c t i o n c a t a l y s t c o n t a i n s c and r - s i t e s a n d the total c o n c e n t r a t i o n of sites, Ct, is the sum of the c o n c e n t r a t i o n s of b o t h types of sites: Ct =

Cto

+ Cc~

It f o l l o w s that the p r o b a b i l i t i e s that a site l o c a t e d on the c a t a l y s t surface is e i t h e r a c - s i t e or a r - s i t e by :

at r a n d o m are g i v e n

163

Poand P,

(8)

C~

C~ C=

-

(9)

S i n c e it h a s b e e n a s s u m e d t h a t the a - s t e s are n o t s u b j e c t to c o v e r a g e b y coke, a a - s i t e c a n be o c c u p i e d by the r e a c t i o n c o m p o n e n t s o r e l s e it w i l l be a f r e e a c t i v e site. In the s e q u e n c e A ! ~ A 2 = A 3 = A4, t h e r e a c t i o n s A I = A z a n d A 3 = A 4 are s u p p o s e d to t a k e p l a c e o n a - s i t e s . A t any p r o c e s s t i m e a o - s i t e m a y thus be c o v e r e d b y o n e of the c o m p o n e n t s At, A2, A3, A 4. T h e r e f o r e , the p r o b a b i l i t y of a o - s i t e b e i n g an o c c u p i e d s i t e is

q~

Po/a =

while

the

probability

Ca

Pf/a -

A r-site

may

the

a a-site

being

a free

site

is

(11)

be

occupied

by A 2 a n d A 3 w i t h

a probability

: (12)

C~

probability

Pf/~ -

for

(1o)

Ct a

Po/~ =

while

+ C~o + c~o + q o C~ a

that

the

r-site

is f r e e

can

be w r i t t e n

:

(13)

Cf~ C~

T h e r - s i t e s a r e n o t o n l y c o v e r e d b y A 2 a n d A3, b u t a l s o by c o k e f o r m e d b y a s i d e r e a c t i o n w h i c h d e a c t i v a t e s them. D e f i n e t h e r a n d o m v a r i a b l e ~t as the t i m e r e q u i r e d for a rs i t e to be d e a c t i v a t e d . T h e n the p r o b a b i l i t y t h a t at t i m e t the site would be deactivated, g i v e n t h a t it is a r - s i t e is Pd/~ = P ( ~ t

< t)

It f o l l o w s f r o m B e e c k m a n phase composition :

(14)

and

Froment

[I]

that

for

a given

(15)

P(~c < t) = [I - e x p ( - r~ w h e r e r~ the r a t e

is t h e i n i t i a l in t h e a b s e n c e

r a t e of f r a c t i o n a l of coke).

gas

site

coverage

(i.e.,

164 From

(14) and

(15)

Pd/~ = [i- exp(- r~

(16)

Two cases are c o n s i d e r e d for the coverage, d e p e n d i n g upon w h e t h e r coke is f o r m e d through a m e c h a n i s m p a r a l l e l or c o n s e c u t i v e to the r e a c t i o n A 2 = A 3. Coke f o r m a t i o n

out of A 2.

The rate e q u a t i o n for the site coverage A 2.: = C c.r is [Froment and B i s c h o f f [2]] :

dC~ dt

through

the

step

_ i kc~ (A2) (Cc~ - Cc) C2" } C~ + C2~ + C3,

(17)

If M e is the m o l e c u l a r w e i g h t of the coke the initial rate of coking, then

dC~ dt Combining

_

~r ~ Ct, - C ~ } _ (aMc){ C~,

(17) and

(18)

precursor

and

re~

r=~

(18)

~zMc

leads to

rc =

(19)

where ~c is the d e a c t i v a t i o n [Froment & Bischoff, 1990]. Further, the initial rate of p r e c u r s o r is r2

=

is

function

for

fractional

the

site

coke

coverage

~g

by

coke

(20)

aM~Ct,

Since r e = rc~

formation

combination

of

(19) and

(20) gives

k ~ (A2 ) C2~

S u b s t i t u t i n g rs~ from the above e q u a t i o n p r o b a b i l i t y that the site is deactivated, site, w h e n coke is formed out of A 2 Pd/, = I - e x p { -

k = (A 2) C2,t

(C~ + ~

+ C3,)

}

into (16) y i e l d s the given that it is a r(21 )

165 Coke

occur is

formation

of

A3

A s s u m i n g a g a i n t h a t b o t h A 2 = A 3 and the c o k i n g r e a c t i o n on s i n g l e sites, the rate of f o r m a t i o n of coke p r e c u r s o r

dCcl dt From that

out

_ { k ~ (A 3) ( C ~ - C=> C 3~ } C~ + C2~ + C3~

(22)

this the p r o b a b i l i t y that the it is a r - s i t e is g i v e n by:

Pd/~ : 1 - e x p { -

MIGRATION DISTANCES PARAMETERS.

site

is d e a c t i v a t e d ,

given

k=(A3)C3~t } C~ + C2~ + C3,)

IN

TERMS

OF

(23)

CONCENTRATIONS

AND

KINETIC

C o m b i n i n g (2) a n d (6) leads to the f o l l o w i n g e x p r e s s i o n the e x p e c t e d n u m b e r of s e g m e n t s in the f i r s t m i g r a t i o n : N, =

P~zoP" + P o / o P . + Po/~P~ + P~/~P,

A c c o u n t i n g for Cto = (24) can be w r i t t e n /71 -

c~

c~

CI~

C~ [--~, + ~ c -

+ C2~

exp(-

Pd/,P,

+

C3~

for

C24)

+ Cc~

and

C t = Cto

k= (A5) C~t

C~ + C2T + C3~ ) ]- 1

+ Ct~ (25)

w h e r e j = 2 for c o k i n g f r o m A a a n d j = 3 for c o k i n g from A 3. A s i m i l a r e x p r e s s i o n for N 2 is e a s i l y derived. Finally,

L~ = 1714

(26)

L~ = N 2 d

(27)

RATE EQUATIONS AND L 2.

CONTAINING

THE

PROBABILISTIC

EXPRESSIONS

FOR

Lz

S u b s t i t u t i n g e q u a t i o n s (26) a n d (27) for L I and L 2 into the rate e q u a t i o n rAt and i n t r o d u c i n g xl, the c o n v e r s i o n of At, leads to the rate e q u a t i o n g i v e n in A p p e n d i x B. for M o d e l 1 :

166 rA~ = f(xl; e~ .... e s) for M o d e l

2 in w h i c h

(28)

the m i g r a t i o n

is ignored: (29)

rA~ : F ( x I ; 8~ .... 86 )

The p a r a m e t e r s 81 .... 86 are r e l a t e d to a d s o r p t i o n , d e s o r p t i o n a n d c h e m i c a l r e a c t i o n , w h i l e 07 and 88 r e l a t e a l s o to m i g r a t i o n .

n.PENTANE

ISOMERIZATION.

PARAMETER

ESTIMATION.

De P a u w and F r o m e n t [1975] s t u d i e d n. Cs i s o m e r i z a t i o n on a P t / a l u m i n a c a t a l y s t in an i s o t h e r m a l t u b u l a r r e a c t o r w i t h p l u g flow, y i e l d i n g n. Cs-conversion, xl, v e r s u s W/Fk~ Therefore, the f o l l o w i n g o b j e c t i v e f u n c t i o n was c h o s e n for the p a r a m e t e r estimation : n

A c c e s s i n g W/FA ~ the c a l c u l a t e d space time, r e q u i r e s s u b s t i t u t i o n of the rate e q u a t i o n s for A I (28) or (29), b u t w i t h c o k i n g o u t of A 3 only, into the c o r r e s p o n d i n g continuity equation and i n t e g r a t i o n b e t w e e n inlet and exit of the i n t e g r a l reactor. The r e s u l t i n g e q u a t i o n s are g i v e n in A p p e n d i x C. The p a r a m e t e r s w e r e e s t i m a t e d f r o m the d a t a at 412~ at w h i c h t e m p e r a t u r e the o v e r a l l e q u i l i b r i u m c o n s t a n t takes on a v a l u e of 2.42. The parameter values and the approximate 95 % confidence i n t e r v a l s are g i v e n in Table i. Parameter

Model

with migration

Model

Estimate

81 82

8

68.98

8

migration

Estimate

+ 13.69

71.09

+ 12.05

9 . 2 1 x 1 0 -3 + 1.14x10 .3

3 . 5 1 x 1 0 -2 + 4 . 6 5 x 1 0 -3

1 . 3 8 X I 0 -3 + 1.62XI0 -4

2 . 8 6 x i 0 -2 + 2 . 1 7 X i 0 -3

,

84

without

133.6

+ 19.67

135.5

+_ 2 0 . 1

67.82

+ 10.67

67.82

+ 8.69

,

,

.

.

.

.

.

.

1 . 9 9 x i 0 -2 + 1.97xi0 -3 ,,

87 Table

1.60

+ 0.29

i. E s t i m a t e s a n d a p p r o x i m a t e the p a r a m e t e r s .

m

95 % c o n f i d e n c e

intervals

of

167

RESULTS

AND

DISCUSSION.

A h i g h l y s i g n i f i c a n t F - v a l u e of 988.7 at the 95 % c o n f i d e n c e level indicates that Model I, accounting for migration, a d e q u a t e l y d e s c r i b e s the p r o c e s s . The t - v a l u e s p e r t a i n i n g to the p a r a m e t e r s of the model all e x c e e d the (I - 0.95) t a b u l a t e d value of 2.201, e s t a b l i s h i n g the s i g n i f i c a n c e of these parameters. Also shown in Fig. 3 is the rat VS W/FA ~ curve for M o d e l 2, i g n o r i n g migration. The F - v a l u e a m o u n t s to 565,2. The fit of this model is less s a t i f a c t o r y , particularly at the h i g h e r W/FA~ i.e. at h i g h e r c o n v e r s i o n s of A I and h i g h e r coke contents, w h e n the i n f l u e n c e on the m i g r a t i o n distances is b e c o m i n g more pronounced. The R2-value of 0 . 9 9 9 9 9 9 for M o d e l 1 as c o m p a r e d with 0.872191 for M o d e l 2 i n d i c a t e s that the v a r i a t i o n e x p l a i n e d by the former is m u c h larger. Fig. 4 shows the e v o l u t i o n w i t h the coke c o n t e n t of the m i g r a t i o n d i s t a n c e s LI and L 2 in a p o i n t in the r e a c t o r w h e r e PA~ has d e c r e a s e d from its inlet v a l u e of 0.757 atm to a v a l u e of 0.675. Initially, i.e. at zero c o k e c o n t e n t L~ and L 2 a m o u n t to 3.31 and 2.47 A r e s p e c t i v e l y , as c a l c u l a t e d from (26) and (27) using the e s t i m a t e of Nz a n d N 2 o b t a i n e d f r o m (24) and (25). The m i g r a t i o n d i s t a n c e s v a r y w i t h the coke c o n t e n t as a c o n s e q u e n c e of two o p p o s i n g e f f e c t s : the c o v e r a g e of r - s i t e s by coke formed out of A 3 that i n c r e a s e s L I and L 2 and the d e a c t i v a t i n g effect of coke on the r e a c t i o n A2r ~ A3r that d e c r e a s e s the c o v e r a g e of r-sites by A3, thus r e d u c i n g L I a n d L 2. The m a g n i t u d e of the v a r i a t i o n s of L I and L 2 d e p e n d s on the local c o n c e n t r a t i o n s of A 2 and A 3 in the reactor.

-t .... [ 9

9

modelwith migration

model without migration experimental

0.01

!

t

I

100

L I

o< ..... _J "o

--

~

0.008

8o

..

~

Q o e-

0.006

A~',

u~

T

"0

- 412"C

Pt = 2.49 arm

~,

0.0t34

60

~

._

~o

E 2O

0.002

0

I

1

:

~

i

20

40

60

80

100

space

Fig. 3. o f A~ v e r s u s s p a c e l e n g t h of 10 h o u r s .

R a t e of d i s a p p e a r a n c e

a run

120

0

time

o.ooos

0.001 coke

t i m e " W / ~ c s( g cat hr / mol ) after

Migration

distance

o.oo~.s

o.oo2

o.o02s

content(9coke

Fig. 4. L~ and L 2 v e r s u s

o.oo3 o.o03s o.o04 /gcat )

coke content

at 412"C.

168

NOTATION Capital letters Ci Cc~ Cn_cs~ C1o , C1~ Car,

C~f

Cta ,

Ct~

Ct

D I, D2 Ki K Lz, Mc

L2

Nz ,N2

gas p h a s e c o n c e n t r a t i o n of Ai, mol cm ~ c o n c e n t r a t i o n of r - s i t e s c o v e r e d by coke, mol goat-I initial gas p h a s e c o n c e n t r a t i o n of n - p e n t a n e , mol cm -3 c o n c e n t r a t i o n of o - s i t e s and r-sites, c o v e r e d by Al, mol gcat-i c o n c e n t r a t i o n of free o-sites and free r-sites, mol gcat- 1

total

Scat -I

concentration

of

r-sites

and

of

r-sites,

mol

total c o n c e n t r a t i o n of sites, mol gcat-I s u r f a c e m i g r a t i o n c o e f f i c i e n t s , cm 2 hr -I equilibrium c o n s t a n t for the i-th e l e n e n t a r y step, cm3/mol overall equilibrium constant migration distances m o l e c u l a r w e i g h t of a coke molecule, g c o k e / m o l e x p e c t e d n u m b e r of s e g m e n t s in m i g r a t i o n

Small letters

i n t e r n a l s u r f a c e area, cm 2 gcat-1 d i s t a n c e b e t w e e n two a d j a c e n t sites, cm or A c o k i n g rate c o e f f i c i e n t , hr -1 a d s o r p t i o n rate coefficient, cm 3 mol -z hr -I d e s o r p t i o n rate c o e f f i c i e n t , hr -I s u r f a c e r e a c t i o n rate coefficient, hr -1 initial rate of coking, g c o k e / g c a t . h r rate of d i s a p p e a r e n c e of A I, mol goat-I hr -l initial rate of f r a c t i o n a l site coverage, hr -I time, hr c o n v e r s i o n of A I

a i

d kc~ ki k~ kl rc ~

r A1

rs ~

t

XAI

Greek symbols

Zf, rf, a 8i

0o to,

free or o c c u p i e d h y d r o g e n a t i o n and d e h y d r o g e n a t i o n site r d free, o c c u p i e d or d e a c t i v a t e d i s o m e r i z a t i o n site c o n v e r s i o n factor, mole coke per mol site i-th p a r a m e t e r of the model

REFERENCES B e e c k m a n J.W. and F r o m e n t G.F., 1980, Chem. Engng. Sci., 35, 805. F r o m e n t G.F. and B i s c h o f f K.B. "Chemical R e a c t o r A n a l y s i s and D e s i g n " 2nd Ed. (1991), John W i l e y and Sons, N e w York. De P a u w R. and F r o m e n t G.F., 1975, Chem. Engng. Sci., 30, 789.

169 APPENDIX

A

R a t e of d i s a p p e a r a n c e

of A1

M o d e l 1, a c c o u n t i n g for m i g r a t i o n s

Let the process represented in Figure 1 proceed through the following steps: (i) Adsorption of A1 from the gas phase on a e-site: kl A1 + o" = A l . o k-1

;

rl = k l C 1 C r

- k-1 C1~

(ii) Reaction on or-site producing the intermediate A2: k2 Al.cr = A2.cr k-2

;

r2 = k2C1~ - k - 2 C 2 a

(iii) Desorption of A: from or-site: k3 A2.o" ~ A2:~ % o" k-3

;

I"3 -- ]c3C2a - ~-3C2Co.

(iv) Migration of A2 over a distance L1 from a a-site to a free r-site: A2(cr) ~ m2(r)

;

Tin1 :

aiD1

(" L1

)(C2a -C2r)

(v) Adsorption of A2 on a r-site: k4 A2 + r ~- A 2 . 7 k-4

;

r4 -- k 4 C 2 C r - ]c_4C2

(vi) Reaction of A2 into A3 on a w-site: ks A2.r ~ A3.r k-s

;

rs -- ksC2r - It-sCar

;

r6 - ]c-6C3r- k 6 C 3 C r

(vii) Desorption of A3 from T-site: k6 A3.T ~ A3 -k r k-6

(viii) Migration of A3 over a distance L2 from a r-site to a cr-site: A3(r) ~ Aa(a)

;

rm2

aiD2

= (-~2)(C3,

-

C3o.)

170 (ix) Adsorption of Aa on a a-site:

k7

A3 + a = A3.a

r~ = k~C3C~ - k-7C3~

(x) Reaction of A3 into A4 on a a-site: ks A3.a = A4.a k-s

rs

--

ksC3a

-

k-sC4~

(xi) Desorption of the product A4 from a a-site"

]r A4.a ~ A4 + a

r9

"-

k9C4a - k-9C4Ca

k-9

The total concentration of a- and r-sites is given by Ct~ = C~s + C~a + C2a + C3a + C4a Ct. = C,j + C2. + Ca, + Co,

In the steady state and if there is no rate determining step rl

- - 7"2 - " r 3

--

rml

- - r 4 - - /'5 - - r 6 - - r r n 2 =

T7 - - F8 m 7.9

so that the rate equation for the disappearance of A1 can be written, after elimination of the unobservable concentrations Ca, Cla, C2a, C3o, C4~, C2~, C2,, C~, C2,, (73,, C3a, and C3~" CA1 - C A 4 / K rA1 = DEN ( A - 1) where K is the overall equilibrium constant and DEN

=

1{11

1 }

C--~ ~ + Klk2 + K1K2k-3

+

l { K3K6 K3K6 K3K6 } Ca K1K~K4Ksk~ + K1K2K4KsK~ks + K1K2K4KsKTKsk_9 ' +

1{/(3

K3

C~ K1 K2 k--------~ + K1 K2 K 4 rfK3

L1

lt, g l g 2 ) ( a - ~ ) " ~ ( g l

/(3} K1 K2 K 4 K s k _ 6 + K3K6 L~ K2K4K5 ) ( aids )}

ks

+

Also, the concentrations in the gas phase and those on the active sites can be related by the Langmuir adsorption isotherms. Cla - KI C1CI~, C2a = KIK2 C1Cla

171 C3~, = (K1 K2

K41(5

K1K2~ (K4/(5 C~, = (K~ K2 K3 ) K4 CI C~.

C3.r = \( KIKaK2 ) K4KsC1C~The corresponding equation for model 2 ignoring migration is easily derived.

APPENDIX

B

Introducing the relations for L1 and L2 and also the conversion 2:1 of A1 into equation (A-l) leads to the following expression for the rate of disappearance of A1

[1- (1+ ~) x,] FAi -- [01 {1 + 02(1 -- xl)} + (~-2~)03 {1 + 04(1 -- Xl) } -F 06IV(1)x]+ 072~[(2)]Xl J with ~)

=

l+04(1-xl)][l+r162 r

1 0s(1_ z04(1 l ) t + 231)) }

-1

and

.N(2) = [1 + where

01 =

O~(1r]z7--[ Xl)

(1)[{1

1 + r

1

+ r/,,,- -

{ ( ~ exp

-1

+ 84(1 - x l )

1 }

)} 1

- 1

C~Ct~, -~1 + K1k2 + K1K2k-3 { K3K6 K3K6 K3K6 } + K1K2K4Ksk7 + K1K2K4KsKTks + K1K2K4KsK7Ksk_9 { KIK2K4Ks K1K2K4K~KTKs} 02 - C~ K1 + K1K2 + K3K6 + KaK6 03

-

-

CfCt, K1K2k4+ K1K2K4ks + K1K2KsK4k-6 {K1K2K4 K1K2K4K~,}

04 --" C;

1(3

+

K3

172

{

for consecutive coke formation 0s = C~ kctK1K2IQKs}, K3 06 =

d

(

C~K1 K2 /

\

07 =

aiD2

Cf K1K2K4Ks

(~

Ct~-]

The parameters 01, 02, 03, 04 and 07 are related to adsorption, desorption and reaction, while 05 and 06 relate in addition to migration. This is the rate equation of Model 1, accounting for the migration. In Model 2 this phenomenon is neglected, so that the corresponding rate equation rA1 does not contain the groups 05 and 06.

APPENDIX

C

The reaction is carried out in an isothermal tubular reactor with plug flow. Accordingly, when the rate equation for the disappearance of A1 is substituted into the integrated continuity equation for A1,

FA = f f ~ ~ the resulting equation is W

FA~

=

[ (~) ( #1+03

+#6 i +

d~

) ( l+~c)]•

l+r/~ +07 I+ r

'7~

~+i) ~~ 1- (1+ ~)~ + 0102+0304 KK +1)

+0406

[

1 (K+I)

l~

1+

{

r

1 1--(l+K)

+0207

1+

rh,,

Xl

n=0 m=0 {

1 K+I

[{(

x

77

( 11 (I+K+04)"+I K+l

0s(1-xl) 1+04(1-zl)

04

)~ (

exp -

(l/}](n_m),( E

+(n+l)

0s(1-_z~.! 1+04(1-Xl)

0~ ~ "\0~ +1] •

)} {( )~ ( )} -

0s 1+04

0s exp - N

This Page Intentionally Left Blank

91997 Elsevier Science B.V. All rights reserved. Dynamics of Surfaces and Reaction Kinetics in Heterogeneous Catalysis G.F. Froment and K.C. Waugh, editors

175

Self-sustained isothermal oscillations m N 2 0 decomposition on Cu overexchanged ZSM-5 P. Ciambelli a, A. Di Benedetto b, E. Garufi% R. Pirone ~ and G. Russo ~ aDipartimento di Ingegneria Chimica e Alimentare, Universita di Salerno 84084, Fisciano (S A), Italy bDipartimento di Ingegneria Chimica, Universita di Napoli "Federico IF', p.le Tecchio, 80125 Napoli, Italy CIstituto di Ricerche sulla Combustione, CNR, p.le Tecchio 80, 80125 Napoli, Italy

Regular periodic isothermal oscillations of N20 outlet concentration are experimentally observed in nitrous oxide decomposition on a Cu over-exchanged ZSM5 catalyst. The addition of 02 in the feed results in disappearing of the oscillations and decreasing of N20 average conversion. The experimental observations have been quite successfully modelled by a network of reactions in which N20 acts either as an oxidising either as a reducing agent for copper catalytic sites.

I. INTRODUCTION The interest in the control of N20 polluting emissions, mainly deriving from adipic acid synthesis and fluidized bed coal combustion, is related to the potemially disastrous effects on the environment, such as global warming and ozone layer depletion in the stratosphere. Many catalysts for N20 decomposition have been investigated such as mixed [ 1] and perovskite-type oxides [2], hydrotalcites [3] and metal exchanged zeolites [4-7]. Cu-ZSM5 is among the most promising catalysts for this process, as recently shown by Li and Armor [5]. Moreover, because of the possible role of N20 as an intermediate in NO decomposition on the same catalyst, the study of this reaction appears to be interesting. A specific feature of Cu-ZSM5 is the ability of introducing an oscillating behaviour in N20 decomposition. As we previously reported [8,9], when a gas mixture of nitrous oxide and He is fed to a flow reactor packed with Cu-ZSM5 catalyst, self-sustained isothermal oscillations of N20 outlet concentration may occur. In a parallel investigation, Lintz and Turek [10] found experimental evidences quite similar to our results, with a gas mixture of N20 and N2 fed to a Cu-ZSM5 catalyst with a different Si/AI ratio (37 instead of 80). Up to now, very few examples of oscillations of monomolecular reactions have been reported [12] whereas this specific behaviour is common to other heterogeneous catalytic systems, such as bimolecular oxidation reactions over metals, metal oxides and zeolites [13]. Oscillations of the outlet concemration of N20 over CuO/MgO catalyst were already observed by Hugo [11 ], but they were caused by reactor thermal instabilities, due to the different feed composition (pure N20).

176

As regards the mathematical modelling of the oscillating heterogeneous catalytic systems, some interesting results have been obtained for the reactions of CO oxidation over Pt [ 14], NO reduction with CO over Pt and H2 oxidation over metallic catalyst [ 13 ]. In this paper, we will compare the main results of our experimental investigation with the preliminary model simulations based on two reaction mechanisms both involving the redox chemistry of copper in Cu-ZSM5 catalyst.

2. EXPERIMENTAL ACTIVITY

2.1 Experimental apparatus The synthesis of H-ZSM5 (Si/A1 = 80) and the preparation of the over-exchanged CuZSM5 catalyst (Cu 3.94 %-wt.) have been elsewhere reported [ 15]. The laboratory fixed-bed microreactor consists of a 60 cm long quartz tube (I.D. 1.0 cm), in which a porous disk supports the catalytic bed. It is surrounded by an electrical furnace supplied with three heated and temperature controlled zones. The temperature is monitored by a Chromel-Alumel thermocouple, placed in another quartz tube, coaxial and internal to the reactor, along the whole length of the catalytic bed (1-2 cm). Four mass flow controllers (Brooks) allow to measure the flow rates of high purity gases: N20 (0.2% vol.) + He, NO (1% vol.) + He, He (99.995% vol.) and 02 (99.995% vol.). The reactor inlet and outlet gas have been analysed by two continuous analysers, to measure respectively NO, NO2 and NO2 (Hartmann & Braunn, URAS 10 E) and O2 (Hartmann & Braunn, MAGNOS 6 G) concentrations. 2.1 Experimental conditions The measurements of the catalytic activity of Cu-ZSM5 catalysts have been carried out after having pre-treated the catalyst in flowing He at 550~ for two hours in order to reduce copper sites in the zeolite from C& 2 to Cu- (pre-reduced catalyst) and by cooling the quartz microreactor down to the desired reaction temperature of 344~ Before the use, the catalyst has been pressed, crushed and sieved to obtain a 400-700 l.tm fraction. N20 inlet concentration to the catalytic reactor has been 300 or 600 pprn, the carder gas being He. The effect of the presence of 02 in the feed has been also investigated in the range from 100 to 5000 ppm. The catalytic bed has been packed with 1.43 g of Cu-ZSM5 catalyst, whereas a 37.5 N1/h total flow rate for the reacting gases has been used, corresponding to a residence time of 1.4-10.3 min. 2.3 Results and discussion Fig. 1 shows the values of N20 and 02 outlet concentrations as functions of time when a gas mixture containing N20 (600 ppm) and He is fed to the pre-reduced Cu-ZSM5 catalyst at 344~

177 400

30o

200

100 d

. ,0-

~

I

0

1

1

1

o

1

I

N20

0

o 300

0 200

100 9

0

I

I

I

1

I

I

200

400

600

800

1000

1200

time, min Figure 1. N20 and 02 outlet concentrations as functions of time after a step change of N20 concentration in the feed (from 0 to 600 ppm in He) on the pre-reduced catalyst at 344~

It can be observed that N20 decomposes and consequently 02 is stoichiometrically produced, but both N20 and 02 outlet concentrations oscillate in a quite regular way after about one hour from the start and during about one day of run. Moreover, the average values of N20 and 02 outlet concentrations have been respectively increasing and decreasing in the first 800 minutes and stabilise. The frequency of the oscillations of N20 outlet concentration is about 30 minutes, i.e. at least four order of magnitude greater than the residence time, while their maximum amplitude is about 100 ppm. It is worthwhile to point out that the catalytic bed temperature axially measured is always constant in the whole experiment (_ 0.1 ~ It has been also evaluated the effect of the presence of 02 on the behaviour of the catalytic activity in N20 decomposition at 344~ by carrying out experimental runs at different inlet concentrations of oxygen (up to 5000 ppm) in a gas feed mixture containing N20 (300 ppm) and balance He.

178 120 d

100

~

80

t2

o

60 9 z ~

40

20 600

1 700

I 800

I 900

I 1000

I 1100

I 1200

1300

time, mm Figure 2. N20 outlet concentration as a function of time at 0 2 inlet concentrations: (a) 0, (b) 100, (c) 1000 and (d) 5000 ppm. T = 344~ N20 inlet concentration -- 300 ppm.

Curve (d) of Fig. 2 shows that the addition of 5000 ppm of 02 makes the isothermal oscillations of N20 outlet concentration to disappear, while in the presence of 02 concentrations, at least up to 1000 ppm, N20 oscillations still exist, although their amplitudes and frequencies decrease with increasing 02 inlet concentration. Moreover, the addition of oxygen in the feed influences the N20 average conversion too. Fig. 2 shows that the timeaverage steady-state value of N20 outlet concentration increases from 45 ppm in the absence of 02 to 120 pprn, when 5000 ppm 02 are added to the feed. These results are different from those already presented by Lintz and Turek [ 10] and more recently confirmed by Turek [16]. They have reported that the addition of 02 up to a concentration of 12% vol. to the feed containing N20 (1000 ppm) in N2 slightly modifies the average reaction rate as well as the region of existence of the N20 outlet oscillations. These differences are likely due to the different Cu-ZSM5 catalyst and the different experimental conditions investigated.

179 3. M O D E L L I N G

3.1 Model equations In order to develop a kinetic model for the observed oscillations of N20 outlet concentration, the following assumptions have been made: (i) the reactor has been assumed as a PF1L i.e. the axial dispersion has not been taken into account; (ii) the external diffusion through the gas film around the catalyst particle and the internal pore diffusion resistance have been neglected; (iii) isothermal conditions have been considered because of the experimental observation and the very low amount of heat produced by the reaction in so high dilute conditions (300-600 ppm of N20 in He). As a consequence, a pseudo-homogeneous isothermal model has been assumed, consisting of mass balance equations on the gas phase and on the active centres. The model equations are the following: Mass balance on the gas phase: 63Ci ~:Ci a v ---~ + v ~ = Ri GZ

(1)

B.C.

for Z -" 0 ~ C i - C~

(1.a)

I.C.

for t = 0 ~ Ci = 0

(1.b)

where Ci is the actual concentration (moL/m3) of the i-th species (i = N20, N2, O2), t is the time (min), z is the reactor axial coordinate (m), a,. is the amount of catalytic active centres per unit volume of the reactor (moist/m3), e is the packed bed void fraction, v is the gas superficial velocity (m/min), Ri is the rate of production of the i-th component due to the reaction (mol/min molcat) and C~n is the inlet N20 concentration.

Mass balance on the solid phase: t~Sj

c3t = Rj I.C.

for t - 0 ~ 8j = ~jO

In particular, ,golcu§I = 1, while ~jo = 0, ~' j ~= [Cu +]

(2) (2.a)

180 where 8j is the molar fraction of j-th sites (i.e. adsorbed species or sites at different oxidation state) and Rj is the reaction of production of sites j (minl). Moreover, the mass balance on the solid is satisfied by the following expression: XjSj= 1

(3)

The mathematical model that derives from equations (1-3) has been solved numerically by means of routines available in the NAG library. The range of parameter values for which oscillations appear has been found by using the software AUTO. 3.2 Model simulation results We have already hypothesised a reaction mechanism based on the redox chemistry of the Cu-ZSM5 catalyst [9]. According to that scheme, we proposed that N20 isothermal oscillations are related to the two possible roles assumed by N20 molecule in its interaction with catalyst copper sites. Based on the well known evidence that in Cu-ZSM5 the copper sites can exist in two different oxidation states (Cu +2 and Cu § depending on the experimental conditions, this assumption expects that N20 could act both as an oxidant to transform Cu + to Cu =2 and as a reductant to promote the reverse process and attributes reaction rate oscillations to the changes of the oxidation state of these sites, according to the following reactions:

N20 + Cu" -+ N2 + [O"Cu .2]

(4)

N20 + [O-Cu +2] -9 N2 + 02 + Cu §

(5)

According to many authors [17-19], it has been assumed that in the oxidised state the catalyst contains an amount of extra-lattice oxygen (ELO) that is held bridged between Cu +2 ions (Cu+2-O-2-Cu+2). The ELO is easily desorbed by only raising the temperature resulting in the reduction of copper sites. Consequently, a reaction step between the adsorbate N20 and a free Cu + site is hypothesised following the N20 adsorption step and producing (Cu+2-O2-Cu § species. The mechanism is reported below: N20 + Cu +

k6 ; ['N20"-Cu +2]

I N 2 0 - C u +2] + Cu "~

k7 > [CR~2-O2-Cu .2] + N2

N20 + [Cu+2-O'2-Cu "2]

k8 > N2 + 02 + 2 Cu=

(6)

(7)

(8)

The reaction mechanism expressed by the steps (6-8) will be indicated as (MI). Its reaction rates have been modelled by supposing that all three reactions are irreversible and described by the mass action law. In particular:

181 R6 = k6.[N20][Cu']

(9)

kT.[N20-Cu+2][Cu -]

(~o)

R8 = ks-[-N20][Cu-2-O2-Cu "2]

(11)

R7 =

where Ri indicates the reaction rate of i-step. Some Hopfbifurcations have been detected with that kinetic model (9-11). Fig.3 reports the trend of N20 outlet concentration as a function of time, resulted from the model simulation performed with a set of kinetic constant, kl, k2, k3, which expects an oscillating behaviour. The boundary condition for N20 inlet concentration has been chosen equal to 300 ppm. 300 E .,.,

275

o 250

h

O U

. 225 O Z 1

I

1

1

E d

75

0 0

50

0

0 0

0

25

O

0 4000

I 4400

I 4800

I 5200

I 5600

6000

t/z

Figure 3. N20 and 02 outlet concentrations as function of time, as expected by M1 reaction mechanism. N20 inlet concentration: 300 ppm.

The results of the simulation carried out with M1 mechanism shows that it can successfully expect an oscillating behaviour. Fig.3 shows also that N20 and 02 outlet concentration

182 oscillations are expected in phase. This result is a consequence of the model assumption, confirmed by the experimental indications [9], according to which the reduced sites Cu" are the most active. In fact, if R6 is higher than Ks, when N20 reacts with the Cu" sites, it reaches its maximum conversion while O2 production is the lowest. In fact, oxygen production is maximum only when step 8 occurs preferentially. Nevertheless, this reaction mechanism cannot take into account the effect of oxygen addition to the feed which has been observed in Fig.2. In order to model these further experimental evidences, another reaction step has been considered and added to mechanism M1. This step, which was also reported in [9], describes the catalyst self-reduction by O~_ desorption, which is known to occur at high temperature on the Cu-ZSM5 catalyst. It is expressed by the following reaction step: 1

--" 02 + 2 Cu" 2

k12 > [Cu*2-O-2-Cu -2]

(12)

" k_12

In fact, when the net rate of step 12 is high enough to compete with R6 the catalyst reoxidation could not proceed via the step (6) and N20 may behave only as a reductant, making N20 oscillations to disappear. Of course, the possibility the rate of step (12) increases is related with a decrease of the reaction temperature and/or an increase of 02 concentration. The reaction mechanism expressed by the steps 6-8 and 12 will be indicated as (M2). The reaction rate of step 12 is given by the following relationship: R12 = k12"[O2]~

2-

k-12-[Cu+2-O2-Cu +2]

(13)

Fig. 4 shows that the effect of the presence of oxygen in the feed (N20 300 ppm), as expected from M2 mechanism. 300

g

iiiila

b

c

./.... ... ..

250

/]j

'~r176 ~ 200

"~..,j/: ".. . .i:: ',. .:.:,J

z o

150 0

I 200

I 400

I 600

t/~

Figure 4. Effect of O2 presence in the feed o n N 2 0 outlet concentration, as expected from M2 mechanism. Feed: N20 (300 ppm), 02 (a 0; b 2400 ppm; c 20000 ppm) and balance He.

183 Fig. 4 shows that the addition of 2400 ppm of oxygen to a feed containing only N20 (300 ppm) in He (balance) is not enough to make N20 outlet oscillations to disappear, although the N20 average conversion decreases. The oscillations amplitude and frequencies decrease as the inlet O2 concentration increases. Moreover, in the presence of an 02 concentration in the feed equal to 20000 ppm not only N20 average conversion decreases, but also N20 outlet concentration oscillations turns off.

4. CONCLUSIONS Isothermal oscillations of N20 outlet concentration in nitrous oxide decomposition over Cuexchanged ZSM5 catalyst have been experimentally observed. The effect of 02 addition to the feed results in disappearing of the oscillations (above 1000 ppm) and decreasing of N20 average conversion. Two reaction mechanisms have been proposed, both based on the two possible roles for N20 interacting with Cu sites either as reductant or as oxidant and succeeding in expecting an oscillating behaviour. The competition between N20 and 02 as catalyst oxidants has been considered in one of the two mechanisms and allowed to take into account the effect of oxygen addition to the feed.

NOMENCLATURE av Ci C'"i ki k12

k-12 Ri t v z

catalyst concentration gas concentration of component i gas inlet concentration of component i kinetic constant of step i (i = 6, 7, 8) kinetic constant of direct reaction 12 kinetic constant of reverse reaction 12 reaction rate of step i time superficial velocity axial coordinate

mol~t/m3 mol/m3 mol/m 3 m3/min mol~t mol ~ m~ mol~ mol/min moGt mol/min moGt min rn/min m

packed void fraction mole fraction of solid component i mole fraction of solid component i at t=0 contact time

~n

Greek letters

% 0~ "C

REFERENCES

1. 2. 3. 4.

J. Christopher and C. S. Swamy, J. Mol. Catal. 62 (1990) 69. C.S. Swamy and J. Christopher, Catal. Rev.-Sci.Eng., 34 (1992) 409. S. Kannan and C.S. Swamy, Appl. Cat. B.; 3 (1994) 109. C.M. Fu, V. N. Korchak and W. K. Hall, J. Catal. 68 (1981) 166.

184 5. 6. 7.

10. 11. 12. 13. 14.

15. 16. 17. 18. 19.

Y. Li and J. N. Armor, Appl. Catal. B 1 (1992) L21. V.I. Sobolev, G. I. Panov, A. S. Kharitonov, V. N. Romannikov, A. M. Volodin and K. G.Ione, J. Catal. 139 (1993) 435. Y.F. Chang and J. G. McCarty, E.D. Wachsman and V.L. Wong, Appl. Catal. B, 4 (1994) 283. E. Garufi, R. Pirone and F. Santagata, Proc. 1995 Annual Meeting of Italian Section of Combustion Institute, CUEN, Napoli (1995) V-20. P. Ciambelli, E. Garufi, R. Pirone, G. Russo and F. Santagata, Appl. Catal. B 8 (1996) 333. H.-G. Lintz and T. Turek, Catal. Lett. 30 (1995) 313. P. Hugo, Proc. 4th Eur. Symp. React. Eng., Brussels 1968, Pergamon, Oxford, 459,1971. S.Y. Hwang and L.D. Schmidt, J. Catal., 114 (1988) 230. M.M. Slin'ko and N.I. Jaeger; Oscillating heterogeneous catalytic systems; Elsevier Science, Amsterdarm, 1994. S.S.E.H. Elnashaie and S.S. Elshini, Dynamic modelling, bifurcation and chaotic behaviour of gas-solid catalytic reactors; Gordon and Breach Publishers, Amsterdarm, 1996. P. Ciambelli, P. Corbo, M. Gambino, G. Minelli, G. Moretti and P. Porta, Catal. Today, 26(1995)33. T. Turek, Appl. Cat. B 9 (1996) 201. E. Giamello, D. Murphy, G. Magnacca, C. Morterra, Y. Shioya, T. Nomura and M. Anpo, J. Catal. 136 (1992) 510. T. Cheung, S.K. Bhargava, M. Hobday and K. Foger, J. Catal., 158 (1996) 301. J. Valyon and W.K. Hall, New Frontiers in Catalysis, Ed. by L. Guczi, F. Solymosi, P. Tetenyi, Elsevier, Amsterdam, Part B, (1993) 1339.

91997 Elsevier Science B. V. All rights reserved. Dynamics of Surfaces and Reaction Kinetics in Heterogeneous Catalysis G.F. Froment and K.C. Waugh, editors

185

Adsorption characteristics of pyridine bases on zeolite(010) examined by atomic force microscopy (AFM) Masaharu Komiyama Department of Chemistry, Yamanashi University. Takeda, Kofu 400 JAPAN

Liquid-phase adsorption characteristics examined by atomic force microscopy (AFM) were compared for two pyridine base molecules, pyridine and/3-picoline, on (010) surfaces of two natural zeolites, heulandite and stilbite. These adsorption systems formed well-ordered, two-dimensional (quasi-)hexagonal adlayers. The 2D lattice structures of the ordered adlayers were dependent on the adsorbate/substrate combinations. Although there existed certain habit in the orientation of the 2D lattice unit vector of the adsorbed phase with respect to the substrate(010) lattice vectors, the molecular arrays were incommensurate with the substrate atomic arrangements.

1. I N T R O D U C T I O N One of the striking capability of scanning probe microscopy such as scanning tunneling microscopy (STM) and atomic force microscopy (AFM) is the direct observation of adsorbed molecule images under various environments such as vacuum, ambient and underwater. Early examples include the observation of benzene rings adsorbed on Rh single crystal surface with STM under ultrahigh vacuum conditions [1]. The resolution of AFM reached that level only recently, and the reports on AFM observations of adsorbed species are scarce. Specifically on zeolite surfaces, only one attempt is known to the authors on the AFM observation of adsorbed molecules. Weisenhorn et al. [2] reported the formation of ordered arr~" of t e r t - b u t a n o l on clinoptilolite(010). In the work the inner structure of the adsorbed molecule was not resolved, and each molecule appeared as a single mass in the AFM images. Recently we successfully obtained i n s i t u molecular AFM images of pyridine base species, pyridine and 3-picoline, adsorbed on cleaved (010) surfaces of natural zeolites, stilbite and heulandite [3-5]. These adsorption systems possessed three adsorption phases: one physically adsorbed, and two chemically adsorbed. One of the latter two adsorption phases consists of monolayer of molecules randomly adsorbed, and the other formed a well-ordered (quasi-)hexagonal array. The present paper compares the adsorption characteristics of these adsorption systems in terms of the array and orientation structure of the adsorbed molecules as determined, for the first time, by AFM. Pyridine adsorption on metal surfaces under vacuum conditions has been studied in detail, in relation to the surface enhanced Raman scattering (SERS) phenomenon. On

186 Ag(111) at 100 K a near-edge x-ray-absorption fine structure (NEXAFS) stud)" revealed that the tilt angle of pyridine ring from the surface plane is 45+5 ~ at low coverages and changes to 70+5 ~ at submonolayer coverages [6]. On Pt(111) NEXAFS showed that below 300 K pyridine at saturation coverage is tilted 52+6 ~ to the surface and above that temperature the angle is 74+10 ~ [7]. Under aqueous environment, there exist only one work on the determination of the adsorbed layer structure of pyridine [8]. Stern et al. adsorbed pyridine on a Pt(111) electrode surface from an aqueous solution under various applied potentials, and examined the adsorbed surface ez s i t u with Auger, LEED and EELS. From the packing density of pyridine obtained by Auger, a tilt angle of 71 ~ to the surface was obtained. A LEED analysis showed that the adsorbed pyridine lattice is incommensurate with the Pt surface and oblique, with lattice vector lengths of 0.332 and 0.474 nm and an inclined angle of 77 ~. In the field of heterogeneous catalysis, pyridine is frequently employed as a molecule probe of the surface acid sites of zeolites and other oxide catalysts, with which the amount and the strength of the acid sites are determined, for instance by infrared measurements of the intensity and the peak position of pyridine adsorbed on zeolite surfaces [9]. Despite its importance in this respect, details of the structure and the molecule orientation within the adsorbed pyridine layer on these nonconductive substrates are not known, due to the lack of techniques to yield such data before the advent of AFM. The series of our present work constitutes the first report on the determination, by means of AFM. of the array structure of the pyridine base adlayers on zeolite surfaces, and the estimation of the molecular orientation within the adsorbed laver.

2. E X P E R I M E N T A L Natural zeolites employed as substrates were stilbite from Bear Brook and heulandite from Ross Creek, both in Nova Scotia. Canada. Their nominal compositions are Na4Cas[A120Si520144]'56H20 for stilbite and Ca4[AlsSi2sO72].24H20 for heulandite. Stilbite and heulandite crystals were cleaved along their (010) fault planes and placed in a sealed AFM liquid cell, which is then filled with membrane-filtered deionized water. The atomic images of the surfaces were first obtained underwater with a contact-mode AFM (.NanoScope II, Digital Instruments) at a typical tip load of ca. 5 nN. After images under pure water were taken, the liquid content in the cell is substituted with a 1 vol% aqueous pyridine base solution for successive AFM scanning on the surfaces.

3. R E S U L T S

AND

DISCUSSION

3.1. A t o m i c a r r a n g e m e n t o n c l e a v e d z e o l i t e ( 0 1 0 ) s u r f a c e s The cleaved zeolite surfaces were extremely fiat. The underwater AFM examinations [3,4] showed terraces more than few hundred nanometers wide, each separated by a ca. 1 nm step, a height corresponding to the lattice spacing of heulandite or stilbite (020) which is 0.90 and 0.91 nm. respectively. When these surfaces were examined at higher magnifications, AFM images shown in Figure 1 are produced. Heulandite(010) (or (020)) surface (Figure l(a)) gives atomic images almost completely corresponding

187

(a)

Figure 1. Underwater atomic level AFM images of (a) heulandite(010) and (b) stilbitr Cr~" ~ 1 ~ is 0.25 . m f~11 sr162 R~p~od~r from r~f~. [3] ~ . a [4]. to the bulk-terminated (010) surface [3]" it shows a oblique 2D unit cell as indicated in the figure, with the unit cell lattice parameters of a - 1.8 nm, c = 0.7 nm and = 116 ~ which exactly correspond to those of the bulk-terminated surface of 1.77 m-n, 0.74 nm and 116 ~ respectively. On the a axis two oxygen atoms per unit cell protrude from the surface. In this underwater AFM image there exist periodic structures that do not correspond to those outermost oxygen atoms, and they are a t t r i b u t e d to the water molecules securely located in the 2D lattice. Stilbite(010). on the other hand, gave somewhat fuzzier atomic level AFM images, an example of which is shown in Figure l(b). It does show ridges with a spacing of 1.1 nm which correspond to that of the topmost oxygen a t o m rows along the c axis as indicated in the figure. The reason for the lower resolution for stilbite t h a n the case of heulandite is not apparent at the moment. 3.2. Pyridine on h e u l a n d i t e ( 0 1 0 ) and s t i l b i t e ( 0 1 0 ) s u r f a c e s The addition of pyridine to the AFM cell drastically changes the surface image of zeolite(010) from its bulk-terminated structure. Figure 2 compares the narrow-area scanning of the pyridine-adsorbed heulandite and stilbite (010) surfaces. It is apparent that the new lattice structure consists of hexagonal arrays of ring structures, each ring size corresponding to that of a pyridine molecule as indicated in the figure. A few observations may be made on the AFM images shown in Figure 2. First, for the first time we are able to determine the array structure of the adsorbed pyridine laver on a zeolite surface from the image. With the images given in Figure 2 and others that are not shown, it was found that pyridine on heulandite(010) forms an almost perfect hexagonal lattice, with its unit cell length of 0.55 rim. On stilbite(010) the ordering is somewhat loose, giving a quasi-hexagonal lattice with a unit cell dimension of 0.51• rim.

While pyridine molecules appear to form a well-ordered s t r u c t u r e on this surface, comparison of the array structure in the adsorbed phase (Figure 2) with that of the

188

(a)

(b)

Figure 2. AFM images of pyridine molecules adsorbed on (a) heulandite(010) and (b) stilbite(010). Gray scales are (a) 0.3 nm and (b) 0.4 nm full scale. Reproduced from refs. [3] and [4]. substrate surface (Figure 1) indicates that the arrangement of adsorbed molecules does not match with the periodic structure of the surface unit cell of either zeolite. This indicates that the pyridine molecules are not adsorbed on specific sites on the surface, but are arranged in a manner similar to the so-called self assembly. Nevertheless, it is noteworthy that the substrate has certain influences over the array structure of this "self-assembled" pyridine phase as embodied in the difference of the 2D lattice unit length. The second observation on the images shown in Figure 2 is that, while the pyridine molecules form a well-ordered (quasi-)hexagonal phase, they appear to be adsorbed on the surface with their molecular axes tilted away from the surface parallel. This is inferred from the observation that the images show only half of each ring clearly,which indicates that the rings are tilted away from surface parallel so that the lower part is not visible with the A F M tip. This inference is further supported by the fact that it is not geometrically possible to arrange pyridine molecules parallel to the surface with the spacing observed by A F M , as may be apparent with the overlaid pyridine skeleton models in Figure 2. A F M cross sectional views also support this inference. A n attempt to determine this tilt angle was done with a help of computer graphics, since A F M observation alone does not yield the information of height variation of the pyridine rings due to the finite size of the tip apex. For this purpose, we first arranged pyridine molecules upright on a surface, with the separations among them set to the values observed by A F M . The molecules are arranged so that every part of a pyridine molecule keeps m a x i m u m distance from other molecules. Then the molecules are tilted toward the surface up to the point where the density of the adsorbed phase equals that of the bulk pyridine solid. This results in tilt angles of 43 ~ and 54 ~ to the surface for heulandite and stilbite,respectively. It is noted that these estimated tiltangles are only the first approximation. However. it is also noted that these values are close to ones observed on pyridine adsorbed on metal

189 surfaces under vacuum conditions, particularly to the low-coverage value on Ag(111) (45+5 ~ [6] or low-temperature value on P t ( l l l ) (52=i=6~ [7].

3.3. 3-Picoline on stilbite(010) The adsorption of/3-picoline on stilbite(010) was characteristically slow. enabling us to observe its adsorption process. Upon contact of the surface with a 13-picoline solution, a fuzzy; cloud-like surface is observed, which is attributed to multi-layer physisorbed ~-picoline molecules. Scraping off the physisorbed layer with the AFM tip revealed the presence of two chemisorbed monolayer phases" one randomly adsorbed and the other well-ordered two-dimensional quasi-hexagonal phase. The development of the latter phase was very slow on this surface, allowing us to estimate its rate to be ca. 6 nm/min. The slow ordering process also allowed us to determine the interrelation between the adlayer and the substrate lattices. It is found that one of the unit vector of the adlayer makes an angle of 90 ~ with respect to the c axis of the substrate(010) plane. The ordered phase is incommensurate with the substrate surface as in the case of pyridine/zeolite(010). Figure 3 shows an AFM image of/3-picoline molecules adsorbed on stilbite(010). From the obtained images the 2D unit cell dimension of the ordered adlayer phase was calculated to be 0.55+0.04 nm. Molecular orientations were also determined using computer graphics as described above. 3.4. C o m p a r i s o n of the three adsorption s y s t e m s Table 1 compares the molecular array and orientation parameters thus obtained for the three adsorption systems examined here. All the three adlayers examined here appears to form a well-ordered (quasi-)hexagonal array on the two zeolite surfaces. Since the surface symmetry of either zeolite (010) surfaces are not hexagonal, this hexagonal

Figure 3. AFM image of i3-picoline molecules adsorbed on stilbite(010). Gray scale is 0.5 nm full scale.

190 Table 1 Structural parameters for the ordered pyridine base adlavers on two zeolite surfaces Unit cell length Pyridine/heulandite Pyridine/stilbite

3-Picoline/stilbite

0.55 nm 0.51+0.03 nm 0.554-0.04 nm

Tilt angle to surface Tilt angle to 2D vector 43 ~ 54 ~ 47 ~

20 ~ 23 ~ 20 ~

order habit of the pyridine base molecules may derive from the nature of the adsorbates. This inference is also supported by the fact that the molecular array is incommensurate with the substrate surface in all the three adsorption systems. Nevertheless. it appears that substrate structure has certain influence over the array structure of the adlaver (cf. Table 1). For instance, pyridine on stilbite gives smaller unit cell length than that on heulandite. Here we would like to note that in the case of pyridine/stilbite system, the substrate surface unit cell dimensions (a = 1.36 nm and c = 1.13 nm) are within 10 % of integer multiples of the dimensions of the adsorbed laver unit cell (0.51-1-0.03 nm) and the angle of the substrate unit cell (128 ~ is also similar to that of the adsorbed layer (ca. 120 ~ The proximity of these lattice parameters may induce a stronger interaction between the two lattices, causing the fluctuations of the pyridine adlaver lattice parameters on stilbite, and also the smaller value for the unit cell length than that over heulandite. When 3-picoline is adsorbed on stilbite, it gives larger unit cell length than pyridine on stilbite as found in Table 1. This is apparently due to the bulkiness of the picoline molecule compared to the pyridine molecule. Thus again interaction among the molecules in the adlaver appears to dominate in determining the array structure of ii. qDr ~ the ordered phase, as we refer to it a "setf-assemalm o. Nevertheless, it is again noted that the stilbite surface unit cell dimensions are within 10 % of integer multiples of the dimensions of the adsorbed picoline layer unit cell, which may have certain influences on the adlayer unit cell dimensions. As for the tilt angles to the surface, we note again that the values listed in Table 1 are the estimates derived assuming solid density in the adlavers. Although this may be a good first approximation for such a close-packed system, it is by no means conclusive. Since AFM measurements by themselves cannot provide absolute values for the height variations as discussed above, experimental determination of the tilt angles my be difficult. Theoretical examinations by means of computer simulations are in progress.

4. C O N C L U S I O N S Adsorption on solid surfaces is an i m p o r t a n t elemental step that leads to various chemical processes including self assembly, catalysis and separation. For this reason the determination of molecular orientation and intrastructure in the adsorbed phase have been the target of numerous studies in the field of surface science. Large body of these works, however, has been limited to metal or semiconductor substrates placed under vacuum conditions, mainly due to the requirements posed by the techniques employed to obtain such information. The present work demonstrates that with AFM similar information can be obtained on nonconductive surfaces, even under liquid-phase

191 conditions. Such knowledge will expand our understanding on adsorption phenomena beyond the limitations of the older analytical techniques.

REFERENCES 1. H. Ohtani. R.J. Wilson. S. Chiang and C.M. Mate. Phys. Rev. Left.,, 60 (1988) 2398. 2. A.L. Weisenhorn. J.E. MacDougall, S.A.C. Gould. S.D. Cox. W.S. Wise. J. Massie. P. Maivald. V.B. Elings, G.D. Stucky and P.K. Hansma. Science. 247 (1990) 1330. 3. M. Komiyama, T. Koyama. T. Shimaguchi and M. Gu. J. Phys. Chem., 100 (1996) 15198. 4. M. Komiyama and M. Gu, ,]pn. J. AppI. Phys., 35 (1996) 3775. 5. M. Komivama and M. Gu. J. Vac. Sci. Technol.. submitted. 6. M. Bader. J. Haase. K.-H. Frank. A. Puschmann. A. Otto. Phys. Rev. Lett. 56 (1986) 1921. 7. A.L. Johnson. E.L. Muetterties. J. Stohr. F. Sette. J. Phys. Chem. 89 (1985) 4071. 8. D.A. Stern, L. Laguren-Davidson. D.G. Frank. J.Y. Gui. C.-H. Lin, F. Lu. G.X. Salaita. N. Walton. D.C. Zapien. A.T. Hubbard, J. Am. Chem. Soc. 111 (1989) 877. 9. J.W. Ward, In ~Zeolite Chemistry and Catalysis." J.A. Rabo (ed.), American Chemical Society: Washington, D.C.. 1976" Chapter 3.

This Page Intentionally Left Blank

91997 Elsevier Science B. V. All rights reserved. Dynamics of Surfaces and Reaction Kinetics in Heterogeneous Catalysis G.F. Froment and K.C. Waugh, editors

193

Transient and Steady-State Studies of the Effect of Water on Cobalt Fischer-Tropsch Catalysts K. F. Hanssen, E. A. Blekkan, D. Schanke '~and A. Holmen Department of Industrial Chemistry, Norwegian University of Science and Technology (NTNU), N-7034 Trondheim, Norway "SINTEF Applied Chemistry, N-7034 Trondheim, Norway"

ABSTRACT The effect of water on the Fischer-Tropsch synthesis over alumina-supported cobalt catalysts has been studied using isotopic transient kinetic methods (SSITKA) in combination with steady-state measurements. Water has been introduced to the catalytic system as a pretreatment procedure as well as u n d e r reaction conditions. The SSITKA results showed a decrease in the number of active surface sites, but no change in the specific site activity.

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

Recent developments in Fischer-Tropsch (FT) technology are focused on the production of higher molecular straight chain waxes that in t u r n can be hydrocracked to products in the middle distillate range. Supported cobalt catalysts promoted by small amounts of a second metal (typically a noble metal) have been shown to be good FT catalysts. The role of t h e second metal is to improve the reducibility of cobalt oxides on the support surface and in certain cases also to aid in keeping the surface clean and increase the coverage of reactive intermediates [1]. We have recently studied the effect of water on the cobalt FT catalysts [2]. Water is the major oxygen containing product in Fischer-Tropsch synthesis (FTS), and in certain reactor types (e.g. backmixed slurry reactors) the concentration of water will be high and this can influence the catalyst surface. Water can deactivate the catalyst, and we have shown that this is mainly due to a surface reoxidation of cobalt. " Present address: Statoil R&D Center, Postuttak, N-7005 Trondheim, Norway

194 Schulz et al. [3] found no changes in the activity of a Co/Zr/aerosil catalyst even at high water pressures but reported a decreasing selectivity to m e t h a n e S(CH4), an increasing S(C~§ and a higher olefin content in the products. Iglesia and co-workers [4] carried out experiments with water and ethylene addition to a CO/H 2 feed over Ru and Co/TiO 2 catalysts and proposed that water inhibits the secondary hydrogenation, especially of cz-olefins. Other results from the literature regarding the effect of water indicate an increasing chain-growth probability [3-5] which can be explained by the inhibiting effect of water on the termination of growing chains by hydrogen addition which is the predominant irreversible termination step. The changes of the properties of Co catalysts in the presence of water have to be taken into account especially when the FTS is carried out in a slurry reactor where the concentration of w a t e r (as the main oxygen-containing product over Co catalysts) will be high throughout the whole reactor due to extensive backmixing. In the present communication we report on the influence of water on the FT synthesis studied by SSITKA and conventional kinetic experiments. Steady-state isotopic t r a n s i e n t kinetic analysis (SSITKA) has proved to be a powerful technique for this work. The technique involves switching between 12CO and 13CO in the feed gas and analyzing the transients with respect to the formation of products containing 12C and ~3C. This technique allows the determination of the true turnover frequency of the active site, decoupled from site coverage. Applied to the FTS over metal promoted cobalt catalysts SSITKA has shown that the true turnover frequency of cobalt always remains the same, regardless of the second metal [6-8].

2. EXPERIMENTAL 2.1. Catalyst preparation Co(17.1 wt.%)/A1203 (88ET=138 m2/g) and Co(17.2 wt.%)/Re(1 wt.%)A1203 (SBzT=138 m2/g) were prepared by standard incipient wetness techniques using either Co(NO3) 2 or co-impregnation of Co(NO3) 2 and HReO 4 on the A1203 (PURALOX, 174 m2/g) support. The catalysts were dried overnight at 393 K followed by calcination in air for 2 h at 573 K. The dispersions were measured by hydrogen chemisorption (assuming a H:Co ratio of 1:1) to be 4.9% (Co) and 9.3 % (Co/Re).

2.2. High pressure steady-state experiments The high pressure steady-state kinetic experiments were conducted in a 10 mm id. stainless steel tubular reactor. 1 g of catalyst (38-75 gm particle size) was diluted with an inert material (non-porous SiC, 75-150 gm) in 1:5 weight ratio to minimize t e m p e r a t u r e gradients. The catalyst was reduced in flowing hydrogen (1 K/min to 623 K, hold time 16 h). After reduction, the catalyst was cooled to 453 K in flowing hydrogen and purged with He before increasing the pressure to 13

195 bar and switching to a feed mixture containing 50 mol% synthesis gas with a H J C O ratio of 2.1 and balance inert gas (N2+He). The reaction t e m p e r a t u r e was t h e n slowly increased to 483 K. After 24 h on stream, w a t e r was introduced to the system by replacing He by an equal flow of w a t e r vapor. The steam was generated by feeding w a t e r by a liquid flow controller (Hi-Tec) to a vaporizer kept at ca. 573 K. The s t e a m was mixed with the synthesis gas j u s t before the reactor inlet. Finally, after 24 h of steam-containing feed, the experiment was t e r m i n a t e d by u s i n g dry feed for a n o t h e r 24 h. Similar e x p e r i m e n t s were performed, where the catalyst w e n t t h r o u g h a w a t e r p r e t r e a t m e n t procedure prior to the Fischer-Tropsch synthesis (directly after reduction). The pressure was increased to 10 bar and the t e m p e r a t u r e to 523 K, before introducing 248 ml/min steam mixed with 24.8 m l / m i n H 2 a n d 177.2 m l / m i n He. On-line GC samples were t a k e n at hourly intervals and analyzed for N 2, CO, CO.~ and CI-C S hydrocarbons on a HP 5890 gas chromatograph equipped with t h e r m a l conductivity and flame ionization detectors. The space velocity was varied to give approx. 10% CO conversion. For further details on the experimental setup, see Hilmen [9].

2.3. L o w p r e s s u r e e x p e r i m e n t s (steady-state and t r a n s i e n t (SSITKA)) Similar e x p e r i m e n t s were performed in a dedicated SSITKA a p p a r a t u s (Figure 1). 100 mg of catalyst (38-75 ~tm particle size) was diluted w i t h non porous SiC (58-75 ~tm) in 1:2 weight ratio. The catalyst was reduced in flowing hydrogen (1 K/min to 623 K, hold time 16 h). After reduction, the catalyst was cooled to 453 K in flowing He before increasing the pressure to 1.8 bar and switching to a feed mixture containing 22 mol% synthesis gas w i t h a H_JCO ratio of 10 and balance inert gas (0.08 mol% Ar + 78 mol% He). The reaction t e m p e r a t u r e was then slowly increased to 483 K. W h e n steady-state conditions were established (after 2-4 hours), a switch from ~2CO to ~3CO in the feed s t r e a m was carried out, followed by the opposite switch 5 m i n u t e s later (see Figure 4). The t r a n s i e n t concentration changes d u r i n g the switches were monitored on-line by a Balzers QMG 420 quadrupole mass spectrometer. The catalytic activity was determined using a H P 5880 GC equipped with a G S - a l u m i n a column (J&W Scientific) connected to a FID detector. After recording the t r a n s i e n t responses, water was introduced to the reactor by replacing a p p r o x i m a t e l y 50% of the He by vapor. Water was vaporized t h r o u g h h e a t e d lines via a peristaltic p u m p and mixed with the feed at the reactor inlet. All lines from the p u m p to the reactor outlet as well as the d o w n s t r e a m lines were heated to 400 K to prevent condensation of water. Due to e x p e r i m e n t a l limitations, it was impossible to record t r a n s i e n t s during w a t e r t r e a t m e n t . The vapor feed was turned off after 16 h and new isotope switches were carried out after 2-3 hours of FTS under dry conditions.

196

OXYTRAP

< P

s

>

(~

MOLECULA?. S ~ V E

IV~C

M A S S ~ _ O W CONTROLL,ZP.

PC

PI~%EE ~JRE- C O N T R O L L E R

MS

M A S S SPECTKO~I'-=/~

C/C

GAS CZ'Z{OMATOGRAPH

I

~ D

J PC!

....

<}C MS

L]gTs

TO VENT

b

L, J;

?UiV~ED CAPIL/U~Y

TO VENT

>

'~IT~C "i ,iV2C~ i 'i M F C I 'M/~Ci i

:'~,

I

i

@

<[

~,~c!

@

JZ,

'3CO~

IHol

• WA~i~ FN

mr

~CT'O~

WKI~=R OUT

Figure 1. Survey of the isotopic transient kinetic apparatus.

As for the abovementioned steady-state experiments, SSITKA experiments were carried out on catalysts pretreated with water vapor. After reduction, the reactor was cooled down to 453 K in flowing He. The temperature was increased to 523 K before introducing a flow of 1.2 g/h of water vapor, mixed with 2.5 ml/min H 2 and 35 ml/min He. After 16 h of water treatment, the catalyst was cooled down to 453 K in flowing He. The further procedure was the same as for the dry catalyst. For further details on the experimental set-up, see Bari~s [10]. 2.4. D a t a t r e a t m e n t The steady-state reaction rate for kinetic experiments is determined as the product of the rate constant (k) and the concentration of active surface intermediates (N): R=k-N

(1)

197 If no readsorption occurs, the 1. order rate constant is equal to the inverse of the average surface residence time (z): k = 1/x

(2)

t h u s giving R = N/z

(3)

In s t e a d y - s t a t e isotopic t r a n s i e n t kinetics, the average surface residence time can be calculated as the area u n d e r the normalized t r a n s i e n t curve [11-13]. E q u a t i o n (3) gives the concentration of active surface intermediates, N. For each experiment, two switches were performed (Figure 4), and by using both the labeled and the unlabelled transients, four parallel values of z and N were calculated. The results are reported as m e a n values and s t a n d a r d deviations.

3. R E S U L T S A N D D I S C U S S I O N

3.1 Activity and selectivity Figure 2 shows the development in the hydrocarbon formation rate as a function of time on s t r e a m for the experiments at elevated pressure. The vertical dotted lines m a r k s the s t a r t and stop of water/syngas co-feeding. The gas hourly space velocity (GHSV) is adjusted after the co-feeding period to give the same level of CO conversion as prior to the introduction of water. It is seen t h a t the promoted c a t a l y s t is significantly more active compared to the u n p r o m o t e d one w h e n no w a t e r p r e t r e a t m e n t has t a k e n place, but this difference is absent on the w a t e r t r e a t e d catalysts. It is proposed t h a t r h e n i u m supports the reduction of cobalt, but the w a t e r p r e t r e a t m e n t leads to a reoxidation of the active metal. The chain length distributions for the different experiments are shown in Figure 3. For the promoted catalyst, the w a t e r t r e a t m e n t procedures do not lead to any significant differences in the chain length distribution. This is not the case for the u n p r o m o t e d catalyst; w a t e r t r e a t m e n t results in a shift towards a higher selectivity to C3-C 6 compounds. In Table 1, the C1-C4 selectivities are listed including the separate values for nparaffins a n d a-olefins. The corresponding values for the atmospheric e x p e r i m e n t s (SSITKA-apparatus) are shown in Table 2. It should be noted t h a t the CO-conversion varies in the range from 10% to 20% for the high p r e s s u r e experiments, a n d from 5% to 15% for the atmospheric experiments. However, for each pair of e x p e r i m e n t s the variation in conversion is small. The t r e n d in the odn-ratio is clear: w a t e r t r e a t m e n t before reaction or under reaction decreases the olefin to paraffin ratio.

198 0.6 o

.~ r

o Co/Re z~Co 9 Co/Re (pretreated with water) 9 Co (pretreated with water)

~176

0.50.4-

o.3-

A

0.20.1-

Water addition i

Dry feed 0

'

0

i

24

48 Time

Dry feed |

! 72

(h)

Figure 2. Observed hydrocarbon reaction rate as a function of time for both catalysts (Co/Re and Co on A1203 support); the effect of water pretreatment and water/syngas co-feeding. T=483 K, P=13 bar.

15

15 Co/A1203

10

~10

-

5 r~

I

1

I

I

0

I

2 3 4 5 Carbon number

6

1

2 3 4 5 Carbon number

6

Figure 3. Comparison of the chain length distribution in the hydrocarbon products for the promoted and unpromoted catalyst. T=483 K, P=13 bar. CO conversion is given in Table 1. Solid symbols: catalysts pretreated with water. Open symbols: FTS directly after reduction (no water pretreatment). Square symbols: before syngas/water co-feeding. Circular symbols: after syngas/water cofeeding.

199

Table

1.

Comparison conversion

of the of syngas

CO-conversion,

X(CO),

and

the

C~-C4 s e l e c t i v i t i e s

over Co/A]20 ~ and Co/Re/A1203. P-13

Catalyst/

x(co)

s(c)

treatment

(%)

(%)

Before H~O

16.8

20.2

After HoO

15.9

s(c~) n-pa

~-ol

bar, T-483

s(c~) ZC~

in

the

K. s(c,)

odn

n-pa

~-ol

ZC~

odn

n-pa

r

ZC.; a/n

1.6 0.12

1.7 0.07

1.4

2.0

3.4

1.4

1.4

2.0

4.1

1.5

24.0

2.7 0.16

2.9 0.06

2.3

2.5

4.8

1.1

2.0

2.4

5.3

1.2

9.7

9.0

0.8 0.39

1.2 0.49

0.7

3.2

3.9

4.7

1.0

3.3

4.7

3.2

11.6

12.4

1.4 0.24

1.7 0.17

1.3

3.4

4.8

2.6

1.9

3.4

6.1

1.8

Co/AI~O~ Not pretreated

Pretreated Before H20 After H20

Co/Re/AI~O~ Not pretreated Before H~O

22.1

10.9

1.0 0.09

1.1 0.09

1.2

2.2

3.4

1.9

1.8

2.0

4.7

1.1

After H20

18.9

14.1

1.4 0.11

1.5 0.08

1.8

2.4

4.3

1.3

2.5

2.2

5.7

0.9

Before H~O

10.4

12.0

1.2 0.35

1.5 0.30

1.1

3.4

4.5

3.0

1.5

3.5

5.7

2.3

After H20

12.6

11.5

1.1 0.28

1.4 0.26

1.1

2.9

4.0

2.7

1.5

3.0

5.2

2.0

and

the

C~-C,

selectivities

in

the

Pretreated

Table

2.

Comparison conversion

of the of syngas

CO-conversion,

X(CO),

o v e r C o / A 1 2 0 3 a n d C o / R e / A 1 2 0 3. P = I b a r , T = 4 8 3

Catalyst/

X(CO)

S(C)

S(C 2)

treatment

(%)

(%)

n-pa

(z-ol ZC~

S(C3) odn

n-pa

a-ol

K. S(C4)

EC3

odn

n-pa

a-ol

ZC~

odn

Co/AI203 Not pretreated Before H20

7.4

53.6

11.1

9.7

3.3 12.9 0.34

2.2

3.2

11.6

1.5

After H20

5.7

47.1

9.4

4.3

6.2 17.3 0.56 13.7 0.46

14.8

3.5 18.3 0.24

2.4

3.7

14.2

1.5

Before H20

5.7

42.1

8.4

3.6 12.0 0.43

18.0

3.3 21.3 0.18

4.8

4.5

18.1

0.9

After H~O

5.5

44.0

10.1

3.6 13.7 0.36

17.8

3.6 21.4 0.20

4.2

3.6

14.9

0.9

Before H.,O

15.5

55.8

12.5

9.0 21.5 0.72

4.1

4.4

1.1

1.9

2.3

9.4

1.2

After H20

10.8

50.8

11.8

7.0 18.8 0.59

8.7

4.1 12.7 0.47

2.0

3.1

11.8

1.6

Before H20

9.8

50.4

12.2

7.4 19.6 0.61

9.1

4.0 13.1 0.44

2.4

3.2

12.4

1.3

After HoO

8.9

49.6

11.0

6.4 17.4 0.58

10.7

4.0 14.6 0.37

2.4

3.4

13.2

1.4

Pretreated

Co/Re/AI203 Not pretreated 8.5

Pretreated

200 3.2 SSITKA

results

In Figure 4, the normalized transient responses for the masses 1 5 , 17 and 40 (12CH4, '3CH~ and Ar, respectively) are shown for an experiment using the Co/Re/A1203 catalyst. The relatively smooth curves make the calculation of the areas under the curves, and thus the surface residence times, straightforward. The two CO transients are omitted from the figure for simplification. Table 3 shows the experimentally obtained values for the CO conversion, the methane selectivity, the number of active surface species reacting to methane and the corresponding surface residence time for all of the SSITKA experiments. The CO conversion and the methane selectivity show the same tendency as for the high pressure experiments. Water treatment reduces both the conversion and the selectivity to methane, either if the water is introduced as a p r e t r e a t m e n t procedure, or co-fed with the reaction mixture. Catalysts pretreated with water show no significant difference in behavior before and after the co-feeding. A comparison of the surface residence times, x, reveals no significant difference for the experiments. From this observation, it can be concluded that the variation in methane yield is due to a difference in the amount of active sites. The intrinsic site activity for the methanation reaction remains unchanged after the water treatment, and it is also unaffected by the presence of the rhenium promotor. It is proposed that rhenium supports the reduction of cobalt, but the water pretreatment leads to a reoxidation of the active metal, particularly the fraction of the cobalt metal that is reduced only when rhenium is present. 1.2

i A

0.8 0.6

o

M=15

9 M=17

0.4

M=40

0.2

-0.2 0

100

200

300

400

500

T i m e (s)

Figure 4. Normalized isotopic transients for '~CH4 (M=15), 13CH4 (M=17) and Ar (M=40) corresponding to a switch from 1~CO to 13CO and vice versa (Catalyst:

Co]Re/A1203).

201 Table 3. Influence of w a t e r t r e a t m e n t on CO conversion, X(CO), the selectivity to methane, S(CH4), the n u m b e r of active surface species reacting to methane, N*, and their surface residence time, x(CH 4) for Co/A1203 and Co/Re/A1203. X(CO)

S(CH~)

N*

~(CH4)

(%)

(%)

(pmol/gca t)

(s)

Before H20

7.4

53.6

21.1 _+0.6

21.4 _+0.6

After H20

5.7

47.1

18.3 _+ 1.0

24.0 +_ 1.3

Before H20

5.7

42.1

16.2 _+0.7

21.3 + 0.9

After H20

5.5

43.9

16.3 _+ 1.3

22.3 _+ 1.8

Before H.O

15.5

55.7

45.8 _+ 0.6

22.2 _+0.3

After H20

10.8

50.8

34.8 _+ 1.2

24.2 _+0.8

Before H20

9.8

50.4

29.1 _+ 1.6

22.3 _+ 1.2

After H20

8.9

49.6

26.9 _+ 1.2

22.7 + 1.0

Catalyst/treatment Co/AL_O/not pretreated

Co/Al~O/pretreated

Co/Re/A120/not pretreated

Co/Re/A120/pretreated

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

Steady-state isotopic transient kinetic analysis (SSITKA) shows that the deactivation of the m e t h a n a t i o n reaction after water t r e a t m e n t is due to a decrease in the n u m b e r of active sites, while the intrinsic site activity apparently remains unchanged. Rhenium promotion of the Co/A1203 catalyst results in a significantly higher activity, both for m e t h a n e and higher hydrocarbons. The promoted catalyst shows no better tolerance for the presence of water. The intrinsic site activity for methane formation is equal to the values found for the unpromoted catalyst.

ACKNOWLEDGEMENTS

The authors gratefully acknowledge financial support from the Norwegian Research Council.

202 REFERENCES

1. C.H. Bartholomew, Stud. Surf. Sci. Catal., 64 (1991) 158. 2. D. Schanke, A,-M. Hilmen, E. Bergene, K. Kinnari, E. Rytter, E. Adnanes and A. Holmen, Energy Fuels, 10(4) (1996) 867. 3. H. Schulz, M. Claeys and S. Harms, 4th Int. Natural Gas Conversion Symp., Kruger National Park, South Africa, 19-23 Nov. 1995. 4. E. Iglesia, S.C. Reyes, R.J. Madon and S.L. Soled, Adv. Catal., 39 (1993) 221. 95. H.H. Nijs and P.A. Jacobs, J. Catal., 66 (1980) 401. 6. E. Iglesia, S.L. Soled, R.A. Fiato and G.H. Via, J. Catal., 143, (1993) 345. 7. D. Schanke, S. Vada, E.A. Blekkan, A.-M. Hilmen, A. Hoff and A. Holmen, J. Catal., 156 (1995) 85. 8. A. Kogelbauer, J.G. Goodwin Jr. and R. Oukaci, J. Catal., 160 (1996) 125. 9. A.-M. Hilmen, Ph.D. Thesis, NTNU, Trondheim, 1996. 10. O.A. Bari~s, Ph.D. Thesis, NTH, Trondheim, 1993. 11. P. Biloen, J.N. Helle, F.G.A. van den Berg and W.M.H Sachtler, J. Catal, 81 (1983) 450. 12. X. Zhang and P, Biloen, J. Catal., 98 (1986) 468. 13. S. Vada, PhD. Thesis, NTH, Trondheim, 1994.

91997 Elsevier Science B. V. All rights reserved. Dynamics of Surfaces and Reaction Kinetics in Heterogeneous Catalysis G.F. Froment and K.C. Waugh, editors

203

Modelling of the dynamics of complex catalytic phenomena based on surface mobility processes and the remote control mechanism P. Ruiz, Y.-W. Li*, E. Gaigneaux, B. Delmon Unit6 de catalyse et chimie des mat6riaux divis6s, Universit6 catholique de Louvain, Place Croix du Sud, 2/17, B-1348 Louvain-la-Neuve, Belgium. ABSTRACT

We report new results concerning the role of surface mobile species in catalytic processes. More precisely, the proposed contribution deals with cooperative effects due to diffusion of some surface species from one to another kind of surface. The understanding of atomic scale phenomena leads to new modelling and chemical engineering developments which will be outlined. With sulfide and oxide catalysts, new catalytic sites are created by the reaction of spiUover species with the surface of a potentially active phase, thus modifying its activity and selectivity. This is the remote control (RC). A most conspicuous result is a strong catalytic synergy between two or several distinct solids. In addition, the active phase is protected against unwanted solid-state transformations leading to deactivation. When a RC operates, some part (phase) of the catalysts dissociates a molecule (02 or H2) to form surface mobile species. These species flow, namely "spiUover", over the other part. In this way, they create active sites, or regenerate sites deactivated by some infrequent but harmful processes occurring in the normal course of catalytic reactions. These may be failure of sulfur atoms to be eliminated by hydrogen in hydrotreating (I-IDS, HDN, etc.) or surface reduction in selective catalytic oxidation. Remotely controlled reactions lead to special kinetic equations, which consist of the product of a function determining the number of sites by another function reflecting the occupancy of sites and the intrinsic reaction rate (e.g. of the Langmuir-Hinshelwood type). Hydrotreating reactions are particularly interesting in this respect, because the remote control causes an interconversion between hydrogenation sites and sites responsible for the hydrogenolysis of carbon-heteroatom bonds. A new model will be shortly presented, which reflects all the specificities of hydrotreating reactions, including transient effects. In selective oxidations, the presence or absence of spillover oxygen has an effect on the facetting and reconstruction of MoO3 crystallites, as demonstrated using SEM and AFM. These results highlight the necessity to keep the surface structure in an optimal configuration during the whole catalytic cycle. A modified Mars-van Krevelen model can account better than previous ones for the reduction-oxidation cycle and the change in number of the active sites as a function of the reaction conditions. The outlooks for the development of new fundamental concepts or for a general chemical engineering approach to catalytic processes will be summarized. RC takes place in numerous reactions. Many aspects are influenced by consideration of the RC concept: activity and selectivity, catalyst formulation and "architecture", ageing processes, kinetic modelling, and process operation. A general philosophy emerging from the results presented is that full consideration of the dynamics of structural changes occurring at the catalyst surface permits conspicuous advances in both the atomic and molecular mechanisms and the macroscopic kinetic modelling. * On leave from State Key Laboratoryof Coal Conversion,Institute of Coal Chemistry,Chinese Academy of Science, Taiyuan 030001, China.

204 INTRODUCTION In spite of the efforts of a few groups in the world, especially the organisers and invited lecturers in this Symposium, only a handful of macroscopic kinetic models incorporate the information which has already been gathered concerning the thermodynamics and kinetics of atomic or molecular processes occurring on the surface of catalysts. Most models used in kinetic analyses are extremely crude, often purely empirical. This explains why industry often does not estimate worthwhile using models much more elaborate than pseudo-first order laws, as a recent survey suggests (1). It is worthless indeed to spend time in elaborate calculations if models are not rooted in the specific details of the reaction mechanisms and, therefore, are unable to represent kinetics in ranges of parameters outside that in which data were gathered. More generally, the vast majority of models presently used in the scientific literature have no predictive value, because they do not take into account the reality of the reaction. Worse, deviations from the results they possibly indicate cannot be analysed for detecting the cause, nor could suggest precise complementary experiments. As will be emphasised in this Symposium, there are several lines for improvement of this situation. Surface science, a fundamental approach to the dynamics of adsorption and desorption, thermodynamic evaluations based on fundamental data, accurate analysis of the role of chemical functions and the corresponding induction or steric effects due to substituents can bring much. In this contribution, we wish to draw the attention to effects discovered in our laboratory in the last 15 years, which lead to completely new kinetic models which account for results in a much more accurate way. These effects are the consequence of spillover. Spillover has been defined as a phenomenon in which a surface (atomic) species is formed from dissociation of molecules on one phase, and moves (spills over) onto a second phase which cannot produce it and where normally this species does not exist. Something had been forgotten in this definition, namely that this mobile species, called spillover (S.O.) species, is often much more reactive than those resulting from the adsorption of the molecules from which they are formed. Typical in this respect is the fact that spillover hydrogen (Hso) can reduce oxides or remove coke poisoning catalyst surfaces in conditions where H2 could not reduce the same oxide or remove that coke. Spillover oxygen is similarly more reactive than 02 for removing coke precursors. In the last 15 years, our work has shown that spillover hydrogen Hso or spillover oxygen Oso, can react with catalyst surfaces to create new sites, or transform catalytic sites poorly active or active in other reactions, to sites very selective for given transformations (2-4). We call this phenomenon the Remote Control (RC), because the phase which dissociates H2 or 02 acts as a sort of "control room" sending a message (Hso or Oso) which modifies the way the "chemical plant", namely the active site, operates. We report here new results concerning cooperative effects due to the diffusion of some surface species from one to another kind of surface. The result is a conspicuous catalytic synergy between two or several distinct solids (2-4). Sometimes, one of them has no activity of its own. We shall outline the main feature of the surface reactions due to spillover species and show that the remote control mechanism corresponds to completely new kinetic laws. These explain phenomena the origin of which had remained obscure for years and permit a much more accurate representation of experimental kinetics than any previous model. The corresponding concept and kinetic models led to the prediction of completely unexpected effects, these having been later verified by experiments (5).

205

H Y D R O D E S U L F U R I Z A T I O N , HYDROGENATION TREATING ON SULFIDED CATALYST

AND

HYDRO-

The hydrotreating catalysts are constituted of sulfides. The basically active, selective, and catalytically stable phase is MoS2 (or alternatively, WS2 in certain cases). It has been empirically discovered a long time ago that Co or Ni sulfides acted as promoters. After having shown that a cooperation existed between the crystalline phases (Co9S8 and MoS2) (6), we ftrst proposed that a remote control explained this synergy (7,8). The arguments proving this mechanism and the irrelevance of the 18 or so alternative explanations proposed in literature are presented elsewhere (9,10). These include several proofs that the popular so-called "CoMpS" phase does not exist. If some unstable association characterized by a Mtssbauer signal of Co detected both when Co is alone or in the presence of Mo, this fades away during activation or when catalysts are in the real conditions of hydrotreatment reactions. The remote control mechanism can be explained as follows. In essence, spillover Hso, emitted by Co9S8 (or other group VIII metal sulfides: NiS, PdSx, PbSx, RhSx) slightly "corrodes" the edges of MoS2, creating 3-fold coordination-unsaturated Mo atoms (3 CUS-Mo) which are active in hydrogenation (HYD) and pairs of sites, probably containing one coordination-unsaturated Mo (CUS-Mo) and a vicinal Mo-S-H moiety, these pairs being active in the hydrogenolysis of carbon-heteroatom bonds, in particular hydrodesulfurization HDS (fig. 1). The H2/H2S ratio and the relative amount of Co9S8, namely Rm=Co9S8/(MoS2+Co9S8) determine the Hso concentration on MoS2, hence the number of HYD and HDS sites. 3-CUS pair site C) S ~Mo | H ',?'~, vacant coordination sites: one up+one dowr 9

,.-,.

'

' single vacant coordination site

. \

-.

--....

\ \ \

"-.

\

t

"-I-t

I i

I

\

\

~\1

S

|

S

t

%

I

Figure 1. Schematic picture of MoS2 slabs viewed from the side and from the top, with representation of the coordinatively unsaturated molybdenum atoms acting as active sites. The movement of spillover hydrogen Hso to create the active sites is also represented. The 3-fold coordinatively unsaturated sites (3-CUS) or x, are active for HYD, the pair sites (CUS Mo + Mo carrying an S-H group) or a, are active in HDS. The kinetics of the whole process are quite particular in this sense that the rate equation is the product of two terms, one N, the number of sites, depending on catalyst composition Rm and gas composition, and a classical one (e.g. a Langmuir-Hinshelwood

206 term, LH) depending on the concentration of the various reactants. This corresponds to a function LH of composition, written LH (composition). The overall rate is: r = N(Rm, H2/H2S) 9LH (composition) Such a model very easily reproduces the synergy curve (7,8), including strange variations of activities when both composition and surface areas of catalysts change (11,12). Taking into account diffusion processes, this approach permitted to predict that catalytic pellets with different, independently varying profiles of Co and Mo, should have very different HYD/HDS selectivities, and this was verified experimentally (9,13). More comprehensive models were established by the group of Froment in Ghent (12, 14-16). This was modified and greatly improved in the model of PiUe and Froment, which accounted exactly for very carefully obtained data in the HDS of thiophene and HYD of the resulting butene in a Berty reactor (17). All this work is summarized in a review paper (5). We recently developed a modified model starting from that of Pille and Froment (17). This new model is aimed at improvements on two lines (18). First, we took into account the most probable structure of the HYD and HDS sites (as indicated above) and incorporated the corresponding stoichiometry with respect to Hso in the equations necessary for the model. Secondly, we took into account more in detail the kinetics of the diffusion of Hso from the promoter to the active phase (MoS2). More precisely, we took into account the kinetics of formation of Hso on the promoter, diffusion from promoter to MoS2, and reaction with the surface of MoS2 to create the HYD and HDS sites: 2Co+H2c:,2HCo

(S1)

[HCo]2 K1 = [Co]2PH2

(1)

HCo r

($2)

K2 = [Co][Hso,Co] [HCo]

(2)

[Cor]2PH2S K3 = [Co]2[Hso,Co]2

(3)

Co + Hso,Co

2Co + 2Hso,Co r

2Cor +H2S

($3)

In these equations, Co represents the sites for hydrogen dissociation and Cor deactivated sites. The relative proportion of the non-deactivated sites, 1-0r, can then be def'med as: 1 + 4K1PH2

1-0r =

(4)

1 + 4KiPH2 + 4 K 3 / P H 2 s [Hso,Co] It is assumed that the main factor limiting the surface diffusion (coefficient Ds) of spillover hydrogen is the transfer of Hso across the point of contact between Donor and Acceptor for unsupported catalysts; for supported catalysts, diffusion on the support surface may constitute the main rate limitation (1,4). The migration step ($5) and the corresponding lumped rate equation in a general form (eq. (5)) are formally identical for both supported and unsupported catalysts. The rate depends on Ds, the specific surface areas of the two partners Sg (see below), and the composition of the catalyst p=Co9Ss/(MoS2 + Co9S8). Hso, Co ~ Hso,Mo

($5)

r4 =

DsSgSCop(1 Or) L - I]pcat([Hso,Co] - [Hso,Co])

(5)

Here, Hso,Co and Hso,Mo represent, respectively, spillover hydrogen on Donor and on Acceptor, the specific area Sg can be defined as the specific surface area of the support for supported catalysts and that 6f the Acceptor phase (MoS2) for unsupported catalysts. The

207 term 13Pcatis a normalizing factor taking into account the mole composition of the catalyst and its apparent density. The HDS and HYD sites are created from inactive sites x by spillover hydrogen present on the Acceptor. The HYD sites, namely 3-fold CUS sites (x) result from step $6, to which eq (6) corresponds: 2n + 2Hso Mo r '

2X + H2S

($6)

K5 =

[x]2PH2s

(6)

[/l:] 2[I-Iso,Mo]2

where x represents the original edge sites (2-fold CUS's). The pair-sites for HDS (o) are formed by 2x + 2Hso,Mo r

o

[~ K6 = [1;]2[Hso,Mo]2

($7)

(7)

For writing the mechanism of HDS (of thiophene) and HYD over the created sites, we took into account data from literature (5,18). It was assumed that the adsorption of hydrogen is rapid, so that the hydrogen atoms needed in HYD and HDS reactions are already on the metal sites where thiophene or unsaturated hydrocarbons are sorbed or, alternatively, on the sulfur sites directly surrounding these metal sites. In addition, we took into consideration the experimental observations that active sites exist even in the absence of a remote control. This is due to the defect structure of the catalyst surface. These are "original" sites, for which no creation process is needed. 0.30

="

o~

=-

80 l Xms=0"0024

XH2S=0"0024

Z O m o3 ew

> z 0 r

.

O.lO

< 0.15

50.11~,. II'--e'--d~'~

' = 0"002

=

experimental: r=P 0

40+ ~ ' " 301

t/h

0.80

~'Z"X"~ . . . . . . .

8

0

I

1 600

L

. . . .

. . . . . 4 0.01 : ~'L-'-""--ll"-'---''t , - " ~ - -"a : 1200 Time

1800

2400

Figure 2. Switch operation: the feed is switched between high H2S and low H2S concentrations (xH2S). The composition of the catalyst r (experimental) or O (theoretical) corresponds to Co9S8/(MoS2 + Co9S8). Left hand: experimental (20,21); fight hand: theoretical. Using the thus built kinetic model, both the dynamic and steady state cases of a continuous stirred tank reactor were simulated. We show here the simulation results concerning transient effects. The case considered is the switching of feeds with different H2S concentrations. The shapes of the trajectories and the variations of activities are generally comparable to the experimental results (20,21), although the latter were obtained under low pressure (fig. 2). It is known that the addition of H2S depresses the HDS activity. Simulation results refine the conclusion. They conf'n'rn the experimentally found phenomenon in (22) that, for catalysts with different compositions, the depression

208 of activity is stronger when the synergy effect is stronger (catalysts in strong synergy range: p--0.3; p--0.1 for simulation) and weaker when synergy effect is weaker (catalysts in weak synergy range: p--0.01 for simulation). Another example concerns the change of HDS activity with a switch of total pressure (fig. 3). It was found that for catalysts with different P, namely different proportions of the promoter (Co9S8), the transient behaviors were different. An interesting feature when the proportion of the promoter is very low, is that the spiUover rate is also very slow and the system needs a much longer time to stabilize than in the case of catalysts with high proportions of promoter (fig. 3).

100 .... a

t9=0.3

D ~176. . . . .

9 . . . . . . . .

80-

~

~ 9

o

-v.-t

-0=0.02

60=

s

s :

J

> 40- ~ . . ~ . . . p ~ .

s

!

S

s

I

(0=0.9

I

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

~.om....,...m,

p=O.O

9

U 20-

5 bar

60 bar

60 bar

r .

m

I

....

t

0

I

600

~

1200 time

.

.

!

.

j

I

1800

2400

Figure 3. HDS activity (theoretical): switching between 60 and 5 bars for various catalyst compositions. Our new model also reproduces the variations of HYD and HDS activities with catalyst composition p, as well as transients. The model also strongly suggests that the interconversion from HYD to HDS sites needs the intervention of Hso, thus removing the ambiguity between the two models found satisfactory by Pille and Froment (5,17). But it should be clear that this argument must still be confh-med experimentally. SELECTIVE OXIDATION General Aspects A simple mechanical mixture of ~-Sb204, which is absolutely inactive, with MOO3, poorly active and selective, exhibits a much higher activity than MOO3. Figure 4 shows that 900 mg of MoO3 crystallites (in the 2-20 gm range) mixed with 850 mg of fine grained a-Sb204 (-< 1 gm) bring about a conversion of isobutene about 7.5 times higher than 900 mg of MOO3. Whereas a-Sb204 and the crystallized MoO3 sample did not produce methacrolein, the selectivity of the reaction to this latter product was substantial when MoO3 was mixed with a-Sb204 (23).

209

Isobutene conversion (%) 30

50 40 30

20

40 30 20

20 10 0Sb204

Methacrolein sel. (%)

Methacroleinyield (%)

60

10 obs. th. MM

MoO3

0 ~b204 MM MoO3 (Theor. = 0.0)

10 0

~b204

MM MoO3 (Theor. = 0.0)

Figure 4. Synergy between o~-Sb204 (850 mg) and MoO3 (900 mg): comparison of catalytic activity between the pure oxides and their mechanical mixtures (MM). The theoretical th. value corresponds to the arithmetic addition of the effect of each oxide measured independently (420 ~ isobutene:O2:He = 1:2:7 in vol.; total gas flow 30 ml.min -1) (23). This result is interpreted as due to the spillover of oxygen from a-Sb204 to MOO3. The effect is the consequence of the excessive reducibility of MOO3. Selective catalytic oxidation consists in a continuous reduction-oxidation of an active phase (e.g. MOO3) or, more precisely, a cyclic succession of the following reactions written in the case of the oxidation of propene to acrolein (where MonO3n suggests the surface of a crystallite): CH2=CH-CH3 + MonO3n ~ CH=CH-CHO + H20 + MonO3n-2

(8)

MonO3n-1 + 02 ~ MnnO3n

(9)

The excessive reducibility of MoO3 implies that the f'u'st step is more rapid than the second. The interpretation of the role of a-Sb204 is that it emits spillover oxygen Oso which assists molecular oxygen in the reoxidation step. Actually, catalysts remain in a higher oxidation state when mixed with a-Sb204, in comparison with experiments in the same conditions, but in the absence of a-Sb204 (24,25). Incidentally, it can be shown that this effect protects catalysts like Fe2(MoO4)3 or ZnFe204 against segregation due to reduction.

Molecular and Structural Aspects of the Catalytic Reduction-Oxidation Cycle In MOO3, the metal is octaedrally coordinated to oxygen ions. In the fully oxidised state, the octaedra are linked by corners. Each time an oxygen atom is removed from the surface, 2 corner-sharing octaedra have to attach to each other by edges. The selective oxidation involves 2 pairs of octaedra. The oxido-reduction process thus corresponds to fig. 5. During the oxidation reaction, a continuous alternation of edgesharing and comer-sharing arrangements takes place. The selective sites consist of pairs of comer-sharing octaedra in a structural environment which allows the changes of fig. 5 to take place without disruption of the lattice. This corresponds to the normal, selective reaction to produce unsaturated aldehydes. Experience shows that when used alone, MoO3 oxide gets more reduced than when in contact with a-Sb204. Simultaneously, the selectivity to the partial oxidation products is kept higher. In the absence of Oso donors (e.g. a-Sb204), the concentration of edge-sharing octaedra pairs becomes larger and these edge-sharing structures can aggregate and bring about the nucleation of shear structures which correspond to non selective sites producing CO2 and H20. The role of Oso is to stop the growth of the rows of edge-sharing octaedra and to maintain these rows

210 below the critical nucleus size corresponding to the formation of the non-selective shear structures.

c, 3 - c,-

c, 2

CH2

=

CH- CHO

.,2 o

o2

Figure 5. Schematic representation of the structural changes at the surface of MoO3 due to the reduction-oxidation cycle corresponding to the selective oxidation of propene to acrolein. The interpretation represented schematically in fig. 5 suggests that, by maintaining an active dynamic swing between corner and edge-sharing octaedra, Oso could permit a progressive reconstruction of the catalyst crystallites in the conditions of the catalytic work. Each step could correspond to the creation of a new arrangement of edge-sharing octaedra, these being in turn transformed to comer-sharing octaedra in eventually a different configuration compared to the initial one. It was therefore logical to expect that this reconstruction could change the proportion of the various crystallographic faces of MOO3. Many authors reported that the (010) face was little active, permitting only total oxidation (this was confm'ned by another set of experiments in which we used millimeter long MoO3 crystals developing principally the (010) face). The same mechanical mixture of 2-20 ~axnMoO3 crystallites with fine grained a-Sb204 as for the results of fig. 4 was subjected to various physico-chemical measurements after the catalytic tests. No trace of mutual contamination of the two phases was detectable, neither in the bulk nor in the surface (using X-ray diffraction and XPS), and no change of surface area was measured (BET). But investigation of the samples by scanning electron microscopy revealed important modifications of the morphology of MoO3 crystallites, but only when MoO3 had been mixed with ct-Sb204 (26,27). The edges between (010) and (100) faces, which exhibited a sharp intersection in the fresh sample, acquired a faceted structure by developing, at the micrometer scale, steps with vertical walls oriented as (100) (fig. 6). This corresponds to changes at the nanometric scale, as shown by atomic force microscopy. The (010) face (smooth and flat when fresh) got fragmented, exhibiting pits and crevices mainly indexed as (100). None of these phenomena was detected when MoO3 was tested alone (28-30). In this case also, the catalytic properties changed considerably. The selectivity to methacrolein rose from 0 to 37%.

211

Fig. 6. SEM micrographs of the edge of a MoO3 crystallite before (top) and after (bottom) the catalytic reaction in the presence of a-Sb204 (conditions in the legend of fig. 4). Black arrows indicate the steps formed by the reconstruction of the (010) face to (100) steps (23,26,27). At the atomic level, the coordination of the Mo atoms on various faces of molybdenum crystallites is different (fig. 7). The (010) face exposes Mo atoms with identical coordination, with Mo=O groups pointing outside, linked together by bridging oxygens, leading to a flat surface. The (100) faces are formed by the sequence of 4 Mo atoms with different coordinations periodically repeated (along the [010] direction). The resulting [001]-directed wide grooves allow the concerted elimination of 2 hydrogen atoms and the insertion of 1 oxygen atom into the isobutene molecule. This is not possible on the (010) face, thus explaining the difference of reactivity of these faces. By triggering the reconstruction, the role of Oso would thus be to promote the formation of a "selective coordination" of Mo atoms (the ones of (100) faces) to the detriment of the "non selective ones" (the ones of (010) faces).

212

(010) netplane ...........i............ ~:~::::ii:i~!i!!!:i'84184 !'i!!~:i::?i:::~i~!~ii~!i~i:~!!:!~!~!i 84 i!!:i!!!!~!i~::!~iii! ....

i!i:i 9 !ii::i 84184 :i!:!!i:ii i!:!:i~i:ii~i... i: ~:iil ~.:~:iii:i:iS i::::i.::ii.i .::: :i i:i~:i::::~/:.ii~:~i:i~i:: :.?..,~i:::~i::i~::i:i:~:i~

::::::::::::::::::::::::::::::::::::::::::::::::::... ::. ::::::::::::::::::::::::::::. ::::::::::::::::::::::::::.::.. :::.: ,.: ::::,.:.::

::::::::::::::::::::::::::::::::::::::::::::::::::::::::::::::::::::::::::::::::::::::::::::::: ::::::::::::::::::::: :,:::,::.:::: ~::::::::::::::::::::::::.:: :.:,.: ::,.:.:~

.....................:::.:::::......................

!!!!i ii!ii!ii Figure 7. Atomic representation of the (010) and (100) faces of MoO3 crystallites (obtained by using the Balzac software, K. Hermann et al., Fritz-Haber Institute, Berlin, Germany). These results have to be related to crystallographic data showing that a suboxide of Mo, namely Mo8023, is formed when the catalyst works in the presence of a-Sb204 (27-30). This work leads to conclusions close to those of Gai-Boyes (31). In the corresponding work, high resolution electron microscopy studies in conditions close to those of the catalytic reaction showed that another suboxide, namely Mo18052, was the real active and selective phase. Both Mo8023 and Mo18052 have structures similar to MoO3 both containing shear planes (with edge-sharing octaedra). These intermediary structures can easily lose and reincorporate two oxygen atoms without any irreversible structural change. Kinetic Models On the basis of this molecular level picture, Rebitsky et al. constructed a complete kinetic model of the remote control effect in selective oxidation, including transient effects (32). This predicts isothermal bistability in certain conditions. Such bistabilities have been found experimentally in a special reaction (33). For a catalyst of composition (by weight) Rm = acceptor/(acceptor+donor), the model gives the fraction a of the potential sites which are really active as a function of the reaction conditions (hydrocarbon and oxygen pressures, PHC and PO2 respectively, temperature T and time t).

a = a (Rm, PHC, P02, T, t)

(10)

It is possible to incorporate in this model the kinetics of the mechanism described in eq. (8) and (9), namely the so-called Mars-van Krevelen mechanism. The reduction and oxidation rates are represented by rred = kredPHC

(11)

rox = koxpno

(12)

213 The rate of the selective oxidation at each of the active sites (those which undergo the reduction-oxidation process, namely the two pairs of octaedra mentioned before) is kredPHC 9koxPO2 n ras = kredPHC + koxPO2 n

(13)

With respect to the hydrocarbon partial pressure, a first order is usually observed. The order n for oxygen corresponds to 1 or 0.5 according to the mechanism of reaction. Let N be the number of potential active sites per unit weight of the active phase. The number of sites really active in the experimental conditions will be tx.N per unit weight of acceptor. The corresponding overall rate is tx.N.ras. The rate rRm per unit weight of a mixture of acceptor and donor of composition Rm can be deduced from this quantity. The fraction, namely the weight of acceptor, being Rm, the rate will be Rma.N.ra: rRm = Rm oNo o~(Rm,PHC,PO,T,t) ~ kredpHC ~ koxpO2n kredPHC + koxPO2n

(14)

Equations of this form (with n--0.5) represent satisfactorily results of careful kinetic measurements in the oxidation of isobutene to methacrolein on mixtures of separately prepared o~-Sb204 and MoO3 of different compositions (34,35). Until this, published kinetic models in selective oxidation were either purely empirical or, if based on the Mars-van Krevelen model, contained fractional exponents which had no meaning in terms of mechanism (34,35). As in the case of hydrotreatments, the distinctive feature of eq. (14) is that the reaction rate is the product of two terms. One reflects the remote control, namely the way the number of active sites varies as a function of catalyst composition and experimental conditions. The other one expresses the intrinsic rate at active sites. The experimental conditions influence both terms. This feature had never been incorporated before in macroscopic kinetic models. OUTLOOKS The phenomena we have described above show that the traditional kinetic models based only on thermodynamics, Langmuir-Hinshelwood and Eley-Rideal equations are insufficient for reflecting adequately the dynamics of the real atomic and molecular processes in heterogeneous catalysis. In a very general way, all structural changes and reconstructions that surfaces undergo must be taken into account, in principle, for obtaining really relevant kinetic models. Even in the field of surface science, only few models yet take really into account the dynamics of the reconstruction of surfaces according to experimental conditions. In a recent article, Delmon and Froment noticed that models incorporating the occurrence of spillover phenomena were still very scarce (5). Spillover hydrogen plays a crucial role in removing carbonaceous deposits, which are coke precursors, in catalytic reforming, catalytic cracking, and hydrocracking. S tumbo et al. showed recently that spiUover hydrogen can enhance by a factor as high as 5 to 10 the number of Br6nsted sites on silica-aluminas in the reaction conditions (36,37). The regeneration of cracking and hydrocracking catalysts involves spiUover oxygen. New mechanisms occurring on zeolites are discovered (36-40). All this can potentially be described by macroscopic models following lines similar to those indicated above. These models in turn will have the same potential to stimulate both fundamental investigations and chemical engineering developments as those concerning the remote control. With respect to surface science, recent examples considering surface mobility phenomena concern new discoveries which, in turn, could lead to macroscopic kinetic models. A typical case is the cooperation between a Cu(111) and a Pt(111) surface in the oxidation of CO (41,42). The presence of islands of epitaxial layers of Cu(ll 1) on

214 Pt(111) influences the bistability observed on Pt(111). Dissociatively adsorbed oxygen on Cu(111) can trigger the reaction of CO adsorbed on Pt(111) with molecular oxygen 02, in conditions where pure Pt(111) is otherwise inactive. A reaction front expands from Cu to the whole Pt surface. The reason is that small amounts of adsorbed O atoms on Cu(111) cross the Cu/Pt border, react with the CO saturated Pt(111) surfaces and thus create vacancies, namely a disorder in the close-packed CO adlayer. This leaves free pairs of adsorption sites on Pt and permits the dissociative adsorption of 02. Once the "disordering" is triggered, the reaction proceeds and a reaction front moves across Pt(111). Other experiments show that Cu atomically attached to Pt(111) increases the rate of front propagation. Doping with Cu has thus two different effects: triggering the reaction and accelerating it when the system is in the low activity bistability branch. Oxidation of CO on Pt(111) is simple in the sense that no reconstruction of the surface takes place. To be representative, any overall modelling would nevertheless be complicated and involve two terms, dealing exclusively, one with the unsteady state behavior, and the other with rates observed in each state. A very strange delay occurring sometimes when the reaction front crosses the border between two domains of different composition (42) will set an additional challenge to modelling. There is little doubt that both fundamental science and chemical engineering are on the verge of exciting discoveries and new developments thanks to approaches taking fully into account the dynamic processes occurring at surfaces. ACKNOWLEDGMENTS We acknowledge the support of the Service f6dfraux des Affaires scientifiques, techniques et culturelles (Belgium) for their support in the frame of the Attraction Interuniversity Pole Katalyse (Promoter: G.F. Froment). We also acknowledge the support of the A. von Humboldt Foundation for a Research Award (BD) and a permission for leave from State Key Laboratory of Coal Conversion, Institute of Coal Chemistry, Chinese Academy of Science, Taiyuan (Y.-W. Li). E. Gaigneaux thanks the Fonds National de la Recherche Scientifique of Belgium for the "Aspirant" fellowship he was awarded. REFERENCES

1.

.

o

4. 5. 6. 7. 8. 9. 10. 11. 12. 13. 14. 15. 16.

A.N.R. Bos, G.B. Matin, L. Lefferts, M.H.G.M. Steins, Taskforce on "Kinetic Research on Heterogeneously Catalysed Processes", of the Working Party "Chemical Engineering in the Applications of Catalysis" of the European Federation of Chemical Engineering. B. Delmon, in "New Aspects of Spillover in Catalysis" (T. Inui, K. Fujimoto, T. Uchijima and M. Masai, eds.) Elsevier, Amsterdam, 1993, 1-8. B. Delmon, Heterog. Chem. Rev., 1 (1994) 219. B. Delmon, Surf. Rev. and Lett., 2 (1995) 25. B. Delmon, G.F. Froment, Catal. Rev.-Sci. Eng., 38 (1996) 69. G. Hagenbach, Ph. Courty, B. Delmon, J. Catal., 31 (1993), 264. B. Delmon, React. Kinet. Catal. Lett., 13 (1980) 203. B. Delmon, Int. Chem. Eng., 20 (1980) 639. B. Delmon, in "Catalysts in Petroleum Refining 1989" (D.L. Trimm, S. Akashah, M. Absi-Halabi, A. Bishara, eds.), Elsevier, Amsterdam, 1990, 1-40. B. Delmon, Bull. Soc. Chim. Belg., 104 (1995) 173. B. Delmon, Bulg. Acad. Sci., Comm. Dept. Chem.: 17 (1984) 107. C.-Y. Yu, G.F. Froment, Chem. Engg. Sci., 46 (1991) 3177. J.L.G. Fierro, J.M. Asua, P. Grange, B. Delmon, Preprints, ACS Div. Petrol. Chem., 32 (1987) 271. I.A. Van Parijs, G.F. Froment, Chem. Engg. Sci., 25 (1986) 431. I.A. Van Parijs, G.F. Froment, Appl. Catal. 21 (1986) 273. G. Villora, A.O. Beyne, G.F. Froment, Ind. J. Technol., 29 (1991) 128.

215 17. 18.

19. 20. 21. 22. 23. 24. 25. 26. 27. 28. 29. 30. 31. 32. 33. 34. 35. 36. 37. 38. 39. 40. 41. 42.

R.C. Pille, C.-Y. Yu, G.F. Froment, J. Mol. Catal., 94 (1994) 369. Y.-W. Li, B. Delmon, Proc. 4th Intern. Conf. on Spillover, Dalian, China, Sept. 15-18, 1997, in press. Y. Okamoto, H. Tomioka, Y. Katoh, T. Imanaka, S. Teranishi, J. Phys. Chem., 84 (1980) 1833. D. Pirotte, P. Grange, B. Delmon, in "Heterogeneous Catalysis" (Proc. 4 th Intern. Syrup. Oct. 1979), Varna, Part 2, 127-132. B. Delmon, Proc. Climax 3rd Int. Conf.: Chemistry and Uses of Molybdenum (H.F. Barry, P.C.H. Mitchell, eds.) Climax Molybdenum Cy., Ann Arbor, Mich., 1979, 73-84. J. Bachelier, M.J. Tilliette, J.C. Duchet, D. Comet, J. Catal., 87 (1982) 292. E.M. Gaigneaux, P. Ruiz, B. Delmon, Catal. Today, 32 (1996) 37. P. Ruiz, B. Delmon, Catal. Today., 3 (1988) 199. L.-T. Weng, B. Delmon, Appl. Catal. A, 81 (1992) 141. B. Delmon, Bull. Soc. Chim. Belg., 105 (1996) 67. E.M. Gaigneaux, P. Ruiz, B. Delmon, 11th Int. Congress on Catalysis, Baltimore, MD (USA), June 30-July 5 1995. E.M. Gaigneaux, P. Ruiz, E.E. Wolf, B. Delmon, Appl. Surf. Sci., accepted. E.M. Gaigneaux, P. Ruiz, B. Delmon, 4th Int. Conf. on Spillover, Dalian (China), Sept. 15-18, 1997, in press. E.M. Gaigneaux, P. Ruiz, B. Delmon, accepted for oral presentation, 3rd World Congress on Oxidation Catalysis, San Diego, CA (USA), Sept. 21-26, 1997. P. Gai-Boyes, Cat. Rev.-Sci. Eng., 34 (1992) 1. T. Rebitzki, B. Delmon, J.H. Block, A1ChE Journal, 41 (1995) 1543. A. Gil, P. Ruiz, B. Delmon, Canad. J. Chem. Engg., 74 (1996) 600. D. Vande Putte, S. Hoomaerts, F.C. Thyrion, P. Ruiz, B. Delmon, Catalysis Today, 32 (1996) 255. S. Hoornaerts, D. Vande Putte, F.C. Thyrion, P. Ruiz, B. Delmon, Catalysis Today, 33 (1997) 139. A.M. Stumbo, P. Grange, B. Delmon, 1l th Int. Congress Catalysis, Baltimore, 1996, J.W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell (Eds.), Studies in Surface Science and Catalysis, 101 (1996) 97-106. A.M. Stumbo, P. Grange, B. Delmon, in "Hydrotreatment and Hydrocracking of Oil Fractions" (G.F. Froment, B. Delmon, P. Grange, eds.), Elsevier, Amsterdam, 1997, 225-235. K. Fujimoto, in "New Aspects of Spillover Effect in Catalysis" (T. Inui et al., eds.), Elsevier, Amsterdam, 1993, 9-16. I. Nakamura, R. Iwamoto, A.I-ino, in "New Aspects of Spillover Effect in Catalysis" (T. Inui et al., eds.), Elsevier, Amsterdam, 1993, 77-84. R. Roessner, U. Roland, T. Braunschweig, J. Chem. Soc., Far. Trans., 91 (1995) 1539. M. Kolodzejczyk, R.E.R. Colen, M. Berdau, B. Delmon, J.H. Block, Surf. Sci., accepted. M. Kolodzejczyk, R.E.R. Colen, B. Delmon, J.H.Block, Appl. Surf. Sci., accepted.

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91997 Elsevier Science B. V. All rights reserved. Dynamics of Surfaces and Reaction Kinetics in Heterogeneous Catalysis G.F. Froment and K.C. Waugh, editors

217

Adsorption and Reactions of Methane on Ferric Molybdate using DRIFTS Technique Surajit Fuangfoo, Anand S. Chellappa and Dabir S. Viswanath* Department of Chemical Engineering, University of Missouri-Columbia, MO 65211, USA The adsorption and reactions of methane on Fe203, MoO3 and ferric molybdate catalysts have been studied using the Diffuse Reflectance Infrared Fourier Transform Spectroscopy (DRIFTS) technique. This study was carried out at the temperature of 703 K and, at the pressure of 30 bar. The Mo/Fe ratio was changed from 1 to 5, thus giving a broad range of atomic ratios. The formation of methoxy (OCH3), adsorbed formaldehyde, dioxymethylene (OCH20), and formate (OCHO) species were observed on the catalyst surface. A mechanism based on the reaction of methane with the surface of the ferric molybdate catalyst to produce methoxides which are further oxidized to dioxymethylene is proposed. The formation of carbon oxides is attributed to either the decomposition of dioxymethylene or to the formation of formate species. This study revealed that MoO3 is the active catalyst for the partial oxidation of methane to formaldehyde and Fe203 gives carbon oxides. Both methanol and formaldehyde can be obtained from the ferric molybdate catalysts. 1. I N T R O D U C T I O N The conversion of methane to methanol and formaldehyde has been receiving greater attention because of the enormous natural gas reserves and the increase in projected methanol demand. Although the partial oxidation of methane to methanol and formaldehyde is a thermodynamically favorable process, the conversions and selectivity required to make this direct process economical have not yet been achieved since it is difficult to activate a stable molecule like methane. Thus, the elucidation of the mechanism of direct partial oxidation of methane is an outstanding problem in chemistry and chemical engineering. Several investigations [1-15] have been carried out to study the active catalysts and to understand the reaction pathways for this process. Among the catalysts studied, vanadium- and molybdenumbased systems appear to be the most promising [1-10]. Mechanisms, based on the Mars-van Krevelen Redox type, have been attempted. Spencer and co-workers [5-7, 11] suggested two different pathways for the partial oxidation of methane over supported molybdenum and vanadium oxides using oxygen as an oxidizing agent. First, methane is converted to formaldehyde and carbon dioxide, and formaldehyde is further oxidized to carbon monoxide. Second, methane oxidation follows a sequential reaction path, in which formaldehyde is formed from methane and is subsequently oxidized to carbon monoxide and carbon dioxide. Both pathways have a common methane activation step where methoxy group is formed. Liu et al. [8] and, Khan and Somorjai [9] carried out the conversion of methane over silica-supported molybdenum oxide using nitrous oxide as an oxidant. They proposed that methane was activated by the reactive catalyst to form methoxide complexes which may decompose to form formaldehyde or react with water to form methanol. The formation of methoxide ions on the catalyst surface was confirmed [8] by the infrared spectrum of the catalyst obtained after an addition of methane at 298 K and a subsequent evacuation. Using tracer isotropic techniques, Banares and co-workers [12-15] showed that in the partial oxidation of methane over MoO3/SiO2 catalyst, the oxygen of formaldehyde comes directly from the molybdenum oxide lattice. Thus, the role of the oxidant is to regenerate the catalyst. They also reported that the reoxidation of the catalyst using nitrous oxide would be less effective than using oxygen. Although various kinetic data on the catalytic conversion of methane are available, surface information and spectroscopic data are still limited. In the present study, we have

218 focused our attention on the characterization of adsorbed surface species under in situ conditions. The use of spectroscopic methods to identify the reaction intermediates should give a better understanding of the molecular interactions occurring on the catalyst surface, and help establish complete reaction mechanism. Using supported ferric molybdate as a catalyst and air as an oxidant, methanol selectivity obtained in the earlier work [ 1] compares well with recent works in the literature. In this work, we attempt to describe the mechanism of methane conversion on the unsupported ferric molybdate catalysts using in situ Fourier transform infrared (FT-IR) spectroscopy. The diffuse reflectance infrared Fourier transform spectroscopy (DRIFTS) technique was used to monitor the species formed on the catalyst surface after the exposure of methane under reaction conditions. The results show that it is possible to observe the principal intermediates involved in the conversion of methane over the iron-molybdenum oxide catalysts. Some adsorption studies on methanol and formaldehyde using spectroscopic methods have been carried out at conditions far removed from the actual temperature and pressure conditions used for partial oxidation of methane. Thus the present study provides more realistic results of methane reactions at the catalyst surface.

2. EXPERIMENTAL Fe203, MOO3, and ferric molybdate catalysts with Mo/Fe atomic ratio of 1.0, 1.7, and 5.0 were used in this study. Fe203 (Aldrich, 99.998%) and MoO3 (Aldrich, 99.99%) were used as received. Ferric molybdate catalysts were prepared by coprecipitation of the solutions of ammonium molybdate, (NH4)6Mo7024.4H20 (Fischer, ACS grade), and ferric nitrate, F e ( N O 3 ) 3 - 9 H 2 0 (Fischer, ACS grade), at a constant pH 2 and 343 K according to the modified method described by Srihari and Vilswanath [ 16]. After filtration and washing, the precipitate was dried at 373 K for 1 h, then calcined at 673 K for 6 h, and gradually cooled to room temperature. The catalysts with different Mo/Fe atomic ratios were obtained by changing the initial amounts of ammonium molybdate and ferric nitrate. A BET sorptometer (Porous Materials Inc.) was used to determine the surface area of the catalysts. Using argon gas as an adsorbate at 77 K, the surface areas of Fe203, MOO3, and the catalysts with Mo/Fe ratio of 1.0, 1.7, and 5.0 were found to be 15, 9, 22, 17, and 8 m2/g, respectively. Infrared studies of the adsorption and reaction of methane on the catalyst surface were carried out using an FT-IR spectrometer (Perkin-Elmer, system 2000) equipped with a DTGS detector and a diffuse reflectance accessory (Spectra Tech, COLLECTORTM), as shown in Figure 1. A high temperature/high pressure (HTHP) chamber with water-cooled zinc selenide (ZnSe) dome windows (Spectra Tech, PN 0030-102) was used as a reactor cell. The catalyst was heated directly by a ceramic heater and a K type thermocouple was in contact with the catalyst for direct temperature measurement. The catalyst temperature could be controlled within +1 K using a temperature controller (OMEGA, CN8500). The HTHP chamber was connected to a vacuum pump which gave a vacuum of 10-2 torr. DRIFF spectra were recorded at a resolution of 4 cm-1, and 64-100 scans were averaged to achieve adequate signal-to-noise ratio. Dried nitrogen was purged through the sample compartment and the detector area of the FT-IR spectrometer in order to minimize moisture and carbon dioxide in the background spectra. The catalysts were ground (<325 mesh) and loaded (about 40 mg) in the sample cup inside the HTHP chamber. Before starting experiments, the catalysts were pretreated at 673 K for 1 h in a flow of air and 30 min in vacuum in order to remove adsorbed water. The temperature was then increased to 703 K and kept for 10 min before the IR spectra of the

219 gr-IR DRIFTS Accessory

Cold Trap Methane

To Vent

PRV

Flowmeter

Air

Helium Temperature Controllerand Readout

Vacuum Pump

Figure 1. FTIR experimental setup. PRV:pressure reduction valve. dehydrated catalysts were recorded and used as the background spectra. The air stream (AG Co., Ultra Zero grade) was dried using a molecular sieve trap before it was introduced to the HTHP chamber. Methane (AG Co., Ultra High Purity grade, 99.99%) was further purified by passing through two molecular sieve traps and an Oxy-trap (Alltech). Methane gas was then added to the HTHP chamber to a pressure of 30 bar and the system was kept static for 2 h. IR spectra were recorded every 15 min during this period. After degassing, the IR spectra were recorded again, and helium saturated with water (HPLC grade) was passed over the catalyst for 1 h at a constant pressure of 1 bar, and the outlet gas was trapped using liquid nitrogen. The condensed liquid obtained from the liquid nitrogen trap was analyzed using a Shimadzu GC14A gas chromatograph, linked with a Shimadzu CR501 integrator and equipped with an FID detector and a Porapak-T column. 3. RESULTS AND DISCUSSION

3.1. Catalyst Characterization The IR spectra of fresh catalysts recorded at room temperature are shown in Figure 2. The spectra of MoO3 and the ferric molybdate catalysts, unlike Fe203, present many bands in the range 1100-800 cm-1. A peak at 1017 cm-1 is characteristic of the octahedron (MOO6) vibration [17]. A band at 985 cm-1 corresponds to MoO3 while a band at 957 cm-1 corresponds to ferric molybdate [ 18]. Two bands at 823 and 787 cm -1 are attributed to the MoO-Mo or Fe-O-Mo stretching modes [18, 19]. The bands between 2000 and 1700 cm -1 may arise from overtones and/or combinations of the intense lattice vibrations [20]. For all catalysts, a broad band around 3400-3300 cm -1 and a fairly narrow band at 1620 cm-1 are due to adsorbed water [21, 22] as the intensifies of these peaks are considerably reduced upon heating. After pretreatment, the remaining bands assigned to OH-stretching vibrations of hydroxy group [23] appear at 3720, 3660, 3630, 3540 and 3470 cm-1 for Fe203, and 3540 and 3405 cm-1 for the catalysts with Mo/Fe ratios of 1.0 and 5.0, respectively.

220 3.04.0-

2.5-

(a~

2.0-

~

o 3.0(b3

1.5-

,

2.0-

< 1.0-

____._

(d)

.(d'~

~.o- (d)

0.50.0 4000

I

i

3500

3000

2500

Wave number (cm-1 )

0.0

,

,

,

,

,

,

,

2000 1800 1600 1400 1200 1000 800 Wave number (cm-1)

Figure 2. Infrared spectra of fresh catalysts at room temperature: (a) Mo/Fe = 1.0, (b) Mo/Fe = 1.7, (c) Mo/Fe = 5.0, (d) MOO3, (e) Fe203, and (') is represented the spectra of the same catalysts after pretreatment at 703 K. Spectra of KBr at the same temperature are used as background.

3.2. Adsorption and reaction of m e t h a n e The IR spectra of gas phase products resulting from the reaction of methane with the catalysts, shown in Figure 3, were obtained by subtracting the spectra at 120 min after the reaction from the spectra at 0 rain (when methane Fast added). The formation of carbon dioxide identified by a typical peak at 2350 cm-1 (doublet), and carbon monoxide identified by peaks at 2170, 2120 cm -1 are seen from these spectra. It can be also seen that among these catalysts, the intensity of carbon dioxide peak decreases with increasing Mo/Fe ratios. In the case of Fe203 and MOO3, the peaks of carbon monoxide could not be observed from the spectra. The gas phase methane shows strong absorption peaks at 3016 and 1304 cm-1 which span a wide range in these regions. Therefore during the reaction, spectra of the adsorbed species could not be monitored directly due to masking of these peaks. The spectra of the adsorbed species were obtained by plotting the difference between the spectra of the reacted catalysts after the excess methane was removed and the spectra of pretreated catalyst before methane was added. The resultant spectra exhibit four sets of peaks as shown in Figure 4. Peaks in the range 3000-2700 cm-1 are associated with C-H stretching vibrations, those at 1474, 1457, 1420(sh), and 1365 cm -1 are assigned to C-H deforming vibrations, others at 1580/1550, 1510, 1300, and 1252 cm -1 are identified as O-C-O stretching vibrations, and that at 1608 cm-1 is identified as C=O stretching vibration. The assignments of these features are based on the previous studies of the adsorption of methanol and formaldehyde on various metal oxides as reported in the literature. There are four different species observed on the catalyst surfaces, namely, (1) two types of methoxy group, one formed on terminal oxygen site indicated by the bands at 2960, 2854 and 1474 cm-1, and the other formed on bridged oxygen site indicated by the bands at 2928, 2826, and 1457 cm -1 [8, 21], (2) a surface dioxymethylene

221 species (OCH20) identified by the peaks at 2900, 1510, 1420, and 1252 cm -1 [24-26], (3) a surface formate species (OCHO) identified by the small bands at 2871, 1580/1550, 1365, and 1.20 2

1.0-

(a) o

L 0

<

CH4

CH4

0.8-

0.6-

0.4-

II'n

0.2-

0.0 4000

I

3500

I

3000

1

2500

I

2000

I

1500

k._ 1000

Wave number (crn-1) Figure 3. Infrared spectra show the formation of gas phase products when the catalysts were exposed to methane at pressure of 30 bar and temperature of 703 K for 2 h: (a) Fe203, (b) Mo/Fe = 1.0, (c) Mo/Fe = 1.7, (d) Mo/Fe = 5.0, and (e) MOO3. The spectra of the catalysts at 0 min (when methane f'wst added) have been subtracted. 1300 cm -1 [24-27], and (4) an adsorbed formaldehyde identified by a small peak at 1608 cm -1 [25, 27], and a shoulder at 2782 cm-1 [28]. In addition to the peaks mentioned above, there exists some other features in the IR spectra around 2000-1700 cm-1 region that have not been assigned. These bands may result from the reduction of the catalyst surface during reaction. All of the observed peaks remained unchanged under prolonged evacuation, but their intensities were reduced when helium saturated with water was passed over the samples at the same temperature of 703 K, as shown by spectra (a') to (e') in Figure 4. The reacted catalysts were reoxidized by heating under a flow of air at 673 K for lh and the spectra taken after outgassing are shown by Figures 4(a") to 4(e"). For Fe203, the spectrum 4(a) exhibits the loss of the peaks at 3720, 3660, 3630, 3540, and 3470 cm-1, seen earlier in Figure 2(e') and assigned to OH-stretching vibrations of hydroxy group. The other two negative bands in the spectrum 4(a) around 1550 and 1190 cm -1 may be attributed to the loss of oxygen from the catalyst surface and they are not able to be regenerated under the reoxidation condition used, as seen in the spectrum 4(a"). The spectra of ferric molybdate catalysts with Mo/Fe ratio of 1.0 and 5.0, Figures 4(b) and 4(d), also show the permanent loss of the peaks associated with the hydroxy group, as discussed earlier in Figures 2(a') and 2(c'). As seen from Figures 4(b") to

222

1.0-

2.0(a'3 0.8_ ~

-

~

~

-

o0.6

. _ _ _ (bl

_

o

O d:a

<

~1 0

(c'3 0

.

4

-

~

(d "')

_

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--'------~ (d")

0.2-

(e")

(e")

_

0.0 4000

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I

3500

3000

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2500

1

I

I

i

2000

1750

1500

1250

1000

Wave number (cm -1) Wave number (cm -1) Figure 4. Infrared spectra of adsorbed species formed on the catalysts at 703 K after exposure to methane at pressure of 30 bar for 2 h and following evacuation: (a) Fe203, (b) Mo/Fe = 1.0, (c) Mo/Fe- 1.7, (d) Mo/Fe = 5.0, and (e) MOO3. ('), and (") are represented the same catalysts evacuated after exposure to a flow of helium saturated with water at 703 K for lh, and air at 673 K for 1 h, respectively. The spectra of the activated catalysts at the same temperature have been subtracted.

223 4(e" 3, the reoxidation reactions under the condition used are quite complete for MoO3 and ferric molybdate catalysts, except for Mo/Fe ratio of 1.0 and Fe203. The condensed liquid obtained by passing helium saturated with water over the reacted catalysts was analyzed for oaethanol and formaldehyde using a GC equipped with an FID and a Porapak-T column. Both methanol and formaldehyde were found to be present in the samples obtained using ferric molybdate catalysts while the sample obtained using MoO3 showed only formaldehyde and that obtained using Fe203 showed only trace amount of methanol. It was also seen that the formaldehyde content in the samples obtained from ferric molybdate catalysts increased with increasing Mo/Fe ratios and the catalyst with the Mo/Fe ratio of 1.7 was the most favorable to methanol. 3.3 R e a c t i o n

mechanism

The results from the infrared studies and from the GC analysis show that the reaction of methane with the ferric molybdate catalysts gives methanol, formaldehyde, carbon dioxide, and carbon monoxide as final products. The IR spectra also indicate the formation of methoxy, surface dioxymethylene, surface formate species, and adsorbed formaldehyde. Based on these observations, a mechanism was proposed to account for all intermediates and final products and is shown in Figure 5. Since the surface structure of the catalysts is not known, the surface is represented by a straight line in the scheme.

CH~

-H ~ Os CH3 I

H

O I

CH2

HCHO ~

v CH30H

/ xCH2 9 9

II

OI

CH

~ CO2+H2

/.

CO

Figure 5. Proposed mechanism of catalytic conversion of methane over ferric molybdate catalysts. According to the scheme, methane reacts with the surface of catalysts resulting in the formation of methoxy groups which are further oxidized to surface dioxymethylene species or adsorbed formaldehyde. The methoxy groups may react with hydrogen radicals or water to form methanol. The adsorbed formaldehyde may desorb or react with other surface oxygen to form surface formate species. The dioxymethylene group may decompose to carbon dioxide and hydrogen or be further oxidized to surface formate which later decomposes to carbon

224 monoxide. This proposed mechanism is in general agreement with the mechanisms proposed from the kinetic studies. 4. CONCLUSIONS Using in situ FT-IR spectroscopy, the gas phase products and the principal intermediates involved in the catalytic conversion of methane over ferric molybdate catalysts were identified and the reaction mechanism was proposed. In the absence of an oxidizing agent, methane reacts with the oxygen of the catalyst to produce methoxy species, which is an important intermediate for methanol formation. Further oxidation of the methoxy groups results in the formation of surface dioxymethylene, adsorbed formaldehyde, and surface formate species. The decomposition of surface dioxymethylene and surface formate species will give carbon oxides and hydrogen. Among the catalysts used in this study for the partial oxidation of methane, the higher methanol content can be obtained from ferric molybdate catalyst with Mo/Fe ratio of 1.7 and higher formaldehyde content can be obtained from ferric molybdate catalyst with higher Mo/Fe ratio. Using MOO3, formaldehyde and carbon dioxide were detected. Using Fe203, the main product obtained was carbon dioxide.

Acknowledgment Partial support from Phillips Petroleum and Monsanto Chemical Company is gratefully acknowledged.

REFERENCES 1. 2. 3. 4. 5. 6. 7 8. 9. 10. 11. 12. 13. 14. 15. 16. 17. 18.

A.S. Chellappa and D. S.Viswanath, Ind. Eng. Chem. Res., 34 (1995) 1933. R. Pitchai and K. Klier, Catal. Rev.-Sci. Eng., 28 (1987) 13. M.J. Brown and N. D. Parkyns, Catal. Today, 8 (1991) 305. D.A. Dowden and G. T. Walker, U.K. Patent No. 1 224 001 (1971). N.D. Spencer, J. Catal., 109 (1988) 187. N.D. Spencer and C. J. Pereira, J. Catal., 116 (1989) 399. N.D. Spencer, C. J. Pereira and R. K. Grasselli, J. Catal., 126 (1990) 546. H.F. Liu, R. S. Liu, K. W. Liew, R. E. Johnson and J. H. Lunsford, J. Am. Chem. Soc., 106 (1984) 4117. M.M. Khan and G. A. Somorjai, J. Catal., 91 (1985) 263. M. M. Koranne, J. G. Goodwin, Jr. and G. Marcelin, J. Catal., 148 (1994) 388. M. D. Amiridis, J. E. Rekoske, J. A. Dumesic, D. F. Rudd, N. D. Spencer and C. J. Pereira, AIChE J., 37 (1991) 87. M. A. Banares, J. L. G. Fierro and J. B. Moffat, J. Catal., 142 (1993) 406. M. A. Banares, I. Rodri'guez-Ramos, A. Guerrero-Ruiz and J. L. G. Fierro, in: Proceedings, 10th International Congress on Catalysis, Budapest, 1992, paper 076, Elsevier, Amsterdam, The Netherlands, 1993, pp.184. M. A. Banares, N. D. Spencer, M. D. Jones and I. E. Wachs, J. Catal ....146 (1994) 204. M. A. Banares, H. Hu and I. E. Wachs, J. Catal., 155 (1995) 249. V. Srihari and D. S. Viswanath, J. Chem. Tech. Biotechnol., 32(1982) 868. G. K. Boreskov, G. D. Kolovertnov, L. M. Kefeli, L. M. Plyasova, L. G. Karakchiev, V. N. Mastikhin, B. I. Popov, V.A. Dzis'ko and D. V. Tarasova, Kinetika i Kataliz, 7 (1966) 144. F. Trifiro, S. Notarbartolo and I. Pasquon, J. Catal., 22 (1971) 324.

225 19. M. A. Banares and J. L. G. Fierro, Prepr. - Am. Chem. Soc., Div. Pet. Chem., 37 (1992) 1171. 20. R. P. Groff, J. Catal., 86 (1984) 219. 21. J. S. Chung, R. Miranda and C. O. Bennett, J. Chem. Soc., Faraday Trans. I, 81 (1985) 19. 22. D. W. L. Griffiths, H. E. Hallam and W. J. Thomas, J. Catal., 17 (1970) 18. 23. C. H. Rochester and S. A. Topham, J. Chem. Soc., Faraday Trans. I, 75 (1979) 1073. 24. G. Busca and V. Lorenzelli, J. Catal., 66 (1980) 155. 25. F. S. Feil, J. G. van Ommen and J. R. H. Ross, Langmuir, 3 (1987) 668. 26. C. Mao and M. A. Vannice, J. Catal. 154 (1995) 230. 27. A. Ueno, T. Onishi and K. Tamaru, Trans. Faraday Soc., 67 (1971) 3585. 28. J. S. Chung, R. Miranda and C. O. Bennett, J. Catal., 11.4 (1988) 398.

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91997 Elsevier Science B.V. All rights reserved. Dynamics of Surfaces and Reaction Kinetics in Heterogeneous Catalysis G.F. Froment and K.C. Waugh, editors

227

DYNAMICS OF MULTI-COMPONENT ADSORPTION WITH I N T E R A C T I O N S 9A M E A N - F I E L D A P P R O A C H

M. Dubel and S.D. Prasad* Physical Chemical Division National Chemical Laboratory Pune - 411008 India. emaih [email protected]

SUMMARY The influenze of adsorbate interactions in Multi-Component Chemisorption is analyzed using a Bragg-Williams lattice Gas Mean-Field Approach (MFA). Using the lumping strategy adopted by Mobil's KINPTR reforming model, all the important reactions in reforming can be modelled, after accounting for the role of adsorbate interactions. There is a 104 order difference in the rates, if we include the role of adsorbate interaction-energies, the numerical magnittude of which is ~ 1-2 kcal/mol. A Monte-Carlo Procedure is adopted to test the feasibility of the method for a 13 component system.

I. INTRODUCTION In petroleum refor.ming a metal such as Pt (or Pt / Re) supported on a silica-Alumina or a zeolite catal,yst is employed for effecting the hydrocarbon conversions(I-5). The idea is to improve the Research Octane Number (RON) or the Motor Octane Nulnber (MON) of the final reformate. The metal function catalyses several dehydrogenation, dehydrocyclization, aromatization reactions, whereas the acid function effects the skeletal isomerization, mainly through a carbocation mechanism(2-5). Under actual conditions of reforming, several species are chemisorbed on the Pt/A1203 surface and they influence the overall dynamics(I-6). Significant adsorbate-adsorbate interactions can be present and they affect the sorption and reaction rates and equilibrium (6-11) markedly, especially for bi-molecular reactions. These also can have pronounced consequences in surface dynamics(7-16). While their roles in surface phase transformation~ have been demonstrated, using ideal single crystal surfaces(12,13), their consequences in surface dynamics are much less understood.(1,12-14) Whether they do matter in the dynamics of complex petrochemical conversions(2-5) where a medley of reactions is involved, is the problem we address in part. S.D. Prasad

Author for correspondence

228 II. FORMULATION

In a reforming scenario, more than 300 distinct chemical species are involved, with a huge array of reaction networks(3,5). Mobil KINPTR (An acronym for kinetic platforming Model) model, which is universally acknowledged to be an excellent one for predictive purposes, identifies 13 kinetic lumps(5). The predominant reactions involving these lumps are all illustrated in Table 1. These include the key comer-stone reactions of the Mobil's KINPTR(5) model and include some others which are also important. These include Paraffin ring closure and isomerization, Napthene- Isomerization and Dehydrogenation ,Cracking and Hydrogenloysis Dehydrocyclization and Aromatization, etc.. Disproportionation and alkyl group exchange form another group of interesting bimolecular steps. Table 1. KEY REACTIONS IN PETROLEUM RFFORMING ACCORDING TO KINPTR AND OTHERS 1. PARAFFIN RING CLOSURE

n C7H16

~

-"

._

n C7H16

~

'N ,,,-~H2C]-I 3 + H2

"/CH3

-"

+H 2

cc ccc C

2. PARAFFIN ISOMERIZATION

C--C--C--C--C--C--C

I C--C--C--C--C

I

3. NAPTHENE ISOMERIZATION

C

4. NAPTHENE DEHYDROGENATION

+ 3H2

229 5. PARAFFIN n CTH16

+

CRACKING

H2-+

C4HI0 +

C3H 8

<>

6. AROMATIC DEALKYLATION

C

H2 +

---

+ CH3

7. D I P R O P O R T I O N A T I O N

Ethyl Benzene

Benzene + p- Diethyl Benzene

8. AROMATIC ISOMERIZATION p-Xylene

-"

m-Xylene

-"

o- Xylene

9. HYDROGENOLYSIS Cyclohexene

~ n hexane

10. DEHYDROCYCLIZATION n hexane

methyl cyclopentane

n heptane

methyl cyclohexane

A wide variety of time scales are associated with these reactions, with Dehydrogenation being the fastest followed by Paraffin Ring Isomerization and Ring closure(I-5). Hydrogenolysis and Cracking are the slowest with Disproportionation also being slow. It is impossible to model all the reaction pathways in toto, so we focus only on a few typical reactions. The influence of adsorbate interactions in surface dynamics has been studied theoretically(7,11,15,16) and experimentally(I,6,12,14). A rough classification of the adsorbate interactions into two types viz. (a) through-space (b) through-bond, has been made for theoretical purposes(13). The former include the dipole-dipole, induced dipole-dipole, ioninduced dipole etc. as well as other multipole interactions. In contrast, the latter is of at least one order higher magnitude than the former, and they are mediated by sub-surface substrate electrons. For simple-system, it has been said that they amount to 1/10th of the total adsorbate-substrate binding energy, as has been shown through theoretical calculations(13). Thus the effects due to substrate mediated adsorbate interaction are best seen in the metal catalyzed reactions, of course, the assumption involved being that metal (Pt) clusters dispersed on the Alumina have electronic properties similar to the bulk platinum.(1)

230 There is considerable agreement among the workers, that the hydrocarbon skeletal isomerizations are carried out through a carbocation (carbenium ion) mechanism(2,3). The most common route for the generation of the carbenium ion is the dehydrogenation of an nalkane to form an olefin(2,3). Subsequent addition of a proton results in a carbocation (released from a Bronsted acid site in the zeolite or Alumina). These carbenium ions undergb rapid skeletal isomerization to give rise to branched paraffins and olefines. There can be repulsive interactions between adsorbed carbenium ions, and the ensuing activated complexes. The fact that partially charged activated complexes can also change the actual activation energy for the surface reaction has been demonstrated earlier(9). A simple Transition Rate Theory (TST) has been employed for this purpose, which give qualitatively correct results(9,10). III. The Model Isotherms : A Mean Field Approach The first step in understanding surface dynamics is to relate the gas and surface phase concentrations of the theoretically important lumps. It is assumed that the adsorption properties of molecules within a lump do not vary much. Further, all the sites involved in adsorption are assumed to be of uniform energy, and that no surface reconstruction(I) occurs following chemisorption. The concentration of unsaturated poly-nuclear aromatic structures (coke) on the surface is minimized by operating at high hydrogen partial pressures, as is customarily done in commercial reformers, i.e., catalyst fouling is ignored. We also ignore the role of hydrogen spillover and the remote control of reforming selectivity(4), so as to retain a tractable theoretical model. There are serious difficulties in the derivation of the multi-component adsorption (which is typical of the reforming scenario), taking into account the role of adsorbate interactions. The only mathematically tractatable way to do it is, by following the BraggWilliams lattice -gas Mean field field Approach (MFA). In this approach, the pair and higherorder correlations between the adsorbed molecules are ignored. With these assumptions, the adsorption isotherms within the MFA format can be easily written(8-11, 17) PA = bOA exp( -QA / RT ) ( 0A / 1- 0A - 0B ) exp[ C ( WAAOA+ WABOB ) / RT ]

(1)

PB = bob exp( -QB / RT ) ( 0B / 1- 0A - 0 B ) exp[ C ( WAB0 A q- WBB0 B ) / RT ]

(2)

W AA , WBB , W AB are the interaction energies for the possible molecular pairs, QA, QB are the heats of adsorptions for components A, B respectively, C is the coordination number, boA , boB are entropy change factors. C = 4, C = 6 for a square and triangular lattice respectively, eg. the (100) and (111 ) surface of an fcc metal. For the n component system, similarly we can write

Pi = boi exp( -Qi / RT ) ( Oi/1-

)"~0i ) exp[ C ( Wii 0 i -i- ~Wij Oj ) / RT]

(3)

231 For computation purposes, let us write all the interaction parameters, in the form of a matrix.

F [_Wni

Wli ........................ Wln l W = /W2i ....................... W2n

(4)

....................... Wan

It is obvious that the interaction matrix is symmetric, as Wij = Wji, being the interaction energy between i, j molecular pairs. The diagonal elements of the matrix W represent the pairwise interaction energy , which can be determined from single-component adsorption measurements. In general, n(n+l)/2 interaction parameters have to be specified for uniquely characterizing W. Since the focus is on the adsorbate interactions, the b0i, Q1 are fixed in the first set of computations. It is hard to obtain exact parameters for a given system, for the Xylenes in ZSM-5 the heat of adsorption is between 15 - 19.5 kcal/mol (7), which is a representative value, b0 = 10 6, and 10 8 ton- are the entropy factors. Adsorption equilibrium constants in the Henry's law range will be almost identical, with the choice of these values. IV. N U M E R I C A L

PROCEDURES

FOR ISOTHERMS

A Newton-Raphson algorithm was used to solve the set of non linear equations (1) - (3), i.e., to determine surface coverages as a function of partial pressures, temperature and interaction parameters. We define 0 , E as 01 02

0 =

"

E =

E1 E2

(5)

two column matrices, 0 ,E. For the former 0 the elements being the surface coverages, and for the latter E , the product of the appropriate row of the experimental interaction parameter matrix W and the 0 matrix, i.e.

Ei

=

exp [ C ( XWij 0j + Wii 0i ) / RT ]

(6)

and the matrix of input partial pressures Pl

p

(7)

P2

pI1

m

232 And the elements of the Jacobian matrix are given by J jk = - Pj- Wjk Ej 0j

j ~: k

(8)

Jii = Jik - Ej

j = k

(9)

and F is a column matrix whose elements are given by 9 Fi = 0i Ei - Pi v

(10)

where v is the fraction of vacant sites (scalar)

given by 9

v = 1 - E Oi

(11)

The procedure is quite straight forward, the 0 matrix is updated with AO matrix as computed by A0 = - J -~F

(12)

on successful convergence A0, F are both below the prescribed limits.

V. D E V E L O P M E N T OF K I N E T I C M O D E L S

In section IV we considered the numerical procedures for the computations of the surface coverages (adsorption isotherms ) as a function of component partial pressures and the interaction energy matrix W. The gamut of reactions presented in Table 1 can be broken down into the following types" (1) A (3)A

~ ""

(5) A + B

B

(2) A

B

(4) A

~ Products

B -"

(6)A + A

B

-"

C

~ Products

i.e.the sequence of reversible first order steps (Hydrogenation - Dehydrogenation) or irreversible second order steps: (eg. Disproportionation) The latter is a very useful diagnostic reaction for the presence of adsorbate-adsorbate interactions, as the rate of a bimolecular

233 reaction depends drastically on the molecular pair probabilities. The pair probabilities in turn sharply depend on the pairwise molecular interaction energies. One more added complication is that, the energy reference states of the reactants and the activated complex are shifted by adsorbate interactions. A quantitative treatment of the same effect using a Transition State Theory (TST) is published elsewhere(17), and we merely quote the results here

A

T

B

,i+_ r'.-:9

B e------ A - - - - B------o A

.+

A e----- A

A + B --~P Figure

o

A --~B

la

T

+-

B m - - - - - B ------o

lb

P o s s i D l e n e i g h D o u r h o o d s of an A-B p a i r and A , B in a s e c o n d o r d e r and r e v e r s i D l e m o n o m o l e c u l a r

B ---~A lc

molecules reactions

The scenario is best visualized with reference to Figure 1.a. Here we consider the activated complex for a bimolecular disproportionation reaction involving dissimilar molecular pairs. The six nearest neighbour sites can be populated by A , B molecules or they can remain vacant. The Figures 1.b and 1.c describe the activated complexes for the forward and backward reactions involving adsorbed A or B molecules. The four neighbouring sites can in turn be occupied by A, B or none. In finding the overall rates of reactions, we have to consider the probabilities of various configurations, and also consider the energetics of the interaction of the activated complex with the surrounding adsorbed molecules. The details have been worked out else where(17), so we give only the final expressions. rAB = k2 0A 0B (-1 - 0A- 0B + ~,10A + ~20B )3 (1 - 0 A- 0 B q- ~20A q- ~30B )3 where kl = exp( Wll / RT ) ;

k 2 = exp( Wl2 / RT ) ;

(13)

)',3 = exp(W22 / RT )

Similarly, for the first order reversible reaction, we have the net rate as the difference between forward and backward reaction. rnet --

rf - r b

(14.a)

234 rf

=

rb --

kfe A ( 1 -0A-0B + ;L10A+ X20B )4

(14.b)

kb OB ( 1 - 0 A- OB 4- X20A 4- X3OB)4

(14.c)

In the above expressions, the terms inside the braces give rise to the corrections of the activation energy due to adsorbate interactions. If the activation energy correction terms are neglected, then ,only the surface coverages reflect the role of adsorbate interactions, and equations (13) and (14) reduce to 0

r AB = k2 0A OB r 0net = k f 0 A

(15)

(16)

k b OB

V l . RESULTS It is implied that we are calculating steady-state surface rates under conditions of adsorption equilibrium, and surface reaction steps are slow compared to, those of adsorption and desorption. q) 2.2 ILl < rY LU > 0 (o It_ n, (D LL 0 0 F
i

!

1PI= l"SE4,P2

0"45

l

=9E4

Torr

2P1= 1 - 5 E 4 ) ~ 1"8

oJ <~

1.4

-1I-v,-7I.--

1"0

m

i

~'~ 0-35

0-25

1 theto 1,2theto 2

2

o~ 0 . 6

i

-0".5

"

I

0

I

0"5

Figure 2 INTERACTiON P A R A M E T E R

1"0

W12

i

o.~5

-1

-0

ts

6

01-5

Figure 3 INTERACTION P A R A M E T E R

1"0

W12

In Figure. 2 the ratio of 02 / 0~ is plotted against the interaction energy for the unlike pair W12 varying from -0.8 to 0.8 kcal/mol at a temperature of 7500 K and PA = 1.5 X 104 torr, PB = 9 X 104 tort, which are typical of reforming conditions. Wll -- lkcal / mol and W22 = 2 kcal/mol, illustrating that repulsive interactions are operative between the like pairs Wl"l and W22. It is obvious that the ratio 02 / 0] increases from 1.23 to 1.9946 as Wl2 changes sign (viz. from attractive to repulsive), all other parameters being the same. There is a dearth of literature on the exact values for a specific system, but the values are typical and the trends too! In Figure. 3 we investigate the ratio of 02 / 01 but for equal partial pressures Pl = P2 - 1.5 x 104 torr. The trends are opposite. The ratios are much different and fall as W]2 becomes more positive from -0.8 to 0.8 kcal /mol, i.e. from 0.74 to 0.64 (As before the

235 interactions change from attractive to repulsive ). The foregoing discussion shows that when interaction effects are operative there can be very complex behavioural pattems. Perhaps it is more instructive to compute the product 01 02, as this will be directly proportional to the rate of the bimolecular reaction, in a first approximation (see equation 15 ). The lowest and highest values for the case envisaged are 0.1575 and 0.5793 for Figure. 2. The corresponding values are 0.054,0.111 for Figure. 3. Thus the products 0 1 0 2 show much more drastic pattems of bahaviour, in comparison to the surface coverages. This fact is exploited later, when we refine the treatment of bimolecular reactions, so as to include activation energy correction factors. 1. The Reversible Monomolecular Reaction Many reactions like the skeletal isomerizations, hydrogenation-dehydrogenation can be described by sequence of first order reactions (see eg. reaction sequence 2, Table 1.). The expressions for the reversible rates with and without activation energy corrections are given by equations (14) and (16) respectively.

Table 2 The Dependance of a Reversible Surface Reaction (Isomerization) with and without Activation Energy Connection on Component Partial Pressures W~= 1 k cal/mol ; W~2 = 0.5 k cal/mol ; W22 -- 2 kcal/mol ; kf = 0.8 x 10 -7 mol/gm.sec ; k b = 0.2 x 10 Tmol/gm.sec. PA torr

PB torr

10 4

5 10 4

10 3 10 2 10

5 10 ~ 5 102 50

0.1

0.5

0A

0.5323 0.5299 0.51007 0.4178 0.25066 9.7318 10.5

0B

0.4673 0.4662 0.4567 0.41048 0.31416 0.18956

Rate with out Rate with activation Activation Energy Energy Correction, 10 -7 Correction, 10 -7 mol/gm.sec mol/gm.sec 0.2389 121.4 0.2374 120.46 0.22538 112.51 0.17008 79.537 7.4865 10-2 34.6708 2.0309 10 -3 8.4838

Table 2 lists the general behaviour, as the component partial pressures are varied over 5 orders of magnitude. The rate constants for the forward and backward reactions are 8 x 10-7 mol/gm.sec and 2 x 10-7 mol/gm.sec as is typical for an isomerization reaction. The interaction parameters chosen are Wll = 1 kcal/mol ; WI2 = W2~ = 0.5 kcal/mol ; W22 = 2 kcal/mol. This corresponds to the presence of repulsive interactions between all possible pairs. Notice the drastic enhancement of the rates, when activation energy corrections due to repulsive interactions are incorporated, by almost 4 orders of magnitude. A simple physical interpretation of this is, that the repulsive interactions push up the reactant molecules on the energy scale, such that the activation energy is lowered ; of course, the interaction energy of the activated complex with the adsorbed molecules is neglected. Both the rates with and

236 without activation corrections decrease monotonically as the partial pressure is decreased. But the order of magnittude difference between r0net, rnet is present at all the pressures.

o

(_) 0 " 2 5

(.9

0-20

L9 .d

J

0 :s: 0 " 1 5 0 n,, o

0 o (lc o

o-~o-

,

,

I--
0.05

o !z

o'.4

0-6

COVERAGE

I

10075 50

z_ ua 25

z LIJ

125

Itl

la.I (/)

-

I--

0

n0-2

0

0"4

Figure 4"SURFACE

0.6

0

o

Figure 5"SURFACE

COVERAGE 1

In Figures. 4 and 5 the r0net , rnet with and without activation energy corrections are plotted against 01, for the pressure range of interest. Once again these plots illustrate what has been said before. Adsorbate interactions are important even in a monomolecular reaction net work. Ideally, we should have similar plots against 02. But since the qualitative features are the same, we defer it here. Table 3 Partial Pressure Dependance of Reversible monomolecular Rate with Attractive non-identical molecular pair-wise interaction energy W12 = -0.5 kcal/mol. The rest of the parameters are like in Table 2. PA tort

10 4 10 3 102 . .

10 . . . 1 0.1

PB

torr

5 10 4 5 10 3 5 10 ~ 50 5 0.5

0g

0B

0.5195 0.5192 0.516 0.49103 0.36724 0.1542

0.48054 0.4802 0.478 0.4617 0.378319 0.2185

Rate with out Rate with activation Activation Energy Energy Correction, 107 Correction, 10-7 mol/gm.sec mol/gm.sec 0.2234 82.543 0.2232 82.447 0.2217 81.506 0.208 73.56 0.14247 42.02 3.59 E-2 10.36

In Table 3 we change the sign of W~2 ---0.5 kcal/mol to make it attractive instead of repulsive, All other parameters are retained. The repulsive interactions between like pairs Wll , W12 still dominate the rates. Once again the correction terms in activation energy are

237 significant. Figure. 6 and 7 give r0net, rnet Vs the surface coverage 01. A quick comparison to Figure. 4 and 6 show that, the rates without activation energy corrections display only the changes in surface coverages 01 , 02 [i.e. terms of the form kf 01 - kb 02 )] and they are marginal. However an Analysis of the Figure. 5 and 7, show clearly that, a fuller accotmting of the subtleties of the interaction effects have to be done through equation (14). The differences between Figures. 4 and 6 are significant, so as to warrant attention. Thus we conclude, that to account for interaction effects, one has to indeed use the more refined model given by equations (14.a - 14.c) o

W or)

s

0"25

|

|

,

o

s

0-20

75

J

,

|

i

0

0 rr L.)

n,"

_o 0.10 z uJ 0 - 0 5 I0

|

W l l =1~ W12 = W 2 1 = - 0 - 5 W

_..I

0 :~ 0-15 0

n,,

100

(13

50

z_ 25 t~

olz

'

0:4

'

Figure 6" SURFACE COVERAGE 1

0-6

a:

0

o!2

'

" o14

'

0-6

F i g u r e 7 SURFACE COVERAGE 1

2. Bimolecular Reactions The influenze of the adsorbate interaction on a bimolecular step is best illustrated for a disproportionation reaction. (see Table.1 and reaction sequence 5, 6). The corresponding rate expressions without activation energy correction and with activation energy correction are given by equations (13) and (15) respectively. Figures. 8 and 9 give the corresponding rate plots Vs partial pressure ( PA = PB ) of either component in the same order. As is obvious from the plots, there are 4 orders of difference between the rates. Thus we conclude that irrespective o f the order o f reaction steps, there is significant variation in the rates spanning 4 orders o f magnitude, if we incorporate the activation corrections due to adsorbate interactions, in a more accurate rate model.

VII. C O N C L U D I N G R E M A R K S The greatest barrier in the application of the Multicomponent Fowler-Guggenheim or Bragg-Williams Lattice gas model to, a practical situation like Pet-reforming, is the absence of experimental interaction parameters. In the simulations of the earlier sections, representative values were used. In general, for an n component system , we need to fix n(n+ 1) / 2 interaction parameters of the symmetric W matrix (91 for a 13 component Model ) .Mobil has used successfully a 13 lump KINPTR model(5), which essentially uses a HougenWatson Langmuir-Hinshelwood approach. This results in a psuedo-monomolecular set of reactions, which is amenable to matrix analysis.

238 The presence of adsorbate interactions make this approach fail totally, as the isotherms are non-linear. The Newton-Raphson algorithm is a practical method of calculation for surface coverages of a 13 component FG isotherm. We make the further assumption that for species within a given lump, the adsorbate interaction energies are identical. In addition, we need to specify only the interaction energy between dissimilar pairs. Thus we have recourse to a Monte- Carlo Procedure(15) The Monte-Carlo Procedure First, we fix the Wii (diagonal elements) of the interaction parameter matrix. This represents the average interaction energy within a lump (i.e. interaction energies of like pairs). Then the off-diagonal elements are generated by the following procedure:

Wij = [ RAND(X) - 0.5 ] scale

(17)

where scale is a scale factor determining the magnittude of interactions. RAND(X) is a uniformly distributed random generator, which returns a value from 0 ~ 1. We also ensure that the W matrix is symmetric, i.e.

w~j = wj~

(I 8)

This method has been used to test the convergence of the Newton-Raphson Algorithm for a 13 component system. In most cases, a few iterations are sufficient to guarantee the convergence. The most interesting simulation corresponds to the l Wij [ = constant = scale, viz. attractive or repulsive. Here only the signs of the Wij parameter, are determined by RND(X). Table 4 gives the influenze of the W matrix, on the surface coverages at a constant temperature and input partial pressures. As is obvious from the Table 4, only a few surface concentrations are appreciable. Since the bimolecular rate processes depend on the product of surface coverages, in principle only a handful of such processes are permitted. Another possibility for the multi-component FG isotherm is the existence of multiple solutions. This is illustrated by the two component FG model with predominantly attractive interactions, p~ -- P2 = 10 torr, W~ = -1 kcal/mol, W~2 = 0.5 kcal/mol, W22 = -0.5 kcal/mol, b0 = 108 torr, RT = 1 kcal/mol, Q = 19.5 kcal/mol. The first set of solution is 01 = 0.3232,02 = 0.6653. The second set is 01 = 0.9968,02 = 2.544 10-3. The possibility of phase separation is real, but we do not pursue it any further.

239

Table 4 Table Illustrating the Monte-Carlo Simulation of the Role o f Adsorbate Interaction in determining Surface Coverages in a 13 lump model.

1

1 10 ~

Trial 1 Surface Coverages 0.25 10-~

2

2 10 3

0.38 10 .02

3

3 10 3

4

Component Name

Partial Pressure, tort

0.10632 0.404 10 ~

0.2518 10 .2

0.635 10 -02

0.8437

10 ~

0.5923 10 -2

4 10 3

0.4186 10 "02

0.1248

10 -~

0.3075 10 1

5

5 10 3

0.4565

10 .02

0.1035

10 -~

6

6 10 3

0.7521

10 ~

0.2277 10 "~

0.2052 10 -1

7

7 10 3

8

8 10 3

0.5546 10 "~ 0.5067 10 ~

0.1075 0.6941

0.1615 10 -1 0.6496 10 -1

9

9 10 3

0.8161

10 11 12

10 10 3

J

o

0.1834

0.1186

11 10 3 12 10 3 13 10 3

0.1815 0.62361 10-1

0.8732 l0 -~

0.1048 0.1687 10 -~

IW,jl = o.5

bo = 108 t o n

0-25

o

ILl

n~

0.8553

10 -1

0.3669

Q = 19.5 kcal/mol

150

.

.

.

.

W11= 1~ W12 = W 2 1 = 0 - 5

~

.

,

0-20 D

0-15-

0-05

o_ z

0

100

0 n,'

~-- 0 " l O -

F

0.4078

0.133645

0 r

z

10 -~ 10 ~

0.2597 10-~

03

s

10 ~

0.425 10 -1

0.2210 0.2325

13

o u.I

Trial 3 Surface Coverages 0.181 10 .2

Trial 2 Surface Coverages

J

62

50

uJ FI

I0 ~

I

10 z

F i g u r e 8 " LOG P R E S S U R E ~ T O R R

<~

104

13::

C)

-i0-I

10 o

Figure 9

~

10 ~

i

10 z

I

103

LOG P R E S S U R E , T O R R

104

240 REFERENCE

1. G.A. Somorjai, Introduction To Surface Chemistry And Catalysis, Wiley-Inter Science Publ. , New York, 1994, pg. 500 2. F.G. Ciapetta, D.N. Wallace, Catalytic Naptha Reforming, Catal. Rev. 5 (1972)67 3. G.A. Mills, H. Heineman, T.H. Milliken, K.G. Obald, Ind. Eng. Chem. 45 (1953) 135 4. B. Delmon, G.F. Froment, Catal. Rev. 38 (1996) 69 5. M.P. Ramage, K.R.Graziani, P.H.Schipper, F.J. Krambeck, B.C. Choi, in "Advances In Chemical Engineering", ed. J. Wei, Academic Press, New York, 13 (1987) 193 6. M.A. Vanhove, R.J. Koestner, J.C. Frost, G.A. Somorjai, Surf. Sci. 129 (1983) 482; C.M. Mate, G.A. Somorjai, H.W.K.Tom, X.D. Zhu, Y.R. Shen, J. Chem. Phys. 88 (1988) 441 7. K.P. Schroder, in "Studies in Surface Science And Catalysis", ed. G. Ohlmann, H. Pfeifer, R. Fricke, Elsevier, Amsterdam 65 (1991) 435 8. A.S. Datar, S.D. Prasad, Chem. Phys. Lett. 181 (1991) 445 9. A.S. Datar, S.D. Prasad, J. Chem. Phys. 96 (1992) 2387 9Ibid. 100 (1994) 1742 10. A.S. Datar, S.D. Prasad, I & EC Research. 31 (1992) 2257 11. A.S. Datar, S.D. Prasad, Langmuir 7 (1991) 1310 12. K. Binder, in "Phase Transitions and Critical Phenomenon", eds. C. Domb, M.S. Green, Academic Press, New York, 1976 13. T.L. Einstein, Crit. Rev. Solid State Sci. 7 (1978) 261 14. G.Ertl, Surf. Sci. 152/153 (1985) 328 15. M. Siverberg, A. Ben-shaul, F.J. Rebentrost, J. Chem. Phys. 93 (1990) 6501 M. Silverberg, A. Ben-shaul, J. Chem. Phys. 87 (1987) 3178 16. H.C. Kang, T.A. Jachinmowski, W.H. Weinberg, J. Chem. Phys. 93 (1990) 1418; H.C Kang, W.H. Weinberg, M.W. Deem, Ibid. 93 (1990) 6841 17. A.S. Dater, S.D. Prasad, Submitted for Publication

91997 Elsevier Science B. V. All rights reserved. Dynamics of Surfaces and Reaction Kinetics in Heterogeneous Catalysis G.F. Froment and K.C. Waugh, editors

241

M o d e l s of a d s o r p t i o n k i n e t i c s on r o u g h s u r f a c e s

Massimiliano Giona" and Alessandra Adrover Centro Interuniversitario sui Sistemi Disordinati e sui Frattali nell'Ingegneria Chimica c/o Dip. Ingegneria Chimica Universit~ di Roma '; La Sapienza" via Eudossiana 18, 00184 Roma. Italy * Permanent address : Dipartimento di Ingegneria Chimica Universit~ di Cagliari. piazza d'Armi. 09123 Cagliari, Italy This article analyzes adsorption kinetics of fractal interfaces and sorption properties of bulk fractal structures. An approximate model for transfer across fractal interfaces is developed. The model is based on a constitutive equation of Riemann-Liouville type. The sorption properties of interfaces and bulk fractals are analyzed within a general theoretical framework. New simulation results are presented on infinitely ramified structures. Some open problems in the theory of reaction kinetics on fractal structures in the presence of nonuniform rate coefficients (induced e.g. by the presence of a nonuniform distribution of reacting centres) are discussed.

1. I N T R O D U C T I O N Many industrial adsorbents and catalysts exhibit fractal properties in their surface structure over a given range of lengthscales [1]. The surface roughness influences both thermodynamic, kinetic and transport properties. Most of the literature on the dynamic properties of solid interfaces is based on the analogy between membrane and interracial kinetics and the electric response of rough electrodes, for which an anomalous impedance scaling with frequency w, referred to as Constant Phase Angle behavior. Z(w) = Ro + c~(iw)-~, 0 < w < 1, i = ~ has been observed [2-3]. The exponent w depends on the roughness of the electrode and ultimately on its surface fractal dimension do. See [4] for a brief critical analysis of the different models proposed for the relationship between w and d~. In a recent paper, Seri-Levy and Avnir [5] propose a model for adsorption kinetics in diifusion-limited conditions. This model is based on an extension of Delahay's analysis [6], which is valid for fiat surfaces, so as to include the effects of fractality. In the article by Seri-Levy and Avnir, the resulting expression for the behavior of the fractional coverage F(t)/Feq vs time t is a modification of the equation derived by Delahay and Trachtenberg [6] in order to include the anomalous scaling of the adsorbed volume with the fractal

242 dimension, which may be related to the electrochemical results obtained in terms of the scaling of Z(w) with w. This article focuses on an approximate mathematical model of adsorption kinetics on rough (fractal) interfaces based on a constitutive equation of Riemann-Liouville type [7]. The sorption properties of interfaces and bulk fractaJs are discussed within the same framework, and some new simulation results on infinitely ramified fractals are also presented. Finally, we close with a brief digression on open problems in the presence of heterogeneity in kinetics. 2. A D S O R P T I O N

KINETICS

ON ROUGH

FRACTAL

SURFACES

In this section we discuss in some detail a macroscopic approximate model for adsorption kinetics on fractal interfaces. Further details can be found in [8]. In diffusion-limited conditions, the balance equations for adsorption on flat surfaces take the form

Oc

OZ =

OJ Oz

z e

(0. oc).

c(~. t) =

c~

r(t) = g c ( 0 , t) = -

~t

g(0,

T)d7,

(1)

where c is the concentration in the bulk fluid phase, J the flux (J - -DOc/Ox), F the adsorbed concentration, and K the Henry constant. The basic model assumption is that the presence of a solid surface modifies the correlation properties in the motion of diffusing/adsorbing particles in the neighbourhood of the interface. This is a consequence of the adsorptive trapping of a fractal adsorbing surface. In macroscopic (approximate) mean-field modeling, in which the fractal surface is mathematically viewed as fiat. i.e. located at x = 0, the anomalies in the adsorption kinetics can be regarded as deriving from a modified non-Fickian constitutive relation. More specifically, it can be shown from scaling arguments (see [8] and references therein), that the flux (J) / concentration gradient (Vc) constitutive equation can be written as J(x. t) = - - ~ f ( t ) 9 Vc(x, t ) .

(2)

where D R , Vc is a Riemann-Liouville (fractional) [7] convolutional operator ~r(t).Vc(x.t)=

0 fOt (t-r)-~Vc(x't)d7" r ( 1Dfo - .) 0~

(3)

where DFo is a constant, F(1 - u) the Gamma function of argument 1 - v, and the exponent u is given by

u = do - d T ,

(4)

where d, is the fractal dimension and dT the topological dimension of the surface, d r = d - 1, where d is the dimension of the embedding Euchdean space. In the case of fractal interfaces, the anomalies in diffusive motion are localized within a thin section located close to the interface and referred to as the fractal layer, since diffusion is perfectly regular (Fickian) in the bulk fluid phase. The existence of a fractal laver close to the interface can be described with reference to two models of increasing complexity [8].

243 0.8

o

.

0.7

r(~)/r~

0.6 0.5 0.4 0.3 0.2

1

l0 s

I

L

1

1

I

210 s 3105 410 s 510 s 610 s 710 s T

Figure 1: F(T)/Fe vs ~- obtained from Monte Carlo simulations (dots) and the twotimescale model (solid line) by Laplace inversion of eq. (6), with Do = 3.5 10 -a and DFo = 0.71. The simulation parameters are: c~ = 4.88 10 -4, K - 20, D = 1. The simpler model (referred to as the two-timescale model) describes the presence of the ffactal layer as an additional anomalous term in the constitutive equation: the constitutive equation is therefore a superposition of a Fickian term and of a RiemannLiouville contribution (of the form of eq. (2))"

J(x, t) = - D o Oc(x" t)

o-------i--- ~ ( t )

Oc(z. t)

,

o~

(5)

where T~p(t) is given by eq. (3), Do is a constant, and the exponent L, equals d,, - d T > 0 eq. (4). In the Laplace domain (here s is the Laplace variable conjugated with the time t), the solution of eq. (1). with the constitutive equation given by eq. (5), is given by

where A

OH(s) = Do + DFos ~

(7)

A

In eq. (6), F(s)/Fr is the Laplace transform of the fractional coverage, and Fe = K c ~ In the second model, referred to as the ffactal layer model, the presence of the fractal layer is explicitly considered. The model is therefore characterized by the presence of three phases: 1) a bulk region, far from the fractal interface, in which the diffusional phenomena are regular, i.e. not influenced by the fractality of the solid surface, and Fick's law applies; 2) a fractal region (the fractal layer), in which anomalous diffusion occurs, described by means of model eqs. (2)-(3); 3) an adsorbed surface phase. In this way, a finite thickness of the ffactal interface LF is introduced, delimiting the region in which anomalous diffusion occurs. The resulting balance equations should therefore read

244 101 A

Do~(s) 100

10-1

10-2 . . . . . . . . . . . . . . . . . . . . . . . . . . . 10 -6

10-4

10-2

,

.

10 0

. .

10 2 8

Figure 2: Comparison of the effective diffusivities De~(S) derived from the models and from Monte Carlo simulation results (o). Simulation conditions: c~ = 3.25 10 -4, K = 2 0 , D = 1. Solid lines represent (a) the fractal layer model, obtained by solving eqs. (8)-(11) in the Laplace domain; (b) the two-timescale model, eq. (6). Fitting parameters: (a) Dfo = 0.95. LF = 7; (b) Do = 0.013. DFo = 0.9. as follows. 1) In the bulk phase ( z E [0. co) ):

Oc 02c O---t= D Oz 2 ,

(8)

where D is the classical diffusion coefficient. 2) In the fractal layer with thickness LF, q(x, t) being the concentration in this phase:

02q

0q & = :DR * Oz2'

(9)

where x E [0, LF]. The spatial coordinates x and z are, of course, related by z = x 3) In the adsorbed phase, characterized by a surface concentration:

r(t)=

~r*

~

LF.

~=0

As in the case of flat adsorbents, the boundary and initial conditions are given by c ( ~ , t) = c ~ r(t) = Kq(O, t) for t > 0, c(z, O) = c~ for z e (0, oc), q(x, 0) = c~ for x e (0, LF). The continuity between the fractal layer and the bulk phase implies that

~(0, t) = q(LF, t)

D

= Z~F 9 ~ z=0

(11) x=LF

The solution of this model in the Laplace domain can be found in [8]. The prediction of Riemann-Liouville approximate models may be compared with Monte Carlo simulation data, obtained on a self-similar fractal interfaces (Koch-like curve) possessing a fractal dimension d~, = log 8/log 4 - 3/2. Figure 1 shows the comparison of

245 Monte Carlo simulation data and the two-timescale model, eq. (6) in the time domain, for which a closed-form expression of the fractional coverage may be obtained [8]. In this figure, ~- is the dimensionless time t D / L 2, where L = 1 1.u. It can be shown from dimensional analysis and from simulation results that the two parameters entering into the model, Do, and DFo, depend upon the bulk ditgusivity as

Do = qoD

Dto = ql D" ,

(12)

where qo. ql are constant. In the case of the fractal layer model, no closed-form expression for the inverse Laplace transform has been obtained so far. Nevertheless. a useful way to compare the model and the simulation results directly in the Laplace domain can be obtained by considering the effective diffusivity in the Laplace domain De~(s) defined as

A s3~2(s)K 2 De~(s) - (1 - sF(s)) z

(13)

Figure 2 shows the comparison of the fractal-layer (solid line a) and two-timescale (solid line b) models with the simulations in terms of effective diffusivity, eq. (13). Both the models furnish a satisfactory level of agreement with simulation data. We may therefore conclude that approximate models based on a Riemann-Liouville constitutive equation are able to furnish an accurate description of adsorption kinetics on fractal interfaces. These models can also be extended to nonlinear problems (e.g. in the presence of nonlinear isotherms, such as Langmuir, Freundlich, etc.). In order to extend the analysis to nonlinear cases, efficient numerical algorithms should be developed to solve partial differential schemes in the presence of Riemann-Liouville convolutional terms. 3. S C A L I N G T H E O R Y OF S O R P T I O N This section analyzes the scaling properties of the uptake curve M ( t ) / M ~ on/across fractals in a single theoretical framework. The (fractional) uptake curve M ( t ) / M ~ is the ratio of the solute quantity entering the structure up to time t and the quantity entering at saturation (i.e. at t -+ oo), i.e.

M(t)/Moo = [ t J(O, r ) d r / [~176 J(O, r)dr. JU

/

Let us first consider sorption properties of fractal interfaces. follows that

M(t)/Moo

~

(14)

JU

From eqs.

(2)-(3) it

(Is)

t (dT+l-d=)/2

It is interesting to compare eq. (15) with the results obtained on finitely ramified fractals by means of Green function renormalization [9-10]. It has been shown that the fractional uptake curve for a structure possessing fractal dimension d;, walk dimension dw, and adsorbing from a reservoir at constant concentration co through an exchange manifold B (which represents the permeable boundary for transfer) possessing fractal dimension d~ scales as M(t)lMoo

~ t(d'-d~ I/~

~ t~ 9

(16)

246

Figure 3: Infinitely ramified fractal structure considered in section 3. Equation (16) also holds for transfer across fractal interfaces. In the latter case, the fractal dimension df refers to the (Euclidean) bulk in which particles diffuse, and is given by d f = d = dT + 1. d~ = 2. while d~ equals the fractal dimension d~ of the interface itself. An open problem in the theory is to ascertain whether eq. (16) also holds for infinitely ramified fractal structures. This problem can be tackled either by means of Green function renormalization or by considering lattice simulations. Green function renormalization can be applied to product lattices by making use of the extension theory [11]. Some preliminary calculations (performed by considering the feed from a single site, i.e. d~ = 0) suggest that eq. (16) may also hold for infinitely ramified fractals as long as (df - d ~ ) / d w < 1, [12]. If (df - d ~ ) / d ~ > 1, then M ( t ) / M ~ ,,~ t. Numerical results can be obtained by solving the diffusion equation with the prescribed boundary conditions for batch-sorption experiments on a lattice approximation of the structure, by making use of finite-difference methods [13]. Figure 3 shows an infinitely ramified fractal possessing fractal dimension df = log 12/log 4. Its walk dimension d~ can be obtained numerically by considering the scaling of the mean square displacement R2(t) = f l x l 2 c ( x , t ) d x starting from an initial Dirac's delta concentration pulse c(x, t)t=o - 6(x). located in the center (of mass) of the structure, x - 0. Figure 4 A) shows R2(~-) vs T = t D / L 2 on the structure depicted in figure 3, from which a fitting exponent 2 / d ~ - 0.87 is obtained, i.e. d~ = 2.30. In the case of sorption from the entire external perimeter, the dimension of the exchange manifold B is given by d~ = log3/log4, and therefore ( d I - d ~ ) / d ~ - 1/d~ = 0.435. Figure 4 B) shows the fractional uptake curve obtained numerically on a 256 • 256 lattice (corresponding to the 4th iterative in the construction process of the structure), compared with the theoretical prediction eq. (16), with/3 = 0.435. The agreement between theory and simulation is satisfactory., and a well-defined power law exists over more than one decade. In any case, the lattice structure considered is fairly small (corresponding only to n = 4), and this may account for the small initial deviations between theory and simulation. The result shown in figure 4 suggests that eq. (16) also holds for infinitely ramified fractal structures, i.e. that [12] Z = max(l, (df - d ~ ) / d w } .

(17)

247 10 3

. . . . . .

,,,..

. . . . .

10 0

I02

M~

I

i01

i0_ l

10 o

I0-~0 6 ..... -

10-5 10-4 i0-3 10-2 7" 10-I

I0 2

10-4

I0-3

I0-2

I0-I 7" 10~

Figure 4: A) R2(T) vs 7 = t D / L 2 on the ffactal structure of figure 3 (dots). The line corresponds to R2(T) ~ ~_2/d~. with 2/d~ - 0.87. B) M ( T ) / M ~ vs ~- for the same structure (dots). The line is the theoretical prediction, eq. (16). with/3 = 0.435. It should also be observed that. in the case of an infinitely ramified ffactal structure fed from all the perimeter sites, an exponent/3 fairly close to fi = 1/2 (i.e. the regular case) is obtained. 0.40 < /3 < 0.50. The "small" difference between the exponent fl obtained for infinitely ramified fractals and in the regular Euclidean case makes it difficult to discriminate between the two scaling behaviors, especially if one makes use of numerical simulations on small lattices. 4. H E T E R O G E N E O U S

REACTION KINETICS-

OPEN PROBLEMS

The study of transport schemes by means of Green function renormalization reveals the analogies existing between sorption properties and the scaling of the effectiveness factor vs the Thiele modulus in a diffusion-limited regime [9-10]. In the case of uniform elementary kinetics (here uniform means that the rate constant is site-independent, and elementary that the reaction rate is of the form - r ( c ) = kc"), the effectiveness factor 77 scales with the Thiele modulus r (in the presence of the same feed configuration, i.e. of the same boundary conditions) as 7/,-~ r

_ r

(lS)

where the exponent/3 is given by eq. (17). The properties of elementary reaction kinetics on fractal media are in no way exhausted by eq. (18). New phenomenologies may arise due to heterogeneity in the distribution of reacting centres. The purpose of this section is to addressing these issues, focusing on interesting new features induced by heterogeneity in the distribution of reacting centres.

248

(o)

(k) 1

3

5

(k) 6

4

2

Figure 5: Construction process of the structure considered in section 4" a one-dimensional line with a distribution of reactive centres localized on a Cantor middle-third set. The two graphs labelled (k) are copies of the structure at iteration n, while the graph labelled (0) is a one-dimensional chain at iteration n with no reactive centres. In real catalysts, the distribution of reacting centres is likely to be spatially heterogeneous. Simple gedanke models can be tackled by means of Green function renormalization and can be investigated in order to understand the influence of spatial heterogeneity on the macroscopic behavior of the system. Fractal models of heterogeneous distributions of reacting centres have been also considered by Gutfraind et al. [14]. In particular, let us consider a fractal structure G (possessing dimension dl) and a distribution of reacting centres localized on a submanifold GR (the reactive manifold possessing fractal dimension dR ~_ dy). The corresponding balance equations attain the form dci dO = ~ g , jcj - pc~ + r ~ x,jcj R + Z E + y_, (Co - cj) , (19) J J J J -

-

where c~ is the concentration at the i-th site, 0 = t D / A x 2 the dimensionless time. D the diffusion coefficient, Ax the distance between nearest neighbouring sites, Hij the entries of the adjacency matrix of G, p the coordination number of G, ~2 = r ' r = ( k L 2 / D ) 1/2 the Thiele modulus, L the characteristic size of the structure, h the dimensionless mass-transfer coefficient, s the set of external sites of the structure, and X~, XZ, ~ the characteristic functions of the manifolds B, ~R, and s i.e. X~ =

1 ff(i,j) eB 0 otherwise.

(20)

and analogously for XR, ~ . In the case dR < d I, the dimensionality of the reactive manifold modifies the macroscopic properties, as expressed e.g. by the scaling of r/vs r in a diffusion-limited regime. To give a simple example, let us take ~ equal to the one-dimensional line (this example corresponds to the classical slab-like model for a cylindrical pore), and let ~R be a Cantor middle-third set, dR = log2/log3 = 0.631. Figure 5 shows the construction process of the structure applied in the renormalization process. The graph of the structure G(,+l) at iteration n + 1 is obtained by connecting two copies of g(") with a one-dimensional chain G(,+I) possessing the same number of sites of G("), i.e. N, = 3", but no reactive sites. The effectiveness factor (in the case of feed from site i = 1, and in the limit for h ~ oo) is given by the known expression [9] 1 - t-(d) "~l, ( r -,1, 9 (r

~) 2)

1 N

( L )2

(21)

249

~

7? 10~

b) ~ = 12 c) n - - 1 4

10-1

"~,~ ~ a)'X~'~

_. )

i

t 10-2

10-1

10 ~

101

10 2

10 a

10 4

Figure 6: Effectiveness factor vs the Thiele modulus for a distribution of reacting centres on the one-dimensional line localized on a Cantor middle-third set. Curve d) is the theoretical behavior, eqs. (18),(23). where ~_(d) V l l is the sorption Green function at site i = 1 [9] ,. evaluated for s = 0 (i.e. at steady state), and N the total number of sites of G . The renormalization recursions for the structure depicted in figure 5 are given by 2

X = x + y ( y x o + xy(yo~ - Xo)) (1 -

~o)

~ -

(xyo)~

y2

r (1 -

~o)

Y~ ~ -

(22) (xyo) ~ '

where x = Gxl-(n). Y = G(~), capital letters X. Y refer to the Green functions at iteration n + 1 a n d Xo, yo are the Green functions Qlx, G21, of the unreactive structure G(n) (the renormalization recursions for Xo a n d yo can be obtained from eq. (22) by substituting x ~ Xo, y --+ Yo, X ~ Xo, Y ~ Yo). The sorption boundary transformations are those typical of the one-dimensional line [9] (zero-flux boundary condition at site i = 2), xCd) ~_Cd) ~-~11 = X -~- y 2 / ( 1 -- X) yCd) _ ~_(d) y / ( 1 - x) and the initial conditions of the renormalization recursions for s = 0 are given by z = y = (p + r Xo = yo = p-1 Figure 6 shows the behavior of the effectiveness factor vs the Thiele modulus for this structure. It can be proved analytically (although the proof is not reported here for the sake of brevity) that in a diffusion-controlled regime the exponent r appearing in eq. (18) is given by =

2dR l§

= 0.7737.

(23)

This result should be compared with ~ = 1. corresponding to a uniform distribution of reactive centres, and indicates clearly that the presence of a heterogeneous distribution of reactive centres localized on a fractal submanifold modifies the scaling properties in the diffusion-limited regime. The relation between the structure of the reactive manifold and the exponent r appearing in eq. (18) is still an open problem with significant theoretical and practical impl_ications in the study of catalytic processes. The expression, eq. (23), for the exponent in the specific case study considered depends on the linear topology of the structure in

250 which particles diffuse. For more complex (fractal) topologies, anomalies occur afortiori, but the analysis is much more complex. The final goal of the fractal theory of heterogeneous distributions is to achieve a general relation expressing the exponent r appearing in eq. (18) as a function of the dimensions df. d~ and dR, and of the walk dimension dw of the structure. The main difficulty of this project (which would complete the fracta] theory of elementary catalytic reactions) is that the exponent ~ most probably does not depend solely on the dimensions df. d~, dR, and on the topology (d~), but also on the relative configuration of the reactive manifold GR, and on the exchange manifold B within the structure ~. A more detailed discussion of this issue will be developed elsewhere. 5. C O N C L U D I N G

REMARKS

Adsorption kinetics on fractal surfaces can be described with sufficient accuracy by macroscopic approximate models based on a constitutive equation (flux/concentration gradient) of Riemann-Liouville type. This article also discusses a general scaling theory for sorption on bulk fracta]s and across fractal interfaces based on the results obtained by applying Green function renorma]ization to finitely ramified structures. Numerical simulations of batch sorption kinetics on infinitely ramified structures confirm the validity of the scaling expression eq. (17). Finally, we have discussed other interesting problems related to the presence of a heterogeneous distribution of reactive centres, localized on a fractal submanifold. A complete theory of the influence of a nonuniform distribution of reactive centres on a fractal structure is still not available. In any case, Green function renorma]ization is the most convenient way to approach this issue rigorously. REFERENCES 1. 2. 3. 4. 5. 6. 7. 8. 9. 10. 11. -12. 13. 14.

D. Avnir. D. Farin, P. Pfeifer, Nature 308. 261 (1984). T. Pajkossy, L. Nyikos, Electrochimica Acta 31, 1347 (1986). T. Pajkossy. J. Electroanal. Chem. 300, 1 (1991). A.E. Larsen, D.G. Grief, T. C. Halsey, Fractals 2, 191 (1994). A. Seri-Levy, D. Avnk, D. J. Phys. Chem 97. 10380 (1993). P. Delahay, I. Trachtenberg, J. Am. Chem. Soc. 79, 2355 (1957). K . S . Miller, B. Ross, An Introduction to the Fractional Calculus and Fractional Differential Equations. J. Wiley & Sons. New York. 1993. M. Giona. M. Giustiniani. J. Phys. Chem. 100. 16690 (1996). M. Giona, W. A. Schwa]m. M. K. Schwalm. A. Adrover, Chem. Engng. Sci. 51, 4117 (1996); ibidem 51. 4731 (1996); ibidem 51, 5065 (1996). M. Giona. W. A. Schwa]m: A. Adrover, M. K. Schwa]m, Chem. Eng. J., in press (1996). W. A. Schwalm. M. K. Schwa]m, Phys. Rev. B 37, 9524 (1988). M. Giona, A. Adrover. W. A. Schwalm. M. K. Schwalm Fractals in press (1996). M. Giona: A. Adrover. A.R. Giona. Chem. Engng. Sci 50. 1001 (1995). R. Gutfraind, M. Sheintuch, D. Avnir. J. Chem. Phys. 95. 6100 (1991)

91997EIsevier Science B.V. All rights reserved. Dynamics of Surfaces and Reaction Kinetics in Heterogeneous Catalysis G.F. Froment and K.C. Waugh, editors

251

I n v e s t i g a t i o n of the structure sensitivity o f n i t r o g e n a d s o r p t i o n on single c r y s t a l r u t h e n i u m clusters u s i n g d e n s i t y f u n c t i o n a l t h e o r y D. J. Dooling and L. J. Broadbelt Center for Catalysis and Surface Science, Department of Chemical Engineering, Northwestern University, 2145 Sheridan Road, Evanston, IL 60208-3120, U.S.A. Ruthenium has long been known to be an effective catalyst for ammonia synthesis. However, compared to the traditional iron-based catalysts, studies on ruthenium-based catalysts are limited. The rate determining step of ammonia synthesis, the dissociative adsorption of dinitrogen, has been shown to be extremely structure sensitive on both iron and ruthenium catalysts. To study this structure sensitivity on ruthenium, density functional theory calculations were performed on Ru(001) and Ru(ll0) clusters. End-on, side-on, and dissociated adsorption states were investigated on both surfaces. While the Ru(110) cluster could stabilize all three adsorption modes, a minimum energy structure for the side-on adsorption on Ru(001) could not be found. It is likely that this side-on mode can provide a low energy pathway to the dissociated state, thereby resulting in faster dissociative adsorption on Ru(110). 1. I N T R O D U C T I O N The promoted iron catalyst used for the industrial production of ammonia from nitrogen and hydrogen has undergone very little change since its development by Haber, Bosch, and Mittasch in the early part of this century [1]. The industrial synthesis of ammonia over this catalyst requires pressures over 300 bar, temperatures above 673 K, and large recycle of unused reactants. These factors lead to high energy demands and thus, less favorable process economics. This has led researchers to search for more active catalysts which would allow the synthesis to be carried out at lower pressures and temperatures [2-5]. Of these, the initial investigations by Aika and coworkers [5] on alkali metal-promoted ruthenium supported on activated carbon (AC), which demonstrated that these catalysts were more active than the promoted iron catalyst, have held the most promise. Indeed, this promise was realized in 1992 when potassium promoted Ru/AC was first used for the industrial synthesis of ammonia [6]. Unfortunately, much of what is known about ammonia synthesis is from studies on iron and promoted iron catalysts [see 7]. Although some information learned from these studies can be applied to the synthesis over ruthenium, it has been shown that, compared to iron-based catalysts, many characteristics of the reactive system are markedly different for ruthenium-based catalysts [8]. For example, Fishel and coworkers [8], Muhler and coworkers [9] and Rambeau and Amariglio [10] have shown that while the reaction order with respect to hydrogen pressure over iron-based catalysts is positive, it is negative, i.e., it inhibits the reaction, over rutheniumbased catalysts. An explanation of this dissimilar behavior was given by Fishel and coworkers, who proposed that the steady state surface coverage of atomic hydrogen is much higher on ruthenium than on iron, where the fractional coverage of atomic hydrogen is less than 0.10 [11]. Similarities do exist between the reaction over iron- and ruthenium-based catalysts. For instance, it has been well established that alkali metals act as chemical promoters on both iron [1,7] and ruthenium [5,12]. This chemical promotional effect is due to the electron donating

252 ability of the alkali metals, which increases the electron density of the catalyst. This increase has two beneficial effects: 1) the catalyst can more effectively donate electrons into the antibonding orbitals of dinitrogen, thereby making its dissociation more energetically favorable and 2) the increased electron density repels ammonia, thereby driving it off the surface. The mechanism through which ammonia synthesis proceeds is also the same on both catalysts. Fishel and coworkers showed that the overall kinetics of ammonia synthesis over rutheniumbased catalysts could be described by the same microkinetic steps used in models for iron-based catalysts [11] involving dinitrogen and hydrogen adsorption and dissociation, sequential hydrogenation of nitrogen atoms and desorption of ammonia. Perhaps most importantly, several investigators have demonstrated that the rate determining step (RDS) for ammonia synthesis over iron-based catalysts, the dissociative adsorption of nitrogen [7], is also the RDS on ruthenium-based catalysts [8,9]. Therefore, as is the case with iron-based catalysts, studies of nitrogen adsorption on ruthenium are paramount to the understanding of ammonia synthesis on ruthenium-based catalysts. Studies of both nitrogen adsorption [ 13,14] and ammonia synthesis [ 15] over single crystal faces of iron have shown both processes to be extremely structure sensitive, lending further support that the dissociative adsorption of nitrogen is rate limiting. Both processes proceed much faster over the open Fe(111) face than over the more tightly packed Fe(100) and Fe(110) faces. In supported catalysts, this structure sensitivity causes the turnover frequency of ammonia to decrease as iron dispersion on the support increases. This behavior results from the decrease in the number of "open" sites needed for dissociative nitrogen adsorption as the iron particles decrease in size. These structure effects are exhibited by ruthenium as well [16], but they have not been as well characterized as those on iron. Therefore, although high dispersions are desired for most supported metal catalysts, a compromise between high dispersion and high activity must be reached for ammonia synthesis catalysts. Although much is now understood about how metal particle size and promoters affect the macroscopic rates of the dissociation step and the overall reaction, the physical and chemical causes behind such effects are not clear. This, coupled with the fact that for ruthenium particles greater than 2 nm in size (a typical particle size for supported catalysts) the particle surfaces are faceted into single crystal planes [17], suggests that a study of the interaction of nitrogen with single crystal ruthenium surfaces can contribute much to the understanding of ammonia synthesis on ruthenium-based catalysts. Unfortunately, experimental techniques cannot easily probe the exact configuration and bonding of the adsorbed species and the transition from molecular to atomic nitrogen. Furthermore, the exact interaction of the promoter with the metal and surface species is uncertain. Thus, theoretical studies which investigate the role of surface morphology and the electronic configuration of the system would complement the current experimental database and increase the understanding of the behavior of ruthenium-based ammonia synthesis catalysts. An efficient theoretical tool for the study of catalytic systems involving transition metals is density functional theory (DFT). DFT is a first principles technique which provides accurate information about energetics, geometries, and transition states that are difficult or impossible to determine experimentally. Because DFT is considerably less computationally expensive than wavefunction calculations, it can be applied to relatively large systems, thereby allowing the study of catalytic surfaces and structures that capture the necessary features of the system under investigation. In recent years, DFT has been successfully applied to a wide variety of catalytic systems including single crystal surfaces [18-20] and zeolites [21-23]. All of these characteristics suggest that DFT is a method which is well suited to study the adsorption and dissociation of nitrogen on single crystal ruthenium clusters. 2. C O M P U T A T I O N A L M E T H O D S The single crystal faces studied, Ru(001) and Ru(110), provided a contrast in surface architecture. The Ru(001) surface is very tightly packed, analogous to Fe(110), while the

253 Ru(110) surface is much rougher and more open, much like Fe(111). Cluster models of the ruthenium single crystal surfaces were used in all calculations. To minimize the effects of cluster size on the calculated adsorbate binding energies [18], clusters were chosen to be large enough to capture the necessary interactions of metal atoms and adsorbates, while remaining small enough to be computationally tractable. A ten atom Ru(001) cluster and a fourteen atom Ru(110) cluster were used. The cluster models used for each crystal face are depicted in Figure 1.

, ....

.

i....~:!:~!.~.

.'i~

9

(a)

i

~

(b)

Figure 1. Model single crystal ruthenium clusters employed: (a) ten atom Ru(001) and (b) fourteen atom Ru(110) clusters. The calculations were performed using the DFT code DMol from Molecular Simulations. Double numeric basis sets with polarization functions were used for all atoms. The 1s orbitals were frozen on nitrogen atoms, and the l s2s2p3s3p3d orbitals were frozen on ruthenium atoms. It has been shown that using only local spin density (LSD) functionals on surface clusters can affect geometries as well as energies [19]. However, non-local (gradient-corrected) functionals are much more computationally expensive than the LSD functionals. As a compromise, the Vosko-Wilk-Nusair LSD functional [24] was used to compute the exchange and correlation energies during the self-consistent field (SCF) procedure. The gradientcorrected exchange and correlation energies were then calculated from the self-consistent LSD density using the functionals of Becke [25,26] and Perdew-Wang [27], respectively. This allowed the gradient-corrected energies to be utilized during geometry optimizations and frequency calculations. Ruthenium cluster geometries were obtained from bulk crystal lattice spacings and were not allowed to optimize, and symmetry was not used for any calculations. When possible, experimentally determined adsorbate geometries were used as starting points for geometry optimizations. Calculations on transition metals are complicated by the small energy gap between occupied and virtual orbitals, resulting in slow or failed SCF convergence. To improve the behavior of the SCF calculations, small amounts (0.03 hartree) of electron density were smeared above the Fermi level in initial calculations. The density obtained from these calculations was used as an initial density for subsequent calculations at lower smear values. This approach was used iteratively until the smear was reduced to a negligible value (< 0.001 hartree) in the final calculation.

254 3. R E S U L T S

AND DISCUSSION

Experimental and theoretical studies on iron [28] have revealed that dinitrogen can exist in two distinct states on the metal surface. At low temperatures, dinitrogen can adsorb with its molecular axis perpendicular to the metal surface on an atop site (end-on). This configuration is labeled the ~, state. When appropriate sites are available, this end-on state can overcome some activation energy and tilt over to a side-on, n-bonded configuration, labeled the c~ state. Either of these molecular species can then overcome an activation energy and dissociate to form two adsorbed nitrogen atoms, referred to as the 13 state. Tomfinek and Bennemann [29] concluded from theoretical studies on the open Fe(111) face that the most energetically favorable dissociation pathway involves initial adsorption into the "f state, followed by tilting over into the state, and finally dissociation into the 13state. They also concluded that other pathways, such as ), directly to 13, are possible but are energetically less favorable. In the present study, the stabilities of the 7, a and 13 states for Ru(001) and Ru(110) were investigated using the procedure outlined in the previous section. It was found that while all three states could exist on the open Ru(110) surface, the Ru(001) surface could not stabilize the a state. Thus, it appears the "appropriate sites" necessary for the side-on adsorption of dinitrogen are not present on the closely packed Ru(001).

3.1. Adsorbate geometries The minimum energy adsorbate structures for nitrogen adsorption and dissociation on Ru(001) and Ru(110) are shown in Figures 2 and 3, respectively. In all representations of adsorbate-surface interactions, the darker colored atoms represent nitrogen, while the lighter colored atoms depict ruthenium. As seen in Figure 2a, the bond distances for both Ru-N and N-N for end-on adsorption on Ru(001) are in good agreement with the experimental values shown in parentheses. Comparison of the geometries for end-on adsorption on the two surfaces (Figures 2a and 3a) shows the similarity of the ~, state on the Ru(001) and Ru(110) surfaces. For both surfaces, the N-N bond distance is only slightly longer than the calculated gas phase value of 1.11 A, which in turn is in good ageement with the experimental value of 1.0975 ,~. One difference between the two geometries is the angle the dinitrogen axis makes with the surface. While dinitrogen is perpendicular to the Ru(001) surface, it makes a 10 ~ angle

~ 1.14 (1.10) 1.99 (2.00)

(a)

3.68 1.1

5

2.21

(b)

Figure 2. Optimized structures of (a) dinitrogen adsorbed end-on and (b) dissociated nitrogen atoms on Ru(001). All values are either N-N, N-Ru or N-surface distances and are in angstroms (A). Experimental values [30] are in parentheses.

255

g 9

l.14

fI: 1.97

....::.k ~

(a)

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~

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~.~<~;.'...~

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~ .G~,"

"~

-

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~

!

,

.:.e

!

l

~

~<*~"~

. .<.<.~z~'~.

(b)

.

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.

~

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. ..

::~i71.

<<.f!7

:#.

~..,...

'li,~'~ ~+>

.<.,. : .. ~i:;~.

. ~ .......

..,.....~~>

(c)

Figure 3. Optimized structures of (a) dinitrogen adsorbed end-on, (b) dinitrogen adsorbed sideon and (c) dissociated nitrogen atoms on RU(lol 0). All values are either N-N or N-Ru distances and are in angstroms (A). The value of 1.96 A in (c) applies to both N-Ru bonds. with the normal to the Ru(110) surface. Although the current set of calculations cannot determine if the inclination is a result of cluster effects or an actual energetic preference of dinitrogen for the nearby hollow, a study of the geometry convergence showed that the potential energy surface for dinitrogen adsorbed end-on Ru(110) was nearly flat. Therefore, there is no strong energetic preference for a specific orientation about the surface normal, allowing the molecule to precess around the surface normal without much increase in energy. This ease of movement about the normal provides a low-energy pathway to the side-on o~ state, which will be discussed further in the following section. The most striking difference between the two surfaces was the inability of the Ru(001) surface to stabilize a side-on o~ state. All attempts either resulted in dissociation or desorption during geometry optimization. In contrast, the open structure of the Ru(110) surface easily accommodated a side-on adsorption state as shown in Figure 3b; the end-on adsorbate is only required to "fall over" and interact with the other ruthenium atom. The inability to stabilize a side-on mode on the Ru(001) surface indicates that this simple "falling over" mechanism may not be available to end-on adsorbates on its surface. For the side-on state on the Ru(110) surface, each nitrogen atom directly interacts with a surface ruthenium atom, and both the N-N

256 and N-Ru bond distances have elongated slightly compared to the end-on state. The elongation of the N-N bond by only 0.03 A indicates that the side-on configuration is a molecularly adsorbed state, in sharp contrast to the dissociated adsorption states depicted in Figures 2b and 3c. The distance between the nitrogen atoms in the dissociated state has elongated, and the distance from the nitrogen atoms to the ruthenium surface (1.15 A and 0.52 A on Ru(001) and Ru(110), respectively) has shortened. It is also clear from Figures 2b and 3c that in the dissociated state, the nitrogen atoms interact directly with at least three ruthenium atoms, whereas each nitrogen atom interacts with only one or zero ruthenium atoms in the molecularly adsorbed states. It is easy to visualize that a simple twisting mechanism can transform the sideon adsorbed state to the dissociated state on Ru(110), but no analogous mechanism seems likely on the Ru(001) surface. The simplest mechanism for dissociation on Ru(001) would most likely be a higher energy process involving the lengthening of the Ru-N bond and subsequent flipping over of the nitrogen atoms into the three-fold sites.

3.2. Adsorption energetics The energetics of each mechanistic step in the adsorption and dissociation of nitrogen on the two crystal faces are summarized in Table 1. The energies for end-on adsorption compare well with typical experimental values for nitrogen end-on adsorption on iron [13,31 ]. The relatively large difference between the adsorption on Ru(001) and Ru(110) would not be expected from a simple comparison of the nearly identical adsorbate geometries for end-on adsorption on the two surfaces. However, closer inspection of the surface morphologies of the two clusters provides an explanation for the stronger adsorption on Ru(110). Since the surface atoms on Ru(110) are more isolated than those on the closely packed Ru(001) surface, the Ru(110) surface atoms cannot interact with each other as effectively as on the Ru(001) surface. This allows the surface Ru(110) ruthenium atoms to more effectively donate electrons to the antibonding dinitrogen orbitals, which results in a stronger N-Ru bond. A simple Mulliken population analysis of the two systems qualitatively supports this picture, revealing that there is a greater charge transfer to the dinitrogen on Ru(110) than on Ru(001). It is clear from the results reported in Table 1 that the side-on adsorption on Ru(110) is actually higher in energy than the end-on configuration. Tom~nek and Bennemann [29] obtained similar results from a theoretical study of nitrogen adsorption on Fe(111). This weaker adsorption is consistent with the simultaneous lengthening of the N-N and N-Ru bonds as the end-on configuration falls over to the side-on state. Therefore, as discussed in the previous section, the main role of the side-on state is not as a stable surface intermediate at reaction temperatures, since even the end-on is more stable, but rather as a pathway from the end-on to the dissociated state. Table 1 Energetics of adsorption and dissociation of nitrogen on single crystal ruthenium surfaces Surface

Reaction

AE (kcal/mol)

Ru(O01)

N2 (g) ~ N2 (7)

-7.5

Ru(110)

N2 (7) ~ N2 (13)

-35.6

N2 (g) ~ N2 (7)

-12.2

N2 (7) --~ N2 (o0

+5.3

N 2 (o0 ~ N 2 (13)

-29.7

Species labels: (g) gas; (7) end-on state; (c~) side-on state; (13) dissociated state.

257 Dissociation was observed to be the most exothermic step on both surfaces. This is expected since the dissociated state allows the nitrogen atoms to interact with at least three ruthenium atoms (Figures 2b and 3c). The overall energy change resulting from the transformation of gaseous nitrogen to two adsorbed nitrogen atoms for Ru(001) and Ru (110) is -43.1 and-36.6 kcal/mol, respectively. These values clearly show that the nitrogen atoms are more thermodynamically stable on Ru(001) than on Ru(110), indicating that the Ru(001) surface would have a higher equilibrium surface coverage of atomic nitrogen than Ru(110). However, since the dissociation of nitrogen is the rate limiting step in ammonia synthesis, this equilibrium state on the catalyst is not achieved. Although energy barriers and the associated kinetics for each step can only be obtained using a full transition state search, it seems likely that the lower energy pathway leading to the dissociated state on Ru(110) would more than compensate for any advantage the higher adsorption energy on Ru(001) offers.

3.3. Promoter effects on Ru(001) The energies in Table 1 can be used to quantify the promotional effect of alkali metals on ammonia synthesis. As discussed earlier, alkali metals are excellent chemical promoters for ammonia synthesis on both iron- and ruthenium-based catalysts. This promotional effect is a result of the ability of the alkali metals to donate electrons to the metal. Single crystal studies of nitrogen adsorption on iron have shown that the promoter effect is more pronounced on the more tightly packed iron faces compared to the open faces, i.e., the less active crystal faces showed a higher activity enhancement in the presence of alkali metals than do the more active faces [31]. This suggests a larger promoter effect would be observed on Ru(001) than on Ru(110). To study the role of the alkali promoter as an electron donor, the overall charge of the Ru(001) cluster and adsorbate was set to -1.00 a.u., increasing the total number of electrons in the system by one. This approach has been successfully applied by Rinc6n and coworkers [32] to study the bonding of dinitrogen to single iron atoms. Geometry optimizations were performed using the optimized end-on and dissociated adsorbate geometries from the uncharged calculations as starting points. The geometry of the end-on species on the charged cluster was identical to that on the uncharged cluster. Optimization of the dissociated state resulted in further separation of the nitrogen atoms as compared to the uncharged cluster, indicating the extra electron increased the adsorbate-adsorbate repulsion. The adsorption energies obtained for the negatively charged cluster are tabulated in Table 2. The adsorption energy for the end-on configuration is slightly lower on the negatively charged cluster than on the uncharged cluster (Table 1). The dissociative adsorption of nitrogen on the charged cluster, however, is much stronger, leading to an overall adsorption energy which is over 10 kcal/mol greater than for the uncharged cluster. Linking these energetic changes to a rate enhancement requires a full transition state search to be performed on both the charged and uncharged clusters and will be the focus of future efforts. Table 2 Energetics of adsorption and dissociation of nitrogen on a negatively charged Ru(001) cluster Reaction

zkE (kcal/mol)

N2 (g) --+ N2 (7)

-6.7

N 2 (~') ~ N 2 (13)

-47.1

Species labels: (g) gas; (~,) end-on state; ([3) dissociated state.

258 These theoretical results are consistent with the experimental contrast provided by studies of iron- and ruthenium-based ammonia synthesis catalysts. Ertl and coworkers [31] concluded that on iron single crystal surfaces, one promotional effect of potassium was to increase the adsorption energy of dinitrogen, thereby increasing the surface coverage of dinitrogen and decreasing the activation energy to dissociation. The ultimate result of this promotion then is a higher rate of dissociative adsorption. Our calculations, however, indicate that the end-on adsorption on the charged cluster is actually weaker than on the uncharged cluster for Ru(001) and suggest a different promotional effect is active. One possibility is the promotional effect of alkali metals through the repulsion of ammonia, thereby driving it off the surface. This type of promotion has been demonstrated experimentally by Muhler and coworkers [9]. They determined that the reaction order with respect to ammonia pressure is negative on unpromoted ruthenium, but is zero on alkali promoted ruthenium. They also showed that the reaction order with respect to hydrogen pressure became more negative in the presence of alkali promoters, indicating that the surface coverage of atomic nitrogen on the alkali promoted catalyst is lowered. Further theoretical investigations studying other adsorbates and the effect of the promoter on their behavior could help to establish the preference for this type of promotion. 4. C O N C L U S I O N S While the adsorbate geometries for end-on adsorption on Ru(001) and Ru(110) are quite similar, the side-on configuration could only be stabilized on Ru(110), suggesting the subsequent dissociation into adsorbed atomic nitrogen is likely facilitated on Ru(110). Both surfaces formed very strong bonds with atomic nitrogen, in agreement with experimental results that this is the most stable nitrogen surface species. Calculations to investigate the role of alkali promoters on Ru(001) showed that while the geometries changed little or not at all compared to the unpromoted catalyst, the energetics were affected. Although the overall adsorption energy of gas phase nitrogen to dissociated atomic nitrogen was more favorable, the energy change upon formation of the end-on adsorbed species was decreased. These calculations are consistent with experimental evidence that indicates that the main effect of alkali promoters on ruthenium-based catalysts is not to increase the surface coverage of nitrogen, but rather to drive ammonia off the surface. Although further investigations of transition state structures and activation energies are necessary to fully describe the adsorption of nitrogen on ruthenium single crystal clusters, the energetics and geometries of the stable surface intermediates provide insight into the differences in the adsorption behavior of dinitrogen on contrasting crystal faces of ruthenium.

REFERENCES 1. M. Grunze, in The Chemical Physics of Solid Surfaces and Heterogeneous Catalysis D. A. King, D. P. Woodruff, Eds. (Elsevier, Amsterdam, 1982), vol. 4, pp. 143. 2. G. Rambeau, A. Jorti and H. Amariglio, J. Catal., 74 (1982) 110. 3. S. Kiperman and M. Temkin, Acta Physicochimica U.R.S.S., 21 (1946) 267. 4. M.R. Hillis, C. Kemball and M.W. Roberts, Trans. Faraday Soc., 62 (1966) 3570. 5. K. Aika, H. Hori and A. Ozaki, J. CataI., 27 (1972) 424. 6. P.G. Menon, AppI. CataI. A, 93 (1993) N 16. 7. G. Ertl, CataI. Rev.-Sci. Eng., 21 (1980) 201. 8. C.T. Fishel, R.J. Davis and J.M. Garces, J. CataI., 163 (1996) 148. 9. M. Muhler, R. Rosowski, O. Hinrichsen, A. Hornung and G. Ertl, Studies in Surface Science and Catalysis, 101 (1996) 317. 10. G. Rambeau and H. Amariglio, J. Catal., 72 (1981) 1. 11. J.A. Dumesic and A.A. Trevifio, J. Catal., 116 (1989) 119. 12. S. Murata and K. Aika, J. Catal., 136 (1992) 110.

259 13. F. Bozso, G. Ertl, M. Grunze and M. Weiss, J. Catal., 49 (1977) 18. 14. F. Bozso, G. Ertl and M. Weiss, J. Catal., 50 (1977) 519. 15.N.D. Spencer, R.C. Schoonmaker and G.A. Somorjai, J. Catal., 74 (1982) 129. 16. J.G. Buglass, S.R. Tennison and G.M. Parkinson, Catal. Today, 7 (1990) 209. 17.A.K. Datye, A.D. Logan and N.J. Long, J. Catal., 109 (1988) 76. 18. M. Neurock, W.D. Provine, D.A. Dixon, G.W. Coulston, J.J. Lerou and R.A. van Santen, Chem. Eng. Sci., 51 (1996) 1691. 19. H. Cheng, D.B. Reiser, P.M. Mathias, K. Baumert and S.W. Dean, Jr., J. Phys. Chem., 100 (1996) 9800. 20. J.-F. Paul and P. Sautet, Phys. Rev. B, 53 (1996) 80158027. 21.J.B. Nicholas, Studies in Surface Science and Catalysis, 101 (1996) 1263. 22.B.L. Trout, A.K. Chakraborty and A.T. Bell, J. Phys. Chem., 100 (1996) 4173. 23.K.J. Farnworth and P.J. O'Malley, J. Phys. Chem., 100 (1996) 1814. 24. S.H. Vosko, L. Wilk and M. Nusair, Can. J. Phys., 58 (1980) 1200. 25.A.D. Becke, Phys. Rev. A, 38 (1988) 3098. 26. A.D. Becke, J. Chem. Phys., 88 (1988) 2547. 27.J.P. Perdew and Y. Wang, Phys. Rev. B, 45 (1992) 13244. 28. H.-J. Freund, B. Bartos, R.P. Messmer, M. Grunze, H. Kuhlenbeck and M. Neumann, Surf Sci., 185 (1987) 187. 29. D. Tom~.nek and K.H. Bennemann, Phys. Rev. B, 31 (1985) 2488. 30. H. Bludau, M. Gierer, H. Over and G. Ertl, Chem. Phys. Len., 219 (1994) 452. 31.G. Ertl, S.B. Lee and M. Weiss, Surf. Sci., 114 (1982) 527. 32. L. Rinc6n and F. Ruette, Hern~ndez, J. Molecular Structure (Theochem), 254 (1992) 395.

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91997 Elsevier Science B.V. All rights reserved. Dynamics of Surfaces and Reaction Kinetics in Heterogeneous Catalysis G.F. Froment and K.C. Waugh, editors

263

O x y d e h y d r o g e n a t i o n o f p r o p a n e on N i M o O 4 catalyst under transient and s t e a d y state conditions S. Pietrzyk, M.L. Ould Mohamed Mahmoud, T. Rembeczky, R. Bechara, M. Czernicki, N. Fatah Laboratoire de Genie Chimique et d'Automatique (LGCA), Ecole Nationale Superieure de Chimie de Lille - Ecole Centrale de Lille, B.P. 48, 59651 VILLENEUVE D'ASCQ CEDEX, FRANCE This work concerns the propane oxydehydrogenation (ODH) reaction 9 C3H ~ + 1/2 O : - C3H 6 + H.O

(1)

and in particular its possible realization in a circulating bed reactor (CBR [1]). The reason for interest in such a subject is the following. The upgrading of cheap alkanes by selective oxidation is an important industrial challenge. The main difficulty is avoiding the consequences of the decrease in the selectivity due to total oxidation reaction, as : CgH s + 3.5 O: - 3 CO + 4 H:O

(2)

5 0 2 - 3 CO z + 4 H:O

(3)

C3H s +

The formation of carbon oxides is usually attributed to <<non-selective oxygen >> i.e. 02, O2(aas), O;, O [2] present when the reaction mixture contains both the hydrocarbon to be oxidized and oxygen (<< COFEED mode ~>). Therefore the application of a CBR reactor fed with propane in the reaction zone (riser) and with oxygen in the regeneration zone (fluidized bed), should improve the selectivity to propene (<>). Other advantages of CBR are discussed elsewhere [3]. When an oxidation process is realized in a CBR the catalyst is alternatively reduced and oxidized in fact its participation in the reaction in both principal parts of the reactor is that of a stoichiometric reactant rather than of a catalyst. Accordingly, the formation of products should be considered as a transient process in which the surface oxygen species and in particular lattice oxygen are alternatively depleted and replenished. Therefore the catalytic solid should be able to support such a cyclic operation without irreversible modifications. In this work, a catalyst is tested from this point of view, both under steady state conditions and in transient conditions simulating those in a CBR. 1. E X P E R I M E N T A L This paper reports the results obtained with a NiMoO4 catalyst, used in propane ODH, in stationary and transient mode operated reactors.

264 1.1 Catalyst The nickel molybdate catalyst used in this study, prepared by the precipitation method described by Mazzocchia [4], has been supplied by Elf Atochem (France). The bulk M o / N i ratio in this solid is practically equal to 1 it can exist as phases c~ and 13, differing by the coordination of Ni ions. Only the results for the c~ form will be reported here. The specific surface area as measured by N2 adsorption is about 40 m 2 / o 1.2 Steady- state reaction tests These experiments have been performed in a classical fixed bed continuos flow reactor system [5]. The catalyst sample (0.2 g) diluted with SiC was charged in a 12.6 mm i.d. tubular stainless-steel reactor, between two layers of SiC. The thermocouple well was placed in the axis of the reactor. The various C3H8 - 02 - N2 mixtures were prepared by means of mass flow controllers and the concentrations of products were determined by gas chromatography, using PORAPAK Q columns (2 m * 1/8") for preliminary separation and CARBOXEN 1000 column (4m * 1/8") for 02, N2, CO, CO2 and an other PORAPAK Q column (2 m * 1/8") for C3Hs and

C3H6. 1.3 Transient reaction tests Transient tests have been performed with a fixed bed of 0.125 g catalyst in a 1 mm id. quartz reactor exposed to a series of concentration steps 9C3Hs-He, He, O2+He, He, C3Hs-He, etc. This was obtained by distributing the three constant gas flows between the reactor and two parallel ~ ballasts ~> with two couples of three way eletrovalves [6]. The concentration of products was determined using a Balzers QMS 420 system based on quadripole mass spectroscopy. The reactor conception was taken from the works of Hoffmann and Mi~ller [7]. In a given experiment, only propane steps or only oxygen steps were analyzed, to avoid interactions between oxygen and hydrocarbon in the QMS analyzer. This was achieved by means of a third couple of three-way electrovalves placed on the downstream side of the reactor. 2. RESULTS AND DISCUSSION 2.1 Activity and selectivity under steady-state conditions The activities and yields of principal products (C3H6, CO, CO2) have been determined between 530 and 560~ in function of the reaction mixture composition, and in function of the contact time for a mixture containing 15% C3H8, 18%02 and 67% N2. The conversions were always under 30%, so the differential reactor approximation is assumed. In figure 1 the yields of C3H6, CO, CO2 are shown as functions of conversion. As shown by Best and Wojciechowski [8], such a representation may be used to distinguish between the primary and secondary reaction products. In this case, it can be concluded that the propene is the only primary product, and carbon oxides are formed as secondary products. This result will be useful in the choice of the reaction mechanism.

265 9%8%

7% 6% (D

5%

~

*-

>= 4% 3% 2%

/~~__~--~.

-CO a C02

1%

0% ~

'

0%

5%

10%

i

15%

20%

25%

Conversion Figure 1. Yields of products vs. conversion. The influence of partial pressures of propane and oxygen was studied at 550~ The results have been correlated using the well known Mars and van Krevelen [9] mechanism 9 n

k 1 k: Pc_~H~ Pc), v

-

(4)

k1 PC.~H~ + k: p~,, with n = 1 or n=0.5 9the obtained values of constants are as follows Table 1 Constants of the Mars and van Krevelen mechanism (n- order of the reoxidation stage, s2 variance in (mol / (min g))2, kl in mol (g bar s) I, k2 in mol (g bar s) -1 for n-1 and in mol (g bar ~ s-1) for n - 0,5) Phase

k~ * 103

k2 * 103

ot

3.2

26.4

1

0.16

ot

3.1

11.0

0.5

0.17

n

s2

2.2 Partial pressures of products in transient experiments As already stated, in each transient experiment including a series of propane steps (15 % C3H8 in He) and oxygen steps (20 % 02 in He) separated by inert (He) steps only the hydrocarbon steps or only the oxygen ones have been analyzed, to avoid interactions between the ~ ~ active gases in the analytical system, giving rise to artifact peaks and to a large variation of response coefficients The typical results obtained for a propane step are shown in figure 2 A typical experiment consisted of 5 to 8 hydrocarbon steps of equal length separated by the adequate number of oxygen steps The evolution of partial pressures in all propane steps obtained at given

266 conditions was practically identical, provided that the propane steps were not too long. In such conditions the properties of catalyst sample changed only very slowly and one sample could be used in many series of experiments.

0,008 -

.'f"", ~ '

0,007 0,006

i

\ ~.-r ....... ""

~

.... \

[

0,005 ~ 0 , 0 0 4

...... ..

"-" . .

9

!

.~.

......

i ] ,

.~ 0 , 0 0 3

................Propene

"-. ,......

""-.... ............. "'""

". . . . . .

"'-~.\

!

e- 0,002 -:

........ "

""-'-x.

COx

........... Water "" --- ,..

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

.... "-----~-,........... f ".', ,i

o,ool,

i,

"

......

". . . . .

..'.-.:"''

--.

9

-

I. . . .

0

i.~.::J"--

0

~

~

~

20

40

60

'

80

100

120

Time (s) Figure 2. Partial pressure of products vs. time in a typical transient experiment. When the reacted mixture corresponding to oxygen steps was analyzed, only small peaks of CO and CO: were observed, the corresponding amounts being of the order of 1 % of reacted propane. The formation of H:O was not observed. It can be concluded that only small amounts of carbon are deposited at the catalyst surface when its reduction is not excessive. When the amount of oxygen withdrawn from the NiMoO4 catalyst approaches 2 % of its total oxygen content, the catalyst undergoes a sudden modification [6]. Its initial properties are not restored anymore by the regeneration step and its catalytic activity involves cracking rather than oxydehydrogenation. An analysis by XRD revealed the presence of MoO:, Ni and NiO in the reduced solid.

2.3 Modeling of transient and steady state results 2.3.1 Reaction mechanism To reproduce the partial pressures of products in transient experiments a rather detailed mechanism must be assumed. We tried several different possibilities [5, 6] .The best agreement with transient experiments was achieved with the following one CsHs+2Os CsHM

k,

s + O s

2OHads

>CsH7,~s + O H a d s

(5)

k:

(6)

) C3H6

+ OHads + O s

k3 > H 2 0 + O s + V

20~ k, > 3 CO4/3 "+- 6 0 H a ds - 8 0 s + 4 V (Os = surface 02. oxygen ion, V - vacancy)

C3H 6 +

(7) (8)

267

The first step is similar to that proposed for ODH of isobutyric acid [ 10], and the equations 6 and 8 express the primary character of C3H6 and secondary of carbon oxides. The ratio 4/3 in the last equation corresponds to the average ratio of CO/CO2 in the products approximately equal to 2. The above equations are sufficient to modelize the formation of products in the propane step, when no catalyst re-oxidation occurs. To apply this mechanism to the steady state reaction, the catalyst re-oxidation stage was written formally as : 0.50: +V

k, > Os

(9)

2.3.2 Application to transient experiments

The calculation of the rate constants k~ was performed using the MATLAB software (Scientific Software). Since the total conversion in considered transient experiments was always lower than 5% the spatial variation of concentrations was neglected [ 11 ]. The equations (5) to (8) give the following system of differential equations 00 C

d t - kl

3oo dt

-

•0OH

~=k3 c~t

PC~Hs00"- --k,_ 0 C 0 o 2 k~ P c.,H~ 0o 2 + k 3 0 o u - - 1 0

(10)

k 4 P<,H~ 002

- k~ Pc.,u~ 0 ~ . k 2 0 c 0 o - 2 k 30ou: + 6 k 4 Pc3H6 0o 2

0oH2 + 4 k 4 Pc~H6 002

(11)

(12)

(13)

where Oo, Oc, 0oH, 6v = surface coverages with oxygen, C3H7, surface hydroxyl groups and vacancies respectively The rate constants have been assumed independent of the coverages; also any participation of lattice oxygen from layers other than the first one was neglected. The corresponding variations of partial pressures were expressed taking into account the specific surface area of the catalyst (A-40 m2/g), 02. maximal surface concentration (N ~ taken as 1019 m-2), mass of the catalyst sample (m), molar gas flow rate (D) and the total pressure

(P): C k~ 0 c 0 o PC:H0-- l + C k 4 0o 2

(14)

2

PH,O -- C k 3 0OH

(15)

Pcox = 3 C k~ 0o 2 PC,H6

(16)

Am C

N ~

NAD

P

NA " Avogadro' s constant

268 The initial coverage by lattice oxygen, 0~ was taken as 0.6, on the basis of an analysis of

C3H6 and COx partial pressures. For a set of the rate constant k~ - k4 the system of ODE (10)-(13) was integrated by the Runge- Kutta method (2/3 order ; MATLAB command <>). The resulting coverages t3~ were used to compute the partial pressures of propane, COx and H20. The object function fm ( k ) - ~

[

w(p,(propene .......

13 propene'~ 2 +. (1 __ VV) (p CO •

d) - - 1" i(comp) ]

i( . . . . . . .

i(comp)) ! a)_ pCO x z

w=0.7

was minimized (<>of MATLAB) to give the optimal values of rate constants. The H~O partial pressure was not used in the computation because its values measured by our analytical system were distorted by adsorption / desorption phenomena. The results obtained for the cz form, between 530 and 562~ are given in table 2. Table 2 Values of the rate constants obtained from the transient experiments T~

k~ (bar s) ~

k2 (l/s)

k3 (l/s)

k4 (bar s) ~

528

0.0593

0.5281

11.237

0.229

542

0.0854

0.6231

16.416

0.337

555

0.1146

0.8680

30.949

0.425

562

0.1244

0.5153

19.106

0.487

The agreement between the calculated curves and the experiment is generally good. (Figure 3) 0,009 r~.....2-

0007

X

0,006 -j = r.~

.-2:"(."..... .,

i

2.:

.i

0,002 ~ _~

"'-X.

...~.:

0 !'*';~" 0

Propene

~,.,: -,,-, ~, ... ,:::-, ,.,,,,.~,, .,, ~, ,,, "~".:N,... '"~g~_,,~, ,,,."~ ",:- .~, ,., .-,, ,. ~g---x..~. . . . . .

b,:'.~ .o, , ., 9 , . O",,-b" ' 20

..'-

COx

.............. Water

..... ..... 2.. -.'-,,, .... -., ,,~g ..... ":-.::I.L

:!

.~ 0,003 "" .~ 0,003 o,ool

"',<x x,~

..;~.

:"

,,.

.....\ ~~, . X

:

g

......

2.

\ :X . . . . . . . . .

:~

0005

~I:~ 0 , 004

................Propene

,2, .:ii:,

0,008

.

.',

""

2.

o

COx

::,9

Water

....... ~

.

~ 'o'. 8"6- 1; o~'d..o ' e , " 6 . o d'.'& 9 ~

i

40

60

80

100

120

Time (s) Figure 3. Experimental (lines) and computed (points) partial pressures of products (T = 550~

269

2.3.3 Application to steady- state results It is interesting to compare the behavior of a catalyst working in the COFEED and in the REDOX modes. Are the advantages for the selectivity of the absence of gas (adsorbed) oxygen real 9 In the case of the propane ODH on the NiMoO4 catalyst, the formation of carbon oxides is still observed when the (oxidized) catalyst is exposed to the propane concentration step, without any oxygen in the gas phase. It should be concluded that the lattice oxygen is also active in the propene oxidation leading to COx and H20. We tried to apply the mechanism chosen to represent the results obtained in the transient mode, with the constants determined as described above, to our steady-state results at 550~ To do this, the differential equations (10 - 13) have been modified by including the terms corresponding to the catalyst reoxidation, equation (9), and transformed into a system of algebraic equations by taking the coverages as constant ( d 8:/d t = 0 ). k1

(17)

OC --k~ Pc.,, S Oo

%-o0 0v = ~ 1 ~ _ k5 v,/Po. --

(18)

p,Ho

(kl Pcd-Is + 7 k 4 PC,Ha 80 -

(19)

"

~176 (20) + 8o ~-7-pc,~, +

+3k a

)+1

-1-0

The partial pressures a r e k: Oc Pc:~ - k~ 0 o Pco~ - 3 C 0 o k~ Pc.~ and the reaction rate and the selectivity to propene are as follows 9 (C ' = N OA/NA, cf. 2.3.1)

r - C' k 1 p c:H80o S =l/(l+Ck4

0~

The <>k5 value is 4.0 (s a bar~ It can be seen that the agreement is rather good for the reaction rates (figure 4) while the selectivities to propene are overestimated (figure 5).

270 1 oo o~o -

10-7 9 "0 * ~a

6

,-. a. E

4 3

8

.,-, >

7 5

0

0

5

9

"5 e-,, E 9 r..)

4~176176 3~176 20% - ...' 109/0 - , , ' 0% " 0%

Mesured rate (* 104)

/"' /

9

70% 60~ 50% -

10

Figure 4. Parity diagram of reaction rates in mol / min/g.

IIOr

80O,o _

=9 2 -~

10 / 0

90% -

/

/

/ / / /f // /

50%

100%

Mesured selectivity

Figure 5. selectivity.

Parity

diagram

of

propene

3. C O N C L U S I O N The model of the ODH of propane on NiMoO~ catalyst, based on a simplified mechanism taking into account the formation of propene and its further oxidation to carbon oxides may be used both with transient and steady state results. In the first case, the evolution of partial pressures of propene and carbon oxides (lumped into COx) is reproduced. When applied to steady state results, the model gives a good representation of the reaction rate, but overestimates the propene selectivity. It should be stressed, that the mechanism proposed takes into account the participation of the lattice oxygen only, both in the formation of propene and in its destruction. Under steady state conditions, where the oxygen gas is present, its participation seems restricted to the propane destruction (but is not significant in the propane activation). In order to develop the equations to model the reactors for catalytic selective oxidation working in the REDOX mode (ex. CBR reactors), the application of transient reactors similar to that used in this study appears to us to be necessary. Acknowledgment

The catalyst samples were kindly provided by Service de Catalyse of ELF ATOCHEM (France). REFERENCES

1. R. M. Contractor, U.S. Patent No. 4,668,802 (1987) 2. J, Haber, Z. Phys. Chem.~ NF, 144 (1985) 69 3. R. M. Contractor, A. W. Sleight, Catal. Today, 1 (1987) 587

27 I 4. C. hlazzocchia, C. Aboumrad, C. Diagne, L e t t , 10, (19?!) 151

E. Tempesti, J. hl.. Hermann, G. Thomas, Catal.

5 . T. Rernbeczky, PhD thesis, University of Compiegne, France (1996) 6. hi L Ould hlohamed hiahmoud, PhD thesis, University of Compiegne, France ( 1996) 7 . E. hluller, PhD thesis, University of Erlangen, Germany ( 1 956). and E hluller, H Hoffmann, Chem. Eng. T e c h , 5s (19S6) 956 S. D. Best, B. Li’. LVojciechowski, J. Catal., 47 (1977) 1 I 9. P. Mars, D. b’.\‘an Krevelen, Chem. Eng. Sci., Special Suppl. 5 ( I 954) 4 1 10. Y. Ono, in J. h l Thomas, K. I. Zamarev (eds), Perspectives in catalysis, Blackiuell, Scientific Publications, London ( 1 997) 1 1 . E’.J. Huang. P. I Lee, J A Schivarz, J. C. Hebirelier, Chem. Eng Cornm , 39 (19Sj) 355

This Page Intentionally Left Blank

91997 Elsevier Science B.V. All rights reserved. Dynamics of Surfaces and Reaction Kinetics in Heterogeneous Catalysis G.F. Froment and K.C. Waugh, editors

273

CATALYTIC I G N I T I O N DURING ~ I E T H A N E O X I D A T I O N ON PLATINU.~I" E X P E R I M E N T S AND _'kIODELING G. \'eser ~'s'*. .J. Frauhammer ~. L.D. Schmidt b. and G . Eigenberger ~ Institut fiir Chemische Verfahrenstechnik, Universit/it Stuttgart. 70199 Stuttgart. Germany b Department of Chemical Engineering and .Materials Science. University of Minnesota. Minneapolis, MN :55455, USA

Abstract The ignition behaviour of methane-air mixtures on a platinum-foil catalyst was studied at atmospheric pressure over the entire range of fuel to air ratios. The measured surface ignition temperatures showed a continously decreasing trend with increasing fuel:air ratio up to very fuel rich mixtures, which could be reproduced using a very simple analytical model indicating strong site competition between methane and oxygen on the catalvst surface. In parallel, steady state experiments during methane oxidation in a Pt coated monolith were modeled with a detailed monolith reactor model, comprising heat-, mass- and impulse-balance and a detailed elementary step surface reaction mechanism. A close fit between experimental and modelling results was obtained. The two approaches were then combined by comparing the results of dynamic simulations of the ignition behaviour of the detailed monolith model with the experimental results on the Pt foil catalyst.

1. INTRODUCTION Due to the known limitations of the world oil reserves, methane oxidation under fuel rich conditions will become increasingly important for the production of synthesis gas. which through methanol synthesis and Fischer-Tropsch reactions is the basis of many important petrochemical synthesis routes. Therefore. catalytic oxidation of methane has again become the focus of much basic and applied catalysis research in recent years. In this context, Schmidt and coworkers were able to show recently, that catalytic direct oxidation of methane over noble metal coated monoliths can yield CO and H2 with very high conversions and selectivities at the desired 1:2 CO:H2 ratio (Hickman and Schmidt. 199:2 and 1993a: Torniainen and Schmidt. 1994: Bharadwaj and Schmidt. 1995). Furthermore. total oxidation of lean and near-stoichiometric m e t h a n e / a i r mixtures has received great interest in the context of catalytically aided combustion of natural gas for power generation (Trimm. 1983: Pfefferle and Pfefferle. 1987). Here. the catalyst helps avoiding undesired side reactions (such as NON formation) by lowering the necessary reaction temperature and stabilising the reaction against fluctuations in the feed composition. "correspon&ng author. G. I gratefully acknowledges financial support through fdlowsh@s by the A.r.Humboldt-Foundation and the Deutsche Forschungsgememschaft (DFG).

274 For all of these processes, a det.ailed knowledge of ignition and extinction limits and an understanding of the underlying reaction mechanisms are of prime importance due to both economic and safety concerns. The present paper focusses on the catalytic ignition and steady-state behaviour during methane oxidation over Pt catalysts. Since catalytic ignition in a monolith reactor is experimentally difficult to investigate, we choose instead to do experiments on a Pt foil catalyst in a very simple stagnation point set-up. Then. we tested a detailed one dimensional model of the monolith reactor, which comprises complete heat-, mass- and impulse balance along with an elementary' surface reaction mechanism, against steady state experiments. With the thus tested reaction mechanism, catalytic ignition in the monolith reactor was calculated and compared to the experiments on the foil catalyst. 2. STAGNATION POINT F.XPERIMENTS 2.1. Experimental Set-up The experimental setup was designed for simplicity and reproducibility'. It is described in detail elsewhere (Veser and Schmidt. 1996) and will therefore be discussed very briefly here. All experiments were conducted at atmospheric pressure in a quartz-glass flow tube reactor (2.5 cm diameter. 20 cm length). The reaction gases were premixed and flowed perpendicular to the catalytic foil in a stagnation point flow configuration (inset fig. 1). All experiments were conducted at total gas flow rates between 1 slpm and 6 slpm, which did not influence the results within experimental error. The high-purity platinum foils were resistively heated and the foil temperature was determined by a chromel/alumel thermocouple spotwelded to the back of the foil. Temperature measurements were reproducible within +10 t~Z on the same foil and within +:30 I( in independent runs with different foils. During an experiment, the gas flow was adjusted to a fixed methane/air ratio and then the power input to the foil was increased until a discontinuous jump in the foil temperature indicated ignition (light-off) of the catalytic surface reaction (fig. 1. point (a)). Upon decreasing the power input, the reaction shows a hvsteretic behaviour, i.e. due to the heat released by the catalvtic surface reaction, extinction occurs at a lower power input than ignition (fig.1. point, (b)). Surface ignition was thus defined by the occurrence of a hysteresis loop in the experimental power-temperature diagram. This procedure was repeated with varying methane content of the reaction mixture, thus determining the ignition-extinction behaviour over the whole range of (catalytic) flammability. To assure reproducibility, of the results, data acquisition and control of the power unit were computer controlled, strict criteria for a stationary point were applied, and the experiments were repeated with several different foils. Throughout this paper, the catalytic ignition behaviour will be discussed in terms of ignition temperature vs a corrected equivalence ratio. While the equivalence ratio (F is usually' defined in the combustion literature as the ratio of the current fuel/air ratio divided by the fuel/air ratio at the stoichiometric composition for total oxidation to H20 and CO2. we prefer to depict the data vs a modified equivalence ratio which we define as: ~F/(1 + ~). This modification has the advantage that it puts equal weight on the fuel lean and fuel rich sides of the ignition curve, i.e. while the usual (F maps fuel lean mixtures on a scale from 0 to 1 and fuel rich mixtures on a scale from 1 to infinity', the modified ratio maps

275

[]

~n

800-

L) %

~oD D

,•

[]

[]

fl+

400- ( b l

I o(a)

\)

9

_

~ I

o

I

5

~w~/W

15

Figure 1" Typical S-shaped ignition-extinction curre, showing a experimental temperature trace t,s external power input for one fixed fuel:air ratio. In addition to surface ignition (a). extinction of the surface reaction (b) is also indicated. The inset show.s the experimental stagnation point set up. fuel lean mixtures on a scale of 0 to 0.5 and fuel rich mixtures onto 0.5 to 1. Thus. the stoichiometric point for total oxidation to CO~ and H20 is at 0.5 on this axis. with typical partial oxidation reactions occurring on the fuel rich side of this axis (~/(1 § ~) > 0.5) and catalytic combustion typically occurring under fuel lean ~/(1 § ~) < 0.5) conditions. Throughout this paper, the term "'equivalence ratio" will be used to refer to this modified equivalence ratio unless stated otherwise. 2.2. Experimental Results F i g u r e 2 shows the results for the catalytic ignition of m e t h a n e / a i r mixtures on a platinum foil. No catalytic ignition is observed for equivalence ratios below 0.2. From there, the ignition temperatures show a continuously decreasing trend with increasing methane/air ratio up to the upper surface ignition limit at ~/(1 + ~) ~ 0.SS. Beyond that ratio, no catalytic ignition occurs any more. Under no conditions any coke formation was observed in the experiments. Ignition temperatures were highly reproducible and independent of the "'history" of the catalytic foil. i.e. independent on whether the foil had previously been exposed to fuel rich or fuel lean conditions. The most obvious result of these experiments is that catalytic ignition occurs at much lower temperatures and over a much wider range of fuel:air ratios than homogeneous ignition. While homogeneous ignition temperatures for m e t h a n e / a i r mixtures lie well above 1000~ catalytic ignition occurs at temperatures of 600 ~ and below, and. while homogeneous ignition is confined to a relativeh: narrow flammability range between an equivalence ratio of 0.33 to 0.6. this range is extended significantly by the Pt catalyst. Also. no visible break occurs in the ignition curve when crossing from very fuel lean to very fuel rich mixtures at ~/(1 + ~)=0.5. Obviously, this point has no special significance for the catalytic surface reaction.

276 700

pc-~zi5' ex~zirr~t

600

O"O

-O

-D

-[:j .rn __

L)

%

O O O - - - - r l -..

500

121. 10 [2

400

C]

o

300

0

'

I

0.2

'

I

0.4

'

i

16

0.8

0.

"~

1

Figure 2: Ignition curve for methane~air mixtures on a platinum foil catalyst. Shown is the catalytic ignition temperature v.s equivalence ratio. (Xot all experimental data points are shown.) The observed continuous decrease in ignition temperatures with increasing fuel content of the reaction mixture is strongly indicative of site blocking by oxygen on the catalyst surface, poisoning the catah'st for the reaction. This hypothesis is supported by a simple analysis based on a very simplified model for the catalytic surface reaction (Veser and Schmidt. 1996), assuming a simple, first-order Langmuir-Hinshelwood (LH) mechanism. in which methane and oxygen adsorb competitively and non-dissociatively on the surface (S) according to: C H4 + S ~ C H4,s 0 2 + S ~ Oz,s 02,S + C H4,s --+ Products + 2S Obviously, the assumption of non-dissociative adsorption as well as that of the reaction being first order in both reactants are gross simplifications, since both methane and oxygen will dissociate prior to the reaction, and the further surface reaction mechanism will involve several consecutive and parallel reaction steps. Nevertheless. the model should allow for a qualitative test of the assumption that site competition between methane and oxygen account for the observed trend in ignition temperatures. The standard kinetic rate equation for the mentioned LH mechanism can be further simplified by assuming that the adsorbate coverages up to the ignition point are determined by adsorption-desorption equilibrium. This also allows us to disregard any mass transport limitations in the gas phase above the catalyst surface, since these onh' become relevant after ignition of the surface reaction. If we finally take into account the hypothesis that a higher oxygen sticking coefficient is responsible for the observed trend and thus assume a mostly oxygen covered surface (at least for the fuel lean side of the ignition curve), we obtain the following overall rate expression which yields a direct proportionality of r with r

--

IX'CH4.o , lc~o 1(o2,o

.e-R.r

E~vv

.(I)

(1)

where /c~.o denotes the kinetic preexponent, ial factor, I(i.o the preexponential of the adsorption/desorption equilibrium constant for methane and oxygen respectively, and

277 3.0 Pt foil/ analytical model 2.5. /")

2.O-

1.5-

"-..?

1.0-

--....

0.5-

........ (> G

0.0700

750

800

T/K

850

900

Figure 3: Model fit (dotted line) for the simplified, analytical model and experimental data

(symbols) for fuel lean methane/air mixtures. The diagramm depicts the (conventional) equivalence ratio d~ t,s foil temperature. denoting the (standard) equivalence ratio. The apparent activation energy E~pp sums the activation energies for the surface reaction as well as the desorption of the hydrocarbon and oxygen by E~pp = ER + AHacI,CH4 - AHad.O~. A heat balance which includes conductive heat losses ( a ( T - T o ) ) , resistive heating (V-I) and the heat evolved by the surface reaction (r AHR) yields an analytic expression for the equivalence ratio as a function of temperature after applying the ignition criterium and substituting the above rate equation: ,~(~)

(

-

1 -

a

Eapp

- ~ . ~---~-

)1

(2)

where "a" is a proportionality factor that lumps several preexponential and geometrical factors. This function can be fitted to the measured ignition curve for methane, yielding apparent activation energy of about E~pp ~ 110 kJ/mol. This seems a very reasonable value for a surface reactions and yields a very close fit between calculated and experimentally measured ignition temperatures for ~/(1 + ~) < 0.6 (~ < 2.5, see figure 3), which strongly supports our hypothesis that site competition between the hydrocarbon and oxygen during adsorption is mainly responsible for the observed trend of decreasing ignition temperatures with increasing methane content in the gas phase. For higher values of ~. this simple model is not capable of describing the increasingly steep drop in ignition temperatures, since the assumption of a predominantly oxygen covered surface in the model breaks down at this point. This emphasizes that, while simplified chemical mechanisms can yield very good fits with experimental data if the chemistry is well understood, their applicability is typically limited to a rather narrow parameter space, and detailed reaction mechanisms are needed for a reliable description over a wide parameter range.

278 3. DETAILED 5'IONOLITH _X,IODEL

In the above, we described experiments using a stagnation point set-up for easy. reliable and well reproducible tests for ignition and extinction in a catalytic surface reaction. along with a very simple analytical model for surface ignition. While these results yield important 'fingerprints' for understanding the catalytic ignition mechanism of the surface reaction, these experiments can not yield data which would be directl~l applicable to ignition in different catalytic reactor configurations. On the other hand. reliable and meaningful measurements of ignition temperatures are almost impossible for a typical reactor set-up. Therefore, we choose a somewhat different approach here: we set up a detailed onedimensional model for a catalytic monolith, which comprises complete heat-, mass- and impulse balances for the reactor, coupled with a detailed surface reaction mechanism. This model was then tested against steady state experiments with a catalytic monolith reactor. After obtaining a good agreement between the model results and experimental data, we did dynamic simulations of catalytic ignition in the monolith reactor with this model. The ignition temperatures we obtained with this model were then compared to the experimental ignition temperatures obtained on the Pt foil. The procedure will be explained in detail in the following paragraphs. 3.1. Reactor Model and Numerics For the simulation a dynamic 1D-model based on mass and energy' balances is used. A complete and detailed derivation of the model is beyond the scope of this paper an will be presented elsewhere (Veser and Frauhammer. to be published). The mass-balance equation contains mass transport by" convection, dispersion as well as adsorption and desorption from the catalyst surface:

Owj _ :.o.

at

ill Owj

-

--~

. 0.:

02wj

~ :. D~:: . 0:2

z

:,:.

(:3t

i

The energy balance of the gas phase, comprising heat transport by' convection and conduction, yields:

:-e-c~

OTa 0,

_

3)

-

.a-~-~+:"~ss~

OT~

02Tj

+~v~

(4)

and the energy balance of the monolith yields" :-o~

OT~ c,- 0t

=

(I-s)-A,-

O2Ts 0:2

av.a.(T,-Tg)+av.~rk',XH~.k

k

(5)

For the pressure drop calculation, we assume laminar flow in the monolith channels. The mass balances on the surface. Eqn. 7 - 19. will be discussed below (see 3.2.). The above shown system of balance equations can be written in the form

0~7 B-:ot =

-~~

0.g

0~7 + b - c - 0.7. +o

-

(61

279

which is the general form of a 1D parabolic differential equation system. Together with appropriate initial and boundary conditions, the system is solved numerically using the fully adaptive method-of-lines (~IOL) algorithm PDEXPACI< (Nowak et al., 1996b). This algorithm has been developed to solve parabolic differential equations of the general form given in Eqn. (6) and is based on the code PDEX (Nowak. 1996a). The numerical algorithm is fully adaptive in time and space. The number of grid points necessary to achieve a certain accuracy for the approximated solution as well as the distribution of the grid points and the length of the time step are automatically adjusted by' the program, based upon the relative errors for the time and the space discretization for every state variable at every gridpoint. The local spatial errors which are estimated at each time step are the base for a local regridding to equilibrate the error in the spatial domain. Based on the local error estimates, global error norms are determined which allow for a separate error control in space and time. This error control procedure and adaptive regridding turn out to be essential to accurately resolve the very steep reaction fronts typically observed on the catalyst surface during ignition processes of catalytic reactions. 3.2. Surface Reaction Mechanism The surface reaction mechanism is based on the mechanism by Hickman and Schmidt (1993b) and comprises the following balance equations for C, CO. CO2. O. H. OH, and H20 on the catalyst surface: OOc Ot

_

--

]Cads,CH4 " P C H 4 "

=

]Cads,CO " P C O

0 e

free -- k/otto,CO

" OC " O 0 n t- lCdeco.CO " OCO " O f tee

(7)

OOco " O f tee - - k c l e s . C O " OCO

8t +kf .... co 000 8t

O,

O00H Ot

=

--

=

9Oc . 0 0 -

ka . . . . c o

9O c o

9Oiree -

kre~c.c02

Oco ' 00

9Ofree -- kre~c,CO~

OCO ' O0

(8)

2 2 . kads,O2 " P o 2 " Ofre~ -- 9-" . lode s 90.~_ " 020

-kfor~,co

Oc 9 . 00 + kdeco,CO " 0 c o

--lee ....

00"

OH"

2" kads.H2

OH Jr- k d . . . . O H " O O H "

" P H 2 " O 2f t e e - - 9- .

lCdes.H2

--~f ....

O H " O0 " OH -J- lgd . . . . O H

--]r ....

H.20 " O O H " OH -~- kct . . . . H 2 0

k f .... oH

Of tee + ]'C. . . . .

O H " O'~H

02H + 4 " kads,CH 4

(9)

PCH4 " 02_free

" OOH " Ofree

(10)

" O H 2 0 " O free

9Oo 9OH -- kdeco,OH " O o n " Of tee - - 2 " k . . . . . O H " O ~ H

- - k f o . . . . H 2 0 " O O H " OH nu lCdeco.H~O

9O H 2 0 "

Of tee

(11)

OOH~_O

Ot

=

tcacS.H~_O 9P H 2 0 " Of tee -- lOdes,H20 " OH20 -qt-kform,H20

9O O H " OH - - k d e c o . H 2 0

" O H 2 0 " O i r e e -'~ k . . . . . O H

" O'~H

(12)

We assume dissociative adsorption for CH4, 02 and H2. For methane adsorption, complete dissociation into C and 4 H is furthermore assumed, since no experimental data is available for individual dissociation steps. While this enters the species balance on the surface, the kinetic rate law for dissociative methane adsorption is formulated with

280

Table 1" K i n e t i c R a t e P a r a m e t e r s reaction So k Is-1] CH4.9 ~ C + 4H 0.01 C + 0 --+ CO 5 91012 C O -+ C + O C O + O --+ C'O~ 9 CO~ ~ CO

02,9 ~ 20 H~.,9 r 2H

5 9 1013

0.84 0.01 0.05

10 TM 1015 1013 10 ~3 5 91013 5 91013 5 910 z3 1013 10 ~3 10 ~3

Eact [kJ/mol] 70 60 180 100 - (90" 50 142.4 215- 19o-60 74 10 18 16 94 50 42

reference Sun and Weinberg Hickman and Schmidt (1993b) Hickman and Schmidt (1993b) Campbell et al. (1980) Campbell et al. (1981) 5latsushima XlcCabe and Schmidt Williams et al. Williams et al. Germer and Ho (thermodynamic constraints) Fisher and Gland Fisher and Gland

0 + H --+ O H OH ~ O + H O H + H -+ H~_O H~_O --+ O H + H O H + O H -+ H , O + 0 H20 9 ~ H20 0.1 all species denote surface species unless denoted with the subscript g

a square dependence on the free surface sites O : ~ . since mechanistically this step will obviously only require two free adsorption sites. Except for m e t h a n e adsorption, all adsorption steps are non-activated, and follow the usual rate law as derived from kinetic gas theory: k~d~.~ =

si

v/2 9:r. 31i . t:l. T

(13)

where si denotes the sticking coefficient and Mi the molar mass of the component i. All other kinetic rate constants follow the usual Arrhenius law. Since our intent was to study catalytic ignition, we abandoned the assumption by Hickman and Schmidt (1993b) that oxygen adsorbes non-dissociatively onto different adsorption sites than the other reactants. This assumption, while valid at the high temperatures during steady, state operation of the reactor, will break down at the comparatively low temperatures of surface ignition processes. Thus. the free surface sites O:r~ obey' the algebraic equation: Of~e~ = 1 - ~ 0 i. (14) i

While the assumption of only' one type of surface sites for all adsorbants turned out to be an important factor in our results, assuming dissociative or non-dissociative adsorption did not change the results much. Obviously. under high t e m p e r a t u r e conditions, free surface sites do not play, an important role in the surface reaction mechanism, while rather the competition for adsorbed oxygen appears to be the limiting factor for the different parallel surface reaction paths. All activation energies and sticking coefficients in the model were taken from experimental investigations published in the literature. The pre-exponential factors, however, have not been measured for most surface reaction steps. Where reliable data was available in the literature, those values were used. For the other reaction steps, we assumed a (pseudofirst order) kinetic constant of 1013 s -1. which can be derived from transition state theory for a surface reaction step in which the transition state complex is not too different from the adsorbed state, i.e. if the ratio of the respective partition functions is close to unity (Zhdanov et al.). Starting from these initial values, we adjusted the pre-exponential factors until a satisfactory agreement between experimental data and model results was

281

obtained for catalyst temperatures, methane conversion and syngas selectivities. T a b l e 1 summarises the parameters used for all simulations. 3.3. Steady State of the Monolith Reactor F i g u r e 4 shows the model results (lines) along with experimental data from steady state experiments (symbols) during methane partial oxidation with air on a Pt coated monolith for two feed temperatures. Obviously, the model very closely reproduces both the catalyst temperature and methane conversion over the entire range of methane/air mixtures studied, while syngas selectivities are in less good agreement with the experiment. The catalyst temperatures, which were fitted using a fixed heat loss coefficient of :32 \V/m•I,2 to describe the heat losses in the experiment, lie within less than 50t,2 (5 5~) from the experimental data. Like in the previous study by Hickman and Schmidt. however. we were not able to reproduce the plateau for methane/O.2 ratios above 1.6. Methane conversion agrees well between the experiment and modelling results, although the model somewhat overestimates conversion rates at low CH4/02 ratios. This is most likely' a result of assuming irreversible methane adsorption, since under the high temperature conditions in this range adsorbed CH4 will start to desorb before dissociating and reacting further and therefore methane conversion will decrease. The model curves for CO and H2 selectivities vs methane/O2 ratio deviate from the experimental data somewhat more strongly. However. both curves reproduce the qualitative trends very well both H2 selectivity curves show the single maximum observed in the experiment, although this maximum is shifted by about 0.2 towards fuel leaner mixtures. The CO selectivity is the quantitaively most strongly deviating feature in the model. being overestimated by up to 10~. In this respect., the new model does not yield an improvement over the Hickman and Schmidt model. It does. however, reproduce very well the main qualitative features in the strong decrease towards the fuel leaner side. and particularly; the cross-over between the selectivities for 25~ and 460~ feed temperature at a methane/air ratio of 1.2. Thus. it appears that. the model describes the essential features of the surface chemistry very well. 3.4. Catalytic Ignition in the Monolith Reactor Based on the good fit between the model results and the experimental data. for the steady state of the reactor, we now attempted to model catalytic ignition in the monolith. For this case, no experimental data is available due to the difficulty of measuring well-defined ignition criteria, and due to the experimental difficulties to externally heat. a monolith reactor to the required ignition temperatures. Experimentally'. ignition is therefore typically achieved by' pre-heating the monolith through a more easily' ignitable catalytic reaction (such as ammonia oxidation), and then switching the feed composition to the desired fuel (i.e. here" methane). However. if ignition is determined by' the chemistry of the surface reaction, the reactor configuration should only influence the absolute value of ignition temperatures, while at least the qualitative trends in the ignition behaviour with the change in feed composition should be identical between the monolith reactor and the simple stagnation-point foil experiment. The results from the simple analytical model (sect. 2.2.) allow us to define a rather simple ignition criterium for the simulations in terms of a vertical tangent of the temperature derivative of the oxygen surface coverage at. the ignition point, i.e. 0 0 o / 0 7 ' -+ oc. It is

282 0.8

.,..,. , ,,- ..................... ~ ,"/ . ,'" ~

0.6-

0.9

o

o

+

+ 0.4-

. . . . . . . . . .; . . . . + . . . . . . .

+ "--....

§ .. - . . .

~....

.......

. --....+

0.8

0.2-

0

018

~

11.2

114 CH4/O2

1'.6

1'.8

1

2

0.7

ola

/,

~12

1

114 CH4/O2

11.6

118

1700

'"'",.,. ,,,,~,,,.

.... 1600 -

0.915000.8-

1400-

',

o

",,, +

o o

' 9, ~.

1300 -

0.7-

o ..

<>

,-~, "-.+.

+ -..

~.

+

"-...

1200 0.61100 0.5

0'.8

1

11.2

114 CH4/O2

"'~.T

1.6

118

1000

018

1

112

114

CH4/O2

116

118

Figure 4: Ezperimental data (symbol.s) and modellin 9 results (lines) for methane o,ridation in a monolith reactor at '25~ (diamonds and dashed lines) and/~60 ~C (crosses and solid lines) feed temperature. Shown are (clockwise from, top,) H2 selectvitiy, CO selectivity, catalgst temperature, and methane conversion us CH4/02 ratio in the feed 9as. (Ezperimental data from Hickman and Schmidt. 1993).

283 90O

850

800 a_:

750

700

650 -.. --600

e x p e r i m e n t a l d a t a / P t foil detailed model/monolith

,',,,l,,,'ri,,, *

i , , , , i , , , , i , , , , l , , , , I , , , , i , , w , l , , , , i

0.5

0.6

0.7

~/(]+~)

0.8

09

Figure 5" Ignition temperatures vs equivalence ratio during CH4-ozidation on Pt" foil

ezperiments (symbols and dotted line) and monolith model (dashed line). (Please note that ezperimental temperatures are given in ft. while model results are in ~ interesting to note that the results of the numerical analvsis showed that this criterium was typically also fulfilled for all other surface coverages. It seems therefore that the hysteretic behaviour observed in the catalyst temperature is also reflected in all surface coverages. A detailed discussion of the ignition criterium is beyond the scope of this paper and will be given elsewhere (Veser and Frauhammer, to be published). The results for ignition in the monolith model are given in figure 5 along with experimental data from the foil experiments (please note that temperatures from the experiment are given in K. while those from the simulations are in ~ Obviously, the model (in which all parameters are unchanged from the steady state simulations) yields a perfect fit for the qualitative trends in the ignition temperatures vs equivalence ratio, though quantitativel 9 the ignition temperatures in the monolith model are about 300K above the ignition temperatures from the foil experiment. The latter is not very surprising, however, since one can expect that the much larger thermal mass of the monolith versus the thin (0.2 mm) catalytic foil as well as the very different flow patterns in these two configurations will lead to a shift in surface ignition towards much higher temperatures in the monolith reactor. The very close agreement in the trend between the model and the experiment, on the other hand. seems all the more remarkable as this trend has been shown to be quite peculiar to methane/air mixtures" Previous experimental investigations in which we compared the ignition behaviour of several different hydrocarbon/air mixtures over Pt foil catalysts showed that for all alkanes except methane, ignition temperatures fell up to an equivalence ratio of ~ 0.S and then started to rise again towards the fuel rich end of surface ignition, while surprisingly for methane an opposite trend, i.e. an increasingly' steep drop in ignition temperatures was observed in the fuel rich region (Veser and Schmidt). We hope that a future extension of this surface reaction model towards ethane (and higher hydrocarbons) will allow to gain a deeper understanding of the differences in the observed ignition behaviour of different hydrocarbons in this range.

284 4. SUMMARY In the present paper, we reported experimental and modelling results for catalytic ignition of m e t h a n e / a i r mixtures over platinum catalysts. Experimentally. the ignition behaviour was tested on Pt foil catalysts in a simple stagnation point set-up. A very simplified analytical model was able to describe the observed ignition behaviour for fuel lean methane/air mixtures very well. demonstrating that mechanistically surface ignition is determined by' site competition between methane and oxygen during adsorption on the surface. A detailed one-dimensional model was then set up for a monolith reactor, comprising complete heat-, mass-, and impulse balances for the reactor along with a detailed elementary step surface reaction mechanism. This model was first tested against stead) state experiments in a monolith reactor, for which it yields a very good qualitative and a satisfactory quantitative agreement. In dynamic simulations with this model, catalytic ignition of the monolith reactor was then compared to the experimental ignition temperatures obtained from the foil experiments. Again. the model yields a very good qualitative agreement with the experiment. Quantitatively, the model predicts ignition temperatures well above the temperatures in the foil experiments, which can be explained by the different reactor configurations. The good agreement between the model and both steady state and ignition experiments indicates that the surface reaction mechanism describes the essential steps of the catalytic reaction very accurately. The next step in developing this model further will now be to fill in the remaining gaps in the surface mechanism for methane oxidation and to extend the mechanism towards C2 chemistry on the catalyst surface.

REFERENCES S. Bharadwaj and L.D. Schmidt (1995), Fuel Proc. Technol. 42. 109 C.T. Campbell, G. Ertl. H. Kuipner, and J. Segner (1980). 3. Chem. Phys. 7:3, 5862 C.T. Campbell, G. Ertl. H. Kuipner, and J. Segner (1981), Surf. Sci. 107, 207 G.B. Fisher and J.L. Gland (1977), Surf. Sci. 94,446 T.A. Germer and W. Ho (1989). Phys. Lett. 188. 449 D.A. Hickman and L.D. Schmidt (1992), 3. Catal. 138. 267 D.A. Hickman and L.D. Schmidt (1993a), Science 259.34:3 D.A. Hickman and L.D. Schmidt (1993b). AIChE J. 39. 1164 T. Matsushima (1985). Surf. Sci. 157. 297 R.W. McCabe and L.D. Schmidt (1977), Surf. Sci. 65. 189 U. Nowak (1996a). Appl. Numer. Math. 20. 129 U. Nowak. J. Frauhammer and U. Nieken (1996b), Comput. Chem. Eng. 20. 547 L.D. Pfefferle and W.C. Pfefferle (1987). Catal. Rev. - Sci. Eng. 29.219 Y.-K. Sun and W.H. Weinberg (1990), J. Vac. Sci. Technol. A8. 2445 P. Torniainen and L.D. Schmidt (1994), J. Catal. 146.1 D.L. Trimm (1983). Appl. Catal. 7. 249 G. Veser and L.D. Schmidt (1996), AIChE J. 42, 1077 W.R.C. Williams, C.M. Marks and L.D. Schmidt (1992), J. Phys. Chem. 96. 5922 \:.P. Zhdanov, J. Pavlicek. and Z. Knor (1988), Catal. Rev. - Sci. Eng. 30. 501

91997 Elsevier Science B. V. All rights reserved. Dynamics of Surfaces and Reaction Kinetics in Heterogeneous Catalysis G.F. Froment and K.C. Waugh, editors

285

Modelling Catalytic Cracking Kinetics Using Estimated Adsorption Equilibrium Constants B. Sowerby a and S.J. Becker u a School of Chemical Engineering, University of Bath, Claverton Down, Bath, BA2 7AY, England b BP Chemicals Limited, Technology and Engineering Department, Poplar House, Chertsey Road, Sunbury-on-Thames, Middlesex, TWl 6 7LL, England

1. INTRODUCTION Fluid catalytic cracking (FCC) is a major unit operation when refining crude oil to produce high-quality gasoline. The need for accurate and fundamentally based process models is becoming more important in design and operability studies as profit margins are squeezed and as new environmental regulations must be met. To date, because of the complexity of the problem, kinetic models that have been developed for FCC risers, strippers and regenerators have been of a simplified form. For example, FCC riser models have traditionally been based on pseudo-components and power law rate expressions. Many workers have implemented Weekman's ten lump cracking model [1]. Such lumped parameter models have been used extensively and successfully for control and qualitative operability studies [2]. The major drawback of such an approach to kinetic modelling is that the models require some retuning for different catalysts and feeds and that it is difficult to describe non-homogeneous catalysts and the effects of inhibition and poisoning. Recent developments in analytical and experimental techniques have made it possible to identify the individual species in FCC feeds and products, to determine the strength and distribution of the catalytic active centres (acid sites) and to study the structure and composition of deactivating coke. In addition, for a variety of model compounds, a wealth of published qualitative and semi-empirical data exists that highlight some of the key features of FCC catalysts. This includes the inhibiting and poisoning effects of various species and deactivation by the progressive reduction in the number and changing distribution of acid sites. To be able to quantitatively describe and predict the aforementioned phenomena and to be able to relate catalyst properties to unit operation performance, a more detailed description of the species involved as well as a better representation of the fundamental processes that are occurring between the bulk fluid and the catalyst surface than that which is currently employed in pseudo-component, lumped parameter, power law models is required. This more fundamental approach to kinetic modelling has been achieved in many other systems where there are only a few components and reactions by using Langmuir-Hinshelwood and/or EleyR.ideal type rate expressions; such expressions are usually developed by considering the

286 individual adsorption, reaction and desorption steps and are validated using rigorously defined experimental programmes. This approach, is not particularly well suited to developing more fundamentally based FCC models because of the large number of species, reactions and thus adsorption and rate constants involved. To develop more fundamental FCC models on the basis of rate expressions that take into consideration the surface of the catalyst a more predictive approach for obtaining the adsorption and rate constants is required. Such an approach must also be able to be validated using a limited but representative range of model compounds. This paper presents a practical approach for estimating adsorption equilibrium constants using an integrated form of van't Hoffs isochore equation. The equilibrium constants can then be used in Eley-Rideal kinetic rate expressions. A summary of how the theoretically derived adsorption equilibrium constants have been used to predict catalyst surface compositions during the catalytic cracking of 2-methyl pentane is presented. For the first time the results of another study in which the inhibiting and poisoning effect of more than twenty five nitrogen and aromatic compounds typically found in FCC feeds are discussed. The work presented in this paper draws on results and observations from a diverse source of published literature and a new approach to modelling of FCC riser and stripper kinetics is demonstrated.

2. ESTIMATION OF ADSORPTION EQUILIBRIUM CONSTANTS In a recent paper [3] we described how adsorption equilibrium constants for a range of paraffinic and olefinic species could be estimated using the following integrated form of van't Hoffs isochore equation, RTInK~= -AH~+TAS~+C1

(1)

where R is the universal gas constant, T is the temperature, Kp is the adsorption equilibrium constant, AH~ is the heat of adsorption, AS~dis the entropy of adsorption, and C1 is a constant of integration. To estimate adsorption equilibrium constants using equation 1 requires sources for the heat of adsorption and the entropy of adsorption for each component. In FCC catalysis, the adsorption process can be considered to consist of two stages, first a physisorption step (akin to condensation on to the catalyst surface) followed by a chemical attachment to an active catalyst site prior to reaction. The physical part of the sorption process has been assumed to be equivalent to condensation and to require the loss of all translational degrees of freedom of the vapour molecule. This first assumption allows published values for the heat of condensation to be utilized and the second gives a feasible approximation to the entropy change on adsorption. Following physisorption, chemisorption can be achieved by exchange of a proton between a physisorbed hydrocarbon species and a catalyst acid site. Thus, chemisorption results in the production of a positively charged hydrocarbon ion; the heat of ionisation associated with this step has been assumed to be the negative of the proton affinity.

287 Applying these assumptions to equation 1 yields, In K~ = :~H + PA - eimm + A_~SSd + C 1 2 RT R

(2)

where He is the heat of condensation, PA is the proton affinity of the adsorbate and PAmm is the proton affinity for the catalyst. In equation 2, it has been assumed that all of the proton charge of an acid site is donated to the hydrocarbon, this may not be absolutely correct but is deemed a valid assumption here. In addition, the two heats associated with the adsorption process, Hc and (PA-PAmm), have been combined together using an analogous approach to that used for combining activation energies of sequential diffusion and rate processes [4] which gives rise to the factor 2 on the denominator. The proton affinity for the catalyst, PAmm, is dependent on the catalyst's acid site strength. Parillo et al [5] related the heat of adsorption to the proton affinity for the adsorption of a series of simple amines on ZSM5 and found that a single acid site strength of 712 kJ/mol adequately described their experimental data. Thus, a value of 712 kJ/mol has been used in this initial analysis. Finally, the integration constant, C1 has been estimated to be 29.1 by using Maatman et als [6] equilibrium data for the adsorption of benzene on silica-alumina. The loss of all the translational degrees of freedom of the vapour molecule has been determined using the Sackur and Tetrode equation that is given in Moore [7] as, AS~d = R In e 5''2 v (2 rc m k T) 3'2 L h3

(3)

where e is a constant, v is gas molar volume, m is the mass of one particle, k is Boltzmanns constant, L is Avogadro's number and h is Plank's constant.

3. A P P L I C A T I O N OF E S T I M A T E D Kp'S TO 2 - M E T H Y L P E N T A N E C R A C K I N G This section summarises our previously published work on using Eley-Rideal kinetic rate expressions to describe the inhibiting effect of reaction product olefins on 2-methyl pentane cracking with the adsorption equilibrium constants being estimated using equation 2 [3]. In the next section the theory is extended to the complete range of species found in typical FCC feeds including poisoning nitrogen and aromatic compounds. For 2-methylpentane cracking the assumed reaction mechanism was based on that proposed by Zhao et al [8]. Reactions are initiated by adsorption of 2-methylpentane feed on to Bronsted acid sites. Subsequent protolysis produces a carbenium ion on the surface of the catalyst (an adsorbed olefin) and a smaller gas phase paraffin molecule. Propagation reactions can then occur by hydride transfer between carbenium ions on the surface of the catalyst and gas phase feed molecules. Zhao experimentally determined certain bimolecular reactions to be more significant than others, the most significant were implemented in our model. Reactions are terminated by desorption of carbenium ions to yield a gas phase olefin molecule

288 and a free acid site. As the isomerisation of carbenium ions adsorbed on the surface will naturally occur an isomerisation reaction was also included. A reaction for coke formation by carbenium ion polymerisation was not included, however catalyst deactivation by coke laydown was accounted for by implementing a time on stream dependent deactivating function as developed by Zhao et al [8]. The reaction rate expressions used in the model were of the following form, rate of reaction j,

rj = kj___~i 12i~ (1 +Y'.Kp~i)

(4)

where kj is the rate constant, Kp, is the adsorption equilibrium constant for component i and p, is the partial pressure of i. Before the rate expressions could be used to predict reactor behaviour, values for the adsorption equilibrium constants Kpi and the reaction rate constants kj needed to be defined. Equation (2) was used to calculate the adsorption equilibrium constants using PA values taken from standard heat of ionization tables [9]. The values used are given in Table 1. The standard heat of condensation at 25~ used for each component is also given in Table 1 along with the calculated K~ values. The reaction rate constants were established by matching model predictions to Zhao et al's [8,10] experimental data where the cracking of 2-methylpentane was studied using a fixed bed gas-phase plug flow reactor packed with HY zeolite catalyst at 400~ and 1 atm pressure. Table 1 Proton affinities, heats of condensation and equilibrium constants Component

PA He Kr,~ kJ/mol kJ/mol

hydrogen methane ethane propane butane pentane 2-Me-pentane C6 isomers ethylene propene butene pentene hexene

628 628 628 683 683 683 683 683 720 720 835 848 855

0.91 8.19 14.72 14.84 20.62 26.29 31.56 29.12 13.55 14.12 19.60 25.07 30.35

0.191 x 0.366 x 0.175 x 0.515 x 0.123 x 0.258 x 0.506x 0.325 x 0.151x 0.463 x 0.302x 0.109 x 0.317 x

]0 -32 10 9 10 -8 10 .8 10 .7 10 .7 10 .7 10 -9 10.4 10 .4 10 .3 10 .2 10 .2

The variation in predicted weight percentage of various products as a function of overall conversion agreed well with the published experimental data. The reactor model developed was also able to predict that for an overall conversion of 10% then over 80% of the surface of the catalyst would be covered with adsorbed C 6 olefin (C6H13" carbenium ion). However, once sufficiently high quantities of product olefins had been formed, the model predicted that the surface would become dominated by product olefins because of their relatively high adsorption equilibrium constants compared to that of the paraffinic feed.

289 4. E S T I M A T I O N OF Kp'S F O R N I T R O G E N C O M P O U N D S The poisoning effect of more than twenty five individual nitrogen and aromatic compounds on cracking catalysts has been studied experimentally by Fu and Schaffer [11]. They deduced that the poisoning effect of a species is a function of it's gas phase proton affinity. Ho et al [12] reanalysed Fu and Schaffer's experimental results and after correlating the poisoning power of each nitrogen compound with twenty four structural variables they found two variables to be significant; heaviness/size and basicity. The former was quantified by a compounds molecular weight and the latter by it's proton affinity. In this section our theory for predicting adsorption equilibrium constants is extended to cover the complete range of species found in a typical FCC feed and comparison of experimental and predicted results show that the poisoning effect of nitrogen and aromatic compounds can be predicted.

4.1 Experimental poisoning power Fu and Schaffer [11] determined the effects of nitrogen and aromatic compounds on cracking catalysts by feeding gas oil and hydrotreated vacuum resids to a micro-confined fluid bed unit that contained around 35 g of cracking catalyst. A catalyst to oil ratio of 6 was used with reaction temperatures ranging from 510 to 565~ A full description of the experimental details is given in their paper. Results for two different sets of conditions were presented. In the first set of experiments, the conversion percentages were measured for eighteen nitrogen and aromatic additives using a West Texas Refinery equilibrium catalyst. The amount of additive was adjusted to give an overall 0.5 wt% nitrogen in the feed for additives containing one nitrogen atom. For additives containing two nitrogen atoms, the amount added was such as to give an overall 1 wt% nitrogen. For additives not containing nitrogen, i.e., benzene, naphthalene and anthracene, the amounts added were the same as the corresponding amounts of pyridine, quinoline, and acridine, respectively. In the second set of experiments, a series of pyridine derivatives and piperdine were used with a metal-free catalyst at 0.3 wt% nitrogen.

4.2 Prediction of poisoning power In Eley-Rideal and Langmuir-Hinshelwood type of kinetic rate expressions, the effect of poisons and inhibitors on the reaction rate is accounted for by allowing part of the catalyst surface to become covered with the poisoning compound and so unavailable for desirable reactions. For example, consider the decomposition reaction of A --~ B + C that occurs in the presence of an inhibitor, I. If it is assumed that the inhibitor does not participate in the reaction but that it does occupy active catalyst sites and if the surface decomposition is rate controlling then the observed rate of reaction in Langmuir-Hinshelwood terms is given by [ 13 ] -rA =

....k~C,___~A_(__~ - Pp PB / Kr 1 +

+

(5)

+

where -r A is the rate of disappearance of A, k S is the surface reaction rate constant, C, is the total number of catalyst sites, P1 is the partial pressure of component i, Kpi is the equilibrium adsorption constant for component i and K e is the equilibrium constant for the reaction. Subscripts A, B, C and I represent components A, B, C and inhibitor, respectively.

290 In this work, the amount of poison that is capable of adsorbing on to the surface of the catalyst has been used as an indication of the poisoning power of that compound. The amount of poison that will adsorb on to the surface of a catalyst has been determined by performing a flash calculation between the vapourised species and the species adsorbed on the surface of the catalyst using a typical FCC feed containing a given amount of poison. Such a calculation yields vapour and catalyst surface compositions. The mass balance for a component between the vapour phase and on the catalyst surface is, F i = V Cbi + W Cs,

(6)

where F, is the total amount of component i in the oil flashed, kmol, V is the vapour volume, m 3, Cb: is the bulk concentration of each component, kmol/m 3, W is the weight of catalyst, kg, and Csi is the surface concentration of each component, kmol/kg. The equilibrium distribution of each component between the vapour and catalyst surface is, K~, =

Cs,__

Cb: x s

(7)

where Kp: is the adsorption equilibrium constant and s is the available surface, kmol/kg. The surface partitioning is then defined by, s =

s

(8)

1 "+" X Kpi Cbi where S is the total amount of catalyst surface, kmol/kg. For a given amount of poison i in an oil feed, if the adsorption equilibrium constants, vapour volume, mass of catalyst, surface of catalyst and total fractional coverage are known then, the surface coverages of the individual components can be determined from the solution of equations 6 - 8. The composition of a typical FCC feed is given in Table 2. The feed composition has been described using sixteen family types, within each family type a range of species of varying number of carbon atoms has been defined. The nitrogen and aromatic compounds used by Fu and Schaffer are given in Table 3 along with the experimentally measured conversions. Tables 2 and 3 also present the adsorption equilibrium constants that have been calculated using equation 2. In the calculations, published values of PA have been used where ever possible, however if no value could be found then a value has been estimated using family type and carbon number. The calculated adsorption equilibrium constants are in the range 10 s to 10 5. At first this appears to be a very wide variation however it is worth remembering that the adsorption

291

equilibrium c o n s t a n t is the ratio o f the adsorption to the desorption rate constant and thus this wide range o f v a l u e s implies that s o m e species will definitely r e m a i n adsorbed on the catalyst surface rather than desorb and will poison the catalyst by r e d u c i n g the observed rate o f reaction by a g r e a t e r extent. F o r e x a m p l e , f r o m T a b l e 3, the p r e d i c t e d adsorption equilibrium constants for aniline and p h e n a z i n e are 35.22 and 8.06 x 104 respectively. These values suggest that o n c e adsorbed, the rate o f p h e n a z i n e desorption will be a p p r o x i m a t e l y 2000 times slower than that o f aniline. So p r e d i c t i n g that p h e n a z i n e will exhibit a greater poisoning effect than aniline. I n d e e d this was f o u n d to be the case by F u and S c h a f f e r ' s in their experimental study.

T a b l e 2 C o m p o s i t i o n and properties o f a typical F C C feed feed feed comp' wt%

PA Hc Kp k J / m o l kJ/mol

n-paraffins C13 0.013 628 C17 0.013 628 C20 0.013 628 C30 6.321 628 C42 8.101 628 cyelohexanes C17 0.060 707 C25 2.320 707 C37 7.651 707 decalins C12 0.020 717 C17 0.020 717 C25 2.320 717 C37 10.15 717 perhydro-phenanthrenes C25 0.520 727 C37 4.050 727 4perhydro-pyrenes C20 4.030 737 C25 0.200 737 C37 4.891 737 benzenes C17 0.040 805 C27 1.120 805 C40 3.240 805 naphthalenes C12 0.010 846 C17 0.010 847 C23 0.450 847 C34 2.560 847

feed feed comp' wt%

65 85 1O0 150 210

3.04 2.06 8.12 6.64 1.04

85 125 185

8.13 x 10.3 0.3011 51.28

60 85 125 185

1.52 x 10-3 1.71 x 10.2 0.6372 108.8

125 185

1.348 230.9

1O0 125 185

0.3047 2.853 490.0

85 135 200

13.21 1.19 x 103 2.98 x 105

60 85 115

24.50 306.7 4.79 x 103 5.68 x 10 s

170

x x x x

10"6 104 10.5 10.3

PA Hc Kp kJ/mol kJ/mol

phenanthrenes C18 0.130 852 C28 1.250 852 pyrenes C20 0.160 858 C23 2.550 858 5 ring aromatic C22 2.740 851 tetralins C12 0.015 859 C 17 0.015 861 C24 0.800 861 C36 3.780 861 decahydro-phenanthrenes C16 0.010 859 C21 0.500 861 C21 0.480 871 C33 4.390 861 C33 2.670 87I tetrahy dro-phenanthrene C16ha 0.030 869 hydropyrenes C20 0.500 859 C30 8.961 861 decahydropyrenes C20 0.350 869 C30 6.371 871 tetrahydropyrene C20 0.250 879 C30 5.911 881

90 140

689.8 6.18 x 104

100 115

2.66

110

3.77

60 85 120 180

68.19 910.4

80 105 105 165 165

480.6 5.69 x 1 0 3 118 x 104 1.08 x 105 2.26 x 105

80

982.9

100 150

3.07 x 103 2.97 x 105

100 150

6.33 x 103 6.21 x 105

100 150

1.31 x 104 1.29 x 106

x 10 3

1.04 x 104 x 10 3

2.20 x 104

3.87 x 106

292

Table 3 Poisoning power and properties of nitrogen and aromatic compounds additive

Fu conv.

PA Hc K~ kJ/mol kJ/mol

aniline benzene pyrrole pyrrolidine pyrazine pyridine piperidine naphthalene indole quinoxaline quinoline 1,2,3,4tetrahyd'quinoline 5,6,7,8-tetrahyd'quinoline anthracene carbazole 1,2,3,4-tetrahyd'carbazole phenazine

0.527 0.518 0.541 0.503 0.526 0.514 0.495 0.522 49.6 0.432 0.392 0.386 0.364 0.543 0.543 0.517 0.420

876 759 868 873 902 924 947 846 749 894 948 918 955 851 880 875 936

45 30 35 25 40 30 30 60 60 60 60 60 60 80 70 70 70

35.22 1.2 X 1 0 .3 5.50 4.11 138.15 336.77 2.153 x 103 18.25 1.02 x 10.2 711.19 4.218 x 104 4.540 X 1 0 3 1.611 x 105 199.09 1.081 X 1 0 3 735.73 8.065 x 104

pyridine 2-methylpyridine 2-ethylpyridine 2-methyl-5-vinylpyridine 2-vinylpyridine 2,4-dimethylpyridine 5-ethyl-2-methylpyridine 2,3-cyelopentapyridine 2-p-tolylpyridine 2,6-di-tert-butylpyridine 3-methyl-2-phenpyridine piperdine 1-ethylpiperdine 2-ethylpiperdine

0.514 0.432 0.426 0.424 0.419 0.411 0.391 0.380 0.372 0.336 0.322 0.495 0.404 0.392

924 938 938 938 938 951 951 945 959 977 955 947 966 953

30 40 45 50 45 45 50 50 55 55 55 30 35 35

336.77 2.653 X 4.440 X 8.204 x 4.649 X 4.789 x 2.252 X 1.395 X 9.967 x 4.690 X 7.359 x 2.153 x 2.035 x 1.759 X

10 3 10 3

103 10 3

The feed composition given in Table 2 has been used in the flash calculations with the relevant amount of nitrogen or aromatic compound in order to determine the amount of surface occupied by each poison and hence it's poisoning power. In Figure 1, the log of the adsorbed equilibrium constant for each poisoning species is plotted against the calculated concentration of each component on the catalyst surface (that is, the adsorbed mole fraction). A log-log plot has been used here purely because of the range of Kv's and surface concentrations seen in this study, 1• .3 to 1.6x105 and 2.6x10 5 to 0.56 respectively. Benzene and indole have not been included in this figure as their calculated Kv's are so small compared to any other species.

10 3 10 4 10 4

104 10 4

104 103 104 10 4

Figure l a focuses in on surfaces coverages in the range 0.05-0.25 mole fraction and illustrates the linear relationship that exists between Kv and surface coverage; this is an inherent assumption in multicomponent Langmuir adsorption.

4.3 Comparison of experimental and predicted poisoning power Fu and Shaffer carried out their experimental study of the poisoning effects of various nitrogen and aromatic compounds in a fluid bed unit. If the flow through such a unit can be approximated by plug flow then the observed rate of reaction is proportional to In (1/(I-X)) where X is the feed conversion. It follows then that In (1/(I-X)) should be inversly proportional to the amount of each poison adsorbed on the catalyst surface.

293 In Figure 2, In (1/(I-X)) is plotted versus adsorbed mole fraction. A log scale has again been used for the adsorbed concentration due to the wide variation in numerical values. Figure 2 shows that generally there is a good correlation between the experimentally observed conversion and the amount of poison adsorbed on the surface of the catalyst. Ho et al, when they reanalysed Fu and Schaffers data, utilised a normalised poisoning parameter, y defined as y= 1 -

X 100

100-X' - X

(9)

X'

where X is the percent conversion to <221 ~ and X' is this value in the absence of additives. (X/100-X) is a widely used activity function in catalytic cracking that is based on second order kinetics. The best empirical model that they were then able to fit to the data was

(lo)

y = 0.075 + 0.7351 MW2PA 2- 0.4067 MW 3

where M W is the molecular weight of the poison divided by 121.68 and PA is the proton affinity of the poison divided by 217.29. In this expression 121.68 and 217.29 are the average molecular weight and proton affinity for the poison additives tested by Fu and Schaffer. The correlation coefficient for equation 10 was 0.788. When the data in Figure 2 is linearly regressed then a correlation coefficient of 0.775 is obtained. This small difference in correlation coefficent indicates that the theoretical approach for predicting poisoning power presented in this paper is almost as good as the best empirical expression. Therefore, the adsorption equilibrium constant is a suitable collective parameter for describing a compounds poisoning potency and as defined here the adsorption constant subtly combines size/heaviness, molecular structure and proton affinity. o~

8

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

. m

i

:. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

~

t~o 9 o3

...................................... ,iiiii i. . . i". .i .

.

.

.

.

.

.

.

.

.

.

'~" 0.01 -5

-

~ 0.001

'

"o.oool 4- .................................................................................

~-4

1E-05 0

1

2

3 log Kp

4

5

Figure 1 log Kp vs log(adsorbed conc.)

6

i

~

0.3

,

~ 0.4

..,

~

.,

~

0.5 0.6 In (1/(1-X))

,

~ 0.7

.i 0.8

Figure 2 ln(1/(1-X)) vs log (adsorbed mole frac.)

294 5. CONCLUSIONS Previously, for 2-methyl pentane cracking we have used Eley-Rideal kinetic rate expressions to describe the inhibition and poisoning effects of species in the feed, as well as intermediate and product species. In order to utilise such kinetic expressions values for the adsorption equilibrium constants are required. A method for estimating adsorption equilibrium constants has been proposed that uses an integrated form of van't Hoff equation. The heats of adsorption have been calculated using proton affinities and heats of condensation. The entropy of adsorption has been calculated using the Sackur-Tetrode expression. In this paper, the methodolog3' has been applied so as to describe the inhibiting and poisoning effect of more than twenty five nitrogen and aromatic compounds typically found in FCC feeds. Catalyst surface coverages have been determined by performing a flash calculation for a FCC feed containing a specified amount of poisoning species over a given amount of catalyst surface. The predicted surface coverages correlate well experimentally measured conversions. Thus, it can be concluded that the adsorption equilibrium constant is a suitable collective parameter for describing a compounds poisoning potency. The adsorption constant subtly combines size/heaviness, molecular structure and proton affinity. The implications of the work presented in this paper, are that now it should be possible to move away from lumped component and gas phase power law kinetic rate expressions for FCC to a situation where the major components are accounted for individually. In future work when the reactivity of the adsobed species has been more clearly quantified, a fundamental description of which species will poison an FCC catalyst as well as which are most likely to act as coke precursors should be possible. Thus, prediction of FCC riser and stripper kinetics without excess parameter retuning will be conceivable.

6. REFERENCES

1. S.Kumar, A.Chadha, R.Gupta and R.Sharma, Ind. Eng. Chem. Res., 34 11 (1995) 3737. 2. V.W. Weekman Jr, Lumps, Models and Kinetics in Practice, AIChE, New York, 1979. 3. B. Sowerby, S.J. Becker, and L.J. Belcher, J. Catal., 161 (1996) 377. 4. G.F.Froment and K.B.Bischoff, Chemical Reactor Analysis and Design, Wiley, 1990. 5. D.J. Parillo, R.J. Gorte and W.E. Fameth, J. Am. Chem. Soc., 115 (1993) 12441. 6. M. Maatman, R.M. Lago and C.D. Prater, Adv. in Cat., 9 (1957) 531. 7. W.J. Moore, Physical Chemistry, Longmans, 4th Edition, 1962. 8. Y.Zhao, G.R.Bamwenda, W.A.Groten and B.W. Wojciechowski, J. Catal., 140 (1993) 243. 9. J. Phys. Chem. Ref. Data, 17 Supplement 1 (1988). 10. Y. Zhao, G.R. Bamwenda and B.W. Wojciechowski, J. Catal., 142 (1993) 465. 11. C. Fu and M.A. Schaffer, Ind. Eng. Chem. Prod. Res. Dev., 24 (1985) 68. 12. T.C. Ho, A.R. Katritzky and S.J. Cato, Ind. Eng. Chem. Res., 31 (1992) 1589. 13. H.S. Fogler, Elements of Chemical Reaction Engineering, (1992) 263.

91997 Elsevier Science B.V. All rights reserved. Dynamics of Surfaces and Reaction Kinetics in Heterogeneous Catalysis G.F. Froment and K.C. Waugh, editors

295

Model discrimination for reactions with stop-effect S. Golay, O. Wolfrath, R. Doepper and A. Renken* Institute of Chemical Engineering, Swiss Federal Institute of Technology CH- 1015 Lausanne (Switzerland)

The stop-effect, a drastic increase of the reaction rate when the feed concentration of a reactant is switched to zero, was studied for the dehydration of ethanol to ethylene on 7alumina at 180 and 200~ Two basic models exist in the literature to describe this phenomenon. They were discriminated on the basis of transient and periodic experiments, coupled with FTIR data of the adsorbed species. The model that best describes these measurements postulates the adsorption of ethanol on two different sites, $1 and $2, with a free $2 site being necessary for ethylene formation.

1. I N T R O D U C T I O N The stop-effect, a drastic increase of the reaction rate when the feed concentration of a reactant is switched to zero, is observed for the catalytic dehydration of alcohols or the deamination of amines on aluminas, zeolites or more generally, on amphoteric metal oxide catalysts in the temperature range of 130 to 340~ [1, 2]. It can be described by different models, which make the hypothesis of different surface intermediates. Two basic models were discussed by Thullie and Renken [3]: Model 1: adsorption of the reactant on two different sites S l and 52, with a consecutive reaction involving an adsorbed intermediate (on $1) and a free site, $2. -

kl k-1

A +S I ~,

x

" , AS~ k,.k, A + S 2< " -,,->AS 2 AS I q-S 2

k3 " B + C + S ~ - I - S 2

(1) (2) (3)

The reactant A is only weakly adsorbed on the site $2 and desorbs rapidly when the feed is switched off, increasing the number of free sites $2 available for the reaction. Model 2: chemisorption of the reactant, with a consecutive adsorption of the reactant on that first adsorbed layer. -

A+S<

k,,k., >AS

*Corresponding author

(4)

296 A + AS < k,., k_2 >ASA

(5)

AS

(6)

k~ ) B + C + S

Again, the second layer is only weakly adsorbed and by switching off the reactant feed the concentration of the chemisorbed complex AS necessary for the reaction is increased. Both models describe the transient stop-effect behaviour and can be distinguished through periodic experimentation. The former is very similar to the reaction mechanism claimed by Koubek et al.[ 1], who associates S1 with an acid site on which the alcohol or the amine is strongly adsorbed and $2 with a basic site on which the reactant is more weekly adsorbed. Koubek takes advantage on the stop-effect to measure the acidic and basic character of aluminas. In this work the stop-effect was measured for the dehydration of ethanol on a y-alumina catalyst. The surface species were monitored using in-situ infrared spectrometry (FTIR). Periodic experiments are compared with transients coupled with FTIR data for the discrimination between the two models.

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

2.1 Catalyst preparation and characterisation y-alumina (Engelhard De Meern B. V., De Meern, The Netherlands) was used in this study. A self-supporting wafer, 13 mm in diameter, was prepared by pressing alumina powder for the infrared transmission cell (7.108 N / m2). Other wafers were pressed and then grinded to particles of 0.35-0.16 mm diameter for the fixed-bed reactor.

2.2 Experimental apparatus The schematic diagram of the experimental set-up is shown in Figure 1. It consists of two separate feed sections converging to a 4-way valve, a fixed-bed reactor, an infrared cell and a gas chromatograph. The carrier gas is argon (Ar> 99.99%, Carbagas, Lausanne, Switzerland) and the ethanol (>99.8%, No. 02860, Fluka Chemie AG, Buchs, Switzerland) feed is provided by a temperature-controlled bubble column fed by Ar. The gas flows were controlled with mass flow controllers (Bronkhorst High-Tech B. V., Ruurlo, The Netherlands). The reaction was carried out at a pressure of 140 kPa in a quartz fixed-bed reactor (internal diameter" l0 mm, length: 200 mm) followed by a stainless steel IR transmission cell (internal volume: 3.8 cm 3, windows: CaF2). The fixed-bed was heated with an electric tape, and the cell with two heating cartridges. The temperatures of the two reaction zones were measured with K-type thermocouples (Philips AG, Dietikon, Switzerland). It was verified that the temperature difference between two opposite axial points in the fixed-bed never exceeded 1~ and that the pressure drop between the inlet of the reactor and the outlet of the cell was less than 2 kPa. Infrared spectra were collected using a Perkin-Elmer Spectrum 2000 spectrometer (PerkinElmer, Rotkreuz, Switzerland) with liquid nitrogen-cooled MCT detector. All spectra were recorded with a resolution of 4 cm -~ (25 scans co-added). The gas phase was analysed with a

297 gas chromatograph (GC) (Hewlett-Packard, Urdorf, Switzerland) equipped with a capillary column (HP-1, 5m x 0.53mm) and a flame ionisation detector.

hood

feed hood

PI

T h

feed

Figure 1. Schematic diagram of the experimental set-up.

2.3 Procedure The stop-effect was measured at 180 and 200~ with an ethanol initial concentration of 0.26 mol / m 3, at total flow rates of 100, 200 and 400 ml (NTP) / rain. Before each experiment the catalyst was pre-treated in-situ at 415~ during 1 hour under inert gas flow, in order to remove the carboneous species adsorbed at the surface and to activate it. The transient experiments were carried out according to the following sequence: 108 minutes ethanol / 830 minutes Ar. An infrared background spectrum was measured before each experiment, after the pre-treatment. A wafer of 27 mg was used in the cell, and the fixed-bed was filled with 483 mg of catalyst.

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

RESULTS

AND MODELLING

3.1 In-situ IR spectrum An in-situ IR spectrum of the catalyst surface at steady-state is given Figure 2. It is very similar to that obtained by DeCanio et al. [4]. Bands attribution is based on previous works dealing with alcohols adsorption on metal oxide catalysts [4-6]. Two types of strongly adsorbed species can be pointed out: surface-bound ethoxides and surface acetates groups. The following IR bands are assigned to the ethoxide species: C-H stretching vibrations of the ethyl group at 2970, 2930 and 2870 cm -1, C-H2 and C-H deformations of the same group at 1450 and 1390 cm -1, C-O stretching vibrations at 1170, 1115 and 1070 cm ~ The bands at

298 1560 and 1480 cm -~ are typical for the acetate species. Their peculiar appearance is due to their presence in the background spectrum, the acetates being not removed during the catalyst pre-treatment. Negative bands at 3760 and 3725 cm -I are due to the perturbation of surface hydroxyl groups on ethanol adsorption. The large shoulder at 3560, 3510 cm -1 is assigned to the O-H stretching vibration of some interacting O-H groups.

2970 - - ~

0.6

1070"I 1115~'~~

2930

3560 1450 "7'

0.4

t" O r ..o <

0.2

0.0

1560

25

1480 3760

-0.2 4000

. 3000

' 2000

' 1500

1000

Wavenumber [cm -1] Figure 2. In-situ infrared spectrum of the catalyst surface under steady-state conditions. (TR = 200~ CA,O= 0.252 mol / m 3, Q = 394 ml (NTP) / min). The surface ethoxides are supposed to be reaction intermediates. Leaute and Dalla Lana [7] have shown that the intensity of the 2970 cm l band is related to their surface concentration and is not influenced by the C-H stretching vibration of the acetate species. Therefore, in this work, the ethoxides concentration will be determined by the height of the 2970 cm -1 absorbance band, with a baseline correction between 3100 and 2600 cm 1.

3.2 Transients experiments In Figure 3, the experimental transient response of the reaction products after a stop in the ethanol supply is shown. The ethylene concentration increases immediately, reaching a maximum after a few minutes, and then decreases slowly. Diethylether shows a completely different behaviour, its concentration decreasing quickly to zero after the stop, whatever are the experimental conditions. Theses observations are in complete agreement with those made by Koubek [ 1]. The position and value of the ethylene concentration maximum are strongly temperature dependent. At 180~ the maximum ethylene concentration is reached eight minutes after the stop and is generally four times higher than the steady-state value, depending on the initial ethanol concentration and the residence time. The ethylene concentration decrease shows two distinct parts, an autocatalytical effect taking place after a regular diminution. The occurrence of this second effect depends on the type of "f-alumina used, and is

299 not described by the models. Hence only the first part of the transient responses will be considered for the modelling.

ETHANOL

~'~ 0.12

-:

E

"

INERT

=:-,

= 0.012 m

0

E ,,_._,

::3-

9 Ethylene

0.10

tO

o

0 . 0 1 0 '
Diethylether

. m

"~ 0.08 tO O

~

0.008 o 0

%

(b (D

t- 0.06

0.006

O x..

,~-,-. O

...c: 0.04 m

(D >,, t-

"6 0.02

. m

E)

0.00

0.004 ~

f

3

___.L~_

0

i

0.002

;o ~

i~ 0

:-<<<'~:<':':~:::x<<<~-<':'<<' :+:':<<<+:<<< ::'i z<(o.<,_-(~<:,u 150

300

.......

450

Time [rain]

600

0.000

Figure 3. Typical response of reaction products after switching the ethanol concentration to zero (stop). (TR = 180~ CA,0= 0.265 mol / m 3, Q = 207 ml (NTP) / min).

3.3 Stop-effect modelling RTD experiments showed that the fixed-bed almost behaves like a plug-flow reactor and the infrared cell like a continuous stirred tmak reactor. This fixed-bed is described by the tanks-inseries model, using 9 tanks for the catalyst compartment. The two kinetic models (Equations 1-6) are able to describe the stop-effect experiments at 180 and 200~ and the considerations made in this work are valid for both temperatures. However, for the sake of clarity, only model discrimination at 180~ will be presented here. In the experimental conditions used here, both models can be simplified: the first adsorption step is considered as irreversible, and instantaneous equilibrium is assumed for the second one. With these hypothesis the total number of kinetic parameters is reduced from five (kl, k-l, k2, k-2 and k3) to three (kl, K2 and k3), and the models can be expressed as follows: Model 1" A +S 1

k, )AS l

rl = kiCA01

Kz ) AS 2

A dr S 2 (

02 =

(7)

KzCA

(8)

K2C A + 1

AS~ + S 2 01 ~---

C

k~ >B + C + S~ + S 2 AS1

CAS, + Cs,

r 3 = k301 ( 1 - 02)

GAs ' -

0a

Z~,~

(9)

CAS., =

"

CAS, + (2s~

CAS 2 -

(10)

Z~,:

300 Model 2: A+S

k, ,AS

A + AS < AS

K.

" ; ASA

k3 ) B + C + S C AS

01=

C ASA "1"C AS -[- C s

-

C AS Z2

r l = k l C A01

(11)

0 2 = K2C A01

(12)

r 3 =k301

(13) C ASA

0,= -

C ASA + C AS + C s

-

C ASA Z,

(14)

The intensity of the IR signal is supposed to be proportional to the number of ethoxide species adsorbed on the wafer and is calculated by the sum of all adsorbed species, using equations (15) and (16) for model 1 and model 2, respectively" I~ = cz 9[(0, Z~,, ) + (02 Z,,2)]- m~a~,c~,

(15)

ItR = o~ 9[(0,Z2 ) + 2(02Z2)] mcat.cel

(16)

Diethylether formation must be taken into account. For simplicity, a quasi-homogeneous reaction term is added: A

k, '>D

r4 = k4CA,,

(17)

The mass balance differential equations for the gas phase and the surface species were solved simultaneously as a function of time and discrete position, using a variable step integration algorithm (Gear) [8]. The total number of actives sites (ZI,~ in model 1 and Z2 in model 2) is determined by dividing the number of ethylene moles produced after the stop by the total catalyst mass in the system. The parameters k4 and n4 were determined independently from stationary experiments. The three remaining kinetic parameters (kl, K2, k3) are then simultaneously optimised (using Nelder-Mead algorithm [8]) to fit the ethylene concentration during a complete experiment (steady-state and transient (stop)). In the last step the proportionality coefficient for In~, or, is adjusted to describe the stationary infrared signal. It was shown by Thullie and Renken that, assuming Zl,l=Zl,2, model 1 with the same set of kinetic parameters as model 2 gives the same transient response. Hence optimisation is first performed with model 2 and compared with the experimental results, Figures 4a and 4b. A good description of the ethylene transient response is achieved, except in the first minutes following the stop: the calculated maximum is lower than the experimental, and is reached 12 minutes after. At the surface, two distinct kinetics are observed: a rapid decrease just after the stop due to desorption of the inhibiting species, followed by a slow diminution due to the consumption of the reacting species to form ethylene. The observed decrease of the ethoxy species is lower than the predicted, indicating a smaller amount of Z~,2 compared to Za,~. Therefore a different number of sites can be assumed for model 1, which is made by introducing a new parameter, q0. It is defined as the ratio between sites 2 and sites 1 (q~ = Z~,2/Zl,1) and is adjusted to describe the initial diminution of the infrared signal. A new optimisation of kl, K2, k3 is performed and it can be seen in Figures 4a and 4b that, by adding this new parameter to model 1, a better fit is obtained for the gas phase and that the surface concentration is now well described.

301

0.007 ]_

ETHANOL

9

E 0.006

INERT

~

0

~

:O

E ,..- 0.005-

_1

Model 1

1

O

E ~D O tO ro E c.-

0.0040.003-

0.002

uJ 0.001 0.000

0.6 "7"

--r-

6

=

c"

o .i.-,

(1)

0.5

0.3

o

0.2

= 0

..Q <

,~.

E THA NOL

'

I

!

100

'

I

150

INER T

,1~...41.

i

"

50 Time [min]

'

200

D.

Model 1 ~i-~i.... Model 2

~

0.4

~-

....,

0

" ' ' ' ' ' ' - . . . . . . . . . . ~

0.1 0.0

0

50

1O0

150

200

Time [min] Figure 4. Experimental and simulated responses of the gas phase and the surface for a stopeffect experiment. (TR = 180~ Ca.0 = 0.265 mol / m 3, Q = 207 ml (NTP) / min). The optimum parameter values are: - Model 1:k1=2.68.10 2 m3/kg.s, K2=10.1 m3/mol, k3=4.53-10 s mol/kg-s, Z~,~=0.80 mol/kg, ~=0.45, ct,.=19000 mol l, k4=1.98-10 2 m8 l/moll7.kg.s, n4=2.7 - Model 2:k1=2.68-10 -2 m3/kg.s, K2=9.60 m3/mol, k3=4.51.10 -5 mol/kg-s, Z],]=0.80 mol/kg, 15000 mo1-1, k4=1.98.10 -2 mSl/moll7.kg-s, n4=2.7

302 3.4

Periodic

operation

A periodic experiment was made under the same experimental conditions as the transient used for model discrimination (TR = 180~ CA 9 = 0.265 mol / m 3, Q = 200 ml (NTP) / min), with a period of forty minutes and a cycle split of 0.5.

INERT

0.006

ETHANOL

INERT ETHANOL

INERT

ETHANOL

,

,

e3

E

_~ 0 . 0 0 5 O

!/

E

) .."

t.0.004 O .,._,

'- 0 . 0 0 3

? )

O r O

o 0.002

(D c" ~D

)

x:: 0.001

Model 1

U..I

Model 2

0.000 15

35

55

,

I 75

Time

INERT

0.6 "r"9

-

,.--

0.4

.9.o "~ .4..,

0.3

-

~ 0

"

INERT

ETHANOL

A

"/a.~..

:

-.A

~ 9

.

A..

:

".

:

"

0.2

..Q x_ O

-Q <

135

[rain]

INERT ETHANOL

'

i 115

0.5

d (.-

ETHANOL

95

o

......

0.1 I

0.0 i 15

35

Model 1

......... Model 2

55

75 Time

95

115

135

[min]

Figure 5. Comparison of the experimental and the predicted behaviour of the gas phase and the surface under periodic operation. (TR = 180~ CA,0 = 0.265 mol / m 3, Q = 207 ml (NTP) / min, period = 40 minutes, split = 0.5).

303 The predicted results of models 1 and 2, calculated with optimum parameters determined on transient experiment, are compared to experimental results for the gas phase and the surface in Figures 5a and 5b, respectively. It must be pointed out that the periods invariance is not obtained simultaneously for ethylene and the surface ethoxides. The response of ethylene becomes reproducible after a different first cycle (not shown here), whereas the infrared signal is not yet stable. Its value still increases during the following periods, due to a parallel slow adsorption of ethanol under a non-reactive form. This adsorption was also observed in the transient experiment, and consequently was included in the determination of cz. Therefore the comparison between models and experiment can only be made on the last period in the case of the adsorbed species. Figures 5a and 5b obviously show that model 1 gives a much better description than model 2. For the gas phase, it describes the ethylene maximum reached under inert gas, whereas model 2 still increases at the end of the inert cycle. For the surface, the last period is also well predicted by model 1, and not at all by model 2.

4. C O N C L U S I O N S By applying transient experiments to the dehydration of ethanol on y-alumina, it has been shown that model 1 and 2 can describe the experimentally observed increase of the reaction rate. The introduction of a new parameter in model 1 % the ratio of the number of sites of type two by the number of sites of type one, enables a better description of the drastic increase of ethylene directly after stopping the ethanol feed. The parallel behaviour of surface intermediates, a sharp decrease followed by a slow consumption, is quantitatively depicted by model 1, whereas model 2 gives only a qualitative description. Periodic experimentation confirms that model 1 is able to describe the gas phase and the surface more adequately than model 2.

5. N O M E N C L A T U R E A B C Ci D I1R k K~ mcat.cei

n4 NTP Q r

Ta Z

Ethanol Ethylene Water Concentration of species i, mol/m 3 for the gas phase or mol/kg for the surface Diethylether Infrared absorbance of adsorbed species Rate constant, various units Equilibrium constant of adsorption step 2, m 3 / mol Wafer mass, kg Apparent reaction order of Equation (13) Normal conditions of temperature and pressure (0~ 1.013-10 5 Pa) Total volumetric flow rate, ml (NTP) / rain Reaction rate, mol / kg-s Reaction temperature Sites concentration, mol / kg

304

Greek letters cx 0 q0

Proportionality coefficient for the infrared, mo1-1 Surface coverage, Ratio of concentration of sites $2 by concentration of sites

Sl

in model 1, -

Subscripts 0

inlet

REFERENCES 1. J. Koubek, J. Pasek and V. Ruzicka, In New Horizons in Catalysis, Elsevier-Kodansha, Amsterdam-Tokyo, 853-862, 1980. 2. J. Koubek, J. Pasek and R. V., In Catalyst deactivation, Elsevier, Amsterdam, 251-260, 1980. 3. J. Thullie and A. Renken, Chem. Eng. Sci., 48 (1993) 3921-3925. 4. E. C. DeCanio, V. P. Nero and J. W. Bruno, J. Catal., 135 (1992) 444-457. 5. G. A. M. Hussein, N. Sheppard, M. I. Zaki and R. B. Fahim, J. Chem. Soc. Faraday Trans., 87 (1991) 2661-2668. 6. H. Kn6zinger and B. Sttibner, J. Phys. Chem., 82 (1978) 1526-1532. 7. R. Leaute and I. G. Dalla Lana, J. Catal., 60 (1979) 460-471. 8. Simusolv, Modeling and simulation software, Reference Guide, Dow Chemical Company, Midland (1990).

91997Elsevier Science B.V. All rights reserved. Dynamics of Surfaces and Reaction Kinetics in Heterogeneous Catalysis G.F. Froment and K.C. Waugh, editors

305

M e t h a n o l Oxidation Over Supported V a n a d i u m Oxide Catalysts: N e w F u n d a m e n t a l Insights About Oxidation Reactions O v e r M e t a l O x i d e Catalysts from Transient and Steady State Kinetics Israel E. Wachs", Goutam Deo"*, Michael V. Juskelis ~, Bert M. Weckhuysen ~ "Zettlemoyer Center for Surface Studies and Department of Chemical Engineering, Lehigh University, Bethlehem, PA, 18015, USA ~W. R. Grace, Washington Research Center, Columbia, MD 21004, USA cCentrum voor Oppervlaktechemie en Katalyse, K. U. Leuven, Kardinaal Mercierlaan 92, B-3001 Heverlee, Belgium

1. INTRODUCTION Metal oxide catalysts are extensively employed in the chemical, petroleum and pollution control industries as oxidation catalysts (e.g., oxidation of methanol to formaldehyde, oxidation of o-xylene to phthalic anhydride, ammoxidation of propylene/propane to acrylonitrile, selective oxidation of H2S to elemental sulfur (SuperClaus) or SO2/SO3, selective catalytic reduction (SCR) of NOx with NI-I3, catalytic combustion of VOCs, etc.). A special class of metal oxide catalysts consists of supported metal oxide catalysts, where an active phase (e.g., vanadium oxide) is deposited on a high surface area oxide support (e.g., alumina, titania, zirconia, niobia, ceria, etc.). Supported metal oxide catalysts provide several advantages over bulk mixed metal oxide catalysts for fundamental studies since (1) the number of surface active sites can be controlled because the active metal oxide is 100% dispersed on the oxide support below monolayer coverage, (2) the oxide support ligands can be varied in order to monitor the influence of the bridging V-O-Support bond and (3) the surface active sites can be directly monitored with in situ spectroscopies (e.g., Raman, IR, DRS, etc.) during oxidation reactions. In the present study, the advantages provided by the supported metal oxide catalysts are combined with transient and steady state kinetic studies in order to obtain additional fundamental insights into oxidation reactions over metal oxide catalysts.

*Supported by the National Science Foundation (grant no. CTS-9417981). tCurrent address: Departmentof ChemicalEngineering,IndianInstituteof Technology, Kanpur 208 016, India. :IcSupportedby the National Fund of ScientificResearch (N.F.W.O.)of Belgiumwith a research assistant grant and a travel grant to visit the Zettlemoyer Center for Surface Studies at Lehigh University.

306 2. EXPERIMENTAL The supported vanadia catalysts were synthesized via the incipient-wetness impregnation method. A vanadium isopropoxide (Alfa, 95-99% purity) methanol solution was used to impregnate the different oxide supports: A1203 (I-Iarshaw, 225 m2/g), TiO2 (Degussa P-25, 55 m2/g), ZrO2 (Degussa, 39 m2/g), NthO5 (Niobium Products Company, 55 m2/g) and CeO2 (Engelhard, 36m2/g). Complete details of the preparation procedure can be found in the literature [1]. Information about the various characterization experiments can also be found in the literature: Raman spectroscopy [2], Infrared (IR) Spectroscopy [3], UV-VIS Diffuse Reflectance Spectroscopy (DRS) [4] and Temperature Programmed Reaction Spectroscopy (TPRS) [5]. The transient TPRS experiments were performed by initially outgassing 100 mg of the catalyst in 5 % O2/He for I hour at 300 C and subsequently cooling the sample to 100 C. The sample was flushed with helium for an additional hour and methanol was adsorbed at the same temperature with helium as the carrier gas. The adsorption temperature of 100 C was selected because it minimized the amount of physisorbed methanol on the catalysts. The sample was flushed with flowing helium for an additional hour to remove any residual physisorbed methanol. The TPRS experiment was then initiated with a heating rate of 10 C/rain in a helium flow rate of 100 cc/min, and product analysis was performed with an on-line mass spectrometer (Hiden Analytical Ltd.) (see reference 5 for additional experimental details). The methanol oxidation studies were performed in a fixed bed reactor at 230 C. The reaction temperatures were only varied to determine the activation energy (additional details about the reactor system can be found in reference 1).

3. RESULTS 3.1 Molecular structures of the calcined surface vanadia species

The molecular structures of the two-dimensional surface vanadia species, after calcination and under in situ conditions, have been determined with the aid of Raman, IR, DRS, solid state 5~V NMR, EXAFS/XANF.S and oxygen isotope exchange experiments [611]. These molecular characterization experiments have revealed that the surface vanadia species possess the foUowing characteristics: (1) the calcined surface vanadia species are present in the V(+5) oxidation state, (2) the surface vanadia species possess only one terminal V - O bond (they are mono-oxo), (3) the surface vanadia species are predominantly present as VO4 species (there is a possibility that a minor amount of surface VO5 or VO6 species may also be present at monolayer coverages), (4) both isolated and polymerized surface VO4 species are present, (5) the ratio of polymerized to isolated surface vanadia species increases with surface coverage and (6) the molecular structures of the surface vanadia species on the different oxide supports (alumina, titania zirconia, niobia and ceria) are essentially identical (even the ratio of the polymerized to the isolated surface vanadia species is the same at comparable surface coverages). The only structural differences among these supported vanadia catalysts are the very minor differences in the terminal V =O bond lengths (corresponding to less than 0.01 Angstrom [12]) and the different oxide support ligands (A1, Ti, Zr, Nb and Ce). The monolayer surface coverages of the vanadia ovedayers, experimentally determined by Raman spectroscopy, and the Raman vibrations of the terminal V =O bonds at monolayer coverages are presented in Table 1.

307 Table 1 Monolayer surface coverages and vibrations of the terminal V = O bonds, at monolayer coverage, for the supported vanadia catalysts

Catalyst

25 % V2Os/A1203 7% V2Os/Nb205 6% V2Os/TiO2 4 % V205/Zl'O 2 3 % V205/Ce02

Monolayer Coverage (V atoms/nm 2)

7.3 8.4 7.9 6.8 5.7

V = O (cm -~)

1026 1033 1030 1030 1028

3.2 Molecular structures of the surface vanadia species and adsorbed intermediates during methanol oxidation During methanol oxidation, the surface vanadia species becomes partially reduced by the reaction environment. In situ Raman studies during methanol oxidation revealed that the surface V(+5) species retained its molecular structure, but that there was about a 40-60% reduction of the Raman signal for the surface V(+5) species [13]. The actual extent of reduction was probably less than 40-60% since the catalysts became slightly darker upon reduction and decreased the amount of scattered laser light to the detector. However, the extent of reduction, within experimental error, was independent of (1) the surface coverage of the surface vanadia species, (2) the specific oxide support, and (3) the reaction temperature (230-350 C). This suggests that the fraction of participating surface vanadia sites, or active surface sites, is relatively constant for all the supported vanadia catalysts under the chosen experimental conditions. In situ DRS studies revealed that the reducext surface vanadia species was predominantly present as surface V(+4) species [14]. In situ IR spectroscopy studies also revealed the presence of surface methoxy, CH30, and surface formate, HCOO, species on the catalyst surface during methanol oxidation over a 5% V2Os/TiO 2 catalyst [15]. The 5% V2Os/TiO2 catalyst corresponds to approximately monolayer surface coverage of the surface vanadia overlayer on the titania support, see Table 1, and minimizes the adsorption of methanol and its reaction products on exposed surface titania sites that are present below monolayer coverage or crystaUine V205 particles that are present above monolayer coverage. 3.3 Steady state kinetics of methanol oxidation over the supported vanadia catalysts The turnover frequencies (TOF), defined as the number of methanol molecules converted to formaldehyde per surface vanadia site per second, are presented for the different supported vanadia catalysts, at monolayer coverages, in Table 2. There is a dramatic variation in the TOFs with the specific oxide support and the variation spans approximately three orders of magnitude at monolayer coverages (the same surface density of surface vanadia species). The TOFs were also relatively independent of surface vanadia coverage [1], which indicates that the reaction rate is first order with respect to the surface vanadia

308 Table 2 Turnover frequencies (TOFs) and activation energies for methanol oxidation over supported vanadia catalysts at monolayer coverages of surface vanadia species

Catalyst

25 % V205/A1203 7% V205/Nb205 6% V205/TiO2 4 ~ V205/ZrO 2 3% V205/C~O 2

TOE (cm-1)

6.8 4.0 1.1 1.7 1.0

x l(r 2 X l(Y1 x 10+~ x 10+~ x 10+1

Activation Energy (Kcal/mol)

20 17 22 18 20

sites. This suggests that methanol oxidation over the supported vanadia catalysts is a unimolecular reaction requiting only one surface vanadia site (the active site). The formaldehyde Selectivities were 90-99 % for the titania, zirconia, niobia and ceria supported vanadia catalysts. For the alumina supported vanadia catalyst, however, the formaldehyde selectivity was only about 50% and dimethyl ether was the major unselective product. The high dimethyl ether formation for the V2Os/A1203 catalyst reflects the presence of surface acid sites for this catalyst. The activation energies for methanol oxidation over the supported vanadia catalysts at monolayer coverage are approximately constant at 19.5 Kcal/mol, with a standard deviation of 2.5 Kcal/mol (see Table 2). Thus, the dramatic variation in the TOFs for methanol oxidation over the different supported vanadia catalysts is primarily related to differences in the preexponential factors of the kinetic expression. 3.4 Transient kinetics of methanol oxidation over the supported vanadia catalysts The major products formed during the TPRS experiments were H2CO (selectivity varied from about 25-55 %), CO (accounting for almost all the remaining carbon containing products), and H20 with minor amounts of CH3OH (trace amounts), CO 2 (typically 2-4 %), and CH3OCH3 (trace amounts only observed for the alumina supported vanadia catalyst). The much lower formaldehyde selectivity during the TPRS experiments relative to the steady state methanol oxidation experiments is a consequence of the absence of significant amounts of gas phase methanol and moisture, which efficiently retard the readsorption and oxidation of formaldehyde via competitive adsorption [16-18]. Formaldehyde is the major product of the decomposition of the surface methoxy intermediate [19] and CO is the major product from the decomposition of the surface formate intermediate [20]. The small mounts of CO2 are due to the further oxidation of the CO product in the catalyst bed. The trace amounts of CH3OH could originate from either some residual physisorbed methanol or from the hydrogenation of the surface methoxy intermediate [19]. The amounts of methanol adsorbed per surface vanadia species on the different supported vanadia catalysts, at approximately monolayer surface vanadia coverage and 100 C, could also be determined from integration of the calibrated product mass spectrometer signals and are presented in Table 3. Thus,

309 Table 3 Number of methanol molecules adsorbed per surface vanadia species at monolayer surface coverages (adsorption temperature of 100 C)

Catalyst

Number of adsorbed methanol molecules/surface V atom

25 % V2Os/AI205 5% V2Os/Nb205 3 % V2Os/ZrO2 5 % V2Os/TiO2 3 % V2Os/CeO2

0.44 0.65 0.57 0.37 0.50

the saturated amount of methanol molecules adsorbed on the different supported vanadia catalysts at 100 C is very similar, 0.51 with a standard deviation of 0.14, and corresponds to approximately one adsorbed methanol molecule per two surface vanadia species. The TPRS experiments also provided additional insights into the kinetics of methanol oxidation over the supported vanadia catalysts. The TPRS experiments only provide kinetics about the surface reaction steps since the adsorption events precede the initiation of the transient temperature ramp. The TPRS peak temperatures for the production of H~CO and CO from the different vanadia catalysts are presented in Table 4.

Table 4 Peak temperatures for the production of H2CO and CO from CH3OH oxidation during the TPRS studies over the supported vanadia catalysts

Catalyst

25 % V2Os/A1203 5 % V2OJNbuOs 3 % V2Os/ZrO2 5 % V2Os/TiO2 3 % V2Os/CeO2

H_2CO(~

C,O(~

210 190 200 220 200

275 285 -270 265

Recall that the formation of H2CO and CO result from the decomposition of the surface methoxy, CH30, and the surface formate, HCOO, intermediates, respectively. The surface formate is produced by re,adsorption of formaldehyde product on the surface vanadia sites [20]. The TPRS spectra reveal that the decomposition of the surface methoxy intermediates

310 (Tp = 190-220 C), as well as the surface formate intermediates (Tv = 265-285 C), possess essentially the same kinetics on all the supported vanadia catalyst. The desorption of H20 exhibited a very broad peak with a maximum at about 300 C for the supported vanadia catalysts. The slight variations among the different TPRS runs are related to the use of several parallel reactor systems to expedite the experiments, and are not due to different decomposition kinetics. Identical peak temperatures were obtained when two different catalysts were studied in the same reactor system. Thus, the TPRS experiments demonstrate that the dramatically different TOFs measured during the steady state methanol oxidation to formaldehyde over the supported vanadia catalyst are not related to kinetic differences in the rate determining surface reaction step, the decomposition of the surface methoxy species [19,20].

4. DISCUSSION

4.1 Molecular structure-reactivity relationships during methanol oxidation to formaldehyde over supported vanadia catalysts The methanol adsorption experiments, at the beginning of the TPRS experiments, revealed that a maximum of one methanol molecule adsorbs per two surface vanadia sites (at monolayer surface vanadia coverage and an adsorption temperature of 100 C). The maximum number of adsorbed methanol molecules at monolayer surface vanadia coverage can be doubled by lowering the adsorption temperature to room temperature, but the increase would be strictly due to the presence of physically adsorbed or weakly interacting methanol molecules rather than to dissociatively adsorbed methanol molecules in the form of surface methoxy species (the reactive surface intermediate) [16,19]. It may be possible that the number of adsorbed methoxy species per surface vanadia sites increases to unity at low surface vanadia coverages where lateral interactions are minimized, but such quantitative data are currently not available because of complications due to competitive adsorption on the exposed oxide support sites. For silica supported vanadia catalysts, where the surface vanadia species are isolated [6,7,21] and competitive adsorption on the oxide support is not a problem at modest surface vanadia coverages [2], the maximum number of surface methoxy per surface vanadia species is close to unity [22]. A careful examination of the methanol oxidation kinetic data reveals that the TOFs were always slightly higher at low surface vanadia coverages, which would be consistent with a higher number of adsorbed methoxy species per surface vanadia site. Thus, the maximum number of adsorbed methoxy per surface vanadia species may not be constant at all surface vanadia coverages and may vary from approximately unity at low surface vanadia coverages to approximately 0.5 at high surface vanadia coverages. It is also important to establish if methanol is directly coordinated to one or two surface vanadia sites, mono-dentate vs. bidentate, or if two surface vanadia sites are required because of lateral interactions among the surface methoxy species at monolayer surface vanadia coverage. Comparative IR studies of adsorbed methoxy on vanadia catalysts with known molecular structural reference compounds reveal that the adsorbed methoxy species is only coordinated to one surface vanadia ~ i e s [20]. This coordination is consistent with the almost constant methanol oxidation TOF as a function of surface vanadia coverage and the insensitivity of the methanol oxidation TOF to the presence of secondary surface metal

311 oxide additives in the surface vanadia overlayer [23]. Furthermore, the essentially identical kinetics for the oxidative dehydrogenation of the surface methoxy species to formaldehyde over the different supported vanadia catalysts during the TPRS experiments (see Table 4) suggests that this rate determining step exclusively occurs on the surface vanadia site without the direct involvement of the oxide support cations or oxygen anions. The molecular characterization studies demonstrated that the oxidized surface vanadia ~ i e s in the different supported vanadia catalysts possess essentially the same molecular structures: predominantly consisting of isolated and polymerized surface VO4 species with the same ratio of polymerized to isolated species at comparable surface vanadia coverages. The surface vanadia species even became reduced to comparable extents during methanol oxidation for all the supported vanadia catalysts. The terminal V = O bond lengths for the different supported vanadia catalysts were also essentially identical, see Table 1, and the minor variations in the V =O bonds did not correlate with the methanol oxidation TOFs, (see Tables 1 and 2). Thus, there are no significant molecular structural differences among the surface vanadia species on the different oxide supports to account for the dramatic variation in the TOFs during methanol oxidation over the supported vanadia catalysts. The only significant difference among the surface vanadia species on the different oxide supports was the oxide support ligands (e.g., A1, Ti, Zr, Nb and Ce). The electronegativity of the oxide support cation affects the electron density on the bridging V-OSupport oxygen: a lower cation electronegativity will result in a slightly higher electron density (more basic oxygen) and a higher cation electronegativity will result in a slightly lower electron density (less basic oxygen). Comparison of the Sanderson electronegativities of the oxide support cations [24] (Ce < Zr < Ti - Nb < A1) with the methanol oxidation TOFs (Ce > Zr > Ti > Nb > A1; see Table 2), reveals an inverse correlation: the lower the oxide support cation electronegativity the higher the methanol oxidation TOF. Therefore, the more basic the bridging V-O-Support oxygen the higher the methanol oxidation TOF. It is necessary to examine the mechanism and kinetics of the methanol oxidation reaction to better understand the relationship between the basic characteristics of the bridging V-OSupport oxygen and the methanol oxidation TOFs over the supported vanadia catalysts.

4.2 Reaction mechanism and kinetics of methanol oxidation to formaldehyde The kinetics of methanol oxidation over metal oxide catalysts were elegantly derived by Holstein and Machiels [16]. The kinetic analysis demonstrated that the dissociative adsorption of water must be included to obtain an accurate kinetic model. The reaction mechanism can be represented by three kinetic steps: equilibrated dissociative adsorption of methanol to a surface methoxy and surface hydroxyl (represented by K,), equilibrated dissociative adsorption of water to two surface hydroxyls (represented by K2), and the irreversible hydrogen abstraction of the surface methoxy intermediate to the formaldehyde product and a surface hydroxyl (the rate determining step, represented by ks). For the case of a fully oxidized surface, the following kinetic expression was derived:

r(~,~

= ~

9

~,

9

PcH, oH

(K2)m (pH2o)

(1)

312 In the above kinetic expression, the Arrhenius rate constant, k~, is modified by the two adsorption equilibrium constants of methanol and water during the steady state kinetic studies. During the transient TPRS experiments, however, only the Arrhenius rate constant, k3, is measured since there is no vapor phase methanol or water to be equilibrated. The similar TPRS results for the oxidation of the surface methoxy intermediate to formaldehyde over the different supported vanadia catalysts reveal that all the catalysts possess the same k3. Consequently, the dramatic differences in the steady state TOFs during methanol oxidation over the different supported vanadia catalysts must be associated with the methanol and water equilibrium adsorption constants. Both methanol and water can be viewed as weak acids that will donate a proton to a basic surface site, but methanol is more strongly adsorbed than water on oxide surfaces, KI > (K2)la [16]. Furthermore, the methanol oxidation TOFs were measured at methanol conversions below 20%, differential reaction conditions, where the partial pressures of methanol were much higher than the partial pressures of water [23]. Thus, the more active catalysts have a higher steady state surface concentration of adsorbed methoxy species, or a higher ratio of surface methoxy to surface hydroxyls, on the catalyst surface during methanol oxidation due to presence of more basic bridging V-O-Support oxygens (Ce > Zr > T i - Nb > A1). The above molecular structure-reactivity relationships provide new insights into the oxidation of methanol to formaldehyde over oxide catalysts. Many investigators have proposed that methanol adsorption occurs at the mono- or di-oxo M = O sites [25,26]. The molecular characterization studies demonstrated that di-oxo functionalities are not present in the supported vanadia catalysts (a similar scenario is also found for supported molybdena catalysts [27]). Furthermore, the mono-oxo functionalities are essentially identical in all the catalysts and, consequently, their characteristics do not correlate with the dramatic variation in TOFs. The current studies demonstrate that the methanol oxidation TOFs correlate with the basicity of the bridging V-O-Support bonds (similar results were also obtained for the analogous supported molybdena catalysts [27]). These new fundamental insights suggest that methanol adsorption does not occur at a M = O functionality, but that the dissociative adsorption step preferentially occurs at the bridging M-O-Support bond to form surface MOCH3 and Support-OH functionalities. The dissociative adsorption of methanol is facilitated in the presence of more basic bridging M-O-Support oxygens. It has also been proposed that methanol adsorption and its oxidation to formaldehyde occurs at coordinatively unsaturated sites, possessing four-fold coordination, rather than coordinatively saturated sites, possessing six-fold coordination [19]. Unfortunately, the surface vanadia species predominantly possess four-fold coordination which prevents this issue from being addressed with the current data. However, supported molylxlena catalysts possesses both four-fold and six-fold coordination and their TOFs for methanol oxidation have been measured [27]. It was found that, contrary to above hypothesis, the coordinatively saturated surface molylxlena species is approximately four times more active than the coordinatively unsaturated molybdena species for titania supported molybdena catalysts. Thus, methanol oxidation proceeds on both coordinatively saturated and coordinatively unsaturated sites at relatively comparable reaction rates (TOFs). The kinetic expression derived by Holstein and Machiels, equation 1, also qualitatively accounts for the rather constant activation energies among the different supported vanadia catalysts during methanol oxidation. The measured activation energy is an apparent activation energy that is a function of the true activation for the abstraction of

313 the hydrogen from the surface methoxy intermediate, the heat of adsorption of methanol and the heat of adsorption of water (see equation 1). However, the heats of methanol and water adsorption partially compensate each other because the methanol equilibrium adsorption term appears in the numerator and the water equilibrium adsorption term appears in the denominator. As a result, the apparent activation energies and the true activation energy are very similar for most catalysts during methanol oxidation (see Table 2). A consequence of this conclusion is that the variation in the TOFs primarily presides in the preexponential term of the kinetic expression of equation 1. Several recent theories have appeared in the literature to explain the dramatic variation in the methanol oxidation TOF over the different supported vanadia catalysts. From ab initio quantum mechanical calculations Weber has proposed that surface methoxy decomposition occurs at the bridging M-O-support bond, the surface methoxy methyl hydrogen initially being extracted as a hydride by the support cation, and that the entropy of the activated complex is affected by the density of ~ s s i b l e electronic states associated with the support cation [28]. However, the TPRS experiments demonstrate that the decomposition of the surface methoxy to product formaldehyde and a surface hydroxyl occurs with the same kinetics for highly active catalysts as well as modestly active catalysts (see Table 4). Thus, this model does not ~ u n t for the experimental observations and should be modified to examine the dissociative adsorption of methanol rather than the decomposition of the surface methoxy species. Steigman assigned the electronic transitions of the vanadia species in a silica matrix and concluded that the bridging V-O-Support bond would be the most active for oxidation reactions [29]. He proposed that the electron density on the bridging oxygen would be enhanced by replacing silica with less electronegative cations like titania that would further enhance the reactivity of the bridging V-O-Support bond. Thus, Steigman's model qualitatively accounts for the observed trends and needs to be examined more carefully on a quantitative basis. REFERENCES 1. G. Deo and I.E. Wachs, J. Catal., 146 (1994) 323. 2. J.-M. Jehng, H. Hu, X. Gao and I.E. Wachs, Catal. Today, 28 (1996) 335. 3. M.A. Vuurman, D.J. Stufkens, A. Oskam, G. Deo and I.E. Wachs, J. Chem. Soc., Faraday Trans., 92 (1996) 3259. 4. B.M. Weckhuysen and R. A. Schoonheydt, J. Phys. Chem., in press. 5. J.E. Swain, M. V. Juskelis, J.P. Slanga, J.G. Miller, M. Uberoi and N.D. Spencer, Appl. Catal. A: General, 139 (1996) 175. 6. I.E. Wachs, Catal. Today, 27 (1996)437. 7. I.E. Wachs and B.M. Weckhuysen, Appl. Catal., in press. 8. G.T. Went, L.-J. Leu, R.R. Rosin, and A.T. Bell, J. Catal., 134 (1992) 492. 9. G. Ramis, C. Cristinai, P. Forzatti and G. Busca, J. Catal., 124 (1990) 574. 10. H. Eckert and I.E. Wachs, J. Phys. Chem., 93 (1989) 6796. 11. T. Tanaka, H. Yamashita, R. Tsuchitani, T. Funabiki and S. Yoshida, J. Chem. Soc., Farady Trans. 1, 84 (1988) 2987. 12. F.D. Hardcastle and I.E. Wachs, J. Phys. Chem., 95 (1991) 5031. 13. G. Deo and I.E. Wachs, to be published. 14. B. Weckhuysen, I.E. Wachs and R.A. Schoonheydt, to be published.

314 15. 16. 17. 18. 19. 20. 21. 22. 23. 24. 25. 26. 27. 28. 29.

L. Burcham, J. Datka and I.E. Wachs, to be published. W.L. Holstein and C.J. Machiels, J. Catal., 162 (1996) 118. W.-H. Cheng, J. Catal., 158 (1996)477. D.S. Lafyatis, G. Creten and G.F. Froment., Appl. Catal. A: General, 120 (1994) 85. U. Chowdhry, A. Ferretti, L.E. Firment, C.J. Machiels, F. Ohuchi, A.W. Sleight and R.H. Staley, Appl. Surf. Sci., 19 (1984) 360. G. Busca, A.S. Elmi and P. Forzatti, J. Phys. Chem., 91 (1987) 5263. N. Das, H. Eckert, H. Hu, I.E. Wachs, J. Walzer and F. Feher, J. Phys. Chem., 97 (1993) 8240. X. Gao, M. Juskelis and I.E. Wachs, to be published. G. Deo and I.E. Wachs, J. Catal., 146 (1994) 335. R.T. Sanderson, Inorganic Chemistry, Van Nostrand-Reinhold, New York, 1967. N. Pernicone, F. Lazzerin, G. Liberti and G. Ianzaveechia, J. Catal., 14 (1969) 293. J.N. Allison and W.A. Goddard, J. Catal., 92 (1985) 127. H. Hu and I.E. Wachs, J. Phys. Chem., 99 (1995) 10911. R.S. Weber, J. Phys. Chem., 98 (1994) 2999. K. Tran, M.A. Hanning-Lee, A. Biswas, A.E. Stiegman and G.W. Scott, J. Am. Chem. Soc., 117 (1995) 2618.

91997 Elsevier Science B. V. All rights reserved. Dynamics of Surfaces and Reaction Kinetics in Heterogeneous Catalysis G.F. Froment and K.C. Waugh, editors

The

Effects

Of

Alkali

Promoters

On

315

The

Dynamics

Of

Hydrogen

C h e m i s o r p t i o n And Syngas Reaction Kinetics On Ru/SiO 2 Surfaces D.O. Uner l, N. Savargoankar 2'3 M. Pruski 3 , and T.S. King 2 1Chemical Engineering, Middle East Technical University, Ankara 06531 Turkiye 2 Chemical Engineering, Iowa State University, Ames IA 50011 3Institute of Physical Research and Technology, Ames Laboratory, Ames IA 50011

The dynamics of chemisorbed hydrogen on unpromoted and promoted Ru/SiO 2 catalysts was studied by means of single pulse and selective excitation 1H NMR spectroscopy. Dynamic NMR studies indicated a reduced mobility of hydrogen in the presence of alkali promoters (Na and K) at high loadings (65 atomic %). On unpromoted Ru/SiO,_ catalysts, the line due to hydrogen-on-metal was homogeneously broadened at pressures above 0.5 Torr Ha. Similar behavior was observed on promoted Ru/SiO 2 catalyst with 66 atomic % K. The line due to hydrogen-on-metal was inhomogeneously broadened at least up to 200 Torr H 2 o n promoted Ru catalyst with 66 atomic % K. A similar behavior was observed on a 65 atomic % Na promoted catalyst up to pressures of 735 torr and temperatures up to 630 K. A homogeneous lineshape indicates that there is fast exchange of hydrogen among different Ru particles whereas an inhomogeneous line indicates that such an inter-particle motion is restricted. The exchange parameter of hydrogen motion was determined from a multisite exchange model. It was determined that this exchange parameter on unpromoted Ru catalysts was 20 fold higher than that on a KJRu catalyst with 66% K at a given hydrogen pressure. The mechanism for this inhibited mobility was postulated as follows: Alkali blocked the low coordination sites needed for dissociative chemisorption of hydrogen and the kinetics of adsorption-desorption was thereby slowed down significantly. Good quantitative agreement was obtained when the exchange parameters are used to determine the effects of alkali promoters on olefin selectivities in Fischer Tropsch synthesis reaction.

1. INTRODUCTION Alkali promoters are frequently employed in syngas reaction catalysts for their effects on improving the selectivity for higher molecular weight hydrocarbons. However, their mechanism of interaction is not clearly understood due to the complexity of the syngas reactions and the diversity of effects caused by the alkali promoters. The effects of alkali promoters on syngas reactions are reported as (i) increased selectivity towards higher hydrocarbons, (ii) decreased rates of the overall reaction, (iii) increased CO2 production and (iv) increased ratios of unsaturated/saturated hydrocarbons. These effects are attributed to one or more of the following mechanisms: (1) Site blocking effects, (2) through -metal

316 electronic interactions, (3) through-space electronic interactions (e.g. electrostatic), (4) direct chemical interactions and (5) promoter induced surface reconstructions[ 1]. In this paper, we postulate that the primary role of an alkali promoter is to reduce the mobility of the chemisorbed hydrogen on the ruthenium surface based on the data obtained from NMR spectroscopy. We will examine the active hydrogen adsorption states (ct and 13) in supported ruthenium catalysts and report the effects of the alkali promoters on the population and the mobility of the adsorbed states.

2. METHODS 2.1. Catalyst Preparation All catalysts in this study were prepared via incipient wetness. The alkali promoters were introduced in the samples via a co-impregnation or a sequential impregnation technique. The details of the preparation was described elsewhere [1 ]. A Na promoted catalyst was also prepared from a ruthenium nitrosyl nitrate salt obtained from Johnson Matthey. Elemental analysis by atomic absorption spectroscopy indicated that this precursor contained 10.7 wt% Na which corresponded to a Na:Ru atomic ratio of 1.75:1in the final catalyst. Two silica supported catalysts were prepared from this precursor with metal loadings of 4 and 10 wt% after the reduction. The catalysts prepared from the Johnson Matthey salts were labeled as NAJM4 and NAJM10, the numbers denoting the metal loading in each catalyst. These Na promoted catalysts were reduced in flowing hydrogen for 2 hours at 623 K. The corresponding unpromoted catalysts were obtained by washing the reduced Na promoted catalysts in boiling water several times. The absence of the Na promoter in the final catalysts were confirmed by the atomic absorption spectroscopy. 2.2. NMR Experiments The details of NMR sample preparation for sealed samples as well as those investigated via in situ techniques are described previously[I,4-7]. A home-built spectrometer with 1H resonance frequency of 220 MHz was used for the NMR experiments on the sealed samples. All NMR measurements on the sealed samples were performed at 294• 1 K. The dynamics of hydrogen was investigated in-situ in a home built spectrometer with a 1H resonance frequency of 250 MHz. About 100 mg of catalyst was placed in 5 mm O.D. NMR tubes and then attached to the sample port of an in-situ NMR probe described elsewhere [4,5]. The in-situ measurements were carried out over a pressure range of 10-5-760 tort and a temperature range of 294-700 K. Quantitative measurements were performed by data accumulation at a recycle time greater than 5Tx of the resonance of interest. The spin lattice relaxation times, T~, of hydrogen on the metal were determined by the inversion recovery technique [3]. Absolute intensifies were obtained by referring to a water sample. The reference sample was sealed in a capillary tube of the length of the catalyst samples to account for field inhomogeneities in the NMR coil [2].

317 Selective excitation was accomplished by using a delays alternating with nutations for tailored excitation (DANTE) sequence of 20 short rf pulses of'r.w=0.6 ~ts duration [4]. The pulse separation used was 20 ~ts resulting in the total duration of the DANTE sequence of "tt=400 ~s and the corresponding spectral excitation width A of the center band of-= 2.5 kHz. The excitation sidebands were separated from the centerband by (xp)-1 = 50 kHz and did not affect the NMR spectrum. After a recovery period,xr, of at least 35 ~ts a final 90 ~ pulse was applied followed by the detection of the free induction decay. 3. RESULTS The 1H NMR spectra of strongly bound hydrogen in a series of co-impregnated, sodium-promoted catalysts are shown in Fig. 1. The upfield peak (-60 ppm) in each spectrum was identified as hydrogen dissociatively chemisorbed on ruthenium and the downfield peak (3 ppm) was assigned to the silanol protons [1,2]. The upfield peak was best fit by an exponential Gaussian function and for the peak at 3 ppm a superposition of one sharp Lorenzian and one broad Gaussian peak was used. The electronic interactions between the Na promoter and the chemisorbed hydrogen was monitored via the 1H Knight shifts of the hydrogen-on-metal resonance. As seen in Fig. 1, the shift in the hydrogen-on-metal resonance did not change with alkali loading. The same behavior was observed for the sequentially impregnated catalysts as well (not shown).

Si

at % Na

~

H

/

o., t~

0 c:

33.3 50 66.6 75 i

200

160

120

80

40

0

-40

-80

-120

-160

-200

Shift (ppm f r o m TMS)

Figure 1. 1H NMR spectra of co-impregnated Na/Ru/SiO 2 catalysts. Each of the spectra shown is obtained by averaging 1000 scans

318

1.0

~

.~

co-impregnation I ---"-- sequential impresnation I[ _ to-one site blocking= 1

9

0.8 0.6 0.4 0.2

o.o

.

!i

" i

' ;

' i

' lb'

1~'

l~

Na/Surface ruthenium atomic ratio

Figure 2. The site blocking effects of the Na promoter on the Ru/SiO 2 catalysts and co-impregnated catalysts. The results were compared with a theoretical one-toone site blocking case as indicated by a solid line in the same figure. In addition to the intensity of the hydrogen-on-metal surfaces, the presence of Na also influenced the intensity of the hydrogen in the support resonances as well. The intensity of the 3 ppm resonance decreased with increasing alkali loading. The effect of Na loading and the preparation technique on the hydrogen-in-support resonance was plotted in Fig. 3.

The effect of Na promoter on the number of available hydrogen chemisorption sites on Ru was also monitored via 1H spin counting of the hydrogen-on-metal resonance. The quantitative information was extracted by comparing the integrated hydrogen-on-metal resonance intensity with a water reference of known spin count. In Fig. 2, the change in the hydrogen-on-metal intensity was plotted against alkali loading for both sequentially

~, o.6 ~~ 0.5

I I

~ 0.4 K o ~ 0~ ~ 0.2 ~ ~~ 0.~ o.o ~6

I

r

to

2o

t

---I'-- co-impregnation --'-- sequential impregnation j

30

40

5'o

6'0

70

80

Na loading (at % of total metals)

Figure 3. Effect of Na on the OH group intensities in the silica support

The presence of Na promoters decreased the amount of the weakly bound hydrogen states significantly. The weakly bound 13state present on unpromoted catalysts [5-7] was not observed on the Na promoted catalysts (Fig. 4). The 13 state of hydrogen was unobservable on both co-impregnated and sequentially impregnated catalysts. This weakly bound state unobservable on a NAJM10 catalyst appeared after the Na promoter was eliminated from the catalyst by washing with boiling water several times. The dynamics of chemisorbed hydrogen was investigated by selective saturation (hole burning) measurements on an unpromoted and a Na promoted catalyst from the coimpregnated catalyst series. Part of the (-60) ppm resonance was excited with a comb of 30 short pulses which was followed by a 90 ~ detection pulse [4]. The effect of temperature, pressure and the presence of Na on the dynamics of the chemisorbed hydrogen was monitored in terms of the ability to create a "hole" in the -60 ppm resonance. Engelke et al. presented the NMR spectra on an unpromoted Ru/SiO2 after a single 90 ~ excitation and after selective excitation measurements [Fig. 1 and Fig 2 in ref.4] as a function of hydrogen pressure.

319

r,ca

"2 eel

100

-100

0

-200

Shift (ppm from TMS) Figure 4. The effect of the Na promoter on the a and 13 states of the chemisorbed hydrogen: (a) Hydrogen on 4 wt% Ru/SiO2, P=120 torr, T=400 K (b) Hydrogen on NAJM10 after washing in boiling water several times, P=160 torr, T=400 K; (c) Hydrogen on NAJM10, P=740 torr, T=400 K

T (K)

,_~jl

\'

\w/\\

365

~5o

loo

io

a

-io

-loo

Shift (ppm from TMS)

Figure 5. Effect of temperature on the hydrogen-on-metal resonance of Na/Ru/SiO2, single excitation spectra. The spectra are temperature corrected by the Boltzmann factor.

150

100

50

0

-50

-100

Shift (ppm from TMS)

Figure 6. Effect of temperature on the hydrogen-on-metal resonance of Na/Ru/SiO2, hydrogen-on-metal resonance was selectively inverted. The spectra are temperature corrected by the Boltzmann factor.

320 Similar measurements done on a Na promoted catalyst with a Na:Ru atomic ratio of 1.85 and a Ru loading of 10 wt% (NAJM10) indicated that the -60 ppm resonance was heterogeneously broadened at higher pressures (up to 740 torr) and higher temperatures (up to 500 K).In Fig. 5, the single excitation NMR spectra of hydrogen chemisorbed on NAJM10 is given as a function of temperature under 740 torr hydrogen pressure. The selective excitation data of identical conditions are presented in Fig. 6. As seen in Fig. 6, the heterogeneous broadening in the hydrogen-on-metal resonance was independent of temperature and pressure over the range covered. The effect of alkali promoters on dynamics of chemisorbed hydrogen was determined in terms of an exchange parameter which was determined from lineshape simulations as described in [4]. The exchange parameters determined for 0 and 66 at. % K promoted catalysts at a hydrogen coverage of 0.8 were listed in Table 1. Table 1. Exchange parameters of dissociatively adsorbed hydrogen with the gas phase for unpromoted and K promoted Ru/SiO2 catalysts. Alkali promotion ievei ~ (atomic % K) 0

66

ExChangeparameter (s-1) 20 000 1 000

.....

4. DISCUSSION 4.1. Electronic Interactions Between the Na Promoter and the Ru Particles The NMR spectra for a series of co-impregnated Na promoted catalysts are shown in Fig. 1. As can be seen, addition of Na did not influence the Knight shift of the adsorbed protons appreciably similar to the effects of K promoters observed previously [1]. The resonance frequencies obtained after deconvolution of the spectra by peak fitting were also uninfluenced from the presence of the Na promoter. This result is not surprising if one considers the fact that alkali promoters do not exist in their zerovalent state except under ultra high vacuum conditions. However, in the supported metal catalysts, alkali promoters usually exist in their oxide or hydroxide forms [8-11]. In their oxide or hydroxide state, alkali promoters can not donate electrons to the metal substrate as in the case of their zerovalent counterparts [12]. In addition, the heat of hydrogen adsorption on Ru/SiO 2 measured by micro calorimetry indicated that the initial heats of adsorption did not change in the presence of alkali promoters, suggesting that there is no electronic or electrostatic effects on hydrogen adsorption due to K addition[ 13]. Therefore, strong electronic interactions between the alkali promoters and the chemisorbed hydrogen can not account for their role as a promoter for the Fischer Tropsch or the ammonia synthesis reaction.

321

4.2. Site Blocking Effects Na promoters blocked the available hydrogen chemisorption sites on Ru irrespective of the preparation technique. The site blocking effects of Na promoters were monitored via 1H NMR spectroscopy by spin counting. The results presented in Fig. 2 suggests that the site blocking effect of Na promoter was more significant for sequentially impregnated catalysts than for the co-impregnated catalysts. We believe that, this effect is both due to the difference in the preparation chemistry as well as due to a decrease in the metal particle size upon co-impregnation. In fact, for K promoted Rh/SiO 2 catalysts, upon co-impregnation, smaller particle sizes were observed via X-ray diffraction techniques [ 14]. The site blocking effects of alkali promoters as monitored via microcalorimetry indicated that K promoters significantly reduced the adsorbed hydrogen amounts and eiiminated the intermediate (50 kJ/mol) and weakly (10 kJ/mol) bound states [13]. The authors concluded that K promoter selectively populated the defect like sites where the weakly bound hydrogen states are present. 4.3. Dynamics of Adsorbed Hydrogen Alkali promoters not only influenced the availability of the chemisorbed hydrogen, but also its mobility as well. The highly mobile, weakly bound 13 state observed on an unpromoted catalyst was not present on the Na promoted Ru/SiO 2 catalysts (Fig. 4). The structure sensitivity of the 13 state was discussed by Bhatia et al. [6-7]. For example, in the presence of Ag, the population of the 13 state was greatly diminished, suggesting the possibility that the presence of weakly bound 13state was correlated with the low coordinated edge and comer atoms. In the presence of Na promoters, the 13 state completely disappeared suggesting the possibility that the Na promoter preferentially occupied the low coordination sites. The absence of the 1~state in the presence of alkali promoters is also in agreement with the absence of the low energy adsorbed hydrogen states measured by microcalorimetry[ 13]. Furthermore, the 13 state observed on unpromoted Ru/SiO 2 catalysts was in fast exchange with the mobile c~ and gas phase hydrogen [7], the absence of it also suggested that the presence of the Na promoter may have influenced the dynamics of chemisorbed hydrogen on the metal surface. Three features in the spectra presented in Fig. 5 are suggestive of the reduced mobility of the chemisorbed hydrogen on the ruthenium surface in the presence of the Na promoter: 1. The linewidth of the resonance does not decrease as a function of temperature. On an unpromoted surface the linewidth of the proton resonance is a strong function of temperature. 2. The signal intensity remains nearly constant after the Boltzrnann correction is introduced. 3. A highly mobile 13 state present on the unpromoted catalysts at these pressures and temperatures does not appear even at a hydrogen over pressure of 740 tOrT. The postulate of the reduced hydrogen mobility in the presence of Na promoters was tested by the selective saturation experiments done on the unpromoted and Na promoted catalysts. One can selectively excite part of an NMR resonance if the resonance is actually a superposition of a distribution of narrower resonances caused by a distribution of environments (heterogeneous broadening). But when the chemical exchange is fast (with

322 time constants shorter than the duration of the selective excitation pulses), the selective excitation process becomes ineffective [4]. On unpromoted Ru/SiO 2 catalysts, chemisorbed hydrogen experiences fast exchange at higher pressures, and the lines are homogeneously broadened [4]. However, in the presence of Na promoters the NMR line due to the chemisorbed hydrogen were heterogeneously broadened as indicated by the fact that the hydrogen-on-metal resonance could be selectively excited even at high pressures (740 torr) and at high temperatures (533 K). This result indicated that the three dimensional interparticle mobility of the chemisorbed hydrogen as a result of adsorption/desorption processes was significantly restricted in the presence of Na promoters even at high pressures and temperatures. As a result, we postulate here that the restricted mobility of hydrogen is one of the prevailing mechanisms of alkali promotion in the Fischer Tropsch synthesis reaction. The hydrogen adsorption studies on clean and K covered Pd(100) single crystals indicate a decrease in the sticking coefficient as a function of alkali coverage [15]. Similar results have been obtained for hydrogen chemisorption on alkali added Ni, Mo, Pt, W and Ru single crystal surfaces[ 16]. A decrease in the sticking coefficients of hydrogen as a result of alkali metal promoters indicate a decrease in the rates of adsorption at constant flux of the gas phase species. The rates of desorption, on the other hand, depend on the energies of interaction between the adsorbed species and the surface. The microcalorimetry data of Narayan et al. [13] indicated that in the presence of alkali promoters, the initial heats of adsorption did not change. If the heats of adsorption, i.e., the activation energies for desorption, of the high energy adsorption states, remain uninfluenced, one can conclude that the desorption rate constants are also not influenced in the presence of alkali promoters. Based on these arguments, we interpret the decrease in the exchange parameter measured by NMR spectroscopy as a decrease in the adsorption rates. The net effect of this on a steady state adsorption process is a decrease in the net rate of adsorption, as well as a net decrease in the ratio of the rate constants for the forward and the reverse rates, ke/kr. With this in mind and a similar postulate about the CO adsorption kinetics, now we will analyze the model given by Kellner and Bell [ 17] using the data of Okuhara et al.[ 18]. 4.4. Effect Of Adsorption Kinetics On Fischer Tropsch Synthesis This restricted hydrogen mobility mechanism is consistent with the effects of the alkali promoters on the Fischer Tropsch synthesis reaction kinetics and on the product distribution [1]. To date, the decreased hydrogenation capacity of the Fischer Tropsch catalysts in the presence of alkali promoters was explained by the decreased availability of hydrogen on the metal surface. But a simple site blocking mechanism is not consistent with the effects of alkali promoters on the product selectivities. Because, the simple site blocking mechanism does not distinguish between hydrogen and carbon monoxide adsorption. Therefore, the net result of a simple site blocking mechanism should be decreased rates of the overall reaction and rates of formation of all the reaction products by the same amount. However, the product distributions of Fischer Tropsch synthesis reaction on alkali promoted catalysts significantly change towards a decreased hydrogenation capacity as indicated by the increased selectivities to c~-olefins. The decreased hydrogenation capacity and increased ot-olefin selectivity of the alkali promoted catalyst could be attributed to the reduced hydrogen mobilities in terms of

323 adsorption-desorption kinetics. According to a model developed by Kellner and Bell[ 17] for Fischer Tropsch synthesis over alumina supported ruthenium catalysts, the olefin selectivities are inversely proportional to the equilibrium constant for the dissociative hydrogen chemisorption step: Rolcf~/~araf~ - kto/{ktv (KH2 P.2 )'/2} (1) where Ro~cf~ is the rate of olefin formation Rwra~ is the rate of paraffin formation kto is the rate constant for olefin termination step ktv is the rate constant for paraffin termination step KH2 is the equilibrium constant for the dissociative hydrogen adsorption step and PI~2 is the hydrogen partial pressure. If one assumes that the rate constants of the termination steps are not influenced by alkali promotion (not unrealistic as long as the energetics at the surface and surface mobilities of the intermediates are not influenced), then the only parameter for the olefin selectivities remain as the equilibrium constant (or the ratio of the forward and reverse rate constants) for the dissociative hydrogen chemisorption, KH:. The experimental data of Okuhara et al. [ 18] on K promoted Ru/AI:O3 is used along with the model given by Kellner and Bell [17] for olefin selectivities, one immediately sees that the kf/kr of dissociative hydrogen chemisorption step must have decreased approximately 10 fold for a 1%K promotion and 40 fold for a 3%K promotion levels. This effect is consistent with our results except the fact that their alkali promotion levels are much lower than ours. This can be explained by the fact that the support material is different in the two cases (SiO2 in our case and AlzO3 in case of ref [ 18]) which may result in different levels of distribution of the alkali promoter between the metal and support surfaces [1]. Therefore, in the case of an A1203 support, higher alkali coverages on the metal surface may be achieved at lower alkali loadings. Also, on Na promoted Ru/TiO: catalysts, a tenfold decrease in ethane formation rate with respect to unpromoted catalyst was observed, while the rates of ethylene increased slightly at maximum Na loadings [ 19], again consistent with our observations. As one extends this analysis to the chain growth probabilities and methane selectivities, an immediate conclusion reached is that, one has to include the effect of the alkali promoters on the CO adsorption kinetics and on olefin reincorporation steps as well. Since, no direct data on CO adsorption/desorption kinetics is available at the moment, we will leave this as a postulate. 5. CONCLUSIONS 1H NMR spectroscopy studies on Na promoted Ru/SiO 2 catalysts indicated that Na blocked hydrogen chemisorption sites on Ru but not on a one-to-one basis. There was no evidence for a through-the-metal electronic interaction between the chemisorbed hydrogen on the metal and the Na promoter. Na promoter also modified the support by exchanging with the protons of the hydroxyl groups. The weakly bound, structure sensitive 15 hydrogen population was not present on Na promoted catalysts, which suggested that Na could block electron deficient defect like sites. In the presence of Na promoters, the mobility of hydrogen

324 on the Ru particles was significantly restricted. The restricted mobility of the chemisorbed hydrogen was evident from (i) the absence of the mobile, weakly bound 13state observable on an unpromoted catalyst at high pressures; and (ii) the heterogeneous line broadening mechanism detected from selective saturation measurements. The absence of the [5 state and the heterogeneous NMR line broadening mechanism of the alkali promoted catalysts prompted the restricted hydrogen mobility as an alternative mechanism in the Fischer Tropsch synthesis: The restricted mobility postulate is consistent with the decreased hydrogenation capacity effect observed in the Fischer Tropsch synthesis reaction. ACKNOWLEDGMENTS This work was supported by the US Department of Energy, Office of Basic Sciences, under contract W-7405-ENG-82 and the National Science Foundation Engineering Research Equipment Grant CBT-8507418. REFERENCES 1. Uner, D.O., Pruski, M., Gerstein, B.C., and King T.S., J. Catal., 146, 530 (1994). 2. Wu, X., Gerstein, B.C., and King, T.S., J. Catal. 135, 68 (1992). 3. Gerstein, B.C., and Dybowski, C.R., "Transient Techniques in NMR of Solids", Academic Press, New York, NY, 1985. 4. Engelke, F., Vincent, R., King, T.S. and Pruski, M., J. Chem. Phys., 101(9), 7262-7272 (1994). 5. Engelke, F., Bhatia, S., King, T.S., and Pruski, M., Phys. Rev. 13., 49, 2730-2738 (1994). 6. Bhatia, S., Engelke, F., Pruski, M., Gerstein, B.C. and King, T.S., J. Catal., 147, 455-464 (1995). 7. Bhatia, S., Engelke, F., Pruski, M., and King, T.S., Catalysis Today, 21,129-140, (1994). 8. Over, H., Bludau, H., Skottke-Klein, M., Moritz, W. and Ertl, G., Phys. Rev. B 46, 4360 (1992). 9. Pirug, G., Ritke, C., and Bonzel, H.P., Surf. Sci. 257, 50 (1991). 10. Aika, K.-I., Hori, H., and Ozaki, A., J. Catal. 27, 424 (1972). 11. P o n e c, V., Ca talysis To day 12, 227 ( 1992 ). 12. Hrbek, J., Surf. Sci. 205, 408 (1988). 13. Narayan, R.L., Savargaonkar, N., Pruski, M., and King, T.S., Stud in S u r f Sci. and Catal., 101-B, 921-930 (1996). 14. Kesraoui, S., Oukaci, R., and Blackmond, D.G., J Catal. 105, 432 (1987). 15. Solymosi, F., and Kovacs, I., Surf. Sci., 260, 139-150 (1992). 16. "Studies in Surface Science and Catalysis" vol 70, ed,. M. M. Kiskinova, Elsevier, 253. 17. Kellner, C.S., and Bell, A.T., J. Catal.,70, 418-432 (1981). 18. Okuhara, T., Tamura, H., and Misono, M., J. Catal.,90, 41-48 (1985). 19. Komaya, T., Bell, A.T., Weng-Zieh, Z., Gronsky, R., Engelke, F., Pruski, M., and King, T.S., J. Catal., 152, 350-359 (1995).

91997 Elsevier Science B. V. All rights reserved. Dynamics of Surfaces and Reaction Kinetics in Heterogeneous Catalysis G.F. Froment and K.C. Waugh, editors

325

Interparticle migration of hydrogen on zeolite and their participation in the hydrogenation of adsorbed species and catalytic reaction I. Nakamuraa, A. Zhanga, Y. Fanb and K. Fujimotoa aDepartment of Applied Chemistry, School of Engineering, The University of Tokyo 7-3-1, Hongo, Bunkyo-ku, Tokyo 113, Japan bDepartment of Chemistry, Nanjing University 22, Hankou Road, Nanjing 210093, China

It was found by FTIR measurement that strongly adsorbed pyridine on acid sites of USY zeolite was hydrogenated over Pt-hybrid catalyst (a physically mixed catalyst with weight ratio of Pt/SiO2 / USY=I:4) to adsorbed piperidine in the presence of gaseous hydrogen at around 473 K, whereas no such phenomena were observed on either USY or Pt/SiO2. The phenomenon revealed the occurrence of hydrogen spillover from Pt site to zeolite acid sites (Br0nsted and Lewis). In isomerization of n-pentane, the Pt-USY hybrid catalyst showed high activity and selectivity equal to those of PffUSY, while USY and Pt/SiO2 had a poor catalytic performance for this reaction. In the absence of hydrogen, the conversion of n-pentane was drastically reduced and the oligomerization reaction became dominant. These result suggest that spilt-over hydrogen (H+, H-) plays an important role in alkane activation and the stabilization of intermediates to give high isomer yield.

I. INTRODUCTION In recent years a number of studies have been devoted to the phenomenon of" spillover", which is defined as the migration of adsorbed species.from one solid phase where it is easily adsorbed, onto another solid phase in contact with the first, where it is not directly adsorbed. One can find many experimental evidences of hydrogen spillover in IR spectroscopy, often being combined with H-D exchange [ 1-3]~ It is well known that supported metal catalysts and zeolite catalysts are the most important catalysts for processing hydrocarbon feedstocks such as hydrocracking and isomerization. The present authors demonstrated that protons on Y-type zeolite were exchanged very quickly with deuterium in the gas phase by the coexistence of physically mixed nickel sulfide supported on alumina or silica-gel and claimed that the dissociated hydrogen on nickel sulfide spills over from nickel site to support material and then migrate to zeolite. This spilt-over hydrogen has been pointed out that to play a key role for the disproportionation of toluene and the isomerization of xylene [4]. Hattori et aL have pointed out that molecular hydrogen dissociates on the platinum to hydrogen atoms which undergo spillover on the SO42--ZRO2and convert to H+ and e- or H-. The H+ acts as catalytic site for acid catalyzed reactions [5-6].

326 However, the behavior or function of spilt-over hydrogen on zeolite is not clear yet. Pyridine, which is a typical organic base, is chemisorbed on either Br~nsted (B) acid site or Lewis (L) acid site to form pyridinium ion or coordinately bonded pyridine complexes with B or L acid sites, respectively, and giving different IR-adsorption bands [7]. Therefore, the change in the IR band of adsorbed pyridine on B or L acid sites could be attributed to the change in the adsorbed pyridine on acid site. In this paper, pyridine adsorption was utilized to investigate the hydrogen spillover phenomenon from metal to acidic centers of zeolite by observing hydrogenation of chemisorbed pyridine on B or L acid sites of USY zeolite using hybrid catalyst composed of USY zeolite and Pt/SiO2. To ~ve insight into spill-over process in the hybrid catalyst system, isomerization of n-pentane, which is one of the typical acid catalyzed reaction, was also studied.

2. EXPERIMENTAL 2.1

Catalyst preparation

Supported metal catalysts Pt/SiO2 , Pd/SiO2 and Ni/SiO2 with same molar loading (0.33umole metal/mE SiO2) were prepared by first impregnating a commercial available SiO2 powder (Aerosil 380 surface area 380mEg-I) with aqueous solution of H2PtC166H20, PdCl2 and NiC126H20, respectively, and then drying at 393 K for 12 h and calcined them in air at 773 K for 4 h. hybrid catalyst were prepared by c%~rinding the mixture of 1 weight part of the above supported metal catalyst with 4 weight parts of commercial available USY (Catalyst & Chemicals Ind. SiO2/A1203 ratio=8.6 ), respectively, and pressure molding the mixture to granules to 20/40 mesh. Pt-supported USY for isomerization of pentane was prepared by ion exchanging the USY with ammine complex ions of noble metals. The ion exchanged USY was washed by water until no chloride ion was detected. 2.2. IR spectroscopy Pyridine chemisorption and hydrogenation on the hybrid catalysts were carried out in a quart-made in situ IR cell with CaF2 windows connected to a vacuum system and gas reservoir of pyridine and hydrogen. All samples were pressed into self-supporting wafer with diameter of l0 mm and weight of 10 mg. Before pyridine adsorption and hydrogenation, the wafer samples were reduced in flowing hydrogen at 673 K for 1 h and evacuated at 773 K for 1 h. The pretreated sample was exposed to 10 torr vapor of pyridine at 473 K for 15 min followed by evacuation at 673 K for 30 min. The pyridine-covered wafer was cooled to given temperature and exposed to 100 ton" gaseous hydrogen. The IR spectra were recorded before and after pyridine adsorption and hydrogenation, and the absorbance #ven minus the background. All IR spectra were taken on Perkin-Elmer 1600 series FTIR spectrometer. Deuterium exchange with OH groups on USY zeolite was carried out in the same IR measurement system. The sample was reduced at 723 K in flowing hydrogen for 3 h, evacuated at the same temperature for 1 h and exposed to deuterium at 573 K. Every 15 min the spectrum was measured at room temperature.

2.3. H2 chemisorption measurement The surface properties of Sit~_ -supported metal (Pt, Pd, Ni) catalyst were studied by using

329 The phenomenon that spiltover hydrogen hydrogenated the pyridine adsorbed on acidic sites of USY can be schematically expressed as a model shown in Fig. 3. The gas phase hydrogen is first dissociated on metal sites and then spillover to SiO,_ support first, then cross the interface between Pt/SiO2 and USY and finally move to acidic sites of USY zeolite through surface mi~ation. Here, pyridine molecules bonded on USY acid sites acted as an acceptor for spiltover hydrogen. Of course, there may be a doubt that the adsorbed pyridine move to metal site and after being hydrogenated, it comes back to acidic site. However, following phenomena can deny this possibility. (1) Weak adsorbed pyridine molecules should be removed during the evacuation at 673 K, therefore, the pyridine remaining after this treatment will be tightly bonded on zeolite acidic sites and should not migrate quickly at 473 K. (2) ff pyridine moved to other site, the LR adsorption wavenumbers of adsorbed pyridine should correspondingly shift to lower or higher regions, but it was not observed.

3.2. Effect of metal properties on hydrogen spillover Hydrogenation of pyridine over USY mixed with different supported metal catalysts were studied to give insight into the effect of metal properties on hydrogen spiHover. As shown in Table 1, which gives pyridine conversion on USY mixed with Pt/SiO2, Pd/SiO2 and Ni/SiO2, the activities of metal catalysts for hydrogenation of pyridine were PIYSiO2 > Pd/SiO2 > Ni/SiO2. Since adsorbed pyridine on BrCnsted and Lewis acid sites were hydrogenated by spilt-over hydrogen species from metal surface, the above results indicate that the rate of hydrogen spill-over is controlled by the nature of metal used in hybrid catalysts. H2 chemisorption was used to characterize the surface properties of the above supported metal catalysts. It can be seen from Table 1 that H2 uptake shows the order as Pt/SiO2 > Pd/SiO2 > Ni/SiO2, which is in a~eement with that of activity for pyridine hydrogenation. As the three metal catalysts have same metal loading (0.33 umol m2-SiO2), it can be proposed that these metals have the same dispersion on silica support. The difference in H2 uptake may suggest different equilibrium concentration of dissociated hydrogen on metal surface. The higher concentration of dissociated hydrogen on Pt surface make hydrogen diffusion onto SiO2 surface and USY surface fast. This can account for the reason that Pt/SiO2 has the highest activity among the three metal catalysts for hydrogenation of adsorbed pyridine on acidic surface of USY via hydrogen spill-over. Table 1 Effect of metallic properties on hydrogenation of adsorbed pyridine over hybrid catalyst Catalyst

USY Ni/SiO 2 + USY Pd/SiO 2 + USY V t / S i O 2 + USY

Pyridine conversion (%)a B-py L-py 0.0 0.0 3.4 57.0

"Reaction temperature 473 K, 1 h

0.0 14.8 14.9 80.0

H 2 uptake of M/SiO 2 (ml/g metal) 0.0 20.3 29.7 5.0

330

3.3. H-D exchange on USY zeolite Deuterium exchange with OH groups on USY zeolite was carried out using Pt-USY catalyst. Change in the IR spectrum of USY zeolite in exposure to deuterium are shown in Figure 4. In the case of Pt/USY, the intensity of acidic OH band at 3650 cm-1 decreased and the band corresponding OD ~ o u p s appeared at 2680 cm-1 increased by exposure to deuterium. On the other hand, acidic OH groups on Pt-free USY zeolite unchanged and new OD ~ o u p s did not appear when it was exposed to deuterium. It was suggested that the observed H-D exchange of acidic OH ~ o u p s on the Pt-USY zeolite attributed to the spiltover deuterium from Pt to br0nsted acid site on the USY zeolite. In other word, spillover hydrogen can exist on USY zeolite as proton.

B acid (H)

=~

J

~9

/1 ',, .t,,

B acid (H)

Pt-USY

B aci (I3) i [li 'X

~.

~,~

,,

3800

~ 3600

USY

9 .',,.'x~,

]]i ', ,'",

"~

["

->

'-~.2800 , . . .9600 . -. 3800 -3-*00 3 6, 0 0

Wave number (cm -I)

B acid (D)

: v , . . . . .2600 .. 3400"2800

Wave number (cm -1)

Figure 4. IR spectrum of USY zeolite in exposure to deuterium at 573 K.

3.4. Isomerization of n-pentane The results of pentane isomerization are shown in Table 2. Table 2 Isomerization of n-pentane on USY catalyst catalyst

USY

Atmosphere Conversion (%) Selectivity (C-mol%) CI-C 2 C3-C4 iso-C 5 C6" , ,,

Pt/SiO2 a

Pt-USY

Pt/SiO 2 % + USY

H2 0.5

H2 1.8

HE 30.5

H2 30.6

N2 1.7

0.0 12.9 87.1

13.3 12.2 66.7 10.5

0.1 0.3 99.3 0.3

2.2 2.2 97.7 0.1

0.8 65.0 22.4 12.0

0.0

l:~./SiO2 b + USY

Reaction temperature 563 K, n-C 5 10 kPa, H 2 or N 2 90 kPa, W/F = 5.0 g-catoh.mol ~, ~Pt/SiO 2 0.20 g, bPt(2.5 wt%)/SiOE-USY = 1:4.

331 It is clear that platinum-supported USY (Pt-USY) or hybrid catalyst containing Pt/SiO2 (Pt/SiO2 + USY) shows high n-pentane conversion and iso-pentane selectivity while both USY or Pt/SiO~_ shows quite low activity and low iso-pentane selectivity. It should be also noted that the catalytic activity of the hybrid catalyst was extremely low under nitrogen atmosphere. This clearly shows that hydrogen gas is essential for the generation of the isomerization activity. 3.5. Reaction model of h y d r o i s o m e r i z a t i o n

of n - p e n t a n e on Pt/zeolite catalyst

The most predominant theory for the hydroisomerization of paraffins over noble metal-solid acid catalysts has been based on the bifunctional theory, which is composed of the dehydrogenation of n-paraffins to n-olefins on the noble metal and its isomerization on the solid acid to iso-olefins and again its hydrogenation on the noble metals to iso-paraffins [8]. However, amount of the intermediate olefin is apparently very small by limitation of equilibrium, the olefin is protnated to be carbenium ion rapidly over bifunctional catalyst. The theory explains well the effects of hydrogen gas and the necessity of noble metals. In isomerization of n-pentane, the Pt-USY hybrid catalyst showed high activity and selectivity equal to those of Pt/USY, while USY and Pt/SiO2 had a poor catalytic performance for this reaction. According to the bifunctional theory, pentene formed in the dehydrogenation over Pt/SiO2 should be equal to the equilibrium composition to explain the i-pentan yield of the hybrid catalyst system. Fig 4. shows the results of dehydrogenation of n-pentane over Pt/SiO2 and the calculated equilibrium conversion. The amount of product pentene in outgas was extremely low, it was far lower than the equilibrium composition. The formation of pentene is not controlled thermodynamically but kinetically. The results shown in Fig suggest that the dehydrogenation activity of supported metals is not essential for the appearance of the paraffin isomerization activity, but the ability of hydrogen activation (dissociation) of the catalyst seems to be essential as well as the acidity. Furthermore, if the whole reaction is controlled by olefin formation, the higJa isomerization reaction activity on Pt-hybrid catalyst can not be expected. 20 []

~15 -

C5= yield

9 Conv.%(equil.)

r.,r

=10 -

1 .o atm

:2 O

r,..) 5 -

m

2OO

m

m

250

n

m n

mm m

--

300 350 400 Temperature (~

450

500

Figure 4. Dehydrogenation of pentane over Pt/SiO2. W/F=5.0 g-cat h mol-1, 513 K, n-C5 10 kPa, N2, 90 kPa.

332 To explain these results consistently, the present authors propose a new concept for the isomerization reaction on UYS catalyst, namely that spilt-over hydrogen from the gas phase onto zeolite plays an important role in generating hydroisomerization activity shown as in Fig. 5. Hydrogen gas is dissociated on the metal such as Pt and spills over onto the zeolite either for the Pt-USY or the Pt/SiO2 + USY hybrid catalyst. The hydrogen transfer between particle has been well known phenomenon (5). The spilt-over hydrogen presumably exists on the zeolite surface as proton and hydride, H2 ~ H++H-. The proton should act as the acid to promote isomerization or cracking. The n-CsHI 1+ is isomerized to iso-C5H11 + and then isoCsHt 1§ is stabilized by hydride addition (iso-CsH11 + + H- ~ iso-CsH12). If the supply of hydrogen from the gas phase is insufficient, as is the case of USY or granular mixed hybrid catalyst, the deficiency of H+ and H- on the USY surface causes low n-pentane conversion and low iso-pentane selectivity. The hydrogen spillover concept is most acceptable for explaining the high activity and the high selectivity of the hybrid catalyst. CH3-CH2-CH2-CH2-CH3

-H +

CH3-CH-CH2-CHz-CH3+ H2

CH3 I CH3-C-CH2-CH3+

H2 tt H H

_

-IV

CH3 !

CH3-C=CH-CH3

= I H_ oligomerization CH3 I CH3-CH-CH2-CH3

cracking, aromatization

Figure 5. Hydrogen spillover model in isomerization of pentane over hybrid catalyst.

REFERENCES

1. P. A.Sermon and G. C.Bond, Catal. Rev. 8 (1973) 211. 2. W.C.Conner and J. L. Falconer, Chem. Rev. 95 (1995) 759. 3. U. Roland, H. Winkler, H. Bauch and K.H. Steinberg, J. Chem. Soc. Faraday Trans. 87 (1991) 3921. 4. M.G. Yang, I. Nakamura and K. Fujimoto, Appl. Catal. A. 127 (1995) 115. 5. K. Ebitani, H. Konishi and H. Hattori, J. Catal. 130 (1991) 257. 6. K. Ebitani, H. Konno, H. Konishi and H. Hattori, J. Catal. 135 (1992) 60. 7. R.H. Thomas and M.W. Harry, J. Phy. Chem, 71 (1967) 2192. 8. B.C. Gates, J.R. Katzer and G.C.A. Schuit, Chemistry of Catalytic Processes, MaGraw Hill Inc., pp.280 (1979).

333

"State-Defining" T A P Pulse R e s p o n s e E x p e r i m e n t s J.T. Gleaves a, G.S. Yablonskii a, P. Phanawadee a, Y. Schuurman b aDepartment of Chemical Engineering, Washington University Campus. Box 1198, Washington University, MO 63130, U.S.A. bCentre National De Recherch6 Scientifique Institut de Recherch6 sur la Catalyse 2 avenue Albert Einstein, 69626 Villeurbanne Cedex, France New theoretical results pertaining to TAP vacuum pulse response experiments are described. The theoretical analysis of a "state-defining" experiment, characterized by an insignificant change in the composition and kinetic characteristics of a catalyst, is given. The analytical results for surface coverage distribution are obtained using an irreversible adsorption + diffusion model. It is shown that a one-pulse experiment is a state-defining experiment. For a multipulse experiment, the number of pulses within a state-defining regime is estimated. 1. INTRODUCTION The TAP reactor system is a powerful tool for unraveling complex catalytic reactions, and other surface processes [1 - 10]. Recently, we developed a new experimental apparatus, called a TAP-2 reactor system that significantly extends the capabilities of the original TAP apparatus, and can be used to perform a special type of kinetic studies that we have termed "interrogative kinetics" (IK) [11]. Interrogative kinetics attempts to systematically probe a variety of different states of a catalyst surface, and to understand how one state evolves into another. An important element of this approach is the rapid-feedback between experiment and analysis that can be likened to a "dialogue" between the researcher and the catalyst sample. The set of experiments that form an IK sequence represents a "question" (e.g. how does a change in the oxidation state of a catalyst change its selectivity, or the activation energy of hydrocarbon conversion) that when answered leads to another question and another IK sequence. Interrogative kinetics involves two types of experiments that are performed in sequence to reveal complex mechanisms and structure-activity relationships. The first type is a "state-defining" experiment that does not significantly perturb the kinetic state of the catalyst, but provides information that can characterize the state. The second type is a "statealtering" (typical) experiment that perturbs the catalyst and changes its composition or structure in some predetermined fashion. To complete the sequence another state-defining experiment is performed to characterize the new state of the catalyst. The TAP-2 system can be used to perform a variety transient response and steadyflow experiments, but is distinguished by a unique type of transient response experiment called a TAP vacuum pulse response experiment. This experiment is performed by injecting a narrow gas pulse into an evacuated, packed bed micro-reactor. The pulse contains a very small number of molecules (10 -10 moles) in comparison to conventional transient response experiments. When the pulse contains a small enough number of molecules it travels through

334 the packed bed by Knudsen diffusion. An important feature of TAP vacuum pulse response experiments is that they can be either "state defining" or "state altering" experiments. The goal of this paper is to describe new results in the theoretical foundation of TAP vacuum pulse response experiments, and to define those conditions under which a TAP experiment can be considered "state-defining".

2. T H E S T A T E - D E F I N I N G R E G I M E

A state defining experiment is one in which the composition and kinetic state of a catalyst does not change significantly during the experiment. A theoretical analysis will show that during a TAP vacuum pulse response experiment, at sufficiently low pulse intensities, the composition change in the catalyst is negligibly small throughout the reactor. Assuming that changes in the composition of a catalyst are due to its interaction with adsorbed species, the amount of change due to adspecies A can be estimated from the fractional surface coverage of A, Oa, calculated using an irreversible adsorption + diffusion model. We will also present the equation to calculate the fractional surface coverage after one pulse, and the number of pulses that are allowed in a state-defining multipulse experiment. 2.1. Mathematical Model for Irreversible Adsorption: One Pulse Experiment

If a reactor is completely filled with a catalyst sample, if adsorption is first order in gas concentration, and ifO A is small (it will be shown that 0 A is very small) so that (1-0 A ) is close to unity, then the mass balance for the gas phase component A and adsorbed species A in dimensionless form [ 11 ] can be described by equations 1 and 2 respectively. ~, t~l"

~2(~

=

A_ k C

03~ 2

a

(1)

A

~0A _ m ~-"~ = k a C A

(2)

The dimensionless variables and parameters are defined as followed: Dimensionless axial coordinate:

~=z L

(3)

CA C a -- IVP ,, a,Eb I A. L,. m

Dimensionless concentration:

Dimensionless time:

z =

(4)

tDeA

(5)

eb L2 -

Dimensionless adsorption rate constant:

ka=ka

eb L2 De A

The apparent adsorption rate constant, k'a (s-1), is defined as

(6)

335

k" = asSv(l _ e.b)ka

(7)

a

6b and OA is defined as m

(8)

OA = a O A

where

a = (1 - e b)AL asSv

(9)

N,,A

i

The variable OA in equation 8 is the product of the fractional surface coverage and the catalyst number, o~ The catalyst number is the ratio of the total number of moles of active centers and the number of moles of gas A in the inlet pulse, or the number of active centers per molecule of gas A in the inlet pulse. In a typical TAP-2 pulse response experiment, the amount of catalyst sample, such as a metal or metal oxide, is =10 -1 g, the total number of active sites is usually between 1017 to 1019, and the number of molecules in a pulse is = 1014. Thus, in a typical TAP-2 pulse response experiment, the catalyst number is 103 to 105. OA is called the pulse-normalized dimensionless surface concentration, and is described by m

"OA --

(10)

CAs i p A I S , otaI

Equations 1 and 2 are solved using the following initial and boundary conditions: Initial conditions: 0_<~<1,

t-O,

0< ~<1,

t =0,

C A =5~

(11)

m

OA = 0

(12)

Boundary conditions: m

~-0,

cgC A = 0

=1,

CA = 0

(13)

m

(14)

2.2. Surface Coverage on a Catalyst after One Pulse The solution for the dimensionless concentration, C A (~, 1:), can be obtained by applying the method of separation of variables. A detailed solution can be found in the literature [11]. Here we are interested in the pulse-normalized surface concentration that is described by equation 15. m

336

( n + 0. 5)2 + 0.5)~)+k.a -- s

OA--=2k a

+ T(a exp(-((n +0.5)

+k )

(15)

When time approaches infinity, equation 15 can be written as -

-

=

cos((n+O.5)zr~)

(16)

0 A,~ = 2ka s (n + 0.5)2 ~ 2 + k-a n---0 m

where 0A,~ is the pulse-normalized surface concentration after one pulse. Equation 16 can also be written in another form [12] as k ~ a sinh[(1 - ~) k ~ a ] 0 A,~ = cosh k ~

(17)

The pulse-normalized surface concentration, 0A,oo is largest at ~ = 0, because the catalyst at the reactor entrance is exposed to the highest gas concentration at the time of injection. The equation for 0 A,~, at ~ = 0 or 0 A,~,~--0 is given by: m

m

0 A,oo,~_---0= k ~ a ta/lh k ~ a When k a is sufficiently large and tanh

0 A,oo,~_--0 ----k ~ a

(18) --- 1 equation 18 can be written as (19)

It should be noted that k ~ a or V[k'aebL DeA is similar to the classic Thiele modulus which appears in diffusion-reaction equations for first order reaction in a cylindrical catalyst pore. Equations 16 or 17 can be used to calculate 0 A,~ at different positions in the reactor. Table 1 presents values of 0Am at 5 different positions along with the corresponding fractional surface coverage after one pulse, OA,~, for a equal to 103 and 105, and/~a = 10. The value of k,, corresponds to a conversion of 91.5% of gas A. Table 1 shows that 0 A has a very small value throughout the reactor including positions close to the reactor entrance. Under these conditions, the assumption (1 - 0 A ) -- 1 is clearly valid. If the change in catalyst composition is due to adsorbed species, then the change can be estimated from 0 m . The order of magnitude of Omoo shown in Table 1 is sufficiently small that the change in a catalyst after one pulse can be considered negligible. Consequently, a typical TAP vacuum pulse response experiment (one pulse experiment) can be considered a state-defining experiment.

337 Table 1 Distribution of 0 A for k a - 10 after one pulse and corresponding Om for ~X - 103 and 105 0.00 0 a,=

0.25

0.50

0.75

0.99

Oaoo (a--lO 3)

3.15 3.15X10-3

1.42 1.42X10-3

0.62 0.62X10-3

0.23 0.23X10-3

0.01 0.02X10-3

0 A o o (0:--105)

3 15xlO-5

1.42X10-5

0.62X10 -5

0.23X10 -5

O.02xlO-5

It is also interesting to determine an expression for the average pulse-normalized surface concentration, 0 A,~,avg, which is described by equation 20. 1

0 A,~,avg "- I ~ A d~ = 1 --

o

1

(20)

c~ k ~ a

The conversion of gas A, X, is described by equation 21 [13].

1

X=Icosh

=1 - m0

(21 )

Equations 20 and 21 show that integration of O A,,,~ with respect to the axial coordinate is the conversion of gas A. This result corresponds to the conservation of mass. Equations 20 and 21 also show that the average pulse-normalized surface coverage after one pulse is simply equal to conversion of gas A. (22)

0 m,~,avg = X

Equation 22 is important and useful for determining the state-defining regime for a multipulse experiment. Another interesting relationship can be obtain by dividing equation 18 by equation 20. The result, given by equation 23, is the ratio of the pulse-normalized surface concentration at the entrance of the reactor and the average pulse-normalized surface concentration. 0A,~*.~=0 ~ A..... g

tanh ~ - ~

1-- c o s h l ~

k ~ mnh k ~

(23)

X

Equation 23 shows that a higher rate constant or a higher conversion gives a greater ratio. For example 0 A,~,,=0/Om,~a~'g is 4.5 when X - 0.975 and 2.0 when X - 0.025.

2.3. Multipulse Experiments and the State-Def'ming Regime TAP vacuum pulse response experiments usually involve a train of pulses. An important question is how many pulses of a reactant are required to cause a significant

338 change in a catalyst. To answer this question, we first define a criterion for a state-defining regime. Our choice for this criterion is that the accumulating Om,oo.avgcaused by the train of pulses does not exceed 0.01. After the first pulse, the accumulation of the surface coverage after each consecutive pulse at each position in the reactor is approximately equal to the surface coverage after the first pulse at the corresponding position provided the pulse intensity and operating conditions remain constant, and the number of pulses injected are not too large. The fractional surface coverage at each position in the reactor after each consecutive pulse will then increase approximately linearly with the number of pulses. Therefore, the quantity OAooavg after the n-th pulse is approximately equal to n times

OA,ooavgafter the first pulse.

m

For example, when X - 0.915, 0 A,oo,avg after first pulse will be

equal to 0.915 which is equivalent to OAoo,avgof 0.915• .5 for a of 105, and OAoo,avgafter 1090 pulses will be equal to 0.01. In this case, the state-defining regime spans the first 1090 pulses. Since conversion is the quantity that is observed experimentally, it is useful to write an equation to estimate the number of pulses within the state-defining regime when the conversion of the first pulse is known. Thus, in terms of the conversion of the first pulse, the number of pulses within the state-defining regime is given by: number of pulses -

O.Ola

(24)

X

Although the conversion changes with the number of pulses, the change is very small within the regime defined by equation 24, and it gives a good estimate of the regime. Table 2 presents results for a number of different conversions and catalyst numbers. It shows that the number of pulses that can be injected into a catalyst bed while remaining in a state-defining experimental regime can range from ten to ten thousand. It should be noted that when the accumulating average fractional surface coverage reaches 0.01, the fractional surface coverage at the entrance will be larger, and can be estimated from equation 23. For example, the fractional surface coverage at the entrance is 0.045 for X = 0.975, and 0.02 for X = 0.025 when the accumulating average fractional surface coverage reaches 0.01.

Table 2 Domain of the experimental parameters for multipulse state-defining experiments X

OA,~avg after first pulse

number of pulses

( 0 A,~,avg ) of first pulse

a-lO 3

0.975 0.750 0.500 0.250 0.025

0.975• -3 0.750x 10-3 0.500x10 -3 0.250x 10-3 0.025x 10-3

a-lO 5 0.975x 10-5 0.750x 10-5 0.50(O10 -5 0.250x 10-5 0.025x 10.5

a=lO 3 10 13 20 40 400

a=lO 5 1000 1300 2000 4000 40,000

Table 2 reports the domain of experimental parameters for a multipulse state-defining experiment in a "one-zone-reactor", or reactor that is completely packed with a single

339

catalyst sample. Similar results are obtained for a three-zone-reactor in which the catalyst sample is sandwiched between two zones that contain inert material. In a three-zone-reactor 0 A,=,avg,

after the first pulse, is also equal to the conversion as a result of the conservation of n

mass. Consequently, Table 2 is also valid for a three-zone-reactor. The definition of 0A in a three-zone-reactor is the same as that in a one-zone-reactor except that L in equation 9 is replaced with the length of the catalyst zone. In a three-zone-reactor, the average fractional surface coverage and the fractional surface coverage at the reactor entrance are closer in value than in a one-zone-reactor at the same conversion. The reason that the two coverages are closer in a three-zone-reactor is that the surface coverage along the catalyst bed is more uniform. The presence of inert-materialzones before and after the catalyst bed decreases the non-uniformity of the gas concentration distribution in the catalyst bed.

3. C O N C L U S I O N The axial distribution of the pulse-normalized surface concentration after one pulse was analyzed using an irreversible adsorption-diffusion model. Expressions for the average and entrance pulse-normalized surface concentrations were given. The average fractional surface coverage was used to indicate the change in a catalyst's composition. It was shown that a one-pulse experiment is a state-defining experiment. For a multipulse experiment, the number of pulses that can be carry out within state-defining regime can range from ten to ten thousand pulses. The relationship between conversion of the first pulse and the number of pulses within the state-defining regime is the same for both one- and three-zone reactors.

NOMENCLATURE a s - concentration of active sites (mol/cm 2 catalyst) A = cross sectional area of the packed bed (cm 2) C A = concentration of gas A (mol/cm 3) a

CA

--

dimensionless concentration defined by equation 4

= concentration of A on the catalyst surface (mol/cm 2 of catalyst) DeA- effective Knudsen diffusivity of gas A (cm2/s) k a = adsorption rate constant (cm 3 gas/mol - s) CAs

apparent adsorption rate constant defined by equation 7 (s -1)

k'a -

k a = dimensionless adsorption rate constant defined by equation 6 L - length of the reactor (cm) m 0 = zeroth moment of the exit flow of gas A. NpA number of moles of gas A in the inlet pulse (mol) Stota l -- total surface area of the catalyst (cm 2) Sv - surface area of catalyst per volume of catalyst (cm-1) t - time (s) X = conversion of gas A z - axial coordinate (cm) a - catalyst number defined by equation 9 5 - delta function e b - fractional voidage of the bed --

340 - dimensionless axial coordinate defined by equation 3 0 m - fraction surface coverage of component A u

O A = pulse-normalized surface concentration defined by equations 8 and 10 Oao. = 0 a after n-th pulse Oa.~,~=O = Omoo at the entrance of the reactor Om~o,~vg = average of Om,.o 0 a,~

-- 0 a after n-th

0 a,=,r 0 A,oo,avg

pulse

= 0 a,o. at the entrance of the reactor -~

average of 0 moo

REFERENCES 1. J.T. Gleaves, J.R. Ebner and T.C. Kuechler, Catal. Rev. - Sci. Eng., 30 (1988) 49 2. D.R. Coulson, P.L. Mills, K. Kourtakis, J.J. Lerou and L.E. Manzer, Stud. Surf. Sci. Catal., 72 (1992) 305 3. F.D. Kopinke, G. Creten and G.F. Froment, Stud. Surf. Sci. Catal, 72 (1992) 317 4. D.R. Coulson, P.W.J.G. Wijnen, J.J. Lerou and L.E. Manzer, J. Catal., 140 (1993) 103 5. O.V. Buyevskaya, M. Rothaemel, H.W. Zanthoff and M. Baems, J. Catal., 150 (1994) 71 6. E.P.J. Mallens, J.H.B.J. Hoebink and G.B. Marin, Stud. Surf. Sci. Catal., 81 (1994) 205 7. G. Creten, D.S. Lafyatis and G.F. Froment, J. Catal., 154 (1995) 151 8. J. T. Gleaves, A. G. Sault, R. J. Madix and J. R. Ebner, J. Catal., 121(1990) 202 9. G. Svoboda, J.T. Gleaves and P.L. Mills, Ind. & Eng. Chem. Research, 31 (1992) 19 10. G. Golinelli and J. T. Gleaves, J. Mol. Cat., 73 (1992) 353 11. J.T. Gleaves, G.S. Yablonskii, P. Phanawadee, Y. Schuurman, Applied Catalysis, 1997 to be published 12. P. Phanawadee, Doctoral Dissertation, Washington University, 1997, to be published 13. G. Svoboda, Doctoral Dissertation, Washington University, 1993

91997 Elsevier Science B. V. All rights reserved. Dynamics of Surfaces and Reaction Kinetics in Heterogeneous Catalysis G.F. Froment and K.C. Waugh, editors

341

Transient and steady-state microkinetic models of catalytic reactions on n o n u n i f o r m surfaces L. J. Broadbelt a and J. E. Rekoske b aDepartment of Chemical Engineering, Northwestern University, Evanston, IL 60208 bDepartment of Chemical Engineering, University of Delaware, Newark, DE 19716 Discerning the extent to which thermodynamic and kinetic surface nonuniformity affect observable phenomena in analyses of heterogeneous catalytic reactions is complicated by the complexity of the reactions and interference from other effects. To eliminate these additional phenomena from clouding the interpretation of the kinetics data, a computational approach was developed and adopted in which transient and steady-state data to be investigated were numerically generated from appropriate models. A representative mechanism was used with defined kinetic and thermodynamic parameters, and temperature programmed desorption and reaction and steady-state kinetics data were generated. The transient data were analyzed assuming the reaction was carried out over a uniform surface, ignoring the known surface heterogeneity, and an optimal set of parameters to describe the transient data were obtained. These parameters were then used in a uniform surface model to predict the steady-state rate that would be observed over a range of reaction conditions. The predicted uniform surface rate was in general a factor of 30 higher than the known nonuniform surface rate, indicating that kinetic and thermodynamic parameters obtained from transient studies alone are insufficient when surface nonuniformity is present. Furthermore, this lack of agreement between the predicted and experimental steady-state data may be one "kinetic fingerprint" of reactions on nonuniform catalytic surfaces. 1. I N T R O D U C T I O N The environmental and economic impact resulting from active and selective catalytic processes is tremendous, with reactions over catalytic materials dominating the chemical and petroleum industries. To improve existing catalysts further and to aid in the development of novel ones, significant activity has been aimed at rational catalyst design [1]. It is clear from decades of research that the composition and structure of a catalyst determine its activity and selectivity. Accordingly, modern approaches for designing improved catalysts seek to establish a fundamental link between catalyst synthesis, or structure, and catalyst performance, using both experiment and theory. However, the complexity of catalysts for heterogeneous reactions has presented a serious challenge to meeting this goal. One inherent property of a catalyst that has obscured the bridge between structure and reactivity is surface nonuniformity. Although the role of surface nonuniformity in catalytic kinetics has been acknowledged by many researchers, a limited number of investigations has sought to quantify its impact on observable kinetics. Quantitative analysis of catalytic processes has advanced significantly by consolidating experimental, theoretical and semiempirical information via microkinetic analysis [2]. Extension of this approach to nonuniform surfaces has been hampered, in part, by the staggering complexity afforded by potential feedstocks interacting with a heterogeneous site distribution. Development of a straightforward modeling approach that rigorously takes surface nonuniformity into account would be of generic value. With such a tool, the extent to which thermodynamic and kinetic surface nonuniformity may affect observable phenomena in analyses of heterogeneous catalytic reactions could be identified.

342 The complexity of heterogeneous catalytic systems and the number of complicating phenomena (e.g., mass and heat transfer, readsorption, side reaction and deactivation) which may be experimentally present hinder the unambiguous determination of the existence of surface nonuniformity effects. To eliminate these additional phenomena from clouding the interpretation of the kinetics data, a computational approach was adopted in which the "data" to be investigated were numerically generated from appropriate models. The use of generated data has two major advantages: (1) the mechanism and kinetic values used in its generation are exactly known, and (2) the phenomena which are included in the generation of the data (e.g., mass transfer limitations and surface nonuniformity) are explicitly specified. Of course, application of kinetic tests devised for the investigation of surface nonuniformity effects in real experimental data would not require this restrictive state of knowledge. Rather, the tests are first developed in the controlled environment of generated data sets to be later applied, and refined, on real experimental data, much in the same way rules for heat- and mass-transfer correlations in catalysis have been developed.

2. SURFACE NONUNIFORMITY APPROACHES IN CATALYSIS: PRESENT STATE OF K N O W L E D G E Considerable effort has produced numerous examples [3-8] where the adsorption of gaseous species on catalyst surfaces proceeds in a fashion which is not consistent with the description of the process of adsorption provided by Langmuir [9]. Though many nonuniformity claims resulting from irreconcilable differences between kinetic investigations and expected surface chemical models are available [3], few studies have investigated the kinetic repercussions of reactive site nonuniformity [ 10, 11 ]. One strategy to overcome the shortcomings of uniform surface models has been to incorporate non-Langmuirian adsorption isotherms to explain the adsorption of gases onto nonuniform surfaces. Corma et al. [12] adopted this approach, incorporating surface nonuniformity solely through a surface coverage averaged over all sites of varying adsorption energy, which allowed the use of a single-valued rate constant. However, rate constants are generally dependent on the overall enthalpy or free energy change of a reaction step as evidenced by the wide application of linear free energy relationships (LFERs) in catalysis [13, 14]. Thus, both the adsorption equilibrium and the rate constant may be affected by the adsorptive characteristics of a catalyst surface [2]. Boudart [15] investigated the quantitative aspects of reaction rates at nonuniform surfaces in comparison to rates estimated from uniform surfaces for some simple two-step reaction mechanisms. The author has shown that a uniform surface model will predict a rate of reaction within a factor of 2-4 of the nonuniform surface rate if one assumes a difference between the maximum and minimum adsorption energies on a typical catalyst to be 10 kcal mo1-1, and the allowed range of energies varies symmetrically about the reactive site of maximum rate. We have recently combined these two approaches and found that the conclusion of Boudart can be further generalized in several ways. Uniform surface models were shown to predict the observed rate from nonuniform surfaces within a factor of approximately 4 for breadths of energy distributions ranging from 0 to 20 kcal mol 1, regardless of the relative positions of the allowed adsorption energy range and the adsorption energy of maximum rate. This suggests that the discrimination of uniform and nonuniform kinetic models solely on the basis of steady-state kinetic data would be difficult. However, it was concluded that for cases in which a factor of approximately 4 in rate is not sufficiently accurate, the only recourse is a siteby-site accounting of the reaction rate as a function of adsorption energy [ 11 ].

343 3. N O N U N I F O R M SURFACE M O D E L I N G CAPABILITIES: A MICROKINETIC A P P R O A C H

The intersection of the known heterogeneity of catalyst surfaces and the currently insufficient modeling capabilities to describe reactions on these surfaces creates a void in the tools which can be applied in the field of catalysis. To begin to fill this void, we have developed a mett;,~dology which directly incorporates surface nonuniformity by allowing heats of adsorption on the catalyst surface to be described by a distribution rather than a single value [16]. Once this distribution of adsorption energies is defined for all adsorbed species, rate constants for elementary steps can be estimated, differential mass balances and reactor design equations can be constructed, and these can be combined and solved in a process which has been termed "microkinetic analysis" [2]. Two complementary modeling methods incorporating these features have been developed which parallel common experimental approaches in catalysis research: transient and steady-state kinetics investigations. The transient model allows simulation of temperature programmed desorption (TPD) and reaction (TPR) results over both uniform and nonuniform surfaces, providing modeling capabilities for two techniques which are widely used to identify underlying reaction fundamentals. The steady-state model allows the incorporation of these reaction fundamentals to correlate and predict typical kinetics measures (e.g., conversion, selectivity, yields, etc.) over an unlimited range of reaction conditions. Our nonuniform surface modeling methodology is best illustrated through an example. Consider the condensation of chemical species A and B to produce C on the surface of a catalyst whose adsorptive sites, denoted by *, vary in their ability to bind all three components in some yet undefined manner. The condensation is presumed to occur via an associative LangmuirHinshelwood mechanism presented in Scheme I below. This scheme will be used throughout the latter sections of this paper. 1.

A

+

2.

B

+

3.

A-* +

*

~

A-*

*

~

B-*

B-* ~

C-* +

*

4. C-* ~ C + * Scheme I: Simple condensation mechanism used in description of methodology. In traditional kinetic analyses, a simplifying assumption would be presented, allowing selection of an appropriate rate expression [17]. Though the concept of a rate determining step may be meaningful for some cases involving surface nonuniformity [18], its applicability is valid only for a restricted subset of all reactions and reaction conditions. In microkinetic analyses, it is preferable to write differential mass balance and reactor design equations [2] without the selection of a rate determining step or other simplifying assumption and allow the solution to these differential equations to suggest possible simplifications. Initially, expressions for the individual rates of the elementary steps in Scheme I are sought and can be written according to mass action kinetics [ 17]. Unfortunately, because of the nonuniformity of the adsorption of all species, the values of fractional surface coverage ( Oi ) and rate constants ( k ) are not single-valued and will depend on the adsorptive properties of each discrete site for t~e three species. In general, the adsorptive properties of any given site toward A, B or C are distinct and independent of the adsorptive properties of that site for another molecule. These complexities have been previously addressed [16]. Briefly, we define the adsorptive property of interest to be the heat of adsorption for each species on any given site, denoted as q A, q B and q c for species A, B and C, respectively. Minimum and maximum values of the heat of adsorption must be defined; these are easily determined by a variety of experimental techniques, including microcalorimetry [19, 20]. Second, a distribution function, o, which describes the density of sites having a given value of the heat of adsorption, is defined for each species, denoted by OA, CYBand o c in the present example. These distributions can be

344 experimentally obtained from differential microcalorimetry experiments [19, 20] or other techniques. It is important to note that it is common for the energies, distributions and adsorptive capacities to be different for each chemical species, and therefore, the methodology developed allows for this most general of cases. Finally, if all adsorbed molecules are assumed to be capable of interaction with species adsorbed on all other sites (i.e., no surface geometric constraints [21, 22]), the rate of each elementary reaction on any given site can be written in terms of a set of rate constants and surface coverages which both depend on the heats of adsorption, q A, q8 and q c. These elementary step rates are then incorporated into traditional steady-state and transient microkinetic analyses as described elsewhere [2, 16]. Much in the same way that the result of a laboratory experiment is a weighted-average of the rate of reaction occumng on all distinct sites, the nonuniform surface model performs a summation (or integration for continuous distributions) over all different sites and weights the rate on each distinct site by its relative population. The mathematical complexity for reactions over nonuniform surfaces does not generally allow a compact analytical expression to be obtained, even for this simple reaction scheme. However, an accurate estimate of the integrals involved can be obtained using numerical integration. Because of its simplicity, accuracy and flexibility, Monte Carlo integration (MCI) was chosen as the numerical integration technique. MCI is considered the technique of choice for dealing with such complicated functionals [23]; moreover, it is easily extended such that any number of adsorbed species can be easily considered in more complex reaction systems [16]. One major obstacle in establishing the impact of surface nonuniformity on observable reaction kinetics remains: no chemical reaction is sufficiently well studied to provide all of the necessary fundamental information. For this reason, we choose to address the role of surface nonuniformity utilizing generated "data" for the reaction sequence in Scheme I. By precisely defining all aspects of the chemical reaction to be studied, several potential sources of error, which may tend to bias the analysis including mechanistic interpretations, kinetic parameters, the presence of masking phenomena (e.g., diffusional limitations), are eliminated. Generation of the nonuniform surface kinetics data proceeds by defining a mechanism (i.e., Scheme I) and the kinetic and thermodynamic parameters corresponding to the occurrence of the elementary steps of the mechanism on a catalyst surface of defined site heterogeneity. Values of these parameters for Scheme I which are used throughout this paper are presented in Table 1. These fundamentals are then incorporated into a kinetic model employing the Monte Carlo integration, allowing generation of data at the desired reaction conditions. Utilizing this ability we have generated nonuniform surface TPD, TPR and steady-state reaction data for a reactive system with precisely defined kinetic and thermodynamic properties, simulating the approach an experimental study would take to probe the kinetics and mechanism. We have used the mechanism given in Scheme I and the properties described in Table 1 in the generation of the pseudo-experimental data. These kinetics data were then analyzed assuming the reaction was carried out over a uniform surface, thus disregarding the built-in nonuniformity. These uniform surface analyses were camed out by adjusting the kinetic and variable thermodynamic parameters (defined thermodynamic parameters such as the equilibrium constant, heat of reaction, etc., were held constant as they are typically known values) until the uniform surface model prediction of the transient and/or steady-state data was optimized. In this manner, the impact of surface nonuniformity on observable kinetics can be easily measured by the ability of a uniform surface model to predict the nonuniform data.

345

Table 1. Parameters used in the generation of nonuniform surface data for the mechanism presented in Scheme I. A ,

Reaction

T o r t -I S-1 o r s -~

A ...................................

B

8000

Adsorption

C

Adsorption i ......................................

Adsorption Desorption

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

kcal

o~

kcal mol ~

mol l

0.0

0.0

i ......................................

8000 1.0

X 1 0 ]l

,o ........................................

8000 1.0 x 1013 ,, ........................................

0.0 0.0 4, ..............................

0.0 0.0 ,, ..............................

0.0

mol l

12.5 - 17.5

0.0

12.5 - 17.5

-5.0

12.5- 17.5

1.0 ,t ...............................................

0.5

Reverse 1.0 x 10 ]~ Parameter Volumetric flow rate Number of sites Heating rate Inlet pressure of reactant A Inlet pressure of reactant B Total pressure Steady-state reaction temperature

16.0

0.5

Forward

0.0

1.0

16.0

reaction

0.0

q i '

kcal

,r ................

1.0 x 1013

Surface

AH/~,

.............. !....0..x...!..0.'..'. ................... .0:.0 ................. .1....0. .................................................................................

Desorption ............................

Eo,

, ..............................................

Value 30 cc min 1

2 x 10 Iv 0.333 K s-1 5 - 100 Torr 10 Tort 760 Torr 275 - 375 K

4. K I N E T I C F I N G E R P R I N T S OF N O N U N I F O R M SURFACES" SIMULATION RESULTS After generation of the nonuniform surface data (as discussed above), a uniform surface model (i.e., with only one adsorption enthalpy parameter per adsorbed species) was constructed and fitted ignoring the surface nonuniformity. In the first phase of the fitting, only the temperature programmed desorption and reaction data were utilized in estimating the kinetic parameter values, as described in Section 4.1. The resultant parameters were then used to predict the nonuniform surface steady-state experimental data as demonstrated in Section 4.2. The resulting prediction lacked sufficient accuracy, and further adjustments to the fitting process using all data simultaneously to estimate the kinetic parameters failed to achieve a compromise between transient and steady-state predictions. 4.1

Temperature programmed desorption and reaction The temperature programmed desorption of A from a nonuniform surface with defined kinetic and thermodynamic parameters (see Table 1) was generated; the resulting curve is presented as the solid line in Figure 1. The peak which describes the evolution of reactant A has a maximum height at 370 K and shows significant broadening, with a full-width at halfmaximum of 130 K. It has been previously demonstrated that this curve shape and broadening are characteristic of first-order desorption from a nonuniform surface with successive readsorption events, as would be expected in TPD experiments from high surface area catalysts

346 [16]. Ignoring this broadening and performing a standard Redhead analysis [24] on the desorption curve results in desorption energy estimates of 21.3.22.9 and 24.6 kcal mol 1 using values of 10 ~1, 1012 and 1013 Sl , respectively, for the desorption preexponential factor. Note that the desorption energy estimate of 21.3 kcal mol ~ corresponding to the preexponential factor of 1011 S1 used in the nonuniform surface data generation is significantly outside the range of adsorption energies values used to generate the nonuniform surface data. 0.0030

'

'

'

'

I

'

'

'

'

Nonuniform .....

0.0025

r

Uniform

I

'

'

'

'

I

'

'

'

'

l

'

'

'

'

surface model

prediction

0.0020

,4~ C~

rrr

0.0015

0 ,4~

D.

'-

0 U} (9

~

0 0" 0 1 0 0.0005 0"0000

tl

t

,"

l"

100

. . . .

~

200

. . . .

i

. . . .

I

. . . .

300 400 Temperature, K

,

500

. . . .

600

Figure 1. Uniform surface model prediction of temperature programmed desorption data from a nonuniform surface for reactant A of Scheme I requires low values of desorption parameters (A d = 3.22 x 107 s ~, E d = 9.01 kcal mol 1) to capture peak breadth. Using this simple, traditional analysis, it is not possible to deconvolute how significantly readsorption and surface nonuniformity effects contribute to this lack of parameter agreement. Therefore, it is more desirable to utilize the entire information contained within the TPD curve, including the characteristic broadening. Accordingly, a first-order desorption microkinetic analysis [2] employing a uniform surface was also performed; the best fit result of this analysis is illustrated by the dashed curve in Figure 1. In this fitting, the preexponential factor for readsorption was fixed at 8x 103 Torr ~ s ~, and the preexponential factor and activation energy for desorption were optimized. This approach was selected since allowing both preexponential factors to vary would result in an infinite number of optimal solutions as the TPD experiment (allowing readsorption) probes only the temperature dependence of the equilibrium constant for the adsorption/desorption process rather than the individual rate constants. The values obtained from this analysis were A d = 3.22 x 10 v s ~ and E d = 9.01 kcal mol 1. These values are both much lower than anticipated. The preexponential factor is nearly four orders of magnitude lower than the value used in generating the nonuniform surface data, while the optimum adsorption energy is outside the range of values used in data generation. Since the adsorption and desorption preexponential factors are linked, the value of both preexponential factors could be raised to obtain a value closer to that used in the nonuniform surface generation, 10 ~ s ~. However, the value of the adsorption preexponential factor in the optimum fit is equal to the value used in the data generation (see Table 1) and raising it would quickly approach the theoretical limit defined by the collision frequency (8 x 105 Torr 1 s ~ for a

347 gaseous species of molecular weight of 1 g mol 1 at ambient conditions [25]). Similarly, the peak maximum could be maintained at the same position by simultaneously decreasing or increasing both values; however, these optimized values were required to provide the proper peak breadth. Since the kinetic and thermodynamic parameters for adsorption and desorption of A and B were identical in the generation of the nonuniform surface data (see Table 1), all of the above arguments for TPD of adsorbed A can be stated identically for TPD of adsorbed B. The same approach was adopted to extract representative uniform surface kinetic parameters for the desorption of product C of Scheme I from a nonuniform surface. In contrast to the TPD data for reactants A and B, however, the generated nonuniform surface data revealed that significant quantities of A and B were formed by the decomposition of adsorbed C as the temperature was ramped. The presence of a nonnegligible surface reaction rate required that both surface reaction and desorption parameters for C were optimized simultaneously to adequately describe the TPD of the product. This can most efficiently be accomplished by concurrently optimizing both the TPD results of adsorbed C as well as temperature programmed reaction (TPR) results of coadsorbed A and B, allowing the kinetics of the surface processes to be probed from each direction. Figure 2 presents the TPD of adsorbed C generated for a nonuniform surface using the kinetic and thermodynamic characteristics described in Table 1. It is evident from Figure 2 that approximately one-half of all of the C molecules which were adsorbed proceed to species A and B by decomposition of adsorbed C, while the other one-half of the adsorbed C molecules simply desorb. 0.0025

'

'

'

'

I

'

'

'

'

I

'

'

i

'

'

'

'

I

--

,

,

,

I

.,.., 0 . 0 0 1 5

r n,c 0

'

N o n u n i f o r m surface, A N o n u n i f o r m surface, C ......... Uniform model prediction, A . . . . . Uniform m o d e l prediction, C

0.0020

:,=

'

--

0.0010

I.'"

.. ~

0

0.0005

0.0000 ,

100

~

n

J n

|

"..%

/,-'"

'... %

I'lb ~ .

200

"

I

!

|

|

I

aoo

z

i

i

J

40;

z

l

,

|

500

,

,

n

600

Temperature, K Figure 2. Uniform surface model prediction of temperature programmed desorption data from a nonuniform surface for A and C of Scheme I. The TPD data presented in Figure 2 and the TPR data for coadsorbed A and B (not shown) were then concurrently optimized to provide a consistent set of uniform surface kinetic parameters which were best able to reproduce these observations. In this fitting procedure, the preexponential factors for adsorption of all species were fixed at 8 x 103 Torr r s ~ using the same logic employed above. Furthermore, the kinetic parameters for desorption of species A and B were fixed at the values obtained above from fitting of the TPD results of reactants A and B. Finally, the value of the preexponential factor for the reverse surface reaction was fixed at

348 1.0 x 1011 sl; as described above, this eliminates the possibility of an infinite number of optimal parameter sets. The values of the kinetic parameters obtained from this analysis which simultaneously describe both the TPD of adsorbed C and the TPR of coadsorbed A and B in an optimal fashion are summarized in Table 2; the four parameters optimized through this process shown in italics. As was observed above for the TPD results of A and B, these parameters are significantly lower than the values used to generate the nonuniform surface data (see Table 1). Table 2. Optimized parameters for a uniform model of desorption of product C and the surface reaction of Scheme I. All other parameters were fixed at the values in Table 1 or from previous optimizations. Reaction

A, Tort ] s~ or s~

Eo kcal tool ~

o~

q;, kcal mol -~

'

A

Adsorption Desorption

8000 2 3.22 x 107

0.0 0.0

0.0 1.0

9.013

B

Adsorption Desorption

8000 3.22 x 107

0.0 0..0

0.0 1.0

9.01

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

C .................................

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

9. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

Adsorption Desorption 9 ................................................

9..............................

8000 3.14 x 1011 t,

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

0.0 0.0 t.

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

,.

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

0.0 1.0 9. .........................

, ...............................................

13.31 | ...............................................

Surface Forward 7.47 x 10 s 9.36 0.5 reaction Reverse 1.0 x 10 l~ 9.3 6 0.5 Optimized TPD (product C) and TPR parametersare shown in bold italics. 2 Parametersfixed at the values used to generatethe nonuniform surface data are underlined. 3 Parameters fixed at values obtained from previous optimizations of uniform model parameters to predict nonuniform surface TPD data for species A and B are shown in bold. Figure 2 compares the nonuniform surface observations (bold lines) to the predictions of the uniform surface model (light lines) employing the optimized parameters of Table 2 for the TPD of a surface saturated with adsorbed C. The observed and predicted behaviors of all species show excellent qualitative and quantitative agreement. Species C is desorbed in one state, centered with a peak maximum temperature of about 230-240 K, in what appears to be a truncated first-order desorption feature. This truncation is due to the fact that decomposition of species C into both A and B is rapidly occurring concurrently with desorption. The uniform surface model predicts that approximately one-half of all adsorbed C molecules desorb intact and one-half are decomposed, which is very similar to the nonuniform surface observation. The observed split between desorption and decomposition is explained by the availability of vacant sites and the kinetics of the two competing processes. Initially, the surface is completely covered by species C, leaving no vacant sites with which C could react to decompose into species A and B via the reverse of the surface reaction of Scheme I. However, after a molecule of C is desorbed, a vacant site is created which immediately reacts with an adsorbed C molecule through the facile decomposition step. This reaction affords one molecule each of adsorbed A and B, resulting in a one-to-one split between the decomposition and desorption pathways. Desorption of species A and B due to the decomposition of species C appears in two states in both the nonuniform surface observations and the uniform surface predictions using the parameters of Table 2. Note that only species A is shown, as the desorption of species B is kinetically identical (see Table 1). Therefore, all observations made below for species A are equally valid for the adsorption and desorption of B. The high temperature state, centered at approximately 370-380 K, resembles in both shape and position the result of the TPD of species A shown above in Figure 1. This feature is ascribed to the desorption of species A from the surface with the occurrence of significant readsorption events. The low temperature feature, centered at 280 K in the nonuniform surface observation and 260 K in the uniform surface

349 model, results from the desorption of species A in which readsorption cannot occur since the surface is still fully occupied by species A, B and C. As the temperature increases, the total surface coverage decreases from unity and readsorption can occur, causing the net rate of evolution of species A to decrease between 280 and 310 K. As temperature increases further, the equilibrium between the adsorption and desorption processes favors desorption, causing the net rate of species A desorption to again increase and to have similar features to results observed for the simple TPD of species A. The unusual behavior demonstrated for the TPD of species C from a nonuniform surface can be adequately described by a uniform surface model, thus ignoring any effects of surface inhomogeneity.

4.2 Prediction of steady-state data with parameters derived from transient experiments The impact of surface nonuniformity on observable kinetics was measured further by quantifying the ability of a uniform surface model to predict the nonuniform surface data. The parameters of Table 1 were used to generate the steady-state rate over a nonuniform surface at a temperature of 325 K and inlet partial pressures of species A, PA, ranging from 5 to 100 Torr. These rates are plotted as a function of PA as the closed circles in Figure 3. The rate passes through a maximum at a partial pressure of A equal to 20 Ton-, although the rate is relatively constant for PA ranging from 10 to 40 Tort and drops off sharply outside this range. 0.0380

'

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I

'

'

'

I

'

'

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I

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'

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I

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i

'

'

'

i

'

'

1

'

9 Nonuniform surface

0

dP

0.0360 O

:6o D

0.0340

n

= 0.0320

o 0.0300

. i

C

r- 0 . 0 2 8 0 0 z 0.0260

9 "',,,

~,,

0.90 0.83

9

c m=. ,==k ,=r

3 O (1)

I I

r

E

-,,

Uniform model prediction

I

N,-!._

t._

.....

~

0.7. 0

! I !

,=r

Q.

0.6~"

.4 I I I ! ! I !

O

O.5-= |

,

,

,

!

20

,

,

,

!

40

t ,

,

I

60

,

,

,

I

,

,

,

l

,

80 100 PA ' Torr

,

,

I

,

120

,

,

I

, O ,

140

160

0.4

Figure 3. Uniform surface model prediction of steady-state rate using optimized parameters of Table 2 is a factor of 30 greater than the nonuniform surface steady-state rate. The uniform surface model predictions over these same reaction conditions were obtained by using the parameters derived from the transient analyses described in Section 4.1 and summarized in Table 2. The variation of the predicted steady-state rate with PA is presented in Figure 3 as the dashed line. The uniform surface model prediction has the same general shape as observed from the nonuniform surface, passing through a maximum in the region of 16 Ton.. However, the maximum in the rate as a function of PA is less broad, and the decrease at higher partial pressures is less pronounced. The most notable difference, however, is the magnitude of the predicted uniform surface rate. The predicted uniform surface rate is a factor of 30 higher than the nonuniform surface rate. This difference is quite striking and leads to a clear conclusion: it is not possible to use kinetic and thermodynamic parameters optimized from transient studies alone to predict steady-state behavior in the presence of surface nonuniformity.

350 A strategy employing both steady-state and transient information to optimize the parameters simultaneously failed to achieve an acceptable compromise between the transient and steadystate uniform model predictions. This inability of optimized parameters obtained from transient studies to predict observed steady-state behavior is one potential kinetic fingerprint of nonuniform surfaces. REFERENCES

1. E.R. Becker and C.J. Pereira (eds.), Computer-Aided Design of Catalysts, Marcel Dekker, Inc., New York, 1993. 2. J.A. Dumesic, D.F. Rudd, L.M. Aparicio, J.E. Rekoske and A.A. Trevino, The Microkinetics of Heterogeneous Catalysis, American Chemical Society, Washington, DC, 1993. 3. S.L. Kiperman, K.E. Kumbilieva and L.A. Petrov, Ind. Eng. Chem. Res., 28 (1989) 376. 4. I.A. Vartanov, M.S. Kharson, M.M. Kostyukovsky, V.G. Kipovich and S.L. Kiperman, Kinet. Katal., 25 (1984) 142. 5. N.V. Voikina, A.K. Avetisov, O.K. Bosdonova and S.L. Kiperman, Kinet. Katal., 1 6 (1975) 1524. 6. N.V. Voikina, A.K. Avetisov, O.K. Bosdonova and S.L. Kiperman, Kinet. Katal., 1 8 (1977) 518. 7. S.L. Kiperman, Some Problems of Chemical Kinetics in Heterogeneous Hydrogenation Catalysis, in Catalytic Hydrogenation, L. Cerveny (eds.), 1986. 8. E. Garrone, V. Bolis, B. Fubini and C. Morterra, Langmuir, 5 (1989) 892. 9. I. Langmuir, J. Am. Chem. Soc., 38 (1916) 2221. 10. V.T. Popa and E. Segal, J. Molec. Catal., 94 (1994) 47. 11.L.J. Broadbelt and J.E. Rekoske, Chem. Eng. Sci., 51 (1996) 3337. 12.A. Corma, F. Llopis, J.B. Monton and S.W. Weller, Chem. Eng. Sci., 43 (1988) 785. 13. I. Mochida and Y. Yoneda, J. Catal., 9 (1967) 386. 14. M.N. Neurock, A. Nigam, C. Libanati and M.T. Klein, Chem. Eng. Sci., 45 (1990) 2083. 15. M. Boudart, Ind. Eng. Chem. Res., 28 (1989) 379. 16. L.J. Broadbelt and J.E. Rekoske, in preparation, (1996) 17. G.F. Froment and K.B. Bischoff, Chemical Reactor Analysis and Design, Wiley, New York, 1990. 18.R. Madix, Chem. Eng. Sci., 23 (1968) 805. 19. N. Cardona-Martinez and J.A. Dumesic, Adv. Catal., 38 (1992) 149. 20. P.C. Gravelle, Adv. Catal., 22 (1972) 191. 21. Y.K. Tovbin and O.V. Chelnokova, Russ. J. Phys. Chem., 62 (1988) 349. 22. Y.K. Tovbin and O.A. Kuznetsova, Russ. J. Phys. Chem., 64 (1990) 224. 23. W.H. Press, B.P. Flannery, S.A. Teukolsky and W.T. Vetterling, Numerical Recipes in C. The Art of Scientific Computing, Cambridge University Press, Cambridge, 1988. 24. P.A. Redhead, Vacuum, 12 (1962) 203. 25. M. Boudart and G. Djega-Mariadassou, Kinetics of Heterogeneous Catalytic Reactions, Princeton University Press, Princeton, New Jersey, 1984.

91997 Elsevier Science B. V. All rights reserved. Dynamics of Surfaces and Reaction Kinetics in Heterogeneous Catalysis G.F. Froment and K.C. Waugh, editors

351

Transient kinetics of methane dehydrogenation and aromatisation" experiments and modelling Y. Schuurman a, T. Decamp a, A. Pantazidis a, Y.-D Xu b, and C. Mirodatos a

aInstitut de Recherches sur la Catalyse, CNRS, 2 Avenue Albert Einstein, 69626 Villeurbanne Cedex, France bState Key Laboratory of Catalysis, Dalian Institute of Chemical Physics, Chinese Academy of sciences, 457 Zhongshan Road, P.O.Box 110, Dalian 116023, China

Isotopic transient experiments combined with in situ DRIFT spectroscopy and TAP experiments were used to study the methane dehydrogenation and aromatisation over a Mo/HZSM-5 catalyst. The stoichiometric reduction of both external and internal molybdenum oxide into carbide-like species was clearly observed in the early stage of the reaction. The active sites for the catalytic dimerization of methane into ethylene are in very low concentration and could be internal carbide-like clusters in tight interaction with the zeolite. The methane molecules are also strongly interacting with most of the OH groups in the zeolite via a reversible activation process leading to the OH/OD exchange. However this activation process is not directly related to the aromatisation process. A large reservoir of benzene precursors is also present in the zeolite, also strongly interacting with the OH groups. Benzene appears as a secondary product formed from ethylene. The aromatisation steps could occur directly on the acid sites of the zeolite, available after the initial reduction and clustering of the Mo ions.

1. INTRODUCTION Non-oxidative methane aromatisation has received much attention recently since it was known that methane can be activated over Mo/HZSM-5 catalysts with conversions around 6 % and selectivities to aromatics of about 70% at 973 K [1-5]. It was also recognised that the acid sites and channel structure of HZSM-5 and the state and location of Mo species are crucial factors for a good catalytic performance [6-10]. Wang et al. [5] and Solymosi et al. [11-12] suggested that during the initial induction period the original Mo+6 ions in the zeolite are reduced and carbided to Mo2C and that the methane activation occurs on Mo2C sites, leading to the formation of C2H4. It was further suggested that oxygen deficient Mo2C-MoO2 may be the active sites [ 11 ]. Since this reaction is carried out at high temperature and in the absence of oxidant, serious deactivation is inevitable due to heavy coking, causing severe difficulty for its mechanistic study. Accordingly, transient kinetic techniques which are able to provide unique information on the actual state of a working catalyst within a very short period of time [ 13,14] were applied to this complex and unstable catalytic system. Non-steady-state and steady-state isotopic transient kinetics (NSSTK and SSITK) combined with in situ diffuse reflectance infrared Fourier transformed spectroscopy (DRIFT) and temporal analysis of product (TAP) were performed in order to analyse some of the above mentioned key steps of the aromatisation process.

352 2. EXPERIMENTAL

2.1. Catalyst preparation. 3Mo/HZSM-5 (3Mo/HZ) was prepared as reported in [1, 2, 10]. 3 wt.% Mo was introduced by impregnating HZSM-5 (SiO2/A1203 mol ratio of 50) with a proper aqueous solution of ammonium heptamolybdate. It was dried at 383 K overnight, calcined in air at 773K for 6 h. The samples were pelletized, crushed and sieved before use.

2.2. NSSTK and SSITK analysis. NSSTK and SSITK experiments were performed with an atmospheric flow system using either a tubular quartz microreactor (70 mg of catalyst) or a catalytic DRIFT cell from Spectratech, allowing the gases to flow through a fixed bed of catalyst pellets (about 30 rag) and able to be heated up to 1173 K. The gas composition was continuously monitored at the reactor outlet by online mass spectrometer and the surface composition was investigated by a FT-IR spectrometer (Nicolet 550) with one spectrum recorded per second. In all cases, the catalyst was pre-treated with He at 1013 K for 40 min. The reacting feed was composed of 10 vol.% methane (12CH4, 13CH4 or ~2CD4) and 90 vol.% He with a total flow rate of 24 ml/min. The reaction was carried out at 1 atm and 1013 K. Ar was used as an inert tracer.

2.3. TAP experiments. A narrow gas pulse of reactants is introduced in a microreactor which is evacuated continuously. The response of this pulse is detected by a quadrupole mass spectrometer at the reactor exit. The shape of the response reflects diffusion, adsorption, desorption and reaction, as extensively described in [15-17]. The TAP-2 reactor system equipped in this laboratory was the first commercially sold unit by Mithra Technologies Inc. Its main features are: 1) four high-speed pulse valves able to deliver 0.25 ms FWHH pulses, 2) a 25.4 mm in length and 4.1 mm in diameter microreactor, 3) a high-throughput liquid nitrogen trapped vacuum system, 4) a slide valve containing a needle valve to switch between high-pressure experiments and vacuum experiments, 5) a quadrupole mass spectrometer located directly underneath the reactor exit, 6) a gas and liquid vapour feed panel to accurately make up blends and 7) a PC based system for data collection and pulse valve control. Typically, a catalyst charge of 25 mg with a particle size of 0.2-0.3 mm was placed in the centre of the reactor between two layers of 0.2-0.3 mm size quartz particles.

3. RESULTS

3.1. Start-up transient experiments A first transient experiment consisted in observing the catalyst behaviour during the reaction start-up at 1013 K by switching from an Ar/He to a CH4/He flow (Fig. 1A and 1B). During this initial period, at least three different reaction steps occurred in sequence: i) from 0 to 200 s, CO (not reported in Fig. 1 due to MS fragmentation uncertainties), CO2 and H20 were the main products with a slight production of H2 and C2H4, ii) from 200 to about 400 s, HE and C2I-I4 formation increased markedly, while a second maximum in CO2 and n20 production was observed, iii) after 400 s, C6H6 appeared at the reactor exit, which corresponded to a strong decrease in C2H4, CO2 and H20 production. TAP experiments were also carried out during this start-up period. The initial CH4 pulses at 973 K showed mainly a formation of CO, H20 and small quantities of CO2. Pulses of CH4 were continued until the formation of H2 and C2H4 was observed. Fig. 2A shows the height normalised responses of CH4, HE, At, C2H4, CO and CO2 from a 9/1 CH4/Ar pulse. The CH4 conversion was smaller then 5% with approximate relative selectivities towards CO, CO2 and C2H4 of 90, 10 and

353 1%, respectively. Similar amounts of H2 and CO were formed. The H2 formation appeared to be instantaneous on the CH4 pulse. Probably due to the low CH4 conversion and the small quantities of CH4 pulsed in the TAP reactor, the formation of benzene could not be observed. Instead, the more reactive ethylene was introduced as reactant. Fig. 2B shows the height normalised responses of C2H4, H2, At, CO, CO2 and C6I-I6 from a 9/1 C2H4/Ar pulse. The C2I-I4 conversion amounts to 94% with approximate relative selectivities towards CO of 30%. However, due to the slow desorption of CO a continuous CO signal was observed which could cause an underestimation of the CO production. Except a similar amount of H2 formed as that for CO, all other products were formed in negligible amounts. Benzene, however, was detected on the C2I-h pulse. The relative slow response of benzene reflects its slow formation and strong interaction with the zeolite.

1

1 i

O.8

0.8

e-. 0.6 ..,.. "o

~ o.s

0.4

~ o.4

In

N

E i,.

O Z

ell)

"

'"

Ar

'

t

"0

0.2

~0"2 t 0,

200

0

Time (s)

400

600

9

-~----

I

200

0

t

400 Time

600

(s)

Fig. 1 Normalised responses for (A) Ar, CH4, CO2, and H20 and (B) Ar, H2, C2H4,and C6I-I6 after switching Ar/He ~ CHVHe at 1013 K over 3Mo/HZSM-5.

~o.,

!o.,

~o.,

_~o.,

0.2

0.2

0.0

~ 0.0

0.0

0.1

0.2 Time (s)

0.3

0.0

0.1

0.2 Time

0.3

(s)

Fig. 2, Normalised responses from a CH4/Ar pulse (A) and C2H,/Ar pulse (B) over 3Mo/HZSM-5 at 973 K. CI-L (m/e=l 5); H2 (m/e = 2); Ar (m/e=40); C2H4 (m/e=26); CO (m/e=28); CO2 (m/e= 44) and C6H6 (m/e=78).

354 3.2. Hydrogen tracing from CH4/CD4 SSITKA CH4/CD4 SSITKA experiments were carried out at 1013 K after 40 rain on stream in the DRIFT cell. Fig. 3A shows the spectra obtained a) initially under CH4 flow and b) under CD4 flow, 5 min after the switch when the spectrum was not changing anymore significantly. A very large part ofthe OH bands at 3715 cm1 and 3583 cm-1 was shifted to OD bands at 2737 and 2641 cm~ when the deuterated methane (C-D bands at 2257 cm]) was substituted to light methane (C-H bands at 3061 cm~). Note that C-H and C-D vibrations of adsorbed carbonaceous species could be hidden under the relatively large bands of the methane gas phase. Fig. 3A shows the norrnalised change with time of the O-H/O-D and C-H/C-D band intensities, obtained from the series of intermediate spectra (around 300) recorded during the 5 rain transient experiment. A clear delay is observed between the C-H/C-D and the O-H/O-D transient curves.

6o

1 m m = o 0.8 "o m u 0.6 o. o "o

9

w

3716r

27'37r

|

,q-------

_~

0.4

o-o

o.~

m

T

E o9- 0 . 2 z

2S41.oln"~

O, .~

~

wavenumber

:~

(cm-1)

1

0

:~

2 3 Time (rain)

4

5

Fig. 3 DRIFT spectra (A) and normalised responses (B) of hydroxyl groups (OH/OD) and gas phase methane (CH/CD) after the isotopic switch CHdHe --+ CDd/-Ie at 1013 K over 3Mo/HZSM5.

Q-C6D 6 ~>'0.8

-->'-0.8 w e-

/ C6HD5

o

_co.s

~

E L._

~ '

9

1

C6H 6

~

/

C6H4D2

/

~o.s

"o N ~0.4

f

0.4

E

~0.2

Z 0.2

1

0

100

Time (s)

200

1

J

300

0

100

Time (s)

200

300

Fig. 4 Normalised isotopic composition of methane (A) and benzene (b) after the switch CHa/He CD4/He at 1013 K over 3Mo/HZSM-5.

355 The changes in the isotopic composition of the gaseous methane at the reactor outlet was also monitored by MS analysis as reported in Fig. 4A. A fast H/D exchange between the light and deuterated methane was attested by the transient formation of CH3D, CH2D2 and CHD3 species during the CH4/CD4 switch. Note that the partly deuterated methane appear and disappear sequentially, which is characteristic of a step-wise exchange process [ 18]. Besides the H/D methane exchange, a similar but slightly slower process of H/D exchange was simultaneously observed for the formed benzene (Fig. 4B). All the intermediate isotopes were sequentially observed during the transient, again indicating a step-wise exchange process.

3.3. Carbon tracing from 1 2 C ] ~ 4 / 1 3 C ] ~ [ 4 SSITKA By switching from 13CH4/He to 12CH4/He at the reactor inlet, a delay of around 5s was observed between the inert tracer transient curve and the methane curves (Fig. 5A). This indicates that an accumulation of reversible methane occurred under the reaction conditions. From the flow rate of non converted methane, this accumulation could be estimated to around 180 lamol/goat. A slight delay was also observed for ethylene, which corresponded to a much lower accumulation of around 1.6 lamol of ethylene precursors/gc, t. In contrast, much larger delays were observed for the transient curves of the various labelled benzene (Fig. 5B). This would correspond to a pool of carbon intermediates leading to benzene formation larger than 125 [amol /go.at. NO precise quantification was achieved due to the impossibility to reach a steady-state isotopic composition, even after 10 min on stream. Thus, before the~3CHj12CI-h switch, i.e. after the initial 1 2 C H 4 / 1 3 C I - - I 4 switch, a non negligible fraction of ~3C5~2CH6 was still observed together with 1 3 C 6 H 6 under 13CH4 atmosphere. Inversely, long after the ~3CHV12Ct-h switch, ~2C5~3CH6was still present at the reactor exit with 1 2 C 6 H 6. Note also a non sequential labelling of the benzene molecules since ~2C6H6 arose before ]2C5~3C6H6, for example. This underlines the complexity of the aromatisation process.

1

,

:>0.8

>,0.8

To.6

"& o.6 m

N

N ,.,.. "~ 0.4

~

0.4

E

o z 0.2 0

I 0

100

200

I

,

300

Time (s)

400

,

I 500

,

600

0

100

200

300

400

500

600

Time (s)

Fig. 5 Normalised isotopic composition of methane (A) and benzene (B) after the switch ~3CHAHe -~ 12CHdHe at 1013 K over 3Mo/HZ.

356 4. D I S C U S S I O N

4.1. Initial reduction of Mo species by C]][4 The reduction of Mo species during the initial start-up period at 1013 K is clearly revealed by the production of COx and H20 either under continuous flow (Fig. 1) or pulses (Fig. 2) of methane. The CO2 and H20 peaks were integrated in Fig. 1 and the corresponding mole production compared to the initial MoO3 content. Thus, for a catalyst loading of 50 mg corresponding to 15.6 ktmol of MoO3, the amount of produced CO2 and H~O was 4 and 32 ktmol, respectively. Within the uncertainties of the calculation (including a rough estimation of the produced CO), these data were found in reasonable agreement with the two reduction equations : 4 MoOs + 5 CH4 -~ 4 MoC + CO2 + 10 H20

(1)

2 MoO3 + 2 CH4 --~ Mo2C + CO2 + 4 H20

(2)

Thus, this start-up transient experiment suggests that the stoichiometric reduction of a large amount of the initial MoO3 into molybdenum carbide MoC and/or MozC is a prerequisite for producing the target aromatic products. The formation of carbide species under methane flow at high temperature was also recently proposed by Wang et al. on the basis of XPS measurements [5]. From the observation of two reduction steps, it could additionally be speculated that the first step corresponded to the reduction of the Mo phase located outside the zeolite, while the internal phase, less accessible and probably more dispersed (stabilised ions), would be reduced more slowly. The existence of external deposits was also deduced from chemical extraction experiments [7,8]. The occupation of internal sites by Mo cations was recently deduced from infrared experiments [ 10]. Within this scheme, it can be noted that the formation of ethylene, residual after the external deposits reduction, was boosted during the reduction of the internal deposits. In contrast the benzene formation started only when the reduction of the internal deposits was completed. Note that the latter observation strongly suggests that ethylene is an intermediate product between methane and aromatics. This statement also agrees with the observation of benzene formation after ethylene pulses in the TAP reactor (Fig. 2B). Further information on this initial reduction period is also provided by the modelling of the TAP data. In TAP experiments the pulse intensities were kept sufficiently small to insure Knudsen flow. The diffusivities for all components in the inert zone were calculated from the argon response over a bed of 0.2-0.3 mm quartz particles. The argon response shown in Fig. 2A was then fitted with the Ar diffusivity in the catalyst zone as the only parameter. From this value the diffusivity for the reactants and products were calculated. These responses were then modelled by trying different reaction mechanisms with only the kinetic constants as fitting parameters. The method used is similar as described by Svoboda et al. [ 16], except that the current one is based on three different zones in the reactor. This method solves the corresponding set of partial differential equations by a Laplace transformation with respect to time. The resulting set of ordinary differential equations is, after integration, transformed back to the time domain by an inverse fast Fourier transform algorithm [17]. The CO and CO2 responses were modelled on the following simplified reaction scheme:

CH4 Tl CHx

+MoO

~

CO

CO2

~'~' CO* +

'~1 CO2"

(3)

357 Methane is irreversibly adsorbed under the form of carbonaceous species and hydrogen. Subsequently, the carbonaceous species react instantaneously with the molybdenum oxide to form CO and CO~, followed by desorption and readsorption of the carbon oxides. The model shows that the formation of CO and CO2 occurs in parallel. A model accounting for the formation of CO2 as an oxidation product of CO failed to describe the TAP response curves of CO and CO2. In agreement with the previous statement, these stoichiometric steps generate reduced Mo sites which serve for the further catalytic activation of methane. Note that the pulse simulation does not specify the final form of the reduced molybdenum oxide species.

4.2. State of the working catalyst The state of the working catalyst, i.e., after the initial stoichiometric reduction step, can be described by means of both H and C tracing experiments. Interaction between the zeolite OH groups and the reacting species. H-tracing transient experiments on hydroxyl groups revealed that a strong interaction established between gaseous methane and zeolite hydroxyl groups in both the external and the internal surface (Figs. 3, 4). This would correspond to a reversible and step-wise methane activation including C-H bond breakage (most likely heterolytic) and hydrogen exchange with the zeolite protons, according to : CD4 + H- .~ ~" CD3H + D"

(4)

However, such an exchange was also observed on unpromoted HZSM-5, in the absence of molybdenum [19]. Furthermore, the H-D exchange in hydroxyl groups on 3Mo/HZSM-5 was found slower than that on the unpromoted HZSM-5. This shows that the introduction of Mo species does inhibit the reactivity of the hydroxyl groups for the exchange process. It can therefore be concluded that the reversible methane activation leading to the HfD exchange cannot be considered as part of the irreversible process leading to ethylene and aromatics formation. The normalised response curves in Fig. 3B were modelled by considering a combination of CSTR pools in parallel or in series, according to the mathematical formalism developed by Soong et al. [20] or Happel et al. [21] (scheme 1). The inlet function of the system was first obtained by modelling the response of argon as an inert tracer (curve a in Fig. 6). Then it was found that a model of two pools in parallel (Model A in scheme 1) described correctly the methane gas phase transient (curve b in Fig. 6), while a model of two pools in series (model B in scheme 1) was more adapted to the response of surface OH groups (curve b in Fig. 6). The calculated curves are reported as solid lines in Fig. 6. The capacity of the pools involved in the two models was calculated from the time constant obtained after curve fitting. Thus, the pool 1 in Model A corresponded to a volume of gaseous methane of 342 ~moL/gc~tand the pool 2 in Model B corresponded to a concentration of exchangeable OH groups of 571 pmo[/gc~t. The latter value has to be compared to the maximum content of OH groups for a Si/A1 = 25 H-ZSM, i.e. 642 ~t mol/gcat. This quantitative approach of the H/D exchange between methane and the zeolite hydroxyls groups therefore indicates that about 90% of the hydroxyl groups content of the working catalyst are exchangeable into OD by interaction with the methane molecules. The same conclusion holds for the benzene molecules, as deduced from the similar exchange process observed in Figs. A and B. Let us recall, however, that this type of reversible interaction would not belong to the overall aromatisation process.

358 Model A

F(t)

u(t)

inlet function

u (t) - exp ( - at) / 1 + exp ( - at) it

outlet function 9 V

F(t) = [Aexp(- t/1:~ ) + ( 1 - A ) e x p ( - t / ~ )]u(t)

'

N-/I

A - (1-x~ / z~)/'c~(1/z~ - 1/x~)

Model B

inlet function "

u (t) = exp (-at)/1 + exp (-at) F(t)

u(t)

Ill

outlet f u n c t i o n

F(t) = ['1~1/('1~1. "1~2) exp(-t/zl) - x2 / (zl- z2) exp(-t/z2) ] u(t)

Accumulation o f reacting carbon species. From the carbon labelling experiments, it was first deduced that a rather large amount of methane (180 lamol/g~) was reversibly adsorbed in the zeolite, which may correspond to the reversible interaction with zeolite protons, as above quantified. In contrast, a very small amount of ethylene precursors was detected (1.6 ~tmol/g~), which indicates that the actual concentration of active sites for the initial methane dimerization is very low. The clustering of the initially dispersed Mo ions into a small number of carbide-like particles within some of the zeolite cages during the start-up period could explain this statement. This would also explain why the appearance of ethylene coincides with the reduction of the internal Mo deposits (Fig. 1). In this respect, the active sites for methane activation could be described as small carbide-like clusters in tight interaction with the zeolite, in agreement with the hypothesis proposed by Solymosi et al. of Mo2C-MoO2 units with an oxygen deficiency [ 11 ].

359

m

c

.9

0.8

"

0.6

-

0.4

-

0.2

-

t'~ t...

c 0 "o N

E

0 Z

0

i

7_

I

0

1

0

t

~

2

i

!

I

3

e

!

4

5

Time (rain)

Fig. 6 Normalised experimental (symbols) and simulated (solid lines) responses of a) argon tracer, b) gas phase methane (CD bands) and c) hydroxyl groups (OD bands) after the isotopic switch CH4/He --~ CD4/He at 1013 K over 3Mo/HZSM-5. A large accumulation of benzene precursors was also present in the zeolite under steady-state conditions (> 125 l.tmol/g). This accumulation may correspond to the storage of benzene precursors on the zeolite acid sites leading to the H/D exchange. It could also come from the existence of the carbon pool constituted by the molybdenum carbide deposits, slowly exchanging with the active sites for methane dimerisation. The appearance of benzene aiter the ethylene formation during the start-up period and the marked difference between the ethylene and benzene precursors concentration strongly suggests that benzene is a secondary product formed from ethylene. The aromatisation steps could occur directly on the acid sites of the zeolite, rendered available after the initial reduction and clustering of the Mo ions. It was thus checked that the aromatisation of ethylene also easily occurred on unpromoted HZSM-5. Thus, the main steps of the aromatisation process can be schematised as : 1.

irreversible dimerisation of methane over Mo clusters in the zeolite. Mo2C/MoOZ ~ 2 H2 + C2I-h

2 CI-h

(5)

2. aromatisation on zeolite protons. W 3 C2H4

~- CrH6 + 3 H2

(6)

3. methane and benzene exchange with zeolite protons as side reactions. CH4 +

W

-~ CH5+

(7)

360 5. CONCLUSIONS Though the exact nature of the active sites in the aromatisation process remains far to be understood, the use of various transient kinetic techniques allowed us to get a more precise picture of this highly complex process. Thus, the following features can be outlined : - The stoichiometric reduction of both external and internal molybdenum oxide phases in the early stage of the reaction was clearly observed with the formation of CO/CO2 and H20, accompanied with the formation of C2H4 and further on of CrHr, either by NSSTK in flow mode or by TAP experiments in pulse mode. The modelling of the TAP experiments showed that CO and CO2 are formed simultaneously. - The specific reduction of the internal deposits coincides with the ethylene production. It is then suggested that the active sites for the catalytic dimerization of methane are carbide-like species in close interaction with the zeolite. The concentration of these sites would be very low compared to the concentration of zeolite acid OH groups - The methane molecules would also be strongly interacting with most of the OH groups in the zeolite, via a reversible activation process leading to the OH/OD exchange. However this activation process would not directly be related to the aromatisation process. - A large reservoir of benzene precursors is also present in the zeolite under steady-state conditions, also strongly interacting with the zeolite OH groups. From the sequence observed during the start-up period, benzene appears as a secondary product formed from ethylene on the acid sites of the zeolite. The aromatisation process can therefore be described as bifunctional, as proposed for the aromatisation of higher hydrocarbons on close systems such as Ga/ZSM-5 [22].

ACKNOWLEDGEMENT

Thanks are due to CNRS and the CAS/CNRS agreements for the stay of Y. Xu at the IRC. REFERENCES .

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

L. Wang, L. Tao, M. Xia, G. Xu, J. Huang and Y. Xu, Catal. Lett., 21 (1993) 35. Y. Xu, S. Liu, L. Wang, M. Xia and X. Guo, Catal. Lett., 30 (1995) 135. F. Solymosi, A. Erdohelyi and A. Szoke, Catal. Lett., 32 (1995) 43. L. Chert, L. Lin, Z. Xu, X. Li and T. Zhang, J. Catal., 157 (1995) 190. D. Wang, J. H. Lunsford and M. P. Rosynek, Topics in Catal., 3 (1996) 289. L. Wang, Y. Xu, M. Xie, S. Liu & L. Tao, Stud. Surf. Sci.& Catal., 94 (1995) 495. Y. Xu, Y. Shu, L. Wang, M. Xie and X. Guo, Catal.Lett., 35 (1995) 233. S. T. Wong, Y. Xu, X. Guo Appl. Catal., A 136 (1996) 7. S. T. Wong, Y. Xu, L. Wang, M. Xie, Catal. Lett., 38 (1996) 39. Y. Xu, W. Liu, S.-T. Wong, L. Wang and X. Guo, Catal. Lett., 40 (1996) 207. F. Solymosi, A. Szoke and J. Cserenyi, Catal. Lett., 39 (1996) 157. A. Szoke and F. Solymosi, Appl. Catal., A, 142 (1996) 361. J. Happel, "Isotopic Assessment of Heterogeneous Catalysis" Academic Press Inc., New York 1986. C. Mirodatos, Catal. Today, 9 (1991) 83. J. T. Gleaves, J. R. Ebner and T. C. Kuechler, Catal.Rev.-Sci.Eng., 30 (1988) 49. G. D. Svoboda, J. T. Gleaves and P. L. Mills, Ind. Eng. Chem. Res. 31, (1992) 19. J. T. Gleaves, G. S. Yablonskii, P. Phanawadee and Y. Schuurman, Appl. Catal, in press. C. Mirodatos, A. Holmen, R. Mariscal, and G.A. Martin, Catal. Today 6 (1990) 601. unpublished results. Y. Soong, K. Krishna, and P. Biloen, J. Catal. 97 (1986) 330. J. Happel, E. Walter, and Y. Lecourtier, J. Catal., 123 (1990) 12. P. Meriaudeau, and C. Naccache, J. Catal., 157 (1995) 283.

91997 Elsevier Science B. V. All rights reserved. Dynamics of Surfaces and Reaction Kinetics in Heterogeneous Catalysis G.F. Frornent and K.C. Waugh, editors

361

The d e s o r p t i o n of C O 2 f r o m the surface as a k i n e t i c a l l y r e l e v a n t s t e p in the C O o x i d a t i o n r e a c t i o n o v e r p l a t i n u m T.A. Nijhuis, M. Makkee, A.D. van Langeveld, and J.A. Moulijn Delft University of Technology, Department of Chemical Engineering, Section Industrial Catalysis, Julianalaan 136, 2628 BL Delft, The Netherlands

ABSTRACT

The platinum catalyzed CO oxidation of carbon monoxide was studied using an advanced transient reactor system, referred to as Multitrack. The experiments indicate, that the reaction is taking place according to the Langmuir-Hinshelwood reaction model. The desorption of CO: from the catalyst surface was shown to be a kinetically relevant step in the reaction. From experiments performed using 1sO2 it was shown that isotopic mixing occurs on the catalyst surface due to the decomposition of CO 2present on the catalyst surface.

1. INTRODUCTION The oxidation of carbon monoxide over platinum was studied as a model reaction to evaluate the performance of the new Multitrack system. Multitrack, developed in our laboratories, is an advanced version of a TAP (Temporal Analysis of Products) apparatus [1]. The advantages of using this reaction as a model reaction are that the reaction is fast, the only possible reported by-product is carbon, and that the reaction has been studied using conventional TAP systems [2,3]. A review on this reaction has been made by Engel and Ertl [4]. In this review they list the following possible relevant reaction steps:

k1

CO < k2 ' ) GOad

(1)

20ad Oad+COa~ k~ > CO 2

(2)

Oad -4-C O

(4)

O 2~

0 2+2COaa

ks ) C O 2

k~ ,~2CO 2

(3) (5)

362 These reactions are based on the following assumptions: only dissociative adsorption of oxygen (2), no recombination + desorption of oxygen, and instantaneous desorption of the carbon dioxide produced (3)(4)(5). In the temperature range of this study (300-650 K) these assumptions were found to be valid by Engel and Ertl. Of the equations given reaction (3) represents a reaction according to the Langmuir-Hinshelwood model and reactions (4) and (5) represent reactions according to an Eley-Rideal model.

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

The CO-oxidation experiments were performed in the Multitrack system. A schematical representation of the Multitrack system is given in figure 1.

"~--~m I

Pulse out, 0.2 s

MS 1 MS 3 Analysis section Intermediate chamber Reactor chamber

//

Time

.--- -" " ""

MS 2

~

%9

",\ . . -.... ~_ Pulse

1 ~.~j Flow

Pulse 2

Time

Figure 1. Schematical representation of the Multitrack system. Shown is the basic layout of the system, together with the typical gas pulse shapes as they are given to the reactor and measured at the analysis section.

The basic system of Multitrack consists of three vacuum chambers. The first contains a small reactor, which can be operated from 300 to 1273 K. The reactor used was 7 m m i.d. with a total bed height of 10 mm. The pressure in this vacuum chamber is about 107 mbar. To the reactor gas pulses can be given using high speed gas pulsing valves, with a pulse width of about 100 ]as. For the experiments described here gas pulses of 1016to 1017molecules were applied. Using a flow valve it

363 is possible to perform continuous flow experiments. The flow valve was used here only to pretreat the catalyst. The second vacuum chamber is an intermediate (differential pumping) chamber. This chamber is used to prevent pressure increases in the reactor chamber from influencing the signal in the analysis chamber. The third section is the analysis section, containing three quadrupole mass spectrometers inline. These mass spectrometers are used to measure up to three components simultaneously as they leave the reactor exit. The system is constructed in such a way that only molecules traveling directly from reactor exit to mass spectrometers are measured (molecular beam). The maximum possible sampling frequency for the mass spectrometers is I MHz. The catalyst used for the CO-oxidation experiments was platinum sponge of 99.99 % purity, supplied by Johnson Matthey. The advantage of using platinum sponge as a catalyst is the macroporous structure, resulting in the absence of pore diffusion limitations, while still having a relatively high surface area for the reaction without the need for using a catalyst support material. The platinum sponge consisted of 2 lam spheres, sintered together to form 40-100 l~m particles. The amount of catalyst used for the experiments was 291 mg. The platinum was diluted in the reactor using 628 mg of 212-250 lain SiC (carborundum) particles resulting in a total bed height of 10 mm. Prior to use the platinum was pretreated in-situ by first oxidation at 773 K, and then reduction at 673 K. The platinum was analyzed using SEM and krypton BET. The surface area calculated from the SEM micrographs was 0.070 m2/g, the krypton BET gave a 0.081 m2/g surface area. Two basic types of experiments were performed. The first type of experiments are of the multiple-pulse type. In this experiment the catalyst was either first pre-covered with oxygen or carbon monoxide. After this the catalyst was subjected to carbon monoxide pulses or oxygen pulses, respectively. By determining the amount of carbon dioxide produced and the amount of CO or 02 consumed the active surface area of the catalyst can be calculated. The second type of experiments are of the single-pulse type. In this experiment type a single gas pulses are given to the reactor. This type of experiment can also be done by alternately pulsing two different gases. By measuring the individual pulse responses at a high sampling frequency and then modeling these pulse responses, information on the reaction mechanism can be obtained. This type of experiment was done using a number of different catalysts surface coverages.

3. RESULTS 3.1. Multiple-pulse experiments Multiple pulse experiments were done in the temperature range of 328 to 423 K. Figure 2A shows the result of one of these experiments as measured directly. As it is difficult to see small variations in pulse size, a different representation for the experiments was used, in which the individual pulse responses are integrated. Figure 2B shows the result of such an integration. In these figures the amount of CO 2

364 A

i

i

, ! i ico2 i! i i l~ i!i . . . . . . . . . "-:-'"

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

"---l-

--'

0

"

--~

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

r . . . . . .

5

'

"-"~- .....

T .................

10

'

-

- 1. . . . . . . .

15

"=-:---':--

"---

'-I 0 2

20

25

T i m e (s)

co

/

/

m

JL ,

~

5

10

:

-"

:

:

,t

,

:

--

~

~.;

:

~

;

~

. . . .

.

15 20 25 30 CO pulse over initially 0 covered Pt (-)

:

:

r..-

~

35

.-

:

;

v

T

40

Figure 2. Results from multiple-pulse experiment. Pulsing 1.0"1017 molecules of CO over oxygen-precovered platinum (333 K): A. as measured directly on Multitrack;. B. alternate representation of same experiment.

produced can be obtained by adding together the individual integrations, along with the amount of used oxygen or carbon monoxide via a similar addition and subtracting the value from the total amount of the gas pulsed. The results of these calculations are given in Table 1.

365 Table 1 Results of series of multiple-pulse experiments over 291 mg Pt sponge. The values given are the total amounts for the entire multiple-pulse experiment. (accuracy +/- 10 %). O: multiple-pulse (CO covered Pt) O, used CO 2 produced ~ (10 ~7molecules) (10 ~' molecules) 328 1.6"10 .2 333 2.6 2.9 338 2.3 2.5 343 3.1 2.6 373 2.1 2.3 393 2.2 2.9 423 2.3 2.2 ~- equals CO adsorbed at start of 02 pulsing. :" equals adsorbed 02 at start of CO pulsing.

T (K)

CO multiple-pulse (O a covered Pt) CO used C O 2 produced: (1017 molecules) (1017molecules) 6.1 3.1 6.1 3.2 5.8 2.9 6.4 3.2 5.9 2.9 5.7 3.0 5.0 3.2

3.2. Single-pulse experiments Although a variety of different single-pulse experiments were performed, only one situation will be discussed here, the steady-state situation of pulsing alternately CO and O: over the catalyst. To obtain additional information on the oxygen source of the CO 2 produced, this experiment was done using 1sO2and C160. Figure 3 shows the results of one of these experiments. In this figure it must be noted that a mass spectrometer measures fragmentation patterns of ionized molecules. The small peak at t=l s in both the m / e signal of 28 and 30 is in this case not caused by C160 and C1SO leaving the reactor exit, but mainly by fragments of ionized CO 2molecules.

4. DISCUSSION 4.1. Reaction model In table I can be seen that at a temperature of 328 K the pulsing of oxygen over platinum precovered with carbon monoxide does not produce carbon dioxide. This can be explained by a very strong adsorption of CO on the platinum, inhibiting the adsorption of oxygen. This is in agreement with the experiments done in the temperature range from 333 K to 393 K (not shown), where when pulsing oxygen over CO pre-covered platinum, there is initially no CO: production and the reaction only starts after some CO has desorbed. The time necessary for the desorption of a small amount of CO is shorter at a higher temperature. Apparently it is necessary to have free sites on the surface for the reaction to start, this indicates that the reaction takes place according to the Langmuir-Hinshelwood model. Pulsing CO over oxygen covered platinum does not show such an inhibition. This can either be because the reaction of CO with adsorbed oxygen can also occur

366

. . . . . .

.....

c18o2

C160180

46

t

44 C1602

~.~

........

~

.

36 180 2 __

30 C180

., 0

!

I

I

!

J

0.5

1

1.5

2

2.5

T i m e (s)

Figure 3. Steady-state results of CO oxidation with 1802 over Pt sponge at 573 K. Experiment done by pulsing 2.3"1016 molecules of 1sO2 at t=0 and 1.0"1017 molecules of CO at t=l s. At t=3.5 s the next 1802pulse is given (equals next t=0).

according to the Eley-Rideal model, or that there are always some free sites available for CO adsorption. Whichever is the case cannot be concluded from the experiments performed. 4.2. N u m b e r

of a c t i v e s i t e s

From the surface areas determined by SEM and Krypton BET the number of sites on the platinum surface can be estimated. This results in 2.7"1017 sites for SEM micrographs and 3.1"1017 sites for the BET measurement. If both the amounts of CO2produced in table I are added and then compared to the amounts of oxygen and carbon monoxide used, it can be seen, that the amount of oxygen and CO used, correspond nicely to the amount needed for the CO 2produced. The number of CO 2 molecules produced during one multiple pulse experiments should be equal to the number of active sites on the catalyst. If the amounts of produced carbon dioxide are compared to the calculated number of platinum surface

367 atoms from the SEM micrographs or B.E.T. measurement, it can be seen that these are in good agreement.

4.3. Isotopic mixing at the catalyst surface Figure 3 shows, that in the steady-state situation when pulsing C160 and 1802 sequentially, three types of CO 2 are produced 9C'60=, C'6OlSO and CXSO=,although only one type, C~60'80, might be expected. The production of the two other CO= isotopes indicates that splitting of the C-O bond takes place. This be explained by: 1. CO decomposes on the surface and produces Cads and O a~; 2. CO2~as produced decomposes to COaas and O ~ The decomposition of CO or CO= in the gas phase is very unlikely at the low reaction temperatures of the experiments and is, therefore, not taken into account. Possibility 1 is not likely to be the main cause because of: 9 no CO 2 production is seen, when only CO is pulsed over the platinum sponge 9 during the '~O2 gas pulse a significant amount of C1~O2is produced, and it is not likely that the decomposition of CO is induced by the presence of oxygen. The second possible explanation for the occurrence of the three types of CO 2 is based on the mechanism proposed in figure 4.

0 II

c

0 II

O~

c

II II ]11 ii i- i ii iiii

CO

CO

d d

II II ..................................................................

\

/

0 II

0 II

II

II

c c G d II

O II

c

"-

II

II i ii

.... R ........................................................................

to=c,o] o *

.

ZIIZIIII i][11111111111.........

.... i, , ..L..O,-,=,] c

0 II C II I[11111

cod

r,,* k . /

iiiiiiiiii

II

002

o"

o

c

o

II ,.* II C U C 0 II II II II ...................................

o [o:o:,o]

co:

Figure 4. Model for the mechanism of the CO oxidation. The shaded areas represent a platinum surface. The oxygen atoms marked with a * denote oxygen-18.

368 Adsorbed CO reacts fast and reversible with adsorbed oxygen to form adsorbed CO 2. As in the decomposition of the CO 2produced both oxygen atoms are identical, both have an equal chance to be split off. The Oad~at the surface can then be either from the CO or the originally adsorbed O atom. By this pathway it is possible to form Clso ,that can then produce C1~O2without CO dissociation. As this mechanism produces adsorbed 160 from adsorbed CX60 during the reaction between C~60 and 'sOa~s, it should also be possible to produce C~O2. The C~BO2is produced simultaneously with both other types of CO,. As C~802 can only be produced by a Langmuir-Hinshelwood model (i.e. reaction only at the surface), it is most likely, that all CO 2produced from the reaction between CO and adsorbed O atoms is produced in a similar manner. To verify the assumption in the proposed mechanism in figure 4, that the isotopic mixing of the carbon dioxide is only caused by the carbon dioxide produced on the surface and not by carbon dioxide re-adsorbing on the catalyst and subsequently reacting with oxygen, an experiment was performed, in which C1~O2and ~sO2were pulsed simultaneously over the catalyst at 523 K. This experiment showed that only very small amounts of C~601~Owere produced (less than 0.1%) and showed no detectable amounts of ClSO2, indicating that the CO2 adsorption on the catalyst is negligible. The ratio between the produced amounts of C ~602, C 16O ~sO, and Clso2 is about 2:4:1 (573K) or 2:6:1 (373 K). If the isotopic scrambling of the CO 2 is very fast compared to the desorption of CO 2, the ratio C~60 9C160~sO 9C~SO2should be I 92 91 on statistical considerations. If the desorption and isotopic mixing of CO, occur at the same rate, it is easily calculated that the ratio C~60:" C160~sO 9C1802 should be I 96 91. From the observed ratios we conclude that the rates are of the same order of magnitude. The slight excess of 160 in these ratios can be explained by very small leaks m the vacuum system and by ~60 present in the ~sO2used. The isotopic mixing of the CO, produced was also found by Huinink [3], who performed comparable experiments on a conventional TAP system with CO and 1~O2 using a platinum sponge catalyst.

5. CONCLUSIONS The CO oxidation takes place via the Langmuir-Hinshelwood model, with the following set of equations: k1

CO< k2 >CO ad

(6)

02

(7)

k~ > 2 0 ad k4

CO~d +O~d < k~ ) COaa d CO2a d

k~ > CO 2

(8)

(9)

369 9 The CO, is produced from CO and O on the catalyst by an equilibrium reaction that is faster than the CO,_ desorption. This implies that the rate of desorption of CO,_ from the platinum surface should be taken into account when modeling this reaction. 9 Multitrack offers the possibility to accurately determine the number of active sites on a catalyst.

REFERENCES

1. J.T. Gleaves, J.R. Ebner, and T.C. Kuechler, Catal. Rev.- Sci. Eng., 30(1), p. 49 (1988). 2. F.H.M. Dekker, J.G. Nazloomian, A. Bliek, F. Kapteijn, J.A. Moulijn, D.R. Coulson, P.L. Mills, and J.J. Lerou, Carbon Monoxide Oxidation over Platinum powder; A comparison of TAP and Step-Response Experiments, Appl. Catal. A, acceptedfor publication (1996). 3. J.P. Huinink, A Quantitative Analysis of Transient Kinetic Experiments: The Oxidation of CO by O,/NO on Pt, Ph.D. Thesis, Eindhoven University of Technology, 1995. 4. T. Engel and G. Ertl, Elementary steps in the catalytic oxidation of carbon monoxide on platinum metals, in D.D. Eley, H. Pines and P.B. Weisz (editors), Advances in Catalysis, vol. 28, Academic Press, New York,1979, p. 1.

This Page Intentionally Left Blank

91997 Elsevier Science B. V. All rights reserved. Dynamics of Surfaces and Reaction Kinetics in Heterogeneous Catalysis G.F. Froment and K.C. Waugh, editors

371

N o n - L i n e a r Steady-State Kinetics of C o m p l e x Catalytic Reactions: T h e o r y and Applications G.S. Yablonskii a, M.Z. Lazman b a Department of Chemical Engineering, Washington University, Campus Box 1198, One Brookings Drive, St. Louis, MO 63130 b 300 Hyprotech Centre, 1110 Centre Str. N., Calgary, Alberta, Canada T2E 2R2 A non-linear theory of steady-state kinetics of complex catalytic reactions is developed. A system of steady-state (or pseudo-steady-state) equations can always be reduced to a so called kinetic polynomial. This polynomial is a function of the steady-state reaction rate and the process parameters (concentrations of the reactants, temperature). The kinetic polynomial is a non-linear implicit equation. The physically meaningful solutions correspond to the different steady-states. Using the kinetic polynomial it is convenient to specify the region of critical kinetic behaviour, e. g. region of multiplicity of steady-states and self-oscillations. As an example of kinetic polynomial CO oxidation was analyzed. Kinetic polynomial can be also applied to the reactor design and reactor control. 1. GENERAL PROBLEM The important stimuli for the development of the non-linear kinetic theory of steadystate catalytic reactions are 1) necessity to explain the critical phenomena that are experimentally observed in the steady-state kinetic experiments and 2) needs of chemical technology to understand and to apply the advantages of non-linear regimes. Mathematical model for the unsteady-state heterogeneous catalytic reaction in a gradientless reactor is of the type [ 1] dC

- - S c f (C,x ) + VoC o - VC dt dx S c --~ = Scg(C,x) = Sc(g +( C , x ) - g - ( C , x ) ) Vg

m

where C is the vector of substance concentrations in the gas phase x is the vector of substance concentrations on the catalyst surface Sc is the surface of the catalyst in the reactor V g is the volume of the gas phase in the reactor Vo and V are the inlet and outlet flow rates of the reaction mixture,respectively Co is the vector of the reactant compositions in the gas phase.

(1) (2)

372

tiC,x) and g(C,x) are the vector-functions of the kinetic relationships f (C,x) = I-'[w(C,x) g(C,x) = FTxw(C,x) (C,x)- w- (C,x)

w(C,x)

= w +

where

F T, F z

are the transposed stoichiometric matrices;

w, w +, w- are the vectors of rates of steps, direct and reverse reaction rates, respectively. w + and w- are expressed in the terms of mass-action-law. If one deals with a steady-state kinetic experiment, the set of differential equations (1) - (2) is transformed into a set of algebraic equations, i.e. a steady-state kinetic model -Scf(c,x) + VoC0 - VC = 0 g+(C,x)- g-(C,x) = 0

(3) (4)

The vector of steady-state reaction rates can be represented as follows:

R=f(C'x)=V~176176 l s rcmZ~at.s

(5)

Common models are mixed algebraic-differential systems, i.e., pseudo-state kinetic models:

Vg dC d-T= _Scf (C,x ) + VoCo _ Vc

(6)

g+(C,x)-g-(C,x)=O

(7)

where an unsteady-state regime of the process is assumed in the gaseous phase and a pseudo-steady-state one is assumed on the catalyst surface, i.e., the rate of formation of the surface intermediate is equal to the rate of its consumption.

2. THE MAIN RESULTS OF THE LINEAR THEORY(BRIEF EXPLANATION) The catalytic reaction mechanisms and corresponding kinetic models can be classified into linear and non-linear models. These terms were introduced by Temkin [2]. For linear mechanisms every reaction involves the participation of only one molecule of the intermediate substance. Therefore the rate of each step depends linearly on an intermediate concentration. Using the principle of pseudo-steady-state concentrations (see equation (7)) we can easily find the solution of linear algebraic equations that corresponds to the linear mechanism and then obtain the values of the pseudo-steady-state (or steady-state)

373 concentrations of the intermediates. Knowing the latter we can easily obtain a general expression for the steady-state reaction rate as a function of the concentrations of the observed substances and temperature. The typical linear mechanism of catalytic reaction is the two-step (Temkin-Boudart) mechanism [3], e.g., the mechanism of water-steam conversion:

1). Z + H 2 0 <--->ZO + H 2 2). ZO + CO ~-+ Z + CO 2 CO+ H 2 0 = CO2+ H2

(8) (9) (10)

where Z and ZO are the intermediate substances on the surface. The steady-state kinetic equation is the following:

R= K+Cc~176

(11)

- K-CH2Cc~

I

where

K + = k1+k2§ ,

K- = klk 2 ,

I = k~Cn20 + klCn2 + k~Cco + k2Cco 2 ,

Cco, CH:o,CH 2, Cc02 are the concentrations of the corresponding gaseous substances,

k{,

k~ and ki-, k2 are the rate parameters for the direct and reverse reactions, respectively. The nominator of equation (11) is so called cyclic characteristic [1], i.e., an expression that is similar to the kinetic equation of the overall-reaction under assumption that the mass-action law is valid for this overall-reaction. If R=0 (equilibrium), the relationship

Keq= CH-----3z'Cc~ is valid, where Keq is the equilibrium constant of the overall-reaction. It

Cco CH~O

means the thermodynamic validity of the kinetic equation (11). The denominator I is the inhibition term. It characterizes the complexity of the reaction, i.e., reflects the retardation of the reaction by intermediate substances on the catalyst surface. Equation (11) can be written in the following form [1, 3].

R = R+ (C)[ 1- f- ( ~ /

~

/ f+ ( C )geq (r)

1

(12)

and

R+( (C)

= Keq

(13)

/ f_(~)

where R+(C) and R-(C) are rates of complex reaction in the direct and reverse directions, <....respectively;

f+(C) and f - ( C ) are kinetic dependencies corresponding to the direct and

reverse overall-reactions; C is the set of concentrations of all the substances" C and C are

374 the sets of concentrations of all the initial compounds and the products, respectively. Equations (12) and (13) correspond to the solution of the well-known Horiuti - Boreskov problem, which can be formulated as follows: how to find the kinetic equation of a complex reaction in reverse direction, from the known expression for the reaction rate in the direct direction and thermodynamic relationships? The steady-state kinetic theory for multi-route linear mechanisms was developed in [1] using graphs theory. Within this approach the steady-state rate of the k-th step is presented as

Rk = ~,c,&/(,

(14)

where Ci is the cyclic characteristic of the i-th cycle that contains the k-th step. Cyclic characteristic is a kinetic equation corresponding to the i-th cycle, assuming that the massaction law is valid for this overall-reaction. Pi is the coupling parameter (see a more detailed explanation in [ 1]). In the general case it is the sum of components (products of rate parameters by concentrations) defined by the reaction graph. Denominator I like the analogous term in equation (11) is the inhibition term that characterizes the complexity of the reaction It should be stressed that both equations (11) and (14) have the Langmuir form that is well-known in heterogeneous catalytic kinetics. The equation (11) is the particular case of the equation (14), i= 1, P= 1. 3. NON - L I N E A R T H E O R Y The analysis of non-linear mechanisms and corresponding kinetic models are much more difficult than that of linear ones. The obvious difficulty in this case is the follows: an explicit solution for steady-state reaction rate R can be obtained only for special non-linear algebraic systems of steady-state (or pseudo-steady-state) equations. In general case it is impossible to solve explicitly a system of non-linear steady-state (or pseudo-steady-state) equations. However, in the case of mass-action-law-model it is always possible to apply to this system a method of elimination of variables and reduce it to a polynomial in one variable [4], i.e., a polynomial in terms of the steady-state reaction rate. We refer a polynomial in the steady-state reaction as a kinetic polynomial. The idea of this polynomial was firstly emphasized in [5]. The polynomial coefficients are, in turn, polynomials in terms of the reaction rate parameters (ks+ and ks ) and reactant concentrations (C), i.e., "Langmuir-type" polynomials, that are similar to the form of the denominators of the equations (11) and (14). The kinetic polynomial for one-route catalytic mechanism is represented as follows

P(R) = BLRL+...+BM RM +BM_IRM-1c+...+Bo CM = 0 where L and M are the integer numbers C is the cyclic characteristic of the overall reaction, see equation (11) and (14)

C = K+f +( C ) - K - f - ( C )

(15)

375 Taking into account the above mentioned "Langmuir" form of the coefficients, the equation (15) can be written as follows

/

P(R) = Z I-[ Kjt" JL

..

+ ~_~ 1-IKjM_~ JM-1

KjM JM

+

(K+f+(C)-K-I-(C))+ +...+Ko(K+ f § ( C ) - K - f - ( C ) ) M = 0

(16)

where K is a product of the rate parameters, see [6]. In the case M=I the kinetic polynomial is represented as:

BLRL +... +B1R + BoC = 0

(17)

or

(18)

The mathematical theory of this problem can be found in our paper [6]. The theory is based on the results of complex analysis, especially multidimensional logarithmic residual theory [7].

3.1 Properties of the kinetic polynomial (K.P.) a) Kinetic polynomial is the most generalized form of kinetic equation for complex reaction. The particular cases of the K.P. are the well-known Langmuir-dependence that is completely corresponding to the linear approximation of K.P. and, obviously, the classical mass-action-law equations. Also it is possible to show that some semi-empirical equations, for example Hougen-Watson equations are the particular cases of the K.P. b) Kinetic polynomial is thermodynamically correct When R = 0 (i.e., equilibrium), one has --)

(.-_

C= K+ f + ( C ) - K - f - ( C ) =

0

(19)

hence a usual relationship follows. c) There can be several physically meaningful of K.P. The kinetic polynomial is a non-linear implicit equation. It is obvious that there can be several solutions of the K.P. The physically meaningful solutions correspond to the different steady states.

376

3.2 The case of the rate determining step Using the kinetic polynomial, one can demonstrate that in the case of the rate determining step the solution is obtained in the form of a series with respect to the small parameter, and a good approximation is the first term of this series which is determined by the formula

R= bk

n 1V/~k 1 - j~__l(gTj)

2o, i=1 j~:k ~"K) .,)

Here k is the index of the rate determining step; bk and vk are the weight and the Horiuti's number of this step, respectively (the weight of the step is equal to the rate of this step at unitary concentranon of the intermediates; Horiuti's number definition see, for example, in o.a.

/

[1,3]); K j = 0J~b7" vj is the Horiuti's number for j-th step; Aji is a cofactor for the ji-th /

J

element of the modified matrix of stoichiometric coefficients which k-th raw is substituted with the stoichiometric coefficients in the left hand side of k-th reaction; Nk is a cofactor to any element of the k-th raw of the same matrix; Pk is the total number of molecules of intermediate substances involved in the rate determining step. The parameter ( 1 / ( . ) is just /v k molecularity according to Boreskov M= y . / vk

3.3 Region in the neighbourhood of equilibrium From the analysis of the kinetic polynomial one obtains ?l

I I ( K j ) D -1 (21) [i=1 j~k

eq

where eq is an index painting to equilibrium. The rest of designations is the same as for the previous equations. The region in the neighbourhood of the equilibrium where the relationship (21) is valid can also be found by analysis of the K.P. 4. AN APPLICATION OF THE KINETIC POLYNOMIAL TO CO OXIDATION Let's analyze the steady-state kinetic model, that corresponds to the well-known adsorption catalytic mechanism (Langmuir-Hinshelwood mechanism): 1)2Z + 02 ~ 2ZO; 2) Z + CO <--->ZCO; 3) ZO + ZCO --->2Z + CO 2, where Z is empty active site of the catalyst, ZO and ZCO are adsorbed substances.

377 The steady-state kinetic model is the following: 2ki~Po2 ( 1 - x - y)2 _ k3xY = 0

(22)

k~Pco (1- x - y ) - k2Y - k3xY = 0 where x, y, (1-x-y) are surface concentrations of substances [Z], [ZO], [ZCO], respectively Po2, Pco are partial pressures of the reactants (02, CO) (Po2 and Pco are parameters of the model (22), POE >> Pco)It is assumed that during our steady-state experiment POE >> Pco. Pco is changing with experiment: k~, k~, k2, k3 are rate parameters. The steady-state rate of CO2 production is equal to: R = k3xy

(23)

Using the above mentioned algebraic technique, it is possible to obtain the kinetic polynomial for non-zeroth values of steady-state reaction rate: (24)

B.3R3 + B2 R2 +B1R + B 0 = 0

At sufficiently high values of rate parameter k3 in the region

(25) k~-Pco < 2k{Po 2 < (kI +k2)2/ -

--

-

/4k~

the analytical solutions for three kinetic branches that correspond to the different kinetic regimes were obtained

R

_ . .

(26)

=

=

r

=

/2k~-Po:

+

The first (upper) and third (low) branches are stable, the second (middle) is unstable [8]. Analyzing the equation (26) it is simple to show that at the "extinction" point Rex t = R 1 = R 2 = 2 k f Po2eXt = k ~ P c o ext and at the "ignition" point Rig n = R 2 = R 3 = k 2. Rext

and Rign are steady-state reaction rates at the "extinction" and "ignition" points, respectively;

378 Po2 ext and Pco ext are the partial pressures of the reactants 02 and CO at the "extinction" point, respectively. It is a "critical simplifying". This fact is experimentally approved for some catalytic systems of CO oxidation [9, 10]. Therefore, the ratio of extreme values

k/'t~'jfeext~'gn~ R~.~Rig" /__

\

is governed by the equilibrium parameter:

k~eco ~

--

K

o

ext

2,eqtCO

where K2,eq is the equilibrium constant of CO adsorption. This case is termed a

"thermodynamics of hysteresis". See more detailed analysis of these phenomena in [ 11]. 5. THE ADVANTAGES OF THE KINETIC POLYNOMIAL The kinetic polynomial (K.P.) can be defined as a state equation of the open isothermal chemical system related to the reaction rate and the parameters of the process (composition of the mixture, temperature). Using the kinetic polynomial it is convenient to specify the regions of critical phenomena, e.g., regions of multiplicity of steady-states, self-oscillations. Also it is possible to determine correlations between critical parameters and to identify coefficients of the model. Kinetic polynomial can be applied to reactor design and reactor control. Modified kinetic polynomial can be used to describe more complicated [12] and realistic cases, especially non-ideal and non-isothermal systems.

REFERENCES

1. G.S.Yablonskii, V.I. Bykov, A.N. Gorban and V.I.Elokhin, Kinetic Models of Catalytic Reactions//C omprehensive Chemical Kinetics, V.32, Elsevier, Amsterdam -N.Y., 1991. 2. M.I.Ternkin, Dokl..AN USSR, 152 (1963) 156. 3 M.Boudart and G.Diega -Mariadassou, Kinetics of Heterogeneous Catalytic Reactions, Princeton University Press, Princeton, N.J., 1984. 4. B .L.van der Warden, Modern Algebra, p 2. Ungar, New-York, 1972. 5. M.Z.Lazman, G.S.Yablonskii and V.I. Bykov, Sov.J.Chem.Phys., 2 (1985) 404. 6. M.Z.Lazman and G.S. Yablonskii, Kinetic Polynomial: a New Concept of Chemical Kinetics, in Patterns and Dynamics in Reactive Media, IMA Proceedings, 1989; SpringerVerlag, Berlin - N.Y., 1991, p. 117. 7. L.A.Aizenberg and A.P.Yuzhakov, Integral Representation and Residues in Multidimensional Complex Analysis (in Russian), Nauka, Novosibirsk, 1979. 8. V.I.Bykov, M.Z.Lazman and G.S.Yablonskii,.J Fiz Chim., 60 (1986) 86//Translated in:Ruissian J.of Phys.Chem. 60 (1986) 49 (ISSN 0036-0244) 9. M.Z. Lazman, G.S.Yablonskii and V.A.Sobyanin, Kinet.Catal., 27 (1986) 57. 10. M.Ehsasi and J.N.Block, Proceed. of Int. Conf. "Unsteady-State Processes in Catalysis", Novosibirsk, 5-8 June,1990;Ed. Yu.Sh.Matros, VSP Netherlands, 1990, p .47. 11. G.S.Yablonskii and M.Z. Lazman, React.Kinet.Catal.Lett., 59 (1996) 145. 12. G.S.Yablonskii, M.Z.Lazman and A.M.Kytmanov, Ukrainian Chemistry Journal,. 58 (1992) 1060.

91997 Elsevier Science B.V. All rights reserved. Dynamics of Surfaces and Reaction Kinetics in Heterogeneous Catalysis G.F. Froment and K.C. Waugh, editors

379

N e u r a l n e t w o r k based model of the kinetics of catalytic hydrogenation reactions E.J. Molga a and K.R. Westerterp b "Chemical Engineering and Process Department, Warsaw University of Technolo~, ul. Waryfiskiego 1, 00-645 Warsaw, Poland b Chemical Reaction Engineering Laboratories, Faculty of Chemical Technology, Twente University, P.O. Box 217, 7500 AE Enschede, the Netherlands

Application of artificial neural networks (ANN) for modelling of the kinetics of a catalytic hydrogenation reaction in a gas-liquid-solid system has been studied and discussed. The kinetics of the hydrogenation of 2,4-DNT over a palladium on alumina catalyst has been described with feedforward neural networks of different architectures. A simple experimental procedure to supply learning data has been proposed. The accuracy and flexibility of the hybrid first principles-neural network model have been tested and compared with those of the classical model.

1. I N T R O D U C T I O N The development of a kinetic model is indispensable to accurate predicting of reactor behaviour as well as to maximizing the yield of the desired products. For a complex reacting system extensive studies have to be performed to obtain all information concerning mechanisms, kinetics and thermodynamics. In industrial practice, particularly in fine chemicals industry, the time required for kinetic investigations is very often longer than the product life time on the market. In this situation a plant manager decides to execute the process without detailed knowledge of the reacting system - that is far-off the optimal trajectory. In a case of exothermic reactions, autocatalytic effects and systems with possible, dangerous side or consecutive reactions this can additionally provoke runaway events leading to serious losses and accidents. A fast development in the application of new special techniques is recently noticed for modelling and simulation of systems when no prior knowledge about the process is available. Amongst others artificial neural networks (ANN) are widely employed for system representation, modelling, control and identification. Models based on ANN exhibit the rule-following behaviour without containing any explicit representations of those rules. The term artificial neural networks is a generic description for a wide class of connectionist representation inspired by models for the activity of the human brain. Using ANN each real, arbitrary complex, system can be approximated with an appropriate set of input-output data. This approach is frequently applied in many areas of chemical engineering such as fault diagnosis, process control, design and modelling- e.g. see [ 1-3]. In the simplest approach the entire object (process) is described as a "black-box" and modelling is based only on the observed input-output signals- e.g. see [4-5]. This "black-box" approach has the very convenient ability to approximate a behaviour of any system without prior information on the

380 process. It simultaneously has also a few serious disadvantages. The most important one is a poor ability for the generalization of the obtained results- e.g. the trained network can not be used successfully for the representation of the same reaction system in another reactor. A more advantageous hybrid approach combines a first principles knowledge, which can be easily represented by a set of differential mass and heat balance equations, with a conjuncted neural network which works as a non-parametric estimator of unmeasured parameters [6-7]. This hybrid model has a better ability for knowledge generalization than the standard "black-box" neural network models and it is able to approximate the system behaviour much more accurately. In this study the application of ANN to modelling of the kinetics of a catalytic hydrogenation reaction in the gas-liquid-solid system is studied and discussed. Also a hybrid first principle-neural network model, where conventional mass and energy balances are supported with a trained network, implemented to represent the reaction kinetics, is proposed and tested. 2. MULTILAYER FEEDFORWARD NEURAL N E T W O R K S The basic feedforward neural network performs a non-linear transformation of the input data in order to approximate the output data. This net is composed of many simple, locally interacting, computational elements (nodes/neurons), where each node works as a simple processor. The schematic diagram of a single neuron is shown in Fig. 1. The input to each i-th neuron consists of a N-dimensional vector X and a single bias (threshold) bi. Each of the input signals x~ is weighted by the appropiate weight Wij, where j = 1-N.

ilIbi

Xl x2 Wi2/~+~ !

XN/WiN :

1

Si

,If(si) i '

2y, B

H2

)

A

,

Figure 1. Schematic diagam of a sinNe neuron.

E

Figure 2. Hydrogenation of 2,4-DNT - reaction pathways.

The i-th neuron calculates the weighted sum Si as: N

S. = E

Wox. + b ,

J"l

Then the output of the neuron is determined following the general relationship:

(1)

381 y, = f (S,)

(2)

where the so-called activation function f (S) usually is taken as a unipolar sigmoid function 1 /Is)

:

-s

1 +e

(3)

Single nodes arranged appropriately in layers form the net- see Fig.3. The number of input and output layers is determined a priori by the nature of the problem being modelled, the manner of input data representation and the form of the network output required. The number of hidden layers as well as the number of nodes in each hidden layer should be optimized. Generally a demand for learning patterns increases considerably while the number of hidden layers increases; a decrease of this number restricts the network ability for non-linear approximation. A set of learning patterns, each consisting of two (X, d) vectors, is used for training of the network where: X = [xl, x2, x 3.... xN]T and d = [ d l , d2, d 3.... dM] "r for N input signals and M output signals, respectively. For P experimentally determined patterns the target function for the learning is defined as follows: P

E(W) = ~1 ~ --

k-1

M

~

Cv[ ) - d,!k))"-

(4)

i-1

During the learning procedure the current values of the weight vector W are changed to minimize the target function E(W) i.e. to minimize the differences between the calculated and target output signals. Usually a backpropagation algorithm, based on gradient optimization, is used but also other optimization procedures are often employed. 3. KINETICS OF THE CATALYTIC HYDROGENATION REACTION In our laboratories extensive studies on the catalytic hydrogenation of aromatic nitrocompounds, as an example of the catalytic three-phase reactions, have been carried out in reactors of different types - e.g. see [8-10]. In all cited cases the time consumed for kinetic investigations had a very significant contribution to the total experimental effort [ 11-13 ]. Particularly for the hydrogenation over palladium on alumina catalyst, the experimental investigations leading to the detection and quantitative description of internal diffusion resistances in catalyst pellets have taken a lot of time. 3.1. L a n g m u i r - t l i n s h e i w o o d model

This approach applied to the description of the hydrogenation reactions has been described in detail in our previous paper [ 13]. For the hydrogenation of 2,4-DNT over palladium on alumina catalyst the complex consecutive-parallel reaction pathways are shown in Fig. 2. The appropriate rate equations can be written as follows [13]: r.j = K k ~ % n 0 ~ 0 ,

(5)

382 where except the kinetic rate constants (k~B, %) and the fractional surface coverages (O H and 03 also the parameters characterizing the influence of internal diffusion resistances (1"1)and the catalyst deactivation (1<:)appear. Quite complex expressions have been derived to describe the kinetics of the studied reaction - see [13]. 3.2. Neural network model Using artificial neural networks (ANN) the reaction system, including intrinsic reaction kinetics but also internal mass transfer resistances, is considered as a black-box and only input-output signals are analysed. With this approach the conversion rate of the i-th reactant into the j-th species can be expressed in a general form as a complex function, being a mathematical superposition of all above mentioned functional dependencies. This function includes also a contribution of the internal diffusion resistances. So each of the rate equations of Eq. 5 can be described with the following function based on the variables which uniquely define the state of the system:

% ~(TR,Pm XA, XB, Xc, XD) =

(6)

The unknown function q~ can be represented (approximated) by a proper neural network. The following problems should be solved to tackle the modelling of kinetics with ANN: -determination of the input-output variables relevant for the investigated reaction, - choice of the net type and net architecture suitable for the representation of the kinetics, - choice of the learning method, planning of the experimental program and the extraction of the data for the learning procedure. The following vectors have been chosen for a system representation: the input variables X = [T R, Pn, XA, XB, Xc, XD]T and the output variables d = [rAB, rBc, rCE, rAD, rDE]~ - see Fig. 3. The first data set is derived from an analysis of the system behaviour and it uniquely defines the state of the reaction mixture. The second data set results from the modelling requirements - it makes a full description of the reacting system possible. The values of rij can not be measured directly but they can be determined from experiments. A feedforward multilayer neural network (MNN) has been chosen for the purpose of this study and the numbers of the input and output nodes equal to 6 and 5, respectively. The number of hidden layers and the number of neurones in each hidden layer have been optimized REACTION KINETICS in order to find a simple but sufficiently TR robust and flexible structure of the net, which accurately represents the kinetics PH ~Bc of the investigated reaction: In this xA ~CE study a Levenberg-Marquard method XB ~Ar has been employed for the learning xC procedure to minimize the target -I ~DE function of Eq. 4. The proposed xD experimental method supplying learning data is described in detail below.

Figure 3. Schematic diagram of the feedforward neural network for the representation of the kinetics of the catalytic hydrogenation of 2,4-DNT.

383 4. M O D E L O F A B A T C H R E A C T O R A batch reactor with a cooling jacket has been chosen for the comparative studies performed in this paper. The catalyst pellets are located in a special basket. Ideal mixing of the liquid phase and a good development of the gas-liquid interface can be assumed due to the installation of a turbine stirrer and baffles. The following mass balance equations have been derived for the hydrogenation of 2,4-DNT over a palladium on alumina catalyst and performed in the batch reactor with the methanol as an evaporating solvent : - for i-th reactant

a(N,. x,) : (r~.- rp m t dt

- for methanol in the liquid phase

d(Nr Xs) - - R v dt

- for methanol in the gas phase

d(N~ Ys) =R dt

- for the gas phase

dNG -

-

R u

-

-

R

(7)

(8)

(9)

(10)

+ R,,,,

dt

- for the liquid

dNL

(11)

dr

- for the hydrogen

d(N~ yn)

act

- for the nitrogen

d(NG YN)

dt

: -R H

(12)

= 0

(13)

The heat balance for this reactor can be written as:

a(NL cpL TR)

d(NG cp G TR) +

dt

d(mcat cp.7 TR) +

dt

5 : ~

dt

r,j A t t 0 moo ' - Q , v

- Qc

(14)

~.l

Following the reaction pathways the hydrogen consumption rate has been determined as: RH = (2rAB + rBC + 3rcE + 3rAD + 3rDE) meat. For the semibatch reactor operating mode, when the hydrogen is continuously supplied into the reactor to keep the reactor pressure constant, the values of R n = 0 have to be introduced into Eqs. 10 and 12, respectively. The solvent evaporation rate has been calculated as: l~v = (1% aGE VL) (Yis - Ys) P/RTR. The appropriate value of the transport coefficient has been found to keep the gas phase saturated with the methanol vapours during the reaction. The rate of heat withdrawal due to solvent evaporation is equal t o Q~v = R~v AHoy and

384 the cooling rate via the cooling jacket has been estimated as: Qc = UA (T R - T j ) . The system of balance equations of 7-14 has been solved for the appropriate initial conditions with a fourth order Runge-Kutta method. 5. R E S U L T S AND DISCUSSION The reactor model presented above has been used in two versions: as a conventional model (CM) and as a hybrid model (HM), respectively. In the conventional model the reaction rates qj appearing in the balance equations of Eq. 7 have been calculated following the expressions for Langrnuir-Hinshelwood kinetics [13] derived in our previous studies. In the hybrid first principlesneural network model (HM) the conventional kinetic subroutines in the CM algorithm have been replaced by the neural network, so in the HM the reaction rates rij have been supplied by the trained network. The abilities of neural networks to represent the reaction kinetics as well as the influence of the experimental program on the accuracy and flexibility of the hybrid model have been studied comparatively at different operating conditions.

5.1. Experimental method proposed to supply learning data In the classical way of kinetic investigations the reaction is carried out at isothermal conditions and during the reaction progress the reactant concentrations vs. time profiles are registered. Further, using the detected or assumed mechanisms, the proposed kinetic expressions are fitted to those concentration profiles and the kinetic rate and adsorption constants are determined. Temperature dependencies for these constants can also be estimated if this procedure is repeated at different temperatures. For a complex reaction system such studies take a lot of time and require a strong experimental effort, particularly for multiphase systems where external and internal diffusion resistances have to be taken into account. This classical modelling procedure becomes extremely difficult or even impossible if the reaction mechanisms are unknown or uncertain. In this case application of NN can be very promissing. For the NN approach instead of many isothermal experiments only a few non-isothermal batch runs are postulated here to supply a sufficient amount of learning data. An arbitrary temperature program can be chosen for the neural network approach to pass an as wide as possible range of the operating temperatures during one experimental run. According to the above proposed architecture of the net the following variables have to be measured during an experimental run: TR, Pn, XA, XB, Xc, XD and also the appropriate, not directly measurable variables: r ~ , rBc, rcE, r ~ , rD~., which have to be determined to complete the set of learning data. The required conversion rates r~ican be determined from the measured reactant concentrations vs. time profiles. For each experimental run on the basis of the time profiles each reactant rate ri can be determined with the relationship:

~,

-

N~o ax, inca t

(is)

dt

and further the conversion rates rij can be found by solving at any moment a set of 5 linear algebraic equations. Depending on the reaction pathways, each equation can be written as follows: r~= r~ + rii and finally the learning data are prepared extracting at any moment the appropriate set of (X, d) data.

385

The aim of this study is to check how much a kinetic experimental program can be shortened in the case of an ANN application and also how the accuracy and flexibility of this approach are related to the classical approach. To this end a comparative study for the neural network and conventional representation of the reaction kinetics has been performed. To avoid an influence of experimental inaccuracies on comparison as well as to check proposals concerning the experimental methods the learning data have been simulated using the conventional model CM. The simulated experiments chosen for the learning are listed in Table 1. It has been found that despite the complexity of the described reaction system only a few experimental runs are necessary to supply a sufficient amount of representative learning data. At isoperibolic operating conditions (constant temperature of the cooling liquid) for exothermic reactions the reactor temperature T R and also the other rate controlling variables (PH, X0 are changed over a wide range. This method supplies a large variability of the learning patterns even from a single measurement, so the representation of even very peculiar system properties is possible. Table 1 Experiments supplying data for learning. Run

Initial conditions

Method

TR [K]

Tj [K]

PH [MPa]

%0 [mol/m 3]

R1

303

303

0.4

219

isoperibolic

R2

303

303

0.1

219

isothermal-isobaric

R3

303

303

0.1

219

isoperibolic-isobaric

R4

333

273

0.4

219

isoperibolic

5.2. Learning procedure From the simulated state variables (TR, PH, Xt) vs. time profiles obtained for the runs listed in Table 1 all learning data have been prepared as described above and further they have been normalized within the range of [0,1 ]. The training data sets (X, d) have been randomly, row by row, mixed before feeding them into the network. As many as 500 learning patterns (X, d) have been prepared in this way creating the set of learning data L 1 - see Table 2. A few nets of different architecture have been tested. The number of hidden layers as well as the number of neurones in each layer have been changed to study the influence of the net architecture on its ability to approximate the investigated kinetics. The obtained minimal values of the target function E(W) are listed in Table 2. Also the root mean square values rms = -f(E(W)/P) are shown to present the quality of data fit. It is visible in Table 2 that significant improvements in quality of fitting can be achieved making the network architecture more complex and introducing more neurones into the hidden layer(s). The nets with at least 2 hidden layers are able to represent the kinetics of hydrogenation reaction quite accurately, and introduction of one hidden layer more does not improve the quality of fitting.

386 Table 2 Results of learnin~ Net code

Net architecture

Number of weights

Learning data set

P

E(W)

rms 102

N1

6-4-5

53

L1

500

0.177

2.66

N2

6-6-5

77

L1

500

0.161

2.54

N3

6-6-3-5

83

L1

500

0.012

0.71

N4

6-6-5-5

107

L1

500

0.002

0.32

N5

6-6-5-5-5

137

L1

500

0.022

0.96

5.3. Accuracy and flexibility of the neural network approach Trained nets have been implemented into the hybrid model. The net N4 - see Table 2 - represents the optimal topology. It assures a good accuracy of the predictions and simultaneously a relative simple architecture. A comparison of the results obtained with the hybrid first principles-neural network model and those predicted with the conventional one (CM) is shown in Figs. 4a-b. 1.01

L..g,;,u,r-H,.~h.,.,oo~ r.o~:,/~

~

0.8~ I

1.0"~ 0 8t ~

Langmuir-Hinshelwood m o d e l / ...... 9 neural notwork model / §

--

0.6~

X" o.4I

X"O. 4

.

t

, ......

+ "

0.2

[ O.OI

0.0

0.2

0.4

~t.lE-] a.

0.6

-

0.8

.

. 1.0

O0

0.0

0.2

0.4

~H [-1

0.6

0.8

1.0

b.

Fig. 4. Comparison of experimental concentration profiles, points- [ 13 ], and predicted ones for isothermal and isobaric runs performed at the following operating conditions: PH = 1.0 MPa and T R = 3 0 8 K (a), 345 K (b), respectively. Both predictions are related to the experimental concentration profiles measured in our previous studies [13]. The results obtained with the I-]M are very close to these obtained with the conventional approach (CM) - see solid and dotted lines in Figs. 4a-b. Only small differences in the distribution of the intermediate compounds B and C are noticed; the maxima on the concentration vs. ~H curves for these compounds are located at almost the same values of the hydrogen conversions ~H. For both, conventional and neural network approaches, some discrepancies between predicted and experimental concentration profiles can be observed for compounds A and B a t T R = 345 K. We observe also a short-coming of the neural network approach. At full hydrogenation of the original reactant, that is at ~H = 1.00, a still not completed conversion is predicted. Further we may observe that in Fig. 4b at ~H = 1.00 still B and D are

387 present, despite 100% conversion into E. Several runs at different reactor conditions have been simulated with the HM to check the flexibility of the neural network approach. A comparison of the obtained results is shown in Fig. 5, where concentration profiles are plotted for isothermal and isobaric runs performed at different hydrogen pressures. A good stability of the obtained solutions is observed in the diagram. At hydrogen partial pressures pn < 0.6 MPa and at the end of the process, where concentrations of intermediate products are 1.0 .. PH [ M P a ] very low, slight instabilities of the obtained o.a 0.8 solutions appear. A significant influence of ...... 1.0 the hydrogen partial pressure PH on the yield 0.6 ,> B .......... 1 . 2 of the intermediate compounds C and B is predicted. Also the special property of the 0.4 ~: investigated reaction system being molecular 0.2 .'J); quenching [ 13] is reproduced by the neural network model- notice in Figs. 4-5 the O0 0 2000 4000 6000 8000 10000 reaction rate being zero-order in DNT at the t[sI beginning of the process up till ~H = 0.2. Fig. 5. Comparison of the concentration profiles obtained with the hybrid model for isothermal and isobaric runs performed at TR = 308 K and at different partial hydrogen pressures PH6. CONCLUSIONS The application of ANN for a representation of reaction kinetics can be a very promising method to solve modelling problems. Besides intrinsic kinetics also internal diffusion resistances can be included into the neural network based model. This approach significantly reduces the time required for experimental studies. Despite that neural networks do not help to understand and develop a real reaction mechanism, they make the prediction of the reactor behaviour possible. This approach can be essential in the case of complex or uncertain kinetics - e.g. for polymerization reactions. In this study the neural network approach has been tested for a batch reactor. A trained network can be successfully implemented into any type reactor model. - This work was supported partially by the Netherlands Organization for the Advancement of Scientific Research (NWO) and partialy by the Polish Committee for Scientific Investigations (KBN Poland) - grant no. PB 0424/P4/94/06.

Acknowledgements

NOTATION A b Cp

d E(W)

-jacket heat exchange surface area, [m 2] - bias (threshold) value, [-] - molar specific heat, [J/mol K] - output signal for training, [-] target function, [-] heat of reaction, [J/moll

-

-

3 8 8

•Igv

-

f k~

-

meat

N % Q

-

-

r

-

R

-

t T W U xi

-

PH

-

= x~/Xao

-

X

-

Yi

-

y

-

heat of evaporation, [J/mol] activation function, [-] rate constant for conversion of reactant A into B, [mol/kg cat s] mass of the catalyst, [kg] number of moles, [moll lq./k~, reduced rate constant, [-] rate of heat withdrawal, [J/s] partial pressure of hydrogen, [Pa] conversion rate, [mol/kg cat s] reaction rate, [tool/s] time, [s] temperature, [K] network weight, [-] overall heat transfer coefficient [W/m 2 K] mole fraction, [-] dimensionless concentration of organic reactant, [-] neuron input signal, [-] molar fraction in the gas phase, [-] neuron output signal, [-]

Greek letters

catalyst deactivation factor, [-] - overall effectiveness factor, [-] - fractional surface coverage by hydrogen, [-] fractional surface coverage by organic species, [-] - function in Eq. 6, [-]

K

-

q 0H 0i q) R

E

-

F

E

R

E

N

C

E

S

1 J.C. Hoskins and D.M. Himmelblau, Comput.Chem.Engng. 12 (1998), 881. 2. A.J. Morris, G.A. Montague and M.J. Willis, Trans. IChemE., 72, Part A (1994), 1. 3 M.L. Thompson and M.A. Kramer, AIChE J, 40 (1994), 1328. 4. N. Bhat and T.J. McAvoy, Comut.Chem.Engng. 14 (1990), 573. 5 Y. You and M. Nikolau, AIChE J, 39 (1993), 1654. 6. D.C. Psichogios and L.H. Ungar, AIChE J, 38 (1992), 1499. 7 I.M.Galvan, J.M.Zaldivar, H.Hernandez and E.Molga, Comput. Chem.Engng. 20 (1996), 1451. 8 K.B. van Gelder, P.C. Borman, R.E. Weenink and K.R. Westerterp, Chem. Eng. Sci., 45 (1990), 3171. 9. K.R. Westerterp, H.J. Janssen and H.J. van der Kwast, Chem. Eng. Sci., 47 (1992), 4179. 10. K.R. Westerterp, E.J. Molga and K.B. van Gelder, Chem. Eng. Process. (1996), accepted 11. H.J. Janssen, A.J. Kruithof, G.J. Steghuis and K.R. Westerterp, Ind. Eng. Chem. Res., 29 (1990), 754. 12. H.J. Janssen, A.J. Kruithof, G.J. Steghuis and K.R. Westerterp, Ind. Eng. Chem. Res., 29 (1990), 1822. 13. E.J. Molga and K.R. Westerterp, Chem. Eng. Sci., 47 (1992), 1733.

91997 Elsevier Science B.V. All rights reserved. Dynamics of Surfaces and Reaction Kinetics in Heterogeneous Catalysis G.F. Froment and K.C. Waugh, editors

389

M i c r o k i n e t i c A n a l y s i s of T e m p e r a t u r e - P r o g r a m m e d E x p e r i m e n t s in a Microreactor Flow System O. Hirtrichsen, E Rosowski, M. Muhler*, and G. Ertl Ffitz-Haber-Institut der Max-Planck-GeseUschaft Faradayweg 4-6, D- 14195 Berlin (Dahlem), Germany Temperature-programmed (TP) experiments provide useful information about the kinetic properties of adsorbates. The TP experiments can be carried out either in ultra-high vacuum (UHV) or at ambient pressure using flow set-ups. Thus, they can serve as a tool to bridge the material and pressure gap between surface science and heterogeneous catalysis by applying microkinetic analysis. The temperature-programmed desorption (TPD) of H2 from Cu is used as an example to discuss the derivation of kinetic parameters from measurements in UHV. N~ TPD experiments from a multiply promoted Fe catalyst are analyzed to demonstrate that the assessed influence of readsorpfion effects depends on the reactor model chosen. The transient continuous-flow stirred tank reactor (CSTR) and the transient plug-flow reactor (PFR) are used for the modeling. Finally, the temperature-programmed surface reaction (TPSR) of adsorbed atomic nitrogen with H2 yielding NH3 over Fe- and Ru-based catalysts is re-examined to illustrate the influence of the reactor model. 1. Introduction

Over the past few years, microkinetic analysis has developed into a powerful tool to elucidate the catalytic surface chemistry of a variety of reactions [ 1]. Using the surface science approach, microkinetic models have been established in case of ammonia synthesis over iron (see references in ref. [2]) and methanol synthesis over copper [3] based on a detailed knowledge about the kinetics of the involved elementary steps which allow to extrapolate the results from single crystal studies to predict the performance of an industrial catalyst operating at high pressure. Following the ideas of Dumesic et al. [ 1], we recently presented a microkinetic model consisting of the full mechanistic pathway for the synthesis of ammonia catalyzed by Cs-promoted ruthenium supported on MgO (Cs-Ru/MgO) [4]. The model based on the Langmuir-Hinshelwood (LH) mechanism was able to reproduce the transient as well as steady-state kinetics for Cs-Ru/MgO and to bridge the pressure gap between ultrahigh vacuum (UHV) and high pressure. Moreover, the model could be applied to the Cs-free Ru/MgO catalyst incorporating only consistent changes in the model in accordance to the transient experiments [5]. In this contribution, we present computer analyses of several selected temperature-programmed desorption (TPD) and temperature-programmed surface reaction (TPSR) experiments in a microreactor flow system operating under atmospheric pressure. The continuous stirred tank reactor (CSTR) and plug flow reactor (PFR) models have been applied for the design equation as *Correspondingauthor

390 the two limiting cases in the degree of mixing within a microreactor. Kinetic parameters derived from experimental studies on single crystal surfaces under UHV condition were used for our modeling. The transient experiments combined with the steady-state kinetics under industrial conditions and the characterization of the catalysts serve as a basis for a detailed understanding of the performance of commercial catalysts in industrial catalytic processes like ammonia and methanol synthesis. The temperature-programmed desorption of H2 turned out to be a new tool for the determination of Cu metal surface areas [6]. For industrial purposes, the interaction of hydrogen with copper surfaces is a significant elementary step in methanol synthesis. Moreover, adsorpfion/desorption kinetics on copper single crystal surfaces is an interesting topic in surface physics [7]. In our present analysis, the derivation of the kinetic parameters from UHV conditions and the importance of the reactor model for the H2 TPD performed with a ternary Cu/ZnO/Alz O3 water gas shift catalyst under atmospheric pressure is discussed in detail. Ammonia synthesis is another important industrial catalytic process. Extended research has been done on various Fe single crystal surfaces [8] resulting in a complete microkinetic model which describes high-pressure ammonia synthesis on commercially used magnetite-based iron catalysts as well as the elementary steps observed on single crystals [2]. Since the dissociation of the adsorbed molecular dinitrogen precursor was found to be the rate-limiting step (rds) in ammonia synthesis, it is necessary to examine this elementary step very closely. Modeling the rds with the wrong reactor model might produce wrong results for the complete microkinetic model. Therefore, in our present analysis it is shown that the assessment of readsorption effects during N2TPD experiments depends on the reactor model. The hydrogenation of atomic nitrogen (N-,) preadsorbed on an iron-based catalyst surface has been studied by Fastrup et al. [9]. For the sake of simplicity, the non-steady-state TPD cell during the TPSR experiment has been treated as a CSTR. In the present study, simulation results are shown using the proper PFR model. Additionally, experimental and simulation results obtained with a Cs-Ru/MgO catalyst are presented to illustrate the influence of the reactor model.

2. Modeling The experimental TPD and TPSR experiments were conducted in a glass-lined stainless steel TPD cell with an inner diameter of 3.9 mm integrated in a high-pressure stainless steel flow apparatus equipped with a calibrated quadrupole mass spectrometer as described in reference [ 10]. The U-tube was filled with catalyst grains of the 300/~m sieve fraction resulting in packed bed heights up to 20 mm. For this microreactor/TPD cell set-up, the following assumptions were shown to be valid [ 11]: (i) No radial and axial temperature gradients are present in the catalyst bed, and correspondingly the gas phase temperature equals the surface temperature of the particle. (ii) External mass transfer effects are negligible, i.e. the concentration of reactants and products at the exterior surface equals to the gas phase concentration. (iii) Because of the small diameter of the U-tube, radial mixing in the bed is perfect. (iv) The desorption/adsorption process or the surface reaction is not limited by intraparticle diffusion according to the criteria given in refs. [ 12-15].

391 The nomenclature in this paper follows the annotations given by Dumesic [ 1]. The characteristics of the microreactor was studied by applying either the model of a transient continuous-flow stirred tank reactor (CSTR) under the conditions of perfect mixing or a plug-flow reactor (PFR) when the adsorbate concentration varies along the bed length. The transient behavior of the reactor was modeled since the impact of the change in the molecular flow of the desorbing species during the heating ramp is not known a piori. For TPD and TPSR experiments without heat and mass transfer limitations, the material balance for the transient CSTR operating at constant total pressure for an ideal gas reduces to the reactor mass balance specified in ref. [ 1]. In the absence of any axial mixing (PFR model), the partial differential equations have been derived. The system of nonlinear differential equations consisting of the steady-state equations for all surface intermediates, the site balance on the catalyst, and the material balances for the gaseous species has been solved numerically. In the case of the transient CSTR, the resulting set of linear or nonlinear ordinary differential equations (ODEs) was solved using either a variable-order, variable-step method for stiff ordinary differential equations implementing the backward differentiation formula method found in the library of NAG [ 16] or the software package DDASAC (Double Precision Differential-Algebraic Sensitivity Analysis Code) developed by Stewart et al. [17] for solving nonlinear initial-value problems containing ODEs and (optionally) algebraic equations. Stewart's code is an extension of Petzold's implicit integrator routine DDASSL (Differential~Algebraic System Solver which uses a variable-step, variable-step predictor-corrector approach developed by Gear [18,19]. Modeling the reactor as a PFR changes the ordinary differential equations to partial differential equations. The set of linear or nonlinear partial differential equations can be solved using a collocation method. An efficient solving method is given in the library of NAG [ 16]. Stewart et al. [ 17] developed the software package PDASAC (Partial-Differential-Algebraic Sensitivity Analysis Code) for differential-algebraic initial-boundary-value problems which turned out to be very useful solving reaction-diffusion problems in catalytic reactors. Using the spatial discretization, the system changes to a set of nonlinear ordinary differential-algebraic equations (method of lines). The resulting system can be solved using the abovementioned techniques.

3. Results and discussion

3.1. I-I2 TPD from a copper-based catalyst On Cu(111), Cu(100) and Cu(110) single crystal surfaces, Anger et al. [7] analyzed the adsorption/desorption kinetics in great detail. The evaluation of the desorption spectra for H2 on Cu(111) was straightforward resulting in a coverage-dependent expression of the desorption parameters. The desorption energy was found to decrease with increasing hydrogen coverage OH associated with a decrease of the preexponential factor with increasing OH. The adsorption kinetics of hydrogen on Cu(110) have been stu6ied by CampbeLl et al. [20]. By means of molecular beam experiments, the activation energy of dissociative adsorption was determined to be (604-5.7) kJ/mol associated with a prefactor of the sticking coefficient of 10~176177176 . Muhler et al. [6] presented H2 TPD experiments with a binary Cu/A12Oa and a ternary Cu/ZnO/A12Oa water gas shift catalyst (Haldor Tops0e LK-821) after reduction as well as after methanol synthesis. Since the adsorption of hydrogen is known to occur directly (not via a precursor-mediated

392 pathway) we write the dissociative chemisorption and associative desorption in one equation

H2+2.

~

2H-,

(1)

where 9 denotes a free active site. The differential equation of the surface intermediate H . is given by d OH

/

2

=

dt

|

9

9~ ,1"1]

\

where ki denotes Arrhenius-type rate constants and | the fractional coverage relative to hydrogen saturation. The kinetic parameters for step 1 (preexponentials factors and activation energies) have been obtained in the following way: for the desorption parameters the results for H2/Cu(111) from Anger et al. [7] were used. Applying the threshold method (TTPD), the coverage dependent activation energy was derived to be Ede, = 77.4 kJ/mol - 29.3 x ~gH kJ/mol (where the actual coverage 19 is defined as O n = N/Nsu,.! with a saturation coverage of half a monolayer (Om,~ = Nma~/N,~,,.! = 0.5) and a frequency factor to be t, = exp(-8 -3 x O ~ ) cm2/(s.particle) for the Polanyi-Wigner equation [7]. Due to a change from actual coverage OH to the fractional coverage 19n = N/N,~a~, (N,~,,JNo,,,.! = 0.5) in our model, we derived a preexponential of Ades = 1.5 x 1011.exp(_ 1.5 x 19H)S- 1 including the condition that one active site consists of two Cu(111) surface atoms and an activation energy of desorption of E d e , = 77.4 kJ/mol 14.6 x 19H kJ/mol. Since the sticking coefficient of H2 on Cu( 111 ) is not known, the prefactor for the sticking and the activation energy of the Cu(110) surface was used as a reliable approximation. From the kinetic gas theory, a preexponential factor Aad, = 9.2 x 105(tort.s) -x was derived for our model. In our calculation we assumed 132 pmol/gr active sites based on theexperimental results ofMuhler et al. [6] for the CturZnO/A12O3 based on a H/. = H:(2 Cu) = 1"1 stoichiometry. Under the conditons of perfect mixing gas phase condition is uniform throughout the bed. The reactor mass balance for a transient continuous stirred tank reactor was used in the form

dpH2 dT

(--PH2 F~ -- po , ( k a p n 2 0 , 2 _ k - 1 1 9 ~ ) ) N,~

For the non-transient CSTR model the design equation reduces to

Ad~, "e~P( -E"'" . 0 ~ nil2

"-" po

+ Aads . e z p (--Eod.). (1 -- 0/-/) 2 RT

Assuming that the composition of the carrier gas varies from point to point along a flow path the following design equation of a plug flow reactor was used:

OF, m

Fo OFo m

OT

N~

with the relation

F6 H2 pO, PH2

F,

OS

393 subject to the initial condition T--T O

19H --

1.0,

PH2

- - O~

and correspondingly the boundary conditions

,S

-- 0

S--1

PH2

--0

OpH2

=0.

aD

#,

600

I

.

.

.

.

E {... o

,

40O

:

',,,~=

/11

II

E .,,,...,

i:~

200

-r

0 2OO

i

i

i

i

I

250

i

i

i

i

I

i

i

i

300

Temperature /

i

l

350

i

~

,

i

I

400

K

Figure 1. H2TPD from a ternary Cu/ZnO/AbOa catalyst. Comparison of the modeling results using the desorption parameters derived by the TrPD method (solid trace B) and using initial values (dashed trace C). The experimental data (trace A) obtained by Muhler et al. [6] are shown in the inset: reactor volume 0.3 cm a, reactor pressure 760 tort, inlet flow rate 100 Nml/min, heating rate 6 K/min, number of sites 2.1-1019.

The inset of Figure 1 shows the experimental data (trace A) obtained with the Cu/ZnO/A12Oa by Muhler et al. [6]. The narrow H2 peak centered at around 300 K with a full width at half maximum (FWHM) of 47 K using a heating rate of 6 K/min clearly shows no significant readsorption as expected from the low sticking coefficient at room temperature. Modeling an H2TPD under the experimental conditions of Muhler et al. [6] with the coverage-dependent kinetic parameters yields trace B (solid line) in Fig. 1, independently of the reactor model used. The desorption sets in at about 210 K and reaches a peak maximum at 260 K. The broad shape of the peak with a FWHM of 68 K represents the adsorbate-adsorbate interaction incorporated in the kinetic parameters. However, choosing the initial value of the activation energy of desorption Eae8 = 77.4 kJ/mol and coverage independent preexponential factor Ades = 1.5 x 1011 molecules/(site-s) results in a symmetric second-order desorption peak with a maximum peak position at 300 K and an FWHM of 32 K (dashed trace C in Fig. 1) in good agreement with

394 the experimental result. Obviously, the TTPD evalutation method used by Anger et al. [7] overestimates the coverage dependence of the kinetic parameters and should be complemented by the complete method [21]. Changing the sites by an order of magnitude or varying the total flow rate of inert gas would result in a scaling of the effluent mole fraction of the desorbing species but not in a change of the shape and the peak position of the curve. This modeling result illustrates that the readsorption of H2 in this example is negligible, and consequently, that the CSTR model is sufficient when modeling the TPD experiments. In this case, it would be justified to derive the desorption parameters from TPD experiments with different heating rates (~1) by plotting l n ( T ~ a ~ / / 3 ) versus 1/Tma~. 3.2. N2-TPD from an iron-based catalyst In previous papers by the authors [9,10], the temperature-programmed desorption of N2 from an iron-based catalyst has been studied experimentally. The microkinetic analysis of these results is based on the kinetic simulation of ammonia synthesis by Stoltze and N0rskov [22-24] using the approach by Dumesic and Trevifio [2]. On Fe single crystal surfaces it was possible to detect a di-tr bound chemisorbed molecular precursor labelled or-N2 - 9 for N2 dissociation [25]. Hence, the dissociation step to atomic chemisorbed nitrogen is written in two equations: N2-4-*

~

N2-*

(1)

N2--*-+-*

~

2N-,

(2)

Fastrup [26] performed N2TPD experiments and temperature-programmed adsorption (TPA) experiments with nitrogen and integrated her results in a microkinetic model which is only valid for high-pressure ammonia synthesis. Table 1 Preexponential factors and activation energies of the Stoltze-NOrskov model for the rate constants ki = Ai e -E~/RT (* ( torr- s) -1 for forward reaction of step 1, s -x for all other elementary steps)[2] reverse rate constant forward rate constant step preexponential activation energy preexponential activation energy factor* (kJ/tool) factor* (kJ/mol) 1 3330 0.0 1.87-1014 43.1 2 4.29-109 28.5 1.32-109 155.0

The kinetic parameters used for the modeling have been taken from Dumesic and Trevifio [2] and are summarized in table 1. Simulating the onset of the experimental trace A (solid line) in the inset of Fig. 2 at 540 K exactly, a preexponential factor A,.ec = 2 x 10amolecules/(site.s) and an activation energy E,.ec = 146.0 kJ/mol for the recombination of N-, were found to be necessary. For the dissociation energy (activation energy of the forward reaction of step 2) a linearly increasing activation energy of Edis = F_~ + (20.0 x ON kJ/mol) was chosen incorporating the UHV results which show that the sticking coefficient strongly decreases with increasing coverage ON. For our microkinetic analysis we solved the following equations

395

describing a N2 TPD experiment: d ON2

dt d|

= =

dt

d PN2 dT OF, N,

OT

(~

-

~_1) 2

(-pN,

(~

9( , ~ -

-

~_~)

,_~)

FO _ v o ,

-

k l p N , O . -- k _ ~ O N ,

=

2

/

9~ k ~ o , , , o ,

-

-- k~ON, O. + k_~O~

k_~o;,)

,_,))

(,, _

N~ F, OF, N,

N,

o~

3o0 E Q.

Ind. Iron C a t a l y s t

~

I

-

"

1 b a r , 5 0 N m l / m i n He, ~ = 10 K/min :/'~

,,

,~ "~"~ ~

N2-TPD

t-

.o

----'-- CSTR - - " PFR

2oo

_.e 0

E

,=., c

9 100

=1:: 0

/

%

ul-";

z

1

5O0

,

,

i

~

L ,

,

8oo

Xempem~JreI 1

600

~,,

i

i

i

.

~ t

~

I

7OO

i

i

Temperature / K

i

K i

I

8OO

J

i

i

i

9O0

Figure 2. N2TPD from an iron-based catalyst. Comparison of the modeling results using either the transient CSTR model (solid line B) or the PFR model (dashed line C). The inset shows the experimental trace (solid line A) obtained by Muhler et al.[ 10]

Fig. 2 shows the modeling results using the design equation of a CSTR (solid trace B) or a PFR (dashed trace C) under the experimental conditions given in table 2 with a linear heating rate of 10 K/min. A detailed interpretation of the CSTR N2TPD results experimentally obtained with the KMIR catalyst was already subject of the paper by the authors [ 10,27]. They discussed their results in terms of coverage-dependent parameters using a simplified CSTR design equation. For both models the desorption sets in at 540 K. In contrast to the H2TPD on the copper-based catalyst, the desorption peak is composed of a first peak at 660 K (CSTR model) or 663 K

396 Table 2 Reactor data for N2TPD and TPSR Reactor volume Reactor pressure Inlet flow rate (N2TPD) Inlet flow rate (TPSR) Number of sites

on an iron-based catalyst and Cs-RtdMgO. 0.19 cm 3 0.19 cm 3 760 torr 760 torr 50 Nml/min He 50 Nml/min He 50 Nml/min H2 50 Nml/min H2 1.7 x 1019 1.47 x 1019

(PFR model) and a shoulder which is dominated by the delayed desorption of nitrogen due to readsorption within the porous catalyst bed. The influence of the reactor model is clearly visible at higher temperatures because of the more effcient elution in case of the PFR model. By heating with linear ramps of 1, 5, and 10 K/min, N2 TPD peaks at about 615, 647, and 663 K were obtained for the PFR model in good agreement with the experimental peaks at 615, 644, and 663 K. Moreover, changing the number of active sites or the total flow rate of inert gas by an order of magnitude resulted in changes of the position and shape of the shoulder. By plotting ln(T2==,/~) versus 1/Tr,.,a=, incorrect values for Aue, and Edes would be obtained, since even the position of the first peak is found to be influenced by the adsorbate-adsorbate interaction. Therefore, it is important to derive kinetic parameters not only from the shape of the experimental TPD trace or from varying the heating rate, but to confirm the result by either changing the reaction conditions such as the number of sites and the total flow rate of inert gas or by different kinetic experiments such as TPA and isotopic exchange reactions [5]. 3.3. TPSR experiments over iron-and ruthenium-based catalysts Fastrup et al. [9] presented a detailed experimental study of the hydrogenation of an iron-based surface precovered with atomic nitrogen using the temperature-programmed surface reaction (TPSR) method. Their modeling results clearly demonstrate the effect of the initial coverage of N - , and the effect of the H2 partial pressure. The kinetic mechanism of the TPSR experiment can be written as depicted below.

N--,+H--,

~

NH--,+,

NH-,+H--, NH2-, + H--, NH3--,

~~ ~

NH2-,+,

NH3-, + , NH3 + ,

~

2H--,

/-/2+2,

(3) (4) (5) (6) (7)

Based on the equations 3-7 we obtain the following differential equations for the coverages of the intermediates dON =

-

(r3

=

--ks|

-

(~

--

k3ON{~},

=

(~,

-

r_3)

dt d {~NH

-

+ k-sONnO, ~_~)

-

(~,

-

dt {~ NH 2

dt

-

k _ 3 { ~ N H { ~ , -- k4{~)NHO n -~- k _ 4 { ~ N H 2 { ~ ,

--

~_,)

~_,)

-

(~

-

~_~)

397

d

ONH s

dt

=

k 4 O N H O H -- k--4ONH2 O . -- kSONH= O H Jl- k - - 5 O N H a O .

=

(~

=

ks|174

=

2 (~

=

2 (k~p,,.o.

-

~_~)

-

(~

~_~)

-

-- k - 5 | 1 7 4

-- k6|

+ k--6PNHs|

doll

dt

-

~_,)

-

~ -

(k4ONHOH --

(~

-

~_~)

k_~o~,)

-

-

(~,

-

(k~o,o.

~_,)

-

-

(~

-

k_~ON,,O.)

~_~) -

k--4ONHaO,) -- (k5{~)NH:OH -- k--SONHsO,)

combined with either the design equations of the CSTR

d PNH8

(--PNH8 F ~ + p o .

dT

(r8 - r_8))

Ns3

dT

No3

where p~= denotes the inlet partial pressure of hydrogen, or the design equations of the PFR _ 3 c ~ F s NHa

O------T--~

=

c~Fs H2 o---~-

F8 c~Fa NH8

Fa

N,

N.

OS

(ro-r_~)

t s c~Fs H2 ~_ t s =

N,

os

~

(~'

-

~-'1"

Table 3 Preexponential factors and activation energies of the Stoltze-N0rskov model using the approach of Dumesic and Trevifio [2] for the rate constants ki = Aie -e~/nT (* ( torr- s) -1 for forward reaction of step 7 as well as reverse reaction of step 6, s -1 for all other elementary steps) step preexponential activation energy preexponential activation energy factor* (kJ/mol) factor* (kJ/mol) 3 1.83.109 81.3 1.15.10 r 23.2 4 1.31.10 aa 36.4 1.38.10 a~ 0.0 5 3.88.1013 38.7 2.33.10 la 0.0 6 3.67.10 a~ 39.2 2.38-105 0.0 7 9230 0.0 3.24.10 la 93.8

Fig. 3A (left half) shows the modeling results using either the CSTR model (solid trace A) or the PFR model (dashed trace B) under the conditions given in table 2 with a linear heating

398

A

!~

TPSR

400

B !,

1600

II II I I I I

i i i I i !

1400

I !

~ 1200

=

II i IronCatalyst

TPSR

I I i I I !

I

Cs-Ru/MgO

~ 1000 t

I I',

~ 20o

s Z

!

At

CSTR

k~

-'--" CSTR

- -

"

PFR

7

100

200 0

0 ,

300

400

Temperature/

500 K

;

300

.

.

.

.

t

.

.

.

.

1

350

Temperature/

~

400 K

Figure 3. N-,TPSR data obtained with an iron-based catalyst (left) and a Cs-Ru/MgO catalyst (fight). The modeling results demonstrate the influence of the reactor model.

rate of 2 K/min. For our calculations, we used the kinetic parameters listed in table 3 for the elementary steps. The initial coverage of atomic nitrogen was chosen to be 0.9. Using the PFR model results in a higher NHa TPSR peak and in in a lower | at higher temperatures compared with the CSTR model. With increasing initial lgN, the influence of the reactor model increases, and changing the number of sites does not scale the shape of the curve linearly in height. The experimental TPSR data obtained with Cs-Ru/MgO under the conditions listed in table 3 with a linear heating rate of 5 K/min are shown as trace A in Fig. 3B (fight half). The surface reactions sets in sharply at about 300 K and reaches its maximum at 325 K. The shape of the onset demonstrates the autocatalytic kinetics: preadsorbed N-, is hydrogenated by dissociated atomic hydrogen forming adsorbed NHa which desorbs thus creating free sites for H2 adsorption. For our calculation, we used the preexponential factors and activation energies given in table 4 under the experimental conditions mentioned above. A detailed discussion about the complete microkinetic model is given in reference [4]. The initial coverage of nitrogen was chosen to be 0.95. Traces B and C in Fig 3B illustrate the influence of the reactor model for Cs-Ru/MgO. In both cases, our model simulates the sharp onset of NH-, formation as well as the broadening of the shape of the TPSR curve at higher temperatures. The broadening is due to the readsorption of NHa being governed by the number of empty sites which is, in turn, determined by the H~ sorption equilibrium. Hence, changing from the CSTR model to the proper PFR model has a

399 Table 4 Preexponential factors and activation energies of the microkinetic model for Cs-Ru/MgO for the rate constants ki = Aie -E~/nr (* ( torr. s) -1 for forward reaction of step 7 as well as reverse reaction of step 6, s -1 for all other elementary steps)[4] forward rate constant step 3 4 5 6 7

preexponential factor* 6.0" 10 aa 4.7.1013 3.3.1013 5.9" 1013 7.3.104

, . .

activation energy (kJ/mo 1) 86.5 60.4 17.2 83.7 0.0

reverse rate constant preexponential factor* 2.8.1014 1.8"1013 9.3.10 lz 2.7.10 ~ 2.3.1013

activation energy (kJ/m o1) 41.2 8.6 64.6 0.0 89.4

significant influence on the height and the width of the TPSR peak.

4. Conclusions The microkinetic analysis of the H2 TPD data from Cu catalysts revealed that the influence of readsorption was negligible rendering the mathematically easier CSTR model appropriate. In this case, the proper kinetic desorption parameters can be extracted from the shifting TPD peak maximum temepratures as function of the heating rate. However, the example illustrated that deriving the coverage dependence from single crystal data is not without pitfalls. The microkinetic analysis of the N~ TPD data from a multiply promoted Fe catalyst showed that the PFR model is mandatory for the the assessment of the influence of readsorption. It is important to derive kinetic parameters not only from the shape of the experimental TPD trace or from varying the heating rate, but to confirm the result by changing the reaction conditions such as the number of sites and the total flow rate of inert gas. For the modeling of the TPSR data obtained with Fe- and Ru-based catalysts, the full set of 5 differential equations combined with the design equation of the PFR had to be solved. For Cs-Ru/MgO, significant changes in the TPSR peak height and shape were obtained when changing from the CSTR model to the PFR model due to the sensitive interplay of the various adsorbates.

Acknowledgement Some results were obtained with the optimization software developed by W.E. Stewart.

REFERENCES 1. J.A. Dumesic, D.F. Rudd, L.M. Aparicio, J.E. Rekoske and A.A. Trevifio, The Microkinetics of Heterogeneous Catalysis (ACS Professional Reference Book, Washington, DC, 1993). 2. J.A. Dumesic and A.A. Trevifio, J. Catal. 116 (1989) 119. 3. P.B. Rasmussen, P.M. Holmblad, T. Askgaard, C.V. Ovesen, P. Stoltze,J.K. N0rskov and I. Chorkendorff, Catal. Lett. 26 (1994) 373.

400 4. O. Hinrichsen, E Rosowski, M. Muhler and G. Ertl, Chem. Engng Sci. 51 (1996) 1683. 5. O. Hinrichsen, E Rosowski, A. Homung, M. Muhler and G. Ertl, J. Catal. in preparation. 6. M. Muhler, L.P. Nielsen, E. T6mqvist, B.S. Clausen and H. Topsc~e, Catal. Lett. 14 (1992) 241. 7. G. Anger, A. Winkler and K.D. Rendulic, Surf. Sci. 220 (1989) 1. 8. G. Ertl, in: Catalytic Ammonia Synthesis, 1st ed., ed. J.R. Jennings (Plenum Press, New York, 1991) p. 109. 9. B. Fastrup, M. Muhler, H.N. Nielsen and L.P. Nielsen, J. Catal. 142 (1993) 135. 10. M. Muhler, E Rosowski and G. Ertl, Catal. Lett. 24 (1994) 317. 11. O. Hinrichsen, Die mikrokinetische Modellierung der Ammoniaksynthese mit RutheniumKatalysatoren (Ph.D. thesis, TU Berlin, 1996). 12. R.A. Demmin and R.J. Gorte, J. Catal. 90 (1984) 32. 13. J.S. Rieck and A.T. Bell, J. CataL 85 (1984) 143. 14. P. Forzatti, E. Tronconi and L. Lietti, in: Handbook of Heat and Mass Transfer, 3, ed. N.P. Cheremisinoff (Gulf Publishing Co., 1988) p. 299. 15. A.R. Balkenende, J.W. Geus, A.J.H.M. Kock and R.J. van der Pas, J. CataL 115 (1989) 365. 16. , in: NAG Fortran Library ed. Numerical Algorithm Group, 256 Banbury Road, Oxford OX27DE and U.K. (Numerical Algorithm Group, 256 Banbury Road, Oxford OX27DE, U.K.,. 17. W.E. Stewart, M. Caracotsios and J.P. Sr Computational Modelling of Reactive Systems (Butterworth, Stoneham, England, 1996) p. in preparation. 18. C.W. Gear, IEEE Trans. Circuit Theory CT-I$ (1971) 89. 19. C.W. Gear, Numerical Initial Value Problems in Ordinary Differential Equations (PrenticeHall, Englewood Cliffs, New Jersey, 1971). 20. J.M. Campbell, M.E. Domagala and C.T. Campbell, J. Vac. Sci. Technol. A 9 (1991) 1693. 21. D.A. King, Surf. Sci. 47 (1975) 384. 22. P. Stoltze and J.K. Nc~rskov, Phys. Rev. Lett. 55 (1985) 2502. 23. P. Stoltze and J.K. Nc~rskov, J. CataL 110 (1988) 1. 24. P. Stoltze, Physica Scripta 36 (1987) 824. 25. E Bozso, G. Ertl and M. Weiss, J. Catal. 50 (1977) 519. 26. B. Fastrup, Top. Catal. 1 (1994) 273. 27. O. Hinrichsen, E Rosowski and M. Muhler, Chem.-Ing.-Tech. 66 (1994) 1375.

91997 Elsevier Science B.V. All rights reserved. Dynamics of Surfaces and Reaction Kinetics in Heterogeneous Catalysis G.F. Froment and K.C. Waugh, editors

401

A M i c r o k i n e t i c A n a l y s i s o f the R e v e r s e W a t e r Gas Shift R e a c t i o n E. Tserpe and K.C. Waugh Department of Chemist~, UMIST Manchester M60 1QD, England

Abstract

The previously published model of the forward and reverse water gas shift reaction on copper (1) has been made more realistic by the incorporation of an energy barrier for the adsorption of hydrogen on to the copper. An activation energy barrier for the adsorption of hydrogen of 12 k cal mol ~ reduces the steady state hydrogen atom coverage of the copper from 90% of a monolayer to <10% of a monolayer. Paradoxically this model predicts an increase in the rate of the reverse water gas shift reaction at all temperatures studied in the temperature range 300 K to 470 K. This results from the lower hydrogen atom coverage producing a greater amound of free copper area on which the CO: can decompose. Reducing the activation energy barrier to hydrogen atom adsorption to 9 k cal mol ~ increases the steady state hydrogen atom coverage of the copper to -30% of a monolayer, a value which is more consistent with the H : Cu ratio of 4 found by Topsoe and co-workers (13). This overall potential energy diagram, incorporating this 9 k cal mo1-1 activation energy barrier now predicts rates of the reverse water gas shift reaction in good agreement with experiment.

Introduction

We have previously shown that the kinetics of the reverse water gas shift (RWGS) reaction can be modelled on the basis of a simple redox mechanism (1). This result contradicted an earlier assertion deriving from a kinetic analysis of the reaction which claimed that the reaction proceeded through a formate intermediate (2). The individual elementary reactions of the redox mechanism, which in combination constitute the overall reaction, are listed in table 1, together

402 with the activation energies and Arrhenius pre-exponential terms for the forward and reverse components of each of the elementary reactions.

The potential energy diagrams for the

component complex reactions CO + O(s ),~ CO 2 4- Cu

and

H 2 0 + Cu ,,.-

H2 +

O(s )

A B

which constitute the redox mechanism are shown in figures 1 and 2. (In reactions A and B, O(s) is an oxygen atom on the surface of the copper.) Although this model gave an exact fit to the experimentally obtained temperature dependence of the rates of the RWS reaction (table 2), its use was merely illustrative, it being well understood that the model was deficient in several respects. The first recognised deficiency was that the activation energy for the adsorption of hydrogen (as atoms) on Cu was set to zero. This was the accepted view at the time of writing the model (3). However it was subsequently shown by Hayden and Lamont that the chemisorption of hydrogen on Cu was activated (4).

The

consequence of this omission on the results produced by the model might reasonably be predicted to be an overestimate of the hydrogen coverage.

A second deficiency relates to the formation of a surface carbonate on the copper. We have shown that the initial interaction of carbon dioxide with clean copper is dissociative, producing carbon monoxide and a surface oxidised copper (5). Subsequent adsorption of carbon dioxide onto this surface oxidised copper produces a surface carbonate species (6). The existence of the surface carbonate will block part of the copper and so reduce the overall rate of reaction. This was not considered in our previous model (1).

403 The extent of blocking of the copper surface by the adsorbed carbonate will be mitigated as a consequence of its being hydrogenated to form a formate species (also not allowed for) which on further hydrogenation produces methanol (6). This reaction sequence also is missing from the model, so that not only is the standing concentration of carbonate omitted from the model predictions, so too are the surface coverages of formate and methoxy species together with predictions of gas phase methanol concentrations. A final deficiency in the model relates to the detailed kinetics of the adsorption and decomposition of carbon dioxide to carbon monoxide and an adsorbed oxygen atom. Experimentally we have shown that the decomposition of carbon dioxide on Cu is precursor state mediated, having an activation energy of 3 k cal mol l for the decomposition of the weakly held precursor state CO 2 (CO2(a)) which is less than the heat of adsorption of the weakly held precursor state by 1 k cal mo1-1 (7). However, the model (table 1 and figure 2) shows an activation energy for the decomposition of the weakly held precursor state of 16.3 k cal mo1-1. This value of 16.3 k cal mol l, required by the combination of the kinetics and thermodynamics involved in the forward reaction 2 CO2(a) + Cu ~ CO(a ) + O(s )

in the model is virtually identical to the value of 16.7 k cal mol -I obtained by Campbell and coworkers for CO 2 decomposition on Cu (11 O) (8) and is consistent with our recent determination of an activation energy of approximately this value for the decomposition of CO 2 on the shift catalyst in the temperature range 300 to 470 K (9). The purpose of this paper is to address two of the omissions alluded to above. Activated adsorption of hydrogen on copper will now be included, as will the formation of a carbonate species on the copper surface. The inclusion of a complete, detailed reaction pathway from

Figure 1

P O T E N T I A L E N E R G Y D I A G R A M FOR T H E R E A C T I O N

CO § O(s) ~

-70

-

-80

-

-90

-

E = 19.1 kcal mol-1

C 0 2 + Cu

E = 16.3 kcal mol-1

Potential

Energy/ k cal mol-1

Potential

CO +

0~,~ CO(a)+;"0(~)

-

AH(CO(a)

= 13.2 kcal mol- 1

Energy/ k cal mol-1

C02(a) +Cu~

~~co~ +Cu AH(CO2(a)) = -4.3 kcal mol-1

POTENTIAL ENERGY D I A G R A M FOR THE R E A C T I O N H 2 0 + Cu ~--~-H2 + O(s)

Figure 2

Potential

Energy/ k cal mol-1

-40

-

-50

-

E = 3.8 kcal mol-1 E = 1.9 kcal mol-1

t

-60 -70 a)+ Cu AH(H20(a)) = -14.1 kcal mol-1

E = 7.2 kcal mol-~

E = 4.8 kcal mol-1

Potential

Energy/ k cal tool-1 'i~

OH(a)+ H(a ) + Cu

/ '

H2 + O(s)

H(a)+ H(a)+ O(~~) AH(H2(s) ) = - 14.4 kcal mol- 1

L~

4~

Table 1

Forward Arrhenius Parameters

Stoichiometry of

Reverse Arrhenius Parameters

Elementary Reactions Reaction NO Iog~o (A/units)

E/cal mol"1

Ioglo (A/units)

E/cal mol4

C O 2 + Cu ,~ CO2(a)

13

4300

1

13

0

2

21

16260

CO2(a) + CU(a ) ,~' CO(a ) + O(s )

21

19130

3

13

13152

CO(a ) ," C O + Cu

13

0

4

13

0

H 2 + Cu ,,,

13

0

5

21

0

H2(a) + Cu ,," H(a ) + H(a )

21

14347

6

21

7174

tt(a ) + O(s ) ,~ OH(a ) + Cu

21

4782

7

21

3826

H(a) + OH(a ),,'' H20(a ) + Cu

21

1913

8

13

14108

H20(a ),- H20 + Cu

13

0

H2(a)

The subscripts (a) and (s) refer to adsorbed and in the surface species respectively.

Tablc 2

Predicted Exit Gas Phase Concentrations and Surface Coverages obtained from a C0,/H2 (5% CO,, 95% H2). Feed Using the Model in Table 1

Exit Gas Phase Molecule ConcentrationIMol % T/K

co

co2

H2O

H,

396

1.795 x 10-3

4.998

I ,795 10-3 7.098 x 0.0 144 0.053 0.084

94.998

413

7.098 x

4.992

423

0.0 144

4.985

450

0.053

4.946

469

0.084

4.9 16

T/K

Surface Coverage I % Monolayer co co,/i 0-4

396

0.02

413

0.056

8.10

423

0.094

8.60

450

0.2 13

9.81

469

0.243

10.56

7.21

94.992 94.985 94.946 94.91 5

H2/10'5

H

OH

0

cu

5.78

97.9

0.016

0.00 1

0.18

8.15

96.8

0.055

0.03

1.8 2.91

0.293

9.80

95.9

0.107

0.07

3.58

0.62

15.3 1

92.6

0.339

0.32

5.96

0.689

20.16

89.8

0.482

0.57

8.2

H2O

0.068

P

s

408 adsorbed carbonate to methanol will be introduced in a later paper. It has been explicitly excluded in this paper since it is extremely complex, the result of which would be that its inclusion would obscure the conclusions to be drawn from the introduction of activated hydrogen adsorption and from the formation of a surface carbonate species. The overall purpose of this type of calculation is to produce a catalytic model which is fundamentally based and which therefore includes the measured kinetic parameters of all known processes occurring on the surface of the catalyst. Such a model ought therefore to be able to predict reaction rates under all conditions of temperature and pressure without the need of the use of adjustable parameters.

Activated Hydrogen Adsorption

Hayden and Lamont showed that the sticking probability of hydrogen (as atoms) on Cu was unity when the hydrogen was in the first vibrationally excited state (4). This state is approximately 12 k cal mol 1 above the ground state level. Additionally, the heat of adsorption of molecular hydrogen on Cu is -3 k cal mol l (10). Incorporation of these changes to the elementary reactions shown in table 1 necessitates setting the activation energy for the desorption of molecular hydrogen,

(H2(a)), E.4, to 3 k cal mol 1

and changing the activation energy for the

dissociation of H2(a)' (Es), to 12 k cal mol "l.

The effect of this on the potential energy diagram is to raise the position of the state H(a) + H(a) from -14.4 k cal mol "l to -5.4 k cal mol "l. Using the same value of-68.1 k cal mol 1 for the heat of adsorption of atomic oxygen (O(s)) as had been used previously (1), the position of the H(a)

409 + H(a) + O(s) state is now at -67.2 k cal mol 1 compared with-76.2 k cal mol 1 used previously. This now makes the reaction H(a) + H(a) + O(s) ,," H20(a ) exothermic by 4.7 k cal mol 1 whereas previously it had been endothermic by 4.3 k cal mol 1. The new potential energy diagram for reaction B is shown in figure 3.

This potential energy diagram has the overall activation energy for the hydrogen reduction of surface oxidised copper at 10 k cal mol l. This value is 1 k cal mol "l lower than that used previously, the value of 11 k cal mol 1 being that found experimentally for the reduction of surface oxidised copper by hydrogen (11).

Table 3 lists the elementary reactions for the overall reaction B together with their Arrhenius parameters.

The potential energy diagram for the forward and reverse shift reactions is

completed by adding figure 1 to figure 3, ie by adding reactions 1, 2 and 3 in table 1 to reactions 9 to 13 listed in table 3. The gas phase concentrations of reactants and products together with the surface composition of adsorbates predicted by this potential energy surface are listed in table 4. As before, these predictions of the extent of reaction are made by solving the Arrhenius parameterised elementary equations (in this case equations 1-3 and 9-13) by Newton Raphson iteration, a convergent solution being deemed to have occurred when the sum of the squares of the mass losses is less than 10-2~ The values of the gas phase and surface species corresponding to that convergent solution are those listed in table 4.

Table 3

Elementary Reactions of the Overall Reaction H20 + G u s H 2+ O(s) Involving Activated H2 Adsorption

Reaction No

Ioglo(A/units)

E/cal mol"1

Stoichiometry

13

0000

10

21

12000

H2(a) + C u ,,." H(a ) + H(a )

11

21

6000

H(a ) + O(s ) ," OH(a ) + Cu

12

23

4000

H(a ) + Ott(a ) ~ H20(a ) + Cu

13

13

14108

1t 2 + Cu ,,.- H 2 (a)

H20(a ) ,-, H 2 0 + Cu

Table 4

Predicted Exit Gas Phase Concentrations and Surface Coverages obtained from a C02/H2 (5% CO,, 95% H2). Feed Using the Model in Table 3

Exit Gas Phase Molecule ConcentrationIMol % TIK

co

co2

H2O

H2

396

0.0369

4.96

0.0369

94.96

413

0.048

4.95

94.95

423

0.056

450

0.08

4.94 4.92

0.048 0.056 0.08

469

0.10

4.90

0.10

94.90

TK

Surface Coverage I % Monolayer co co2 H2O

396

8.4

0.014

413

6.5

0.014

423 450 469

5.5 3.5 2.5

0.013 0.01 1

0.009

94.94 94.92

H2

H

OH

0

cu

28.4

0.054

6.8

13.3

2.6

40.4

20.9

0.054

7.03

12.8

3.3

49.4

0.053

7.04

12.3

3.8

54.1

0.048

6.71

10.4

4.9

64.5

0.043

6.29

9.0

5.5

69.8

17.2 10.1 6.9

P

c c

Figure 3

THE POTENTIAL ENERGY D I A G R A M FOR THE R E A C T I O N H 2 0 + C u , , " H 2 + O(s) WITH ACTIVATED H2 A D S O R P T I O N

10-

O

14.4

.~ ..... ~---

_

- 5.4 Potential

Energy/ k cal mol-1

,_L

-10-

H(a ) + H ( a ) + C u

,,....

AH(O(s)) = 61.8 k

-55 -57.8 . . . . . . -60 -

_

~,~o

8.7/ H(a) \~r+ H(a) + O(s)

"q~-- -~,-~r/ 6 v H(a) + H H20(a ) + Cu -65.9

-69.9

-67.2

413 The exit CO concentrations predicted by this model are in reasonably good agreement with experiment and with model 1 (table 2) at 423,450 and 469 K. They are however too high at 396 and 413 K. The overall activation energy predicted by the model is only 5 k cal mol l which is much lower than that found experimentally.

Examination of the surface coverages predicted by the activated hydrogen adsorption model is particularly helpful in revealing the major difference produced by including an energy barrier to the dissociation of H2(a). The hydrogen atom coverage drops from 90% to 7% and the amount of free copper surface increases to 40% to 70%. Qualitatively this was expected; the extent, however, was surprising.

At low temperatures (396 to 324K) the predicted hydrogen atom

coverage increases from 6.8% to 7.04% of a monolayer showing it to be kinetically controlled in this regime. Above 423K the coverage falls with increasing temperature due both to its increased rate of consumption by reaction and to its being at adsorption equilibrium. The low coverages are due therefore not directly to its adsorption being activated but to the lowered heat for atomic hydrogen adsorption of 5.4 k cal mol "1. This results from having maintained the desorption activation energy constant at the experimentally obtained value of 14.4 k cal mol l (12) (the previous heat of adsorption) this value being lowered by the difference between the activation energy to adsorption (12 k cal mo1-1) and the heat of adsorption of the molecular hydrogen precursor state (3 k cal moll).

Curiously, in spite of the hydrogen atom coverage being reduced by a factor of 10 and in spite of the detailed kinetics of reaction (A) remaining unchanged, the rates of reverse shift reaction predicted by this model are higher at all temperatures than those predicted by the non-activated

414 hydrogen adsorption model. The difference derives from the amount of free copper predicted by the two models. The larger amount of free copper (40% - 70% of a monolayer) predicted by the activated hydrogen adsorption allows for a greater degree of CO 2 adsorption and decomposition and hence the higher overall rates.

Topsoe and co-workers have found a maximum H : Cu ratio on dosing hydrogen alone on to copper of 0.4 at a dosing temperature of 248 K (13). H:Cu ratios of roughly this value can be obtained by ranging the activation energy to hydrogen adsorption due cognisance being taken of the higher temperature at which these experiments are conducted. An activation energy of adsorption of 9 k cal mol l (heat of adsorption of 8.4 k cal mol"l) gives a hydrogen atom coverage of 33% of a monolayer (H : Cu = 0.33) at 396 K. Lowering the energy barrier by 1 k cal mol 1, ie increasing the heat of adsorption to 9.4 k cal mol I increases the hydrogen atom coverage at 396K to 49% of a monolayer. On this basis the activation energy of 9 k cal mol "1 for hydrogen adsorption appears to be more consistent with experiment.

Carbonate Formation

The calculations described above clearly show that, at steady state, during the reverse shift reaction, part of the surface of the copper which is catalysing the reaction is covered with oxygen atoms to an extent which is dependent on the detailed kinetics of the adsorption of hydrogen. The adsorption CO 2 on to this area of surface oxidised copper will produce a surface carbonate (6). In temperature programmed desorption experiments this carbonate species desorbs as CO 2 at a peak maximum temperature of 333 K. The desorption/decomposition activation energy, Ed,

415

of the carbonate can be obtained by solution of the Redhead equation (equation C)

E.

A

e

-EJRT,

C

for an assumed value of 1013S"i for the desorption pre-exponential term and for a heating rate, 13, of 5 K "1 at the peak maximum, Tm, of 333 K (14). The value so obtained is 22 k cal mol 1. This can be incorporated into the detailed kinetics of the forward and reverse shift reaction (equations 1-3, 9-13) by the inclusion of reaction number 14. Reaction N ~ LQ-gl0 (A/units) 14

13

log 10 (A/S'I) E/cal mol "1

E/cal mo1-1 0000

C O 2 + O(s ) ,., CO2(a)

13

22000

Reaction 14 assumes that the formation of the carbonate species is non-activated.

The prediction obtained by including the non-activated formation of a surface carbonate species having a heat of formation of 22 k cal mol 1 is that, at steady state, the surface will be totally carbonated and that no reaction will proceed. This is an important conclusion that the carbonate species having a heat of formation of 22 k cal mol 1 will adsorb totally on the surface with the effect of poisoning its activity. It cannot act merely as a spectator. The carbonate species once formed must therefore be removed by further reaction, ie by hydrogenation to methanol. It is a corollary of this work then that methanol synthesis must always accompany the forward and reverse shift reactions.

416 References

1

Waugh, K.C.,1996, Chem. Eng. Sci., 51, 1533.

2

T. van Herwijen, T. and de Jong, W.A., 1980, J. Catal., 63,83.

3

Hadden, R.A., Vandervell, H.D. and Waugh, K.C. 1989, Proc. 9th Int. Cong. Catal., 4, 1835.

4

Hayden, B.E. and Lamont, C.L.A., 1989, Phys. Rev. Lett.,63, 1823.

5

Hadden, R.A., Vandervell, H.D., Waugh, K.C. and Webb, G., 1988, Catal. Lett., 1, 27.

6

Millar, G.J., Rochester, C.H., Howe, C. and Waugh, K.C., 1991, J. Mol. Phys., 76, 833.

7

Elliott, A.J., Hadden, R.A., Tabatabaei, J., Waugh, K.C. and Zemicael, F.W., 1995, J. Catal., 157, 153.

8

Nakamura, J., Rodriguez, J.A., and Campbell, C.T., 1989, J. Phys. Condens. Matter, 1 SB, 149.

9

Elliott, A.J., Sakakini, B.H., Yabatabaei, J. and Waugh, K.C. in preparation.

10 Hayden, B.E. private communication. 11 Chinchen, G.C., Spencer, M.S., Waugh, K.C. and Whan, D.A., 1987, J. Chem Soc. Faraday

Trans., 83, 2193. 12 Wachs, I.E. and Madix, R.J., 1979, Surf Sci.,84, 375. 13 Muhler, M., Nielson, L.P., Tornqvist, E., Clausen, B.S. and Topsoe, H., 1992, Catal. Lett. 14, 241. 14 Redhead, P.A., 1961, Trans. Faraday Soc., 57, 641.

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91997 Elsevier Science B. V. All rights reserved. Dynamics of Surfaces and Reaction Kinetics in Heterogeneous Catalysis G.F. Froment and K.C. Waugh, editors

419

Selectivity Enhancement in consecutive reactions using the Pressure Swing Reactor A. J. Kodde and A. Bliek * Department of Chemical Engineering, University of Amsterdam, Nieuwe Achtergracht 166, NL-1018 WV Amsterdam, The Netherlands Pressure swing operation can potentially be used to enhance yields and selectivity in packed bed catalytic reactors with either equilibrium limited or competitive sequential reaction paths. It will be shown that selectivity enhancement for a competitive sequential reaction is indeed feasible, but then using alternative cycles to the classical Skarstrom variant. Using mathematical modeling preferred operating regimes are identified and discussed. A classical plug flow reactor and a pressure swing reactor are compared on the basis of their selectivity-conversion behaviour and productivity. 1. I N T R O D U C T I O N Reactor performance improvement can be achieved by i n - s i t u separation. Equilibriumlimited reactions can be driven to completion by separation of the products and the removal of the intermediate product from the reactive phase in consecutive reactions may improve the selectivity towards this product. These principles have been demonstrated in membrane reactor technology[I]. I n - s i t u separation can also be achieved by selective adsorption on an adsorbent. Compared to a membrane, a mixture of catalyst and adsorbent offers a much larger interracial area although the component is not permanently removed from the unit and is obtained by regeneration. Regeneration is possible by lowering the overall pressure. With the alteration of high and low pressure a pressure swing cycle is imposed on the unit and the pressure swing reactor (PSR) is operated in cyclic mode. The PSR, membrane reactors, simulated- and countercurrent moving beds are all equivalent competitive techniques for i n - s i t u separation in catalytic reactors. Chatsiriwech et al.[2] have shown that the conversion of equilibrium-limited reactions can be increased in a PSR. In the present contribution we explore the possibilities to enhance the selectivity in consecutive reactions through pressure swing operation. Potentially interesting applications include hydrogenations which are catalyzed at temperatures suitable for adsorbents. For cyclically operated systems, such as the PSR, the number of operating variables is substantial and its characteristic cycle invariant state(CIS) cannot be determined in a straightforward manner but requires some iterative method. Thus determination of individual parameters requires excessive simulation work. In earlier work a general framework has been developed for the design of PSA cycles for adsorber columns[3]. Phenomena like *corresponding author E-mail: [email protected]

420 separation reversal[5], where the product composition is determined by reaction kinetics rather than adsorption equilibria, indicate that different design rules apply for PSR as compared to PSA units. YY~ will show that for a consecutive reaction a traditional PSA cycle does not yield selectivity enhancement compared to a steady state plug flow reactor(PFR) in contrast with the use of a cycle never reported in PSA design. The influence of the key operating parameters is presented.

2. M O D E L D E V E L O P M E N T Factors which influence the performance of a PSR unit are treated in the section below. Each is evaluated on its relevance to this study. Although all pressure swing cycles have common characteristics, there can be one or more pressure-equalization steps, co- or countercurrent, blowdown-, purge- and backfillsteps and arrangements ranging from 1 to 12 beds. In this exploratory study the starting point is the single bed version of the classical four step Skarstrom cycle[4], where part of the product gas is stored in a stirred storage tank and is used as purge gas later on. An alternative cycle with co-current blowdown and purge steps has also been investigated. The packing of the unit includes both the ratio of adsorbent and catalyst and their packing through the bed. A non-uniform packing is easily realized but its influence is hard to predict a priori. A perfect mix of the catalyst and adsorbent is assumed. The adsorption window is demarcated by the adsorption equilibrium within the boundaries of the imposed pressure cycle and is determined by the multi-component adsorption isotherm and the pressure range. The Langmuir type isotherm is assumed. In pressurization and depressurization steps, the amount of gas leaving or entering is determined by the pressure change which was set at a constant rate. The flow rate during the adsorption step determines the conversion range and was set to ensure a low conversion. The purge gas flow rate is expressed relative to the adsorption gas flow rate. Fast sorption rates were assumed, thereby closely approaching the instantaneous equilibrium case. Mass transfer is described by the linear driving force model. The reaction kinetics dictate the intrinsic selectivity of the system We opted for a simple model system in which both reactions are first order in the reactant concentration and both elementary reaction rate constants are equal. Physically this implies that the catalyst is always in quasi steady state. The extreme case of adsorbent selectivity, where solely the intermediate product adsorbs is investigated here. Heat effects have a very complicated influence on the reactor performance as the adsorption isotherms and reaction kinetics are strongly non-linear functions of temperature. Isothermal operation is assumed to facilitate interpretation. In the previous section the following assumptions have been made" (1) Perfect mixing of the catalyst and adsorbent particles, (2) The adsorption isotherm is of the Langmuir type, (3) Mass transfer resistance is restricted to the external film, (4) Reactions are first order in reactants concentrations, (5) The intermediate product adsorbs exclusively, (6) Isothermal operation. Supplementary to the choices described above, the following assumptions are made: (7) Ideal gas law is obeyed, (8) Axially dispersed plug flow in the bed, (9) Negligible pressure drop over the bed, (10) Bed of constant voidage, bulk density

421

Table 1 Boundary conditions in the PSR models Skarstrom cycle PFR & depres, purge pres. adsorption z=L z=O z=L z=O z=L z=O z=L z=O

(~)

(c)

(~)

(c)

(b) OP

(d)

(b)

(e)

0t -- Cup

P-

Phigh

Equations: Oyi

(~) Dz 52z

(c)

(c)

(b) (d) aP -- Cdown Ot

(c)

(~)

(b)

(e)

P-

Plow

Oy -- ?2(Yi- Yi,in) (b)" equation 3 (c) Ozz

-

alternative cycle depres, purge z-O z=L z=O z=L

(c)

(c)

(d) (b) 0P --- Cdown 0t

(~)

(c)

(e)

(b)

P-Plow

0 (d)" v - 0 (e)" v - Vfixed

z

and particle size, (11) No radial concentration gradients, (12) Feed gas consists of pure A. The PSR is compared to a P F R with identical catalyst weight and hourly space velocity. Reactor characteristics are quantified in terms of the conversion of A, the selectivity towards B and the productivity. The following equations hold.

(~ + ~ ( 1 - ~ ) ) ~

-

Rgas T

+

Dz -~j

Rgas T ] + PbedJi

+

R/

(1)

i=I...N-1 YN

=

1-

N-1 ~ y~ i--1

(2)

1 OP (~b + c~(1 - ~b)) Rg~T Ot -

3Di,m

S~ -~b - 5 2

N P Ov ~- E ~,Pbed~)i /'^ cg--+b nu Pc/) R g ~ T Oz i=1

.

(3)

(4)

(~ - ~ )

TparticleCP

(5)

mbByBP riB-- I + bByBP RA -- - k y A P

RB -- + k y A P -

kyBP

R c = +kyBP

(6)

The equations describing the solid phase are also valid at z=O and z=L, but for the gasphase the boundary conditions differ for each process step. At the end of the reactor three different cases occur: there can be a feed flowing inwards, the end can be closed or there can be an outlet flow. Each case is described with the appropriate Dankwerts conditions and the pressure ramp is imposed. The resulting equations are listed in table 1. The conversion is given by X A __ f7 ~--~streamsnvyAPdt f7 VfeedYA,feedPfeed dt

(7)

422 Table 2 Operating parameters for the PSR in all Parameter Unit Value bs Pa-1 1.36 10 .6 DB,m m2s -1 10 -8

Dz

m 2 s -I

I0 -s

k m

mol m - l P a -1 mol kg -1 Pa Pa m

10 -6 4.458 3 105 1 105 10 - 3

Phigh Plow rparticle

figures unless explicitly noted. Parameter Unit Value T K 298.15

tII

S

10

tw Vfeed Vpurge Cb ~p Pbed

s

I0

m s -1 m s -~ kg m -3

5 10 -a 5 10 -2 0.35 0.85 592.62

where n is +1 for in going and -1 for out going streams. The selectivity is defined as

f~- ~"~streamsnvysPdt SB = ~ Estreams nvyAPdt

(8)

The productivity is given by"

QA -- --7" 1 f VfeedYA,feedRga~T P dt

(9)

The P F R is described by equations 1 to 6 and the boundary conditions for the adsorption step with the accumulation terms set to zero. The entire conversion range was covered by varying the superficial gas velocity. The model was solved with the 9PROMS package running on a IBM RS6000. The axial axis was discretized using second order orthogonal collocation on 20 elements. The resulting set of ODE's was solved by numerical integration. Starting from an initial state where the gas phase was either filled with a pure inert or component A, the CIS was calculated by successive substitution and convergence checked by the overall material balance. The CIS was typically reached after 50-100 cycles. 3. R E S U L T S

Proper comparison between the PSR and PFR takes into account conversion, selectivity and productivity. In the present contribution selectivity will be emphasized as improvement of this parameter is sought. 3.1. S e l e c t i v i t y S k a r s t r o m cycle We started out to analyze the PSR operated according to the classical four step Skarstrom cycle as depicted in figure 1 for various cycle times. The selectivity decreases compared to the PFR. The highest conversion is reached in the steady state limit (tn,tw ~ oo). In the lower limit (tiI,tiy $0) hardly any sorption of B takes place. All shown cases have similar characteristics and the behavior of the PSR under the Skastrom cycle was further explored by investigating a single case. In figures 3 and 4 the trajectory of component B is shown. During pressurization the feed purges the feed end

423 HP

1 O0

roduct

__~__ t i

f 1

S

95

l

90

85 Ill

IV

"*',.,,.~

80 i

0 Feed

LP product

Figure 1. Single bed version of the Skarstrom cycle. Cycle steps: I: pressurization, II adsorption, III depressurization, IV purge, S: stirred storage tank. The PSR is filled with a mixture of catalyst (o) and adsorbent (-).

10

1

20

30

X A / [%] Figure 2. Selectivity towards B versus conversion of A in the PSR operated under the Skarstrom cycle. Line: PFR, A: PSR, Parameters: table 2, Vpurge= 5 10 -3 m s -1 and from left to right: t i i = t i y = O, 5, 10, 30.60. 120, 240, 500, cos

and concentrates B at the product end of the reactor. This process continues during the adsorption step as B has to be formed before it can adsorb and a linear increase of yB over the unit is reached. As opposed to the ordinary PSA profile B desorbs over the first two-third of the reactor as shown in figure 4. The countercurrently operated blow down and purge steps are very inefficient as B has to pass the relatively clean adsorbent at the feed end on its way to the reactor exit and migration of B is retarded by equilibration. As a result B is concentrated in the middle of the reactor and the decreased selectivity towards B results from the increased residence time of B. As B is concentrated at the product end of the unit a co-current blow down may remove it more effectively. Co-current purging will further prevent readsorption at the feed end of the unit and thus reduce the amount of B purged during the subsequent pressurization. In PSA adsorbers co-current depressurization steps are used to recover the mechanic energy of the high pressure gas and increase the purity of the adsorbed component, but never to obtain the adsorbed component. This alternative cycle employing pressurization, adsorption, co-current depressurization and co-current purge steps has no analogue in PSA separation. Alternative cycle The selectivity of the PSR under this alternative cycle can indeed by lifted above the steady state limit as shown in figure 5 and 6. The characteristics of all cases shown is qualitatively similar and one case will be explored further. Selectivity enhancement is not reached in all cases which shows that the choice of operating parameters is critical for this cycle. The influence of individual operating parameters will be discussed to identify the 1The small maxima in YB at the start of the cycle in figures 3 and 7 result from the gPROMS reinitialisation algorithm.

424 0.25 -~

0.35 0.30

0.20-

t..___l

0

0.15

~,----~ "

0.150.10-

0.10 0.05 E

0.05

O.O0~-'er--7 0.0 0.2

(

, 0.4

: 0.6

,.., 0.8 1.0

z/L / [-]

Figure 3. The trajectory of component B under the Skarstrom cycle in the gas phase. Parameters: table 2, t z i = t z y = 30 s, Vpurge= 5 10 -3 m s -1 Legend: +: Start cycle 1, end of A: pressurization, o: adsorption, ~: depressurization, D: purge.

0.00 0.0

,,

0.2

0.4

0.6

0.8

1.0

z2L / [-]

Figure 4. The trajectory of component B under the Skarstrom cycle in the solid phase. Parameters: table 2, t ; I = t i v = 30 s, Vpurge= 5 10 -2 m s -1. Legend: End of o: adsorption. []: purge.

operating window for selectivity enhancement. The nearly linear increase of YB shown in figures 7 resembles the P F R case. The purge effect of the feed is clearly illustrated by the drop of nB during pressurization in figure 8 and is much smaller than under the Skarstrom Cycle. During the adsorption step the increase of n s and the limited rise of YB point at effective removal of B from the gas phase. During depressurization, desorption of B causes an increase of the gas phase fraction which is effectively removed by co-current depressurization and purge steps. Readsorption of B from purge gas is limited as illustrated by the marginal increase of nB near the feed end. The simulations show that during the adsorption step, nearly over the whole reactor length, approximately 55 % of the B formed on the catalyst is adsorbed on the sorbent. As the mass transfer rate during the adsorption step hardly decreases over the adsorption interval (figure 9) and "the reaction rate of the first reaction hardly drops, this percentage adsorbed on the surface remains fairly constant. The system remains far from saturation as the gas fraction rises to 5% at the reactor end compared to 15% if no adsorption would take place. Thus the capacity of the adsorbent is much larger than the amount formed by reaction. As a result adsorption takes place right after formation over the entire reactor length instead of over a limited region in the case of a moving concentration front. Figure 9 further shows that the amount adsorbed depends on the coverage of the adsorbent at the start of the high pressure half-cycle as a more efficient purge gives a higher mass transfer rate during the adsorption step. Series A and B in figure 5 represent cases where the sorbent is regenerated by depressurization only. The highest selectivities correspond to the highest mass transfer rate. In series F and G shown in figure 6, the adsorption and purge times are varied at constant purge ratio of 1/3. The highest conversion corresponds to infinite time intervals.

425 98-

97.5

97-

95.0

96-

92.5-

r~

9594 <..

,

.

7

8

9

.

. 10

. 11

90.087.5 4,..,

12

5

10

X a / [%]

,

,

7

15

20

25

x ~ / [%]

Figure 5. Selectivity towards B of the PSR under the alternative cycle. Line" PFR, Markers: PSR. Parameters: table 2; Series: A(A)" tsss=10 s, t s v = 0 s, DB,m=lO-9,10-Sm2s-1; B(~)" ttsi=20 s, tlv=O s, DB.m=10-9,2 10-9,10-8m2s -1"

Figure 6. Selectivity towards B of the PSR under the alternative cycle. Line- PFR, Markers: PSR Parameters: table 2, Series: F(+)" Vpurge=5 10-3m s-l, tzv=tss=30,60,120,240,ec s; G(E:]):

Qv=O.1.tiI=3,6,12,24,48,96,oc s

c(+)- tzv=0,3,6,9 s; D(C~). Phigh= 3,5,9 105 Pa; E(I~)- R = 1 / 3 , t~v=3,6,12,30 s

0.15

0-127

_..~

0.08

,----, 0.10,.2..,'

~176 ~176 2e-

0.05

0.02 .~~-', 0.00 0.0

0.00 c 0.2

0.4

0.6

0.8

1.0

z/L / [-]

Figure 7. Trajectory of YB in the PSR under the alternative cycle, Parameters table 2, Legend: +: Start cycle, end of A: pressurization, o: adsorption, O: depressurization, [:3: purge

0.0

0.2

0.4

0'.6

o.s

1'.0

z/L / [-]

Figure 8. Trajectory of nB in the PSR under the alternative cycle, Parameters table 2, Legend: end of A: pressurization, o: adsorption, [:3: purge

426 0.0004e~o

0.0003

o

<>I 0.0002 A

L.....a

0.250.20o

0.150.10-

O.0001

=

al

0.0000 0.0

0.2

0.4

0.6

0.8

1.0

z/L/[-] Figure 9. Mass transfer rates during the adsorption step in the PSR under the alternative cycle, Parameters: two cases from series C (figure 5) 9 R = l / 3 , A: R=I; Legend marker: open: start adsorption step, closed: end of adsorption step

0.05~ 0.00 0.0

. 0.2

. . . 0.4 0.6

. 0.8

1.0

z/L / [-] Figure 10. Limits of solid phase loading in the PSR under the alternative cycle. Parameters: (O): one case from series C (figure 5), R=1/3, A, as O but m=8.916 tool kg -a. Legend marker: open: end of purge step, closed: end of adsorption step

The short residence time during the purge step (series G) is necessary to lift the selectivity above the steady state limit. In series C the purge ratio is varied between zero and one. Both the conversion and selectivity increase with the purge ratio, but an increasing fraction of the product is obtained at low pressure. In series E, the purge time is increased and simultaneously the gas velocity is lowered to keep the purge ratio constant. A longer purge time improves regeneration, but the increased residence time decreases the selectivity. The latter effect is illustrated here as this series and the PFR show a similar trend regarding their conversion-selectivity behavior. When the difference in the equilibrium loading at high and low pressure increases, the driving force for both adsorption and desorption is enhanced. Doubling the saturation capacity nearly doubles the available adsorption window but does not significantly alter the performance of the unit (figure 10). Increasing the pressure range will increase the selectivity, but conversion drops(series D). Since the conversion is related to the amount fed at high pressure, the increased pressure difference results in a lower conversion during the low pressure half-cycle. Furthermore the amount of feed gas needed to pressurize the unit to the upper pressure increases and thus leads to lower conversion. As the selectivity hardly depends on the equilibrium loading window of the sorbent, the increased selectivity may be attributed to the improved regeneration during the adsorption step as the amount of gas blown out during the depressurization increases with the pressure range. The influence of the operating parameters discussed before illustrates that. selectivity

427

2.0-

r~

eq

i

1.5-

E o

1.o-

< 0.5-

cy 0.0

0

5

10 X A/[6~

]

7

,

i

15

20

25

Figure 11. Productivity of the PSR versus PFR. Legend Line: P F R Markers: PSR, O: Skarstrom cycle (figure 2), Alternative cycle: A - s e r i e s A, ~" series B. +: series C, ~ series E from figure 5, [:3: series F, V: series G from figure 6

improvement requires a fast purging of the gas and high mass transfer rates to realize an efficient regeneration with a minimal increase of the residence time. In that case the selectivity improvement realized during the high pressure half-cycle by adsorption is not counteracted during the regeneration half-cycle. 3.2.

Productivity

The productivity of the PSR is lower than the PFR (figure 11). During the regeneration there is no feed to the unit and the purge gas velocity is set to minimize reaction. If expressed per units of reactor volume the difference will even be greater as the PSR also contains adsorbent. 4. C O N C L U S I O N S When a PSR unit is operated according to a Skarstrom cycle, the feed and product end of the reactor are periodically cleaned as they are purged with feed and high pressure product gas respectively. As a result the intermediate component is concentrated in the middle part of the reactor and its residence time is increased which in turn decreases the selectivity of the overall unit towards this product. When using co-current depressurization and purge steps, the concentration of the intermediate product increases towards the product end over the whole cycle. Selectivity enhancement is obtained provided that effective regeneration can be accomplished without significant increase of residence time. The catalyst loading based productivity of the PSR is always lower then the PFR as there is no continuous feed to the unit. 5. N O T A T I O N

bi Cdown Cup

Di,m Dz

adsorption affinity of component i rate of pressure rise during pressurization step rate of pressure decline during depressurization step gas-phase diffusivity of component i in the mixture axial dispersion coefficient

Pa -1 Pa s -1 Pa s -1 m2s -1

m2s -~

428 k m N ni P

QA rparticle

Rga~ /~ S g->b SB t ti T v

XA yi z ~b cv ~/9bed

reaction rate constant of reaction step I and II saturation capacity number of components solid phase loading of component i pressure productivity (def: eqn.9) particle radius universal gas constant mol based reaction rate of component i sorption rate of component i selectivity towards B (def: eqn.8) time cycle step time of step i temperature superficial velocity conversion (def: eqn.7) molar gas fraction of component i axial distance interparticle void fraction intraparticle void fraction total cycle time packed adsorbent density

mol j-is-1 mol m - 3 mol kg -1 Pa mol m-2s -1 m J mol-1 Kmol kg~-dls s -1 mol m -3 s -~ s s K m s -1

m

s kg m -3

Subscripts i I feed in

component cycle step feed stream inflowing stream

Superscripts 9

equilibrium

REFERENCES 1.

S. Agarwalla and C. R. L. Lund. Use of membrane reactor to improve selectivity to intermediate products in consecutive catalytic reactions. J. Membrane Sci., 70:129144, 1992. 2. D. Chatsiriwech. E. Alpay, L. S. Kershenbaum, C. P. Hull, and N. F. Kirkby. Enhancement of catalytic reaction by pressure swing adsorption. Cat. Today, 20:351-366, 1994. 3. D. M. Ruthven, S. Farooq, and K. S. Knaebel. Pressure Swing Adsorption. VCH Publishers, 1994. 4. C . W . Skarstrom. Method and apparatus for fractionating gaseous mixtures by adsorption. U.S. Patent 2,944,627 to Esso Research Engg., 1960. 5. G. G. Vaporciyan and R. H. Kadlec. Periodic separating reactors: Experiments and theory. AIChE Y., 35(5):831-844, 1989.

91997 Elsevier Science B.V. All rights reserved. Dynamics of Surfaces and Reaction Kinetics in Heterogeneous Catalysis G.F. Froment and K.C. Waugh, editors

429

Experimental studies of transient thermal effects during catalytic oxidation in a packed-bed reactor S. Marengo, P. Comotti. S. Scappatura and M. Vasconi Stazione Sperimentale per i Combustibili Viale A. De Gasperi 3, 20097 San Donato Milanese, MI. Italy

The dynamic phenomena associated with the rhodium-catalyzed oxidation of carbon monoxide, methane and propane have been studied by in-situ infrared thermography. Highresolution temperature maps of the reacting catalyst revealed the mobility of the reaction front during ignition and extinction of the CO oxidation, and the development of thermokinetic oscillations. The catalytic oxidation of methane and propane produced weaker dynamics. Chemisorption and kinetic experiments suggest that the competitive adsorption of the reactants and the occurrence of self-inhibition, represent key factors in the development of the observed transient effects.

1. INTRODUCTION The dynamics of exothermic catalytic reactions has been the object of investigation for almost half a century comprehensive updated studies have been recently reported [1, 2]. Whereas various models have been proposed in a large number of papers, describing mathematical and theoretical approaches, experimental data concerning these phenomena are relatively scarce, due to the difficulty of performing accurate measurements under reaction conditions. Novel in-situ techniques capable of providing reliable spatial information on reactor behaviour can thus represent a valuable tool for extending our knowledge in this field and improving the efficiency and the safety of a number of catalytic processes which are carried out in the ignited state [3, 4]. More efficient experimental methods may also favour the development of new techniques of reactor operation, based on a periodic change of the process parameters [5]. Infrared imaging was utilized in several studies of spatial effects in exothermic catalytic reactions over model catalysts, such as isolated particles, wafers, plates, discs [2]. Our approach has been to characterize the catalysts directly in a packed-bed microreactor, under realistic reaction conditions. In-situ measurements by infrared thermography of the adsorption properties of catalytic materials have been previously reported [6]. In the present study, the catalytic oxidation of compounds having different chemical properties was investigated by the same technique, with the aim of obtaining comparative data useful to better understand the factors governing the complex phenomena associated with catalytic combustion.

430 Experimental data describing the dynamics of ignition and extinction will be reported in detail, and the principal factors determining the behaviour of the different reactants will be analysed.

2. E X P E R I M E N T A L The catalyst was Rh on T-alumina in the form of particles with diameter of 0.1-0.3 mm. In some measurements, supported Pt or Pd were also utilized. The catalyst was packed in a specially designed, 8-mm i.d. tubular reactor, between two layers of inert material. A forced ventilation oven allowed control of reactor temperature. The expressions "reactor inlet" and "reactor outlet" adopted in the text to describe the reaction front motion, are referred to the catalyst bed only, without taking into account the inert layer. High-resolution thermal maps of the surface of the catalyst bed (8-15 mm high) were recorded with an Agema Thermovision 900 apparatus, equipped with optic for close-up view and with a data processing unit for real-time image analysis. An axial multiple thermocouple placed in the catalyst bed allowed the measurement of internal temperature. Catalytic experiments were carried out at 0.1 MPa total pressure, with molar fraction (y) of CO, CH 4 or C3H 8 between 0.01 and 0.04, oxygen slightly above stoichiometric value, and Ar or He as balance gas. Reaction products were monitored by on-line analysis with a gas chromatograph and a quadrupole mass spectrometer. The peculiar aspect of this study is represented by the in-situ measurement of thermal effects produced throughout the reactor under real operating conditions. In this respect, it is important to verify that the thermal map of the surface of the catalyst bed, obtained by thermography, describes reliably the phenomena occurring in the bulk of the bed. For this purpose, the external temperature profile was compared with the profile obtained by an axial multiple thermocouple placed inside the catalyst bed. It was observed that both in the steady and in the transient state, the two profiles have similar shape, although the temperatures are not identical due to axial gradient.

3. RESULTS

3.1. Ignition It is known that the catalytic combustion of hydrocarbons exhibits ignition-extinction phenomena which depend on feed composition [3]. The ignition experiments reported in this study were performed with inlet fuel concentration comprised in the region of surface flammability, in which hysteresis effects occur; typical values were: YCH40.04, YC3H80.02, YCO 0.03. The superficial velocity (u) of the reactant gas was between 1 and 5 cm/s. In order to minimize the effects of external factors on the dynamic behaviour of the reaction system, no external perturbation was introduced during the tests, with the exception of a slow change in the heating rate. Ignition was obtained by gradually raising the reactor heating rate, while keeping feed composition and flow rate constant. In the CO-O2 reaction, light-off was revealed by a fast temperature rise near the reactor outlet; in the same time, CO conversion jumped to about 100 %. Figure 1 shows the

431 temperature map of the surface of the catalyst bed and the temperature profile along a vertical line on this surface.

i~ ~~ i~iiiiii!iiiiiiiiiiii!i!iiliiiiiiii i ii!ii i iil ii i ii! 1 1 ~~ ~iiii~'.ii~i!:i~ilii~~ii:il!i-:;i~,ii!,i!li,];:i;,~!i~i~ii~i=!!i~!:i!i~illiiiII!i~!i!=i:ii~!-iil:~,il/!~;iI il!i=iii!lX!illI,ii ! !

Figure 1. Temperature map and profile of the catalyst bed upon ignition of the CO-O2 reaction. Downward flow, u=4.8 cm/s; Yco =0.03.

Figure 2. Temperature map and profile in the ignited state for the CO-O2 reaction.

Figure 3. Temperature map and profile during extinction of the CO-O2 reaction.

432 Ignition was followed by upstream creeping of the reaction front, with a linear velocity comprised between 0.8 and 1.2 mm/s. In the ignited state, the reaction front was stabilized near the inlet (Figure 2), where a sharp temperature profile and significant radial gradient (internal catalyst temperature was 20 ~ higher than at the surface) were measured. When external heating was lowered, the maximum temperature in the bed decreased, the reaction front broadened and moved downstream, eventually reaching the end of the bed, where extinction occurred (Figure 3). Such a behaviour seems to be a characteristic of the CO-O2 reaction, since it was observed in our study with different active metals (Rh, Pt, Pd), different loadings, various catalyst shapes and with different superficial velocities and flow directions (upward and downward). Similar dynamics has been reported for the CO oxidation on Pt/alumina in a larger tubular reactor [7]. Varying inlet CO concentration did not alter substantially this dynamics; however, for Yco below 0.02, the reaction front after ignition moved only up to the middle of the bed, were it was stabilized or moved back again, depending on heating rate. Between ignition and extinction, temperature oscillations were observed in the active portion of the catalyst bed. The oscillating patterns were complex and multiform, depending to a large extent on the catalyst composition and reaction conditions. When, starting from the ignited state, external heating was lowered, the amplitude of oscillations increased and the period decreased, until extinction took place. Real-time analysis of the temperature map revealed synchronization of oscillations in a portion of the reaction front (Figure 4, points 1, 2). In the downstream zone (point 3), oscillations in opposite phase were detected, revealing influence of mass transfer on the dynamic phenomena. In the same time, CO2 and O2 concentration in the effluent gas oscillated in opposite phase, with the same period as the slow temperature fluctuations.

Figure 4. Local temperature oscillations during CO oxidation over 3% Rh/AI203. Yco =0.03, u=4.8 cm/s. Points 1, 2, 3 are on a vertical line at 1-mm distance.

433 In the combustion of methane, a different transition to the ignited state was observed. An incipient reaction front formed near the reactor inlet, due to release of the heat of reaction. Upon ignition, the reaction front developed slowly, holding its position and showing no significant dynamics. Propane exhibited a behaviour similar to methane. However, the initial reaction front formed slightly downstream of the reactor inlet, and shifted upstream to the very initial part of the bed (Figure 5). The radial gradient in the active zone was remarkable (40 ~ between axial and external position). With both methane and propane, lowering heating caused extinction near the inlet, without any significant shift of the reaction front.

320 1 310 G" o

300

290 280 E

~.

s

270 260 250 0

2

, 4

, 6

Axial position (mm) Figure 5. Temperature profiles during the ignition of the C3H8-O 2 reaction over 3% Rh/AI203. YC3H8=0.02, Yo2 =0.12, u= 3.3 cm/s.

In propane oxidation, only weak kinetic oscillations were revealed by the spectrometer signal of CO2 at reactor outlet. The combustion of methane under similar operating conditions showed no oscillating behaviour. 3.2. Adsorption and kinetic measurements

To investigate the factors controlling the catalytic combustion in the reaction systems investigated, adsorption measurements of the different reactants were performed at room temperature on freshly reduced Rh/A1203 (Rh 1%). A series of 3-gmol pulses were injected into an Ar stream flowing through the reactor with u=l cm/s. The pulses of CO produced a significant thermal effect in definite portions of the catalyst bed (Figure 6). The rise of the catalyst temperature was about 2 ~ after each pulse. Interestingly, a peak at mass 2 was observed by the quadrupole analyser during CO adsorption, indicating displacement of H2 by CO. Pulses of oxygen on the reduced catalyst in the same conditions, produced much stronger thermal effects than CO (Figure 7). The temperature rise after each pulse was about 8 ~ On

434 the other hand, the portion of catalyst bed involved in the adsorption of one pulse of 02 was larger than with CO, as revealed by the position of the temperature peak.

i!iiii~ Jii~i:iliJiii~!}i ii~i~:!i:li~i:i}ii~}i}iiiiii!: }iJi ii!i}:i!i};iiiiiiil~'i!iliil!ili~;l~i;iii~iiiiliiiiiii!iiili~ili;itiiiili;}iii!i}i~i!i!ii!iiiii!iiiiiiii!{iiiiiiili~ii!iiiiiiiiii i!iiiii iiiiiiiiiiiiiilliiilliiiiiiii:i~!i!i:!}i!i}}iii!i}iiii i~i~i!i:iiill~iiiiiiliiil~}iiiiili}i~ili!iii!}ii~!i!iiiili}ii

::[i]iiii::!i~ii)iii~1iiiiiiiiiii::iii::iiiiiiii::::::;i)iiiiilili i;iiiiiilliliiii: il}i::iiiiiiiii::!i;::}::!::i! i:il)iiiigiiiiiii::i::iiii;iiiii!::2i::iii::;ii:::::: :ii•i•g•ii•i:;ii:i:•:i:•i••i:•::;•i:•i:i::•i••i:•i•:i:•:!:ii}:•:i:ii!:•i:::i:i•:iii:ii:::::i:. i:•i:.i:::::::::::::::::::::::::::::::::::::::::::::::::: :::::: i•:i!i:•i:!i::::i::::i:::~:: ii::::: iii::i!ili~!i]i::i~:: iiiiiiii~;!

Figure 6. Thermal effect produced by injection of a 3-gmol pulse of CO onto reduced RMA1203.

Figure 7. Thermal effect produced by injection of a 3-gmol pulse of 0 2 o n t o reduced Rh/AI203.

These pulse experiments reveal a higher heat of adsorption of 02 compared to CO. Instead, the sticking coefficient of CO on rhodium appears higher than that of 02, as can be deduced from the volume of catalyst saturated by one pulse of each reagent under the same conditions. This result is in qualitative agreement with data reported for the competitive adsorption of the same species on Pt [8]. The dynamics of adsorption of alkanes is totally different. Both methane and propane were detected at reactor outlet already after the first pulse, revealing a very low sticking probability compared to CO and 02. Methane produced no thermal effect on the catalyst, whereas propane caused a weak temperature rise (about 0.2 ~ Activity measurements were carried out on 3% RMA1203 under isothermal conditions; the rate dependence on the fuel partial pressure was determined by gradually raising the concentration of one reactant, while keeping other parameters constant. In the CO oxidation, the effect of partial pressure of CO is described by a typical "volcano" curve, as reported in the literature for Pt group metals [9-11 ]. At low partial pressure, CO exhibits a strong positive effect on the reaction rate; for C0/02 ratio higher than about 0.3, a self-inhibiting effect takes place (Figure 8). The oxygen dependence is 1.2, revealing relatively weak adsorption in the presence of CO. In methane oxidation, a positive effect of methane partial pressure was measured in a wide range of feed composition; the reaction order is 0.7 in CH4 and 0.1 in 02. In propane oxidation, the rise of fuel partial pressure promotes the reaction only at low reactant/oxygen ratio; at higher values, surface saturation seems to take place (Figure 9). The oxygen dependence is 0.8, suggesting stronger competition by propane for surface sites, compared to methane.

435

60 450 e-

er" O

E

~" 50 350

r: "~ 40

, m

E o

E E

250

E

k_

Od

0 L)

150

50

r

30

20

10 0

0.5

1

1.5

2

2.5

Pco (kPa) Figure 8. CO dependence of CO2 formation rate on 3% Rh/A1203 . T=90 ~ Po2 =2 kPa.

0

0.5

1

1.5

2

2.5

3

PC3H8 (kPa) Figure 9. Propane dependence of CO2 formation rate on 3% Rh/A1203 . T=200 ~ P 02 =10 kPa.

4. DISCUSSION During the oxidation of CO, CH 4 and C3H8, the ignited state is characterized by a reaction front stabilized in a thin portion of the bed near the reactor inlet. This condition, corresponding to a diffusion-controlled reaction, is predicted by the known models of exothermic catalytic reactions [4]. The chemical factors determining this dynamics are the heat of reaction and the activation energy. For all of the reactants considered in this study, a similar behaviour in the ignited state is observed. Marked differences between the reactants are exhibited in the transition from the low activity state to the ignited state. The dynamic phenomena occurring in the catalytic oxidation of CO can be explained by assuming that the main factors governing ignition are the inhibiting effect of the reactant and the strong dependence of the desorption rate on temperature. In the pre-ignition state, the surface is predominantly covered by CO, as suggested by the high sticking coefficient of the latter and by literature data [12]; the rate of reaction is therefore controlled by the adsorption of oxygen. Initial CO conversion takes place along the catalyst bed rather than at reactor inlet, where surface CO coverage is presumably slightly higher. Heat is transferred towards the reactor outlet, where the concentration of CO is lower due to partial conversion, and the desorption rate from the surface is higher [13]. As a consequence, the surface CO coverage continues to drop in this region, eventually reaching the level corresponding to the sharp transition to the high activity state. Ignition near the outlet is followed by an upstream front motion, according to the dynamics described in numerous studies [1, 14, 15]. The downstream front motion accompanying extinction of the CO-O2 reaction, is controlled by the same factors determining ignition.

436 In the oxidation of C H 4 and C3H 8, no inhibition by the reactant is operative, and initial conversion is higher near the reactor inlet, where favourable conditions for ignition are stabilized. This relationship between dynamic behaviour and adsortion-desorption properties, finds support in the literature. Site competition between hydrocarbons and oxygen has been recognized as a key point in the mechanism of catalytic oxidation [3]; in the same study, a substantially different behaviour in catalytic combustion was reported for alkanes and ethylene, and this was related to the strong adsorption properties of the latter. The mechanism of self-sustained oscillations is not completely understood at present, despite the remarkable number of studies reported in the literature [ 11, 16]. It has been shown that in the CO-O2 reaction over Pt catalysts, oscillations occur in the transition region between the high and low activity state, corresponding to predominantly CO- or O- covered surface [ 11 ]. The existence of such a region of sharp transition appears a basic requisite for oscillating behaviour to take place. This statement applies also to the Rh/A1203 catalyst of this study, for which a similar transition region in the CO oxidation has been revealed by kinetic experiments. The weaker oscillations observed with propane and the absence of any significant dynamics in methane oxidation can similarly be related to the adsorption properties of the alkanes. This conclusion is supported by the fact that oscillations have been rarely reported for hydrocarbon oxidation, and in the case of methane, their occurrence has been excluded for a Pt/A1203 catalyst [ 17].

5. CONCLUSIONS High-resolution infrared imaging of the working catalyst provided information on the dynamics of combustion reactions with unprecedented detail. The capability of exhibiting competitive adsorption in the presence of 02, together with self-inhibiting effect, represents a key factor in the development of dynamic phenomena. This assumption is able to explain the different behaviour of CO during catalytic oxidation, with respect to light hydrocarbons. The same chemical factors that determine the mobility of the reaction front, seem also to play a role in the mechanism of oscillations. These results can help define general criteria for predicting unstable behaviour in processes carried out in the ignited state.

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

J.E. Gatica, J. Puszynski and V. Hlavacek, AIChE J., 33 (5) (1987) 819. M. Sheintuch and S. Shvartsman, AIChE J., 42 (4) (1996) 1041. G. Veser and L. D. Schmidt, AIChE J., 42 (4) (1996) 1077. H.J. Viljoen and R. C. Everson, AIChE J., 42 (4) (1996) 1088. P. Silveston, R. R. Hudgins and A. Renken, Catal. Today, 25 (1995) 91. S. Marengo, G. Raimondini and P. Comotti, in "New Frontiers in Catalysis", k. Guzci et al. (eds.), Elsevier, Amsterdam, Part C, (1992) 2573. 7. J. Puszynski and V. Hlavacek, Chem. Eng. Sci., 39 (4) (1984) 681. 8. Y. Nishiyama and H. Wise, J. Catal., 32 (1974) 50.

437 9. S.H. Oh and C. C. Eickel, J. Catal., 112 (1988) 543. 10. G. Ertl, Adv. Catal.. 37 (1990) 213. 11. M. M. Slin'ko and N.I. Jaeger, "Oscillating Heterogeneous Catalytic Systems", Studies Surf. Sci. Catal., Vol. 86, Elsevier, Amsterdam (1994). 12. F. Qin. and E. E. Wolf, Ind. Eng. Chem. Res., 34 (1995) 2923. 13. S. H. Oh, G. B. Fisher, J. E. Carpenter and D. W. Goodman, J. Catal., 100 (1986) 360. 14. G. Padberg and E. Wicke, Chem. Eng. Sci., 22 (1967) 1035. 15. D. A. Frank-Kamenetskii, "Diffusion and Heat Transfer in Chemical Kinetics", 2nd ed., Plenum, New York (1969). 16. F. Schtith, B. E. Henry and L. D. Schmidt, Adv. Catal., 39 (1993) 51. 17. T. R. Baldwin and R. Burch, Appl. Catal., 66 (1990) 337.

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91997 Elsevier Science B.V. All rights reserved. Dynamics of Surfaces and Reaction Kinetics in Heterogeneous Catalysis G.F. Froment and K.C. Waugh, editors

DYNAMIC OPERATION

OF TRICKLE

439

BED REACTORS

H.W. Piepers and A.A.H. Drinkenburg Department of Chemical Engineering, Eindhoven University of Technology, P.O. Box 513, 5600 MB Eindhoven, The Netherlands

1. ABSTRACT Conventional trickle bed reactors work under steady state conditions, whereby components in the.liquid, mostly via gas-to-liquid mass transfer are converted by the solid catalyst. Such a process is highly non-linear and thus it is questionable whether steady operation will provide an optimal conversion and selectivity in particular. Operation in the dynamic mode will provide an extra parameter to optimise the production. Trickle beds happen to show just naturally a dynamic flow regime in which gas/liquid discontinuities occur: the pulsing flow regime. This study concerns experimental investigations to expand the natural region of the pulsing flow mode in the flow map by varying the gas or liquid feed in time. It was found that it is possible to expand the pulsing flow regime to lower average gas and liquid flow rates. By pulsing the gas feed it is possible to predetermine the pulse frequency. However, the pulse frequency is always lower than the natural transition pulse frequency. The work fits in a program directed to establish new process technology, in general working with non-steady-state operations in a continuous manner.

2. INTRODUCTION Developments in the field of chemical engineering were very innovative in the third quarter of this century. Thereafter many improvements were made in existing technology, e.g. by better modelling and control. Nevertheless, in the seventies and eighties complete new instruments came to our disposal in the forth of new computer technology, without leading very much to new process technology. This is the more remarkable since in reality processes have a very high non-linear character, e.g. heterogeneously catalysed processes. Partly, since linear partial processes are lumped together non-linearly and partly since the partial processes are non-linear. In view of this non-linearity it is to be expected that steady-state processing, in which an output parameter like conversion or the selectivity of a reactor is connected to one steady set point of an input parameter, will never lead to an optimum result. This means that dynamic operation of reactors can show promising results. Haure et al. (1989) and Lange et al. (1994) convincingly provided examples of this principle. Through modem model-supported feed-forward control such action can be taken.

440 The need for internal recycling of energy and material in actual processes asks for continuous processing as well. Hence the combination of the two was in unsteady-state continuous processing. Although the concept is sometimes seen as revolutionary, we must realise that the dynamic operation in itself is only distinguishing on a time scale. On a molecular level all processes show unsteady-state behaviour, one time step higher we promote mass and heat transfer by unsteady turbulence. In this work we deliberately perform experiments in a trickle bed reactor in the dynamic mode. Trickle bed reactors are normally operated in the so-called gas-continuous regime, where in fact both gas and liquid percolate cocurrently and in continuous phases through the bed. Beyond certain gas and liquid loads the bed shows an unsteady behaviour, pulses of liquid are seen to develop and pass the bed with velocities of about 0.5 to 1 m/s (Rao and Drinkenburg,1983; Tsochatzidis and Karabelas, 1995a). These pulses are not distinctive masses of liquid, but are, in fact, waves of liquid-rich hold-up that pass the bed as can easily be seen from liquid tracer experiments (Blok et al., 1983). The waves are initiated as liquid flow instabilities and then grow to pistonlike slugs, although the waves are permeable to gas, penetrating the pulse at the rear of the wave in the form of bubbles and leaving the pulse at the front end. Of course then the linear gas flow must be higher than the pulse wave velocity (Blok and Drinkenburg, 1982; Tsochatzidis and Karabelas, 1994.). The pulses and the penetrating gas bubbles cause a number of beneficial features in the chemical reactor viz: 9 enlargement of mass transport from gas to liquid and liquid to solid (Seirafi and Smith, 1980; Rao and Drinkenburg, 1985; Tsochatzidis and Karabelas, 1995b ). 9 capture of the stagnant liquid hold up and redistribution after the pulse has passed. This leads to less axial mixing and better plug flow behaviour. 9 complete wetting of the particles at very short intervals, which is the more important when runaway exothermic reactions may be expected between gas and particles that have fallen dry (safety). 9 high levels of radial heat transfer enabling the design of internal heat transfer equipment. Moreover, once the pulses arrive in approximately the time constant of the overall reaction rate resonance in the cycle transfer - adsorption - reaction - desorption occurs, as was calculated by Wu et al. (1995), be it as yet with a rough model. Such resonance can lead to substantially better selectivities and conversions. In this paper experiments are described that expand the region of the pulsing flow mode from its naturally occurring window in the flow map to lower average gas and liquid flow rates. Hanika et al. (1990) showed that such is possible by varying the liquid flow rate in time.

3. EXPERIMENTAL A flow scheme of the equipment is shown in Fig. 1. Experiments were carried out in a cylindrical column made of Plexiglas. The inner diameter of the column was 0.11 m and the height was 1.0 m. Along the column height, pressure taps were mounted in order to measure local pressures. The trickle bed was packed with fairly uniform Raschig rings of glass with characteristic dimensions di,du,1= 5x6x10 mm.The packing material is supported, at the bottom of the bed, by a rigid nickel screen. The column was carefully packed to ensure a uniform porosity of the bed. The porosity of the bed was e = 0.65 and the specific surface area of bed was S = 930 m1.

441

>

Air was the gas phase and water was the liquid phase. To get an even liquid distribution over the packing the liquid phase was fed into the > top of the column through a distributor containing 11 tubes. The gas phase was introduced cocurrently with water at the top of the trickle bed between the liquid distributor head and the packing. Both the central gas and liquid feed were divided up in two feed lines. Air and liquid flow rates in each feed line were measured by calibrated rotameters. In one feed line of the liquid and in one feed line of the gas a magnetic valve was installed. The magnetic valve was controlled by an electronic timer. air With this timer the opening and closing time of air the valve could be adjusted. With this set-up it was possible to cycle the liquid and gas flow water ~/ rate between two finite values. water exit The flow rates used ranged from ca 0.01 to 0.03 m3/m2.s for the liquid and from 0.08 to 0.7 Figure 1. Flow sheet of the equipment. m3/m2.s for the gas. The packing was completely prewetted by operating the bed at high gas and liquid rates, afterwhich the desired gas and liquid rate were adjusted. The total liquid hold-up was measured indirectly by a tracer technique which measured the mean residence time of the liquid with which the hold-up could be calculated. The transition from gas continuous to pulsing flow was determined as follows. At various gas flow rates the liquid feed was increased in steps until regular pulsing was observed. The pulse frequency was determined either by a pressure transducer connected to a fast fourier analyser or visually with help of video recordings. In forced pulsing of the liquid feed the base liquid flow rate was adjusted to a fixed value of ca 0.01 m3/m2.s during a time tb, the so-called base time. The pulsed liquid feed flow rate was fixed during a time tp, the pulse time, to a value of twice the base value: ca 0.02 m3/ma.s. E Starting from a gas and base liquid flow rate combination whereby no pulsing occurred, the base time tb and tp were adjusted. Then the gas flow rate was increased in steps until pulsing was observed. For the base time tb 5, 15 and 30 sec. was used. The liquid feed pulse time tp was varied between 0.2 to 15 sec. Not only the liquid feed could be pulsed, but also the gas feed. Starting from a liquid flow rate in the trickling flow area a base gas flow Figure 2. Example of periodic changes in rate, the base time tb and the pulse time tp were liquid/gas feed. adjusted.

.i..I ' ],[[

magnetic valve

~~

,.~

baseflowpLJIsfloew

v

0 L_

< tb~p~

timeIs/

442 Subsequently the pulse flow rate was increased in small steps to determine the transition to pulse flow and then this procedure was repeated for another liquid flow rate. At the same time also the frequency of the resulting pulses were determined. The frequency of the feed (liquid or gas) is given by:

f:-/~tb+tp) 4. RESULTS AND DISCUSSION 4.1. Steady state experiments The results of the hold-up measurements are shown in Fig. 3. In this figure also the correlation for the hold-up proposed by Rao (Rao and Drinkenburg, 1985) is given. Rao proposed the following correlation for the total liquid hold-up:

/3-0.037 a,,~(Re, \~egJl

~

(1)

where a, is the specific surface area of the packing and Re~, Re g are the Reynolds numbers based on the superficial liquid or gas velocity. For our system Eq.(1) can be rewritten as:

/ 3 - 0.59

r,

(2)

In Fig. 3 the correlation of Rao is shown by the solid line. The agreement with our results is altogether reasonable. All the data lie within the + 15% error limits of the correlation.

0.5

v

i

i

0.4 0.3

0 ..C

t

i

i

i 9

O,J~r T I

0.2 0.1

|

0

0.04

0.08

0.12

0.16

0.2

ul/Ug (-) Figure 3. Comparison of liquid hold-up with the correlation of Rao (1985); Eq.2.

443 The results of the transition from trickling to pulsing flow are shown in Fig. 4. The results show that with increasing gas velocity the superficial liquid transition velocity decreases. According to Blok (Blok et al., 1983) the transition from trickling to pulsing flow occurs, for each system and type of packing, at a constant Froude number: 2

Fr, -

v~.,

ga,

where

vl. t

_ cons tan t

(3)

is the linear liquid velocity at the transition point. This means that at the transition

the linear velocity yr., - u~'t//~sj3 must be constant. The linear liquid velocity is also plotted in Fig. 4. As can be seen from this figure the transition to pulsing flow indeed occurs at nearly constant linear liquid velocity (so constant Froude number). However there is still no reasonable explanation for the fact that the transition occurs at a constant Froude number. Sicardi (1979) suggested that pulsing could occur only above a certain liquid hold-up. Preliminary experiments and modelling points also to a minimum liquid hold-up that is needed for the transition from trickling to pulsing flow. In Fig. 5 the results of the pulse frequency measurements are shown. From their measurements Rao (Rao and Drinkenburg, 1983) found that the pulse frequency could be correlated with the difference between the linear liquid velocity v~and the linear liquid velocity v~,t at the transition from trickling to pulsing flow:

where c is a constant depending on the packing characteristics.

0.11 0.09 j

e• k

,

'

!

,

I

9

~" 0.07 E

-,

i

j 0.03

pulsing !

0.025

!

-.

!.'.,.

o

. . ,

0.05 0.03 -

0.015 if trickling

0.01

i

0.01 0

0.1

0.2

0.3

0.4

0.5

, 0.6

0.005 0.7

Ug (m/s)

Figure 4. Steady-state transition from trickling to pulsing flow as a function of superficial gas velocity. (e superficial velocity;. II linear velocity)

444 _

!

i

,-, 2 . 5 -

s

co)

v

>,

o t-

Q)

2 -

,~

9

0

w

O

-

1.5

1,..

1 2I

~

0.5-

0

1

2 (vl- vl,t).lO0

3 (m/s)

4

5

Figure 5. Natural pulse frequency as a function of the difference in linear liquid velocity. (. . . . . correlation of Rao, Eq. 4)

From Fig. 5 it can be concluded that our measurements are also reasonably fitted by Rao' s correlation, although the deviations are sometimes large. This is due to the fact that the pulse frequency is never very stable. Deviations of the average value up to 1O-15% can occur. 4.2 I n d u c e d pulsing experiments

The results of the experiments in which the liquid feed is cycled are given in Figures 6,7 and 8. The typical behaviour of the influence of the pulse liquid feed on the transition from trickling to pulsing flow is shown in Fig. 6. In this figure the results are depicted for a base flow time of tb = 5 s and a pulse time ranging between 0.3 and 15 s. Experiments with base flow times of 15 and 30 s resulted in analogue pictures. From the results it is obvious that by pulsing the liquid feed the transition from trickling to pulse flow shifts to lower liquid velocities. However, pulses can only be generated by the pulsing liquid feed if the pulse flow rate has a value which lies in the natural pulsing flow area. From Fig. 7 it can be seen that first the transition gas velocity u~t decreases with increasing pulse time tp and then becomes practically constant at a gas velocity that equals the minimum steady-state transition gas velocity at the higher liquid flow. This is an indication that at higher pulse times the hold-up in the bed is adjusted to a value corresponding with the higher pulse flow. For pulse times smaller than ca 0.2 s no pulses can be generated in the bed because the hold-up is not enough increased. This effect can also be seen in Fig. 8, in which the resulting bed pulse frequency is plotted as a function of the adjusted liquid feed frequency for three different base flow times. The bed pulse frequency is always smaller than the steady-state transition frequency (0.5-0.7 Hz). With increasing feed frequency, smaller tps, the pulse frequency becomes almost constant, which means that at these small tps one feed cycle generates one pulse. For smaller feed frequencies, high bs, one cycle can generate more than one pulse.

445 0.8

base time tb=5s

0.7

9 0.6

tp=O.3s

0.5-

E

i

0.4=

!

~b

iX_

0.30.2-

tp=

0.10 0.005

0.015

0.01

0.02 u,

0.025

0.03

(m/s)

Figure 6. Influence of liquid feed pulsing on the transition from trickling to pulsing flow. ( 9 steady state; A cycle feed, base flow 0.0095 m/s, pulse flow 0.019 m/s)

0.7-

i

0.6-

t I

0.5r

E .,-.=

I

0.4,~

0.3-

0.2

I ,

i

I

0.1 .

0

~

" r~

'

[

2

4

6

I

'

!

i

t

'

:

-

,

"~

~

T

!

!

8

10

12

14

16

t n (s)

Figure 7. Transition gas velocity as a function of the liquid pulse time tp for various base times. ( 9 tb = 5 S; 9 tb = 15 S; A tb = 30 S; base flow 0.0095 m/s, pulse flow 0.019 m/s)

446 0.35 ~" 0.3

-

>, 0.25

0 ctl)

=

!

0.2

(1) I,i

"-'0.15 (1)

0.1 "13

-Q 0.05

-

v

0 0

0.04

0.08

0.12

0.16

0.2

liquid feed frequency (l/s)

Figure 8. Relation between induced liquid feed frequency and pulse frequency. ( 9tb = 5 S; 9 tb = 15 S; A tb = 30 S; base flow 0.0095 m/s, pulse flow 0.019 m/s)

By pulsing the gas feed it appears that a pulse time tp of minimal 1 s was necessary. Pulse times smaller than 1 s did in fact induce pulses in the bed but these were not stable and disappeared before the pulses reached the end of the bed. If the pulse time was 1 s the resulting bed pulse travelled along the entire bed. Experiments were done with a pulse time tp = 1 S and base times tb of 2, 3 and 4 s, resulting in a frequency of the gas feed of 1/3, 1/4 and 1/5 Hz, respectively. The superficial base gas flow was varied between 0.06 m/s and 0.3 rn/s. The value of the pulse flow whereby a stable pulse in the bed was generated was practically equal to the steady-state transition gas flow, corresponding with the adjusted liquid velocity. The resulting bed pulse frequency was equal to the gas feed frequency, but was always smaller than the steady-state transition frequency (0.5 -0.7 I/z). As an example is in Fig. 9, for two base gas velocities and a gas feed frequency of 1/3 Hz, the averaged transition gas velocity plotted as a function of the liquid velocity. The results for a feed frequency of 1/4 and 1/5 Hz were similar. From the results it appears that the transition gas velocity can be shii~ed to almost every value of the gas velocity in the area below the steady-state transition velocity and between the broken lines (see Fig. 9) by choosing the feed frequency and the base gas flow.

5. CONCLUSIONS AND SUGGESTIONS

Periodic operation of a trickle bed reactor has many advantages. Over a wide range of gas and liquid flow rates pulsing conditions can be induced namely by variation of the liquid flow at constant gas flow or the other way around. A more intricate pattern of simultaneous variation of gas and liquid flow rates has so far not been investigated. The important feature of the periodically operated bed is the establishment of pulses with a predetermined frequency and

447

0.80 . 7

tb=2S

-

to-1 s

0.6 t/) 0.5 ,

E 0.4

I

i

0.3 i

0.2 0.1

I r

_

0.005

~

0.01

0.015

i

T

i

0.02

0.025

'

0.03

u~ (m/s) Figure 9. Lowering of the gas transition velocity by gas feed pulsing. ( steady state transition; base gas flow: 9 0.08 m/s, A 0.27 m/s )

therefore time constant. Most promising is the pulsing of the gas flow, pinpointing the pulse frequency. It will be worthwhile to compare the efficiency of the pulsed trickle bed reactor with that of a trickle bed reactor in the gas continuous flow mode and with a monolith reactor. Based on its good heat and mass transfer characteristics the suggestion is that for reaction with a relatively small adiabatic temperature rise and a low reaction rate the trickle bed reactor in the gas continuous mode can be used. For high reaction rates and low thermal effects the monolith reactor is promising. In between the two a trickle bed reactor with shell-catalyst establishes itself For reactions at a relatively low reaction rate which are, however, highly exothermic a pulsed flow trickle bed shows good features and induced pulsed flow is then a wanted design criterion. For both high reaction rates and high adiabatic temperature rises at the moment slurry reactors are favourite.

NOTATION dp

g Frt

ff f~ S

particle diameter (m) acceleration due to gravity (m/s 2) Froude number at transition (-) feed frequency (s-1) pulse frequency (s 1) specific packing surface (m l )

448 tp

tb Ul Ul,t Ug Ug,t Vl Vl,t

pulse time (s) base time (s) superficial liquid velocity (m/s) superficial liquid velocity at transition (m/s) superficial gas velocity (m/s) superficial gas velocity at transition (m/s) real liquid velocity (m/s) real liquid velocity at transition (m/s) liquid hold-up, defined as a fraction of void volume (-) porosity of the bed (-)

REFERENCES

Blok, J.R. and Drinkenburg, A.A.H., 1982, Chem. Eng. J., 25, 89 Blok, J.R., Varkevisser, J. andDrinkenburg, A.A.H., 1983, Chem. Eng. Sci., 38, 687. Hanika,J., Lange, R. and Turek, F., 1990, Chem. Eng. Process., 28, 23. Haure, P.M., Hudgins, R.R. and Silveston, P.L., 1989,AIChE J., 35, 1437. Lange, R., Hanika, J., Stradiotto, D., Hudgins, R.R. and Silveston, P.L., 1994, Chem. Eng. Sci., 49, 5615. Rao, V.G. and Drinkenburg, A.A.H., 1983, Can. J. Chem. Eng., 61, 158. Rao, V.G. and Drinkenburg, A.A.H., 1985,AIChE J., 31,1010. Rao, V.G. and Drinkenburg, A.A.H., 1985,AIChEJ., 31,1059. Seirafi, H.A. and Smith, J.M., 1980, AIChE J., 26, 711. Sicardi, S., Gerhard, H. and Hofrnarm, H., 1979, Chem. Eng. J., 18, 17 Tsochatzidis, N.A. and Karabelas, A.J., 1994, lnd. Eng. Chem. Res., 33, 1299. Tsochatzidis, N.A. and Karabelas, A.J., 1995a, AIChE J., 41, 2371. Tsochatzidis, N.A. and Karabelas, A.J., 1995b, Chem. Eng. Commun., 137, 85. Wu, R., McCready, M.J. and Varma, A., 1995, Chem. Eng. Sci., 50, 3333.

91997 Elsevier Science B. V. All rights reserved. Dynamics of Surfaces and Reaction Kinetics in Heterogeneous Catalysis G.F. Froment and K.C. Waugh, editors

449

Simulation of a catalytic converter of automotive exhaust gas u n d e r dynamic conditions A.J.L. Nievergeld, E.R.v. Selow, J.H.B.J. Hoebink" and G.B. Marin Emdhoven University of Technology, Laboratorium voor Chemische Technologie P.O. Box 513, 5600 MB Eindhoven, The Netherlands

ABSTRACT An adiabatic, one-dimensional model of a catalytic converter of automotive exhaust gas was used to simulate the behaviour during cyclic feeding and warmingup. Both the oxidation of CO, C_oH2 and CzH4, and the reduction of NO are considered. Accumulation of mass in the bulk gas phase, in the pores of the washcoat and on the catalyst surface is accounted for, as is accumulation of energy in both the gas and the solid phase. Exhaust gas components are converted in a fixed sequence and the light-off temperature of individual components is rather irrelevant for the behaviour of a real exhaust gas because of m u t u a l interactions in a mixture. Forced concentration cycling below the light-off temperature of CO, Coil2 and C_~H4can reduce the emissions of the individual components, but the optimal feed temperature is not the same for each compoment.

1. I N T R O D U C T I O N Car a n d catalyst manufacturers put much effort in reducing the emission of pollutants in order to meet the more stringent standards. Both engine construction improvements and new catalyst developments are necessary to achieve ultra low emission vehicles. Optimisation of the existing catalysts is facilitated by knowledge of the underlying reaction kinetics, while insights obtained from reactor modelling can be used to optimise the performance as a function of the operating conditions. Due to the fuel control system the converter operates dynamically. The delay time in the feedback loop of the control system leads to an oscillating exhaust composition around the stoichiometric point at a typical frequency of 1 Hz which is of the same order as the turnover frequency of the reactions. Therefore in reactor modelling the dynamic behaviour of these reactions has to be considered meaning that detailed kinetic models based on elementary steps have to be used. Furthermore accumulation both in the gas phase and on the catalyst surface should be taken into account. to whom correspondence should be addressed

450 2. K I N E T I C M O D E L The kinetic model, see Table 1, for the reactions between CO, 02, NO, C2H4 and Table 1. " E l e m e n t a r y " steps and reaction pathways, indicated by the stoichiometric numbers o i. 01

02

03

04

05

CO.

2

2

2

0

0

20*

1

0

0

3

5

0

2

2

0

0

C oH 4.*

0

0

0

1

0

c,,g,,****_ _

0

0

0

0

2

+ O.

0

1

2

0

0

COo + 2 ,

2

2

2

2

0

No.+O,+,

0

1

0

0

0

0

0

1

0

0

0

0

0

1

0

0

0

0

0

2

k a.CO

C O + 9 ,~ kd, cO ka.o 2

02 + 2*

-" ka.NO

N O + 9 ,~ N O . kd,NO

k~,.c.,.H4

Cell 4 + 2 *

,~ kd.c H ka,c 2H2

C_oH2 + 4 .

kd, C H

kd~

NO.

+ 9 -. N .

CO,

+ O.

NO.

+N,

kco.z -. kN 2.1

--

kN2.2

2N.

-

N2+2. kH20.1

C2H4.. +40.

- 2C0. +2H20+4.

C~/2.*** +50. o 1" 2 C 0

kH.~O,2 - 2C02+H,,_0+9.

+ 02-. 2C02

02, o 3 : 2 C 0

+ 2NO-.

2 C 0 , , + No

o4: CJ-/t + 302 - 2COo_ + 2 H 2 0 o~: 2Coil 2 + 502 - 4 C 0 2 + 2 H 2 0 * s t a n d s for catalytic site

451 C_oH2 was constructed from literature data. Basically the steps involving CO, NO and 02 were taken from Oh et al. [1]. Ethene oxidation was based on the work of Sant et al. [2] for low oxygen concentrations. As data for ethyne oxidation were not available at all, the same model but for high oxygen concentration was applied u n d e r the adaptation of the sticking probability and the surface reaction rate constant kH2o., to account for a higher C~H.o reactivity [3]. The global reactions shown in Table 1 correspond with the minimal number of reactions which has to be taken into account. From a mathematical point of view they can be regarded as a base for the space including all possible reactions. The reaction between e.g. NO and Coil 4 or C.~-I2 can also take place and is a linear combination of the global reactions given in Table 5.1. The adsorption of oxygen is assumed as irreversible and dissociative with a rate proportional to the fraction of vacant sites [1]. In contrast to Sant et al. [2] adsorption of both ethene and ethyne is also assumed to be first order in the vacant sites. The rate of the surface reactions between adsorbed ethene and adsorbed oxygen, and between adsorbed ethyne and adsorbed oxygen, is considered as proportional to the product of the involved surface coverages. The adsorption rate coefficients are obtained from the kinetic gas theory, while Arrhenius-type expressions are used for the rate coefficients of desorption,dissociation and the surface reactions. The kinetic parameter values used in this study are shown in Table 2. Table 2. Kinetic p a r a m e t e r values and references used in the modelling. Sco ~

0.5

[1]

Ed.C2H,

57.3

[2]

EN2,1

87.8

[1]

A~,co

1.6"10 '4

[1]

S~ ,_H , ..,

1.0

[4]

AN2,2

3-10 '0

[ 1]

E~co

112

[1]

Ad.CzH~

6.5"10 '3

[2]

EN2,2

120

[5]

S O2 O

0.01

[1]

Ed.~2H 2

115.0

[2]

AH2O,'

6.0"1013 [2]

SNO~

0.5

[1]

A~.ss

3" 10 '~

[1]

EH2O.1

62.7

[2]

Ad,NO

5"10 '3

[1]

E~.~

79.4

[1]

AH2O.2

10 '2

[1]'

Ea,NO

108.7

[1]

Ac%

1" 10 '2

[1]

EH2O,2

54.3

[2]

SOoH,

0.35

[2]

Eco._'

60

[ 11

Ad,C2H,

1.0"1013

[2]

A%.1

2"109

[1]

~) The frequency factor for ethyne oxidation is increased from 8-107 s" (reported for ethene [2]) to 10 '-~ g' (reported for CO [1]), reflecting the much higher oxidation rate of ethyne compared to ethene [3]. 3. R E A C T O R M O D E L The adiabatic model of the monolithic converter is based on the model reported by Lie et al. [6]. For both the gas phase and the solid phase, only axial concentration

452 and temperature gradients are considered. Heat and mass transport resistances are located at the interface between bulk gas phase and the washcoat. Axial mass diffusion in both the gas phase and the washcoat as well as axial heat conduction in the gas phase is neglected. Pore diffusion limitation is also neglected [7]. Gas p h a s e concentrations are expressed as C/pf in order to correct for density changes due to heat generation. The fractional coverage of the surface species is defined by the ratio of the sites occupied by a species to the total number of active sites m e a n i n g t h a t e.g. for the adsorption of one molecule ethene four active sites are needed. The continuity equation for reactant i (i = CO, 02, NO, C~H4, or C,Hz) in the bulk gas phase is given by: c~

= - (bsup c~

- Pfkf.iav

(1)

a n d in the pores of the washcoat by:

~wPf

4edw 0 d---b- at

= pfkf.ia~ pf ]

. _ Cs.i|~ _ ac, tR i pr Pf )

(2)

The continuity equation for surface species j (j = CO*, O*, NO*, N*, Cell4**, or CeHz****) is given by:

LtaOJ 0t : Rj

(3)

a n d for the vacant sites by: o

- 1 - ~

oj

(4)

J The energy equations for the gas phase and the solid phase are given by: dTf sup ~Tf ~:pfCpf at - -~)m Cpf aX - aav(Tf - T )

(5)

dTs s (l-e) a2Ts 4 ~ (-5rH)kr k (1-e)psCps a t - ~ ax-~ + a a v ( T f - T s ) § ac, tk--1

(6)

'The h e a t production is calculated from the reaction enthalpies of the global reactions given in Table 1. This is allowed, since temperature oscillations due to the the time dependent heat production as a result of the changing surface coverages are negligible [8]. B o u n d a r y conditions for the bulk gas concentrations are given by the inlet concentrations:

453 ~----irl

Cf.i(0,t) : Cf.i[1 + B s i n ( 2 = f t + ~)]

t ~ 0

(7)

w h e r e ~ = 0 for CO, C2H2 a n d C_oH4, a n d ~ - 7~ for 02 a n d NO. For t h e gas phase t e m p e r a t u r e the b o u n d a r y condition at the inlet is defined by: Tf(0,t) = T ~

t ~ 0

(8)

while for the solid phase zero fluxes are a s s u m e d at t h e reactor inlet a n d outlet: aT (0,t) dx

aT (L,t)

= 0,

dx

= 0

t >0

(9)

The initial conditions are imposed by the a s s u m p t i o n t h a t the r e a c t o r operates in steady state before the feed oscillation starts. Hence, for the concentrations in the bulk gas p h a s e and in the solid p h a s e they are given by: SS

SS

Cf.i(x,0) = Cr.i (x), C~.i(x,0 ) = C~.i (x)

(I0)

0 _< x _< L

a n d for the surface coverage of species j by: $S

Oj(x,O) = Oj (x)

(II)

0 < x _< L

T h e initial conditions for the b u l k gas p h a s e t e m p e r a t u r e and the solid phase t e m p e r a t u r e are: Tf(x,0) : T;S(x), T (x,0) = T~S(x)

0 _< x _< L

(12)

Table 3. Reactor p a r a m e t e r s at T - 500 K [5]. e

0.60

av

2.4"103

a

141

~ms"p

5.83

db

1.0-10 .3

(-ArH) 1

283" 103

pf

0.680

d~

2.5-10 ~

373-10 z

ps

2.5-103

L

0.15

(-ArH) 2 (-hrI-I)3

Sh, Nu

3.66

A~.

6.0.10 .3

(-ArH) t

1254" 103

kf.co, kf.%

0.18

c~f

1.081-103

Lt

2 . 7 1 0 .5

kf.NO

0.21

cps

1.015-103

ac.~t

1.25-104

kf.c2H4, kf, c2s2

0.14

),f

3.86-10 .2

B

0.15

ew

0.40

~,s

1.675

1322.103

454 4. R E S U L T S AND D I S C U S S I O N Both steady state and cyclic feeding simulations have been performed. The latter allow the calculation of the oscillatory effects on the fi'me average performance. Time average inlet concentrations and the corresponding equivalence ratio JC are given in Table 4. Table 4 Equivalence ratio X and the corresponding time average inlet concentrations. )~ CO O~ NO C,Hy X CO 02 NO C,I-Iy vol% vol% ppm ppm vol% vol% ppm ppm 0.94 0.95 0.96 0.97 0.98 0.99 ") C x H y

-

1.41 1.21 1.03 0.87 0.73 0.61 C2H2 or

0.32 0.34 0.36 0.39 0.42 0.44 Coil4

637 678 727 782 846 919

615 593 572 552 534 516

1.00 1.01 1.02 1.03 1.04

0.50 0.41 0.33 0.27 0.22

0.48 0.51 0.54 0.58 0.62

1000 1088 1182 1280 1379

500 485 472 461 451

4.1. S t e a d y s t a t e r e s u l t s Figure 1 shows the hght-off curves of all four pollutants in the mixture. Ethyne is oxidized first, followed by CO above 480 K. When most CO has reacted the conversion of ethene and NO starts simultaneously. From this figure it is clear t h a t the reactions take place in the sequence: C,J-I~ < CO < NO+C,J~,~. This is also shown in F i g u r e 2 in which the bulk gas and solid phase concentration profiles at 500 K are depicted. In the first part of the reactor ethyne is converted. When most ethyne has reacted, the CO oxidation starts 100 followed by both the NO reduction and the ' /."NO C_J-I4 oxidation. The calculated sequence is 80 /,'/ in agreement with the experimental work t t): .~ 6o / ] i of Mabilon et al. [3]. Pattas et al. [9] l! l: 40 .'" showed that 20% of the hydrocarbons is i ~ CoHo co/ ..:7 converted before CO oxidation starts, o 20 " I / :I:CoH 4 which is in line with the alkyne oxidation i --.-,,-, -, .... ,.;,~ i.,/ 480 490 ~o ~;o 470 ~oo as the alkyne content of exhaust gas is typical 20% of the total organics. As Ti, [KI depicted in Figure 2 significant mass transfer limitations are absent. The strict Figure 1. Outlet concentrations reaction sequence can be understood from versus the feed temperature. the corresponding surface coverage profiles shown in Figure 3. In the first part of the Conditions: k = 1, Tables 3 and 4.

=

t

455 reactor mainly ethyne and CO compete for the active sites and inhibit the adsorption of oxygen and ethene. There is also some NO adsorption and dissociation. After all C2H2 and CO has reacted, vacant sites become available for C2H4 adsorption and also more NO adsorption and dissociation can occur. When almost all pollutants are converted, the surface becomes oxygen covered. The axial t e m p e r a t u r e profiles are not shown, but simulation results showed t h a t the solid temperature is slightly, about 10 K at most, higher t h a n the bulk gas temperature due to the heat of reaction. The maximum temperature rise across the reactor equals the adiabatic temperature rise amounting to 94 K.

0.50~

> o.4o ."r

~'~"

1.00 r

, , , ~ CO

1 0.80 [

030 / c~n~ "9 2 ~, 0.20[ \ c,,H, ~,:,~,, 9 ~

008.00

. . . .

'~,

"7,_

1

" " 0.60 I

':. / /' / -......

C.J-I,

/

.....

'................. "--"" ~ ~-~-~-4-'-::- ~- -1.. . . . 0.08.00 0.03 0.06 0.09 0.12

0.15

x [m]

Figure 2. Concentration versus axial coordinate at a feed temperature of 500 K. Thin line: bulk gas phase; thick: line pores of the washcoat. Conditions: k = 1, Tables 3 and 4.

0 <--t ,. ,//

:....~-

o2oi o

o.09 0.)2

axial coordinate

........................ :..

CO .... ,

0.40 ~- " ] .... ..C..2H2

., " :.-k-:.-.:..

0.03 --o~oG

N :-

axial

coordinate

"....... 0.15

x [m]

Figure 3. Surface coverage versus axial coordinate at a feed temperature of 500 K. Conditions" k = 1, Tables 3 and 4.

4.2. L i g h t - o f f b e h a v i o u r ""

F i g u r e 4 shows the conversion of the pollutants as a function of time as a result of a step increase of the feed t e m p e r a t u r e from 300 K to 550 K. In this figure a similar conversion sequence of the pollutants is visible as in Figure 1. When the ethyne conversion starting after about 10 s is considerable, the CO oxidation begins. After 50 s conversion of both NO and Cell4 is completed. The shoulder in the NO conversion is caused by strongly mcreased NO adsorption when the CO oxidation is complete.

I~176 80 .~ oo

C..H~y

>~

/ ....

Go [

:/

40 =

/.::.

T

20[I 0

0

,'" C

?/

,"'" .~........."" - / ~ ~ " "':~" " 10 20

time

/

/

/ C2H4 /' 30

40

[s]

Figure 4. Conversions versus time after a step increase of the exhaust gas temperature from 300 K to 550 K. Conditions: k = 1, Tables 3 and 4.

456 4.3 Cyclic f e e d i n g r e s u l t s

/\

1OF

0.40

C0 / / ~ . , ,

'-~ 0.30

CO

> 0

=c

C2H '

o2o1o

-~

'

05

:!.!.

o.

....

!:i....i

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

.

O

0.00 8.0

\8.4

100 C2H~

8.8

9.2

9.6

10.0

t i m e [s]

Figure 5. Outlet concentrations versus time during 1 Hz cycling. Conditions: ~ = 1, Tt~ = 485 K, Tables 3 and 4.

~

0

../

~.~.

/

.\C2H~ .\

... / /

~.

-10

.c -20

~ N -30 c 470

,

475

I

i

480

feed temperature

L

485

,

~

490

[K]

Figure 6. Conversion improvements due to 1 Hz cycling versus feed temperature. Conditions: ~ = 1, Tables 3 and 4.

Figure 5 shows the gas phase outlet concentrations as a function of time when the inlet concentrations are cycled at a typical frequency of 1 Hz. The curves shown in this figure are part of a stable oscillation; the transient part of the response to reach this stable oscillation is not shown. The amplitude of all concentrations except for 02 increases towards the outlet of the reactor as reported in the literature for the reaction between CO and O2 [6], and for the reaction between CO, O2 and NO [5]. When the concentration of the reducing components increases, the higher degrees of CO and C2H2 coverage lead to stronger inhibition. The resulting lower degrees of coverage of other components lead to a decrease of the reaction rates. In contrast, when the concentration of the reducing components decreases the lower CO and C2H2 coverages result in higher surface coverages of the other species and higher reaction rates. The out-of-phase oscillation of both oxygen and NO leads to a further acceleration of this effect. As shown in Figure 5 most ethyne is converted. During the net oxidizing part of a cycle period complete conversion of both CO and C2H2 occurs. The noble metal surface becomes then almost totally covered with molecular NO and with O and N adatoms, and the 02 and the NO gas phase concentrations start to increase. When the concentration of the reducing reactants increases due to the inlet oscillation, the reaction with adsorbed oxygen begins and the O2 gas phase concentration decreases. Several experimental studies show that cyclic feeding leads to a higher time average performance [10-12]. In general this is caused by the non-linear behaviour of the occurring reactions [13]. To show the effects of the oscillation on the reactor performance, calculations have been performed for different temperatures, frequencies and feed compositions. Oscillations are only beneficial at a relatively low temperature and the optimal cycling frequency depends strongly on the

457 temperature. The effects as a function of the feed temperature at a frequency of 1 Hz are summarized in Figure 6. From this figure it is clear that the beneficial effects of the individual components do not coincide at any temperature. For a feed temperature of 485 K Figure 7 shows the effect of the air/fuel ratio on the performance of the reactor during both steady state and cyclic feeding. At this temperature steady state operation is preferable and the so-called X-window should be as narrow as possible to obtain high conversions of all pollutants. Cyclic feeding is only beneficial at non-stoichiometric feeds. At values of X above unity, i.e. net oxidizing conditions, NO conversion is reduced tremendously and cyclic feeding becomes beneficial. Ethene oxidation is also slightly reduced since ethene = 0.60? /,.+ /. \a 1 adsorption has to compete with oxygen for o.,o # the vacant sites. Under net reducing > I N :, conditions ethene is not converted since it o _ o/ _+ is a much weaker competitor for oxidizing '8.;4-o.;G"o.;s' l.oo' .o21.o4 species than ethyne and CO. The time l a m b d a 1-] average ethene conversion is then higher than the steady state conversion. Under deeper net reducing conditions NO Figure 7. Effects of air to fuel ratio on reduction is decreased since its adsorption is inhibited by adsorbed CO and adsorbed both the steady state and time N. Cyclic feeding is then beneficial for CO average conversions. Full lines: steady state; dotted line: cyclic and ethyne. The simulation results shown feeding, o = C2H4; + = CO; 0 = ~ I-Z_; in Figure 7 are in qualitative agreement = NO. Conditions: T~" - 485 K, f - 1 with experimental results from the literature [ 10]. Hz, Tables 3 and 4.

!o-"

!

CONCLUSIONS A dynamic model based on first principles is used to predict the behaviour of an exhaust gas converter as a function of the process conditions. The simulation results are in qualitative agreement with the experimental results from the literature and can be used m the optimisation of the reactor performance. First ethyne is converted, then carbon monoxide, and finally nitrogen oxides and ethene simultaneously. Cycling feeding can either improve or deteriorate the performance of the monolithic reactor and the effects are strongly influenced by the process conditions. The conversion improvement of a pollutant increases with increasing temperature until periodically complete conversion occurs. At a higher temperature the time average conversion becomes lower than the steady state conversion. Above the light-off temperature of all components steady state reactor operation at the stoichiometric point leads to the highest performance.

458

NOTATION

Roman letters a~t a~ A As B % C db d~ E, f A~H kf.i L I~ Nu r R So Sh t T x

catalytic surface a r e a geometric surface area frequency factor cross-sectional area fractional amplitude specific heat concentration channel diameter washcoat thickness activation energy forcing frequency reaction enthalpy mass transfer coefficient reactor length concentration active sites Nusselt number reaction rate gas constant sticking probability Sherwood number time temperature axial coordinate

G r e e k letters mNM2 mR3 n%2 mR3 s~ m~J k g ~K tool mr3 mR mR kJ mol ~ Hz J mol ~ mf3 mi- s I mR mol mNM' 9

a heat transfer coefficient c void fraction of monolith eW washcoat porosity 0 fractional surface coverage ~, thermal conductivity equivalence ratio

W m 2 K~ m~ mR3 m~3 mw3 mol mO1NM1 W m ~ K~

v stoichometric coefficient p gas density o stoichiometric number phase slfift Omsuperficial mass flow

kg mf3

(A/F)/(,VF)~,c~

rad kg mR2 s 1

.9

.9

-1

mol m-NMS kJ mol ~ K~ s K mR

Subscripts

Superscripts

a d diss f s stoich sup

in ss

adsorption desorption dissociation bulk gas phase solid phase stoichiometric superficial

inlet steady state

REFERENCES 1. 2. 3. 4. 5. 6. 7. 8. 9. 10. 11. 12. 13.

S.H. Oh, G.B. Fisher, J.E. Carpenter and D.W. Goodman. J. Catal. 100 (1986) 360 R. Sant and E.E. Wolf. AIChE J. 35 (1989) 267 G. Mabilon, D. Durand, and Ph. Courty. Catalysis and Automotive Pollution Control III 96 (Studies in Surface Science and Catalysis), Amsterdam, (1995) 775 R.L. Palmer. US NTIS AD Report A040579 (1977) A.J.L. Nievergeld, J.H.B.J Hoebink and G.B. Marin, Studies in Surface Science and Catalysis 96 (1995) 909 A.B.K. Lie, J.H.B.J. Hoebink and G.B. Marin. Chem. Eng. J. 53 (1993) 47 R.E. Hayes and S.T. Kolaczkowski. Chem. Eng. Sci 49 (1994) 3587 E.R.v. Selow. Graduate Report, Institute for Continuing Education Eindhoven University of Technology, (1996) 63 K.N. Pattas, A.M. Stamatelos, P.K. Pistikopoulos, G.C. Koltsakis, P.A. Konstandinidis, E. Volpi, and E. Leveroni. SAE Paper 940934 J.C. Schlatter, R.M. Sinkevitch and P.J. Mitchell. Ind. Eng. Chem. Prod. Res. Dev. 22 (1983) 51 K.C. Taylor and R.M. Sinkevitch. Ind. Eng. Chem. Prod. Res. Dev. 22 (1983) 45 H. Muraki. H. Shinjoh, H. Sobukawa, K. Yokota and Y. Fujitani. Ind. Eng. Chem. Prod. Res. Dev. 24 (1985) 43 Y.S. Matros. Catalytic processes under unsteady-state conditions, Studies in surface science and catalysis, 43, Elsevier, Amsterdam, 1989

91997 Elsevier Science B. V. All rights reserved. Dynamics of Surfaces and Reaction Kinetics in Heterogeneous Catalysis G.F. Froment and K.C. Waugh, editors

459

Effect of Variables on the Periodic Operation of a Trickle Bed Reactor

Leonardo Gabarain ~b, Jorge Cechini b and Patricia Haure b• a CIC. Comision de Investigaciones Cientificas de la Provincia de Buenos Aires, Argentina. b INTEMA. Fac Ingenieria. UNMdP. Av. J.B. Justo 4302. Mar del Plata 7600. Argentina.

1. INTRODUCTION Previous works on the periodic operation of a Trickle Bed Reactor (TBR) have demonstrated that for the hydrogenation of alpha methyl styrene (AMS) to cumene on Pd/A1203 catalyst, reaction rates are increased up to 400% with respect to the steady state results (CasteUari and Haure (1995)). In this mode of operation, the liquid phase is switched on and off while the gas passes continuously through the reactor. Silveston et al. (1995) pointed out that due to the complexity of the unsteady operation and the probability of higher costs than the conventional procedure, cycling is recommended if the rate enhancement is at least higher than 100%. The performance of a TBR under cycling is extremely complex, specially for the gas-limited, liquid volatile, exothermic reaction situation. It depends of the switching between wet and dry operation (Gabarain et al. 1996). In the non-wet cycles, wetting of the packing is incomplete. The reaction rate can be greater or smaller than the rate observed over completely wetted packing. This depends on whether limiting reactant is present only in the liquid phase or in both gas and liquid phases. If the reaction is gas limited, rates will be higher because the the gas reactant can access the catalyst pores from the externally dry area. When the reaction is liquid limited and the liquid reactant is non-volatile, a decrease in the wetting efficiency will cause a decrease in the reaction rate. But if the liquid reactant is volatile and significant heat effects are also present, then, a gas phase reaction can occur on the dry catalyst resulting in higher rates and temperatures(A1-Dahhan and Dudukovic, 1995). The associated increase in the local reaction rate and the temperature rise may be a feature of concern because it can lead to the formation of uncontrolled hot-spots and the subsequent sintering of the catalyst, runaway conditions and undesirable side reactions, as pointed out by Hanika et al. (1975).

to whom correspondence shouldbe addressed

460 Proper selection of the cyling variables is strongly recommended to exploit the vaporization phenomena avoiding related problems. The occurrence under steady-state conditions of vaporization during reaction in multiphase reactors at the single pellet level is well documented (Hanika et a1.(1976), Watson & Harold (1993,1994)). Watson and Harold (1993) studied the Pd-catalyzed hydrogenations of AMS (to cumene) and of cyclohexene (to cyclohexane) in a single pellet catalytic reactor. The key difference between the two reactions is the volatility of the less volatile reactant. The experimental data reveal a interplay between the exothermic reaction, the endothermic vaporization and internal and external transport processes. For the case of drying with hydrogenation, they observe departure from traditional drying theory, specially in the initial period due to the influence of the latent heat of vaporization, the heat conduction and the heat generated by the exothermic reaction. In general, hydrogenation accelerates the drying process because reaction on dewetted sites speeds up the pore emptying process. The rate of drying depends on the volatility of the liquid and, if sufficient time elapses that a fraction of the catalyst becomes exposed to the gas, a much more rapid gas-phase catalytic reaction occurs with an accompanying temperature excursion. Gabarain et al. (1996) proposed a phenomenological model that explains the events associated with periodic interrumption of the liquid phase to a TBR Experimental and model results compare reasonably well. Cycling deliberately creates hot spots during the "dry" cycles due to the onset of gas-phase reaction and results in higher conversions compared to the convencional operation. The overall reaction rate is an average between the diffusioncontrolled liquid phase reaction and the more rapid gas-phase catalytic route. The reactor operates at higher average temperatures compared to the steady state. However, it is the phase change responsible for the large magnitude of the rate increase. The extent of pore filling depends on several variables, such as catalytic activity, support permeability, heat of reaction, volatilities of the liquid components, overall reaction energies, gas flowrate and composition, etc (Watson and Harold, 1994). Once the gas phase reaction commences, the AMS is rapidly consumed. The optimum conditions for cycling are a strong function of the "depletion time" of the liquid reactant and can be achieved by a proper selection of splits and periods providing that all other conditions remain constant. The objective of this work is to experimentally examine the effect of several variables such as gas flowrate, gas and liquid composition and liquid feed temperature in the periodic operation of a TBR The above mentioned parameters influence the degree of external wetting, liquid holdup and transport and reaction mechanisms. The couplings between the extent of pore emptying, partial external wetting, transport and reaction can be assessed by means of the model previously developed by Gabarain et al (1996).

2. EXPERIMENTAL DETAILS The hydrogenation of AMS to cumene has been extensively studied by others ( Morita and Smith, 1978; Herzkowitz et a1.,1979; Cini and Harold, 1991). Reaction rate is rather fast

461 under mild conditions and cumene is the only measurable product Properties of the system are shown in Table 1.

Table 1 System Properties Liquid Gas Catalyst

-AH~ (40~

AMS 99% purity (Aldrich Chemical Co.) hydrogen 99.9 %, nitrogen 99% Pd/AI203 (commercial) 109 KJ/mol

Properties of the catalyst and operating and analysis procedures are given in detail by Castellari and Haure, (1995). Bed characteristics are given in Table 2.

Table 2 Bed Characteristics Diameter Type of inert bed Pre packing depth Catalytic bed depth Post packing depht Catalyst weight Bed void fraction

2.54 cm ID glass beads 12 cm 1.5 cm 50 cm 3.48 gr 0.48

A sheathed thermocouple was inserted axially in the middle of the catalytic bed. Local bed temperatures were recorded continuously by a data acquisition device. A three way solenoid valve (Jefferson model 365) activated by timers was used to generate liquid flowrate variations. Hydrogen was presaturated with AMS. Experiments were performed randomly. Operating conditions used in most of the experiments are given in Table 3

Table 3. Operatin~ Conditions Split 0.5 Period 20 min Gas Flowrate 400, 900 and 1500 cc/min Mean liquid Flowrate 2.27 cc/min Reaction rates were evaluated following the procedure described by Castellari and Haure (1995) and are expressed in mol of cumene/s gr. of catalyst. 3. RESULTS The behaviour of the catalytic bed during periodic operation can be explained by means of a simple phenomenological model (Gabarain et al, 1996). Consider a catalytic pellet, as shown in Figure 1 During the wet cycle (Zone A) the pellet is completely wet, the dissolved gas

462 species is the limiting reactant and the overall rate is external mass transfer controlled, as in conventional TBR operation. Once the liquid flow is halted, the bed partially drains and the catalytic reaction proceeds between the flowing hydrogen and the liquid holdup. Heat generated is not easily dissipated by the flowing gas and evaporation of the liquid phase, followed by a much rapid gas-solid catalyzed reaction may occur in Zone B 1. The overall reaction rate has then two contributions : one from the wetted areas or liquid-solid catalyzed reaction and the other from the non-wetted surfaces. The holdup diminishes with time, due to evaporation. Once the external holdup is depleted, the reaction takes place between the evaporated internal holdup and the gas reactant (Zone B 1). Here we assume that reaction takes place via gas-solid catalysis. When the internal holdup is depleted, or depletion time (td), bed temperature decreases. (Zone C) Experiments were carried out to observe the performance of the reactor while varying liquid feed temperature, gas flowrate and gas and liquid composition. These variables influence the drying rate of the liquid film, approximated by the transport of AMS through a boundary layer: rd =

ka (pv (TA -pb) ZONE BI

ZONE A e,'
R~- Rd+Rw ZONE B2

C)

ZONE C

t

Rd

Figure 1 : Phenomenological Model after Gabarain et. al (1996) 3.1. Liquid feed temperature

Although liquid feed temperature have not an important influence at steady state conditions where mass transfer resistances are of significant importance (Castellari & Haure,1995), it is an important variable while operating a TBR in circumstances where vaporization and gas-solid

463 catalized reaction are possible. Figure 2 shows temperature profiles obtained at three different liquid feed temperatures and gas flowrate 900cc/min. Table 4 compares experimental and theoretical results. As liquid feed temperature was increased, maximum temperatures reached were higher and a displacement between the maximum is found. Profiles have approximately the same shape, indicating the occurrence of similar phenomena. Temperature raise enhances the reaction rate, as expected. The liquid's vapour pressure is higher, increasing the drying rate according to equation (1). It also allows the gas phase catalytic reaction to start before and accelerates the pore emptying process, modifying the depletion time. Comparison between experimental and predicted results is presented in Table 4. Reasonable agreement is found Temperature '

I

100

(o C ) '

I

'

I

'

I

'

I

'

Bath T e m p e r a t u r e

---41oc I .... 33oc /

80

/ /

"--"---"

_ ,"--,

/

!

\

60 I

._

I

,f

40 ~ .

20 6800

,

, 7000

,

, 7200

,

i

7400

Time

,

l 7600

,

, 7800

,

8000

(s)

Figure 2. Temperature profiles for different liquid feed temperatures.

Table 4. Gas Flowrate 1500 cc/min; Split 0.5 and Period 20 min. Cumene Concentration < 15%. T bath (~ r * 10 6 (model) r * 10 6 (experimental) tD (rain) 31 2.1 1.48 8.7 40 3.41 4.51 5.3 53 4.45 5.95 4.05

3.2. Gas Flowrate

Temperature profiles for three different flowrates of pure hydrogen (400, 900 and 1500 cc/min ) are plotted in Fig. 3. The maximum temperature rise and the depletion time varied. The gas flowrate influences the mass transfer coefficient in (1), and the heat transfer coefficient. Mass transfer and heat removal are enhanced and they have an opposite effect.

464 The more rapid initial drying obtained with the higher hydrogen flowrate resulted in a more rapid temperature rise and vaporization of liquid and subsequent pore emptying. Similar results were reported by Watson and Harold (1993). However, the heat transfer coefficient is also enhanced, allowing the bed to remove heat more efficiently. Maximum temperatures differ in less than 20 C.

Temperature (~C) 80

9

I

I

'

'

I

'

I

'

I

"

Gas Flowrate 400 cc/min

70

9

900 cc/min ........

9

1500 cc/min

~,,

/

9

I jr

:

\

".

I

60 50

-f

o.

40

.

S

30 ,

6800

I

7000

,

I

7200

~

I

7400

,

I

,

7600

I

!

7800

8000

Time (s) Figure 3. Temperature Profiles for different Gas Flowrates. Liquid Feed Temperature 41 ~ Table 5.Liquid Gas Flowrate 400 900 1500

feed Temperature = 41 ~ and Cumene Concentration,< 15%. r * 10 6 (experimental) r* 10 6 (model) tD (min) 2.48 3.1 8.125 3.25 3.1 6.04 4.517 3.1 5.3

Model fails to predict experimental results probably due to differences in the calculated heat and mass transfer coefficients.

3.3. Liquid Concentration The effect of the concentration of cumene was also studied. Temperature profiles in Fig 4 show that a pure liquid reactant feed reaches higher temperatures and reaction rates. The dilution effect of the liquid reactant should be further investigated.

465 Temperature 80

'

I

(o C) '

I

'

Cum____eneCon__centration 70

--

--

_

< 15% r= 4.52 10-6 > 15% r = 2.53 10-6

I

'

I

'

I

,

/ ~ / !

,-. \ / \-,\

60 50 40 30 ,

750

I

1000

,

I

,,

1250

I

1500

,

1750

2000

T i m e (s)

Figure 4. Temperature Profiles for different feed compositions. Table 6. Liquid Feed Temperature 40~ Gas Flourate 1500 cc/min. Period 20, Split 0.5. Cumene Conc r * 10 6 (model) r * 10 6 (experimental) tD (min) < 15 % 3.45 4.52 5 > 15 % 2.89 2.53 5.58

3.4. G a s C o m p o s i t i o n

Hydrogen concentration (,Vh2)in the gas feed was varied. Mixtures of hydrogen (high purity) and nitrogen (high purity) fed continuosly through the reactor. Temperature profiles in Figure 5 reveal that the maximum temperature reached is the same, but there is a displacement in the depletion time. Maximum at Yh2 25% is found after 25 min (not shown). If Yh2 is below a critical value, temperature remains constant. This can be explained as follows : once the gas phase reaction commences, it just depends on gas phase AMS concentration, given by the holdup, provided that there is enough hydrogen to sustain the zero order dependence. Although H2 is not pure, its concemration is enough to be considered in excess in the gas phase step. As hydrogen concentration becomes lower, mass transfer of hydrogen becomes limiting. Experimental rates were also monitored. In general, the higher the fraction of hydrogen present in the gas feed, the higher the rate. Below a critical value of hydrogen no reaction is observed. When Yh2 is higher than 25% maximum temperature reached are identical. Once the gas phase reaction commences, it just depends on gas phase AMS, provided that there is enough hydrogen to sustain the zero dependence. Although H2 is not pure, its concentration

466 As hydrogen concentration

is enough to be considered in excess in the gas phase step. becomes lower, it becomes limiting. Temperature ( ~C) 0

,

1

'

,

I

,

I

Gas Composition

70

'

I

/ ~

100% H2 50%H2

60

I

/

'

.--,,

/ \

~

,'"

/

50 40 30 20 3600

,

i

3800

,

i

4000

,

I

4200

,

!

4400

,

1

4600

,

4800

Time (s)

Figure 5.Temperature profiles. Split = 0.3. Period = 20 min. Gas Flowrate =1500cc/min. Table 7. Period 20, Gas Composition 100 % 50 % 25 % 17 %

Split 0.3. Gas Flowrate 1500 cc/min. Liquid Feed Temperature 40~ r *10 6 (model) r * 10 6 (experimental) to (min) 3.62 3.94 7.17 2.28 3 13.5 1.35 1.8 25 no reaction observed no max.

4. CONCLUSIONS Periodic operation of TBR deliberately modifies the wetting condition of the catalytic bed, allowing to create dry areas. If properly controlled, it can benefit reactor performance, specially in gas-limited - liquid volatile reactions with important heat effects. In the dry cycles, reaction proceeds between the liquid holdup (internal and external) and the flowing gas. The liquid holdup diminishes with time, until is fully depleted. Gas flowrate and composition and liquid feed composition and temperature have an important effect on reactor performance during cycling.

467 REFERENCES

A1-Dahhan M.H., and M.Dudukovic. "Catalyst Wetting Efficiency in Trickle Bed Reactors at High Pressure". Chem. Engn.Sci. 50, 15,2377 (1995). Cini P., and M. P. Harold. ''Experimental Study of The Tubular Multiphase Reactor". AIChE J.., 37, 7 (1991) Castellari A.T. and P. M. Haure. "Experimental Study of the Periodic Operation of a Trickle Bed Reactor". AIChE J. 41,6 (1995) Gabarain L. , A. T. Castellari, J. Cechini, A. Tobolski and P. Haure. " Analysis of Rate Enhancement in a Periodically Operated Trickle Bed Reactor". In Press. AIChE J, Nov. 1996 Hanika J., K. Sporka, V. Ruzicka and J. Krausova, "Qualitative Observations of Heat and Mass Transfer Effects on the Behaviour of a Trickle Bed Reactor", Chem Eng Comm., 2, 19 (1975). Hanika J. ,K. Sporka, V. Ruzicka and J.H. Rstka " Measurement of Axial Temperature Profiles in an Adiabatic Trickle Bed Reactor". Chem Engn J. 12, 193 (1976). Herzkowitz, M., R.G. Carbonell and J.M. Smith. "Effectiveness Factors and Mass Transfer in Trickle Bed Reactors". AIChE J., 25,272 (1979). Morita, S., and J.M. Smith "Mass Transfer Limitations in a Trickle Bed Reactor" I&E Fund. 1,113, (1978). Silveston P.L., R.R. Hudgins and A. Renken. "Periodic Operation of Catalytic Reactors introduction and overview". Catalysis Today 25, 91-112 (1995) Watson P.C. and M.P. Harold. "Dynamics Effects of Vaporization with Exothermic Reaction in a Porous Catalytic Pellet". AIChE J. 39,6 (1993). Watson P.C. and M.P. Harold. "Rate Enhancement and Multiplicity in a Partially Wetted and Filled Pellet: Experimental Study". AIChE J. 40, 1 (1994).

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91997 Elsevier Science B. V. All rights reserved. Dynamics of Surfaces and Reaction Kinetics in Heterogeneous Catalysis G.F. Froment and K.C. Waugh, editors

469

Oxidative coupling of toluene under periodic conditions on Pb/Li/MgO: A selective pathway to 1,2-diphenylethane S. Dubuis, M. Lorenzi, R. Doepper and A. Renken* Institute of Chemical Engineering, Swiss Federal Institute of Technology, 1015 Lausanne, Switzerland

Toluene oxidation over Pb/Li/MgO is investigated in a tubular fixed bed reactor at low conversions. A periodic process alternating toluene oxidation in absence of oxygen in the feed and regeneration of the catalyst by oxygen presents important changes in product distribution in comparison with steady-state oxidation. Periodic conditions permit to increase the selectivity in favour of the coupling reaction which lead to 1,2-diphenylethane and to avoid formation of oxygenated products.

1. INTRODUCTION The catalytic oxidation of toluene over metal oxides to benzaldehyde and benzoic acid are well-known industrial processes and minor amounts of coupling products are detected among the products. It is shown that the oxidative coupling of toluene is favoured under anaerobic conditions and that metal oxides of the groups III to V of the periodic table catalyse methylmethyl coupling [1 ]. Pb/Li/MgO is chosen for this study because it is known as a selective catalyst for the oxidative methylation of toluene with methane [2-4] and as good benzyl radical producer [ 1]. Activity and selectivity in heterogeneously catalysed reactions may be enhanced during the transient period which follows a forced concentration perturbation at the inlet of the reactor. Under transient conditions, the concentrations of reagents and intermediates as well as the catalyst activity differ from steady-state conditions. These changes can induce the disappearance of a reacting route by the lack of a reagent or permit the formation of new intermediate species. To perform a process under transient conditions, it is possible to change the concentrations of the inlet flow periodically or to carry out the different reaction steps at different positions in the reactor. As the oxidative coupling of toluene is favoured under anaerobic conditions, a periodic process with a reacting time interval which permits the toluene coupling and a regenerating interval during which the catalyst is oxidized, will increase the 1,2diphenylethane yield and selectivity. A practical example of this kind of process is the selective oxidation of butane to maleic anhydride [5].

Corresponding author

470 2. EXPERIMENTAL

2.1 Catalyst Preparation The Pb/Li/MgO catalyst is prepared by wet impregnation method with an atomic ratio of 1/3/16 [2]. MgO Powder (Fluka) is added to an aqueous solution of LiNO3 (Fluka) and Pb(NO3)2 (Fluka). The slurry formed is dried by evaporation and further dried at 120 ~ in air. It is then calcinated at 750 ~ for 12 h. The calcined catalyst is crushed and grains of 355-500 /an are collected by sieving.

2.2 Experimental set-up and method A tubular stainless steel reactor (I.D. 104 mm) heated by an electrical oven at atmospheric pressure is used for the oxidation of toluene [Fig. 1]. The toluene is dosed with an HPLC pump (LKB2150) to an evaporator at 320 ~ and then mixed to the 02 and N2 flows which are controlled with mass flow controllers (Bronkhorst High-Tech B. V.). Nitrogen is used as diluent. The catalyst fixed-bed preceded by quartz beads is maintained between quartz wool. The temperature of the fixed-bed is measured with a K-type thermocouple (Philips AG). The outlet gases are cooled in three consecutive condensers. The liquid products are collected and analysed by gas chromatography with a flame ionisation detector for quantification (PerkinElmer Autosystem gas chromatograph, capillary column Supelco SPB-1, 30 m x 0.53 mm I.D. x 0.50/an film thickness) and with an electron ionisation detector for identification (HewlettPackard, G1800A, GCD System, capillary column HP-5, 30 m x 0.25 mm I.D. x 0.25/an film thickness). The experiments are carried out at a conversion less than 5 per cent.

vent

k

02 N2

Ill h0

Toluene

Q I Figure 1. Schematic experimental set-up.

,I 1

vent

471 For the calculation of yields (Y0, the amount of condensed products collected during a time interval (At) is referred to the toluene feed [Eq. 1]. Selectivities (Si) are calculated for the condensed products [Eq. 2]. For the periodic experiments, an average yield is calculated over the period to allow comparison with steady-state conditions. Finally, the modified residence time is defined as the catalyst mass divided by the toluene molar flow at the inlet of the reactor [Eq. 3 ].

Yi ~

ni fi tol.0 "At

Si -

ni Enj J

(1)

(2)

mcat fitol,0

(3)

3. RESULTS AND DISCUSSION

3.1 Steady-state experiments The oxidation of toluene over Pb/Li/MgO leads to different products which selectivities are greatly dependant on temperature and residence time. The steady-state experiments are carried out with a toluene/oxygen ratio of 3.8 which corresponds approximately to the stochiometric ratio for the oxidative coupling of toluene to 1,2-diphenylethane. 0

Figure 2. Reaction scheme for the oxidative coupling of toluene. Figure 3 presents the products at x' = 18 kg~ts/mol as a function of temperature. For temperatures below 450 ~ the only detected product is benzaldehyde. With an increase of temperature, other products appear especially 1,2-diphenylethane and benzene, but also Estilbene, styrene and fluoren-9-one. Oxidation of the methyl group is favoured at low temperatures to produce benzaldehyde and oxidative coupling needs higher temperatures. Zhu and Andersson [6] explained this by a lower activation energy of direct oxidation reactions compared to the activation energy of coupling reactions. Moreover, the coupling reactions need benzylic radical species in the gas phase and benzyl desorption does not occur at low temperatures [6,7].

472 1.0

u

0

0.90.80.70.60.5CO

0.4-

Z~ o +

benzene benzaldehyde

9 1,2-diphenylethane E-stilbene

x a

fluoren-9-one styrene

0.30.20.1 A

0.0

r"

35O

'

I

4OO

'

'

45O

+

I

I

5OO

55O

I

60O

650

T [~ Figure 3. Selectivities as a function of temperature. ~' = 18 kg~ts/mol, Toluene:oxygen ratio = 3.8. 1.0 0.90.8-

A

benzene

o

benzaldehyde 9 1,2-diphenylethane

0.70.6"7' 0.5-

O9

0.40.30.2

A

-

o

0.1A

0.0

10

I

20

I

A

o I

30

40

x'

[kg~ts/mol ]

o I

50

Figure 4. Selectivities as a function o f modified residence time. T = 527 ~ Toluene:oxygen ratio = 3.8, CtoL0= 7.5 mol/m 3.

60

473 Benzene begins to be observed at 450 ~ and its selectivity increases with temperature. It could be produced by direct oxidation of toluene or also by oxidation of benzaldehyde [6,8]. Trace amounts of other compounds, such as E-stilbene and styrene are only observed at high temperature. The influence of the residence time on the selectivities is studied at 527 ~ Figure 4 shows that benzaldehyde is favoured at short residence times and 1,2-diphenylethane at higher residence times. The highest selectivity for the 1,2-diphenylethane, obtained in the experimental domain, is 0.61.

3.2 Transient experiment The toluene oxidation over Pb/Li/MgO is carried out in absence of oxygen. The catalyst is first oxidized under an oxygen in nitrogen flow. At a given time (t - 0), the oxygen flow is switched off and replaced by toluene in nitrogen. The instantaneous yields of detected products are reported in Figure 5. Only 1,2-diphenylethane, benzene and E-stilbene (not shown) are detected. E-stilbene is only present in traces. The rapid decrease of the reaction rate suggests that toluene is activated on oxidized sites at the surface of the catalyst which generates benzyl radicals or adsorbed benzyl groups. The disappearance of a detectable amount of benzaldehyde among the products in absence of oxygen indicates that gas phase oxygen is necessary to obtain this product. The lattice oxygen or oxygenated sites is not active for the formation of oxygenated products such as benzaldehyde and fluoren-9-one. The absence of oxygen in the feed does not permit to regenerate the active sites. As there is no oxygen in the feed the catalyst deactivates completely during a period of 4000 s. As the oxidative coupling of toluene can be performed on the catalyst in absence of oxygen, the use of periodic conditions is considered to enhance its activity and its selectivity in favour of 1,2-diphenylethane. 0.015

i

0.010

9 1,2-diphenylethane zX benzene

>0.005

0.000 0

I 1000

I 2000

1 30OO

t[s] Figure 5. Yields as a function of time in absence of oxygen. T = 527 ~ z' = 40 kgcatS/mol, CtoL0= 7.5 mol/m 3.

4000

474

3.3 Periodic experiments Periodic experiments which alternated reacting and regenerating time intervals are carried out at 537 ~ over 8 g of Pb/Li/MgO. The toluene flow is set to 310 .4 mol/s during the reacting intervals and the oxygen flow is set to 2.410 .4 mol/s during the regenerating intervals [Fig. 6]. The split ratio between reacting and regenerating interval is 3:1. A short purge time is inserted respectively between reacting and regenerating intervals and at the end of the period. The mean flows of toluene and oxygen all over a period and the total flow, kept constant using N2 as an inert, correspond to the flows of a steady-state experiment. The comparison of the periodic results is done with the corresponding steady-state experiment. 5x10 ~ total flow

4xI0 ~-

toluene

3x10 "40

02

E

=____, o

t::: 9 2x10"4N 2 ..........................................................................

I

-1

t

1xl 0 "4-

0

0

'

I

300

'

I

600

'

I

900

'

I

1200

'

I

1,500

'

1800

I

2100

'

I

2400

2700

t[s] Figure 6. Molar flows for a period of 2700 s and a split ratio of 31 between reacting and regenerating intervals. Benzene and 1,2-diphenylethane yields during a periodic experiment are plotted as a function of time in Figure 7. It is interesting to remark that benzaldehyde is no longer detectable, corresponding to a benzaldehyde selectivity below 0.01. The production rate decreases rapidly during a period in absence of oxygen. After three periods, the cycles are time invariant. The oxygen flow is sufficient to regenerate the catalyst to a constant activity, but the average yields are lower than the yields during the first period. Under periodic conditions, the 1,2-diphenylethane selectivity compared to steady-state conditions is increased from 0.6 to 0.9 by suppressing gas phase oxygen [Fig. 8]. The interaction between the catalyst and the toluene methyl group produces benzyl intermediates which react selectively to form 1,2-diphenylethane and benzene.

~>I gO0O = *~2a ' [ E o[.le,t l!ids '3o L~:S = i SOI.11.A.IIOO|gSptre spIo.[,~ olels-,~peols ptre o[.pouod jo uos.uedtuo D 8 zart~t.d [S] p o ! J e d OOLg

OOgl.

O'O--

006

C)9I~

alels-Kpem9 -

bO-

"7"

0000 LO0O

gO-

~7

gO-

800"0

~O-

~ooo -.'<

_co s o 1....i

"7"

90LOg0-

A%!A!IOalaS auezuaq

~l!A!1oeles eueqlal~uaqd!p-E' l

~7

9

1~00"0

pig!,( auazugq K'~_~ pig!,( gulaq~,al/~uaqd!p-E' 1.

900"0 I

60O'b-

9000

~ I gO0O = *~-'tu ' [ s = o p e a l ! i d s 'Do LZ;S = l ' s OOLZ; = p o u o d suop,!puoo o[.popod aopun atu[.1jo uol.lOUrL3 e se sppt.A L oaru~[.:I

[s]~ ... iv.. v-.!..? ? I 9 ~I

. .$ . .v..i. . .~ .~ I I

$

~ I ~__~...

Ix

~I

9

9

-v ' '~v

"I

I

',

OOZE I

:

,

x ~..~-. 9 v

0 0000 EO0O

9

"'I

-lX30O 9

9

9

i

-9000

t I

Ple!A e6eJeAe auazuaq

-<

9000 . ' -OLOO

i

ple!X auazuaq

~

-E 1,00

Pla!A e6gJaAR aueqlalXuaqd!p-E' L - -

Pla!A aueqlals

L --I--

-i;' 1,00 9 I,OO

~L~

476 The yield of 1,2-diphenylethane is also higher under periodic conditions compared to the steady-state experiment. It increases with shorter periods, due to a shorter reacting interval which leads to less deactivation of the catalyst. The catalyst is therefore globally more active at short periods.

0.004 ,-

1.0 9

9

-0.9

-0.8

0.003 -

>'-

-0.7 ~ ~

0.002-

~

1,2-diphenylethane yield benzene yield

:

i- o.6 .~.

i '-0.5 00-0.4

lb21~lziePnheenyl;;~ia;~ d selectivity

-0.3

0.001 -

"0.2 -0.1

0.000

1"1

i

31

i

6.81

,

0.0

Split [-] Figure 9. Yields and selectivities comparison as a function of the split ratio between reacting and regenerating intervals. T = 527 ~ Period = 1800 s, rn~t = 0.008 kg. The influence of split ratio is studied at a period of 1800 s. The mean flows are kept constant but the ratio of the reacting and regenerating intervals are varied from 1:1 to 6.8:1. As shown in Figure 9, no significant differences are observed in yield or selectivity compared to the split 3:1. The higher toluene concentration during the split 1:1 experiment induces a faster deactivation during a shorter reacting interval, therefore the selectivities and average yields do not differ from those of the split 3:1. The lower toluene concentration during the experiment at the ratio 6.8:1 leads to a slower deactivation during a longer reacting interval, therefore the selectivities and average yields correspond again to the values obtained at the ratio 3:1.

4. CONCLUSION The oxidation of toluene under periodic conditions permits to obtain 1,2-diphenylethane at a selectivity of 0.9 and to avoid the production of benzaldehyde and other oxygenated products. This reaction is an example which illustrates how a process under periodic conditions may change the selectivities and yields of an heterogeneously catalysed reactions compared to steady-state conditions.

477 5. NOMENCLATURE Roman letters

Ctol.O I.D. Si T Yi rr~t ni

toluene concentration at the inlet of the reactor internal diameter selectivity of product i temperature yield of product i catalyst charge in the reactor mole of product i condensed

[mol/m3]

[mm] [-] [oc] [-] [kg] [mol]

molar flow of reagent i at the inlet of the reactor

[mol/s]

time

[s]

modified residence time time interval

[kg~ts/mol]

At Indices cat tol 0

catalyst toluene at the inlet of the reactor

ni.o t Greek letter

[s]

REFERENCES

1. 2. 3. 4. 5. 6. 7. 8.

S.T. King, J. Catal., 131 (1991) 215-225. H. Kim, H.M. Suh and H. Paik, Appl. Catal. A, 87 (1992) 115-127 H.M. Suh, H. Kim and H. Paik, Appl. Catal. A, 96 (1993) L7-L 11 Y. Osada, N. Okino, S. Ogasawara, T. Fukushima, T. Shikada and T. Ikariya, Chem. Lett., (1990) 281-282 R.M. Contractor and A.W. Sleight, Catal. Today, 3 (1988) 175-184 J. Zhu and S.L.T. Andersson, J. Catal., 126 (1990) 92-100 J.P. Bartek, J.M. Hupp, J.F. Brazdil and R.K. Grasselli, Catal. Today, 3 (1988) 117-126 S.L.T. Andersson, J. Catal., 98 (1986) 138-149

This Page Intentionally Left Blank

91997 Elsevier Science B. V. All rights reserved. Dynamics of Surfaces and Reaction Kinetics in Heterogeneous Catalysis G.F. Froment and K.C. Waugh, editors

479

Reduction-Oxidation-Cycling in a Fixed Bed Reactor with Periodic Flow Reversal

H. Seiler and G. Emig Lehrstuhl fiir Technische Chemic I. (:niversitat Erlangen-,Viirnberg, D-91058 Erlangen, Germany

Simulations have been carried out for a modified steam-iron-process as a test reaction for processes where reduction-oxidation cycles will occur in two separated steps in a fixed bed reactor. The particularity of the used reaction svstem is the combination of a relatively fast exothermic gas-solid reaction in one half-cycle with a slow endothermic one in the other half cvcle. A simple pseudohomogeneous, one-dimensional dispersion model yields a complex dynamic behaviour due to the interaction of gas concentration. temperature and oxygen content inside the bed. The simulations show that a periodic change of the flow direction is necessary for utilizing the heat released by the exothermic reaction for the endothermic step and for maintaining a. higher conversion rate in the case of a separated equilibrium reaction. Wether the flow reversal should be after a half cvcle or a full cycle depends on the reaction system. For the evaluation of these operation alternatives the conversion rate, the time for reaching the periodic steady state and the maximum temperature during the whole process were used.

1

S c o p e of the W o r k

By transient o p e r a t i o n of chemical reactors the efl:iciency in the utilization of mass or energy may be enhanced. There are m a n y possibilities in realizing such d y n a m i c states, e. g. by periodic change of the inlet conditions [1] or the flow direction [2]. A useful technique is s e p a r a t i n g heterogeneous catalytic reactions into two gas-solid reactions, a reduction and an oxidation step [:3-9]. This may lead to higher yields of the desired product, because s i m u l t a n e o u s side reactions with active gas phase oxygen on the catalyst ~urface are avoided. A n o t h e r a d v a n t a g e is t h a t each gas-solid reaction may be o p e r a t e d under its respective o p t i m a l conditions. \Vhile fluidized beds are c o m m o n l y used for these reduction-oxidation cycles fixed bed reactors have some advantages like lower mechanical stress on the particles or the well-

480 defined residence time. Energy may be saved by using an adiabatic fixed bed reactor where the heat released during the exothermic reaction step is stored in the bed and utilized by the endothermic reaction directly without losses in heat exchangers. This is possible because the heat cal~acity per unit volume of the solid is much higher than that of the gas, so the velocity of the temperature front is much slower than the gas velocity. A wellknown way to trap energy is changing the flow direction before the hot spot leaves the reactor [2]. Operating an adiabatic fixed bed reactor with this periodic flow reversal, the inlet temperature may be decreased without, extinguishing the reaction, which is useful in the catalytic purification of polluted air [10,11]. In the simulations three basic alternatives of using an adiabatic fixed bed reactor for reduction-oxidation cycles (Fig. 1) were investigated. In the first alternative the flow Alternative 1" Red.

(

Alternative 3" ~

)..._.~

Red.._..~( [ ~ ~ ~ ~

) r

I>NN

)

)

Alternative 2:

Red._( )_ox

(DNN

Red.

_.. Ox.

Fig. 1. Basic alternatives for reduction-oxidation-cycling in a fixed bed reactor direction is not reversed which has the advantage that only few switching valves are needed for such a process. In the second alternative the flow is reversed after each half cycle, i. e. each reaction step. In the remaining alternative it is reversed after each full reduction-oxidation cycle. So it is obvious that different aspects from well-known processes like adsorption, periodic flow reversal and two step approach for heterogeneous reactions are combined here. The solid is now a reservoir for both heat and mass which leads to complicated dependencies between the state variables. In this case simulation studies prior to real experiments provide fast information on the behaviour of the system and to investigate the influence of several mechanisms like dynamics, kinetics etc.

2

2.1

Test R e a c t i o n and M o d e l

Reaction System

For the simulation studies a test reaction was chosen where the separation in a reduction and an oxidation step is necessary. The steam iron process was used in the twenties for hydrogen production from coal [12]. It includes a reduction of an iron oxide (mostly Fe304)

481

by lean or synthesis gas and a subsequent reoxidation of the reduced iron oxide (normally FeO) which yields hydrogen [1:3] (Fig. 2). The advantages of this reaction system for basic

H2, C O ~ .

Reduction

H2

H20, CO2

H20

Oxidation

Fig. 2. Scheme for the steam iron process studies are the small number of reactions which can occur and the possibility of having endothermic and exothermic steps - depending on the active metal used. Using hydrogen as only reducing agent reduces the complexity of the system and allows a closer look at the reversible reaction. For the modified steam iron process metal components are required where the equilibrium constant is nearly one. otherwise hydrogen would not be removed from the lean gas or steam would not be converted. Other relatively cheap metals like nickel may substitute iron in the steam iron process if an economic operation at lower temperatures (T < 800K) is possible. Producing hydrogen from industrial reducing waste gases or cleaning hydrogen for low temperature fuel cells are potential applications. For the simulations, nickel oxide was used as active metal component (modified steam iron process).

~.~

Reduction:

NiO +

H2 k+.,~Ni +

Oxidation:

Ni-t- H20

k_

H20

> NiO + H2

Mathematical Descrip*ion

The process is described by an one-dimensional, pseudohomogeneous, non-steady state dispersion model for an adiabatic fixed bed reactor. The kinetics are modelled by a reversibll reaction system where each reaction step follows a power law with a reaction order of one in the gas and in the solid component. The t e m p e r a t u r e dependency of the reaction rate constant follows the Arrhenius law. The equilibrium constant is set to be independent of temperature. Material balance for hydrogen:

0---7 = - u ~ z z + e D~. ~-~z2 + ( 1 - e)os

-

-

-

(1)

482 Material balance for oxygen in the solid phase:

O@ox fo~ [k_(Co-c)(i-(9o=)-k+c (9o~] Ot : 3Is 7~M~

(2)

fMe is the fraction of the fully oxidized solid at time zero and .1"o~-is the amount of active

sites (oxygen) needed to convert one mole of the gas component i.

C0 - - C -nt- C H 2 0

fMe

--

rZt~

sites

72tot,catalyst

,

fo:c - -

~ i

|

Enthalpy balance" )'o~ is the amount of active sites (oxiygen) which will be needed to react with one mole of the gas component i.

OT ~IJ i)2T c)T ( 1 - c) os %,s 0t = - 0 6 u ~ + 5 ~ Oz 2 + (1 - c),os I~HRI [-

~HR,red = -

~HR.ox

,

k_(co -

(4)

c)(1 - Go~)+ k+c @o~]

k + - k+0 -exp

--

,

k_ = K c

Balance of the active gas components"

Oco O~ = -

Oco u ~

02co + e D~

Oz 2

(6)

This balance is used instead of a total gas balance because the reaction has no influence on the mole fraction of the inert component. The boundary conditions are used according to Danckwerts [15].

3

Results and Discussion

All the results which will be discussed in this section were obtained by solving the set of differential equations of the model derived above with the simulation tool PDEX [14]. Unless mentioned otherwise the reactor temperature at the beginning of the calculations is set to 573 K which is equal to feed temperature. The inlet concentration during the oxidation (8.0mol/m 3) is twice the concentration of the reducing component during the reduction step. This is because the oxidation reaction is ten times slower than the reduction reaction which is fixed by the equilibrium constant. The inner diameter of the reactor is set to 25 mm and its length to 0.8 m.

483

3.1

:\'ecessit9 of the Flow Reversal

The first, alternative of Fig. 1 (no flow reversal) is investigated with a period less than the time needed for. full reduction or oxidation of the bed. A coupling between the two reaction steps via the solid phase is hence possible. During the initial oxidation steps. the hydrogen formed in the first, part of the bed is consumed by the downstream part which is still oxidized (Fig. :3 a). Almost no hydrogen is released until the process has ,_,6

,_,6 E

- - - 1st P e r i o d - - - 2 n d Perio~

"5

=7 4

- - - 3rd Period

s

- - 4th P e r i o d

cO 0 tO

"'"''",, ~ ....

O0

,

0.0

"',..._ - - - 7

0.2

.

.

.

~x .

I

0.4 Reactor

Length

0.6 [m]

T

0.8

\

""

-- - 0 7 1 t _ p e r " ~.. ..~

-.

,.,,

8 0 0.0

E

i'

0.2

0.4

i

I

I

- - - 0 . 8 1 t per ] m 1.00 t _ p e r

I

~-;,....,:.,....~...2.,."

~2

6th Period

~2

f" /"

o

- - 5th P e r i o d

c-

.o

-- - 0 . 5 2 t _ p e r - - - 0.62 t_per

._.._

J

06

1

0.8

Reactor Length [m]

Fig. 3. Concentration profile inside the reactor without flow reversal (alternative 1) for hydrogen in the oxidation step: (a) at the end of each period during the transient and (b) in the periodic steady state (oxidation: 0.5
484

/ "~ ;'-'-'-..... : l-Z; '2"'.., ........

.'T'~.O.8

/f/'/'. i;'

\

//,ii~/" ~ X~ / I - ' : "~\", /F t

~0.6 s

t

r

,'1

.o 0.4 r ._ x

0 0.2

/

/

I

.~

*

"'"',

9x

I--0.33 t_perI -.

'

o

600

". -.

:

550 .~

9

0

I ;

l

I- ' ~

"-,,,,,

700 650

600 $ i-

. 4 ~

S

550

0 0.2 0

50O

,

0.1 0.2 Reactor Length [m]

0.6

"~'-~0

e

"....~

0.0

650

J'" "0.50t per I

"",.."-..

x ,.

'/ =08~ / 7 " - .

700

]--0.17 t_perJ

0.3

500 0.0

0.1 0.2 Reactor Length [m]

0.3

Fig. 4. Profiles of oxidation degree (thick lines) and temperature (thin lines) inside the reactor with flow reversal (alternative 2) for the first period: (a) reduction step (0.0 < t_per < 0.5): inlet: left (b) oxidation step (0.5 < t_per <_ 1.0): inlet: right heat conductivity. The change of the flow direction yields the expected effects in the beginning of the process. The observed behaviour for the periodic state will be discussed in the next section.

2.2 9

Flow Reversal after one Half Cycle

As mentioned before, in this operation mode the flow is reversed upon each change of feed composition (see alternative 2 Fig. 1)" during the reduction step the flow direction is always from the left side and during the oxidation step from the right side. The process is started with the reduction of the fully oxidized bed at time zero for each case study. The temperature profiles at the end of each oxidation step during the transient (Fig. 5 a) shot, the shift, of the reaction front from left to right. The system has reached the 700

700

t",, / \ -'t',, " 29 ,/ .,,- , . '~k.

-- - 1st Period

-

T 650

K 600-

650

--- 100th Period

"'""~"~.........

,~--i

-- 1000thPeriod

0.0

i~. 600

,

,

0.2

0.4 Reactor Length [m]

w

0.6

"-.. "'-.

El.

#.

550

- - - 0.33 t_per - - - 0 . 5 0 t per -- - 0.83 t_per - - 1.00 t_per

~

- - - 10th Period

0.8

550

t , /

0.0

.."2.

,

i

i

0.2

0.4

0.6

Reactor Length [m]

0.8

Fig. 5. Temperature profile inside the reactor with flow reversal (alternative 2) (a) at the end of each period during the transient and (b) in the periodic steady state periodic steady state when a certain fraction of the bed is reduced which is only a part of the whole bed. So the period time can be increased before hydrogen leaves the reactor during the reduction step. During the oxidation step in the periodic steady state the gas cools down the heated bed first before the reaction starts which can be seen by the parallel lines from 0..35 m to 0.8 m in figure 5 b. Then the endothermic oxidation reaction

485

decreases the t e m p e r a t u r e of the hot spot while it is shifted to the outlet (0.0 m). The local minima seen in figures 5 a and b mark the region of the reactor where the rate of the oxidation reaction has its highest value. These positions are determined by the optimum of temperature, oxidation degree (amount of nickel) and water concentration. One of the interesting process parameters is the integrated conversion which is defined as

.~~ =

1

~end

--

9

~beg

1 -

dt

(7)

Ci.O

tbeg

where i marks hydrogen or steam, respectively. For determing the integrated conversions which can be reached during each step, the optimal period time were evaluated. Optimal means with respect to the shortest transient time. The values of the process variables were found to be 9,000 s for the period time and 0.1 m / s for the gas velocity. The dependency of the integrated conversion for the oxidation and the reduction step from the fraction of the adiabatic t e m p e r a t u r e rises shows a contrary tendency for the two parts of the cycle (Fig. 6). This behaviour of the system can be explained with heat and mass balances. In .,':R. ,:...- 1 . 0 o ._

0.8

> to 0.6 0

g,~ 0 . 4 o') I:::

O.2 0.0

i

-0.6

-0.4

w

-0.2

i

0

Ig ( Tad.re~ / Ta~.o,

i

0.2

)

i

0.4

0.6

[-]

Fig. 6. Integrated conversion vs. fraction of the adiabatic temperature rises for the reduction and the oxidation step in the periodic steady state for alternative 2 (period time: 9.000 s" gas velocity 0.1 m/s) the case of a steady state adiabatic fixed bed reactor the equation for the temperature rise is [16]"

T - To + ( - - X H R ) . q , o

X - To + A T ~ . X

(8)

From the definition of the periodic steady state it follows, that on one side the oxygen removed from the solid by the reduction must be replaced during the oxidation step. On the other side the heat released by the exothermic reaction must be used by the endothermic reaction or leave the reactor by convection. Since equation 8 is a combination of heat and mass balance and t e m p e r a t u r e is a measure for heat it can be used for the

486 reduction and the oxidation step respectively by integration over time. Because of the conservation law of energy these equations must be equal t_per

t_per

0

t_per

2

t_per 2

T0.red + :-kT~d.red " -~'red- t _ p e r . To.ox +/kT~j.ox Xox 2

(10)

The inlet, temperatures are equal here for the two reaction steps which yields"

Xox = I"--XT~d'redt 9Xred

(11)

The conversion for the oxidation step decreases with increasing inlet concentration of the oxidizing agent. It decreases faster when the conversion for the reduction has reached one. The reason for not reaching a conversion of one for each step is because the equilibrium constant is not unity. This constant is responsible for the asymmetric shape of the curves in figure 6. too. The other effect which can be seen here is that a hydrogen yield of about 70 % at 100 c~, conversion during the reduction step can be reached by using the periodic flow reversal.

3.3

Flow Reversal after one Cycle

The advantage of the third alternative for operating reduction-oxidation cycles in fixed bed reactors (flow reversal after each complete cycle) should be an even use of the bed. This is confirmed by the simulations. In addition transient times are relatively short which is helpfull for a technical application. A disadvantage as compared to alternative 2 is the higher process t e m p e r a t u r e during the transient (Fig. 7). 500

300

400 2OO

~" 3oo 200

................ l i o om ............. - - - U = 0.1.6

I00

100 ,

0

20 000

,

I--

o.o5 m/s I

I ' " ~ 1 7 6 rn/sJ

40 000 60 000 Period Time [s]

80 000

0

,

0

5 000

,

,

10 000 15 000 Period Time [s]

m/s Ii

o.o5

m/s I

0.02

m/s I

20 000

Fig. 7. Dependence of the maximum difference between inlet (To) and process temperature (T) on period time and gas velocity: (a) with flow reversal after each cycle (alternative 3) and (b) with flow reversal after a half cycle (alternative 2)

487

4

Conclusions

The reduction-oxidation-cycling in fixed-bed reactors described here displays complex behaviour which was investigated and characterized by simulation studies. Flow reversal is necessary for a good efficiency in m a t t e r and heat for such processes. Higher conversions than at thermodynamic equilibrium may be reached. The optimal point of flow reversal will depend on the reaction system and cannot be determined with the simple model presented here. Further investigations with more detailed models as well as experimental studies are currently being carried out.

Acknowledgements This report is part of a project supported by the German Ministry of Education. Science. Research and Technology (BMBF) under the support number 03D00:3:3C. The authors are responsible for the contents of this publication.

References [1] P. Silveston, R. R. Hudgins, A. Renken. Periodic Operation of Catalytic Reactors Introduction and Overview. Cat. Today 25 (1995) 91 [2] Yu. Sh. Matros. Catalytic processes under unsteady-state conditions. Vol. 43 of Studies in Surface Science and Catalysis. Elsevier. Amsterdam (1989) [3] M. N. Schevendam, R. McBride. R. Deuter. M. Isaacs Porous Metal Filters Prevent Fluid Bed Catalyst Loss. Chem. Processing 46(1983) 100 [4] E. Miiller-Erlwein. J. Guba, Experimentelle Untersuchung zum periodischen Reaktorbetrieb bei der heterogen katalysierten Oxidehydrierung von Isobutyraldehyd zu Methacrolein. Chem. Ing. Techn. 60 (1988) 1072 [5] X. Lang, R. R. Hudgins, P. L. Silveston. Application of Periodic Operation to Maleic Anhydride Production. Can. J. Chem. Eng. 67 (1989)635 [6] R. Burch, R. Swarnakar, Oxidative Dehydrogenation of Ethane on Vanadium-Molybdenium Oxide and Vanadium-Niobium-Molybdenium Oxide Catalysts, Appl. Catal. 70 (1991) 129 [7] R. M. Contractor. DuPont's New Process for n-Butane to Tetrahydrofuran. Appl. Catal. B 6 (1995) N3

[8] L. Weismantel. J. St6ckel. G. Emig, Improvement of Selectivity with a Two-Step Process for the Oxidation of Isobutyric Acid. Appl. Catal. A 137 (1996) 129 [9] M. Mortenson. R. G. Minet. T. T. Tsotsis. S. Benson. A Two-Stage Cyclic Fluidized Bed Process for Converting Hydrogen Chloride to Chlorine Chem. Eng. Sc. 51 (1996) 2031 [10] U. Nieken. Abluftreinigung in katalytischen Festbettreaktoren bei periodischer Str6mungsumkehr, Fortschrittberichte VDI. Reihe 3: Verfahrenstechnik, Nr. 328. VDI-Verlag, Diisseldorf (1993)

488 [11] L. van de Beld. Air Purification by Catalytic Osidation in an Adiabatic Packed Bed Reactor with Periodic Flow Reversal. Ph. D. Thesis. University Twente (1995) [12] Fr.-R. Block. R.-Fr. Speicher. G. Kollers. Der Dampf-Eisen-Prozefi zur Wasserstofferzeugung. Report from the Institute for Producing Iron at the Rheinisch-\'\'estf~ilische Technische Hochschule Aachen Arch. Eisenhiittenkunde 54. Nr. 4 (1983) [13] M. S. Casper, Hydrogen Manufacture by Electrolysis. Thermal Decomposition and Unusual Techniques, Noyes Data Corporation, Park Ridge. New York (1978) [14] U. Nowak. U. Nieken. G. Eigenberger, Full)' Adaptive Algorithm for Parabolic Partial Differential Equations in one Space Dimension. Computers and Chemical Engineering 20 5 (1996) 447 [15] P. V. Danckwerts. Continous Flow Systems, Chem. Eng. Sci. 2 (1) (1953) 1 [16] K. R. Westerterp, 'v\:. P. van Swa.aij, A. A. Beenackers. Chemical Reactor Design and Operation, John W~ley ~" Sons. Chichester (1984)

Used Symbols c

rmoZ] tin3 J

cp

concentration

Co

[k--f-~]

specific heat capacity

D==

EA

[~--~]

activation energy

k

[k-~-~]

Ms

J

[moZ] m 3 1 [m 2 -

tota] concentration of H20 and H2

-7-]

axial dispersion coefficient

AHR

[-4;7]

reaction enthalpy

reaction rate constant

I(G

[--]

equilibrium constant

[~k--~got]

molar mass of the solid

n

[moll

amount

R

[,~J K]

common gas constant

t

[s]

time

t_per

Is]

period time

T

[K]

temperature

AT~d

[K]

adiabatic temperature rise

u

[~]

mean gas velocity

X

[--]

conversion

z

[m]

axial space coordinate

e

[--]

porosity of the bed

)~:~f

w [PUg] effective heat conductivity in axial direction

u

[--]

stoichiometric coefficient

o

[m~]

005

[--]

oxidation degree

density

Subscripts i

component A.i

oz.-

oxidation step

red, +

reduction step

s

solid

0

inlet

i•i••i•i••i••i•i•••i••i•i•!•i•i•i•••i•i•i•i••i•i••i••i•

iiiiiiiiiiiiiiiiiiiiiiiiiil~ iiiiiiiiiiii !iiiiii!ii

i!iiiiiiiiiiiiliii~iiiiiiiiiiiiii iiii ii!i!ii!iiii~i!iiiiiiiiiiiil !iii iiiiil

i!i!iiiiiiiiiiiiiiiiii iiiiiiiiiiiiiiiii i!i!i!!ii iiiiiliiii

iiiiiiiiiiiiii!iiiiiili!iiiiiiiiliiliiiiiiiiiii

This Page Intentionally Left Blank

91997 Elsevier Science B. V. All rights reserved. Dynamics of Surfaces and Reaction Kinetics in Heterogeneous Catalysis G.F. Froment and K.C. Waugh, editors

491

Ether decomposition activity of CaNi5 hydrogen storage alloy H. Imai, T. Hosoya and S. Futami Department of Materials Science & Technology, Faculty of Engineering, Toin University of Yokohama, Kurogane-cho, Aoba-ku, Yokohama, 225, Japan.

The catalytic activity of CaNi5 hydrogen storage alloy is studied for decomposition of diethyl ether in the temperature range between 570 and 720 K by an atmospheric flow method. The alloy showed a catalytic activity at temperatures higher than 570 K. Following mechanism is proposed on the basis of the effect of W/F on the composition of products. Ethane and acetaldehyde are the initial products, while methane and carbon monoxide are the secondary products, formed by the decomposition of the acetaldehyde.

1. INTRODUCTION Hydrogen storage alloys have a number of potential applications in many fields of energy technology. One of the applications in the field of chemistry is the use in the catalyst system. Many hydrogen storage alloys showed characteristic catalytic activities, vg) In the previous paper 8), we studied the methanol decomposition activities of several hydrogen storage alloys, and found that each alloy showed a characteristic selectivity depending on its chemical composition. In the present paper, the catalytic activity of CaNi5 hydrogen storage alloy is studied for decomposition of diethyl ether. 2. EXPERIMENTAL The CaNi5 hydrogen storage alloy was obtained from Ergenics, a division of MPD Technology Corporation. The alloy was crushed and sized (32-60 mesh) for use in the measurements. The BET surface area of the alloy was 0.0639 mZ/g. The decomposition of ether was carried out in the temperature range between 570 and 720 K by an atmospheric flow method. 9) After the catalyst was reduced by hydrogen for 3h at 623K, the temperature was changed to a reaction temperature, and the helium gas containing 185.1 Torr of diethyl ether vapor was flowed through the packed bed. The reaction products were analyzed by gas chromatography with 3-m EGA (15%)/Chromosorb W and 1-m grease (2%)/ activated charcoal columns. XRD spectra were recorded on a Shimadu XD-610 powder X-ray diffi'actometer with monochromic filtered CuK a radiation. Diethyl ether (Grade S) was obtained from Wako Pure Chemical, and dried with Molecular Sieve 4A before use in the reaction. Hydrogen (7N) and nitrogen (5N) were obtained from

492 Nippon Oxygen Co. Helium obtained from Japan Helium Center was purified through a Supelco gas purifier and an OMI purification tube. 3. RESULTS AND DISCUSSION

The alloy showed the catalytic activity for decomposition of the ether at temperatures higher than 570 K. The activity decreased at an initial period, but remained almost constant after 60 min. So the activity was measured after 60 min in this paper. The rate of formation increased with increasing temperature, as shown in Fig. 1. Products were carbon monoxide, methane, ethane and acetaldehyde. 1.0

o L_.

> 0.5 o o o

/ I

i

600

650 760' Temperature (K)

,

750

Figure 1. Reaction of diethyl ether on CaNis. W/F=0.085 g min/m/

.~ 50 o

E

0

0

.2 0

E 0

o

-O~

0

0.1 W/F (g-min/m/)

0.2

Figure 2. Effect of W/F on the composition of products. O ethane, ~ methane, I ) CO, O acetaldehyde

493 Figure 2 shows the effect of W/F on the composition of products at 623 K. Ethane and acetaldehyde were produced at the lowest W/F. In the higher W/F range, methane and carbon monoxide contents were increased with increasing W/F. These facts suggest that ethane and acetaldehyde are the initial products, while methane and carbon monoxide are the secondary products. Methane and carbon monoxide may be formed by the decomposition of acetaldehyde on the catalyst surface. The results show that the reaction occurs via following mechanism.

(1) (2) (3) (4) (5) (6)

C-C-O-C-C (g) ~ C-C-O-C-C (a) C-C-O-C-C (a) ---, C-C (g) + O=C-C (a) O=C-C (a) ---, O=C-C (g) O=C-C (a) --, CO (a) + CH4 (a) CH4 (a) -~ CI-h (g) CO (a) ---, CO (g)

Figure 3 shows XRD spectra of the alloy before and after the reaction. No new crystalline peak appeared after the reaction at 623 K for 2 h, although relative intensities of the peaks varied a little. Changes in relative intensities became more remarkable, at 723 K (2 h). This suggests that an amorphous oxide film was formed on the alloy surface during the reaction.

4

~

4 -1 ~

~

CD 3

2

1

o

10

20

30

40

50

60

70

2 0 (deg) Figure 3. X R spectra of CaNi5 before and after the reaction. (a) original alloy, (b) after the reaction at 623 K for 2h.

80

494 REFERENCES

1. W.E. Wallace, Chem. Tech., (1982) 752. 2. H. Imamura, Shokubai, 25 (1983) 202. 3. F.P. Daily, J. Catal., 89 (1984) 131. 4. T. Imamoto, T. Mita and M. Yokoyama, J. Org. Chem., 52 (1987) 5695. 5. H. Imamura, S. Kasahara, T. Takada and S. Tuchiya, J. Chem. Soc., Faraday Trans. 1, 84 (1988) 765. 6. M.P. Sridhar Kumar, B. Viswanathan, C.S. Swamy and V. Srinivasan, Indian J. Chem., 28A (1989) 19. 7. V. Padmasubhashini, I.A.P.S. Murthy, M.P. Sridhar Kumar and C.S. Swamy, Indian J. Chem., 29A (1989) 1104. 8. H. Imai, T. Tagawa and K. Nakamura, Appl. Catal., 62 (1990) 348. 9. H. Imai, T. Tagawa and K. Nakamura, Report of Res. Lab. of Eng. Mat'ls, Tokyo Institute of Technology, [15] (1990) 61.

91997 Elsevier Science B. V. All rights reserved. Dynamics of Surfaces and Reaction Kinetics in Heterogeneous Catalysis G.F. Froment and K.C. Waugh, editors

Chemical Kinetics of a two component system. A simple rate model

495

phase segregated

Amir A. AI-Haddad* and Johnson Mathew Chemical Engineering Department Kuwait U n i v e r s i t y - PO Box 5969 Safat 13060- Kuwait 1. ABSTRACT Considerable research has been performed on polymeric systems that exhibit liquid crystalline or mesomorphic behavior. The principal reason for this effort was that these materials might be developed as ultrahigh strength materials. Jackson and Kuhfuss from Tennessee Eastman demonstrated that liquid crystalline behavior existed in copolymers based on poly(ethylene terephthalate) (PET) and para-hydroxy benzoic acid (ABA). However, intricate details pertaining to the polyesterification kinetics have remained unexamined. Transesterification reactions between poly(ethylene terephthalate) PET, and acetoxybenzoic acid (ABA) were conducted using the melt polymerization technique to understand the transesterification kinetics of a phase segregated system. The transesterification kinetics of two compositions PET 20 / 80 (ABA) and PET10 / 90 (ABA) have been studied at 260, 275, 290 and 305~ using dibutyl tinoxide (0.1 mole percent) as a catalyst. Homopolymerization of acetoxy benzoic acid was also studied at similar temperatures and catalyst concentration. In the present experimental work moles of acetic acid found experimentally is computed using a standard procedure. The rate constant k is determined. The role of the catalyst is also evaluated. Keywords:Copolymerization kinetics, transesterification reactions, melt polycondensations. 2. I N T R O D U C T I O N Aromatic polyesters undoubtedly represent the most important class of thermotropic nematics. 1-5 Fully aromatic rod-like homopolymers such as poly(poxybenzoate) or poly (p-phenylene terephthalate) melt at temperatures which are too high to form a stable nematic mesophase. 4 However if the regular chemical structure of the homopolymer is disrupted, the melting temperature is reduced and it is possible to obtain thermotropic nematics. 1"5

*Author for correspondence e-mail:[email protected] Fax #: (965) 483-9498

@

m

~,,~~

:r

m

:r

@

@

0

l-,,I ~

@

@ @

O

~o

G

n o

-11

NO rn rn

m ~ -4 Z Z

~

tn rn 0 7)

Jo

0 7)

cn

~

~

rn

m

"o

0 0 0

@ @|174174

t

'

t'

.• ''.

.... '..

,

,

'.-

'.'.

497 Copolyesters of poly(ethylene terephthalate) and 4-acetoxybenzoic acid (PET / oxybenzoate) were synthesized by Jackson et al. through high temperature melt transesterification. H o w e v e r i n t r i c a t e d e t a i l s p e r t a i n i n g to t h e polyesterification kinetics have remained unexamined. In this system, insertion of 4-oxybenzoate moieties, with stiff rod like conformations, into flexible PET chains fosters the development of thermotropic character within a definite range of copolyester composition. Optical and electron microscopy examinations, 6 coupled with data from x-ray and conventional electron microscopy, endorse the existence of ordered domains or lamellar 4-oxybenzoate blocks in the PET 20 / 80 ABA copolyester. The c o o l i n g of a n a n n e a l e d m e l t is k n o w n to t r i g g e r a b i p h a s i c s t r u c t u r e . 6 The foregoing discussion clearly shows t h a t while there is information available on the structure property aspects of PET / ABA systems, but there is a total lack of kinetic information. This work is an a t t e m p t at formulating a plausible phase segregated kinetic model for the PET 20 / 80 ABA and PET 10 / 90 ABA melt copolyesterification reaction. There is no published literature on the melt polyesterification kinetics of a phase segregated two component system like PET / 80 ABA and PET / 90 ABA. Here we explore the kinetics of a two component system wherein many parallel reactions take place simultaneously. Precipitation has been observed during the synthesis of these polymers. 3. E X P E R I M E N T A L 3.1. M a t e r i a l s

4-acetoxybenzoic acid was prepared by sodium hydroxide catalyzed reaction of 4-hydroxybenzoic acid with .acetic anhydride and was recrystallized using methanol. The yield was around 70% and had a melting point of 184~ Poly(ethylene t e r e p h t h a l a t e ) (1800 ~m), of intrinsic viscosity 0.62 (Aldrich chemical company), was used as received. 3.2. R e a c t o r

A 250 ml glass reactor as shown in Figure 1 was used for the melt transesterification kinetic investigations. 3.3. P r e p a r a t i o n o f C o p o l y e s t e r s

Polyesterification kinetics were investigated for i

.

iiiii-

ABA homopolymerization PET / 80 ABA and PET / 90 ABA.

P E T m e l t e d a r o u n d 260~ and degraded around 326~ on k e e p i n g isothermally for 20 minutes. The reaction temperatures chosen were 260, 275, 290 and 305~ A dry nitrogen b l a n k e t was m a i n t a i n e d t h r o u g h o u t the experiments to prevent oxidative degradations. The rate of byproduct evolution (acetic acid) was monitored as a function of time. Sodium acetate (0.1 mole

498 percent) was used as catalyst for the melt transesterification reactions. This h a s h i t h e r t o not been e v a l u a t e d as a catalyst for the s y n t h e s i s of P E T / ABA systems.

4. R E S U L T S A N D D I S C U S S I O N In our earlier study 7 it was observed t h a t second order kinetics was valid for polyesterification of ABA. It is assumed t h a t in the present analysis the s a m e m e c h a n i s m is valid for homopolyesterification. For t r e a t i n g the second reaction it was a s s u m e d t h a t the 4-acetoxybenzoic acid (ABA) monomers approach a P E T homopolyester, followed by a reaction. This also can be t r e a t e d as a second order reaction between a P E T segment and an ABA oligomer. In the h o m o p o l y e s t e r i f i c a t i o n of ABA 7 a tacit a s s u m p t i o n was m a d e t h a t oligomers upto D P = 5 were in the melt.

4.1. Kinetics and M e c h a n i s m The reactions which occur w h e n PET and 4-acetoxybenzoic acid are h e a t e d together or m a i n t a i n e d isothermally at t e m p e r a t u r e s of crystalline P E T m a y be depicted below. 0

II

~

3

0

II

II -OH+

~

o

0

/

H

O

~

II

__-_~

3

o OH+CH

(A)

COOH 3

Oligomer 2

II -OCH C H O 2 2

+CH

- C

---

OH

3

PET

Oligomer

0

+ CH 3 C O O H

Copolymer

Scheme

1

(B)

499 There are m a n y possible ways by which acetic acid can be g e n e r a t e d t h r o u g h route (A), like the d i m e r reacting with a n o t h e r monomer or t r i m e r or t e t r a m e r or other h i g h e r oligomers. W h e n a reactive polymer like P E T is incorporated into this reaction the kinetics becomes even more complicated. Considering reaction A, one finds t h a t for the formation of an even or an odd oligomer w i t h a n y n u m b e r of r e p e a t units, two different functional groups i.e. O II

m C O O H or - O - C - C H

3 r e m a i n as end groups.

Thus one can note the existance of m a n y variables (rate c o n s t a n t s ) for t h e s e independent reactions. It is difficult to solve these m a n y variables analytically. Hence the following a s s u m p t i o n s are m a d e to simplify the kinetic picture.

_

34-

Major p a r t of the ABA homopolymerization reaction leads to the formation of oligomers (dimers). A dimer of ABA can react with P E T molecule. Since dimers are a s s u m e d to be the major p a r t during the initial stages of the reaction, f u r t h e r reaction of ABA can only be with a P E T molecule. Processes exist wherein higher oligomers of ABA are formed, but their rate of reaction is a s s u m e d to be too slow for the generation of acetic acid a n d hence of no consequence in the m a s s balance of ABA. Hence the above two reactions (A & B) are given prime importance.

If Z denotes the d i m e r formed by ABA reaction the following steps are a s s u m e d for the polymerization reaction X -+" X (a-x) (a-x)

kl

)Z-+-HA z x

(I) Z + (x-y)

P

k2 > Z P + H A

(p-y)

y

(y)

where X, Z, P, ZP, HA, a, x, z, p and y denotes ABA, dimer of ABA, poly(ethylene t e r e p h t h a l a t e ) , copolymer, acetic acid, initial concentration of ABA, n u m b e r of moles of ABA converted, initial concentration of dimer, initial concentration of P E T segments and n u m b e r of moles of P E T segments converted. Rate of dimer formation can be given as dx dt

-dX ~=kl(a-x

)2

dt

(1)

From m a s s balance we have

or

x =z +y z -- x - y

(2a) (2b)

Rate of copolymer formation is V~

l~g]

"-__.e_~= ---.__.__v_r= k2 z ( p - y) dt

dt

(3)

500 wi

•dp

dt

dt

-__e_,=

= k2 ( x - y ) ( p - y)

(4)

Total r a t e of acetic acid p r o d u c t i o n can be given by the algebraic s u m of e q u a t i o n s 1 and 4 as d(HA.__...~)= d(x + y) = kl (a - x) 2 + k 2 (x - y) ( p - y) dt

(5)

dt

A s s u m e r a t e of (z) is r a t e controlling step and the incorporation of d i m e r (z) into the polymer is fast so t h a t the concentration of (z) at any time is very small, t h u s production of (z) can be set equal to zero. Thus dz

d ( x - y)

dt

dt

= k 1 (a

- x) 2 - k 2 (x - y) ( p - y) = 0

(6)

T h u s equation (5) can be given as d(HA_____~)= 2 k I (a - x) 2

(7)

dt

Now, we take into account t h a t the reacting system is n o n h o m o g e n e o u s a n d s e p a r a t e s into two different phases i.e. a polymer rich (PET) and poor p h a s e w i t h c o n c e n t r a t i o n s d' a n d d". It is a s s u m e d t h a t the reactions proceed in both the p h a s e s by the scheme shown in (I). Consider a new distribution coefficient (k) which is r e p r e s e n t e d as C'

K .

.

.

C"

m' v'

(8)

.

m" v"

w h e r e m' and m" are the corresponding masses and v ' / v " are two p h a s e volumes r e p r e s e n t e d by 7It is noted t h a t m

m'= ~ m and m" = ~ I+TK 1+~ 1 7K

(9)

For both the p h a s e s we a s s u m e equation to be valid. d(nA) 1

= 2 k I (a'-x' )e

Thus (10)

dt

for the second p h a s e we have d(HA)

1

d(HA)'

1

d(HA)"

dt

1+- 1 7

dt

1 + ~,

dt

(11)

501 Equation (7) can be replaced by d(HA----~) I I += 2dkl t

T(1 x+ 7-)K) K2) 2 1 (a 1-

(12)

where the notations like a, x and HA have similar meaning to that in equation (7). For K = 1 i.e. no phase segregation, equation (12) reduces to equation (7). The new term I1 + T(1-K2_) (1 + 7 K) 2 J1 is an additional term to replace the rate constant kl which is affected by phase separation. Let the new term obtained be represented 1 by kl. Solution of equation (12) under condition of equation (6) and solution of equation (9) and (10) leads to 2a 1 =l+k]at 2a-y-x 1-P

(13)

Figure 2 shows experimental data points for the catalyzed ABA homopolymer system at different temperatures and fitting curves according to equation 13. This figure also indicates t h a t the reaction rate model is adequate. Rate constants and activation energies are listed in Table 1. It is obvious that the catalyst sodium acetate plays a very marginal role. Arrhenius plots for catalyzed and uncatalyzed ABA homopolyesterification reaction is indicated in Figure 3. Kinetics in systems comprising 80 to 90% of ABA were also studied and evaluated according to equation 13 both for uncatalyzed and catalyzed reactions. 1 The quantity is taken as the average degree of polymerization. Number (l-p) of reactions were carried out between 260 - 305~ in steps of 15~ interval. Typical examples are presented in Figure 4 and 5. It is obvious that the experimental data points can be modelled by equation 13. The rate constants for different reactions are given in Table 1. Figure 6 and 7 depicts a typical Arrhenius plot for uncatalyzed and catalyzed PET20 / ABA 80 and PET 10 / ABA 90 composition. Table 1 reveals that no remarkable changes in rate constants and activation energies occur with rising ABA content. Figures 4 and 5 shows slight periodic deviations of experimental d a t a points from the s t r a i g h t line for higher temperatures. This could reflect periodic phase dissolutions and phase separations during the course of the reactions. 5. C O N C L U S I O N Kinetics of a two component system, PET and ABA in which phase separation occurs has been investigated. To retain simplicity of the analysis few assumptions were made. A generalized scheme in which acetic acid is produced through two channels is considered valid for PET rich and poor phase. Kinetically both these reactions were assumed and shown to be of second order with respect to reactants. Steady state approximation has been considered. Parameters were chosen such that the least squares deviation between moles of

Table

tg~

1

t,9

Rate constants from second order plots of ABA and PET / ABA systems

System

Composition

ABA 100

PET / ABA

20 / 80

PET / ABA

10/90

Temperature oC

Uncatalysed rate constant Lit mol 1 sec -1

Catalysed rate constant Lit mo1-1 sec 1

260 275 290 305

0.0118 0.029 0.070 0.082

0.023 0.038 0.080 0.098

260 275 290 305

0.002 0.010 0.018 0.029

0.004 0.029 0.038 0.032

260 275 290 305

0.006 0.018 0.030 0.038

0.01 0.029 0.04 0.055

Uncatalysed EOA (Kcal / mol)

Catalysed EOA (Kcal / mol)

17.1 +_3

15.8 +_2

18.2 +_3

8.8 +_3

19.6 +_4

10.1 +_3

503 19

17 [] A o

15

15

"-"

11

I

--

9

1

0

4

8

12

16

20

TIIME(min)

Fig.2. S e c o n d - o r d e r p l o t i l l u s t r a t i n g t h e effect of t e m p e r a t u r e for s o d i u m a c e t a t e (0.1 mol% c o n c e n t r a t i o n ) c a t a l y z e d r e a c t i o n s for ABA

260~ 275~ 290~ 305~

504

9 (cat)

[] (uncat)

-2

t2

-3

-4--

-6

~

~

!

t

~

l

~

~.

"

~

i

0.00177 0.00178 0.00179 0.0018 0.00181 0.00182 0.00183 0.00184 0.00185 0.00186 0.00187 lfr

Fig. 3. Arrhenius plots for u n c a t a l y s e d and catalyzed ABA reactions

e-t.-

~

I

0

*#

~

OZ.,

O8

:3

--4

b

ot

Po

Ob

03

[

FO !

7 J . m i

FO ! !

pO

1/(1-P) r'd !

cb

13

9 o

po

~n 0

U~ 0

~ u~ Po po Po ~" o to --4 m

II

I

I

L~

o L~

0

el'-

~0

e-F~

I

ee-

0

0

Crl

E

j~

PO 9

GO

-

I

i

0 0

1

I

I

.

i

i

o~

-~9 i

n

i

1/(1-P) i

I

~

PO I

I

l

l

i

~

5 ' 0 to -.4 @ ~. t.n 0 t.~ o

loo.o

l

o~

PO ~ I

c)J 9

OX

L/I

597

-e4 I, u..rdct~T 9 Cs

_2.6l -2-8 -5"0 y c'"

-3"2 -3-4

0

-3"6 -3"8 1

0-00175

l

1

1

0"00179

1

0-00183

I

O 1

O.00157

1/T

Fig. 6. A r r h e n i u s plots for u n c a t a l y z e d a n d c a t a l y z e d P E T 20 / 80 ABA

508

-0.5

9 (eat)

[]

(uncat)

-1

-1.5

-2

-2.5

-3

-3.5

-4

-4.5

i

~

'

0.00172 0.00174 0.00176 0.00178

;

0.0018

!

I

1

0.00182 0.00184 0.00186 0.00188

1/T

Fig. 7. A r r h e n i u s p l o t s for u n c a t a l y z e d a n d c a t a l y z e d P E T 10 / 90 A B A r e a c t i o n s

509 acetic acid produced and theoretically predicted is a minimum.

Plots of

1

(l-p) versus time were found to generate linear fits. The kinetic order for the dimerization steps is determined independently. The polymerization catalyst is found to play a very marginal role. ACKNOWI~DGEMENT The authors would like to acknowledge the generous funding from the research administration of Kuwait University under project number EC 065 without which this work could not be initiated. REFERENCES

1-

Preston, J. Angew, Makromol. Chemie 1982, 109 / 110, 1.

2-

Dobb, M.G. and McIntyre, J.E. Adv. Polym. Sci. 1984, 60 / 61, 61.

_

McFarlane, F.E., Nicely, V.A and Davis, T.G. Contemporary topics in polymer science, Vol. 2 (Eds. E.M. Pearce and J.R. Schae'fgen), Plenum Press, New York, London, 1976, p. 109.

4-

Jackson, W.J., Jr. Br Polym. J. 1980, 12, 154.

5-

Jin, J.I., Antoun, S., Ober, C. and Lenz, R.N. Br. Polym. J. 1980, 12, 132.

6-

Zachariades, A.E., Economy, J., Hogan, A . J . J . Appl. Polym Sci. 1982, 27, 2009.

_

Mathew, J., Bahulekar, R.V., Ghadge, R.S., Rajan, C.R., Ponrathnam, S., and Prasad, S.D. Macromolecules, 1992, 25, 7338.

This Page Intentionally Left Blank

91997 Elsevier Science B.V. All rights reserved. Dynamics of Surfaces and Reaction Kinetics in Heterogeneous Catalysis G.F. Froment and K.C. Waugh, editors

511

Influence of S o m e P h e n o m e n a O c c u r r i n g on the Surface and in the Active Phase of the V a n a d i u m Catalyst on the R e a c t o r D y n a m i c s Krzysztof Gosiewski Institute of Inorganic Chemistry, ul Sowiflskiego 11 44-101, Gliwice, Poland 1. INTRODUCTION During the simulation of the process of SO 2 oxidation, carried out by the present author for a number of years, it became apparent that certain modifications of the Boreskov equations [1,2], introduced following the procedure described in ref [3], lead to a representation of the effective reaction rate that is sufficiently accurate from the practical standpoint. It might seem, therefore, that there is no need to look deeper into the processes occurring in the active layer of a catalyst. This would indeed be the case if the problem could be reduced to calculating the oxidation rate in the mathematical model. The studies on dynamic models conducted by Gosiewski (e.g. [4, 5]) have revealed that the phenomena taking place on the surface and in the active phase of a catalyst, independently of the kinetics itself, can have a profound influence on the accumulation of heat in the catalyst bed. This conclusion has also been corroborated by experimental observations. In reference [5] a review has been presented of the dynamic models employed to describe the stationary catalyst bed. Independently of the degree of their mathematical complexity, the models used so far reduce the thermal capacity of the bed to a term that contains only the specific heat of the catalyst, with all other ways of accumulating heat omitted from the analysis. The present study is concerned with certain phenomena that take place on both the external and internal surface of a pellet of the industrial vanadium catalyst during SO2 oxidation. These are phenomena specific to this type of catalyst and reaction. We cannot exclude, however, the possibility of similar phenomena manifesting themselves, albeit on a different scale, in other catalytic processes. The oxidation of SO 2 over the vanadium catalyst takes place in the liquid melt phase of active species (mainly V20 5 and K2SO4). This observation, described extensively in the literature (e.g. [6, 7, 8]) seems to have been proved beyond any doubt, and will be discussed here only with regard to consequences it may have in the description of the bed dynamics. In recent years two papers have appeared [11,12] which support an earlier observation made by the present author, namely, that the phenomena occurring in the active phase of a catalyst can have a profound effect upon the dynamic parameters of the model. In this study the available information concerning this problem is reviewed and substantiated by the author's own measurements.

2. ACCUMULATION OF HEAT IN THE VANADIUM CATALYST BED EFFECTIVE THERMAL CAPACITY OF THE BED The absorption and desorption of gases from a porous catalyst surface, along with the dissolution of gaseous species in the liquid phase of the melt of active components, have a considerable impact on the effective thermal capacity (of. [4,5]). These phenomena affect the

512 measured accumulation of heat in the bed, as they lead to the increased uptake of heat from the heating gas when the bed is warmed up, and the enhanced release of heat when the bed temperature drops. In order to obtain an agreement between the time scales of the actual and simulated transients, a simple way of including these phenomena into the mathematical model has been proposed in ref. [4] instead of the specific heat of the catalyst, C k, a certain effective value C k eft is introduced that can be identified by comparing the actual dynamic transients in a reactor with those simulated using the model. It can be easily shown that, in practical situations, in the models discussed the time scale is affected only by the product (Pu Ck)- The suggestion that the time scale can also be influenced by the thermal capacity of the reactor body was dropped based on appropriate calculations which included this parameter (cf. Appendix A in ref. [4]). The ratio Ck eft / Ck at which the time scale predicted by the model becomes equal to that corresponding to the actual start-ups of industrial catalytic oxidation plants is Ck eft / Ck ~ 1 - for the warming up of a reactor with the hot air free of SO 2, at a temperature of up to 400 oC; Ck eft / Ck ~ 2 - for the warming up with the SO2-containing process gas, at a temperature above 400 ~ following the initial warming up with hot air; Ck eft / Ck ~ 3 to 5 - for the warming up with the SO2-containing process gas, during the start-up following short breaks in the reactor operation, before which no purge was carried out. This approach, which is formally equivalent to identifying an unknown (or difficult to determine) parameter has raised some doubts. It seems therefore reasonable to try to study this question in more detail. A short discussion of this problem has been presented in ref. [4]. It has to be noted that the heat of physical adsorption of SO 2 over silica gel at 273 K is about 34-103 kJ/kmol; for other gases it is in the range 5-100.103 kJ/kmol. The heats of chemisorption for gases are of the order 50-500 kJ/krnol. In ref [5] the heat effects that accompany these phenomena have been estimated as equal to at least 100 kJ/(1 mol) of the gas desorbing from the surface. The thermogravimetric studies carried out by the present author (cf. [5]) over the range 300 - 600 oC for several catalyst samples led to the change in the mass of a sample from 1 to 5% of the initial mass, and to the value of Ck eft from 2.33 to 4.8 J / (g. K). Similar studies of Br6tz will be discussed in section 3.1 3. ADSORPTION OF GASEOUS SPECIES IN THE LIQUID MELT PHASE OF THE ACTIVE CATALYTIC COMPONENTS 3.1 Thermogravimetric studies of Bri~tz et aL [10] Some interesting results of thermogravimetric studies were published by Br6tz et al. in 1979 [ 10]. It follows from these studies that samples of the vanadium catalyst pellets, heated up to above 400 oc and treated with gases of varying compositions (from the pure inert gas, N2, to a mixture of the inert with the active components, SO2/O2/N2) show the change in mass from 2 to 11% of the initial mass (depending on the composition of the gas used). The results presented in [ 10] provide a valuable source of information on the nature of the phenomenon. The change in mass of the activated catalyst can be observed even in these cases when the sample is treated with gases which do not react chemically (N 2 or the mixture O2/N2). An increase in mass at temperatures of 440 and 520 ~ was 2% for the pure N 2 and 5% for the mixture O2/N 2. These changes are, however, much more pronounced for the mixture SO2/N 2 (7 - 8%), and reach a maximum for the mixture SO2/O2/N2, i.e. the one which enables the oxidation to SO 3 to occur. The increase in mass observed in this case was 10 -11% of the initial mass of the sample. The studies were preceded by careful activation of

513 the catalyst, which consisted in the sulphatation of the activator K20 to K2SO 4. Thus, it seems obvious that the changes in mass do not result from the activation process itself. Of particular interest are the transients of change in mass of the samples recorded in the study of Br0tz et al.. [ 10], after the samples were subjected to a cyclic saw-tooth temperature variation from 375 to 575 ~ Such a variation leads to the cyclic changes in the catalyst mass, of the amplitude up to 4.6% depending on the composition of the gas mixture. These transiens reveal that the variation in mass is a totally reversible process and that it becomes significant only above the melting point of the active species, i.e. above 390 - 400oc. It seems certain that the phenomena found by Brotz et al. [10] cannot be explained otherwise than by the absorption (and desorption) of the gaseous species into a catalyst sample. The authors explain these phenomena by the dissolution of the gaseous components in the liquid melt phase. Although, in the context of the results presented in ref. [ 10] this is only a hypothesis, the possibility that this is indeed the case remains very high. 3.2

Discussion of the model of Bunimovich et al.

[11]

In ref. [ 11 ], Bunimovich et al. have made an attempt to include the process occurring in the liquid phase into the mathematical description of the bed. They stress an "extremely high capacity" of this phase with respect to SO 3, which leads to a high time constant for the accumulation of this component in the bed. The authors relate these phenomena especially with the absorption of SO 3 into the liquid phase of the catalyst, and support this idea by Gas

Surface reaction

SO [g/l~m2] 0.4 ,, "

9

b

, <35 ~

-\ 0,3 r \ X

07_

0,1S ' - ~ .

Fig. 1 Schematicdiagram of the reaction zone

:i

i

:

0,25,\

0,1

I

,

~

,

i

i

i

i ~

I

i

: I

i :

: '

; .

1

t i at the outlet!from pass iii ,,,y . . ~

.~N,

~ ' " , d. ~

!. ~.

a~ the 9udet frorri pass iV

~ i~/

Oi , I I 0 10 20 30 Chemical analysis results: 9

,~

I

III pass

9

! _ ~9

i

I

IA

40

50

Z, .

== 60 Time

70 [hours]

IV pass

Fig. 2 Contentof Sulphur oxides expressed as SO cor~centration at the reactor outlet during the purge of the bed. pointing to a very high difference between Henry's law constant for SO 3 and for the other gaseous components. In reference [11] the equations of a mathematical model are also presented; the model assumes, however, the absorption of all the gaseous species into the liquid phase. Also, the computer simulation results are Nven, with emphasis placed upon the accumulation of the mass of SO 3 in the melt. Bunimovich et al. do not discuss the associated heat phenomena although the model proposed in [11] includes the description of such phenomena. Despite some formal objections which can be raised with regard to this model, both the reasoning and results presented in [11] seem interesting, even if they contradict to some extent the expenmental results of Brrtz et al. [10]: if the absorption of the gaseous species in the liquid melt concerns only SO 3, then the change in mass of the catalyst, found in

514 ref. [ 10] even when SO 3 could not possibly be produced, becomes difficult to explain. It may be that some other surface phenomena come into play. The model does, however, require some discussion. If we assume that it is indeed SO 3 (i.e. the reaction product) which is mainly absorbed in the liquid phase, then the natural consequence of such an assumption should be the occurrence of the reaction mostly on the surface of the liquid phase (cf. Fig. 1) If, however, all the gaseous species are assumed to dissolve in the melt, then the reaction can take place in the bulk of the melt. 4. DESORPTION OF SULPHUR OXIDES DURING THE PURGE OF THE REACTOR BEFORE A ROUTINE BREAK IN OPERATION

Before switching off a catalytic SO 2 oxidation reactor to carry out routine overhauls, the reactor is purged with the hot dry air to remove sulphur oxides from the bed. The large mass (or molar) capacity of the catalyst bed is well illustrated by the fact that, in order to obtain the outlet concentration of sulphur oxides below 0.001~m 3, the reactor has to be purged for several dozen hours. To estimate this capacity the measurements were carried out of the outlet concentration of sulphur oxides during the purge. Typical concentration of sulphur oxides at the outlet from passes III and IV of an industrial reactor during the purging procedure are shown in Fig. 2. The first three passes were purged independently of the parallel pur~ng of pass IV. The reactor was purged with a total of about 40,000 m3(STP)/h of hot air (inlet temperature--- 440 ~ It is difficult to determine the actual molar capacity of the catalyst based solely on the measurements performed since, due to the heat losses during the purge, only pass I remained at a temperature above the melting point of the melt (about 400 oc). Moreover, during the flow of the gas through cold heat exchangers a large part of SO 3 blown out of the bed condenses to form H2SO 4 or even oleum and undergoes corrosive reactions. It follows, however, from the measurements performed that during the purge more than 250 kg of SO 3 is removed from the reactor, the major part of which probably originates from pass I which contains abut 13,000 kg of the catalyst; these amounts are therefore quite considerable. 5.

TIME DELAY OF THE MAXIMUM EMISSION OF SO 2 DURING THE

START-UP OF THE REACTOR

As Mishra and Sebastian [12] have pointed out, during the start-up of the reactor, following the heating of the bed from an external heat source (the so-called "cold" start-up) a well marked time delay (of about 45 - 60 minutes) is observed in the appearance of the maximum SO 2 emission at the outlet, compared with the transient predicted by mathematical models that neglect the absorption in the active liquid phase. The studies carried out by the present author [4] on the simulation of the "cold" start-up of an industrial installation also reveal a delay in the appearance of the maximum SO 2 emission by almost an hour. Although this phenomenon has been explained by the slow attainment by the sulphur burner of its full operational capacity, it is also possible that the conclusions presented in ref. [ 12] may, at least in part, be correct. It is suggested in [12] that delay is associated with the solubility of the gaseous species in the liquid melt (cf. Brotz et al. [10]) Mishra and Sebastian regard the dissolution of SO 2 (rather than SO 3, as suggested by Bunimovich [ 11 ]) as being of primary importance. If the melt absorbed only the product of the reaction, the delay in the appearance of the maximum SO 2 emission would be difficult to explain. In order to estimate the "molar capacity" of the bed that might produce an almost one hour's delay a number of simulations were carried out. The simulations were based on a dynamic model of a reactor bed, in which, together with the already mentioned effective

515 accumulation of heat, a term C m has been introduced which artificially allows for the accumulation of SO 2. The equations describing model of the bed are therefore

Pu Ck eff

~ Tk ~0 2 Tk + a S ( Tg - Tk ) + ( - AH ) r Pu & =2eft 0l 2

Cm

~Xsoo Ot -

= - ng

OXsoo Yl _ _ r . P u

(2)

(3)

Since the delay found in references [4] and [12] concerned the start-up of the whole oxidation plant that included a multi-pass catalytic reactor, the simulations were carried out using the model of the whole plant comprising not only the reactor itself, but also the installations for the recovery of the heat of reaction. For the assumed values of C m the startup was simulated from the state in which the installation was thoroughly cooled down. The SO 2 concentration transients at the outlet of the reactor are shown in Fig. 3. 6.

CONCLUSIONS The phenomena occurring on the surface and in the liquid active 0.0351 =0 phase of the vanadium catalyst have undoubtedly a major effect on the 0 . 0 3 0 ~ C m= 1 [ transients calculated by the dynamic models of catalytic reactors. The physicochemistry of these phenomena has not, so far, been satisfactorily elucidated. The introduction of an effective specific 0.010[ ~\ l___ [ heat and effective molar capacity of the bed as parameters to be identified by comparing the actual transients with those predicted by 0 0.5 1.0 1.5 2.0 2.5 3.0 3.5 4.0 Time [ hours ] the model is only a temporary measure. It cannot be expected that Fig. 3 Simulation of the SO 2 concentration at the the phenomena occurring on the outlet from an oxidation plant with a reactor packed with 82.4 m 3 of the vanadium catalyst, for the different values catalyst surface can be thoroughly of the mass accumulation, Cm and for explained without expensive and tedious studies. The results of the Ck eft = 2.1 kJ/(kg.K) present investigations can be summarized as follows: (1) The heat effects associated with the phenomena taking place in the active phase of the catalyst can lead to severalfold prolongation of the time scale of the dynamic mathematical model of the bed. (2) If the absorption (and desorption) of gases in the amount of about 5 10% of the catalyst mass (about 1 to 2 kmoles of gases / m 3 of catalyst) results from the change in temperature by 200 K, the accumulation of heat in the bed will change by several kJ / (kg- K). For instance, the heat effect of these phenomena expressed in SCb molefraction [- ]

~176176l/f >--l~ /

I t

516 terms of the effective specific heat will be 2.6 to 6.8 kJ / (kg 9K) for the physical adsorption and 7.6 to 20 kJ / (kg- K) for chemisorption, with the physical specific heat of the catalyst as low as about 1 kJ / (kg- K) ! (3) Based on the thermogravimetric studies [10] it can be concluded that all the gaseous species are absorbed into the catalyst to a similar extent, although it is not clear whether this is due only to the dissolution of the gases in the liquid melt phase. However, it follows from these studies that this process is fully reversible. (4) About 4 0 - 60 minutes' delay in the appearance of the maximum emission during the "cold" start-up of an SO 2 oxidation reactor cannot be plausibly explained if only the absorption in the liquid melt phase is taken into account. This would require the "molar capacity" of the bed, C m, of about 10 kmol / m 3. At such an assumption, the amount of SO 2 retained in the bed during the first 40 minutes of the start-up would correspond to the change in mass of the catalyst of 8 - 10%. Therefore, it seems little probable that it is solely the absorption of SO 2 in the bed which is responsible for such a considerable shift in the maximum emission in the unsteady state. Nomenclature C~, specific heat of the gas [kJ/(krnol .K)] Ckeff -effective thermal capacity of the catalyst [kJ/(kg.K)]

c~ (-a~) l-

~g

-

r-

xi -

-

effective molar capaci~ for SO?

[kmol/m 3 ] heat ofreaction [kJ/kmol] distance coordinate along the bed [m]

S -

extemal specific surface area of the catalyst [m2/m 3] Tg, Tk gas and catalyst temperatures, respectively [K] t-

time [s]

or-

Greek letters individual heat transfer coefficient [kJ/(m 2- K-s)]

molar gas flux [kmol/(m2. s)]

~eff-

effective heat conductivity. K)]

effective rate of oxidation of SO9. to SO 3 [kmol/(kg 9s)]

Pu -

bulk density, of the catalyst [kg/m 3]

[W/(m-

mole fraction of component i

References 1. G.K. Boreskov., R.A. Buyanov, A.A. Ivanov, Kmet. Katal. 8 (1967) 153 - 159. 2. G.K. Boreskow, M.G. Slinko, V.S. Beskow, Khim Prom. 3 (1968) 13-18. 3. R. Sztaba, K. Gosiewski, Ir~. i Ap. Chem. No 6 (1981) 3 - 7. 4. K. Gosiewski, Chem. Engn and Proc. 32 (1993) 111 - 129. 5. K. Gosiewski, ~eria Chem. i Proc. 3, (1994) 393 -413. 6. H. Livbjerg, J. Villadsen, Chem. Engn. Sci. 27 (1972) 21 - 38. 7. G.H. Tandy, J. Appl. Chem. 6 (1956) 68-74. 8. G.K. Boreskow. W.W. Illaronov, R.P. Ozierov, E.W.Kildisheva, Zumat Ob. Chim. vol. XXXIV (1964) 23-29. 9. S.J. Gregg, The Surface Chemist~ of Solids Chapman & Hall Ltd London, 1951 10. W. Br6tz, B. SchOnbucher, H. Issler, Ger. Chem. Eng. 2 (1979) 108. 11. G.A. Bunimovich, N.V. Vemikovskaya, V.O. Strots. B.S. Balzh aev. Yu.Sh. Matros, Chem. Eng. Sci. 50 (1995) 565 - 580. 12. J.C. Mishra, M.T. Sebastian, Proc. Int. Conf. Unsteady State Processes in Catalysis USPC-1 Novosibirsk (1990) 659 - 664

91997 Elsevier Science B. V. All rights reserved. Dynamics of Surfaces and Reaction Kinetics in Heterogeneous Catalysis G.F. Froment and K.C. Waugh, editors

517

A Detailed Kinetic Model for the Hydrogenolysis, Isomerization and D e h y d r o g e n a t i o n of n-Butane. Iv~in Machfn, Trino Romero, M. M. Ramfrez de Agudelo. INTEVEP S.A., Apdo. 1070A Caracas Venezuela.

A quantitative kinetic model, denominated TC4, for the catalytic conversion of n-butane is proposed. The model considers 56 elementary reactions, six of them were chosen to occur in heterogeneous phase. The TC4 model can be used to predict the product distribution and the heterogeneous rate constants for a wide range of conditions and on different catalyst types. The model can fit also the experimental data from the isobutane dehydrogenation reaction. A plot, that we have denominated "the graphic's performance of a catalyst", is proposed for the evaluation of the maximum yield of a catalyst with a minimum of experimental data.

1. INTRODUCTION The production of light olefins represents for the petrochemical industry the starting point for the generation of high value products such as: polymers, gums, resins etc. As a consequence, an enormous increase in the demand of light olefins has been brought about. In order to face this demand, besides the developed processes such as steam-cracking and FCC which generate large quantities of C4 and C5 hydrocarbons by-products, new processes are also needed to add value to these light hydrocarbons, and to mantain or increase the fuel pool. The aim of developing a kinetic model might be to optimize the conversion of the reactants toward a given desired product, for example, or, to have the possibility of screening catalytic conditions in order to minimize experimental trials. This would led to a reduction in new catalysts development costs. A summary of some of the available kinetic models for the catalytic conversion of light saturated hydrocarbons (n-butane, isobutane) is shown in Table 1 (including de TC4 model proposed in this study). We use the definitions of kinetic model and reaction scheme given by Knrzinger [1] and Weller [2]. The TC4 kinetic model was based on formal reactions. The proposed intermediates not necessarily have physical sense in this model. We only pursuit to propose a predictive kinetic model with the objective of estimating products yields for a given reaction condition. In this way, the TC4 model describes the product distribution for the catalytic conversion of n-butane. The following products were considered: hydrogen, methane, ethylene, propane, propylene, n-butane, isobutane, iso,cis,trans, l-butene. The TC4 model is the first quantitative model based in free radicals and is applicable to a diversity of data from different catalysts obtained in our laboratory as well as by other authors (reported in the open literature).

518

Table 1 Summary of available kinetic models for the catalytic conversion reaction of light saturated hydrocarbons (n-butane, isobutane).. Reaction Kinetic Model # reac. Intermediary # param. Ref. Type Type Isobutane dehydrogenation Qualitative 5 * [5] n-Butane disproportionation (Classical-mechanism)

Qualitative

7

Caboniumion (octyl ion)

-

Isobutane conversion (Steady-State Approx.)

Quantitative scheme of reactions

6

Carboniumion

3

Propane, ethane, isobutane and n-butane hydrogenolysis

Quantitative scheme of reactions

8

-

8

Propane, ethane and n-butane hydrogenolysis

Quantitative scheme of reactions

8

*

10

[9]

Isobutane dehydrogenation (various shemes of reactions)

Quantitative scheme of reactions

5

*

6

[lo]

n-Butane conversion

Quantitative

21

Carbonium ion

8

[11]

n-Butane and isobutane conversion

Quantitative (TC4 model)

55

Free radicals

6

This work

# reac. # param.

= = =

[6]

Number of elementary reactions considered in the model. Number of fitting paramers in the model. Half-hydrogenated fragments adsorbed onto catalyst.

2. T H E T C 4 K I N E T I C M O D E L The reactions considered in the TC4 model are shown in Table 2. The model involves 6 adjustment parameters associated with heterogeneous kinetic constants (see reactions 1 to 6 in Table 2). The differential equations associated to the rate laws of the elementary reactions proposed in this study were solved by using a Fortran program developed by Braum and coworkers [3]. The kinetic constants from the homogeneous phase reactions were obtained from the National Institute of Standards and Technology (NIST) database [4] which the following kinetic constant definition holds: Kc = A T b exp(-E/RT)

(7)

where T and R are the Kelvin temperature and the ideal gas constant (R = 2 cal./mol Kelvin), E is defined as Ea/R and A and b are the empirical constants.

519

Table 2 TC4 reactions for the catalytic conversion of butane. N~

REACTION

A

1

n-C4H 10 + catalyst ---) sec-C4H9 + H

-

B -

-

2

n-C4H 10+ catalyst ~ C2H5 + C2H5

-

-

-

3

n-C4H10 + catalyst ~ CH3 + C3H7

-

-

4

n-C4H 10 + catalyst ~ i-C4H 10

-

-

-

4a

i-C4H 10+ catalyst ---) n-C4H 10

-

-

-

-

-

i

-

I

-

5

i-C4H10+ catalyst ---) i-C4H8 + H2

5a

i-C4H8 + H2+ catalyst ---) i-C4H10

6

C2H6 + H2 + catalyst ~ CH4 + CH4

-

-

H + sec-C4H9 ~ n-C4H 10

1.xl08

0.

n-C3H7 +CH3 ~ n-C4H10

2.09x1013

C2H5 + C2H5 ~ n-C4H 10

1.08x1013

10

n-C4H9 + H2 ~ n-C4H 10 + H

3.8x1013

0.

11728

11

sec-C4H9 + H2 ~ n-C4H 10

3.98x1013

0.

8907

12

sec-C4H9 + H ~ n-C4H 10

1.xl013

0.

1.29x1013

0.

13

sec-C4H9 ~

1-C4H8 + H

~

i 0.

.

.

.

o

18300

14

sec-C4H9 ~ cis-C4H8 + H

4.17x1012

0.

17500

15

sec-C4H9 ---) t-C4H8 + H

4.57x1012

O.

17100

16

n-C3H7 ~ C3H6 + H

1.xl014

18763

.

i

17

C3H6 +H --~ n-C3H7

7.24x1012

0.

18

C3H6 + H ~ iso-C3H7

4.x1012

0.

481

19

n-C3H7 + H2 ~ C3H8

2.6x1012

0.

7448

20

iso-C3H7 + H2 ~ C3H8

2.6x1012

o.

7448

21

iso-C4H9 + H ---) iso-C4H 10

1.xl013

0.

22

iso-C4H 10 ---) iso-C4H9 + H

lxl013

46000

23

iso-C4H 10 ~ iso-C4H9 + H2

1.98x1014

4715

24

iso-C4H8 + H ~ iso-C4H9

1.26x 1011

25

C3H6 + CH3 ~ sec-C4H9

1.78x1011

26

C3H8 ~ n-C3H7 + H

1.xl013

27

C3H8 --~ iso-C3H7 + H

1.xl013

28

CH3 + C2H5 ~ C3H8

2.83x1013

-0.5

29

C2H5 + CH3 --~ CH4 + C2H4

1.13x1012

-0.5

30

C2H5 + C2H4 ~

31 32 33

i

1-C4H8 + H

0.

1459

0

3548 49000 47500

5.2x107

C2H4 + H --) C2H5

2.64x1013

C2H6 ~ CH3 + CH3

1.18x1018

-1.79

45834

CH3 + CH3 ~ C2H6

2.64x1013

-0.64

0

1087

34

C2H6 ---) C2H5 + H

1.xl013

0.

49000

35

C2H5 + H2 ~ C2H6 + H

2.48x109

3.6

4253

520 Table 2 TC4 reactions for the catalytic conversion of butane. N~ REACCION

B

E

36

CH3 + H2 --+ CH4 + H

1.52x1010

3.12

4384

37

C2H6 + H --+ C2H5 + H2

2'53x1011

3.5

2600

3S

C3H6 + CH3 ~ sec-C4H9

1.78x1011

0.

3548

39

iso-C4H 10 + H --+ ter-C4H9 + H2

5.11x1013

0.

3030

40

H + iso-C4H8 ~ ter-C4H9

1.6xl013

0

75s

41

ter-C4H9 --+ iso-C4H8 + H

4.68x1014

0.

19827

42

ter-C4H9 + H ---> iso-C4H10

1.xl013

0.

43 44

ter-C4H9 + H2 ---> iso-C4H10 + H

2.3x1012

0.

8540

CH3 + C3H6 --+ iso-C4H9

2.x1011

45

2.3x107

0. 0.

4390 0.

46

C2H5 + C2H4 --+ C3H6 + CH3 C2H5 + C2H4 --->n-C4H9

1.58x1011

0.

3674

47

C2H4 + H2 --+ C2H5 + H

1.02x 1013

0.

34300

H + C2H5 --+ C2H6

3.61x1013

0.

0

CH3 + CH3 --+ C2H5 + H

8.x 1014

0.

0

50

C2H6 --->C2H4 + H2

2.4x1016

O.

44010

51

C2H6 + H --->CH3 + CH4

5.4x1013

0.

5852

52

ter-C4H9 + H ---->H2 + iso-C4H9

0.

0

53

iso-C3H7 + CH3 --+ iso-C4H 10

0.

0

54

iso-C3H7 + H --+ C3H8

3.55xK42 2.09x1013 2.x1013

0.

0

55

n-C3H7 + H --+ C3H8

2xlO 13

0.

0

0. .

49

3. RESULTS AND DISCUSSION. This section will discuss the results obtained firstly through a comparison of the kinetic model with experimental results taken from different sources and its application to the construction of the yield curves of specific products as a function of temperature and contact times (the catalyst performance curves).

3.1. Comparison of the TC4 results with experimental data. A plot of calculated results vs. the experimentally observed values is shown in Figure 1. The predicted product distribution obtained from the TC4 model was used to evaluate the selectivity towards hydrogenolysis and isomerization reactions.

3.2 The catalyst performance curves. Figure 2 is a plot of isobutene yield against temperature at various contact times during the conversion of the n-butane catalyzed by the system Pt,Sn/H-Mor. The contact time is the number at the top of each bar. The last bar for each temperature, named EQ, represents the value of the isobutene equilibrium yield evaluated for a mixture composed by n-butane, isobutane, isobutene and the lineal butenes (1-butene, cis and trans-butene).

522 One only needs to evaluate the kinetic parameters for a given contact time and temperature and the kinetic model TC4 can generate (extrapolate or predict) the isobutene yield values for a set of contact time values. The maximum isobutene yield which can be obtained from a given catalyst might be evaluated. Therefore, the performance of any catalyst can be determined with a minimun of experimental data.

4. CONCLUSIONS The agreement found with the considered data confirms the predicting capabilities of the proposed model.

REFERENCES

1. 2. 3. 4.

5. 6. 7. 8. .

10. 11.

12. 13. 14.

Kn6zinger H., Kochlefl K., Meye W., J. Catal., 28, (1973) 69. Weller S. W., Catal. Rev.-Sci. Eng., 34, (1992) 227. Braum W., Herron J., Kahaner K., Int. J. Chem. Kinetics, 29, (1988) 51. Westley F., Herron J., Cetanovic R., Hampson R., Hallard W.; "NIST Chemical Kinetics Database, Version 2.0 July 1990", National Institute of Standards and Technology, U.S., Department of Commerce, Gaithersburg, Maryland 20899. Horiuti J., Polanyi M., Trans. Faraday Soc., 30, (1934) 663, 1164. Fuentes G., Gates B., J. of Catalysis, 76, (1982) 440. McVicker G., Kramer G., Ziemiak J., J. Catal., 83, (1983) 286. Leclercq G., Leclercq L., Bouleau L., Pietrzyk S., Maurel R.; J. Catal., 88, (1984) 8. Bond G., Hui L.; J. Catal., 137, (1992) 462. Kao Juo-Yu, Piet-Lahanier H., Walter E., Happel J., J. Catal., 133, (1992) 383. Dumesic J., Rudd D., Aparicio L., Rekoske J., Trevifio A.; "The Microkinetics of Heterogeneous Catalysis", Editorial Professional Reference Book, Am. Chem. Soc., W~.vshington, DC 1993, p.259. INTEVEP unpublished results. Maness J. Jr., Dooley K., J. Catal., 117, (1989) 322. Bond G., Hui L.; J. Catal., 137, (1992) 462.

91997 Elsevier Science B. V. All rights reserved. Dynamics of Surfaces and Reaction Kinetics in Heterogeneous Catalysis G.F. Froment and K.C. Waugh, editors

523

Limitation of Metal Particle Size to Carbon Chain Growth in Fischer-Tropsch Synthesis Yannan Yang, Kechang Xie and Xiaoying Li Institute of Coal Chemical Engineering, Taiyuan University of Technology, Taiyuan, Shanxi 030024,P. R. China

1.INTRODUCTION In 1951 Anderson[l] established the production distribution formulation of the FischerTropsch synthesis fiTS), which is called Anderson-Schulz-Flory (ASF) formulation .Since then, for a long time it is almost always possible to describe FTS product distribution by ASF formulation, which has the following mathematical expression: @n=~lOt"-1 where @, is the number of mole of hydrocarbon with carbon number n and ot is the probability of chain growth. However, in recent years, a lot of negative deviations from ASF distribution were observed on highly dispersed catalysts [2-11].Some authors have suggested that the negative deviation might be caused by external origin, such as wax deposition in transient condition .Some deviations are due to these external reasons,but many others are not ,especially for the recent experimental results by Snel,who claimed that all the external effects were eliminated in his experiment and the negative deviation was only chemically induced [8] .A number of papers suggest that the carbon number distribution is related to the size of metal crystallite, particularly, that the length of the growing chain is limited by the dimensions of the metal crystallite. Aiming at explaining the phenomenon ,Nijs and Jacobs did lots of work ,and many interesting results were obtained [3,12]. They postulated that the simple chain growth scheme occurs,but that the chains are terminated at certain carbon length that is proportional to the size of metal crystallite in the catalysts.Later in 1992,Y.Yang et al [13], based on the restriction of the dimension of the metal crystallites to the carbon chain growth ,established a new FischerTropsch product distribution formulation, which extended ASF model .This new model has two

524 outstanding advantages.One is that it can return to ASF formula for FTS on the catalysts of normal size. Therefore, to some extent, the ASF model can be regarded as a special case of this new model .The other advantage is its simplicity and that using a formula the deviation can be described quantatively.The Yang's formulation can be expressed as: O. = ~lOt'lq(A.) In Yang's model, a restricting factor q(A,) was introduced,which reflects the cut -off effect on chain length by metal particle or crystallite size .Thus, according to this model ,on a given size metal particle not carbon chain of any length can be formed.In other words, for a molecule with certain carbon number can only be produced on parts of the catalyst metal surface, not on the whole.when the model was used to fit the FTS results on small particle catalysts, good agreements were achieved.[ 13 ] The physical meaning of q(A~) has been discussed in detail in Yang's article[13].It has direct relationship with the metal particle distribution, and can be regards as an integral of the size distribution function. In this article, it is approached that the functional relationship between particle size and the maximum hydrocarbon chain length formed on it, and the functional properties of q(A,) is also investigated.The results support the viewpoint that the negative deviation from ASF distribution is caused by the cut-off effect of small particle size.

2. T H E O R Y In 1980,Nijs and Jacobs [3] postulated that a linear relationship existed between particle size and the maximum hydrocarbon chain length which can be formed on it .Here, we assume that ,when the particle size increase a given value, the maximum molecule formed on it adds one carbon number .Thus, it is easy to figure out that a linear relationship exists between particle size and maximum product formed on it ,which can be expressed as: x=bn+a

(1)

where x is the particle size, n the carbon number of the maximum molecule formed on it, b and a are constants. Generally the particle or crystallite size distribution is described by a Skewed Gaussian function[14], but unfortunately the indefinite integral of this function can not be obtained, so it can not be used to meet our mathematical needs here. O f the other distribution functions, Releigh function has the kind of property, which is very similar to the particle size distribution, and its indefinite integral can be obtained.The Releigh function seems rather simple to be used to describe the particle size distribution, but to a great extent it will serve our purpose well here.

525 The Releigh function is expressed as: x2

(2)

p(x)=~-e

where g is a constant, (rc/2)la the geometric mean of the distribution, and (4-rt)l.t2/2 the variance of the distribution.. The equation (2) is integrated as follows: x2

p(x)dx =

e

x:

2

:

=e

X2 2g:

(3)

where the x represents the particle size. From the definition of the q(A~) [ 13 ], x2

j-f

---

q(An)= p(x)dx : e 2,:

(4)

substitute x with equation (1), l_.(bn+a)2

q(A.)=e

2g"

(5)

lnq(A.)=-+(bn

+ a) 2

(6)

[-lnq(A.)] v z = [ + ( b n + a )2 ],/:

[-lnq(A.)]

I/2._.____~

a

/.t,~ n +/.t---~-

(7)

(8)

The result implies that if the deviation is caused by the cut -off effect of the particle or crystallite size,a linear relationship will exist between [-lnq(A.)] vz and n.The lnq(A.) is the deviation value from ASF.

526 3.

APPLICATION AND DISCUSSION

The equation (8) has been used to fit the experimental FT product distributions published by several authors[6-11]. In these experiments apparent negative deviations from ASF distribution at higher carbon numbers were observed .The results are shown in fig. 1-8. Good agreements are achieved.It can be seen obviously that linear relationship exist between[-lnq(A~)] v2 and carbon number n. The equation (8) is based on the two ideas. (a) The negative deviation is due to the limitation of the particle size to carbon chain growth. (b) A simple linear relationship exists between the metal particle size and maximum molecule chain length formed on it. So,the results confirm the viewpoint that for FT synthesis on the catalysts with highly dispersion, the negative deviation is the result of cut-off effect of the small particle or crystallites size.The results also confirm the assumption that the maximum product chain length depends on the particle size linearly. We do not deny that the deviations may be caused by some secondary effects and other external reasons such as diffusion, secondary cracking, wax deposition in the catalyst bed, operation in transient conditions, etc. But, for these experimental results used here, it was claimed that the other reasons had been excluded and the improved selectivity were due to primary effect, at the level of active site.

2.0 c"

~1,o (jC

v

t

J

J

2.0

J

c-

~I.0 oc

i

0.0

0.0

carbon number n Fig. 1 The plot of [ -Inq(An)] 1/2 versus n (ref. [8]).

carbon number n Fig.2 The plot of [-Inq(An)] lr2 versus n (ref. [6]).

7

9

2.0

2.0 ~le,I

//

C

t--

$1.o

~..%.1.0

UT C

C

I

!

o.o

0.0

Fig.3

8

carbon number n The plot of [-Inq(An)] 1/2 versus n (ref. [9])

f

~

~

carbon number n Fig.4 The plot of [-Inq(An)] 1/'2 versus n (ref. [11], fig. 2, curve 4).

527 2.0

2.0

,,-.,~'1.0 1D" v

el

v~- 1.0 era

J

0.0

v

8

9

|

Y

0.0

10

12

13

carbon number n Fig.6 The plot of [ -Inq(An)] lj2

versus n (ref. [10], fig. 2, curve 3 ).

versus n (ref [11], fig. 6 ).

2.0 ~le,

11

carbon number n Fig.5 The plot of [-Inq(An)] 1'2

3,0 ~le,

i

c"

i

r"

<-<-.1.o

~EF2 . 0

i ,.,_,.,

i ,.,__,.,

c"

c

o o

1'1

1'2

1.o

9

,

10

,

11

carbon number n Fig. 7 The plot of [ -Inq(An)] lr2

carbon number n Fig. 8 The plot of [ -Inq(An)] 1'2

versus n (ref. [7], for Fe/C (10%)).

versus n (ref. [7], for Fe/C (3%)).

ACKNOWLEDGMENTS Financial support from Province Youth Nature Science Foundation of Shanxi of China is highly appreciated.

REFERENCES

(1) Anderson, R. B. ; Friedel, R. A. and Storch, H. H .J. Chem. Phys. 1951, 19, 313. (2) Vanhove, D. ; Makambo, p. and Blanchard, M. J. Chem. Soc.Chem. Commun. 1979, 605 (3) Nijs, H. H. and Jacobs, P. A. J. Catal. 1980, 65, 328 (4) McDonald, M. A. ; Storm, D. A and Boudart, M. J. Catal. 1986, 102, 386 (5) Espinoza, R. L. and Snel, R. J. Chem.Soc. chem. Commun. 1986, 1796 (6) Mitsudo, T. ; Boku, H. ; Murachi, S.;Ishihara, A. and Watanabe, Y.Chem. Lett, 1985,1463 (7) Jones, V. K. ; Neubauer, L. R. and Bartholomew, C. H. J. Phys. Chem. 1986, 90, 4832 (8) Snel, R. Catal. Lett., 1988, 1,327 (9) Vanhove, D. ; Zhang, Z. ; Makambo, L. and Blanchard, M. Appl. Catal. 1984, 9, 327-342

528 (10) (11) (12) (13) (14)

Ballivet-Tkatchenko Danielle and Tkatchenko Igor, J. Mol. Catal. 1981, 13, 1-10 Liu Fu and Bartholomew, C. H. J. Catal. 1985, 92, 376-387 Jacobs, P. A. and Van Wouwe, D, J. Mol. Catal. 1982, 17, 145-160 Yang, Y; Pen, S and zhong, B. Catal. Lett. 1992, 16, 351-357 Anderson, J. structure of Metallic catalysts, Academic Press. London, New York, San Francisco, 1975;. P. 369

91997 Elsevier Science B. V. All rights reserved. Dynamics of Surfaces and Reaction Kinetics in Heterogeneous Catalysis G.F. Froment and K.C. Waugh, editors

529

Network Simulation of Catalytic Cracking Reactions C.I.C. Pinheiro, F. Lemos, F. Ram6a Ribeiro Centro de Engenharia Biol6gica e Quimica~ Departarnento de Engenharia Quimica, Instituto Superior Trcnico, Av. Rovisco Pais, 1096 Lisboa Codex, Portugal.

In this work we present a new semi-lumped model for cracking of paraffins. The model involves simulation of the ensemble of reactions taking place during catalytic cracking of paraffins. The reaction scheme is detailed to the carbon atom number level, i.e., in this approach all molecules having the same carbon atom number are lumped together. A large number of reactions between all these species are taken into account, ranging from cracking, hydride transfer and chain growth - eventually leading to coking. The rate constants for all these reactions are described in a nine parameter system of equations that describes the way the kinetic rate constants change with the carbon atom number of both reactant and, when necessary, products. The significance of the nine rate constant parameters is discussed.

1. INTRODUCTION Modelling complex network reaction schemes is a difficult task and, for the particular case of catalytic cracking, several approaches have been used [1, 2]. The main difficulties are related both to the number of species and reactions involved and to the high degree of interrelation between all the species Catalytic cracking is a particularly complex system, because, not only its reaction scheme is very complex, but also some of the products, namely coke, interfere with the activity of the catalyst and lead to a significant reduction in its performance. Simulations covering both a description of the product distribution and the deactivation process are very tmcommon. The availability of desktop computers with very high computation power can, however, change the way we deal with this kind of reactions in the near future. In fact, it is now possible to perform routine simulations of very complex systems in a desktop computer and this will certainly bring new fight to the processes being studied and will allow more complex theories to be tested against experimental results. There is, however, another difficulty that arises in complex network reaction schemes and that can not be lifted by increased computational power. The high number of reactions involved usually requires a high number of kinetic parameters to be determined. It may eventually be necessary to reduce the number of parameters in order to obtain feasible interpretations of the experimental results.

530 2. M O D E L DESCRIPTION

The model we describe in this work has the objective of constituting a preliminary study for the development of models that are able to explain, at the same time, both the product distribution and the decay of activity that is observed in the catalytic cracking reactions; the description of the decay of activity implies also that on is able to estimate the amount of coke deposited on the surface of the catalyst. To do this, the model is developed so as to describe the evolution with time of the partial pressures of the relevant species in gas phase, as well as their surface concentrations, when adsorbed onto acid sites. Several simplifications are made, to ensure that a feasible solution is accessible. These simplifications can be neglected in other developments of this model. The main simplifications that are used are: 9 All the active sites are considered to be of the same strength. 9 The species are lumped together according to their number of carbon atoms, irrespective of being paraffins, olefins or aromatics. 9 The reaction rates for each kind of reaction depend only on the number of carbon atoms of the species involved as reactants and products. 9 According to the second assumption, no reactions involving hydrogen transfer were considered. 9 The maximum number o carbon atoms considered was taken to be 21; this figure although arbitrary, is similar to the number of carbon atoms observed in coke species deposited inside Y zeolites during the transformation of n-heptane [3]. In fact, the amount of compounds with 21 (or more) carbon atoms was taken as the coke content of the catalyst. Three sets of reactions were taken into consideration, embracing the possible reaction types when the above mentioned simplifications are taken into account: 9 Chain transfer reactions - whereby a gas phase molecule exchanges place with an adsorbed carbocation;

g On

+

a C m

. . . . . . . t~.

g C m

+

a C n

9 Chain reduction reactions - these are the cracking reactions, where an adsorbed carbocation is split into a smaller carbocation, still adsorbed on the surface site, and a gas phase molecule with the remaining number of carbon atoms; this reaction was only considered for species with more than six carbon atoms;

531 a

a

CI i

r-

9 Chain growth reactions - the reverse of the cracking reactions and is considered to involve the attack of an adsorbed carbocation to a gas phase molecule, leading to an adsorbed carbocation with a molecular weight which is the sum of those of the two reactant species;

+

c,

a _ _ ~

C a .

.2

It us further considered that all the species present could participate in all of the above reactions, creating a complex network where all the species interacted with all the others by means of the three sets of reactions considered. Since we considered the existence, both in gasphase and on the surface, of species with a number o carbon atoms ranging from 1 to 21, the chain transfer reactions alone represent over 400 individual reactions being considered. Similar numbers exist for the other types of reactions. All these reactions are taken to proceed with simple kinetics. Since the number of individual reactions possible is so great, where the rate constants were considered to be a function of the number of carbon atoms of the species involved. For these kinetic rate constants simple, albeit arbitrary equations, were used to compute them for each set of reactant/products considered. Each expression was parametrised with three parameters. Below are given the expressions that were used for each of the reaction types that was considered. 2.1 Chain Transfer

For the chain transfer reactions we considered that the transfer would be faster as the weight of the adsorbed molecule decreased and that of the gas phase molecule increased. Accordingly, the following expression was used:

kt(i,j) = kto

exp(-ktl

i)tanh(kt2J )

The parameter k~0 represents the intensity of the chain transfer reactions, increasing it will result in a general increase of the chain transfer reactions rate. The parameter ktl represents the way the reaction rate decreases with the increase of the number of carbon atoms of the adsorbed molecule, while kt2 represents the way the rate changes when the number of carbon atoms of the gas phase reactant increases. 2.2 Chain R e d u c t i o n

The expression for the cracking reactions was taken to depend on three parameters so as to observe the following general considerations: 9 The cracking rate must increase as the number of carbon atoms of the reactant increases;

532 9 The maximum cracking rate for each species corresponds to the situation where the molecule is split in half The expression used was

k:(i,j) : kco exp[- kcl(j-

i/2) 2 ]{1 +

tanh[kc2(i-

7)]}

which reflects these considerations, kco is the parameter related to the intensity of the cracking reactions, while the term associated with kcl reflects the fact that the maximum cracking rate occurs for a splitting in half of the molecule, with a more or less accentuated decrease when we deviate from this situation, and the parameter kc: reflects the way how the cracking rate increases with the increase of the reactant size. In this case the hyperbolic tangent was used to prevent an indiscriminate increase of the rate for heavy reactants.

2.3 Chain Growth The expression for the chain growth kinetic rate constant was used as

kr(i, j) = kro exp(-krl j)tanh(kr2 i) This for was chosen so that we could control the intensity of the chain growth reactions (k~0). It was further considered that the chain growth reactions should be easier if the gas phase molecule was smaller (adjustable with krl) and the adsorbed molecule was heavier (adjustable with kr2).

2.4 Model Implementation This kinetic model was applied to experiments carried out in a continuous micro-reactor. The transient material balances for all the species considered, 21 gas-phase species and 21 adsorbed ones, were written and solved using a 4th-order Runge-Kutte scheme. From the partial pressures and surface concentrations that were computed the expected product distributions at the exit of the reactor as well as the coke content at the surface and the global activities, can be obtained with this model. Global product distributions of both gas phase and surface adsorbed species, can be computed as a function of time.

3. E X P E R I M E N T A L RESULTS AND DISCUSSION The model was run to simulate the transformation of n-heptane over a Y zeolite. Several sets of parameters were tried till a set resulted in a simulation which fairly coincides with the observed experimental results. The experimental results, as well as one of the simulations that were made, are presented in figures 1 and 2. As it can be seen, the main features that are observed in the actual experiment are obtained. The activity increases sharply from during a very short time interval, corresponding to the filling of the reactor volume and the adsorption of the first species on the surface of the catalyst. After this induction period, which is very short and depends mainly on the rate of the

533 chain transfer reactions, the activity decreases steadily with time, in a fashion similar to the one observed in the experimental data, although the activity decreases to zero, while the experimental one retains a residual activity after the decay process. This discrepancy can result simply from the fact that we did not consider any other form of transformation of the reactant, except the acid catalysed transformation over the surface sites; this does not correspond to reality and some transformation at the external surface of the catalyst, over the coke itself or in gas phase is expected. 25 = 20 @

.i..i

15

i\,

= @ 10

g

5

9

|

I

110

20

30

40

50

60

t (min)

Figure l. Comparison between simulated (--) and experimental (1) profiles of n-heptane conversion as a function of time on stream.

As for the product distributions, it can also be seen that the results are in good accordance with the experimental results. The asymmetry between the amounts of C3 and C4 produced and that is observed in the experimental results for the initial times is also observed in the simulated ones. This asymmetry is reduced as the deactivation proceeds, both in the experimental and in the simulated results. 50

50

t = 5 min

t., 30

k 30

~q o

o 20

zo 10

10 0

t = 60 min

40

40

1

3 5 7 9 11 13 Carbon atom number

1 3 5 7 9 l l 13 Carbon atom number

Figure 2. Comparison between simulated (m) and experimental ( ) product distributions for two different times on stream.

534 The simulated results show a slight tendency to favour heavier products, a fact which still may be corrected by adequate changes in the kc2 parameter.

4. CONCLUSIONS From the results presented above it is clear that the network model presented in this paper is able to reproduce most of the characteristic features that are observed in these reactions. The increased computational power of modem desktop computers will allow the addressing of these complex mechanisms with increasing detail, allowing a new insight into processes which involve an enormous amount of possible reactions and species.

5. REFERENCES

1. W. Feng, E. Vynckier, G.F. Froment, Ind. Eng. Chem. Res., 32 (1993) 2997. 2. R.V. Shendye, R.A. Rajadyaksha, Chem. Eng. Sci., 47(3) (1992) 661. 3. J.M. Lopes, F. Lemos, F. Ram6a Ribeiro, E. Derouane, P. Magnoux, M. Guisnet, Appl. Catal. A, 114 (1994) 107.

91997 Elsevier Science B. V. All rights reserved. Dynamics of Surfaces and Reaction Kinetics in Heterogeneous Catalysis G.F. Froment and K.C. Waugh, editors

535

Non-Catalytic Carbon Gasification Modelling I. Santos Silva~, C. Palma ~, F. Lemos b, F. Ram6a Ribeiro b, J. Sousa Lobo a aDepartamento de Quimica, Faculdade de Ci6ncias e Tecnologia, Universidade Nova de Lisboa, Quinta da Torte, 2825 Monte da Caparica, Portugal. bCentro de Engenharia Biol6gica e Quffnica, Departamento de Engenharia Quimica, Instituto Superior Trcnico, Av. Rovisco Pais, 1096 Lisboa Codex, Portugal.

In this work we present a simple model for non-catalytic carbon gasification, which can explain the induction periods that are often observed in experimental results, by accounting for the possibility of generating new active sites as the result of the gasification process of other active sites as well as for the deactivation, by a generic deactivation step.

I. INTRODUCTION Carbon gasification is a very complex process involving the transformation of a solid reactant into gaseous products by means of a solid/gas reaction. These reactions involve significant changes in the carbon material, no only at a chemical level, but also in the porosity and texture of the material. Modelling of carbon gasification has been performed by a variety of methods [ 1, 2], but the adequate description of all the characteristics of this process is not usually achieved, in particular in the respect to the existence of induction periods, that are commonly observed, namely for gasifications performed at lower temperatures, and the decrease in gasification rates for high conversions. It is clear that when the gasification process occurs, the reactivity is strictly restricted to the carbon that is located at the surface of the particles, and that the gasification of the surface carbon exposes additional carbon species that can eventually be gasified. However, it is also possible that the gasification of the more active carbon sites can also induce an increase in reactivity of other sites. The fact that surface sites can present widely different reactivities is well accepted and is directly observed by the differences in reactivity that can be obtained with carbons having a similar surface area.

536

2. THEORY In the model described below we consider the existence of two main types of carbon sites: potentially active and inactive sites. Inactive sites have no gasification ability and may be considered to include the ashes of the carbon material. Although the ashes sometimes present some catalytic activity in relation to the gasification process, we will not consider any catalytic activity in this model, since we are applying it to an ashless carbon, in this context, then, the inactive material should be regarded only as the carbon material which is difficult to gasify. The potentially active sites correspond to the effectively gasifiable carbon material, and will comprise the majority of the mass of the carbon material. Nevertheless, for the reasons presented above, we will consider that only a part of them, which may be associated to the ones at the surface of the particles, are readily active for gasification on the fresh sample. The readily active sites can be gasified, and in this process they may generate a new set of active sites, either because they have uncovered new sites that laid beneath them, or because their gasification provided a new access to sites having a low activity, rendering them more active. We will consider, then, that for each carbon site that is gasified, a certain number of potentially active sites is converted into readily active sites; of course, this generation depends upon the availability of potentially active carbon. We will also consider the poss~ility of the transformation of the active sites into inactive ones, by some chemical process that is capable of decreasing their reactivity. The above-mentioned set of assumptions can be translated into the following equations, describing the way the amount of each of the types of species evolve with time: dmA

_

- k v mA + k m (kv mA )(mAm - mA ) - k i mA

dt

(1)

dmAm

_ - kv m A

dt dml dt

_ ki

mA

- ki m A

(2)

(3)

where mA represents the amount of readily active carbon (in mass), mam, the amount of potentially active carbon (which also includes the readily active one), and m/the amount of inactive carbon, kv, ki and k,, are, respectively, the gasification and inactivation rate constants, and the average number of sites that is generated by the gasification of one site, for a fresh sample. This model can be fitted to experimental data, by numerically integrating equation 1 through 3, with the initial condition that for t = 0, the amount of sites of each kind is mAo, mAmo and mio, respectively for the readily active, potentially active and inactive sites. Both the set of kinetic constants, kv, ki and km, and the initial condition parameters, mAO, mamo and mzo, are taken as fitting parameters for each particular experiment. This model can equally be applied to isothermal and to TPR data, provided that a suitable form of variation of the various kinetic constants with temperature is used. In our case, we assumed that both kv and ki followed an Arrhenius expression and k,, remained constant.

537 3. E X P E R I M E N T A L RESULTS AND DISCUSSION Carbon gasifications were carried out in a C.I. Electronics MK II microbalance with continuous recording of change in weight. All the system was purged with a stream of nitrogen before heating to the desired temperature (in constant temperature gasifications). After the temperature was stabilised, the nitrogen was replaced by air and the weight of the sample was recorded as a function of time. A constant air flow of 4 cm3/s was maintained. In Temperature Programmed Reaction (TPR) experiments, the sample was heated in a flow of air from the beginning of the run, at a linear rate of approximately 10 ~ A delay of 24 s between opening the switch to change the gas flow from nitrogen to air and the effective arrival of air to the reaction chamber was computed and taken into account in the interpretation of the results. In this study a sample of BDH 33033 carbon was used. The fitting of the experimental data to the model described above was carried out in a commercial worksheet programme (Excel 5 - 9 Microsoft) using a least-squares method. The equations were integrated using the Euler method with a suitable time step, and the sum of the squares of the residuals for all data points in an experiment was minimised using the Solver tool in the software described above.

1 _ 0,9 ~ ' ~

"

0,8 0,6 -~u

E 0,5 -?in ' EO,4+ 0,3 ~-

.~

0,2 i n ~

o,1

,

0

-

~

i

200

400 t/s

-

600

-

800

Figure 1. Comparison of experimental ( , - 625 K, 9 - 675 K, J, - 700 K, 9 - 725 K) and simulated (--) data for the isothermal gasification of BDH 33033 carbon.

Figure 1 shows the experimental results, and corresponding fittings to equations 1 through 3, for isothermal gasification experiments. The parameters obtained in the fitting are given in table 1. As it can be seen, the fittings are quite good through-out the whole curve for all the temperatures that were tested, fitting both the gradual increase in the gasification rate that is observed for low temperature gasifications at the beginning of the run, and the slow decrease

538

in rate at the end of the rtm. Nevertheless, for the char trader study it was verified that the inactivation process did not occur and m1 and k~ were always fitted to zero (see table 1).

Table 1 Values obtained from the fitting of equations 1 through 3 to the experimental data for the gasification of BDH 33033 carbon under isothermal and temperature programmed reaction conditions.

k~

k,

k,,,

m.4/mA,,

mi /m,4,,,

625

9.0 • 10-3

0

5.0

7.4 x 10-3

0

675*

2.6x 10-2

0

6.0

2.6x 10-2

0

700*

3.7 x 10.2

0

4.9

5.1x 10.2

0

725

6.0 x 10.2

0

6.1

7.0x 10.2

0

-

0

2.7

1.6x 10.2

0

Temperature

TPR (10 K/min)

* Values presented are the average of two distinct measurements.

As it can be seen, the values of k,, are approximately independent of the temperature and are probably linked to more structural properties of the carbon material. Both the gasification rate constant and the fraction of readily active sites increase with temperature. The increase of the fraction of readily active sites may indicate that the surface of the coal is highly heterogeneous in reactivity, as one would, in fact, expect, and that, as the temperature is increased, the number of sites that can be readily gasified increases, due to the availability of sites with a lower reactivity.

Table 2 Arrhenius parameters obtained from isothermal and TPR gasification experiments.

kvo

Ea / kcal/mol

Isothermal

7.7 x 105

32.5

TPR (10 K/min)

3.8 x 105

32.5

From the values of kv as a function of the temperature, the activation energy was computed, by fitting the data to an Arrhenius equation. A nice plot was obtained leading to an activation energy for the gasification process around 32 kcal/mol. The results can be seen in table 2 and in the Arrhenius plot in figure 2.

539

-2

7

-2.5

i

-3 -3.5 -4 -4.5 T I

-5

'

0.001

0.00115

0.0011

0.00105

1 / T / K "1 Figure 2. Arrhenius plot for the values of kv as a function of temperature. Experimental ( , ) and fitted line (--).

The same model was also applied to temperature programmed reaction data. In this case only k~. was considered to vary with temperature, according to an Arrhenius law, and the preexponential factor and the activation energy were directly estimated by the fitting, along with km and the initial amounts of each kind of carbon species; The value of kz was taken as zero from the start, in view of the results obtained in the isothermal gasification experiments. Both experimental data and the corresponding fitting are depicted in figure 3. The relevant model parameters are also given in tables 1 and 2.

1.4 T 1.2 1 0.8 0.6 0.4 0.2 0 j 0

~ 500

1000

1500

2000

2500

t/s

Figure 3. Comparison of experimental ( , ) and simulated (m) data for the temperature programmed gasification of BDH 33033 carbon.

540

It can be seen that the fitting is also very good and that the activation energy value that is obtained is entirely coincident with the one obtained from the Arrhenius plot for the kv values from the isothermal gasification experiments. However, the pre-exponential factor, k,, and mA/mA,, parameters present significant differences. These differences may be related to the different progression of the gasification process in the two different cases. Gasification through isothermal or temperature programmed conditions will certainly differ in the way they use-up the various sites that are available. In fact, the temperature programmed reaction experiment shows a k,, value much lower than the ones obtained in the isothermal experiments, indicating perhaps the large importance that the initial rates of gasification have on the estimation of this parameter, in connection with the fact that in these experiments the gasification process starts very slowly at rather low temperatures. The amount of sites readily active has a value that falls within the range of those obtained for the isothermal experiments. Although its value has a very clear meaning in the isothermal experiments, the amount of carbon that can be gasified immediately at that temperature, due to the heterogeneous nature of the char surface, its meaning is not so clear in a TPR experiment, where the gradual increase in temperature must also lead, not only to the increase of the rates of the reactions, but also to the increase of available sites at the surface due to this increased reactivity imparted by the raise in temperature. Several remarks can be made regarding the results presented above. 9 The model is able to fit quite well all the experimental results that were obtained, including the ones in temperature programmed gasification 9 Application of the model to isothermal and TPR data supplied activation energy values coinciding within 1 kcal/mol. 9 The fraction of carbon material that is readily available for gasification increases markedly with the temperature; a fact which is consistent with the known heterogeneity of the carbon materials; at larger temperatures, more material will be readily available to gasification. 9 The induction period can be explained by the chain reaction-type mechanism, since the value of km is always significantly higher than 1, indicating that, namely at the beginning of the run, for each site that is gasified, several sites may become available. As a conclusion, we think that the use of this type of models for carbon gasification can provide a significant insight to the way gasification of carbon occurs and, eventually, lead to new mechanisms for the quantitative characterisation of carbon reactivities.

4. REFERENCES

1. N. Yasyerli, T. Dogu, G. Dogu, I. Ar, Chem. Eng. Sci., 51(11) (1996) 2523. 2. A.A. Lizzio, H. Jiang, L.R. Radovic, Carbon, 28(1) (1990) 7.

91997 Elsevier Science B.V. All rights reserved. Dynamics of Surfaces and Reaction Kinetics in Heterogeneous Catalysis G.F. Froment and K.C. Waugh, editors

541

Effect of solubility parameter on the MTBE synthesis kinetics C. Fit6, J. Tejero, M. Iborra, F. Cunill and J.F. Izquierdo Department of Chemical Engineering. Universiw of Barcelona. Marti i Franqu6s 1. E-08028 Barcelona. Spain

1. I N T R O D U C T I O N Because of the environmental impact of the gasoline, it has been necessary to reduce and/or substitute some of its components. MTBE is one of the preferred compounds to add in order to fulfill the specifications of reformulated gasolines without reducing its quality. The main advantages of MTBE are very good anti-knocking behaviour, high octane number and low Reid vapor pressure (Rvp). Finally, it allows to introduce oxygene into the gasolines, in such a way that its use reduces the unburned hydrocarbons, volatile organic compounds and carbon monoxide emissions. Industrially, MTBE is synthesized from methanol and C4 cuts, mainly containing isobutene, 1-butene. n-butane, cis-, trans-2-butene and isobutane. The involved reaction is the addition of methanol (MeOH) to isobutene (IB) to form MTBE. Both reactants are relatively cheap and the reaction can be carried out under mild conditions (313-363 K and 1.6 MPa). Reactor feed contains IB and MeOH in a molar ratio IB:MeOH 1:1.1 to minimize the extent of side-reactions and to achieve vet? ~high IB conversions, because chemical equilibrium is shifted to the product formation. A sulfonic macroporous ion exchange resin with a polystyrene-divinylbenzene matrix is used as the catalyst. The physical properties of the reaction medium highly influence the resin structure. In very polar media the resin presents a swollen state, due to the permeation of polar molecules into the resin backbone. In such a way, most of active centers can be accessible to the reactants. However, in this polar media the sulfonic groups would be dissociated and the solvated protons would be the catalytic species. Thus, inside the porous the catalytic process would become a pseudo-homogeneous one. In a non-polar medium, the sulfonic groups form a network of hydrogen bonds that collapses the resin structure avoiding the inner active centers to become accessible to the reactants. The reaction in such media would involve the protons bonded to the sulfonic groups in a pseudo-heterogeneous catalysis, that generally implies a higher reaction rate. The most widely used kinetic equation for the MTBE synthesis reaction is probably that proposed by Rehfinger and Hoffmann [1 ]. It was obtained from experiments made in similar conditions than those of the present work, but in the use of inerts (1-butene and/or n-butane) as

542 solvents, that were introduced into the system to simulate C4 cuts obtained from FCC units in the refinery industry. The reaction was carried out in a continuous stirred tank reactor and rate data were obtained at the stationary state. Mixtures in the absence of product were fed to study the forward reaction; pure MTBE was fed to determine the kinetics of the decomposition. In that work, the kinetic equation obtained for the whole range of compositions is based on a LHHW mechanism in which IB and MeOH, both adsorbed on the resin, react to form MTBE. The rate-controlling step is surface reaction. As a simplification in the activities calculation, the authors consider that the reaction mixture is composed by three compounds: methanol. MTBE and a C4 pseudo-compound that includes all hydrocarbons. In the absence of product, the proposed kinetic equation is

r = k'

aiB

(1)

a MeOH

where k' and ai are, respectively, the apparent rate constant and the liquid phase activity of compound i. As mentioned before, the catalytic behaviour of the resin strongly depends on the nature of the medium. In order to take into account which is the influence of the medium in the reaction kinetics, experiments should be carried out in conditions wherein the medium polarity varies significantly. It can be done performing experiments in the absence of diluents.

2. EXPERIMENTAL SETUP Reaction rates were measured in the absence of MTBE in a thermostated packed-bed reactor. The reactor was fed with pure isobutene (IB) and methanol (MeOH), with a molar ratio IB/MeOH (Rim) between 0.5 and 2. It was operated isothermically in a differential regime, in the absence of mass transfer control. The experiments were carried out at four different temperatures in the range 318-363 K, and the pressure was kept at 1.6 MPa to assure that all the compounds involved in the reaction are in the liquid state. The sulfonic macroporous resin with a styrene-divinylbenzene matrix Bayer K2631 was used as the catalyst. The reactor inlet and outlet were analyzed by a gas chromato~aph with a FID detector. Reaction rates were determined from these compositions at the steady state.

3. RESULTS AND DISCUSSION Our rate data are of the same order of magnitude as those of Rehfinger and Hoffmann [1 ]. Experimental values obtained at each temperature were fitted to Eq.1. Depending on the composition of the reaction medium, two different trends were observed: i) At RI/A > 1 calculated values of reaction rate were lower than the experimental ones. ii) At RVA < 1, calculated values were higher than the experimental ones.

543 In addition, at each temperature the observed change in the reaction rate was about of 45%, while the predicted change by Eq. 1, namely the change in the quotient a~/aMeoH, was only of about 6%. This fact suggests that Eq.1 explain just approximately experimental data, as Fig.1 shows. At the other temperatures the same behaviour was observed. As Eq.1 cannot predict reaction rates accurately, a reformulation of the kinetic equation should be undertaken.

0.4

0.3

o

E

0.2

0.1

0.1

0.2

0.3

0.4

re• (mol/kg s) Figure 1 Experimental (rexp) and calculated (rcalc) reaction rates from Eq. 1 at 60~

The main differences between both experimental procedures are: i) The reaction system, which does not influence the intrinsic kinetic equation. ii) Rehfinger and Hoffmann used Amberlyst 15 and a self-prepared similar one as catalysts, but the authors indicate that the only property that influences the reaction rate is the resin capacity. iii)The absence of inerts compounds in the present work. So, the main reason for the observed inaccuracy is assumed to be structural changes in the resin due to the medium properties. MeOH and IB are very differents as regards to their polarity. In the absence of diluting compounds, the reaction medium shifts from strongly polar when MeOH is in excess to nearly apolar in excess of IB. The presence of inert compounds similar to IB reduces the MeOH contribution to the medium polarity in such a way that its influence cannot be seen.

544 Qualitatively, accessibility, to active centers has been related to the medium solubility. parameter, 8m, and reaction rates depend on the number of accessible active centers [2]. The near the ~m value to that of the resin solubility parameter (Sp) is, the higher the affinity resinmedium is. In the present syste m, the solubility parameter value of an IB-MeOH medium increases with the methanol content. In example, at 60~ the 8 values for IB and MeOH are, respectively, 12.04 10-3 and 27.78 10-3 (j/m3) 1/2. In this case, the reaction mixture permeates easily into the resin porous and there are more accessible active centers. As an opposite effect, when 8m and 8p are similar, the medium is highly polar, and the kinetic mechanism tends to a pseudo-homogeneous one. that usually shows a lower reaction rate. The modification of the kinetic equation should take into account both effects. To verify the relationship between ~m and the lack of fit of Eq. 1, weighted residuals were plotted versus ~m (Fig.2). It can be seen that the effect of 8m is not random. Therefore, the solubility parameters should be included in the kinetic equation. To do it, the rate constant has been considered as a product of two factors: the intrinsic rate constant, that only depends on the temperature, and the accessibility and activity of the active centers, that depend on the medium properties and its interaction with the resin. So, the rate equation should be

r

=

k'

f((~m)

(2)

am a MeOH

0.25 0.15 "N 0.05 I1) .,.a

$

-0.05

,,-, (D

-0.15 -0.25 14

16

18

20

~m 10"3 (j/m3) 1/2

Figure 2: Dependence of weighted residuals from Eq.1 on the medium solubility parameter (Sm) at 60~

545 Fig.2 shows that at low am values, experimental rates were higher than calculated ones. Opposite trend is observed at high ~rn values. Therefore, in Eq.2 the function fshould exhibit an inverse relationship with 8m. The empyrical function f(6"m) = 1 / ~m improves significantly the fit of r to the experimental data and the bias observed in Fig.1 and 2 is strongly reduced. Consequently, the searched function f(c~m) with a physico-chemical meaning must be similar to the function 1/Sin. Similarly to the expression for the activity coefficient of a nonelectrolyte solute dissolved in a solvent [3], it has been considered the following expression:

r =kexp

f-

R

T

P

,t

(3) aM~on

B

where k is a global rate constant, V m the molar volume of liquid, am its solubility parameter, ~p the porosity of the resin, and 8p its solubility parameter. The plot of the exponential factor vs am, by considering the mean experimental value for Vm, the resin porosity supplied by the manufacturer (~p = 0.5) and the resin solubility parameter estimated by group contribution methods (Sp = 22.73 10-3 (j/m3) 1/2) [4], presents the same form that the corresponding for the function I/Sin. It is worth noting that the exponential factor does not depend on the absolute value of 8m, but on the difference between the reaction medium and the resin solubility parameter: it equals to unity only when ~m = ~p and it is greater than 1 otherwise. The inclusion of flSm) in the kinetic equation implies the introduction of one new parameter to fit, related to the resin solubility parameter, at each temperature. It has been observed a smooth lineal dependence of am with the temperature. By analogy, 8p should show also this kind of temperature dependence. At each temperature, the fit of Eq.3 implies a 65-fold average reduction in the sum of squares of weighted residuals with regard to those obtained with Eq.1. This fact is not only imputable to the addition of one parameter in the equation to fit, but to the improved kinetic equation. Fig.3 shows the calculated reaction rates at 60~ with Eq.3. As it can be seen, the bias of the residuals has been drastically reduced. Identical behaviour has been observed at all temperatures. The temperature dependence of rate constant yields an apparent activation energy of the reaction of (81.8 + 1.69) kJ mol -~, that agrees fairly well with values quoted in literature [5]. In addition, the smooth lineal temperature dependence of 8p has also been observed, with a mean value of 22.45103(j/m3) 1/2, very similar to the previous estimated value 22.73 103 (j/m3) 1/2. It is worth noting that no restrictions were imposed in the value in the fitting of the parameter related to 8p.

546 0.4

0.3

~ o.2

g~

0.1

0.1

0.2

0.3

0.4

rexp(mol/kg s) Figure 3" Experimental (rexp) and calculated (rcalc) reaction rates from Eq.3 at 60~

4. CONCLUSIONS Experimental rate data for the MTBE synthesis reaction were determined from IB-MeOH mixtures with a molar ratio RVA between 0.5 and 2 in the temperature range of 318-363 K. These values could not been explained by the probably most used kinetic equation for that reaction [ 1]. It has been attributed to the fact that in the present work no diluents were used and the variation in the medium polarity influences significantly reaction kinetics. A term concerning the medium properties must be added in the kinetic equation. Taking into account the termflSm) (Eq.3) strongly improves the fit to the experimental data, reducing drastically the sum of squares. Values of obtained apparent activation energy and resin solubility parameter agree fairly well with those published in the literature and estimated by group contributions methods, respectively.

REFERENCES

1. 2. 3. 4. 5.

A. Rehfinger and U. Hoffmann, Chem. Eng. Sci. 45 (6) (1990) 1605-1617 C. Buttersack, React. Polym_ 10 (2-3) (1989) 143-164 C. Reichardt, 'Solvents and Solvent Effects in Organic Chemistry', VCH, Weinheim, 1988 S. Sourirajan and T. Matsuura, 'Reverse Osmosis / Ultrafiltration. Process Principles', National Research Council Canada, Ottawa, 1983 J. Tejero, F. Cunill, J.F. Izquierdo, M. Iborra, C. Fit6 and D. Parra, Appl. Catal. A 134 (1996) 21-36

91997 Elsevier Science B.V. All rights reserved. Dynamics of Surfaces and Reaction Kinetics in Heterogeneous Catalysis G.F. Froment and K.C. Waugh, editors

547

H y d r o g e n s p i l l o v e r effect over t h e oxide s u r f a c e s in s u p p o r t e d nickel catalysts

V. Almasan~,T. Gaeumann b, Mihaela Lazar ~, P. Marginean ~, N. Aldea ~ ~Institute of Isotopic and Molecular Technology P.O. Box 700, 3400 Cluj-Napoca Romania bEcole Politechnique Federale Lausanne, CH- 1015 Lausanne, Switzerland

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

The aim of this paper is to determine the quantity of chemisorbed hydrogen on nickel catalysts and in hydroxyl groups on the metallic and support surfaces, respectively [1]. The proposed method is the isotopic exchange reaction between hydrogen stream and catalyst surface. The hydrogen from hydroxyl group is isotopically exchanged through spillover phenomenon. Four supported nickel catalysts:

Ni/Cr203, Ni/Si02, Ni/MgO and Ni/Zr02 were

investigated by Isotopic Transient Kinetics Method [2]. The catalyst samples were prepared by coprecipitation method. B.E.T. surface area and metal dispersion were measured by krypton adsorption and hydrogen chemisorption, respectively [3, 4]. The main measured values of the metallic surface area and the total surface area of the catalysts are reported in Table 1. Table 1: The surfaces areas measured for the nickel catalysts

(rn2/9)

Catalyst

Ni surface area

B.E.T. surface area

Ni/Cr203 Ni/Zr02 Ni/Si02 Ni/MgO

29.7

160

19.4

85

61.9

194

41.4

98

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

A typical reaction system for isotopic exchange H/D studies was employed. It consists of a gas supply unit and a differential-bed-plug-flow reactor (PFR) followed by a

548 quadrupole mass spectrometer(Fig.l)[2, 5]. I

GAS SUPPLY UNIT

QMS

_1 CATALYTIC REACTOR

I Figure 1" Experimental set-up The input feed stream is composed of an effluent gas (argon) and one of the hydrogen isotopes, H or D, at a time. The output species from reactor were monitored by the mass spectrometer, with a constant inner pressure of 6 x 10.7 torr. The gases used in all experiments are of ultra high purity (min 99.99 % vol.). 100

I

I

I

IS]'" [.-~. . ~ - ~ .

90 80 -

EIEI"szD.o.. cr

.'~"

. ['7]. . [~] " " [ ~ "

" [~"

.["7]..

"

"e H2 "@.HD+.D2 .El..

70 60 Rel. units

. ~'~.

50 9

40-

!

o9

9

..-+.

o-i 2~ 4

10L: 0

+:%.

4 .~ ~v~

§

[ r~" [ .........

0

5

. I

I

10 time [min]

15

20

Figure 2: MS diagram for HAD experiment on Ni/Zr02 at 250~ The catalyst sample was initially reactivated the 10 hours in situ, in a hydrogen stream at 350~

The temperature was afterwards reduced to the desired value 9 The system was

isothermally operated at seven different temperatures ranging between 21~

and 300~

When the reactor temperature was stable the hydrogen stream was changed to argon (150 cm3/min), which washed the catalyst surface until the mass spectrometer did not detect

549 any signal for mass 2 (H2). At this moment D2 (flow rate 3 cm3/min) was injected in the effluent gas (Ar). Simultaneously the MS diagrams corresponding for the 2, 3, 4 masses, were recorded, with the scanning speed of one mass unity in 3 seconds. The experiments were coded with the initial letter of gases(e.g,

the experiment is

coded HAD when the passing gasses are hydrogen - argon - deuterium). When all catalyst surface was covered with deuterium, the D2 pulse was replaced with H2 pulse, and the similar mass spectrometer diagrams were recorded (the DH experiment). The time origin was taken as the moment when the last gas was introduced. Same ways were performed the other two symmetrical experiments coded HD and DAH, respectively [6]. A typical MS diagram is presented in fig. 2. 3. R E S U L T S

AND DISCUSSION

From the transient kinetic analysis [8] we obtained the kinetic parameters and the quantity of H2 (D2) which spilt over from the Ni surface to the oxide surface.

20 18

I

14 12 H2

I

-

-

[cm~]

I

21 o C

100~ ~ 2000 C + ' " ~/7 300 o C .rn.. ~ /

C ~

10 8 -

~/

9

.,. +. +. +. +. +- -k-. •

~:~:~: ~. ~. ~. ~. ~.

,

D. ~.

,

+- ,+

~.

o

§247247247247247247247 ~.o.

~.

a

6 4

2 0 0

I 5

i 10 time [min]

i 15

Figure 3 Kinetic desorption curves of strongly adsorbed hydrogen on

20

Ni/Zr02

in the

HAD experiment The phenomenon of isotopic exchange between the spilt over species from metal and

550 the OH (OD) removed from the oxide surface is based on the reaction:

D(H) + OH(OD) - H(D) + OD(OH)

(1)

which takes place on the oxide surface. The MS diagrams (fig.2) were transformed in kinetic desorption curves for strongly adsorbed hydrogen. An example is shown in fig.3. The plateau values of the kinetic curves are considered the maximum values of H2(D2) amount which result from isotopical exchange reaction between spilt over species and the strong chemisorbed hydrogen isotope on the oxide surfaces. The amounts of hydrogen isotopes desorbed from lm 2 of oxide surface for Ni/Zr02 catalyst are illustrated in table 2. A strong variation of the number of OH (OD) exists as a function of the temperature and of the nature of the support, as it can be noticed in Table 3. For each experiment considered in Table 2, the number of the OH (OD) groups was obtained by taking into account the averaged volumes of the hydrogen isotopes desorbed from 1 m 2 of support. The high variation from 37 % for Ni/Cr203 up to 60 % for Ni/Zr02 of the desorbed quantities of the hydrogen isotopes when the temperature is decreased can be explained by the fact that more hydrogen remas strongly adsorbed on the surface at low temperatures. For this reason two hypotheses are suggested: Table 2" The amounts of hydrogen isotopes desorbed from lrn 2 of oxide surface in the case of Ni/Zr02 catalyst (cm 3) Temperature

D AH

DH

HAD

HD

Aver age

21

0.31

0.35

0.28

0.32

50

0.29

0.33

0.28

0.29

100

0.29

0.25

0.24

0.24

0.25

150

0.24

0.24

0.24

0.23

0.24

200

0.23

0.17

0.14

0.17

0.17

250

0.14

0.15

0.19

0.15

300

0.13

-

0.16

0.14

(oc)

0.12 0.12

0.32 0.30

9 The first hypothesis is based on the supposition that the number of OH from the oxide surface remains unchanged with temperature and the metal surface retains some amount of strongly adsorbed hydrogen.

551 Table 3: The average values (cm 3) of hydrogen strongly adsorbed from 1 m 2 of support

(VH2) and (~

(NON) Ni/Cr203 Ni/Zr02 Ni/Si02 Ni/MgO vm NoH VH~ -~oH ~)~ i\bH Vm -~oH

21

0.178

9.57

0.315

16.94

0.183

9.84

0.372

20.00

50

0.178

9.57

0.298

16.03

0.178

9.57

0.353

18.98

100

0.154

8.28

0.250

13.44

0.152

8.17

0.304

16.35

150

0.155

8.33

0.233

12.53

0.118

6.34

0.273

14.68

200

0.138

7.42

0.179

9.63

0.095

5.11

0.208

11.18

250

0.123

6.61

0.150

8.06

0.091

4.89

0.202

10.86

300

0.112

6.02

0.134

7.21

0.095

5.11

0.203

10.91

the corresponding number of hydroxils groups

Temperature

9 The second hypothesis assumes that spilt over hydrogen creates new sites of strongly adsorbed of hydrogen on the oxide surfaces. The difference between the amount of the adsorbed and desorbed hydrogen on Ni at 300~

and 21~

is within the error limits is and therefore cannot explain the large

variation of the amount of hydrogen isotopes desorbed during the experiments DAH, DH, HD, HAD. Therefore the second hypothesis is more likely to be true, the strongly adsorbed hydrogen is retained on the oxide surface and not on the metal surface.

It seems to be

reasonable to assume that the mobile bond between adsorbed hydrogen and nickel surface is maintained on the oxide surface, too. It is very likely that spilt over hydrogen forms adsorption bonds with oxygen surface from the oxide, up to the individual OH formation. 4. C O N C L U S I O N S The transient kinetic analysis of the isotopic exchange H/D is a direct and useful method to investigate the transformation of the catalytic surfaces under different reaction conditions. These experiments prove that there is a strong variation among the OH population over the oxides in the hydrogen stream as a function of the temperature and of the nature of the oxidic support of Ni catalysts.

552 REFERENCES

1. Z. Paal and P.G. Menon (eds.), Hydrogen effects in catalysis, Marcel Dekker inc. New '~%rk, 1988. 2. S. L. Shannon and J.Goodwin jr., Chem. Rev., 95(1995), 677. 3. P. Marginean and A. Olariu, J. Catal., 95 (1985) 1. 4. P. Marginean and A. Olariu, Appl. Catal., A: General, 140 (1996) 59. 5. G.M. Pajonk, S. J. Teichner and J. E. Germain (eds.), Spillover of adsorbed species, Elsevier, Amsterdam, 1993. 6. V. Almasan, N. Aldea and P. Marginean, The IIF d national symposium of catalysis, oct 1993, Bucharest. 7. W. C. Conner jr, and J.L. Falconer, Chem. Rev., 95(1995), 759. 8. H. W. Chen and J.M. White, J. Mol. Catal. 35(1986), 355. 9. T. Inui, K. Fujimoto, T. Uchijima and M. Masai (eds.), Studies in Surface Science and Catalysis 77, Elsevier, Kyoto, 1993.

91997 Elsevier Science B. V. All rights reserved. Dynamics of Surfaces and Reaction Kinetics in Heterogeneous Catalysis G.F. Froment and K.C. Waugh, editors

553

Transient Investigation Of The Catalytic Activity Of Copper In NO Decomposition Over Cu-ZSM5 R. Pirone a, P. Ciambellib, E. Garufi a and G. Russo c alstituto di Ricerche sulla Combustione (CNR), p.le Tecchio 80, 80125 Napoli, Italy bDipartimento di Ingegneria Chimica e Alimentare, Universita di Salerno, Fisciano, 84084 Salerno, Italy CDipartimento di Ingegneria Chimica, Universita. di Napoli "Federico II", p.le Tecchio 80, 80125 Napoli, Italy The transiem interaction of NO with Cu-exchanged ZSM5 at 80~ has been studied to investigate the redox chemistry of Cu in these catalysts. The unsteady-state formation of gaseous N20 upon NO adsorption on He pretreated zeolite (550~ has been observed and associated to the reoxidation from Cu § to Cu .2 of a fraction of Cu sites. A similar dependence on the Cu content for NO steady-state decomposition runs at 500~ and N20 transient production at 80~ has been obtained for Cu-ZSM5 catalysts at Si/A1 = 25 and 80, showing the ability of this transiem analysis in titrating Cu sites active in NO decomposition.

1. INTRODUCTION The catalytic decomposition of NO is very attractive in principle as an effective process for the control of polluting emissions. Among all catalysts investigated, Cu over-exchanged ZSM5 zeolite exhibits the highest activity for this reaction [ 1]. It is known [2-5] that copper sites in Cu-ZSM5 zeolites could be easily reduced (from Cu +2 to Cu § by only raising the temperature (above 400~ under vacuum or in inert gas flow, due to spontaneous desorption of oxygen, according to the following stoichiometry: 2 (Cu+2-O2-Cu § ~ 4 Cu ~ + 2 02

(1)

Moreover, it has been also proposed that NO decomposition over Cu-ZSM5 proceeds via a redox mechanism in which the Cu ~ site acts as the active centre [ 1,6-7]. Consequently, the process is effective above 400~ since at lower temperatures the oxygen produced by the reaction cannot be desorbed neither the reduced state of catalyst restored. We have shown [2] that an investigation of the transient interaction between NO and CuZSM5 at temperatures so low that the catalyst has no steady-state activity should contribute to elucidate the redox chemistry of copper in these systems, that is strictly related to the mechanism of NO decomposition. In this work the study was extended to Cu-HZSM5 catalysts at different copper content and Si/A1 ratio.

554 2. EXPERIMENTAL The parent zeolite H-ZSM5 has been synthesized according to the procedure reported in [8]. The Cu-ZSM5 catalysts have been prepared by ion exchanging two parent H-ZSM5 with different Si/AI ratios (25 and 80) in aqueous solution of copper acetate. The Cu content of the samples ranges from 0.38 to 3.94 wt.%, corresponding to a degree of exchange from 41 to 640%. The samples have been labelled as CuHZx(y), where x represents the Si/AI ratio and y the degree of copper exchange. Before each experiment, the catalyst has been treated in He (25 N1/h) for 2 h at 550~ in order to obtain the reduced state and then quenched down to the desired temperature. The kinetic measurements have been carried out at 500~ by feeding a gas mixture containing 0.5 % vol. of NO diluted in He. Contact times as low as to keep the NO conversion smaller than 15% have been chosen in order to calculate the decomposition rates in conditions of differential reactor. In the transient experiments a gas mixture containing NO (600 ppm) and He (as a balance) has been fed to the catalytic reactor at 80~ at which temperature CuZSM5 catalysts exhibit no steady-state activity. The gas analysis has been continuously carried out by measuring and recording the inlet and outlet concentrations of NO, NO2, N20 and 02 with two Hartmann & Braunn analyzers (URAS 10 E and MAGNOS 6 G).

3. RESULTS AND DISCUSSION

Fig. 1 shows the typical patterns of NO and N20 outlet concentrations after a step change of NO concentration in the feed (0-600 ppm) to the pre-reduced CuHZ80(640) catalyst. 600 . . . . . . . NO inlet concentration

f

E =~ 400

o

. ,,.,~

.-,., N20." ."

o 200

r-,

_ ,

0

! L

~

"..\

, 10

...... ,...... 20

30 time, min

500

..... 1000

Figure 1. NO (--) and N 2 0 ( ..... ) outlet concentrations of as functions of time in NO adsorption run on pre-reduced Cu-ZSM5(640). T = 80~ NO inlet concentration = 600 ppm.

555 Since the catalyst has no steady-state activity at 80~ towards reactions involving NO, as shown by our previous studies [8], the steady-state values for NO and N20 outlet concentrations are the inlet ones. Notwithstanding, it can be observed the transient appearance of N20 in the gas phase, after an initial period in which no nitrogen containing species are present in the outlet gas mixture. We have explained the observed phenomena as follows. In the first period, NO is irreversibly adsorbed on the reduced catalyst surface and re-oxidizes copper sites, according to the following stoichiometry: 2 NO + 2 Cu § ~ N20 + Cu§

"2

(2)

Exchange level, % -~

0

200 ""1

~

"3 1 5 -

l

40o '

I

60o '

I

~ii 12

~

6

.~

3

& 2:

o

~

160 -

r

> 120 -

Z

b

/ "2

o ~~..~; v l 1 I 0 1 2 3 Cu content, % wt.

4

Figure 2. a) Amount of N20 produced as a function of Cu content in NO transient adsorption on pre-reduced CuHZ80 at 80~ (NO inlet concentration = 600 ppm); b) NO decomposition rate (R~ as a function of Cu content on CuHZ80 at 500~ (NO inlet concentration = 5000 ppm).

Only when the whole surface is covered by nitrogen-containing species, the appearence in the gas phase of N20 produced by reaction (2) is possible. Since no 02 is detected in the gas phase during the whole run, the measure of the total amount of N20 produced in such transient experiments allows to titrate copper sites that can undergo their self-reduction at high temperature, according to the stoichiometry of reactions (1) and (2). In particular, during the transient experiment on CuHZ80(640) the total amount of N20 produced was 1.3-10 .4 mol/gc~t, corresponding to a 2.6-10.4 mo[/g~t of copper re-oxidized by NO. Since the total amount of copper contained in CuHZ80(640) is 6.2-10.4 mol/g~t, this result shows that only a fraction (40%) of the whole copper can be reduced in He flow at 550~ as measured by its reoxidation with NO at 80~ On the other hand, it has been already reported [9] that not the whole copper in Cu-ZSM5 is associated to A1 framework, neither all Cu .2 species are reduced to Cu § in the absence of a reducing agent. Consequently, we have compared the transient behaviour of NO interacting with Cu-ZSM5 zeolite at low temperature to the catalytic activity towards NO decomposition, assuming that only Cu sites which are able to move through the two states of oxidation, succeed in decomposing NO into N2 and 02. In order to verify the possibility of correlating transient experiment results of NO adsorption at low temperature with steadystate measurements of catalytic activity for NO decomposition, CuHZ80 samples with different copper content have been tested (Fig.2).

556 Unsteady-state runs show that the amount of reducible copper grows more than linearly with the total Cu content, even at values much higher than 100% exchange (Fig.2a). For example, the amount of N20 increases one order of magnitude from 147 to 640%. On the other hand, the values of NO decomposition rate (Rg in Fig.2b) show that the same trend can be assumed for both the catalytic activity at 500~ in NO decomposition and the amount of N20 produced by NO adsorption at 80~ as function of copper content in CuHZ80 catalysts. A different effect of copper overexchange was found in a parallel investigation on the selective reduction of NOx with hydrocarbons at spark ignition engine exhaust. In that case the conversion of NOx increased with copper Exchange level, % content up to about 200% exchange level and 50 100 150 then became constant [10]. Since the Si/AI I ' I ~ . I ratio in these catalysts is relatively high (80), E 15 the large atomic distances among Al sites (there are about 1 A1 per unit cell of zeolite) " 12 could make the probability that four Cu sites interact to undergo the self-reduction of _= 9 reaction (1) very low, if the exchanged ionic 9 copper is only that associated to the AI ions, according to the usually assumed ion9 6 exchange stoichiometry (1 Cu for 2 A1). On z the contrary, being the maximum value of the ~ 3 atomic ratio of [Cu]/[A1] in CuHZ80 about 3.2 for CuHZ80(640) and the presence of "-" 0 bulk copper oxide not observed in any 2: CuHZ80 catalyst [10], a different exchange's stoichiometry must be hypothesized. The ~600 presence of [Cu(OH)] § species and their > related oligomers such as [CuxO2(OH)2x.5]+ o instead of Cu +2 as exchanging cations in Cuc400 ZSM5 had been already suggested [2,11-13 ]. Such species could explain both the copper over-exchange and its self-reduction in the -~ 200 E catalysts with Si/A1 ratio as hio=Ja as 80. In 9 fact, for x > 4, the very small distances among z , AA/ i I i ~ 0 Cu atoms, all associated to the same Al site, 0 1 2 3 allow a local thermal self-reduction to occur. For example, for x = 4, the stoichiometry of Cu content, % wt. Cu reduction is Figure 3. a) Amount of N20 produced as a function of Cu content in NO [Cu402(OH)3]+ -+ [Cu4(OH)3] + + 02 (3) transient adsorption on pre-reduced CttHZ25 at 80~ (NO inlet In order to investigate the effect of Si/A1 concentration = 600 ppm); b) NO ratio, transient runs were performed with decomposition rate (Rg) as a function of CuHZ25 samples with different copper Cu content on CuHZ25 at 500~ (NO content (Fig.3). inlet concentration = 5000 ppm). m

Ue~.

--

~9

a

557 Fig.3a shows that the amount of N20 formed on CuHZ25 samples in the transient NO adsorption runs sharply increases with the Cu content, much more than for the catalysts at Si/A1 = 80, becoming three times greater by only passing from 70 to 100% exchange. Moreover, the fraction of the reducible copper measured by these transient experiments are much greater than those resulted on CuHZ80 samples, up to achieve the re-oxidation of a fraction of more than 90% of the whole copper exchanged in the transient runs carried out on the catalysts at the highest Cu content. For example, over the CuHZ25(107) sample, we have obtained an amount of 1.46-10 -4 mol/g~ of N20, which corresponds to the re-oxidation of 2.9-10-4 mol/g~ of copper, i.e. 92% of the total copper content. Also the values of NO decomposition rate (Rg), reported in Fig.3b, appear much higher than those reported for CuHZ80 (Fig.2b). It can be observed that even the activity of a 60% copper exchange sample is higher than the most active CuHZ80 catalyst (640% exchange). Moreover, as for the CuHZ80 samples, the values of Rg depend on copper content in a quite similar way with respect to the transient results, increasing steeply with the Cu content around 75% copper exchange. These different behaviours can be probably attributed to the different Si/Al ratio of the CuHZ80 and CuHZ25 catalysts. In fact, at Si/Al = 25, in the presence of an average value of 3.84 Al atoms per unit cell, the higher Al content allows copper self-reduction to occur according to the usually assumed stoichiometry (1), being not necessary to achieve copper content as high as corresponding to 640% exchange, as it can be observed for the CuHZ80 catalysts. On the other hand, the transient analysis camed out at 80~ seems to succeed to titrate the copper sites that are active in NO decomposition, whatever the Si/Al ratio of the framework. This occurrence confirms the redox nature of NO decomposition mechanism and indicates that only the Cu sites able to undergo the thermal self-reduction above 400~ can have a catalytic role.

4. CONCLUSIONS By investigating the transient behaviour of NO interaction with Cu-ZSM5 catalysts with different copper contents in a range of experimental conditions for which no steady-state activity is exhibited, a significant contribution to understand the copper redox chemistry, which appears to be hardly involved in NO decomposition mechanism, can be attained. In particular, an interesting correlation between the transient production of N20 and the catalytic activity for NO decomposition has been found. A similar trend is observed for both the catalytic activity at 500~ in NO decomposition and the amount of copper re-oxidized by NO adsorption at 80~ as functions of copper content in CuHZ zeolites. As a conclusion, NO transient adsorption experiments allow to titrate copper sites which undergo self-reduction from Cu § to Cu +, by desorbing molecular 02 above 350~ Consequently, considering the redox nature of the reaction mechanism, these sites are probably involved in NO decomposition.

5. ACKNOWLEDGEMENTS This work was supported by ENEL/CRAM Milano (Italy).

558 REFERENCES 1. M. Iwamoto, H. Furukawa, S. Kagawa, New Developments in Zeolite Science Technology, Ed. by Y. Murakami, A. Iijima, J. W. Ward, Elsevier, Amsterdam (1986) 943. 2. R. Pirone, E. Garufi, P. Ciambelli, G. Moretti and G. Russo, Catal. Lett., (1997) in press. 3. E. Giamello, D. Murphy, G. Magnacca, C. Morterra, Y. Shioya, T. Nomura and M. Anpo, J. Catal. 136 (1992) 510. T. Cheung, S.K. Bhargava, M. Hobday and K. Foger, J. Catal., 158 (1996) 301. 5. J. Valyon and W.K. Hall, New Frontiers in Catalysis, Ed. by L. Guczi, F. Solymosi, P. Tet6nyi, Elsevier, Amsterdam, Part B (1993) 1339. M. Iwamoto, H. Yahiro, N. Mizuno, W.-X. Zhang, Y. Mine, H. Furukawa and S. Kagawa, J. Phys. Chem., 96 (1992) 9360. W. K. Hall and J. Valyort, Catal. Lett., 15 (1992) 311. 8. R. Pirone, P. Ciambelli, G. Moretti and G. Russo, Appl. Catal. B, 8 (1997) 197. 9. D. J. Parrillo, D. Dolenec, R. J. Gorte and R. W. McCabe, J. Catal., 142 (1993) 708. 10. P. Ciambelli, P. Corbo, M. Gambino, G. Minelli, G. Moretti and P. Porta, Catal. Today, 26(1995)33. 11. G. Moretti, G. Minelli, P. Porta, P. Ciambelli, P. Corbo, M. Gambino, F. Migliardini and S. Iacoponi, Studies Surf. Sci., 105 part B (1996) 1525. 12. G.D. Lei, B.J. Adelman, J. Sarkany, W.M.H. Sachtler, Appl. Catal. B, 5 (1995), 245. 13. K.C.C. Kharas, D.-J. Liu and H.J. Robota, Catal. Today, 26 (1995) 129. .

.

.

91997 Elsevier Science B. V. All rights reserved. Dynamics of Surfaces and Reaction Kinetics in Heterogeneous Catalysis G.F. Froment and K.C. Waugh, editors

559

T h e r m o d y n a m i c T r a n s i t i o n - S t a t e Theory a n d E x t r a t h e r m o d y n a m i c C o r r e l a t i o n s for the L i q u i d - P h a s e Kinetics of Ethanol D e r i v e d Ethers Ravindra Datta, Kyle Jensen, Prakob Kitchaiya and Tiejun Zhang Department of Chemical and Biochemical Engineering, The University of Iowa, Iowa City, Iowa 52242, USA. The thermodynamic transition-state theory (TTST) is utilized for the elementary steps within the Langmuir-Hinshelwood-Hougen-Watson (LHHW) framework to develop rate expressions for liquid-phase catalytic reactions in terms of activities for the family of tertiary alkyl ethyl ethers. The TTST formulation also provides a rationale for the extrathermodynamic correlations (ETC) observed. 1. INTRODUCTION The liquid-phase chemical reaction studied here is the ethanol-derived family of tertiary ethers, namely ethyl tert-butyl ether (ETBE) from isobutylene, tert-amyl ethyl ether (TAEE) from the Cs tertiary olefins 2-methyl-l-butene (2MIB) and 2methyl-2-butene (2M2B), and tert-hexyl ethyl ethers (THEE) from the C 6 tertiary olefins, i.e., THEE1 from 2-methyl-l-pentene (2MIP) and 2-methyl-2-penene (2M2P), THEE2 from 2,3-dimethyl-l-butene (2,3-DMIB) and 2,3-dimethyl-2-butene (2,3-DM2B), and THEE 3 from cis-3-methyl-2-pentene (C3M2P), trans-3-methyl-2pentene (T3M2P), and 2-ethyl-l-butene (2EIB). The liquid-phase etherification reaction takes place under mild conditions in the presence of an ion-exchange resin catalyst with sulfonic acid groups such as Amberlyst-15 and is accompanied by isomerization reaction of tertiary olefins C5 or higher. When blended with gasoline, tertiary ethers enhance the gasoline octane number, replace aromatics and are effective in carbon monoxide reduction. In the US, the Clean Air Act Amendments have created a huge market for such oxygenates for blending purposes. Should the supply of isobutylene become limiting, ethers produced from the Cs and C 6 tertiary olefins will become of interest. The liquid-phase reaction kinetics of this family of tertiary alkyl ethyl ethers can be consistently described within the LHHW formalism [1], with the TTST being applied to the elementary mechanistic steps. The systematicity of reaction kinetics is discussed here based on ETC, which describe relation between kinetics and thermodynamics. Further, it is shown that the rate expressions for this nonideal liquid-phase reaction system should be written in terms of activities.

560 2. THEORY

It is assumed that the composite catalytic reaction involves several elementary steps, e.g., adsorption, surface reaction, and desorption, which may individually be treated according to TTST, i.e., each step is assumed to possess its own transition state. For example, for the adsorption of A, the forward step is represented by, A + S ~ [XA]*--+ A'S. The free energy changes of activation associated with each step may, of course, be substantially different p r o v i d i n g justification of the c o m m o n assumption of the rate determining step (rds). The rate of the forward elementary step i is proportional to the universal frequency, v = k J / h [2], a transmission coefficient, ~r varying between 0 and 1 [3], and the concentration of the transition-state complex (TSC) F, =Ic

ksT h

C

(1)

tx, f

From the equilibrium constant of the formation of the TSC from reactants, j = 1, ..., r,

k :, =

(I

j=l

; =

-

(2)

RT

the link between the thermodynamics and the kinetics of the reaction is obtained, based on the assumption of psuedo-equilibrium postulate of the TTST. Thus -v

r-, = k lZla; '

(3)

j=l

where the activity of species j, a j = 7'ix j, and the rate constant --

k,-x'----ex h 7i~;f

-

RT )

._

Ir

h

7tx, f

R

ex -

d~-/iT

RT

(4)

Since the concentration of the TSC is vanishingly small, its activity coefficient may be assumed to be constant, corresponding to its value at infinite dilution. Thus, even though the rate of reaction on a catalyst is assumed to be proportional to the surface c o n c e n t r a t i o n of the transition-state complex (TSC), this in t u r n becomes related to the species a c t i v i t i e s in the bulk fluid due to the assumed pseudo-equilibrium between TSC. Further, the ratio of the forward and reverse rate constants in k, k,

1 (zX4:~ zX4:o RT

1

o

(5)

561 Comparing eq 4 to the experimental Arrhenius form k , - .4, exp(-/~,/RT)

L - ~ , : ; +RT

(6)

and A,-~c--~ex h YE.v,f

tT

1+

(7)

R

Thus, for the forward and reverse reactions in A,_ = ~ : ctT - ~ : ~tT = AS tTc

(8)

A,

and ~ , - ~, - a~9 :~, - a ~ :~,: - aH;~

(9)

If it is assumed that the heat capacity of the TSC is a linear combination of those of reactants and products, the linear free energy relation results 5 ( A ( ~ ) - a,5(A(~;r)

(10)

3. R E A C T I O N K I N E T I C S

Through the use of partially deactivated catalysts to vary the total n u m b e r of active catalyst sites, Jensen and Datta [4] and Kitchaiya and Datta [5] d e t e r m i n e d that a total of three sites are involved in the surface reaction rate-determining step of the etherification reaction (A + B ~ D), with two adsorbed ethanol sites reacting with one adsorbed tertiary olefin, while a total of two sites are i n v o l v e d in the surface rate determining step of the olefin isomerization (B ~ C), with an adsorbed olefin reacting with a vacant site, where A represents ethanol, B and C iso-olefins and D the ether product. Applying the TTST formulation to each of the elementary steps of the mechanism, along with the assumption of pseudoequilibrium for the adsorption and desorption steps with surface reaction as the rate determining step and ethanol the most abundant surface species a s s u m p t i o n (MASSA) owing to the fact that K a >> K 8, Kc, and K~, yields the following rate expressions for etherification

562

r,

(1+ KAaA)3

(11)

while for isomerization

k~3(aB-at~K3) r3-

(12)

(I+KAa.4)2

where the ethanol adsorption equilibrium constant KA for ethanol on Amberlyst 15 was determined by independent adsorption experiments by Kitchaiya and Datta [5]. The above rate expressions were found to fit the entire family of the tertiary alkyl ethyl ethers studied [4 - 6]. There has been some question as to the validity of activities in rate expressions. Thus, if the activity coefficients are assumed to be unity, the species activities a i in the rate expressions may be replaced by mole fractions x i. The resulting expressions are compared in Figure 1 for the case of TAEE. It is seen that the etherification rate expression in which the activities have been replaced by mole fractions predicts initial rates that are higher than the observed rates by up to a factor of ten.

e.,0

.d

,E ~/

o

E

.

.

9 o.

.

Activity f Mole Fraction j/ . .

F

e,.

o

o

ooo;

_./ <> // ~

-_+

t~

r

o

1

o

?

J /

0.1 ~ 9

....... ,

........ ,

........ ,

....... t

0.1 1 10 100 1000 Experimental Rate of Reaction ( m m o l / h . g )

Figure 1. Comparison of the rate expression using mole fraction or activities for TAEE synthesis.

563 Figure 2 compares the ratio of the forward and reverse rate constants for the three reactions involved in TAEE synthesis with the corresponding equilibrium constants for the reactions determined independently [6]. These are consistent, as expected.

60

i

,

~

i _j /

F

50,:_ ~

:

p

1

k

40

. . . . .

K 2

.........

K 3

.m

? 2

?

J

3O

4

"b-... I0

........................ .<>...~ ._i '

~ ,

310

_q

o

?

315

320

325

330

,

i

335

340

345

T (K) Figure 2. Comparison of the ratio of forward and reverse rate constants with the equilibrium constants for etherification (K 1 and K2) and isomerization (K3) in the TAEE system. The linear free energy relation of the form of eq 10 is also shown to be valid in Figure 3 for the case of the six C 6 tertiary olefins that form tertary hexyl ethyl ethers. 4. CONCLUSIONS It is shown that rate expressions in terms of activities are appropriate for the liquid-phase tertiary alkyl ethyl ether synthesis system. The rate expressions are based u p o n the application of the thermodynamic transition-state theory to the elementary steps within the LHHW formalism. Extrathermodynamic correlations that relate the kinetics to the reaction thermodynamics can also be rationalized within this framework and are experimentally observed for this family of tertiary ethers.

564

80

PL' '

78

'

I

'

'

'

I

'

'

'

I

'

'

'

i

'

'

'

i

' ' 2' , 3' - D! M ' 2 B ' o '

[

A i '

AG ~

~76 (kd/mol) 74 b E 72 70

~% -14

,~o.,p~

o

,,

[]

-12

-8

-6

-4

o

~_ -601

_

-10

4o0

-2

0

2

O

AG ,-r (kJ/mol)

Figure 3. Linear free energy relationship for the THEE reaction system.

REFERENCES

1. O.A. Hougen and K. M. Watson, Ind. Eng. Chem., 35 (1943) 529. 2. M. Boudart, Kinetics of Chemical Processes, Prentice-Hall, Englewood Cliffs, NJ, 1968. 3. I. Levine, Physical Chemistry, 3rd Ed., McGraw-Hill, New York, 1988. 4. K. Jensen and R. Datta, accepted Ind. Eng. Chem, Res., 1995. 5. P. Kitchaiya, and R. Datta, submitted to Ind. Eng. Chem. Res., 1995. 6. P. Kitchaiya, Ph.D. Thesis, The University of Iowa, Iowa City, IA, 1995.

91997 Elsevier ScienceB. V. All rights reserved. Dynamics of Surfaces and Reaction Kinetics in Heterogeneous Catalysis G.F. Froment and K.C. Waugh, editors

HYDRODEMETALLATION FRACTIONS.

KINETICS

OF

565

RESIDUAL

PETROLEUM

M.T. Martinez*, M.A. Callejas', E. Carb6 § and A. Hem/mdez § *Instituto de Carboquimica. P.O. Box 589, Zaragoza, Spain +Repsol Petroleo. Valle de Escombreras s/n. Cartagena (Murcia), Spain

A pre-treated Maya residue has been hydrotreated catalytically in a continuous hydroprocessing unit provided with a perfectly mixed reactor. A commercial catalyst, Topsoe TK-751, has been used and kinetic study of hydrodemetallation reactions has been performed. The conversion data fit second order kinetics for Ni removal reactions and first order kinetics for V removal reactions. The pressure dependence of the rate constants at 375~ was 1.39 for Ni and 1.82 for V and the activation energies 94.30 and 45.53 kcal/mol, respectively.

1. INTRODUCTION The basic chemical concept of the hydrodesulphurization (HDS), hydrodenitrogenation (HDN) and hydrodemetallation (HI)M) processes is to remove selectively the heteroatoms and metallic species from the organic moiety. Demetallation results from chemical transformations of the metal-beating compounds on the catalyst surface coupled with diffusional transport to and from the active sites. Unlike sulphur, nitrogen and oxygen which are removed as gaseous products (H2S, NH3, and H20), elimination of nickel and vanadium results in their deposition on the hydrotreating catalyst and deactivation of the catalyst through fouling and pore blockage. The resulting metal deposition ultimately determines total metals adsorption capacity and useful operating lifetime of the catalyst. Residue hydroprocessing requires the diffusion of multiringed aromatic molecules into the pore structure of the catalyst prior to initiation of the sequential conversion mechanism. The observed diffusion rate may be influenced by adsorption interactions with the surface and a contribution from surface diffusion. Since chemical reactions between species inherently imply transfer of mass and energy, the amounts of mass and energy required by a reaction will become increasingly difficult to transfer physically as the reaction rate increases. Usually it is necessary to obtain the intrinsic kinetic information by running experiments under conditions where transport limitations are eliminated or can be neglected.

566 2. EXPERIMENTAL A pretreated (1) residue Maya has been hydrogenated in a hydroprocessing unit provided with a continuous stirred tank reactor (CSTR). The feedstock characteristics are indicated in Table 1. It has 33.5 vol.% of compounds with boiling point above 540~ and high content of metals, heteroatoms, asphaltenes and Conradson carbon. Table 1 Feedstock characteristics IBP-540* C (Vol. %) Sulphur (wt. %) Nitrogen (wt. %) Nickel (ppm) Vanadium (ppm) Asphaltenes (wt. %) .Conradson Carbon. (wt. %)

.

66.5 2.25 0.27 30.26 121.17 5.6 8.51

The CSTR has 1000 c.c. of capacity and it is provided with a static basket and a kinetic basket for 185 c.c. of catalyst. The reactor is equipped with three furnace control thermocouples. The inside temperature is measured by two thermocouples one at the bottom of the reactor and other in the head. The utilisation of a CSTR in the kinetic study provides a tool for working in absence of intrareactor gradients but another type of gradients can be present. Gradients at the boundary between different phases (gas~quid/solid) in the reaction medium. - Gradient within the catalyst particle. In order to identify the interphase gradients, two diagnostic tests have been applied: a) Change in stirring speed until the speed does not affect the reaction rate. b) Increase the gas flow rate until no effect in the rate of reaction is observed. For working in absence ofintraparticle gradients the criterium generally utilised is to decrease the catalyst particle size until no effect in the rate of reaction is obtained. In our kinetic study, working conditions in absence of interphase gradients were investigated. The stirring speed was varied from 2000 to 3500 r.p.m, and the gas flow rate from 6000 to 10000 std.cu.ft./bbl. The increase of gas/liquid ratio above 6000 std.cu.flPobl and the stirring speed above 2500 r.p.m, did not influence the S, Ni and V conversion. Nevertheless in order to assure the absence of interphase gradients, the kinetic study was carried out at 10000 std.cu.f~bl gas/liquid ratio and 3500 r.p.m, stirring speed. The conditions of temperature, pressure and liquid hourly space velocity (LHSV) used are indicated in table 2. The catalyst used was Topsoe TK-751. It is a general purpose HDS catalyst for residual feedstocks. Their functions are desulphurization, demetallation and asphaltenes and Conradson carbon reduction. It is recommended for HDS of residua with moderate metals content and for 2"a stage catalyst in composite fillings. It has good HDS activity, good HDM selectivity and capacity for metals uptake. It is Ni/Mo type. The catalyst was presulphured for 10 hours at 3 50~ and atmospheric pressure in a flowing stream of hydrogen containing 10 (v/v) % of H~S. In order to avoid intraparticle gradients, the catalyst has been crushed at a range of particle size between 53 and 530 l.tm. To avoid -

567 misleading results by losses, catalyst has been located inside bags of 35 mm sieve aperture. The metallic bags have been introduced in the static basket of the reactor. After 180 hr of running the hydroprocessing unit, steady state conditions for the catalyst have been reached and the runs indicated in Table 2 have been started. Table 2 Operating conditions and results from the kinetic runs Temperature Pressure LHSV V (~ (MPa) (l/h g~t) (ppm) 375 10.0 2.24 49.68 375 10.0 1.62 76.31 375 10.0 0.81 85.57 375 12.5 2.30 44.15 375 12.5 1.44 57.53 375 12.5 0.73 67.13 375 15.0 1.86 21.68 345 15.0 1.43 45.25 375 15.0 0.51 54.78 400 12.5 2.28 15.15 400 12.5 1.45 32.97 400 12.5 0.67 42.43 415 12.5 2.07 5.28 415 12.5 1.37 11.59 415 12.5 0.62 34.56

Convers. (%) 59.03 37.07 29.44 63.59 52.56 44.64 82.12 62.69 54.83 87.51 72.81 65.01 95.65 90.44 71.50

Ni (ppm) 18.56 22.22 25.41 17.62 20.98 21.97 13.52 18.86 19.91 9.24 15.62 19.55 2.15 5.58 16.35

Convers. (%) 38.66 26.57 16.03 41.77 30.67 27.40 55.32 37.67 34.20 69.46 48.38 35.39 92.89 81.56 45.97

For a heterogeneous perfectly mixed system working continuously and after steady state conditions have been reached, the rate equation assuming constant density Co - C

1

-r

LHSV

being Co = inlet concentration. C = outlet concentration LHSV = Liquid hourly space velocity -r = Metals or heteroatoms removal rate = K* pm H: C"

(1)

(2)

being K* = intrinsic rate constant m = order dependence on H2 pressure n= kinetic order P m = hydrogen pressure To obtain the kinetic order n, it has been worked at constant pressure, 12.5 MPa, and it has been defined a pseudokinetic rate constant K.

K-- K* P,7~

(3)

568 3. RESULTS AND DISCUSSION The nickel and vanadium concentrations in the products and the conversions obtained from the kinetic runs are indicated in table 2. The data of table 2 fit first order kinetics for V removal reactions and second order kinetics for Ni removal reactions. Figure 1 and Figure 2 illustrate first order kinetic plot for V removal reactions and second order kinetic plot for Ni removal reactions, respectively. 15

4O

20

I05 1

1 0 o

4150C

7.5 5.0 2.5

400"C

/

.

m

~_._.~.

C

b 0 r,.) ~"

0.4

0* 12

0.2 0.0

3750C 0.0

0.5

1 .0 I/LHSV

I .5

2.0

Figure 1.- First order kinetic plot for V removal reactions at 12.5 MPa H2 pressure

o

.o

"

o'

0.0

0.5

! .0 I/LHSV

1.5

2.0

Figure 2.- Second order kinetic plot for Ni removal reactions at 12.5 MPa H2 pressure

The linear equations (4) and (5) represent the pressure dependence of the rate pseudokinetics constants at 375~ for V and Ni, respectively. Ln K= 2.42 10-4 + 1.82 Ln Pm

(4)

Ln K - 4.19 10.5 + 1.3 9 Ln Pm

(5)

The pseudokinetic rate constants and the intrinsic rate constants of the Ni removal reactions at 375, 400 and 415~ are indicated in Table 3. The activation energies for Ni and V removal reactions were obtained from the semilogarithmic representation of In K versus 1/T by applying the Arrhenius equation, Figure 3. The values of m, n and Ea are indicated in Table 4. There are discrepancies in reaction kinetic order with respect to the total metal concentration. DemetaUation kinetics have been described in the literature by using first-order, second-order or in between first and second order rate laws (2). Measurements of intrinsic kinetics of metal removal from model oil comprised of metalloporphyrins have been reported with values ranging from 0.1 to 1.1 and hydrogen pressure dependence greater than first order. Nevertheless, hydrodemetallation kinetic studies of petroleum oils and residues indicate higher kinetic orders. Riley (3) reported first order kinetics for both nickel and vanadium removal in the hydrotreating a Safaniya atmospheric residue, Oleck and Sherry (4) found a better description of the reaction system with second order kinetics for nickel and vanadium removal from Lago-medio (Venezuelan) atmospheric residue, Van Dongen (5) indicated that vanadium removal can be described satisfactorily by an apparent reaction order of 1.0 in a CSTR and 1.5 in a plug flow reactor and Inoguchi et al. (6) found that both first and second-order descriptions equally fit the data from their plug-flow reactors. The apparent HDM reaction orders greater than one have been

569 attributed to the presence of more than one class of metal compounds reacting with different rates (4), (7). Just as in hydrodesulphurization, the simultaneous occurrence of several first-order reactions with different rates can lead to an apparent reaction order greater than unity (8). Wei and Hung (9) theoretically demonstrated conditions whereby two first-order reactions give rise to apparent sexond order kinetics. Table 3 Intrinsic (K*) and pseudo kinetic (K) constants corresponding to HDM reactions. 375~

400~

415~

NICKEL

( 1 m./ KI25MPa, . . . ,gc., aht . "pp

3.06.10 2

1.46.10 1

3.10

0.97

0.9

0.87

4.19.10 5

1.77.10.4

3.77.10 .3

K125MPa (h. g l . ppml

1.38

4.53

12.96

Correl. Coefi.

0.90

0.99

0.96

2.42.104

6.92.10 .4

1.98.10 .3

Correl. Coeft.

K, h.~. ~(~/~m') ~ VANADIUM

K*[

1

h'gcat-ppm'(Kg~/cm2)I~2

-3

-6 V an ad iu m ,.d

-9

N ickel

-12

-15

..

1 .4 4 x 1 0 " 3 91 .4 8 x| 1 0 " 3 "1 .5 2 x| I 0 - 3 "1 .5 6 x| 1 0 - 3 "l .6 0 x 1 0 -3

1/T Figure 3.- Arrhenius plot for the rate constant of the V and Ni removal reactions.

570 Table 4 Kinetic orders and activation energies coresponding to the HDM reactions NICKEL

VANADIUM

Kinetic order .......................................... 2 Order of depencence on H2 pressure ....... 1.39 Correl. coeft ........................................... 0.95 Ea (kcal/mol) .......................................... 94.30 Correl. coeft ........................................... 0.93

1.82 0.96 45.53 0.99

1

Most investigations (10) have revealed that vanadium removal rates exceed those as nickel in petroleum residua. In our case, the values of the rate constants are not comparable since they fit different kinetic orders. Nevertheless the percentages of nickel removal are in all cases lower than those of vanadium removal. The higher value of activation energy of the nickel removal reactions indicates a higher temperature dependence, it is for this that the nickel converions at high temperature (415~ approach to those of vanadium conversions.

ACKNOWLEDGEMENTS The authors wish to thank the European Community (Contract JOU2-CT92-0206) and the DGICYT (AMB93-1137-CE) for financial support.

REFERENCES

1-M.T. Martinez, Hydroprocessing of heavy petroleum fractions. Final Report Contract JOU2- CT92-0206. (1996) 2. K.L. Riley, ACS Prepr. Div. Petrol. Chem., 23 (1978) 1104. 3. C.J. Pereira, Ind. Eng. Chem. Res., 29 (1990) 512. 4. S.M. Oleck and H.S. Sherry, Ind. Eng. Chem. Process. Des. Dev., 16 (1977) 525. 5. R.H. Van Dongen, D. Bode, H. Van der Eijk and J. Van Klinklen, Ind. Eng. Chem. Process Des. Dev., 19 (1980) 630. 6. M. Inoguchi, H. Kagaya, K. Daigo, S Sakurada, Y. Satomi, K. Inaba, K.Tate, tL Nishiyama, S. Onishi and T. Nagal, Bull. Jpn. Petrol Inst., 13 (1973) 153. 7. R.R. Cecil, F.Z Mayer and E.N. Cart, AIChE Meet., Los Angeles 1968. 8. A. De Bruijn, Proc. Int. Cong. Catal. 6th, 2 (1976) 951. 9. J. Wei, and C.W. Hung, Ind. Eng. Chem. Process. Des. Dev., 19 (1980) 197. 10. J. Quann, R.A. Ware, C.W. Hung and J. Wei, Advances in Chemical Engineering, 14

(1988) 95.

91997 Elsevier Science B. V. All rights reserved. Dynamics of Surfaces and Reaction Kinetics in Heterogeneous Catalysis G.F. Froment and K.C. Waugh, editors

571

D e v e l o p m e n t o f a c o m p u t a t i o n a l tool for the t r a n s i e n t kinetics o f c o m p l e x chemical heterogeneous reaction systems G. A. Carrillo Le Roux a, I. Bergault b, H. Delmas b and X. Joulia b aLaborat6rio de Simulag~o e Controle de Processos - DEQ - EPUSP - P.O. Box 61548, 05424970 - $5.o Paulo - Brazil - e-mail: [email protected] bLGC - INPT / ENSIGCT - 18, Chemin de la Loge - 31078 Toulouse CEDEX 4 - France - email: Xavier.Joulia@ensigct. fr

We present a program for the modelling of the transient of complex heterogeneous chemical reaction systems. The philosophy of the program is based on the formulation of a generic model for heterogeneous reactions derived from a phenomenological description of the system, where all the reaction steps should be detailed. The challenge of dealing with the complexity of the model thus generated is addressed by the robustness of the numerical tools employed for its simulation and parameter estimation. The package is able to obtain a fit even for complex models. This fact justifies a meaningful discussion on the usual philosophy of model analysis which consists of steps of reduction and discrimination. The possible issues are illustrated with a study on the hydrogenation of acetophenone over Rhodium catalyst. 1. I N T R O D U C T I O N The usual approach for the generation of models for heterogeneous reacting systems, derived from Langrnuir-Hinshelwood Hougen-Watson (LHHW) equations, involves stages of model reduction by means of hypothesis of rate-controlling steps, quasi steady-state approximations, fast reversible reactions hypothesis and the implementation of lumping of species (1). The approach is based on model discrimination techniques, which involves fitting and comparison of a high number of candidate models, generated by the combinatorial of all pertinent hypothesis (e.g. the dehydrogenation of 1-butene to butadiene (1) which involves the discrimination of 15 candidate models, all based on elementary rate-determining steps assumptions). For complex systems this task can be very tedious and its results disappointing. However, in most cases, the necessity of this step can only be justified by the pragmatic argument that it is very difficult to fit models too complex to data. In this work, we propose an approach to the resolution of the equations derived from the modelling of heterogeneous systems and their fit to experimental data, that permits to relax the necessity of formulation of some of the restrictive hypothesis cited above. In consequence, the worth of the arguments that justified the approach based on steps of reduction and discrimination is reviewed. We present some ideas that can contribute to the development of alternative paths.

572 2. M O D E L

IMPLEMENTATION

The model is formulated in terms of mixed differential and algebraic equations. It is suitable for liquid-solid and gas-liquid-solid reactors. The main assumptions are that the liquid volume and solid mass inside the reactor are constant and that the effects of mass and heat transfer can be neglected. It was generated from balances on both the liquid and solid systems. 2.1 Material Balance in the Liquid Phase The material balance for the liquid phase is expressed as follows:

dc dt

- FIlr r ( 0 , c , | 1 7 4 2.... ,|

(1)

+ FI,er

where c is a no-vector (where nc is the total number of species) which groups the concentrations (in mole per unit volume of liquid) of all the species, in the liquid phase, 0, is a r~-vector of parameters which assembles all the kinetic and equilibrium parameters and | (with i varying from 1 to ns, the number of different sites types) is a n~+l-vector which groups the surface concentration (in mole of species per total number of active sites of type i) of all the species in the site of type i, including the surface concentration of free sites of this type, i noted | f. r is a real n~-valued function by means of which all the reaction velocities are evaluated, without distinction of whether they take place in liquid or solid phase. However, as the dimension of r is mole per unit volume of liquid phase, per unit of time, if a chemical reaction taking place on the solid phase is considered, the reaction terms are to be multiplied by the catalyst concentration in mass of catalyst per unit volume of liquid phase, such that the expression be consistent with equation (1). FI~ is a nr x n~ matrix, each of its columns corresponds to a reaction and contains the stoichiometric coefficients associated to each specie in the liquid phase. Algebraic constraints, in a number of ne, are used to represent fast reversible reactions and equilibrium terms like those arising from adsorption phenomena. They will be presented further. ~ is a n~-vector of displacement velocities. They are slack variables (2) that have the same unit as r. They can be understood as the velocity by which each specie must vary, while the system is reacting, in order to keep the algebraic equations that expresses equilibrium constraints satisfied. FI~ is a nc x n~ matrix which multiplies ~, in order to keep elemental balances satisfied. 2.2 Material Balance in the Solid Material balances are built for each type of site over the catalyst. For a given site of type i, we write down balances for each of the j = 1..... no, adsorbed species:

w C is d|at - FI~, r(O,c,O~,O 2,...,|

+ ~r

(2)

573 and for the free sites of type i we use a material balance expressed in algebraic form: nr

(3)

O j~+ O f ~ --1 j=l

where | j~ is the surface concentration of each of the adsorbed species (j = 1, ..., no, currently grouped in |

and |

f

that of the free sites, w is the catalyst concentration in mass of catalyst

per unit volume of liquid phase and C~s the concentration of active sites of type i, in mole per unit mass of catalyst. ~ r and ~ e are the stoichiometric coefficients of the species adsorbed on the site of type i corresponding to reactions and displacement velocities and are matrices of size nc x n~ and n c x nr respectively. The algebraic constraints, which can express fast reversible reactions or adsorption equilibria correspond to a set of ne equations which can be expressed by:

0=feq(0,C, Ol,O2,...,O ns)

(4)

where feq is a ne-valued real function. A complex heterogeneous reacting system can be thus described by a set of nc (1 +ns) differential equations (equations (1) and (2)) and ne+n s algebraic equations (equations (3) and (4)). It is not necessary to the user to write down the differential-algebraic equations because the package generates the corresponding system on the basis of a detailed description of the reaction mechanism and the kinetic assumptions. It is not necessary to formulate physical hypothesis in order to obtain a simple expression for the overall reaction, however, some usual physical hypothesis can be easily implemented, whose parametrization is equivalent to that of the reduced one. For instance, if it is supposed that sites can be considered in pseudo-steady state it suffices to set the parameters C~sto zero, so that the differential equations (2) are transformed into algebraic ones. If simple adsorption equilibrium is to be considered for all the species over all the sites, this would be equivalent to write down equations (4) as: 0 = c ~ - K I|

i = 1 .... ,n~,

j = l .... ,n~

(5)

Evidently, each of these approaches generates specific implementation problems in the treatment of the initialization of the EDA system and integration error control, but this was generically previewed in the package (3). 3. N U M E R I C A L

TOOLS:

FACING

COMPLEXITY

The EDA system corresponding to the model is solved by a modified version of the LSODI routine, which is based on Gear's method. The version implemented performs the solution of the EDA system concomitantly to the evaluation of the parameters' sensitivities based on the decoupled direct method (4). As a matter of fact, the simulation of the system is

574 very robust, such that it can be carried out with success for very different parameter values while the sensitivities obtained are very accurate. In the estimation procedure most of the techniques described by Farris & Law (5) (essentially the rotational discrimination) and some other features (3) are implemented. This procedure, in conjunction with the great accuracy of the sensitivities, allows a fit to be obtained for models with complex parametrization. However, in most cases a large amount of computational effort is necessary, but it does not worth to remember that this kind of resource is getting less expensive everyday. The main properties of the solution proposed by the algorithm are: it does not correspond to the set of parameters that performs the best fit of the data. If it was the best one, it would very probably be useless for prediction because of the inflation of the estimate due to a possible over-fit of the experimental data (6). it can be said to be inspired on "biased estimation" techniques (7), but it cannot be classified as an estimator. T h e solution proposed is strongly dependent on the initialisation of the parameters values and on other numerical settings, so that it cannot be considered in oneto-one relation with an experimental data set, as an estimator should be. If the model performs a good fit of the data, with parameter values physically consistent, it should be considered as a serious candidate to represent the system. It cannot be considered inadequate because of the great uncertainty of some parameters, neither because of the fact that the solution proposed is not an estimate. -

-

4. D I S C U S S I O N The objective in which the conventional approach is founded is the elucidation of the intrinsic phenomenology of the system, for which the stages of discrimination-reduction are thought to be adequate. The necessity of reduction because of the incapacity of the numerical tools and hardware and the hideous results obtained with conventional estimation procedures on models with a large number of parameters are some related arguments that help to strengthen the necessity of this approach. However, the only scientific proof proper to demonstrate that a model is the "one" that represents a system resides in showing that all the other candidate models are inadequate to represent it. Any different approach does not correspond to a proof, and should be classified as selection procedures, for which preference criteria are designed. For instance, for nested models, a great deal of criteria have been developed like information criteria (8) and likelihood ratio tests (9). In this case the paradox the selection procedure is thought to solve, consists on the fact that while complexity increases, the ability of a model to represent a system also increases. It is usual that mechanisms be assumed without serious physical considerations, an overall expression being chosen from a LHHW table. These models would have the same worth of black-box ones (10) while in the "quest for the true mechanism", the possible fit obtained can lead to the illusion that the proposed mechanism is validated, while a more serious experimental analysis would have been necessary. Our approach permits that a potential candidate model be fitted to the data, independently of its complexity. This does not means that it can fit any model to any data set, however such a feature allows that more attention could be paid to the phenomenological aspects, rather than to the mathematical manipulations necessary to deduce a single expression

575 for the overall rate. It allows also that discrimination on fundamental aspects can be carried on, as many physical hypothesis can be considered in the model without much effort. Furthermore, after a fit is obtained, the Principal Components Analysis (11) can be applied in order to suggest reduction steps a p o s t e r i o r i , which can allow the model to assume a parametrization in accordance with the usual approach. 5. A P P L I C A T I O N We studied the hydrogenation of acetophenone, where a great deal of side-reactions and production of intermediaries take place in a three-phase reactor, with liquid batch and gas continuously fed, at constant pressure and temperature. The catalyst is Rhodium (3%) over activated carbon. The solvent is Cyclohexane. Samples are taken at different instants and analysed by gas chromatography. The species for which measures are available are: acetophenone, AC, phenyl-ethanol, PE, methyl-cyclohexyl-ketone, ethyl-benzene, EB, ethylcyclohexane, cyclohexenyl-ethanol, CNE, methyl-cyclohexenyl-ketone and cyclo-hexylethanol, CE. The proposed reaction mechanism is presented on figure 1. Seventeen runs were performed at 80~ at a pressure of 25 bar. The initial concentration of commercially available species (AC, PE, MCC and EB) was varied between runs.

AC ~ 2 H 2 PE

Y

.-.,

MCNC

X'8 7 EB

EC

CNE

,v

MCC

CE

Figure 1. Reaction Mechanism for the Hydrogenation of Acetophenone An Eley-Rideal mechanism was assumed for the hydrogen: as the pressure was the same for all the runs, the data was non-informative about its adsorption. Three different approaches were used for the analysis of the data: I. In this approach it was assumed that all the species are in adsorption equilibrium and that the sites can be considered in quasi-steady state. II. Independent kinetics were supposed for the adsorption and desorption of all the species. No hypothesis was made on the concentration of sites. III.PCA was applied to eliminate some parameters from the model fitted through approach II. In consequence, it was assumed that the sites are in quasi-steady state and that AC, EC, MCC and CE are in adsorption equilibrium. It was also assumed that kinetic constants for reactions 4 and 8, and 1 and 2 are not independent and are related through a linear relation.

576 The results obtained, summarized on table 1, illustrate that the PCA is able to decrease the ill-condition of the information matrix of the parameters. The criterion presented on the table was deduced from the hypothesis that the covariance matrix of measurements is diagonal and unknown, so that they can be negative. It can be noticed that posing a model with few restrictive assumptions (II) and performing reduction a posteriori, leads to a better fit. This can be an evidence that some of the assumptions made a priori could not be proved experimentally. Table 1 Summary of the results obtained for the hydrogenation of acetophenone Approach I Approach II Number of Parameters 17 26 Final Criterion -2,129 -2.429 Condition Number of the 2 . 1 0 .5 4 . 1 0 .6 Information Matrix

Approach III 21 -2.383 2. 10 -3

6. CONCLUSIONS The great advantage of the approach here presented is that it flees the analyst from having to make simplifying assumptions since the beginning of the modelling process, which can be a dangerous approach (1). The extension of the approach to other systems, following the same philosophy, is intended.

REFERENCES 1. G.F. Froment, K. B. Bischoff, Chemical Reactor Analysis and Design, John Wiley & Sons (1990) 2. V.G. Dovi, Use of slack variables in the mathematical modelling of reaction equilibria in complex chemical kinetics, Comput. Chem. Eng., 19 (4), 489-491 (1995) 3. G.A. Carrillo Le Roux, Stratrgie d'Identification de Modrles Algrbro-Diffrrentiels. Application aux Systrmes Rractionnels Complrxes, Thrse de Dr. Ingrnieur de l'Institut National Polytechnique de Toulouse, France (1995) 4. A.M. Dunker, The decoupled direct method for calculating sensitivity coefficients in chemical kinetics, J. Chem. Phys., 81 (5), 2385-2393 (1984) 5. R.H. Farris, V.H. Law, An efficient computational technique for generalized application of maximum likelihood to improve correlation of experimental data, Comput. Chem. Engng., 3, 95-104 (1979) 6. S. Vajda, H. Rabitz, E. Walter, Y. Lecourtier, Qualitative and quantitative analysis of nonlinear chemical kinetic models, Chem. Eng. Comm., 83, 191-219 (1989) 7. D.W. Marquardt, Generalized inverses, ridge regression, biased linear estimation and nonlinear estimation, Technometrics, 12 (3), 591-612 (1970) 8. L. Ljung, System Identification, Prentice-Hall Inc., N.J. (1987) 9. D.M. Bates, D.G. Watts, Nonlinear Regression Analysis and its Application, Wiley (1980) 10.J.B. Butt, Reaction Kinetics and Reactor Design, Prentice-Hall, New Jersey (1980) 11.S. Vajda, P. Valko, T. Turanyi, Principal component analysis of kinetic models, Int. 3.. Chem. Kinet, 17, 55-81 (1985)

91997 Elsevier Science B. V. All rights reserved. Dynamics of Surfaces and Reaction Kinetics in Heterogeneous Catalysis G.F. Froment and K.C. Waugh, editors

577

Methods of elimination and the problem of nonuniqueness

of inverse problem solutions in models of non-stationary chemical kinetics

S.I.Splvak and R.M.AsadullLn Department o f Mathematics, The Bashkir State University, Ufa, 450074, Russia I. INTRODUCTION

Consider

the

mathematical

m o a= ~l

chemical reaction

of

a

fi~s . +~ order

complex

c=Ac

c(O) = c o

(!)

where c :

(ci,..., c q ) are the concentrations of reactants, participating in reactions. The elements of the matrix A are linearly dependent on the rate constants k• k = (k] k ) The problem of the determination of rate constants based on kinetic measurements is said to be an inverse problem (IP). Much attention is paid to the analysis and solution of such problems at present (see e.g. the review in [! ] ). The reason is the following feature of models of complex chemical reactions - as a rule, only a part of the klnetical curves

are

reactants fact,

that

experimentally

.q.nd products linear

of the reaction.

systems

investigated well enough,

specific

problems

in

measurable:

most

And

of differential

often

initial

in spite of

equations

(I)

the

are

the above mentioned feature creates

decision.

The

main

problem

is

a

[2], this problem has been investigated

in

nonunlqueness of solutions of the IP. Considered for the first time

in the work

578

works of various authors [I ,3-6]. The main conclusions,

following:

that were made Lu these workm are the

a) a shortage of kinetical data Is the main reason

of the occurrence of nonuniqueness solutions

can be separated

isolated solutions

unidentifiability)

(In terms and

nonuniqueness,

an

infinite

set

of

functional

c)

in the

combinations

of

or

solutions,

in a space of constants

or unldentiflabillty) ; some

set of

[3] - a global nonuniqueness

representing some hypersurface nonuniqueness

; b) a nonuniqueness of IP

into two types - a finite

(local

case of

constants

LP

may

happen to be able to be determined in a unique way. A number and a form of functional type

of

an

IP

combinations permits

nonunlqueness;

algorithms

permit

in a number

priori

understood

d)

of

the

to determine a

developed

cases

methods

to carry

out

and

an a

priori analysis of models with respect to IP nonuniqueness. is

in

the

sense,

that

nonunlqueness before numerical calculations. IP nonuniqueness

unreasonable

can result

conclusions

in incorrect

we

can

show

A

the

And as far as an

and technologically

[7] and the numerical solution of an

IP is a long and complicated procedure,

thus the necessity of

an a priori analysis is clear. The method, considered in this paper, permits to solve a uniqueness problem of an IP and it also transforms the initial model to a form convenient for the estimation of functional combinatlons.

The

algorithm

has

a

simple

realization

computer algebra systems (such as, e.g. REDUCE [8] ).

in

2. A I ~ R I T ~ Write the initial model (I) In the form:

(2) (3) (4.)

x=Ax+By

y=

Cx+Dy

x(O) =x o, y(O) = Yo"

where

x

reactants,

=

(x I,...,

xn)

Y = (Yl ' ' ' "

are

measurable

concentrations

of

Ym ) are nonmeasurable concentrations;

579

the matrices

A, B, C and D depend on the Fate

only and have the dimensions respectively.

n : n,

n , m,

constants

m - n and

k•

m * m

We consider the ideal situation, i.e. we assume,

t~hat the measurable variables x are given without experimental

ezrors. In other words, the question about principal

wlth

the

structure

of

the

initial

model)

nonuniqueness of IP solutions is analyzed. We

make

in

system

transfonmat ions: (E

d_

(2),

(3)

the

(connected

uniqueness

following

or

matrix

A)x = By

(5)

Cx = (E

d_ dt

(6)

where E

and E

n dt

m

the

n

operator

determinant of

D)y m

are identity matrices. Let L be the adJolnt to (E d ~dt

_

D);

det(E d ~t

the operator (E d _ malt

_

D)

be

the

formal

D). Multiplying system (5)

by det(E d _ D) and the system (6) from the left on L and on B mdt we obtain det(E

m

d

dt

_ D)(E

n

BLCx = BL (E d _ mdt

d__ _ A)x = det(E dt

D) y

m

d_ dt

D)By

(7) (8)

where L is an operator to be de~ine below. With regard to the ~ormula E detH = H L [9], where H is a n o n s i ~ a r

L is its adJolnt matrix, we have L(E

m

d_

D) = det(E

dt

m

d_ dt

matrix, and

D)

and, therefore det(E X(O)

m

d

dt

_ D)(E

= X 0,...,

n

d_ _ A)x = BLCx dt

X Cm)(O)

= Xo,m.

(9) (10)

580 det(E

d____ D) is a linear differential operator of m-th dt order, L is a matrix of operators of (m-!)-th order, therefore Eq. (9) represents the system of n differential equations of (m+1)-th order of the measurable variables x and their derivatives only. The initial data (IO) are calculated by differentiation of system (2) and substitution of right hand sides (3) instead of y. In a real situation the order of the majority of equations (9) is less than (m + I ). Write the system (9) in the form m

n

qiJ

=

XOi v

-7. 7. a• k x j (k) J=Ik=O (V)(O) , i = I

~•

Xi

=

,''"

(11)

0

,n" v = I ,

,---

(12)

,m.

where qis' Pl ~ N, 0 <~ qij< Pi ~ m+1 ; aij k = (~iJk(k) are the functions of rate constants. System ( 1 1 ) can be considered

as

n

linear

algebraic

non-homogeneous equations for alj k. The initial data (I2) are dependent on k and y(O) and they can give additional equations for the determination of functional combinations, provided that the measurable variables x are exactly given and the initial conditions for y• are kmown. In the case, when the initial data for unmeasurable variables are unka~own, they should be added to the vector of unknown constants and in this case we shall consider the problem of rate constants and unknown initial data estimation as the IP to be investigated. Now carry out an analysis of nonuniqueness of the l_P, using the system (11),(12). Suppose, that we have obtained r (r = n n

+ +j=lqijT,)

polynomlal equ,atlons on k•

aiJk(k) = Cij k

(13) I

where Cljk iS determlned by x, x,... ,x (m+ I ) The system (13) is a finite set F of polynomials

in mary

581

variables.

The number of equations

(I3) can be more or less

than the number of constants. Obviously, the problem of rate constants determination has an infinite set of solutions when r s (a number of functional combinations is less than a constants number ). Let r > s. A part of equations in (12) can

be of in in

functionally dependent on other ones. Therefore, the problem determination of functionally independent polynomials of k i the set F and, further, the determination of exact solutions F arises. To s o l v e t h e s e p r o b l e m s t h e r e s u l t a n t m e t h o d [10] o r t h e

Groebner

basis

method

second method is realization

[9]

more

in several

According

to

[8],

I over the r i n g

be

used. as

far

computer algebraic the

polynomials of n variables ideal

can

convenient

Groebner

It

appears

as

it

has

that

the

numerical

systems.

basis

for

a

system

is the system of generators

o f p o l y n o m i a l s KEx 1

of

G of the

. x ] concerning to

some variables ordering. The method of Groebner basis permits not only to show the functional dependence of polynomials of a given set, but also to find exact solutions of the set F . S. EXAMPLW~ Consider the application of the algoritD~n to concrete systems, arisir~ in modelling first-order reactions. The first example represents the reaction, corresponding to the scheme (the example is t~ken from [11 ] ): A I --,

A2

;

A 2 -~

A3

+ A4

;

A I -~

A3 + Aa

.

(A)

For example, the reaction of ethylene c.hlorlnation under the condition

scheme:

of significant

excess of chlorine

p r o c e e d s on

such a

C12

1

~ C2H~C12

'

-~

!

c2H. --!! ci 2

I

C2H SCI+HCI

(B)

582

Let +.h~ " concentration oZ HO! be measured. We denote it by x. The concentration Calla and C~H~Cl~ are denoted by Yl and Y2" The mathematical model with the preceding notation has the form

y~ = - ( k ~

+ ~1

y~,

:'a = t~y~ - k2y 2,

(14)

x ( o ) = o, y~ (o) = 1, y..._,(o) = o.

(15)

According to (2)- (3) we introduce the matrices A

=

O,

B

=

Then we f lnd

l~

~-~

,

C =

d de t( E - - -

o

,

D

=

d D ) = ( --+

k~

-~j

d k I + k3)(--+

.

i%). We

needn't to find the matrix L, because the right hand side (9) is equal to 0 because C is zero matrix. Then (9)- (!0) can be rewritten in the form" .

,-

I

i n

m

(16)

z(o) = o,

x ( o ) : ~ y ~ (o) + ~ y 2 ( o )

: ~,

i i o ) :(k,k~ - ~ ( k ; + %))y~ (o) - ~ k1~ - %(k 1+%)"

y~(o) : (17)

Equation (I 6) together with the initial data (17) permits to flnd four functions of rate constants, provided that the exact solution x is given

% = (k, + %)t%, %=~., (18) A Groebner basis,

constructed for the system of polynomials

583

(18) shows the existence of two solutions, because for the constant F~ one of the po!~rnomials in the basis .has the form 2

(19)

i.e. a global nonuniqueness of the IP solutions occurs. Consider this example more carefully from the point of view

of

initial data representation. We assumed above, that the initial data y(O) had been given, and that has allowed us to use the set of initial data (17) to determine the (~i~k values. Let the Initlal data of the variables y• be unknown. Then the system (18), consisting of four equations will contain five tmFmown values k• i=I ,...,S and Yl (0), y2(O), that, obviously, results in an infinite set

of IP solutions. The second example gives general results for a wide class of reactions. Consider the consecutive reaction, in which (m + 2) reactants participate:

k, x~-~

~ ~ km kin,, Y~-~Y~-~...-~ Ym -4 X~

(C)

The corresponding mathematical model has the form X 1 = -kIx I

x~ = ~ + l Y m Yl = k I X 1 -

(21)

k2Yl

x I(0) = I , xe(O) = YI(O) = ...= Ym(O) = O. Write the eliminate

system the y•

(21) in the form of (2)-(4) and then variables. The system of equations,

corresponding to (9) for this example is

584

X I = -kIZ I

Cm~+ c ~ + 2

(m-1)

.

+ . km+ ., ),~.

+ i~k~ .... km+ , ~

+.

~%+

(m-2)

.

+ k m k m + 1 ~,~

: k,k~...km+ ,~,

+

....

(22)

From equations (22) it is possible to determine the following functions of constants" C~I = k I

%

: ~~+

...+ km km + l

% , 2 = % k2"'" ~ + ~ am+ 2 is expressed by a I ~ am+ I

(2a) and can be discarded.

functions aa,... 'am+ I are elementary

symmetric polynomials

The of

k• and the calculation of the Groebner basis for this set results in an equation of m-th order relatively any constant. The expressions for x92(0),..., x aC m - 1 IO) also contain only symmetric polynomials. In this case the IP has ml solutions. And finallly the example, which is the scheme of many experiments (23), is considered

XI

(2S)

--* k_ I

The mathematical model has the form X I =-kIX

I + k_ix 2

x 2 : klx I - (k_1 + I<21X 2 X 3 = k2x 2

(24)

x I (0) = X~o, x2(O) = Xao, x3(O) = Xao.

(25)

We consider three schemes of such an experiment- a) the concentration x~ is measured; b) the concentrations x~ and x 3

585

measured; c) the concentration x 3 is measured. For each case we consider various Luitlal data. Omitting all intermediate calculations , we get the table, where for each scheme of experiment the number of solutions of the IP has been shown (Tab. I). are

Table 1

Measured

X1

x3

Inltlal

Combination of

Number of IP

data a r e

constants

solutions

given

%=

not given

C[1 =

given

kl+ k_l+ k~,

unique

k I+ k_

infinite

%= %~, %= k~. 1+

k2 '

az= k~ k~. Depending on initial data

I) at= k1+ k i+ k~, ao= k I k~r t._

Inflnlte set

2 ) at= k I+ k_1+ k 2,

unique

not given

(z2 (o)~o) a~ = k~ + k_ ~+ ~ ,

infinite

given

at= k I- k_1- k 2,

unique

(z2 (o)=o)

%= k~,

"1

" X3

set

not given

%= ~.

%= k ~ .

%= k~l,,~, %= ~+k~. " "

set

unique

By this way, having only the scheme of a reaction, it is possible to plan an experiment so that the IP will have a unique solution. Similar tables can be costructed for any first order react ion. Note, that the

increase

of

a

reactants

number

and

a

586

complication o$ reactions (convertible stages, cycles) results only in increasing of analytical calculations. The considered above eliminations technique remains the same . Nevertheless, the analysis o$ such situations requires the special investigation, as it can result in the combinations o~ the rate constants, admittir~ a specific chemical interpretation. REFERenCES I.

A.G. Pogorelov, An inverse problem of non-stationazw chemical kinetics, Science, Moskow, 1988. 2. M.V. Klibanov, S.I.Spivak, V.I.Timoshenko, M.G. SllnF~o, Dokl. Akad. Nauk SSSR, 208 (1975 ) 1387. 3. S . I . S p l v a k , V.G.Gorskll, Dokl. Akad. Nauk, 257(1981 )412. 4. V.N.Lukashenok, Teoret. Osnov. Khim. Techu., 14 (1980)86. 5. S.VaJda,, Ma~y Kern. Foly, 83 (I977 )74. 6. P.Erdy, J.Toth, Mathematical models of chemical reactions, Princeton University Press, Princeton, 1989. 7. M.G.Slinko, Kinetlka i Katallz, 22(1981 )5. 8. Computer algebra. Symbolic and Algebraic Computation, Edited by B.Buchberger et al, Sprlnger-Verlag, Wien, New York, 1982. 9. V.P.Maslov, Operator methods, Science, Moskow, 1973. IO. B.L.Van tier Waerden, A l ~ b r a I, Springer-Verlag, Berlin, 1967. 11. L.V.Bistrov, V.G. Gorskii, S.I. Splvak, Teoret. i eksperiment. Khlm., 6 (I985 )701.

587

[i!i~~~i!i!i~!~i~i l ~!i l ~i!ii~ i ~ i

Adrover A. AI-Haddad A.A. Aldea N. Almasan V. Asadullin R.M. Bechara R. Becker S.J. Bergault I. Blekkan E.A. Biiek A. Broadbelt L.J. Bunimovich G.A. Callejas M.A. Carb6 E. Carrillo Le Roux G.A. Cechini J. Che M. Chellappa A.S. Ciambelli P. Clausen B.S. Comotti P. Cunill F. Czernicki M. Datta R. Decamp T. Delmas H. Delmon B. Deo G. Di Benedetto A. Doepper R. Dooling D.J. Drinkenburg A.A.H. Dubel M. Dubuis S. Eigenberger G. Emig G. Ertl G. Fan Y. Fatah N. Fite C.

Frauhammer J. Froment G.F. Fuangfoo S. Fujimoto K. Futami S. Gabarain L.

~ H I I ~

i~ S ! !

iil~ii ~ i

~ ~ i E ~ i i i ! i l ilii~i~i!iii;iiii!iii~!~iiiii!;iiil;il

241 495 547 547 577 263 285 571 193 419 251,341 141 565 565 571 459 91 217 175, 553 121 429 541 263 559 351 571 203 305 175 295, 469 251 439 227 469 273 479 389 325 263 541

273 159 217 325 491 459

588 Gaeumann T. Gaigneaux E. Garufi E. Gelten R.J. Giona M. Gleaves J.T. Golay S. Gosiewski K. Hanssen K.F. Haure P. Hern~ndez A. Hinrichsen O. Hoebink J.H.B.J. H~jlund Nielsen P.E. Holmen A. Hosoya T. Iborra M. Imai H. Izquierdo J.F. Jensen K. Joulia X. Juskelis M.V. King D.A. King T.S. Kitchaiya P. Kodde A.J. Komiyama M. Lambert J.-F. Lazar M. Lazman M.Z. Lemos F. Li X. Li Y.-W. Lorenzi M. Mach~'n I. Makkee M. Marengo S. Marginean P. Marin G.B. Marti'nez M.T. Mathew J. Matros Y.Sh. Mirodatos C. Molga E.J. Moulijn J.A. Muhler M. Nakamura I. Neurock M. Nievergeid A.J.L. Nijhuis T.A. N~rskov J.K. Ould Mohamed Mahmoud M.L.

547 2O3 175,553 61 241 333 295 511 193 459 565 389 449 121 193 491 541 491 541 559 571 305 79 315 559 419 185 91 547 371 529,535 523 203 469 517 361 429 547 449 565 495 141 351 379 361 111,389 325 3 449 361 121 263

589 Ovesen C.V. Palma C. Pantazidis A. Phanawadee P. Piepers H.W. Pietrzyk S. Pinheiro C.I.C. Pirone R. Prasad S.D. Pruski M. Rami'rez de Agudelo M.M. Ram6a Ribeiro F. Rekoske J.E. Rembeczky T. Renken A. Romero T. Rosowski F. Ruiz P. Rupprechter G. Russo G. Santos Silva I. Savargoankar N. Scappatura S. Schanke D. Schmidt L.D. Schuurman Y. Seiler H. Singh S. Somorjai G.A. Sousa Lobo J. Sowerby B. Spivak S.I. Strots V.O. Tejero J. Tops~e H. Tops~e N.-Y. T0rnqvist E. Tserpe E. Uner D.O. v. Selow E.R. v.d. Runstraat A. van Langeveld A.D. van Santen R.A. Vasconi M. Veser G. Viswanath D.S. Wachs I.E. Waugh K.C. Weckhuysen B.M. Westerterp K.R. Wolfrath O. Xie K.

121 535 351 333 439 263 529 175,553 227 315 517 529,535 341 263 295,469 517 111,389 203 35 175,553 535 315 429 193 273 333,351 479 159 35 535 285 577 141 541 121 121 121 401 315 449 61 361 61 429 273 217 305 401 305 379 295 523

590

Xu Y.-D. Yablonskii G.S. Yang Y. Zhang A. Zhang T.

351 333, 371 523 325 559

591 S T U D I E S IN SURFACE SCIENCE A N D CATALYSIS Advisory Editors: B. Delmon, Universite Catholique de Louvain, Louvain-la-Neuve, Belgium J.T. Yates, University of Pittsburgh, Pittsburgh, PA, U.S.A. Volume

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Preparation of Catalysts I.Scientific Bases for the Preparation of Heterogeneous Catalysts. Proceedings of the First International Symposium, Brussels, October 14-17,1975 edited by B. Delmon, P.A. Jacobs and G. Poncelet The Control of the Reactivity of Solids. A Critical Survey of the Factors that Influence the Reactivity of Solids, with Special Emphasis on the Control of the Chemical Processes in Relation to Practical Applications by V.V. Boldyrev, M. Bulens and B. Delmon Preparation of Catalysts I1. Scientific Bases for the Preparation of Heterogeneous Catalysts. Proceedings of the Second International Symposium, Louvain-la-Neuve, September 4-7, 1978 edited by B. Delmon, P. Grange, P. Jacobs and G. Poncelet Growth and Properties of Metal Clusters. Applications to Catalysis and the Photographic Process. Proceedings ofthe 32nd International Meeting ofthe Societe de Chimie Physique, Villeurbanne, September 24-28, 1979 edited by J. Bourdon Catalysis by Zeolites. Proceedings of an International Symposium, Ecully (Lyon), September 9-11, 1980 edited by B. Imelik, C. Naccache, Y. Ben Taarit, J.C. Vedrine, G. Coudurier and H. Praliaud Catalyst Deactivation. Proceedings of an International Symposium, Antwerp, October 13-15,1980 edited by B. Delmon and G.E Froment New Horizons in Catalysis. Proceedings of the 7th International Congress on Catalysis, Tokyo, June 30-July4, 1980. Parts A and B edited by T. Seiyama and K. Tanabe Catalysis by Supported Complexes by Yu.l. Yermakov, B.N. Kuznetsov and V.A. Zakharov Physics of Solid Surfaces. Proceedings of a Symposium, Bechyhe, September 29-October 3,1980 edited by M. L~iznieka Adsorption at the Gas-Solid and Liquid-Solid Interface. Proceedings of an International Symposium, Aix-en-Provence, September 21-23, 1981 edited by J. Rouquerol and K.S.W. Sing Metal-Support and Metal-Additive Effects in Catalysis. Proceedings of an International Symposium, Ecully (Lyon), September 14-16, 1982 edited by B. Imelik, C. Naccache, G. Coudurier, H. Praliaud, P. Meriaudeau, P. Gallezot, G.A. Martin and J.C. Vedrine Metal Microstructures in Zeolites. Preparation - Properties- Applications. Proceedings of a Workshop, Bremen, September 22-24, 1982 edited by P.A. Jacobs, N.I. Jaeger, P.Jin3 and G. Schulz-Ekloff Adsorption on Metal Surfaces. An Integrated Approach edited by J. Benard Vibrations at Surfaces. Proceedings of the Third International Conference, Asilomar, CA, September 1-4, 1982 edited by C.R. Brundle and H. Morawitz Heterogeneous Catalytic Reactions Involving Molecular Oxygen by G.I. Golodets

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Methane Conversion. Proceedings of a Symposium on the Production of Fuels and Chemicals from Natural Gas, Auckland, April 27-30, 1987 edited by D.M. Bibby, C.D. Chang, R.F. Howe and S. Yurchak Innovation in Zeolite Materials Science. Proceedings of an International Symposium, Nieuwpoort, September 13-17, 1987 edited by P.J. Grobet, W.J. Mortier, E.F. Vansant and G. Schulz-Ekloff Catalysis 1987. Proceedings ofthe 10th North American Meeting ofthe Catalysis Society, San Diego, CA, May 17-22, 1987 edited by J.W. Ward Characterization of Porous Solids. Proceedings of the IUPAC Symposium (COPS I), Bad Soden a. Ts., April 26-29,1987 edited by K.K. Unger, J. Rouquerol, K.S.W. Sing and H. Kral Physics of Solid Surfaces 1987. Proceedings of the Fourth Symposium on Surface Physics, Bechyne Castle, September 7-11, 1987 edited by J. Koukal Heterogeneous Catalysis and Fine Chemicals. Proceedings of an International Symposium, Poitiers, March 15-17, 1988 edited by M. Guisnet, J. Barrault, C. Bouchoule, D. Duprez, C. Montassier and G. Perot Laboratory Studies of Heterogeneous Catalytic Processes by E.G. Christoffel, revised and edited by Z. Paal Catalytic Processes under Unsteady-State Conditions by Yu. Sh, Matros Successful Design of Catalysts. Future Requirements and Development. Proceedings ofthe Worldwide Catalysis Seminars, July, 1988, on the Occasion of the 30th Anniversary of the Catalysis Society of Japan edited byT. Inui Transition Metal Oxides. Surface Chemistry and Catalysis by H.H. Kung Zeolites as Catalysts, Sorbents and Detergent Builders. Applications and Innovations. Proceedings of an International Symposium, WL~rzburg, September 4-8,1988 edited by H.G. Karge and J. Weitkamp Photochemistry on Solid Surfaces edited by M. Anpo and T. Matsuura Structure and Reactivity of Surfaces. Proceedings of a European Conference, Trieste, September 13-16, 1988 edited by C. Morterra, A. Zecchina and G. Costa Zeolites: Facts, Figures, Future. Proceedings of the 8th International Zeolite Conference, Amsterdam, July 10-14, 1989. Parts A and B edited by P.A. Jacobs and R.A. van Santen Hydrotreating Catalysts. Preparation, Characterization and Performance. Proceedings of the Annual International AIChE Meeting, Washington, DC, November 27-December 2, 1988 edited by M.L. Occelli and R.G. Anthony New Solid Acids and Bases. Their Catalytic Properties by K. Tanabe, M. Misono, Y. Ono and H. Hattori Recent Advances in Zeolite Science. Proceedings of the 1989 Meeting of the British Zeolite Association, Cambridge, April 17-19, 1989 edited by J. Klinowsky and P.J. Barrie Catalyst in Petroleum Refining 1989. Proceedings of the First International Conference on Catalysts in Petroleum Refining, Kuwait, March 5-8, 1989 edited by D.L. Trimm, S. Akashah, M. Absi-Halabi and A. Bishara Future Opportunities in Catalytic and Separation Technology edited by M. Misono, Y. Moro-oka and S. Kimura

594 Volume 55

New Developments in Selective Oxidation. Proceedings of an International Symposium, Rimini, Italy, September 18-22, 1989 edited by G. Centi and F. Trifiro Volume 56 Olefin Polymerization Catalysts. Proceedings of the International Symposium 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 Volume 58 Introduction to Zeolite Science and Practice edited by H. van Bekkum, E.M. Flanigen and J.C. Jansen Volume 59 Heterogeneous Catalysis and Fine Chemicals II. Proceedings of the 2nd International Symposium, Poitiers, October 2-6, 1990 edited by M. Guisnet, J. Barrault, C. Bouchoule, D. Duprez, G. Perot, R. Maurel and C. Montassier Volume 60 Chemistry of Microporous Crystals. Proceedings of the International Symposium on Chemistry of Microporous Crystals, Tokyo, June 26-29, 1990 edited by T. Inui, S. Namba and T. Tatsumi Volume 61 Natural Gas Conversion. Proceedings of the Symposium on Natural Gas Conversion, Oslo, August 12-17, 1990 edited by A. Holmen, K.-J. Jens and S. Kolboe Volume 62 Characterization of Porous Solids I1. Proceedings of the IUPAC Symposium (COPS II), Alicante, May 6-9, 1990 edited by F. Rodriguez-Reinoso, J. Rouquerol, K.S.W. Sing and K.K. Unger Preparation of Catalysts V. Scientific Bases for the Preparation of Heterogeneous Volume 63 Catalysts. Proceedings of the Fifth International Symposium, Louvain-la-Neuve, September 3-6, 1990 edited by G. Poncelet, P.A. Jacobs, P. Grange and B. Delmon Volume 64 New Trends in CO Activation edited by L. Guczi Volume 65 Catalysis and Adsorption by Zeolites. Proceedings of ZEOCAT 90, Leipzig, August 20-23, 1990 edited by G. Clhlmann, 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, Balatonf~red, September 10-14, 1990 edited by L.I. Simandi Volume 67 Structure-Activity and Selectivity Relationships in Heterogeneous Catalysis. 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 Volume 68 Catalyst Deactivation 1991. Proceedings of the Fifth International Symposium, Evanston, IL, June 24-26, 1991 edited by C.H. Bartholomew and J.B. Butt Volume 69 Zeolite Chemistry and Catalysis. Proceedings of an International Symposium, 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

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Volume 73 Volume 74 Volume 75 Volume 76 Volume 77

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Volume 79 Volume80 Volume81 Volume82

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Volume85 Volume86 Volume 87

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 R Ruiz and B. Delmon Progress in Catalysis. Proceedings of the 12th Canadian Symposium on Catalysis, Banff, Alberta, Canada, May 25-28, 1992 edited by K.J. Smith and E.C. Sanford Angle-Resolved Photoemission. Theory and Current Applications edited by S.D. Kevan New Frontiers in Catalysis, Parts A-C. Proceedings of the 10th International Congress on Catalysis, Budapest, Hungary, 19-24 July, 1992 edited by L. Guczi, E Solymosi and P.Tetenyi Fluid Catalytic Cracking: Science and Technology edited by J.S. Magee and M.M. Mitchell, Jr. New Aspects of Spillover Effect in Catalysis. For Development of Highly Active Catalysts. Proceedings 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.E Howe New Developments in Selective Oxidation II. Proceedings of the Second World Congress and Fourth European Workshop Meeting, Benalmadena, Spain, September 20-24, 1993 edited by V. Cortes Corberan and S. Vic Bellon Zeolites and Microporous Crystals. Proceedings of the International Symposium on Zeolites and Microporous Crystals, Nagoya, Japan, August 22-25, 1993 edited by T. Hattori and T. Yashima Zeolites and Related Microporous Materials: State of the Art 1994. Proceedings 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 II1. 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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Volume 101 Volume 102 Volume 103 Volume 104 Volume 105

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 ofthe International Symposium on Acid-Base Catalysis II, Sapporo, Japan, December 2-4, 1993 edited by H. Hattori, M. Misono and Y. Ono Preparation of Catalysts VI. Scientific Bases for the Preparation of Heterogeneous Catalysts. Proceedings of the Sixth International Symposium, Louvain-La-Neuve, September 5-8, 1994 edited by G. Poncelet, J. Martens, B. Delmon, P.A. Jacobs and P. Grange Science and Technology in Catalysis 1994. Proceedings of the Second Tokyo Conference on Advanced Catalytic Science and Technology, Tokyo, August 21-26, 1994 edited by Y. Izumi, H. Arai and M. Iwamoto Characterization and Chemical Modification of the Silica Surface by E.F. Vansant, P. Van Der Voort and K.C. Vrancken Catalysis by Microporous Materials. Proceedings of ZEOCAT'95, Szombathely, Hungary, July 9-13, 1995 edited by H.K. Beyer, H.G.Karge, I. Kiricsi and J.B. Nagy Catalysis by Metals and Alloys by V. Ponec and G.C. Bond Catalysis and Automotive Pollution Control II1. Proceedings of the Third International Symposium (CAPoC3), Brussels, Belgium, April 20-22, 1994 edited by A. 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. Da,browski and V.A. Tertykh Catalysts in Petroleum Refining and Petrochemical Industries 1995. Proceedings of the 2nd International Conference on Catalysts in Petroleum Refining and Petrochemical Industries, Kuwait, April 22-26, 1995 edited by M. Absi-Halabi, J. Beshara, H. Qabazard and A. Stanislaus 1lth International Congress on Catalysis - 40th Anniversary. Proceedings ofthe 1lth ICC, Baltimore, MD, USA, June 30-July 5, 1996 edited by J. W. Hightower, W.N. Delgass, E. Iglesia and A.T. Bell Recent Advances and New Horizons in Zeolite Science and Technology edited by H. Chon, S.I. Woo and S.-E. Park Semiconductor Nanoclusters - Physical, Chemical, and Catalytic Aspects edited by P.V. Kamat and D. Meisel Equilibria and Dynamics of Gas Adsorption on Heterogeneous Solid Surfaces edited by W. Rudzinski, W.A. Steele and G. Zgrablich Progress in Zeolite and Microporous Materials Proceedings 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 ofthe 1st International Symposium / 6th European Workshop, Oostende, Belgium, February 17-19, 1997 edited by G.E Froment, B. Delmon and P. Grange Natural Gas Conversion IV Proceedings of the 4th International Natural Gas Conversion Symposium, Kruger Park, South Africa, November 19-23, 1995 edited by M. de Pontes, R.L. Espinoza, C.P. Nicolaides, J.H. Scholtz and M.S. Scurrell Heterogeneous Catalysis and Fine Chemicals IV Proceedings of the 4th International Symposium on Heterogeneous Catalysis and Fine Chemicals, Basel, Switzerland, September 8-12, 1996 edited by H.U. Blaser, A. Balker and R. Prins Dynamics of Surfaces and Reaction Kinetics in Heterogeneous Catalysis. Proceedings ofthe International Symposium, Antwerp, Belgium, September 15-17, 1997 edited by G.E Froment and K.C. Waugh Third World Congress on Oxidation Catalysis. Proceedings of the Third World Congress on Oxidation Catalysis, San Diego, CA, U.S.A. 21-26, September 1997 edited by R.K. Grasselli, S.T. Oyama, A.M. Gaffney and J.E. Lyons Catalyst Deactivation 1997. Proceedings ofthe 7th International Symposium, Cancun, Mexico, October 5-8, 1997 edited by C.H. Bartholomew and G.A. Fuentes

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