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

15 downloads 854 Views 28MB Size Report

This content was uploaded by our users and we assume good faith they have the permission to share this book. If you own the copyright to this book and it is wrongfully on our website, we offer a simple DMCA procedure to remove your content from our site. Start by pressing the button below!
Report copyright / DMCA form

DOWNLOAD PDF

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

This Page Intentionally Left Blank

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

iiiii ii !iiili!iil!ii!iili!iiiiii! i i! ii ii

iiii!ii ii !i!iili!iiii!i!i!iiii!ii!l

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


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]


~;ii i iiil

i i ill iiiii!i iiiiiii ilii i ii i i i

~i~i!i::i::::!i!iii::ii~i::!::::ii::iii::::g!::::ii!::::ii::ii::~i::i::i!gii::i!i!i::~i!~::~!i::g::i::giii!!iii!::i::::igi::i::::~::ii::!i::ii::::i~::ii::ii!i::!i::!i::i~ iii!~ii~i!iiiiii~i!i!iii~i~i~!~i~i~!!!~ii!~ii!iiiii!!!i!!i~i~i~i~iii!iii!iiii!!iiiii!~!!!iii!~ii!:!:~i~ii!:!i!i:~!!ii~!~!i!i!i:!

i ii :iNiiill :

i il !imi iiiiiillii

!i iii! ii iNiiiii iiiiiii

iii iiiiiiiiiimiiiiii i iii

ill iliiii~ i! !

ii:iii:iii!i,ii!ii ii !iii'!

i!~i~i!i~ii!!iiiiiii:~ii~!~i~i~i~i~i~i~i~i~ii~::.~:~:~i:::~:i~iiii:i:~::!~i~i!:!:!;i:i!ii:i!!::ii:~i!i:i

ii

iii•i•ii•i•!iiiiiii!ii•i•iiiii•i•ii!iii!i!ii•iii•ii!•!i!i!i!i•i!i••i•ii•••!!!•!ii!iii•i!i•i•!ii•ii!!!ii!!i!!i!i!!!!i!!•i•!i!iii!i•ii!

i I!i•i!i!iii•i!••ii!ii!ii!iiiiiiiiiii!i!i!iiiii!iii!iiiiiii•!iii!ii!iiiii•i!i!i!i!i•iiiii!i!iiiiiii•i•i•i!iii!iii!iii!i•i•i!i!i!ii!ii

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


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.

"4

0 0 ,~11r

.."'-'..Donation

.-" :.

..

.

.

-..

"-44. tl

"" Backdonation

i~

Pauli Repulsion

Metal Surface

Ad-Molecule

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

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.


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.



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

[

ML/s

""-.

"...

..... ...

.......

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

,

I

,

,

,

,

_

I

,

,

,

,

I

,

,

L

,

_

r,.., / MLs-' x 5.0

r.,.o

A

,

,

,

I

,

,

,

~

/ 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).



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 .

References (1)

Che, M. In Proc. lOth Intern. Cong. Catal, Budapest 1992.; L. Guczi et al., eds; Elsevier, Amsterdam 1993; pp 31-68.

(2)

Schwarz, J. A.; Contescu, C.; Contescu, A. Chem. Rev. 1995, 95, 477-510.

(3)

Zamaraev, K. I. Topics in Catalysis 1996, 3, 1-76.

(4)

Dyrek, K.; Che, M. Chem. Rev. 1997, 97, 305-331.

(5)

Che, M.; Bennett, C. O. Adv. Catat. 1989, 36, 55-172.

(6)

Zecchina, A.; Spoto, G.; Coluccia, S.; Guglielminotti, E. J. Phys. Chem. 1984, 88, 2575-2581.

(7)

Zecchina, A.; Spoto, G.; Borello, E.; Giamello, E. J. Phys. Chem. 1984, 88, 2582-2587.

(8)

Kn6zinger, H.; Ratnasamy, P. Catal. Rev. Sci. Eng. 1978, 17, 31.

(9)

Tsyganenko, A. A.; Mardilovich, P. P. J. Chem. Soc. Faraday Trans. 1996, 92, 48434852.

(10)

Garofalini, S. H. Diffus. Defect Data 1987, 53-54, 21.

(11)

Contescu, C.; Jagiello, J.; Schwarz, J. A. Langmuir 1993, 9, 1754-1765.

(12)

Contescu, C.; Hu, J.; Schwarz, J. A. ]. Chem. Soc. Faraday Trans. 1993,89, 4091-4099.

(13)

Atsumian, R. D.; Sasaki, M.; Yasunaga, T.; Schelly, Z. A. 1. Phys. Chem. 1981, 85, 3832-3835.

(14)

Toney, M. F.; Howard, J. N.; Richer, J.; Borges, G. L.; Gordon, J. G.; Melroy, O. R.; Wiesler, D. G.; Yee, D.; Sorensen, L. B. Surf. Sci. 1995, 335, 326-332.

(15)

Danielewicz-Ferchmin, I.; Ferchmin, A. R. J. Phys. Chem. 1996, 100, 17281-17286.

(16)

Danielewicz-Ferchmin, I. J. Phys. Chem. 1995, 99, 5658-5665.

(17)

Bassetti, V.; Burlamacchi, L.; Martini, G. J. Am. Chem. Soc. 1979, 101, 5471-547.

(18)

Benjamin, I. Chem. Rev. 1996, 96, 1449-1475.

109 (19)

Luthra, N. P.; Cheng, W.-C. J. CataI. 1987,107, 154-160.

(20)

Sarrazin, P.; Mouchel, B.; Kasztelan, S. J. Phys. Chem. 1989, 93, 904-908.

(21)

Sarrazin, P.; Mouchel, B.; Kasztelan, S. J. Phys. Chem. 1991, 95, 7405-7408.

(22)

Ong, S. W.; Zhao, X. L.; Eisentahl, K. B. Chem. Phys. Left. 1992, 191, 327-335.

(23)

Bockris, J. O.; Conway, B. E.; E. Yeager, eds. Comprehensive treatise of electrochemistry - Vol.l: The double layer; Plenum Press ed.; New-York, 1980.

(24)

Brunelle, J. P. Pure Appl. Chem. 1978, 50, 1211-1229.

(25)

Colquhoun, H. M.; Stoddart, J. F.; Williams, D. J. Angew. Chem. Intnl. Ed., 1983, 25, 487-582.

(26)

Bargar,J. R.; Towle, S. N.; G.E. Brown, J.; Parks, G. A. Geochim. Cosmochim. Acta 1996, 60, 3541-3547.

(27)

Setoyama, T.; Che, M.; unpublished results.

(28)

Lambert, J.-F.; Hoogland, M.; Che, M.; submitted to Inorg. Chem..

(29)

Chen, Y.; Hyldtoft, J.; Jacobsen, C. J. H.; Christensen, D. H.; Nielsen, O.F. Spectrochim. Acta 1994, 50A, 1879-1889.

(30)

Lepetit, C.; Che, M. J. Mol. Catal. 1996, 100, 147-160.

(31)

Fokkink, L. G. J.; Keizer, A. D.; Lyklema, J. J. Coll. Interj. Sci. 1990, 135, 118-131.

(32)

Yiacoumi, S.; Tien, C. J. Coll. Interfi Sci. 1995, 175, 333-346.

(33)

Hachiya, K.; Sasaki, M.; Saruta, Y.; Mikami, N.; Yasunaga, T. J. Phys. Chem. 1984, 88, 23-27.

(34)

Hachiya, K.; Sasaki, M.; Ikeda, T.; Mikami, N.; Yasunaga, T. J. Phys. Chem. 1984, 88, 27-31.

(35)

Johnson, R. E. J. J. Coil. Interf. Sci. 1984, 100, 540-554.

(36)

Habib, M. A.; Bockris, J. O. In Comprehensive treatise of electrochemistry VoI.l: The double layer; P. Press, Ed.; New-York, 1980; pp 138 sq.

(37)

K(~kicheff,P.; Marcelja, S.; Senden, T. J.; Shubin, V. E. J. Chem. Phys. 1993, 99, 60986113.

(38)

Spanos, N.; Vordonis, L.; Kordulis, C.; Lycourghiotis, A. J. Catal. 1990,124, 301-314.

(39)

Le Bel de Penguilly, V.; Lambert, J.-F.; Che, M.; unpublished results.

(40)

Spielbauer, D.; Zeilinger, H.; KnOzinger, H. Langmuir 1993, 9, 460-466.

110 (41)

Ludwig, C.; Schindler, P. W. J. Coll. Interj'. Sci. 1995, 169, 284-290.

(42)

Bergaoui, L.; Lambert, J.-F.; Suquet, H.; Che, M. J. Phys. Chem. 1995, 99, 2155-2161.

(43)

Olivier, D.; Bonneviot, L.; Cai, F. X.; Che, M.; Gihr, P.; Kermarec, M.; LepetitPourcelot, C.; Morin, B. Bull. Soc. Chim. Fr. 1985, 370-380.

(44)

Bonneviot, L.; Cai, F. X.; Che, M.; Kermarec, M.; Legendre, O.; Lepetit, C.; Olivier, D. J. Phys. Chem. 1987, 91, 5912-5921.

(45)

Garrot, J.-M. Ph.D. Thesis, Universit6 Pierre et Marie Curie (Paris VI), 1996.

(46)

Clause, O.; Kermarec, M.; Bonneviot, L.; Villain, F.; Che, M. I. Am. Chem. Soc. 1992, 114, 4709-4717.

(47)

Kermarec, M.; Carriat, J. Y.; Burattin, P.; Che, M.; Decarreau, A. J. Phys. Chem. 1994, 98, 12008-12017.

(48)

d'Espinose de la Caillerie, J. B.; Kermarec, M.; Clause, O. J. Am. Chem. Soc. 1995, 117, 11471-11481.

(49)

Scheidegger, A. M.; Lamble, G. M.; Sparks, D. L. J. Coll. Interf. Sci. 1997,186, 118-128.

(5o)

Stumm, W. Chemistry at the solid-water interface: processes at the mineral-water and particle-water interface in natural systems.; J. Wiley & Sons, New-York, 1992.

(5~)

Wehrli, B.; Wieland, E.; Furrer, G. Aquatic sciences 1990, 52, 1-114; quoted in (50).

(52)

d'Espinose de la Caillerie, J.-B.; Kermarec, M.; Clause, O. J. Phys. Chem. 1995, 99, 17273-17281.

(53)

Louis, C.; Cheng, Z. X.; Che, M. J. Phys. Chem. 1993, 97, 5703-5712.

(54)

Carrier, X.; Lambert, J.-F.; Che, M.; unpublished results.

(55)

Chisholm-Brause, C. J.; G.E. Brown, J.; Parks, G. A. Physica B 1989, 158, 646-648.

(56)

Chisholm-Brause, C. J.; O'Day, P. A.; G.E. Brown, J.; Parks, G. A. Nature 1990, 348, 528531.

(57)

Che, M.; Bonneviot, L. In Successful design of catalysts; ; T. Inui, ed.; Elsevier, Amsterdam, 1988; pp 147-157.

(58)

a. Burattin, P. Ph.D. Thesis, Universit6 Pierre et Marie Curie (Paris VI), 1994; b. Burattin, P.; Che, M.; Louis, C.; submitted to J. Phys. Chem.

(59)

Meyer, T. J.; Taube, H. In Comprehensive coordination chemistry; G. Wilkinson, ed.; Pergamon, Oxford, 1987; Vol. 1; pp 331sq.


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.


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

~ 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)



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] [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 ( 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).


i!!ii!i!i!ii!~i!!!i!i!ii!


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 [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.


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.



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!

....

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

:'~,

I

i

@

0 (o It_ n, (D LL 0 0 F 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)

~

ii~t

"

~

~ .g

~.~ 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 PAS 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
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.


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 '-0.8 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.


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.



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, 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:

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 ~

/ ....

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


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,' 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).



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 -.'
'-

-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



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 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



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.


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.


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

,

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

Chemical Kinetics and Reaction Dynamics

Chemical Kinetics and Reaction Dynamics

Chemical Kinetics and Reaction Dynamics

Catalysis and automotive pollution control IV: proceedings of the Fourth International Symposium (CAPoC4), Brussels, Belgium, April 9-11, 1997

Zeolite Chemistry and Catalysis Proceedings of an International Symposium Prague

Equilibria and Dynamics of Gas Adsorption on Heterogeneous Solid Surfaces (Studies in Surface Science and Catalysis)

Equilibria and Dynamics of Gas Adsorption on Heterogeneous Solid Surfaces (Studies in Surface Science and Catalysis)

Handbook of Asymmetric Heterogeneous Catalysis

Handbook of Asymmetric Heterogeneous Catalysis

Isotopes in Heterogeneous Catalysis

Isotopes In Heterogeneous Catalysis

Equilibria and Dynamics of Gas Adsorption on Heterogeneous Solid Surfaces

Scientific Bases for the Preparation of Heterogeneous Catalysts, Volume 175: Proceedings of the 10th International Symposium, Louvain-la-Neuve, Belgium, ... (Studies in Surface Science and Catalysis)

Theoretical Aspects of Heterogeneous Catalysis

Theoretical aspects of heterogeneous catalysis

Principles and Practice of Heterogeneous Catalysis

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

Heterogeneous Catalysis in Organic Chemistry

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

Mechanical Catalysis: Methods of Enzymatic, Homogeneous, and Heterogeneous Catalysis

Dynamics of fractal surfaces

Heterogeneous Catalysis Research Progress

Heterogeneous Catalysis in Industrial Practice

Mass Transfer in Heterogeneous Catalysis

Heterogeneous Catalysis in Organic Chemistry

Mechanisms in Homogeneous and Heterogeneous Epoxidation Catalysis

Mechanisms in Homogeneous and Heterogeneous Epoxidation Catalysis

Catalysis by Zeolites: International Symposium Proceedings (Studies in surface science and catalysis)

reaction kinetics and reactor design

Heterogeneous Catalysis and Fine Chemicals II: Proceedings of the 2nd International Symposium, Poitiers, October 2-5, 1990 (Studies in Surface Science and Catalysis)

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

Chemical Kinetics and Reaction Dynamics

Chemical Kinetics and Reaction Dynamics

Chemical Kinetics and Reaction Dynamics

Catalysis and automotive pollution control IV: proceedings of the Fourth International Symposium (CAPoC4), Brussels, Belgium, April 9-11, 1997

Zeolite Chemistry and Catalysis Proceedings of an International Symposium Prague

Equilibria and Dynamics of Gas Adsorption on Heterogeneous Solid Surfaces (Studies in Surface Science and Catalysis)

Equilibria and Dynamics of Gas Adsorption on Heterogeneous Solid Surfaces (Studies in Surface Science and Catalysis)

Handbook of Asymmetric Heterogeneous Catalysis

Handbook of Asymmetric Heterogeneous Catalysis

Isotopes in Heterogeneous Catalysis

Isotopes In Heterogeneous Catalysis

Equilibria and Dynamics of Gas Adsorption on Heterogeneous Solid Surfaces

Scientific Bases for the Preparation of Heterogeneous Catalysts, Volume 175: Proceedings of the 10th International Symposium, Louvain-la-Neuve, Belgium, ... (Studies in Surface Science and Catalysis)

Theoretical Aspects of Heterogeneous Catalysis

Theoretical aspects of heterogeneous catalysis

Principles and Practice of Heterogeneous Catalysis

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

Heterogeneous Catalysis in Organic Chemistry

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

Mechanical Catalysis: Methods of Enzymatic, Homogeneous, and Heterogeneous Catalysis

Dynamics of fractal surfaces

Heterogeneous Catalysis Research Progress

Heterogeneous Catalysis in Industrial Practice

Mass Transfer in Heterogeneous Catalysis

Heterogeneous Catalysis in Organic Chemistry

Mechanisms in Homogeneous and Heterogeneous Epoxidation Catalysis

Mechanisms in Homogeneous and Heterogeneous Epoxidation Catalysis

Catalysis by Zeolites: International Symposium Proceedings (Studies in surface science and catalysis)

reaction kinetics and reactor design

Heterogeneous Catalysis and Fine Chemicals II: Proceedings of the 2nd International Symposium, Poitiers, October 2-5, 1990 (Studies in Surface Science and Catalysis)

Recommend Documents