ADVANCES IN CATALYSIS VOLUME 39, Volume 39

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ADVANCES IN CATALYSIS VOLUME 39

Advisory Board

M. BOUDART Stanford, California

V. B. KAZANSKY Moscow, Russia

G. A . SOMORJAI Berkeley, California

G. ERTL BerlinlDahlem, Germany

A. OZAKI Tokyo, Japan

W. 0. HAAG Princeton, New Jersey

W. M. H. SACHTLER Evanston, Illinois

J. M. THOMAS London, U . K .

ADVANCES IN CATALYSIS VOLUME 39

Edited by D. D. ELEY The University Nottingham, England

HERMAN PINES Northwestern University Evanston, Illinois

PAULB . WEISZ

University of Pennsylvania Philadelphia, Pennsylvania

ACADEMIC PRESS, INC. Harcourt Brace & Company

San Diego New York Boston London Sydney Tokyo Toronto

This book is printed on acid-free paper.

@

Copyright 0 1993 by ACADEMIC PRESS, INC. All Rights Reserved. No part of this publication may be reproduced or transmitted in any form or by any means, electronic or mechanical, including photocopy, recording, or any information storage and retrieval system, without permission in writing from the publisher.

Academic Press, Inc. 1250 Sixth Avenue, San Diego, California 92101-431 1 United Kingdom Edition published by ACADEMIC PRESS LIMITED 24-28 Oval Road, London NW1 7DX

Library of Congress Catalog Number:

49-7755

International Standard Book Number:

0-12-007839-2

PRINTED IN THE UNITED STATES OF AMERICA 9 3 9 4 9 5 9 6 9 7 9 8

QW

9 8 7 6 5 4 3 2 1

Contents CONTRIBUTORS .................................................................... .................... PREFACE............................................................................................................ MIKHAIL 1. TEMKIN. 1908-1991 ................................................................................

ix xi xiii

Application of Percolation Theory to Describing Kinetic Processes in Porous Solids V. P . ZHDANOV I. I1. 111. IV. V. VI .

Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Elements of Percolation Theory . . . . . . . . . . . . . . . . . . . . . . . . . . . . Desorption from Porous Solids . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Mercury Penetration into Porous Solids . . . . . . . . . . . . . . . . . . . . . . . . Catalytic Deactivation by Site Coverage and Pore Blockage . . . . . . . . . . . Conclusion . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

i

5 16 36 43 48 48

Oscillatory Reactions in Heterogeneous Catalysis F. SCHUTH.B. E . HENRY.AND L . D . SCHMIDT I. I1. I11. IV. V. VI .

Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . A Survey of Oscillatory Reactions . . . . . . . . . . . . . . . . . . . . . . . . . . . Experimental Methods . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Models of Oscillatory Behavior . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Synchronization. Spatial Patterns, and Chaos . . . . . . . . . . . . . . . . . . . . Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

51 54 62 70 105 118 118

Zeolite-Supported Transition Metal Catalysts WOLFGANG M . H . SACHTLER AND ZONGCHAO ZHANG

I. 11. 111.

IV. V. VI . VII . VIII .

Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Preparation of Metal Zeolites . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Carbon Monoxide-Induced Reorganizations of Metal Atoms . . . . . . . . . . . Alloy Particles . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Metal Redispersion . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Metal Particles from Neutral Complexes . . . . . . . . . . . . . . . . . . . . . . . Formation of Metal Clusters by the Ship-in-a-Bottle Method . . . . . . . . . . . Metal-Proton Adducts . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . V

129 133 153 160 164 169 173 175

vi

CONTENTS IX . X.

Catalytic Reactions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Prospectives for Future Applications . . . . . . . . . . . . . . . . . . . . . . . . . . References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

181

208

Selectivity Control and Catalyst Design in the Fischer-Tropsch Synthesis: Sites. Pellets. and Reactors ENRIQUE IGLESIA. SEBASTIAN c. REYES.ROSTAM f. . MAWN.AND STUART L . SOLED 1.

V. VI . VII .

Introduction

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

Design of Fischer-Tropsch Catalysts . . . . . . . . . . . . . . . . . . . . . . . . . . Conclusions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Nomenclature . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

221 223 230 231 278 295 296 298

Catalysis by Metal Ions Intercalated in Layer Lattice Silicates YUTAKA MORIKAWA I. I1 . 111. IV. V. VI . VII .

Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Catalytic Materials . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Catalytic Activity of Metal Ion-Exchanged Fluorotetrasilicic Mica . . . . . . . Dehydrogenation of Methanol over Copper Ion-Exchanged TSM . . . . . . . . Reaction of Methanol over Ti4+-TSM and Related Catalysts . . . . . . . . . . . Immobilization of Catalyst Solution . . . . . . . . . . . . . . . . . . . . . . . . . . Conclusions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

303 305 306 309 312 319 324 326

Catalytic Synthesis of Chlorofluorocarbon Alternatives L . E . MANZER AND V . N . M . RAO I. 11 . I11.

Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Catalytic Synthesis of Key CFC Alternatives . . . . . . . . . . . . . . . . . . . . Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

329 334 348 348

CONTENTS

vii

Molecular Mobility Measurement of Hydrocarbons in Zeolites by NMR Techniques J. CARO,M. BULOW,H. JOEC, J. KARGER,AND B. ZIEROWIUS

I. 11. 111. IV. V.

VI.

VII. VIII.

IX . X.

INDEX

Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . ........ Basic Principles of NMR Self-Diffusion Studies . . . . . . . . . . . . . . . . . . . . Principles of NMR Techniques to Detect Molecular Reorientations . . . . . . . . Principles of Other Mobility Measurements (Comparison with NMR Data) . . . Combined Application of Different Experimental Techniques to Study Molecular Translations and Rotations of Hydrocarbons in Zeolites . . Structure-Related Molecular Self-Diffusion in Zeolites by Pulsed-Field Gradient NMR: Influence of Pore Diameter, Si/AI Ratio, and Concentrations of Internal OH Groups and Cations; Self-Diffusion of Mixtures Location of Diffusion Obstacles Inside the ZSM-5 Framework by Gradient NMR . . . . . . . . . . . . . . . . . . . . . .......... Detection of Spatial Distributions of Diffusion Obstacles over the Crystal Size by Pulsed-Field Gradient NMR . . . . . . . . ............... Self-Diffusion in ZSM-5 Particles of Differe ................. Conclusions . . . . . . . . . . . . . . . . . . . References

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

35 1 353 362 366 369

389 394 399 406 409 410

415

This Page Intentionally Left Blank

Contributors Numbers in parentheses indicate the pages on which the authors' contributions begin.

M. BULOW,The BOC Group, Murray Hill, New Jersey 07974 (351) J. CARO,Zentrum fur Heterogene Katalyse, D-1199 Berlin-Adlershof, Germany (35 1) B. E. HENRY,Department of Chemical Engineering and Materials Science, University of Minnesota, Minneapolis, Minnesota 55455 ( 5 1) ENRIQUE IGLESIA,Corporate Research Laboratories, Exxon Research and Engineering Company, Annandale, New Jersey 08801 (221) H. JOBIC, lnstitut de Recherches sur la Catalyse, 69626 Villeurbanne, France (35 1 ) J. KARGER, Fachbereich Physik der Universitat Leipzig, D - 7010 Leipzig, Germany (351) ROSTAM J. MADON,Corporate Research Laboratories, Exxon Research and Engineering Company, Annandale, New Jersey 08801 (22 1) L. E. MANZER, Corporate Catalysis Center, Central Research and Development, Du Pont Company, Experimental Station, Wilmington, Delaware I9880 (329) YUTAKA MORIKAWA, Research Laboratory of Resources Utilization, Tokyo Institute of Technology, Yokohama 227, Japan (303) V. N. M . RAO,Corporate Catalysis Center, Central Research and Development, Du Pont Company, Experimental Station, Wilmington, Delaware 19880 (329) SEBASTIAN C. REYES,Corporate Research Laboratories, Exxon Research and Engineering Company, Annundale, New Jersey 08801 (221) WOLFGANG M. H. SACHTLER, lpatieff Laboratory, Department of Chemistry, Center for Catalysis and Surface Science, Northwestern University, Evanston, Illinois 60208 (129) L. D. SCHMIDT, Department of Chemical Engineering and Materials Science, University of Minnesota, Minneapolis, Minnesota 55455 ( 5 1) F. SCHUTH,Institut f i r Anorganische and Analytische Chemie, Universitat Mainz, 0-6500 Mainz, Germany ( 5 1 ) STUART L. SOLED,Corporate Research Laboratories, Exxon Research and Engineering Company, Annandale, New Jersey 08801 (22 1) ZONGCHAO ZHANG,lpatieff Laboratory, Department of Chemistry, Center for Catalysis and Surface Science, Northwestern University, Evanston, Illinois 60208 (129) V. P. ZHDANOV, Institute of Catalysis, Novosibirsk 630090, Russia ( 1 ) B. ZIBROWIUS, Department of Chemistry, University of Manchester Institute of Science and Technology, Manchester M601QD, England (35 1) ix


Preface The diverse efforts of authors and the wide scope of coverage undertaken in Advances in Catalysis-indeed, the breadth of catalysis itself-are well characterized by the seven chapters of this volume. These chapters were written by authors from departments of chemistry, chemical engineering, inorganic and analytical chemistry, materials science, and physics; from several universities; from two major industrial laboratories; and from at least three institutes dedicated specifically to catalysis! Similarly, the breadth of subjects covered is rather remarkable. The subjects covered in the present volume include Principles of the physical structure of solid catalysts: V. P. Zhdanov reviews modern concepts concerning pore structures that go beyond the old, lumped parameters of pore size and tortuosity. Catalytic phenomena: F. Schuth et al. have collected many observations, theories, and ideas concerning oscillatory rate behavior. Structure-rate relationships: W. M. H. Sachtler and Z. Zhang present a view of many aspects of catalysis and catalysts utilizing transition metals in zeolites; E. Iglesia et al. discuss catalysts, mechanisms, and performance in the Fischer-Tropsch reaction; and Y. Morikawa makes us aware of a class of intracrystalline catalysts other than zeolites. Specific product orientation: L. E. Manzer and V. N. M. Rao describe catalytic pathways to the modern product challenge of alternatives to chlorofluorocarbons. New analytical techniques for internal catalyst properties: The unique collaborative team of J. Car0 et al. describes developments of NMR techniques for the study of the motions of guest molecules in zeolites. Altogether, one must marvel at the many “disciplinary” aspects that interplay to generate the knowledge and applications of catalytic science.

xi

Mikhail 1. Temkin

Mikhail I. Temkin, 1908-1 991 Mikhail 1. Temkin, born in Bielostock, graduated in 1926 from the Lepeshinsky School in Moscow. At that time, the law prescribed two years of work prior to admission to a university. According to Temkin, these years spent at chemical plants shaped his future interests in linking theory to practice. He graduated in 1932 from Moscow State University and joined the Karpov Institute of Physical Chemistry, where he began studies on thermodynamics and kinetics, with applications to catalysis. After a visit in 1935 to the laboratory of Michael Polanyi in Manchester, Temkin returned to the Karpov Institute, where he started the Laboratory of Chemical Kinetics that he headed for 50 years. All his life, Temkin contributed to science in many areas, such as diffusion of heavy water into ordinary water, fugacity of gas mixtures, theory of mixtures of molten salts, and mass transfer in chemical engineering. But he left his indelible mark in the fundamentals of catalytic kinetics, on a par with C. J. Christiansen and J. Horiuti. It all started in 1938 when Temkin first applied transition state theory to heterogeneous catalysis. Soon after, he published with V. Pyzhev one of the most frequently cited papers in catalytic ammonia synthesis. Since both Mikhail Temkin and Paul Emmett had a profound influence on the theory and practice of this famous reaction, it seems proper to quote here Emmett’s assessment of the 1939-1940 paper of Temkin and Pyazhev.” Numerous studies of the kinetics of ammonia synthesis and decomposition have been made. With a few exceptions, work has tended to show that the slow step in the synthesis of ammonia is the chemisorption of nitrogen and the slow step for the decomposition is the desorption of nitrogen. Furthermore, it turns out that the decomposition and synthesis of ammonia usually involve in the rate expression a term Ph3/Pa2,where y / x is close to 1.5. In 1940, Temkin and Pyzhev derived an equation consistent with both of these observations [M. I. Temkin and V. Pyzhev, Acta Physiochim. U.R.S.S. 12, 327 (1940)]. It has formed the basis for most of the kinetic treatments of ammonia synthesis and decomposition in recent years. Temkin assumed a heterogeneous surface and set up equations for the adsorption equilibrium of nitrogen on iron, for the rate of adsorption, and for the rate of de-

*Reproduced with permission from “The Physical Basis for Heterogeneous Catalysis” (E. Drauglis and R . 1. Jaffee, eds.) Plenum, New York, 1975.

...

Xlll

xiv

MIKHAIL I. TEMKIN, 1908-1991

sorption. Specifically his three pertinent equations are:

1 8 = -In@‘

f

where t9 is the fraction of the surface covered, P’ is the equilibrium pressure or the “virtual pressure” of nitrogen, u is the rate of adsorption, w is the rate of desorption, andf, ao, k,, kd, g, and h are constants. These equations are constructed to conform to the idea that the rates of adsorption and desorption of nitrogen depend exponentially on the fraction of the surface covered with nitrogen. At high coverage, adsorption is slow and desorption fast. Incidentally, it may be noted that measurements by Emmett and Brunauer [P.H. Emmett and S. Brunauer, J . Am. Chem. SOC. 56, 35 (1934)l showed that up to 50 atm partial pressure, the adsorption of nitrogen increased as (PN2)’I6 regardless of whether the nitrogen was by itself or equilibrated with a 3: 1 H2:Nz mixture. In applying these equations, the authors assumed that the adsorption of nitrogen on the iron catalyst in the presence of an ammonia-hydrogen mixture is the same as it would be when at a nitrogen pressure equivalent to the existing partial pressure of ammonia and hydrogen in the gas mixture. Thus, since the equilibrium constant for ammonia synthesis is

K = (PNHJ)~/(PHz)~(PN~) the value of P’ can be represented by equations becomes

(PNH3)2/K(PH2)3,

(7) and the first of the Temkin

As an illustration, the application of these equations to the decomposition of ammonia would take the form

Love and Emmett [K. S. Love and P. H. Emmett, J . Am. Chem. SOC. 63, 3297 (1991)] found experimentally that over a doubly promoted catalyst the rate of decomposition is proportional to (PNH3)”’6/(P&)o.9. This would correspond to fi having a value of 0.3.

The three seminal ideas in this early work of Temkin are quite general. The first is that adsorption of nitrogen is rate determining. The second is the virtual pressure or fugacity of adsorbed nitrogen, a concept of great importance to the understanding of catalytic cycles at the steady state. The third idea is the kinetic description of the catalytic surface as a nonuniform one. The last was systematized later by Temkin’s school, both in theory and in application, to a

MIKHAIL I . TEMKIN, 1908-1991

xv

large number of important catalytic reactions. The importance of Temkin’s theory of kinetics on nonuniform surfaces is not so much in its formation but in the deeper kinetic understanding of how any catalyst works and how to select the catalyst with the fastest turnover rate. Like the kinetic concepts of Christiansen and Horiuti, those of Temkin were far ahead of their common acceptance by the catalytic community. Even today, more than 50 years after the Temkin-Pyzhev paper, the idea of virtual fugacity is not well understood by the majority of workers in catalytic kinetics. It is safe to predict that many of the other ideas of Temkin, like that of average stoichiometric number or reaction routes, will influence younger catalytic kineticists who now have access to powerful computers. The legacy of Temkin is a rich one. While strong on theory, Temkin was also a gifted and exacting experimentalist. He and his co-workers proposed and built in 1950 a continuous flow reactor that operates in a “gradientless” manner to measure directly the rate of reactions catalyzed by solids. Jean-Paul Sartre is reported to have said “I shall die twice: the first time physically, and the second time when no one shall read my works.” Mikhail Temkin will live a long, long second life, as his name will remain known by new generations of kineticists who will not even need to read the original writings, since the main ideas of Temkin are already in all textbooks and monographs on heterogeneous catalysis. MICHELBOUDAIZT


ADVANCES IN CATALYSIS, VOLUME 39

Application of Percolation Theory to Describing Kinetic Processes in Porous Solids V. P. ZHDANOV Institute of Catalysis Novosibirsk 630090, Russia

1.

introduction

Porous solids are widely encountered in industry and everyday life, and their importance has long been recognized by scientists and engineers (1). A considerable amount of work on porous solids has been undertaken in both academic and industrial laboratories ( I 2’). The properties of porous solids, e.g., adsorptive capacity, chemical reactivity, and catalytic activity, are dependent on their pore structure. In general, porous materials may have extremely complex pore structures, which are difficult to represent by simple geometrical models. As a rule, however, the actual pore structures belong to either corpuscular or spongy classes ( 3 , 4 ) .Corpuscular systems are formed by particles of various shapes connected to one another. In this case, the pores represent interstices between particles (see, e.g., Fig. 1). In spongy structures, the pore space can be treated as a lattice of voids interconnected by necks in a three-dimensional network (Figs. 2 and 3). In turn, the latter class of porous solids can be divided into two groups: (1) structures with the pore volume concentrated primarily in voids, whereas the necks possess no volume of their own (Fig. 2), and (2) structures with the pore volume concentrated in necks, whereas the void volume is negligible (Fig. 3 ) . There exist, of course, mixed structures wherein, for example, particles of a corpuscular system may have inner spongy porosity. A feature of special interest for many purposes is the width of pores, e.g., the diameter of a cylindrical pore, or the distance between the sides of a slit-shaped pore. A convenient classification of pores according to their 1 English translation copyright 0 1993 by Academic Press, Inc. All rights of reproduction in any form reserved.

2

V . P . ZHDANOV

FIG. I . Corpuscular porous structures. (a) The element of the packing of spheres; (b) the model of round disk packing; (c,d) models of the packing of round rods.

PERCOLATION THEORY AND KINETICS IN POROUS SOLIDS

3

4

V . P . ZHDANOV

FIG. 2. Spongy porous solids with regular (a) and irregular (b) structures. The pore volume (white areas) is primarily concentrated in voids, whereas the necks do not possess volumes of their own.

average width originally proposed by Dubinin (5) and now officially adopted by IUPAC (6) contains three size ranges corresponding to characteristic adsorption effects as manifested in the nitrogen isotherms. In micropores (of widths less than 20 A), the interaction potential is significaiitly higher than in wider pores owing to the proximity of the walls, and the amount adsorbed at a given ressure is correspondingly enhanced. In mesopores (between 20 and 500 ), capillary condensation with its characteristic hysteresis loop takes place. In the macropore range (more than 500 A), the pores are so wide that it is virtually impossible to map out the isotherm in detail because the corresponding relative pressures are close to unity. In the latter case, the pore-size distribution can be studied by employing mercury penetration. Most of the pore structures (e.g., spongy structures) consist of extensive three-dimensional networks in which there is a profusion of interconnections between voids within the structure. The latter interconnections affect considerably the kinetics of various processes in porous solids. This effect can adequately be described by employing the ideas developed in percolation theory (7-13). In the framework of this theory, the medium is defined as an infinite set of sites interconnected by bonds. Percolation theory can be applied to porous solids via identification of network sites with voids, and bonds with necks. Thus, the theory is applicable primarily to spongy porous structures but in some cases also to corpuscular structures. During the past decade, percolation theory has been successfully used to analyze condensate desorption from porous solids (14-34), mercury penetration into porous solids (35-43), and the kinetics of catalytic deactivation

R


5

FIG. 3. Spongy porous solids with regular (a) and irregular (b) structures. The pore volume (white areas) is primarily concentrated in necks, whereas the void volume is negligible.

by site coverage and pore blockage (44-46). Our aim, to survey the main results obtained in this field, is of particular interest to the heterogeneous catalysis community.

II. Elements of Percolation Theory Historically, percolation theory goes back to World War 11, when Flory and Stockmayer used it to describe the formation from small branching molecules of larger and larger macromolecules as increasing numbers of chemical bonds formed between the original molecules (12).Usually, however, the origin of percolation theory is associated with a 1957 publication by Broadbent and Hammersley (7), which introduced the terminology and dealt with the matter more mathematically using geometrical and probabilistic concepts. During the past 15 years, percolation theory has become very fashionable in various fields of physics and chemistry. The theory has been comprehensively discussed in several reviews (8-1I ) and monographs (12,Ij).In the following discussion, we briefly outline the main ideas and results of percolation theory, which are of interest from the point of view of describing the kinetic processes in porous solids.

A . PERCOLATIONMODELS In contrast to many other modern research fronts, percolation theory is a problem that is, in principle, easy to define (8-13). In general, a percolation model is a collection of points (or sites) distributed in space, certain

6

V . P . ZHDANOV

pairs of which are assumed to be linked by a random mechanism, the details of which depend on the phenomenon under consideration. A path is considered to exist between two points of the distribution if a sequence of points may be found, beginning from the first and ending with the second, such that successive points in the sequence are linked. The points of the distribution may be aggregated into clusters such that pairs of points in the same cluster are connected but there is no path between points of different clusters. Most systems to which percolation theory is applicable contain so many points that the surface effects may be ignored and the system may be replaced by a model with an infinite number of points. In the latter case, with increasing numbers of linkages, the cluster size may become infinite at some critical density of linkages. Above this density, the system is said to be in a percolating state. The transition from a nonpercolating state to a percolating one is a kind of phase transition. The probability that a given point belongs to an infinite cluster is known as the percolation probability. The main aim of percolation theory is to calculate the percolation probability for various models. There are many percolation models (ZZ,Z3). The most typical and important ones are discussed in the following sections. 1.

Bond Percolation on Regular Lattices

In this case, the medium is defined as an infinite set of sites located in the vertex positions of a regular lattice and interconnected by bonds that are directed along the edges. The random characteristic of the medium is introduced by blocking some of its bonds at random, so that any bond, independently of all other bonds, has a fixed probability q of not being blocked (see, e.g., Fig. 4).The percolation probability 9 b is defined as the probability that a single site is connected with an infinite number of other sites with unblocked bonds (47).The percolation probability introduced characterizes the sites. For the bond problem, one can also employ the percolation probabilities YPbland P b Z , which characterize bonds. The former probability is defined to be the probability that a given unblocked bond belongs to an infinite cluster, and the latter is the probability that an arbitrary bond is included in this cluster. The probabilities 9 b l and g b 2 are obviously connected as

2. Site Percolation on Regular Lattices In this model, the vertices of a lattice (i.e., sites) rather than its edges are blocked, i.e., each site is declared to be open with the probability q and is

-..

::m

2,:

A.

-..I -...

.... .-. .....l J " x : ! , l J ~ ~ 3% ... ;;z; d.... 1,..-..... .... 1.. -1.-

I - . .

_I.

-

] . . . I

......L

::FJ ::=: !I :L! I : :I :I

.-

1.1

... ... ... 'I .I .1 ... -. .111

(a) q = 0.25

FIG.4. Realization of bond percolation on a 50 X 60 section of the square lattice for four different values of q. The diagrams have been created using the same sequence of pseudorandom numbers, with the result that each graph is a subgraph of the next. Attentive readers may verify that open paths exist joining the left to the right side when q = 0.51 but not at q = 0.49. (From Ref. 13, with permission.) (continued)

8

V . P . ZHDANOV

(c) q = 0.51

FIG.4. (continued)


9

closed otherwise, independently of all the other sites. The percolation probability GPSlis defined as the probability that a given open site belongs to an infinite cluster. In addition, one can use the probability Ps2,which is introduced as the probability for an arbitrary site to be in the unbounded cluster. By analogy with Eq. (l), we have PS2

= qPsl

(2)

Site percolation is known to be more general than bond percolation in the following sense: every bond process may be reformulated as a site process on a different lattice, but there exist site processes that do not arise from the bond processes ( 1 3 ) .In particular, for any given lattice, one can construct a covering lattice so that each edge of the original lattice corresponds to the vertex of the new lattice. Two such vertices are adjacent if the corresponding edges of the original lattice have an end vertex in common (see, e.g., the honeycomb lattice and its covering Kagome lattice in Fig. 5). It is evident that there is no difference between studying bond percolation on the original lattice and studying site percolation on the covering lattice.

3. Mixed Percolation In the framework of this model, both the edges and the vertices of a regular lattice may be open or closed, with different probabilities. The lattice

FIG. 5 . The solid circles are vertices of the hexagonal honeycomb lattice, and the open circles are vertices of its covering lattice.

10

V . P . ZHDANOV

formed by blocking some of its sites at random is often said to be a randomized lattice. The study of bond percolation on randomized lattices is of particular interest because these lattices are in fact the simplest examples of irregular lattices.

4. Percolation on Bethe Lattices The Bethe lattice (or tree) is a lattice containing no closed loops (see, e.g., Fig. 6). The latter results in a simple analytical solution of the bond and site problems for these lattices. In addition, the general percolation properties for Bethe and regular lattices are often close. For these reasons, the Bethe lattices are rather popular in applied science publications, although these lattices have no physical sense. 5 , AB Percolation

In this case, a given lattice site can be occupied at random by particle A with the probability q or by particle B with the probability 1 - q. Each edge is declared to be open if its end vertices are occupied by particles of different types (the idea is that dissimilar particles bind together, whereas similar ones repel each other). Of principal interest is the probability of the lattice origin to be an infinite cluster of open edges. The properties of the models just described are governed by the topology of the underlying lattices. These models have been generalized in several directions, including (1) anisotropy of the probabilities of being open,

FIG.6 . The Bethe lattice.

11


(2) long-range bonds, and (3) correlated percolation, in which the states of different sites or bonds are not independent, and so on (13). In addition, there are continuum versions of percolation theory. For example, recall the following classical problem of stochastic geometry: ascertain the minimal density of the unit disks in the plane (or spheres in the space) that guarantees the existence of an infinite cluster. All these problems are, however, far beyond the scope of this review.

B. PERCOLATION PROBABILITIES The percolation probability 9 ( q ) for the lattice models is defined as the probability that a given site (or bond) belongs to an infinite open cluster (47). It is fundamental to percolation theory that there exists a critical value q. of q such that 9 ( q ) = 0 at q < q,, and 9 ( q ) > 0 if q > 4.. The value q. is called the critical probability or the percolation threshold. Mathematical methods of calculating this threshold are so far restricted to two dimensions, consistent with the experience in the field of phase transitions that three-dimensional problems in general cannot be solved exactly (12,13). Almost all quantitative information available on the percolation properties of specific lattices has come from Monte Carlo calculations on finite specimens (8,11,12). In particular, Table I summarizes exactly and approximately known percolation thresholds for the most important two- and threedimensional lattices. For the bond problem, the data presented in Table I support the following well-known empirical invariant (8) zq, -- d / ( d - 1)

(3)

TABLE I Percolation Thresholds and Dimensional Invariants" Bond problem Latticeh

Z

9c

Site problem

Z4C

q

4.

17%

Honeycomb Square Triangular

3 4 6

0.653 0.500 0.347

1.96 2.00 2.08

0.61 0.79 0.91

0.696 0.593 0.500

0.424 0.468 0.455

Diamond

4 6

0.388 0.249 0.180 0.119 0. I24

1.55 1.49 1.44 1.43 1.49

0.34 0.52 0.68 0.74 0.74

0.428 0.312 0.246 0.198 0.204

0. I46 0. I62 0.167 0. I47 0.151

sc

BCC FCC HCP

8

12 12

"From Refs. 8 and 12. Lattice arrangements: SC, simple cubic; BCC, body-centered cubic; FCC, face-centered cubic; HCP, hexagonal close packing.

12

V . P . ZHDANOV

where z is the coordination number and d is the dimensionality of the network. To construct an invariant for the site problem, one may circumscribe spheres (or circles) about each site of the lattice, with the radii of half the nearest-neighbor distance. Then, the empirical site-problem invariant can be introduced as (8) vc = 774.

(4)

where q is the filling factor for the lattice, defined as the ratio of the volume of a circumscribed sphere to that of a unit cell (0, -- 0.45 for d = 2, and vc = 0.15 for d = 3). The percolation probabilities obtained by Monte Carlo simulations on common lattices are shown in Fig. 7. Taking into account these results and the invariant [Eq. (3)], one can consider in applications that the percolation probability p b for the bond problems is a universal function dependent only on two parameters, i.e., d and zq. In particular, for three-dimensional lattices this probability can be represented as pb(X)

=

0, x < 1.5 1 . 5 4 ( ~- 1.5)04/[1 1, x > 2.7

i

-F

0.606(~- 1.5)0.4],

1.5 < x < 2.7 (5)

where x = zq. On the other hand, the percolation probability for the site problems is often assumed to be a universal function of d and qq. The universal percolation probabilities for regular lattices are expected to be applicable also to irregular lattices. It is useful to consider two examples supporting this important statement. 1. The simplest way of constructing the irregular lattice with a variable coordination number is to break bonds of a regular lattice at random. In this case, the average coordination number in the irregular lattice is z = q o z o , where qo and zo are the fraction of unbroken bonds and the coordination number of the regular lattice, respectively. Let us now analyze the bond problem on the irregular lattice constructed. Blocking bonds of the irregular lattice, we expect the percolation probability for this lattice to be ppb(zq), where q is the probability for bonds in the irregular lattice to be unblocked and pb(x) is the same function as for regular lattices. Taking into account that the probability for bonds in the original regular lattice to be unblocked is ql = 4.4, we obtain zq = qozoq = z o q l , i.e., the parameter zq for the irregular lattice is just the same as that for the original regular lattice. The latter justifies the assumption that the percolation probability for either lattice is described by the same function.

13


a TRIANGULAR

HEXAGONAL

m HCP

m rn

;:i .20 40

b

O3

0.4

I

.04 I2 .20.28.36 44.52.60

TRIANGULAR HEXAGONAL I SOUARE I

.40

TETRAHEDRAL SIMPLE CUElC

48 .56 .64 .72 .80.88 16 .24 .32 40 48 .56 .64

t

/

,////I /

t

,

0.21 /

o k 0

I

I

0.2

,

0.4

0.6

0.8

5.

1 1

FIG.7 . (a) Percolation probabilities P,, and 8,)as a function of q for two (left-hand panels) and three (right-hand panels) dimensions. HCP, hexagonal close packing; FCC, facecentered cubic. (b) Percolation probability Pb2versus q for the simple cubic lattice. (From Ref. 8, with permission.)

14

V . P . ZHDANOV

2. The mixed bond-site percolation problem yields a more complex way of constructing irregular lattices. In the framework of this problem, q b and q. (fractions of open bonds and sites of the original lattice) can have any values in the range from 0 to 1 simultaneously. If one analyzes, for example, the bond problem on an irregular lattice formed by blocking sites of the original regular lattice, the average coordination number for the irregular lattice is z = qszo, where zo is the coordination number for the original regular lattice. Employing the universality hypothesis, one should expect that the percolation threshold for the irregular lattice constructed takes place at zqb = z o q s q b = 1.5 (for d = 3), i.e., at the percolation threshold the product z o q s q b should be approximately constant. The latter expectation is indeed in good agreement with the results obtained by Monte Carlo simulations (Fig. 8). Finally, it is reasonable to present the equations describing percolation on the Bethe lattices (Fig. 6). The percolation probability can be calculated exactly for this model and is the same for bond and site problems (10). For example, the probability that all open walks from a chosen site are of the finite length, 1 - Yb, can be represented as 1 - Yb = [Q(q>I.

Infinite cluster region d

't 0

Finite cluster region

0.2

0.4

0.6

0.8

9s

1

FIG.8. Parameter zoq,qb at the percolation threshold as a function of qs for the simple cubic lattice. The drawing is constructed by employing the Monte Carlo results (33).


15

where Q(q) is the probability that all open walks in a chosen direction are finite. The step from a given bond to z - 1 nearest-neighbor bonds yields the following recurrent relation: Q(q)= 1 - q

+ qQ(q)'-*

(7)

Equation (7) always has a solution Q ( q ) = I corresponding to 9&) = 0 at q < 4. For z 2 3, there is a second solution corresponding to 9&) > 0 at q > q c . An elementary analysis yields q c = 1/(z - 1)

(8)

Comparing Eqs. (8) and (3), one can conclude that the Bethe model is appropriate for describing the three-dimensional lattices only at z = 3 and 4. If z 2 5, the percolation probability for the Bethe lattice differs considerably from that for regular lattices.

C. CRITICAL EXPONENTS AND SCALING Percolation media can be characterized not only by the percolation probability but also by other quantities (Table 11)-for example, by the correlation length, which is defined as the average distance between two sites belonging to the same cluster. Near the percolation threshold, all these quantities are usually assumed to be described by the power-law equations (Table 11). All current available evidence strongly suggests that the critical exponents in these equations depend only on the dimensionality of the lattice rather than on the lattice structure (12). Also, bond and site percolations have the same exponents. TABLE I1 Critical Exponents for Three-Dimensional Lattices" Exponent

Value

Equation

Quantity ~~~~

a

P Y V U

7

-0.62 0.41 1.80 0.88 0.45 2.18

M, - [ P - P c ) 2 - " 9- ( P - P,)P

-(P- P,(p 5 - IP P c ( - " s

-

Q. (9) Eq. (9)

"From Ref. 12. with permission.

~~~~~~~~~

Singular part of the mean number of clusters per site Percolation probability Mean number of sites in a finite cluster Correlation length Cluster numbers Cluster numbers

16

V . P . ZHDANOV

Different critical exponents are not independent. The relationships between them can be obtained by employing the scaling assumptions. In particular, one can assume that the average number of s clusters (i.e., clusters containing s sites) per site can be represented at arbitrary qand s as (12) ns(q> = s-‘%:C(q - qc)sU1

(9)

where T and u are the cluster number exponents. The precise form of the scaling function 4 ( x ) may be determined by computer experiments. However, even if the precise form of 4 is unknown, it is possible to express the exponents a , p , and y through 7 and (T by employing only general analytical properties of %(12). The latter yields the following scaling relationships:

2-

ff

= 2p

+y

= (7 - 1)/U

(10)

Using similar ideas, one can obtain (12) dv=2-a The scaling law, Eq. (1 I), where the dimensionality d becomes valid, is often called “hyperscaling.” Finally, it is of interest to discuss briefly percolation on finite lattices. In the latter case, the percolation probability can be defined as the probability that an arbitrarily selected site (or bond) belongs to the largest cluster of the system. The scaling expression for 8 is usually assumed to be (12) 8 = L-O’”G[(q

-

qc)L””]

(12)

where L is the lattice linear dimension. Because for a finite system the largest cluster always has a finite number of sites, the scaling function G(x) is a positive function of its argument and is analytical everywhere, even at the percolation threshold where x = 0. The behavior of G ( x ) for the simple cubic lattice at L = 10-50 has been studied in detail by Kirkpatrick ( 4 8 ) (see also Ref. 3 4 ) .

111.

Desorption from Porous Solids

In the following discussion we will consider the application of percolation theory to describing desorption of condensate from porous solids. In Sections II1,A-II1,C we briefly recall types of adsorption isotherms, types of hysteresis loops, and the Kelvin equation. The matter presented in these sections is treated in more detail in any textbook on adsorption [see, e.g., the excellent monographs written by Gregg and Sing (6) and by Lowell and Shields ( 4 9 ) ] .Sections III,D-II1,H are directly connected with percolation theory. In particular, general equations interpreting the hysteresis loop are


17

presented in Sections II1,D-II1,F. The effect of various factors characterizing the pore structure on the desorption process is demonstrated in Section III,G. Typical experimental data are discussed in Section II1,H.

A. ADSORPTION ISOTHERMS The study of the pore structure of porous solids is closely connected with interpretation of adsorption isotherms (6). The majority of those isotherms resulting from physical adsorption may conveniently be grouped in five classes (Fig. 9), as was originally proposed by Brunauer et al. (50-52). The type I isotherm corresponds to the Langmuir case when adsorption is confined to a monolayer. The multilayer physical adsorption of gases by nonporous solids, in a vast majority of cases, gives rise to a type I1 isotherm, which can be described by the Brunauer, Emmet, and Teller (BET) equation (6,51). The type IV isotherm corresponds to adsorption and desorption in porous solids. In particular, the mesoporous range of pore sizes usually gives rise to this type of isotherm (6). In the low-pressure region, the type IV

RELATIVE PRESSURE, P/Po FIG.9. Five types of adsorption isotherms, I to V, using the classifiction of Brunauer er al. (50) together with type VI, the stepped isotherm. P is the vapr pressure (Po corresponds to saturation). The amount adsorbed, U, is in arbitrary units.

18

V . P . ZHDANOV

isotherm follows the same path as the corresponding type I1 isotherm, but then it begins to deviate upward until at higher pressures its slope decreases. Isotherms of types 111 and V are rare and will not be discussed further. A characteristic feature of type IV isotherms measured on porous solids is the hysteresis loop. The exact shape of the loop varies from one adsorption system to another, but the amount adsorbed is always larger at any given relative pressure along the desorption branch than along the adsorption branch. The first classification of hysteresis loops was proposed by de Boer (53), who has identified five types (A-E) of loops that may be correlated with various pore shapes. Subsequent experience has, however, shown that types C and D hardly ever occur in practice (6). A new classification of hysteresis loops recommended by IUPAC (6) consists of four types, i.e., H I , H2, H3, and H4 (Fig. 10). The first three types correspond to types A , E, and B, respectively, in the original de Boer classification. As pointed out above, certain shapes of hysteresis loops are usually assumed to be associated with specific pore structures (6). Thus, type H1 loops are often obtained with agglomerates of spheroidal particles of fairly uniform size. Some corpuscular systems (e.g., certain silica gels) give rise to an H2 loop. Types H3 and H4 have been obtained with samples having

RELATIVE PRESSURE, P/Po FIG.

10. IUPAC classification of hysteresis loops; U is in arbitrary units.


19

slit-shaped pores or platelike particles. The shape of hysteresis loops is also strongly dependent on the pore-size distribution. In particular, changing the neck-size distribution, one may obtain various types of hysteresis loops even if the void-size distribution is fixed.

B. THEKELVINEQUATION The study of mesoporous solids is closely related to the concept of capillary condensation and its quantitative expression in the Kelvin equation (54). This equation is the basis of virtually all the various procedures for calculating the pore-size distributions from the type IV isotherms that have appeared over the past 80 years, beginning with the first papers of Zsigmondy (55), Anderson (56), and Foster (57) (see also Refs. 6 and 49) A convenient form of the Kelvin equation is ln(P/Po) =

-

2yVL/r,RT

(13)

where P / POis the relative pressure of vapor in equilibrium with a meniscus having a radius of curvature r,, Po is the saturation vapor pressure corresponding to rm = a,y is the surface tension, and VL is the molar volume of the liquid. The derivation of Eq. (13) is presented in any textbook on thermodynamics (see, e.g., Ref. 58). The Kelvin equation is also discussed in detail by Gregg and Sing (6) and by Lowell and Shields (49). For nitrogen as adsorbate at its boiling point of 77 K, the Kelvin equation can be rewritten as rm = 9.4/ln(Po/P)

(14)

where rmis in angstroms. Applying Eqs. (13) and (14), one must take into account that during isotherm determination the pore walls are already covered with an adsorbed film of thickness t (6). Thus, capillary condensation actually occurs not directly in the pore but rather in the inner core (Fig. 1l), and we have the following connection between the pore and core radii: rp = rk

+t

(15)

In addition, the core radius is related to rmas rk

= rmcos

8

(16)

where 8 is the angle of contact between a liquid and a solid surface. The accurate measurement of a contact angle is very difficult even for macroscopic systems. For this reason, Eq. (16) is usually employed with 8 = 0 (cos 8 = 1).

20

V.P. ZHDANOV

FIG. 1 1 . Cross section of a cylindrical pore of radius r,,.

The thickness of the adsorbed film can be estimated by using the BET isotherm. The appropriate results can be described closely by the Halsey equation, which for nitrogen is written as (49,59) t = 3.54[5/ln(PO/P)]'/' where t is in angstroms. Typical values of r,, Table 111.

c.

rk,

(17) and t are given in

DESCRIPTION OF ADSORPTION

As pointed out in Section I, the pore space can generally be treated as a lattice of voids interconnected by necks in a three-dimensional network (Figs. 2 and 3). It is often possible to consider that the pore volume is concentrated in voids, whereas necks do not possess a volume of their own (Fig. 2). In the framework of this model, the filling of every void on the adsorption branch of the isotherm is determined only by the individual void characteristics and does not depend on the neck-size distribution. In particular, voids with radii lower than the Kelvin radius, r < r,, are completely filled and those with r > r, are filled only partly via the reversible sorption TABLE I11 Values of rk, t , and rp as a Function of P/POfor Nitrogen at 77.4 K

PIP* 0.3 0.4 0.5 0.6 0.7 0.8 0.9

rk

(A)

7.8 10.2 13.6 18.4 26.4 42.1 89.2

t (A)

r, (A)

5.7 6.2 6.8 7.6 8.5 10.0 12.8

13.5 16.4 20.4 26.0 34.9 52.1 102.0


21

mechanism (multilayer adsorption). Recall that in our review the Kelvin radius at a given pressure P is defined by Eq. (15), i.e., it incorporates the width of the adsorbed film. If, for example, the void shape is spherical, the fraction of the pore volume filled by adsorbate at a given value of rp on the adsorption branch of the isotherm is represented as Uad(rp)= ( [ r 3 f ( r ) dr

=1

-

+ [[r3

- (r

-

t ) 3 ] f ( r )d r ) / I

‘P

110

(I - t)3f(r) dr/I

(18)

‘P

with

I

=

[

r ’ f ( r ) dr

where f ( r ) is the radius distribution function for voids. The analysis of the adsorption branch of the isotherm permits one to obtain the void-size distribution (6). D. DESCRIPTION OF DESORPTION (APPROACH 1) During desorption, as the relative vapor pressure is reduced, pore solids in which capillary condensation occurs often show a hysteresis loop. The simplest interpretation of this phenomenon is given by the ink-bottle model (6). In the framework of this model (Fig. 12) the adsorption and desorption processes are controlled, respectively, by the void and neck sizes. Thus, desorption from a given pore occurs at a lower pressure than adsorption.

FIG. 12. Ink-bottle pores with cylindrical and tapering bodies.

22

V . P . ZHDANOV

However, as Everett (60) pointed out, the analogy of a pore as a narrownecked bottle is overspecialized, and in practice a series of interconnected pore spaces rather than discrete bottles is more likely. In the latter case, a void with the radius r > r, is empty during the desorption process only if it is connected with the outer surface by a chain of voids and necks with r > r,. Thus, the emptying of a given pore is not solely determined by its immediate characteristics. Hence, a correct analysis of the desorption branch of the isotherm should take into account the three-dimensional interconnection of various voids. This problem can be solved by using percolation theory, as has been done by Wall and Brown (14), Kheifets and Neimark (15-17), Mason (Z8-21), Fenelonov et al. (22-25), Palar and Yortsos (26,27),Mayagoitia et al. (28-32), Yanuka (33), and Seaton (34). Application of percolation theory to describing the desorption process from porous solids is based on the identification of network sites with voids, and bonds with necks. A bond is considered to be unblocked if the neck radius r > r,. Unblocked sites belonging to the percolation cluster correspond to voids containing nitrogen vapor. Generally speaking, a strict description of the desorption process demands the solution of problems more complicated than the ones that have been solved in modern percolation theory, because, in the case of very porous materials, the coordination number is not constant and the voids are not equivalent. In addition, the size distribution of voids and necks can be overlapping. For a rather large group of porous solids, however, a suitable description of the desorption process can be obtained on the basis of the available percolation theory data by using the mean coordination number and assuming that the percolation probability for nonregular media is the same as that for regular ones. As pointed out above, the desorption process is dependent both on the void- and neck-size distributions, f(r) and cp ( r ) . If the radii of all the voids are larger than those of all the necks (i.e., B > C ;Fig. 13a) and the void and neck arrangements are random (the latter term means that the probability for an arbitrary void or neck to have a given value of the radius does not depend on the sizes of the neighboring voids and necks), the desorption process is mathematically equivalent to the bond problem in percolation theory. In particular, the probability that an arbitrary void is empty at a given value of the Kelvin radius during desorption is equal to the percolation probability 9,,(z,q) for the bond problem. Thus, the volume fraction of emptied voids under desorption [I - Udes(rp)]can be represented as the product of the fraction of pore volume that may be emptied in principle at a given value of rp [l - Uad(rp)] by the percolation probability %,(z,,q), i.e., [1 - Udes(rp)]= [l - Uad(rp)]9b(Z,q)

(20)


23

Pore radius FIG. 13. Scheme of the size distribution of voids and necks without (a) and with (b) overlapping.

where u a d ( r p ) and u d e s ( r p ) are the fractions of pore volume filled by adsorbate during adsorption and desorption, respectively, zo is the mean coordination number for all the voids defined by

zo = 2 and

4=

[

q ( r ) d r / [ f ( r ) dr

f4Pk)

(21)

B

d r / [ cP(r) dr

(22)

‘P

is the fraction of necks with r > rp (the parameters A , B , C , and D are clarified in Fig. 13). Multiplying Eqs. (21) and (22), we have

zoq = 2

f [

q ( r ) dr/

f ( r ) dr

(23)

‘P

Here and below, the radius distribution of voids, f ( r ) , is assumed to be defined so that f ( r ) d r is the number of voids with the radius from r to r + dr per unit volume (or mass) of the sample. The radius distribution of necks, cp(r), is defined analogously. Equations (20) and (22) were first derived by Wall and Brown (14) and were then employed by Kheifets and Neimark (15- 17). In practice, the case in which the size distributions of voids and necks do not overlap is rare. In general, these distributions may be overlapping. The latter always results in correlations in the arrangement of voids and necks. For example, if the arrangement of voids is random, the distribution of

24

V.P. ZHDANOV

necks is correlated because high-size necks cannot be put between low-size voids. On the other hand, if the neck distribution is random, the void arrangement will necessarily have some correlations. At present, data on the correlation effects in the void and neck arrangements are rather scarce. For this reason, the existing approaches to describing desorption from porous solids usually take into account only the simplest effect, i.e., the fact that the neck between the two nearest-neighbor voids should be lower than or equal to the size of the smaller void. The latter fact may be incorporated into the theory in two slightly different ways. The first approach outlined in this section was originally proposed by Zhdanov et af. (23). The second approach, which is presented in Section III,E, was developed by Mason (18-20) and Palar and Yortsos (26,27). Following Zhdanov et af. (23), we will apply the percolation theory data only to those voids that can, in principle, be emptied at a given value of the Kelvin radius, i.e., to those with the radius r > rpand to their directly connected necks. By analogy with Eq. (20), we can write [I -

udes(rp)]

= [I -

uad(rp)]yb(Zq)

(24)

where p b ( Z q ) is the percolation probability for voids with r > r p . The parameter zq in Eq. (24) should be calculated also only for voids with r > rp. Employing Eq. (24), we may assume that the percolation probability for the sublattice of voids with r > rp is the same as the universal percolation probability for the bond problem (Section 11). The latter probability was originally calculated only for regular lattices. The sublattice of voids with r > rp is not regular. In addition, some of the voids with r > rp are not connected with the other voids having r > r,. However, the numerical results obtained by Yanuka (33) for a “randomized” cubic lattice (see also the discussion in Section 11) support the hypothesis on the universality of the percolation probability for both regular and irregular lattices. To apply Eq. (24), we must determine the parameter zq for voids with r > rp and for the connecting necks. The number of voids with r > r p is

I

D

n(rp> =

f ( r )dr

(25)

‘P

Let the total number of their connecting necks be equal to Nt(rp).Then the mean coordination number for voids with r > rp is determined as Z(TP)

=

2Nk-P)hkP)

The number of necks with r > rp is

(26)


25

These necks correspond to unblocked bonds in percolation theory. Thus, the fraction of unblocked bonds for the sublattice of voids with r > rp is determined by the relationship drp)

=

W P ) / W P )

(28)

Multiplying Eqs. (26) and (28), we get zq = WrP)/4rP)

(29)

Substituting Eqs. (25) and (27) into Eq. (29) yields zq = 2

p(r) d r / p f ( r ) dr P'

(30)

'P

Equations (24) and (30) solve the problem; they relate the fractions of pore volume emptied from the condensate during adsorption and desorption with the percolation probability and the radius distributions of voids and necks. If these distributions do not overlap, i.e., rp < C (Fig. 13a), one can replace rp by C in the denominator of Eq. (30). In this case, Eqs. (24) and (30) transform into Eqs. (20) and (23). Note also that the original treatment (23) of the desorption process contains an additional parameter A 3 1.0 (the empirical coefficient taking into account that a void with the radius r has necks with radii lower than r/A). In the present discussion, we put A = 1. Taking into account the fact that the mean coordination number for all the voids is defined by Eq. (21), we can rewrite Eq. (30) as zq =

where

zo W-P)/Fb-P)

1;

(31)

cP(r) d r / [ cP(r) dr

(32)

F ( r J = p ( r ) d r / & f ( r ) dr

(33)

W.P)

=

and

P'

are the fractions of necks and voids with r > r,. Equation (31) is useful in practical calculations. Finally, it is reasonable to discuss the relationships between fractions [Eqs. (32) and (33)]. Assuming that the neck arrangement is random and taking into account that the fractions of necks and voids with r < rp are defined by 1 - @(rp) and 1 - F ( r p ) [cf. Eqs. (32) and (33)], we have the following obvious constraint for the neck and void distributions: [1 -

@(rP)]'o

3

1 - F(rp)

(34)

26

V . P . ZHDANOV

The underlying principle for Eq. (34) is that no void has a size smaller than its associated zn necks. In practical calculations (18-20,26,27), Eq. (34) is often replaced by the relationship (35) [I - @ ( r p ) b= 1 - F(rp) If the void arrangement is random, the probability for at least one of two nearest-neighbor voids to be smaller than rp is 1 - F 2 ( r p ) This . probability should be lower than the fraction of necks with r < r,, i.e., 1 - F2(r,) < 1 - @(rp),or W P ) > @(rp) (36) E. DESCRIPTION OF DESORPTION (APPROACH 2) Mason (18-20) and Palar and Yortsos (26,27) have employed another way of describing desorption from porous solids. Their approach is based on the assumption that the neck arrangement is random, i.e., the probability for an arbitrary neck to have a given value of the radius does not depend on the sizes of adjacent voids and necks. In this case, one can apply the percolation theory data obtained for the bond problem to all the voids. In particular, the probability for an arbitrary void to be empty during the desorption process is precisely gb(Zoq), where the parameter znq is given by Eq. (23). The latter probability is calculated for all the voids. We, however, know for a fact that voids with r < rp are filled. Thus the probability for a void with r > rp to be empty is just 9t,(z,q)/F(rp), where F(rp) is the fraction of voids with r > rp [Eq. (33)]. Then, by analogy with Eq. (20), we derive

[I - Udes(rp)] = [I - u,d(rp>]gb(z,q)/F(rp) (37) If the void- and neck-size distributions are not overlapping, i.e., rp < C (Fig. 13a), one can replace r, by C in the numerator of Eq. (33). In this case F = 1 and Eq. (34) transforms to Eq. (20). To calculate the percolation probability in Eq. (34), Mason (18-20) and Palar and Yortsos (26,27) have employed the Bethe tree model. This model is known (see Section 11) to yield quantitative results only at 3 d zo S 5 . In general, however, one can use in Eq. (32) the universal percolation probability calculated for regular lattices (see Section 11). Mason (20) and Palar and Yortsos (2627) have also considered in detail adsorption scanning curves starting on the boundary desorption isotherm (Fig. 14) and desorption scanning curves starting on the boundary adsorption isotherm (Figs. 14 and 15). Mayagoitia et af.(28-32) have analyzed the desorption process from the pore space described by a joint site-bond radius distribution with a correlation function that carries structure information about the network. The integral equations derived (29) are based on the Bethe tree model.

3

5

I

0.35

I

0.45

I

I

0.55

I

0.65

I

0.75

I 0.85

RELATIVE PRESSURE, P/P, RELATIVE PRESSURE, PP, FIG. 14. Adsorption (a) and desorption (b) scanning curves for xenon on Vycor porous glass at 151 K; U is in arbitrary units. (From Ref. 20, with permission.)

28

V.P. ZHDANOV

FIG. 15. Scheme of desorption starting on the boundary adsorption isotherm. Vapor is in white pores; liquid is in black pores. (a) At initiation of the desorption process; (b) at a later stage. (From Ref. 26, with permission.)

Palar and Yortsos (27) have studied in detail nucleation effects during desorption. In this case, desorption was treated as a growth problem with continuous generation of voids occupied by vapor. It was concluded that homogeneous nucleation may be neglected in typical experimental conditions. Heterogeneous nucleation, being pore-size specific, has a higher likelihood to affect the desorption isotherms. Deviations from a percolation behavior attributed sometimes to compressibility and finite lattice size can likewise be explained by nucleation.

F. DESCRIPTION OF DESORPTION (APPROACH 3) Analyzing adsorption (Section II1,C) and desorption (Sections II1,D and III,E), we assumed that the pore volume is concentrated in voids, whereas necks do not possess volumes of their own (Fig. 2). Seaton (34) has recently considered an alternative model of porous solids assuming the pore volume to be concentrated in necks (Fig. 3). In the framework of this model, the adsorption process is described by analogy with Eq. (18) (one should only replace the void radius distribution by the neck radius distribution), and consequently the analysis of the adsorption branch of the isotherm allows one to obtain the neck-size distribution. The desorption process can be described by using the same ideas as in Sections III,D and III,E because this process is mathematically equivalent to the bond problem in percolation theory, even if the pore volume is concentrated in the necks. In particular, the volume fraction of emptied necks under desorption [ 1 - Udes(rp)]can


29

be represented as the product of the fraction of pore volume that may be emptied in principle at a given value of r, [l - Uad(rp)], by the percolation probability p b l (zoq), i.e.,

[1

-

udes(rp)] = [I

-

uad(rp)]9bl(Zoq)

(38)

where zois the mean coordination number for all the voids and q is the fraction of necks with r > rp [Eq. (22)]. Recall that the probability ??bl in percolation theory is defined (Section 11) as the probability that a given unblocked bond belongs to an infinite cluster. Unblocked bonds correspond to necks with r > rp. Thus 9 b l is just the probability that a neck with r > r, is empty. Taking into account the relationship given by Eq. (l), one can rewrite Eq. (26) as

[I

-

udes(rp)]

= [1

-

~ad(~p)]~bZ(zoq)/q

(39)

where p b 2 is the probability that an arbitrary neck is empty. It is of interest to note that in the case under consideration (i.e., when the pore volume is concentrated in necks) the desorption process is described [Eqs. (38) or (39)] by employing only the neck-size distribution and the mean coordination number for voids. The neck-size distribution can be calculated from the adsorption branch of the isotherm. Thus, the analysis of the desorption branch of the isotherm allows one to obtain the mean coordination number. In fact, zo can be obtained from the position of the desorption knee. In addition, using Eq. (39) and the scaling expression for 9 b 2 [Eq. (12)], it is possible to estimate the average linear dimension L of the microparticles in porous solids (34). G . RESULTS OF SIMULATIONS

The results of simulations demonstrating the effect of various factors on nitrogen desorption from porous solids are presented in Figs. 16-18. To describe the desorption process, we use Eqs. (24) and (30). The adsorption branch of the isotherm is described by Eq. (18). The size distributions of necks and voids are assumed to be lognormal: cp (r) = (z,N,/2) e ~ p { - [ l n ( r / ~ ~ , > ] ~ / 2 u ~ ) / [ r ~ ~ ( 2 r r ) ”(40) ~I

and f ( r ) = N , exp{-[1n(r/~~)]~/2a2}/[ru,(2rr)”~]

(41)

where T and u are the neck and void median and dispersion, zo is the coordination number for all voids, and No is the number of voids per unit volume (or mass) of the catalyst.

30

V . P . ZHDANOV

1 0.7 ct L

0.6 0.5 I

0.1

'

0

0.1

0.2

0.3 0.4 0.5 0.6 0.7 0.8 RELATIVE PRESSURE, P/Po

0.9

1

0.07 0.06 0.06 0.05 822 0.05

2

0.04 0.04

8

a vzvz

22 22

0.03 0.03 0.02 0.02 0.01 0.01 00 00

10 10

20 30 40 20 30 40 PORE RADIUS RADIUS [A) [A) PORE

50 50

60 60

FIG. FIG.16. 16. (a) (a)Adsorption Adsorption and anddesorption desorption curves curves atatz,z, = = 4, 4, 8, 8, and and 12. 12. (b) (b) Pore-size Pore-sizedistribudistribution: solid solid line, line, the the void void radius radius distribution,f(r)/N,, distribution,f(r)/N,, with with r, tion: r, == 30 30 AA and and uV uV= = 0.5; 0.5;dashed dashed line, line, the the neck neck radius radiusdistribution, distribution, 22pp((rr))//zz,,NN,,,, with with 7, 7, == 15 15 AA and and (T(T,, = = 0.5. 0.5.

31


0.1

0.09

0.01

0

0

10

20

30

PORE RADIUS

40

50

60

(A)

FIG. 17. (a) Adsorption and desorption curves at z, = 8, un = 0.5, and 7" = 10 and 15 A. (b) Pore-size distribution: solid line, the void radius distribution (the same as in Fig. 16); dashed line, the neck radius distribution with f. = 10 A and crn = 0.5.

32

V . P . ZHDANOV

1 0.9

0.8

3

0.7

m

!x 0.6

'

0.5 0.4

0.3 0.2 0. I

0.1

0.09 2

0.08 0.07

2

0.06 0.05

8

0.04

2 2

0.03 0.02 0.01 0 0

10

20

30

PORE RADIUS

40

50

60

[A)

FIG. 18. (a) Adsorption and desorption curves at zo = 8, 7" = 15 A, and a. = 0.3 and 0.5. (b) Pore-size distribution: solid line, the void radius distribution (the same as in Fig. 16): dashed line, the neck radius distribution with 7, = 15 8, and a, = 0.3.


33

All of the isotherms shown in Figs. 16- 18 have been calculated for the same void radius distribution (Fv = 30 A and crv = 0.5) and different neck radius distributions. The desorption branch of the isotherm is seen to shift to higher pressures with increasing zo (Fig. 16), T;” (Fig. 17), and a, (Fig. 18).

H. TYPICAL EXPERIMENTAL DATA The general equations presented in Sections II1,D-II1,F allow one to obtain structure information on porous solids from the adsorption/desorption data. In particular, if the pore volume is concentrated in voids, the void radius distribution can be determined from the adsorption branch of the isotherm. Then, postulating the value of z,, one can calculate the neck radius distribution from the desorption branch of the isotherm. To illustrate this approach, we have calculated (23) the void and neck integral distribution functions F(r) and @(r) [Eqs. (32) and (33)] for various porous solids using typical nitrogen isotherms. The void radius distribution has been derived on the basis of adsorption data employing the method of Broekhoff and de Boer (61)for spherical voids. The neck radius distribution has been calculated taking into account the overlapping of the void- and neck-size distribution [Eqs. (24) and (30)] or neglecting the overlapping [Eqs. (20) and (22)]. In addition, we have assumed that z, = 6. Figure 19 shows the results for thoria gel prepared by Wall and Brown (14). In this case, the overlapping of the void- and neck-size distribution is not very important. Figures 20 and 21 show the isotherms and the pore-size distribution for silica and titania gels (23). In the latter cases, the overlapping of the size distributions of voids and necks is significant. In particular, if the overlapping is not taken into account, the neck sizes are strongly overestimated. One can also note that an analysis of the desorption isotherm allows one to obtain the neck-size distribution only for large necks. The detailed data on the neck-size distribution for small necks cannot be determined, because with decreasing relative pressure the adsorption and desorption branches of the isotherm coincide (i.e., 9% becomes close to unity) when the Kelvin radius is larger than the radii of the lowest necks. Mason (20) has analyzed in detail the isotherms for xenon on Vycor porous glass (Fig. 14). The void radius distribution has been derived employing the adsorption branch of the isotherm. The void and neck radius distributions were assumed to be connected by the relationship given by Eq. (35). The latter assumption allows one to obtain from the desorption data the value of zo (for Vycor glass z, = 3.3) and to calculate the neck radius distribution (Figure 22).

34

V . P. ZHDANOV

FIG. 19. (a) Nitrogen adsorption and desorption isotherms for a sample of sintered thoria gel. (b) Pore-size distribution. Line I , according to Eqs. (20) and (22); line 2, according to Eqs. (24) and (30).

3 RELATIVE PRESSURE, P/P,

7

5 PORE RADIUS

(A)

FIG.20. (a) Nitrogen adsorption and desorption isotherms for silica gel. (b) Pore-size distribution. Line I , according to Eqs. (20) and (22); line 2, according to Eqs. (24) and (30).

35


1.o

,

,

,

,

[

0.0 0.0

0.4

0.2

0.6

0.8

1.(

t

2

RELATIVE PRESSURE, P/P,

6 PORE RADIUS

(A)

FIG. 21. (a) Nitrogen adsorption and desorption isotherms for titania gel. (b) Pore-size distribution. Line 1 , according to Eqs. (20) and (22); line 2, according to Eqs. (24) and (30).

6

E: 2

I

0

2

4

0

2

I

4

6

PORE RADIUS (nm) FIG. 22. Neck and void radius distributions for Vycor glass. The neck distribution has been calculated employing Eqs. (35) and (37). (From Ref. 20, with permission.)

36

V.P. ZHDANOV

If the pore volume is concentrated in necks, the neck radius distribution can be calculated from the adsorption branch of the isotherm. Then employing Eq. (38) [or Eq. (39)], one can determine zo from the desorption branch. For example, Seaton (34) has obtained zo = 5-7 for various alumina, silica, and alumina/silica samples.

IV. Mercury Penetration into Porous Solids We now consider application of percolation theory to describing mercury intrusion into porous solids. First we briefly recall the main physical principles of mercury porosimetry (in particular, the Washburn equation). These principles are treated in detail in many textbooks [e.g., Lowell and Shields (49)].The following discussions (Sections IV,B and IV,C) introduce general equations describing mercury penetration and demonstrate the effect of various factors characterizing the pore structure on this process. Mercury extrusion from porous solids is briefly discussed in Section IV,D. A. THEWASHBURN EQUATION The method of mercury porosimetry requires evacuation of the sample and subsequent pressurization to force mercury into the pores (49). This technique was originally developed to enable pore sizes to be determined in the macropore range, where the gas adsorption method breaks down for practical reasons (6). Application of mercury porosimetry is based on the Washburn equation (62,63), r, = -27 cos 8 / P

(42)

where y is the surface tension, 8 is the contact angle, P is the applied pressure, and r, is the corresponding capillary radius (Fig. 23). According to

FIG. 23. Mercury penetration into a cylindrical pore.


37

the Washburn equation, mercury can at a given value of P penetrate only into capillaries with a radius of Y > r,. Taking into account that for mercury y = 0.48 Nm-' and 8 = 140", one can rewrite Eq. (42) as r, = 7.3 x 104/P

(43) where r, is in angstroms and P is in atmospheres (1 atm = 0.101 MN m-'). The latter equation shows that a capillary of sufficiently small radius will require pressures considerably higher than 1 atm for mercury penetration (see Table IV).

B. MERCURY INTRUSION (APPROACHES 1 AND 2) During mercury intrusion, a given void or neck with r > r, can be filled by mercury only if it is connected with the outer surface by a chain of voids and necks with r > r,. Thus, mercury intrusion into porous solids is equivalent to the bond problem in percolation theory [Androutsopoulos and Mann (35), Wall and Brown (14), Chatzis and Dullien (36), Lane et al. (37), Zhdanov and Fenelonov (38), Tsakiroglou and Payatakes (39-41), Day et al. (42), and Park and Ihm (43)l. The equivalence is based on the identification of network sites with voids, and bonds with necks. A bond is considered to be unblocked if the neck radius r > r,. To describe mercury intrusion, one can use the same approaches as were employed in Section I11 for simulating condensate desorption from porous solids. In particular, if the pore volume is concentrated in voids (this model, shown in Fig. 2, has been analyzed in Refs. 14,37-41), the fraction of pore volume filled by mercury, Uin(rp),can be represented as (cf. Section II1,D or Ref. 38) Uin(rp)

= Uo(rp)%(Zq)

TABLE IV Pore Radius as a Function of Applied Pressure for Mercury Intrusion

P (atm)

rp(4

1 2

7.3 x 104 3.6 x 104 7.3 x 103 730 360 146 73 36 15

10

100

200 500 lo00 2000 5000

(44)

38

V.P. ZHDANOV

where U,(rp) is the fraction of pore volume corresponding to voids with r > rp, and %,(zq) is the percolation probability [the parameter zq is defined by Eq. (31)]. If, for example, voids are spherical, we have r 3 f ( r )d r / [

V o ( r p )=

r 3 f ( r ) dr

(45)

‘P

wheref(r) is the void radius distribution (Fig. 13). The other approach (Section II1,E) yields

/ F (rp) (46) where F(r,) is the fraction of voids with r > rp [Eq. (33)], and the parameter z,q is given by Eq. (23). The results of simulations demonstrating the effect of various factors on mercury intrusion into porous solids are shown in Figs. 24-26. All the intrusion curves presented have been calculated for the same void radius distribution [Eq. (41) with Tv = 3000 A and crv = 0.51. The mercury intrusion process is seen to start at higher pressures with decreasing zo (Fig. 24), F,, (Fig. 25), and a,, (Fig. 26). Equation (44) [or Eq. (46)] allows one to determine the neck radius distribution provided that the void radius distribution is known from independent experiments. For example, Fig. 27 shows the integral radius distributions for necks and voids, @ ( r ) and F(r) [Eqs. (32) and (33)], obtained for a model porous structure formed by a dense random package of glass balls having diameters of =250 pm (38).The account of interconnection of various pores is seen to lead to an essentially different distribution of necks over radii compared to the independent cylinder model: the distribution shifts toward small radii and its slope is less steep. It is also seen that the neck radius distribution is significantly affected by the overlapping of the void and neck radius distributions. From the data presented in Fig. 27, we can conclude that the approach based on percolation theory permits one to obtain the neck-size distribution only in a relatively narrow range of radii. This is due to the fact that the percolation probability p b ( Z q ) has a threshold ( ? h ( z q ) = 0 at zq < 1.5) and increases from 0 to 1 in a relatively narrow range of 1.5 < zq < 2.7. Equations (44) and (46) take into account the overlapping of the neck and void radius distribution. However, these equations have been derived employing the assumptions similar to those used in the mean-field approximation in statistical physics. In particular, Eqs. (44) and (46) ignore the effect of pore-size correlations on mercury intrusion. The latter effect has been recently studied in detail by Tsakiroglou and Payatakes (41) employing the Monte Carlo method. Simulations were made on a square lattice uin (rp)

=

uo (rp ) p b ( z o ~ )

39


1

5

0.4

0.2

0

50

100 PRESSURE

0

1

2

3

150

(ATMI

4

5

6

PORE RADIUS (1000 A)

FIG. 24. (a) Mercury intrusion curves at z, = 3, 4, and 6 . (b) Pore-size distribution: solid line, the void radius distribution,f(r)/N,, with F, = 3 F O A and cv= 0.5; dashed line, the neck radius distribution, 2 p ( r ) / z , N 0 , with 7. = 1000 A and crn = 0.5.

40

V.P. ZHDANOV 1

0.8

B

m

w 9 5

0

0.6

0.4

4 0.2 n 50

100

150

PRESSURE (ATMI FIG.25. Mercury intrusion curves at z, = 6, void radius distribution is the same as in Fig. 24.

a m

(T"

=

0.5, and F, = 500 and lo00

A. The

0.8 -

0.6 -

v)

m U

& 0.4

-

0.2 -

50

100

PRESSURE (ATMI

150

FIG. 26. Mercury intrusion curves at z, = 6 , 7" = loo0 A, and on= 0.3 and 0.5. The void radius distribution is the same as in Fig. 24.


0.0 1

1

I

I

1

1

30

40

50

60

70

PORE RADIUS

41

(pm)

FIG. 27. Integral radius distributions of necks and voids, @(r) and F ( r ) [Eqs. (32) and (33)] calculated from mercury penetration and photomicrographic data, respectively. Curve 1 , the model of independent cylinders with a length proportional to their radii; curves 2 and 3 , Eq. (44) without and with taking into account the overlapping of the neck and void radius distributions.

with three types of networks: uncorrelated, void- neck correlated, and void-void and void-neck correlated. It is found that, whereas the effect of void-neck correlation on mercury porosimetry curves is relatively weak, the effects of void-void and void-neck correlation are strong. The latter correlation widens the intrusion curve, extending it in the ranges of both low and high pressures.

C. MERCURY INTRUSION (APPROACH 3) In the previous section, we assumed that the pore volume is concentrated in voids, whereas the necks do not possess volumes of their own (Fig. 2). Androutsopoulos et al. (35,42,43)have considered an alternative model of porous solids, assuming the pore volume to be concentrated in necks [the combined model has been analyzed by Chatzis and Dullien (36)]. If the void volume is negligible, the fraction of pore volume filled by mercury can be described as [cf. Eq. (39)]

uin ( r p ) = Uo ( r p ) P b z (2, q)/q

(47)

42

V.P. ZHDANOV

where Uo(rp)is the fraction of pore volume corresponding to necks with r > r,, 9 b 2 ( 2 4 ) is the percolation probability for bonds (see, e.g., Fig. 7), z,, is the mean coordination number for all the voids, and q is the fraction of necks with r > r, [Eq. ( 2 2 ) ] . From Eq. (47) we can conclude that the effect of interconnection of various pores on mercury intrusion is very strong for large necks at relatively low pressures (when 24. < 2.5). Mercury penetration into large necks is blocked by small necks. For small necks at relatively high pressures (when z,q > 2.5), the percolation probability 9 b 2 ( z o q ) is close to q (see, e.g., Fig. 7b), and Eq. (47) yields the same result as the model of independent cylinders. D . MERCURY EXTRUSION Cumulative volume curves generated by intruding mercury into polrous solids do not coincide with those obtained as the pressure is lowered and mercury extrudes out of the pores. In all cases, the depressurization curve lies above the pressurization curve and the hysteresis loop does not close even when the pressure is returned to zero, indicating that some mercury is entrapped in pores (49). A general explanation of hysteresis in mercury porosimetry is based on the fact that mercury intrusion into large voids is often blocked by the void necks or by other neighbor small voids and necks and consequently takes place at relatively high pressures. On the other hand, mercury extrusion from large pores can occur only at low pressures corresponding to the pore sizes. If the porous solid is completely filled by mercury, the extrusion process will be quite difficult to initiate. For this reason, the initial states of mercury extrusion are dependent on the traces of residual air that remain in the smallest necks and voids. With decreasing pressure, voids containing menisci that must move under the existing conditions will be emptied. The latter is possible if these voids are connected with the outer surface through a chain of voids and necks filled by mercury. During these stages, many new menisci are formed. In addition, some filled voids or groups of filled voids become surrounded by empty voids. Mercury extrusion from these filled voids is impossible, i.e., mercury is entrapped [this mechanism of mercury entrapment was first analyzed in detail by Androutsopoulos and Mann ( 3 5 ) ] .An additional factor resulting in the formation of new menisci and mercury entrapment is the rupture of mercury in small necks with decreasing pressure. The latter effect has been discussed in detail by Tsakirogiou and Payatakes (39-41 ). Monte Carlo simulations of mercury extraction have been carried out on square ( 3 5 , 3 9 4 1 ) and cubic ( 4 2 ) networks. So far, however, universal


43

quantitative results in this field are absent. Some qualitative conclusions on mercury extrusion can be formulated as follows (39-41). (1) With decreasing mean coordination numbers, the degree of hysteresis and the quantity of trapped mercury increase. (2) The effect of trapped mercury on the extrusion process becomes more important with an increasing ratio FV/Tn,where r, and T;, are the mean void and neck radii. (3) As the degree of correlation between sizes of neighboring voids and neck increases, the quantity of trapped mercury decreases. (4) The intrusion and extrusion curves are strongly affected by the values of the intrusion and retraction contact angles. Exact determinations of these parameters are needed to obtain more adequate information on the pore structure.

V.

Catalytic Deactivation by Site Coverage and Pore Blockage

Many catalytic processes are accompanied by side reactions (for example, by coke formation) that decrease the catalyst activity. The deactivation of pore catalysts by coke formation is a complex phenomenon that includes coverage of active sites, simultaneously with coke growth and pore blockage (64-67). The latter phenomenon can be described employing percolation theory (44-46). A.

GENERAL EQUATIONS

For the analysis of deactivation kinetics, the following assumptions are customarily used (44-46): 1. The pore space is treated as a lattice of voids interconnected by necks in a three-dimensional network. The pore volume and surface are assumed to be concentrated either in voids (46) or in necks (44,45). Active sites are uniformly distributed on the catalyst surface. 2. The deactivation of a catalyst by coverage of active sites and pore blockage occurs under kinetic control. The conditions at the catalyst surface do not change with time. The process is isothermal. 3. The rate of disappearance of active sites is described by the first-order law

dc,/dt = -kc,

(48)

where csis the concentration of sites and k is the deactivation rate constant. 4. Pore blockage is assumed to be caused primarily by blockage of necks. The latter process is described by the time-dependent critical radius, rc(t). At time t , any neck with Y < r c ( t ) is assumed to be completely

44

V.P. ZHDANOV

blocked [by coke deposition on this neck or in other necks with r < r c ( t ) ] , and necks with r > r,(t) are blocked only partly. The simplest time dependencies of the critical radius are (46) rc(t) = ut

(49)

and r,(t) = w Vi

Equation (49) is applicable, for example, to coke growth through a polymerization process, the rate of which is independent of the degree of polymerization. Equation (50) is correct, e.g., if coke formation is limited by diffusion of reactants through the coke coverage. The rate constants k and o (or w) may be, in principle, interconnected if the disappearance of active sites is caused by formation of the coke precursor. However, we do not take into consideration any interconnection between these rate constants. Percolation theory can be applied to describe the deactivation process via identification of network sites with voids, and bonds with necks. A bond is considered to be unblocked if the neck radius r > r,. The deactivation of a catalyst by site coverage and pore blockage can be expressed in terms of relative activity, R, which is the ratio of the rate of the catalytic reaction at time t to the rate at time zero. Supposing the pore volume to be concentrated in voids (Fig. 2) and using assumptions 1-4, we have (cf. Section II1,D) R =

%(rc)9b(zq)

exp(-kt)

(51)

where 8 b ( z q ) is the percolation probability [the parameter zq is defined by Eq. (31), where r, should be replaced by r,], and %(rc)is the fraction of pore surface belonging to voids with r > r,. If voids are spherical, the latter value is defined by

%(rc)=

r 2 f ( r )d r / [

r2f(r) dr

(52)

wheref(r) is the void radius distribution. The other approach (Section II1,E) for the same problem yields R =

%(r,)9b(zoq)

exp(-kt)/F(r,)

(53)

where F(r,) is the fraction of voids with r > r, [Eq. (33)], and the parameter z,q is given by Eq. (23), where r, should be replaced by r,. If the pore volume is concentrated in necks (Fig. 3), the reaction rate can be represented as [cf. Eq. (3911 R =

%(.c)pb2(Zoq)

exp(-kt)/q

(54)

PERCOLATION THEORY AND KINEIJCS IN POROUS SOLIDS

45

where 9b2(zoq)is the percolation probability for bonds (see, e.g., Fig. 7b), q is the fraction of necks with r > r, [Eq. (22)], and 9 ( r c )is the fraction of pore surface belonging to these necks [the latter fraction is calculated by analogy with Eq. (52)].

B.

RESULTS OF SIMULATIONS

To demonstrate the effect of various parameters on the deactivation kinetics, we have used Eqs. (51) and (52) and a lognormal radius distribution of necks and voids [Eqs. (40) and (41), with tr, = a, = a , and Fig. 281. The deactivation kinetics for a linear dependence of the critical radius on time [Eq. (48)] are presented in Figs. 29 and 30. In particular, Fig. 29A shows that the disappearance of active sites makes a dominant contribution to the deactivation process at kF,,/u > 0.5. To stress the effect of pore blockage on the deactivation kinetics, we assume further that the contribution of the disappearance of active sites to the catalyst deactivation is negligible and set kTn/V = 0. Figure 29B shows the influence of the mean void radius on the deactivation process. An increase in the mean void radius results in a sharper threshold, below which the network loses its global connectivity. The effect of the mean coordination number on the deactivation kinetics is shown in

I

0

1

I

I

2 PORE RADIUS,

3

4

r/ Fn

FIG. 28. Lognormal size distribution of necks and voids at Fv/Fn= 2 and u = 0.5.

46

V.P. ZHDANOV

1.0

R 0.8 0.6

0.4

0.2

0.8 0.6

0.4

0.2

0

0.2

0.4

0.6

0.8

1.0

1.2

1.4

1.6

v t/ T" FIG. 29. Relative activity as a function of time at various values of the deactivation rate constant (A) and at various values of the mean void radius (B). The time dependence of the critical radius is described by Eq. (49).

Fig. 30A. An increase in the coordination number leads, of course, to a delay of pore blockage. The deactivation process is delayed also by an increase in the dispersion of pore-size distribution (Fig. 30B). If the time dependence of the critical radius is described by Eq. (50),the deactivation kinetics is not so sharp as in the case of Eq. (49) (cf. Figs. 29A and 31). However, the effect of various parameters on the deactivation process is qualitatively the same in both cases.

1.0

R 0.8

0.6

0.4

0.2

0.8 0.6 0.4

0.2

I

FIG.30. Relative activity as a function of time at various values of the mean coordination number (A) and at various values of dispersion of pore-size distribution (B). The time dependence of the critical radius is described by Eq. (49). 1.o

R 0.8

0.6 0.4

\\\

0.2

0

1

W 2 t /r,2

2

FIG. 31. Relative activity as a function of time at various values of the deactivation rate constant. The time dependence of the critical radius is described by Eq. (50).

48

V.P. ZHDANOV

VI. Conclusion Modern percolation theory contains many interesting problems that usually are easily formulated. Solving these problems is customarily connected with extensive Monte Carlo simulations. However, the final results are often universal and can be employed in various fields of physics and chemistry. The data discussed in the present review reveal that applications of percolation theory have been particularly successful for describing the kinetic processes in porous solids, including desorption of condensate, mercury penetration, and catalytic deactivation by site coverage and pore blockage. All these processes are of fundamental importance for heterogeneous catalysis. For this reason, knowledge of the main ideas and results of percolation theory seems to be essential for all scientists working in the field of adsorption and heterogeneous catalysis. To solve the latter problem was the main goal of the present review. ACKNOWLEDGMENTS My interest in percolation theory was inspired by Prof. V. B. Fenelonov (Institute of Catalysis, Novosibirsk). Jointly with V. B. Fenelonov and D. K. Efremov, we have published original papers (22,23,38)devoted to the application of percolation theory to describing desorption of condensate from porous solids and mercury penetration into porous materials. My study (46) of the catalytic deactivation problem was initiated by Prof. G. Zgrablich (University of Saint Luis, Argentina); Prof. D. Stauffer (Cologne University, Germany) supplied the data in Tables I and I1 prior to publication of the second edition of his book ( 1 2 ) . All of these contributions are gratefully acknowledged. I also thank all of the authors and publishers who have permitted me to reproduce their data. REFERENCES

I. “Studies in Surface Science and Catalysis. Vol. 39: Characterization of Porous Solids” (K. Unger, J. Rouquerol, K. S. W. Sing, and H. Kral, eds.). Elsevier, Amsterdam, 1988. 2 . “Studies in Surface Science and Catalysis. Vol. 62: Characterization of Porous Solids 11” (F. Rodriguez-Reinoso, J. Rouquerol, K. S. W. Sing, and K. K. Unger, eds.). Elsevier, Amsterdam, 1991. 3. Karnaukhov, A. P., in “Characterization of Porous Solids” (S. J . Gregg and K . S. W. Sing, eds.), p. 301. Soc. Chem. Ind., London, 1979. 4 . Karnaukhov, A. P., in Ref. 2, p. 105. 5 . Dubinin, M. M., Chem. Rev. 60,235 (1960). 6. Gregg, S. J., and Sing, K. S. W., “Adsorption, Surface Area and Porosity.” Academic Press, New York, 1982. 7. Broadbent, S. R., and Hammersley, J. M., Proc. Cambridge Philos. Soc. 53, 629 ( 1957). 8. Shante, V. K . S., and Kirkpatric, S., Adv. Phys. 20, 315 (1971). 9. Kirkpatrick, S . , Rev. Mod. Phys. 45, 574 (1973). 10. Essam, J. W., in “Phase Transitions and Critical Phenomena” (C. Domb and M. S. Green, eds.), Vol. 2, p. 197. Academic Press, New York, 1972.


49

1 1 . Essam, J. W., Rep. Prog. Phys. 43, 833 (1980). 12. Stauffer, D., “Introduction to Percolation Theory.” Taylor & Francis, London, 1985. (2nd Ed., 1992.) 13. Grimmett, G . , “Percolation.” Springer-Verlag, Berlin, 1989. 14. Wall, G. C., and Brown, R. J. C., J. Colloid Interface Sci. 82, 141 (1981). 15. Kheifets, L. I., and Neimark, A. V., “Multiphase Processes in Porous Media.” Khimiya, Moscow, 1982. (In Russ.) 16. Neimark, A. V., ColloidJ. 46, 927 (1984). (In Russ.) 17. Neimark, A . V., in Ref. 2, p. 67. 18. Mason, G., Proc. R. SOC.London Ser. A 390, 47 (1983). 19. Mason, G., J . Colloid Interface Sci. 95, 277 (1983). 20. Mason, G., Proc. R. SOC.London Ser. A 415, 453 (1988). 21. Mason, G., in Ref. 1 . p. 323. 22. Zhdanov, V. P., Fenelonov, V. B., and Efremov, D. K., React. Kinet. Catal. Lett. 32, 185 (1986). 23. Zhdanov, V. P., Fenelonov, V. B., and Efremov, D. K., J. Colloid Interface Sci. 120, 218 (1987). 24. Efremov, D. K., and Fenelonov, V. B., React. Kinet. Catal. Lett. 40, 177 (1989). 25. Efremov, D. K., and Fenelonov, V. B., in Ref. 2 , p. 115. 26. Palar, M., and Yortsos, Y. C., J . Colloid Interfare. Sci. 124, 162 (1988). 27. Palar, M., and Yortsos, Y. C., J . Colloid Interface Sci. 132, 425 (1989). 28. Mayagoitia, V., Rojas, F., and Kornhauser, I., J . C. S . Faraday I84, 785 (1988). 29. Mayagoitia, V., Rojas, F., and Kornhauser, I., J. C. S. Faraday 184, 801 (1988). 30. Mayagoitia, V., Cruz, M. J., and Rojas, F., J . C. S. Faraday 1 8 5 , 2071, 2079 (1989). 31. Mayagoitia, V., in Ref. 2 , p. 51. 32. Zgrablich, G., Mendioroz, S., Daza, L., Pajares, J., Mayagoitia, V., Rojas, F., and Conner, W., Langmuir 7, 779 (1991). 33. Yanuka, M., J . Colloid Interface Sci. 134, 198 (1990). 34. Seaton, M., Chem. Eng. Sci. 46, 1895 (1991). 35. Androutsopoulos, G. P., and Mann, R., Chem. Eng. Sci. 34, 1203 (1979). 36. Chatzis, I., and Dullien, F. A. L., Int. Chem. Eng. 25, 47 (1985). 37. Lane, A., Shah, N., and Conner, W. C., J . Colloid Interface Sci. 109, 235 (1986). 38. Zhdanov, V. P., and Fenelonov, V. B., React. Kinet. Catal. Lett. 33, 377 (1987). 39. Tsakiroglou, C. D., and Payatakes, A. C., J. Colloid Interface Sci. 137, 315 (1990). 40. Tsakiroglou, C. D., and Payatakes, A. C., in Ref. 2 , p. 169. 41. Tsakiroglou, C. D., and Payatakes, A. C., J. Colloid Interface Sci. 146, 479 (1991). 42. Day, M., Parker, I. B., Bell, J., Thomas, M., Fletcher, R., and Duffie, J., in Ref. 2, p. 75. 43. Park, C.-Y., and Ihm, S.-K., AIChE J. 36, 1641 (1991). 44. Sahimi, M., and Tsotsis, T., J . Catal. 96, 552 (1985). 45. Yortsos, Y. C., and Sharma, M., AIChE J. 32, 46 (1986). 46. Zhdanov, V. P., Catal. Lett. 9, 369 (1991). 47. The conventional symbols for the bond occupation probability and the percolation probability are p and P, respectively, rather than q and 9”;these symbols are not used in the present review to avoid confusion with the symbol for the pressure. 48. Kirkpatrick, S., in ‘‘Ill-Condensed Matter” (R. Balian, R. Maynard, and G. Toulouse, eds.), p. 321. North-Holland Publ., Amsterdam, 1979. 49. Lowell, S . , and Shields, J. E., “Powder Surface Area and Porosity.” Chapman & Hall, London, 1984. 50. Brunauer, S . , Deming, L. S., Deming, W. E., and Teller, E., J. Am. Chem. Soc. 62, 1723 (1940).

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51. Brunauer, S., Emmett, P. H., and Teller, E., J . Am. Chem. Soc. 60, 309 (1938).

52. Brunauer, S., “The Adsorption of Gases and Vapors.” Oxford Univ. Press, London,

1945. 53. de Boer, J . H., in “The Structure and Properties of Porous Materials” (D. H. Everett and F. S. Stone, eds.), p. 68. Butterworth, London, 1958. 54. Thomson, W. T., Philos. Mag. 42, 448 (1871). 55. Zsigmondy, A., Z. Anorg. Chem. 71, 356 (1911). 56. Anderson, J . S . , Z. Phys. Chem. 88, 191 (1914). 57. Foster, A. G., Trans. Faraday Soc. 28, 645 (1932). 58. Landau, L. D., and Lifshitz, E. M., “Statistical Physics.” Pergamon, Oxford, 1980. 59. Halsey, G . D., J. Chem. Phys. 16, 931 (1948). 60. Everett, D. H., in “The Solid-Gas Interface” (E. A. Flood, ed.), Vol. 2, p. 1055. Dekker, New York, 1967. 61. Broekhoff, J .C. P., and de Boer, J. H., J. Card. 10, 153 (1968). 62. Washburn, E. W., PhysRev. 17, 273 (1921). 63. The Washburn equation is a special case of the Young-Laplace equation. 64. Buyanov, R. A., “Coking in Catalysts.” Nauka, Novosibirsk, 1983. (In Russ.) 65. Butt, J. B., and Petersen, E. E., “Activation, Deactivation, and Poisoning of Catalysts.” Academic Press, San Diego, 1988. 66. “Studies in Surface Science and Catalysis. Vol. 34: Catalyst Deactivation 1987” (B. Delmon and G. F. Froment, eds.). Elsevier, Amsterdam, 1987. 67. “Studies in Surface Science and Catalysis. Vol. 68: Catalyst Deactivation 1991” (C. Bartholomew and J. B. Butt, eds.). Elsevier, Amsterdam, 1991.


Oscillatory Reactions in Heterogeneous Catalysis F. SCHUTH Institut fiir Anorganische und Analytische Chemie der Universitat Mainz 0-6500 Mainz, Germany

B. E. HENRY AND L. D. SCHMIDT Department of Chemical Engineering and Materials Science University of Minnesota Minneapolis, Minnesota 55455

1.

Introduction

Since the development of nonequilibrium thermodynamics in the late 1940s, initiated by the work of Prigogine ( I ) , numerous reports have appeared dealing with the possibility of oscillations in reaction systems far from equilibrium. Initially the main focus of these studies was the Belouzov-Zhabotinskii liquid-phase reaction (2), but since the discovery of oscillating reactions in heterogeneous catalysis in the late 1960s (3-7), over 300 publications have described research in this field as well. This review focuses on this emerging and important area of research. There are four primary reasons for studying oscillatory catalytic reactions. ( 1) Many surface reactions oscillate; probably all bimolecular oxidation reactions on transition metals will exhibit oscillations under some conditions. (2) Although studied for about 20 years, they are still intriguing and poorly understood phenomena, with no universally acknowledged mechanism. There is, however, a great deal of insight to be gained from the study of these phenomena, and the analysis of unstable steady states promisesanalogous to investigations of transient behavior-to reveal information about factors influencing catalyst performance that have not been discovered through the study of stable states alone. (3) Oscillatory states of reactors are potentially dangerous in chemical plants. In order to avoid these states, it is important to learn about the conditions favoring the occurrence of oscilla51

52

F. SCHUTH et

al.

tions. Together with stability analysis, experimental characterization of oscillating reactions is a valuable tool in understanding stable reactor operation. (4)Finally, in spite of the potential hazards of oscillatory states, there are in some cases possible benefits to operating reactors in unstable regimes [for a recent review on this topic see Matros (8-10)]. The first attempts at unstationary reactor operation were made when it was realized that feed cycling might yield better conversions or selectivities for certain reactions. This is, for instance, predicted theoretically for systems with a large Thiele modulus, wherein the effectiveness factor of the catalyst can be influenced by small changes in reactant pressures (11), or for a system described by a nonlinear Eley-Rideal expression, wherein pressure oscillations are predicted to enhance the conversion by up to 35% (12). Additionally, recent modeling results by Thullie and Renken (13) have shown that operation of a catalytic reactor to which the flow of a reactant is periodically halted may result in rates more than twice as high as the maximum rate under optimum steady-state conditions. Experimental data also show that unsteady operation can increase reactor performance. Studies were carried out in the SO2/O2/CZO5 system showing that pressure cycling of the feed could increase the SO3 yield by up to 25% (14). The CO oxidation reaction over V205showed a resonance peak of the reaction rate when sweeping the pressure-forcing frequency, during which the conversions were almost twice as high as those observed for the steady state (15). A system in which feed cycling occurs naturally is demonstrated in the automotive catalytic converter. Cutlip et al (16) found that pressure cycling during the CO/Oz reaction over Pt catalysts could yield higher conversions and were able to explain this behavior with an elementary-step model. Similar results were reported by Svensson et al. ( 1 7 ) . Experiments with actual catalytic converters and model feeds consisting of NO, CO, and O2 were carried out by Hegedus et al. ( l a ) , who found an appreciable increase in the conversion for cycling frequencies on the order of 1 Hz. Similar experiments were then performed with real exhaust gas, and it was shown that conversions of all three major pollutants (NO,, CO, and CH,) were higher under unsteady conditions as long as the reaction temperature was below the light-off point (19). In some cases, however, fluid dynamic effects can negate the kinetic advantage of cycling, as was shown by Baiker and Richarz for the C2H4/H2reaction on Ni (20). The examples above illustrate the benefits gained by unsteady operation. They are, however, only partially related to the phenomena dealt with in this review. The instabilities described above are externally introduced by forcing operation parameters, whereas oscillatory states in heterogeneous catalysis are inherently unstable. Because these autonomous oscillations usually arise as a Hopf bifurcation, wherein the stable state is completely lost,

OSCILLATORY REACTIONS

53

there is no stable state with which to compare the oscillatory rate. Thus there are few examples in the literature in which an autonomous oscillatory state can be described as advantageous when compared to a steady state. One of these examples is again the technically important catalytic converter. If sulfur is present in the engine exhaust gas because of impurities in the gasoline and the motor oil, all the SO2 formed during combustion is converted to SO3 over the catalyst-potentially leading to health hazards in urban areas. Olsson and Schoeoen (21) found that SO3 formation on Pt and Rh monoliths was largely suppressed when the CO/SO* oxidation was carried out under conditions in which autonomous oscillations occur. In this case, oscillatory operation has a positive effect on the selectivity of the catalyst. Operating in the oscillatory regime produced higher yields than did the steady state for the CO oxidation reaction over Pt (22) and Pd (23),and for the endothermic methylamine decomposition over Pt wires (24). In these cases comparison between oscillatory and steady states was possible because the unstable states coexisted with the respective stable states of the systems. For these reasons-their ubiquitous occurrence, lack of mechanisms, potential hazards, and possible performance enhancement of catalytic systems-the study of these phenomena has become an important part of catalysis research. Several reviews have appeared on the topic of oscillatory reactions in heterogeneous catalysis, although most have only dealt with the CO oxidation reaction (25-33). The most extensive reviews are those by Sheintuch and Schmitz (34) and in 1986 and 1987, by Razon and Schmitz (30, 35). The 1986 review, however, dealt only with CO oxidation on Pt, and the 1987 review focused only on the theoretical aspects of oscillations in heterogeneous catalysis. Other general reviews (28, 36-42) are either not very comprehensive or relatively old and lacking description of recent important results, such as the discovery of oscillations on single-crystal surfaces. However, two recent reviews have covered exclusively single-crystal surfaces. One by Imbihl (33) discussed oscillations of the CO/O2 system, and the other by Ertl(43), which appeared in the 1990 volume of Advances in Catalysis, discussed oscillations of both the CO/O2 and NO/CO reactions. Finally, a recently published book written by Gray and Scott (44) deals with the theoretical aspects of chemical oscillations and instabilities, with one chapter discussing heterogeneous systems. It is therefore appropriate to give an updated and extensive review on oscillations in heterogeneous catalysis. This work will focus mainly on the mechanisms that explain oscillations, considering both experimental support for these mechanisms as well as the mathematical models that describe them. The last section will consider work on spatial patterns and synchronization.

54

F. SCHUTHet

al.

II. A Survey of Oscillatory Reactions Oscillations have been found in numerous catalytic systems. At least 24 different oscillating reactions are reported in the literature, involving catalysts ranging from noble metal single-crystal surfaces to zeolites. Table I lists the reactions that have been observed to oscillate and also briefly indicates the catalyst, the pressure regime, and the type of the reactor in which the oscillations have occurred. to lo-* In Table I the high-vacuum (HV) range means a pressure of Torr; entries designated by Torr mean pressures between 0.1 and 10 Torr; flow refers to an unspecified steady-state flow pattern. It is apparent from Table I that there is a great diversity in the different oscillation conditions and catalytic systems. The pressures under which oscillations have been observed vary from lo-’ Torr for the CO/NO reaction on Pt(100) (141, 142) to atmospheric pressure for a large number of systems. The reactors used in these studies include ultrahigh-vacuum (UHV) systems, continuous stirred tank reactors (CSTRs), flow reactors, and reactors designed as infrared (IR) cells, calorimeters, and ellipsometric systems. Most of the reactions investigated are oxidation reactions, and the greatest number of publications is found for the CO oxidation reaction on noble metal catalysts, the most common being platinum. In most of the systems oxygen is involved, but there are some reactions that oscillate without the presence of oxygen, such as C2H4hydrogenation and CH3NH2decomposition. All reported reactions are exothermic, except for the CH3NH2decomposition reaction, which is endothermic by approximately 150 kJ/mol. Most of the publications report oscillations on noble metal catalysts, but numerous other catalysts have also been shown to display oscillations in certain reactions, i.e., other transition metals (Fe, Co, Ni, Ce, and Cu), oxides (Vz05. CuO and La203-BaO-MgO), and zeolites (KY, NaX, and ZSM-5). The catalysts were used in various forms, such as single crystals [Pt(100), P t ( l l l ) , Pt(llO), Pt(210), Pt(13,1,1), and Pd(llO)], wires, foils, plates, and powders and supported on oxides and zeolites. The wide range of reaction systems, catalysts, and reactors that exhibit oscillatory reaction rates reinforces the motivation for research in this field. Oscillations may be “lurking” in every heterogeneous catalytic system (one might speculate that every heterogeneously catalyzed reaction might show oscillations under the appropriate conditions), and it is crucial to know about this possibility when engineering a chemical process. Just as there are numerous systems that display oscillations, so too are there numerous ways in which the oscillatory behavior manifests. The observed oscillation patterns include approximately sinusoidal time series, relaxation-type oscillations (the reaction remains in one state for most of the

55

OSCILLATORY REACTIONS TABLE I Oscillutory Reactions in Heterogeneous Catalysis Reaction

Catalyst

Form

Pressure"

Reactor

co/02

Pt

( 100) (1 10) (2 10) (1 11)

HV HV HV HV/atm HV/atm Atm Torr Atm Atm Atm Atm Atm Torr Atm HV Atm Atm Atm Atm Atm Atm Atm Atm Atm Atm Atm HV HV Atm Atm Atm Atm

UHV system UHV system UHV system UHV system UHV system Flow CSTR Flow Flow Flow Flow Recycle CSTR Recycle UHV system Flow CSTR Flow Flow Flow Flow Flow Flow Flow Flow Recycle UHV system UHV system Flow Recycle Flow Flow

(23,79-83) (84) (81,85) (26,86-90) (22) (3,91-1 15) (I16- I 18) (I19,120) (I16,I 1 7) (121) (23) (122) (123-129) (17,124,130-134) (135,136) (137) (23) (138) (127) (127) (139) (140-145) (146,147) (91,148) (91) (148) (149-152)

Atm

Flow

(153,154)

Atm Atm Atm Atm Atm Atm Atm Atm

CSTR Flow Flow Flow Flow Flow CSTR Flow

C0/O2 C0/O2

Pd/Pt Pd

co/02

Rh

co/oz

Ir

co/02 co/o* co/oz CO/NO

Pd/Ce Pd/Co CUO Pt

CO/NO COINO/O2 H2102

Pd Pt Pt

H2/02 H2102 Hd0~

Pd Rh Ni

CH,OH oxidation

Pt

(13,l.l) Wire Wire Film Foil Foil Supported Supported Supported Supported (1 10) Wire Wire Supported Pd-zeolite Supported Ribbon Wire Supported Supported Supported Supported (100)

Wire Supported Supported Supported Wire Polycrystalline Polycrystalline Supported Wire Supported Foil Film Film Supported

Ref. (45-58) (45,46,58- 74) (75- 78)

(56) (561

(cont.)

F. SCHUTH et

56

al.

TABLE I (continued) Reaction

Catalyst

Form

Pressure”

Reactor

CH,OH oxidation CH,OH oxidation CzHSOH oxidation CzHSOH oxidation City gas combust ion C2H4 oxidation C2H4 oxidation CzHh oxidation

Pd

Supported Pd-zeolite

K-Y-zeolite

-

Atm Atm Atm

Flow Flow Flow

Pd

Supported

Atm

Flow Flow

cu

Supported

Atm

Flow

Pt

Wire Supported Supported

Atm Atm Atm

CSTR Flow Flow

Wire Supported Ribbon -

Atm Atm Atm Atm

Flow Flow Flow Flow

Pd

Powder Foil Supported

Atm Atm Atm

Flow Flow Flow

K-Y-zeolite

-

Atm

Flow

NaX-zeolite

-

Atm

Flow

Nay-zeolite

-

Atm

Flow

Pt Pt La@BaO-MgO Pt

Supported Supported

-

Atm Atrn Atm

CSTR CSTR Flow

Supported

Atm

Flow

Wire Wire Supported Wire -

Torr Atm Atm Torr HV Atm

CSTR Flow Flow CSTR UHV system Flow

Fe-zeolite Supported Supported

Atm Atm Atm Atm

Flow CSTR Flow Flow

Rh Pt

C& oxidation C& oxidation 1-Hexene oxidation C yclohexane oxidation Cyclohexane oxidation Cyclo hexane oxidation COIOZICdH, COIOzlC3Hs Oxidative CH4 coupling Partial propene oxidation by NO NHJ02

cue

NHjIOz NO/NH, NOIH? Methanol to gasoline CO/HZ

Pt/Rh Pt Pt ZSM-5

CO/Hz

Ag

Pt

Fe Pd

(100)

Ref.

57

OSCILLATORY REACTIONS TABLE I (continued)

Reaction CZHJH~ CzHJHz CZWGHZ CH3NH2 decomposition CH,NHz decomposition CHSNH2 decomposition Propylene oxide/Oz PhNO2 hydrogenation PhNO2 hydrogenation

Catalyst

Form

Pressure" Reactor

Ref.

Pt Ni Pd Pt

Supported Supported Supported Wire

Atm Atm Atm Tom

Flow Flow Flow CSTR

(223) (224) (225,226) (24,143,227,228)

Rh

Wire

Torr

CSTR

(24)

Ir

Wire

Torr

CSTR

(24)

Ag

Film

Atm

CSTR

(229)

cu

Supported

Atm

Flow

(230)

Ni

Supported

Atm

Flow

(230) ~

~~~~

"HV, High vacuum; UHV, ultrahigh vacuum; CSTR, continuous stirred tank reactor; Atm, atmospheric pressure.

time and deviates from this state in ignition or extinction spikes), period doubling sequences, quasiperiodicity , and chaotic time series. Figure 1 shows examples of these types of oscillations. Reaction rate oscillations may be accompanied by temperature oscillations [temperature fluctuations of up to 500 K have been reported (24)] or they may be isothermal. Isothermality occurs either because the catalyst can conduct heat away much faster than the rate at which it is produced by the reaction, as is the case in UHV studies, or because isothermal conditions are forced on the system by anemometry, as described in the work of Luss and co-workers (157). Oscillation frequencies can range from more than 10 Hz (24) up to periods of several hours (217,219). Often there is evidence for several time scales in a single oscillating system. Relatively regular highfrequency oscillations may be superimposed over relaxation oscillations (93,98), with the two types of oscillations caused by different changes on the catalyst surface. The details of oscillation patterns, especially for chaotic oscillations, are often very sensitive to changes in the catalyst sample caused, for example, by variations between catalyst batches or differing pretreatment conditions. However, it is usually possible to observe clear trends in the frequency and oscillation amplitude with the variation of particular reaction parameters. For instance, increasing reaction temperature normally leads to increased frequency and decreased amplitude of the oscillations if other parameters

58

F. SCHUTH

et a / .

1.8'

L

2

1

3

Time (min)

(b) T,IKI

1 f

460-

H

- t

10min

(c) 100

50

t (sec)

FIG. 1. Some examples of heterogeneous oscillations observed experimentally. (a) Sinusoidal time series for H2/02 on a Pt wire (from Ref. 41). (b) Relaxation oscillations for CO/Oz on Pt/AI2O3 (from Ref. 100). (c) Oscillations after a single period doubling for C O / 0 2on Pt(l10) (from Ref. 231). (d) Model and experimental (inset) quasiperiodic oscillations for NO/CO on supported Pd (from Ref. 232). (e) Deterministic chaos produced by a period doubling sequence for CO/O2on Pt(ll0) (from Ref. 231).

59


I 8.0 n (Y

0 0

6.0

8

Y

g!

4.0

3

2.0

0.0

I

;

0

I

I

I

800

1600

2400

Time [s]

I

3200

41 30

150

are kept constant. The influence of variation in the reactant inlet pressures, however, is more complex and has to be evaluated for each reaction investigated. For most of the reactions listed in Table I no bifurcation diagrams with respect to reactant pressures are presented in the references cited. However, for the frequently studied CO/Oz reaction on noble metals, many bifurcation curves with respect to CO pressure have been published. Figure 2 shows examples of such diagrams for atmospheric pressure oscillations and for high-vacuum oscillations. The atmospheric pressure oscillations arise as Hopf bifurcations with very small amplitudes at low CO pressures. With increasing CO pressure, the amplitude increases until the reaction reaches a single steady state on the low-reaction-rate branch. Similar behavior was also found for the CO/NO reaction on supported Pt catalysts (91). A more detailed discussion of different bifurcation diagrams would, however, exceed the scope of this review. The reader is therefore referred to the original publications for further discussion.

F. SCHUTH

et al.

Yo c02 /

/

/ / / /

/

i

2

1

I

3.9

3.0

3.7

-

3.6

PCO

3

3.5 1 1 0 . ~ Torr)

%CO

3.1

*


61

In the early studies of oscillatory reactions in heterogeneous catalysis, the question frequently arose as to whether oscillations were an effect inherent in the reaction dynamics, or if they occurred due to transport effects and impurities (233). It would appear, however, that this question has now been partially answered. Several experiments under thoroughly controlled CSTR conditions (7,22,93,97,111,234-241) and the observation of oscillations under UHV conditions rule out transport effects as the origin of oscillations. Experiments with high-purity gas streams suggest that impurities are not essential for the observation of oscillations. Furthermore, work performed under HV conditions clearly demonstrated that impurities in the reactant gas could quench oscillations (76). Sustained oscillations could only be achieved by carefully controlling and maintaining the purity of the reactant gases. However, even if impurities and transport effects are not at the root of the oscillations observed in heterogeneous catalysis, they may still strongly influence the macroscopically observable oscillation behavior. Relatively recent investigations have shown that impurities from the catalyst [i.e., Si (56) or C (26,86-89)] can drastically change the oscillatory behavior or may even induce oscillations. Other investigators, however, did not observe changes in the oscillatory behavior with different crystals having a different content of impurities (76). Transport effects can also play a significant role as high pressures and high space velocities are approached. The effects of mass transfer limitations on macroscopically observable behavior has been studied with theoretical investigations by Jensen and Ray (242,243) and those of heat transfer with theoretical investigations by Schiith et al. (232). These effects have also been experimentally observed, as will be discussed in Section V. The effects discussed here demonstrate that oscillations in heterogeneous catalysis are very complex phenomena. The closer an experimental system is to an industrial catalytic process, the more factors there are that have to be taken into account in order to truly understand the mechanisms leading to oscillatory reaction rates. Table I1 shows which factors influence the macroscopically observed oscillatory behavior of a catalytic system in different pressure regimes.

FIG.2. Bifurcation diagrams describing the behavior of the C O / 0 2system as CO pressure is varied. (a) COz effluent concentration (which is proportional to the reaction rate) as a function of CO inlet concentration at four different temperatures in an atmospheric reactor over a pulverized Pt/silica/alumina catalyst. Oscillation existence regions are indicated by vertical hatching (from Ref. 98). (b) Work function maxima and minima plotted as a function of CO pressure at 540 K on Pt(l10). The periodicity of the oscillations (as indicated above the curve) is seen to increase as CO pressure is decreased. (From Ref. 231.)

62

F. SCHUTH et

a/.

TABLE I1 Importance of Rate-Limiting Mechanisms in Different Pressure Regimes" Mechanism Surface reaction Mass transfer Heat transfer a

High vacuum

Torr

Atmospheric

+

+

+ + +

-

-

+I-

Significant mechanisms are indicated by a "+ ."

111.

Experimental Methods

The rich phenomenology of oscillatory reactions and the inherent differences in systems necessitate different experimental methods in order to probe each case. Not all of the experimental tools developed for the study of surface processes are useful for the study of oscillations on catalysts, because the means for studying these phenomena must meet certain requirements. First, the methods should be in situ methods, because it is difficult to ascertain otherwise that the catalyst is in the same state during the measurements as it was when the oscillations occurred. This means that the modern high-vacuum surface analytical techniques can only be applied to the study of reaction systems that exhibit oscillations in the high-vacuum range-a field opened by the discovery of oscillatory states during the CO/02reaction on Pt(100) in 1982 by Ertl et al. (55). Second, the in situ techniques should not disturb the oscillatory system. If they did somehow affect the surface processes, it would be difficult to distinguish between artifacts and the true autonomous oscillatory behavior (with the exception of forcing experiments). Methods that might lead to local sample heating, for instance, could trigger an oscillation wave starting at the heated spot and are thus not suitable (see, e.g., Ref. 244). A third important requirement is that the analytical technique must be sufficiently fast. If information about the oscillation mechanism is desired, different phases of an oscillation cycle have to be analyzed. This means that the analysis should provide a sufficiently high signal-to-noise ratio in about one-tenth of the time needed for the completion of an entire oscillation cycle. This restriction can be partially overcome by the use of multiplexing techniques, wherein information from different phases of the oscillation cycle are accumulated in separate channels. The long-term drift of the analytical system has to be low, however, and extremely regular oscillations are necessary in order to apply a multiplexing technique.

63


As early research on oscillatory reactions in heterogeneous catalysis began, little attention was given to the state of the catalyst surface. These first studies recorded the reaction rate by analysis of the product concentrations (see, e.g., Refs. 3,81) or by measurement of catalyst temperatures (3,162). Later, however, attempts were also made to monitor the catalyst surface during the oscillations, first by measurement of the work function (81), and later by methods such as infrared (IR) spectroscopy (108) and low-energy electron diffraction (LEED) for HV oscillations (245). Table I11 lists methods employed to study oscillations. In the following sections these methods will be discussed in more detail, beginning with generally applicable techniques, followed by a discussion of methods for HV and a brief description of forcing methods. However, no

TABLE 111 Experimental Techniques Used in Oscillatory Studies Experimental Method Work function/contact potential difference measurement

Reaction System H2/02, C0/02, NO/CO

Optical detection of thermal waves LEED

Desorption and titration methods Solid electrolyte potentiometry Auger electron spectroscopy Photoemission electron microscopy Rutherford backscattering Ellipsometry Thermogravimetry

(45,46,50,51,53-55,57,59,64, 65,66,68,76,78,81,121,144,

145,I54,I65,166,I68,170, 231,246-248) (18,80,86,87,89,90,91,93,10I, 108,112-114,119,123,125, 126,135,138,148,158,180, 200,251-254) (24,136,155,164,194,228,258)

Infrared spectroscopy

Pressure forcing

Ref,

(45,46,50,53,54,62,66,77,78, 144,145,231) ( I 7,18,45,47,63,65.67,92,103, 130,134,206,231,255,256)

C0/02, NHdOz, CO/N0/02, COICsHd02 c0/02

(60,71,86,98,11I )

H2/02, CO/Oz

(171, I 72,188,189,249,250)

NO/CO

(141,142)

c0/02

(57,58,73,74)

co/oz

(54)

H2/02 c0/02

(150,163) (257)

64

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a/.

special attention will be given to the techniques used to follow the conversion, because the standard tools for studying reaction kinetics (e.g., gas chromatography, dispersive and nondispersive IR techniques, temperature measurement with thermocouples, and mass spectrometry) are well documented.

A. TIME-RESOLVING TECHNIQUES The first method used to obtain surface information was the measurement of the work function (81). It is still an important technique today, although the focus has shifted from obtaining surface information to simply monitoring oscillatory states (59).The experimental setup is typically the Kelvin arrangement, in which the sample is one electrode of a capacitor, and a vibrating ring in front of the sample, usually made from oxidized tungsten, forms the other electrode. A change in the work function of the sample causes a change in the capacitance, which is then measured. In addition to the Kelvin setup, the measurement of photovoltaic currents on Si/SiOJPt electrodes has been successfully applied (81, 258). In oscillation studies, the most important effect on the work function is that caused by oxygen adsorption or complete oxidation of the sample, which leads to appreciable changes in the work function. However, the measurement of the work function reveals only limited information on the surface processes, thus the Kelvin probe is primarily used as an oscillation monitor because more specific methods are now available. Another method that basically measures the oxygen activity on the catalyst surface is solid electrolyte potentiometry, which has been applied in several studies (171,172,229,249,250,259).Figure 3 shows a typical setup for such measurements. Usually ZrO2 or Ti02 stabilized with 8% Yz03 is used as the solid electrolyte, with metal electrodes on both sides. The electrodes are either evaporated (260) or applied in the form of a metal paste (250).The leads are made of inert noble metals such as gold or silver in the form of a net or a wire. One electrode is exposed to air as a reference gas; the other electrode is exposed to the reaction gas mixture. The chemical potential of oxygen on the reference side is given by u r ( 0 2 )= u ! ( G )

+ RT In P ( 0 2 )

(1)

with P(02)being 0.21 atm. The chemical potential on the reaction side is ~ ( 0 2 = ) ~ ' ( 0 2 )

+ RT In ab

The activity a. is squared due to the dissociative adsorption of metals of interest. Combination of Eqs. (1) and (2) yields a. =

(0.21)0.5exp(2FEIRT)

0 2

(2) on the (3)

65


Reference electrode Reaction gases

Y,O,-doped Ti0,

Optional working electrode

Pt catalyst film

FIG.3. A schematic of a typical solid electrolyte potentiometry apparatus used to observe oscillations by electrochemical detection of oxygen activity.

Thus the electromotive force (EMF) measured in such an arrangement is related to the oxygen activity on a platinum catalyst. CO adsorbed on the metal surface also influences the chemical potential and therefore the EMF via the relationship given in Eq. (3). With a slightly different arragement, Okamoto et al. (261) were able to measure local currents, which allowed them to obtain independent information about CO and oxygen coverages. The physics behind these measurements is described in the references cited above. Another variation of the basic arrangement described in Fig. 3 was introduced by Yentekakis and Vayenas (249). By using a third electrode on the reference side they were able to adjust the oxygen activity at the reaction electrode by applying an external voltage. They found a strong influence of the external voltage on the oscillatory behavior of the CO oxidation on the Pt electrode. There are some.studies that did not use in situ measurements. These include surface titration and desorption experiments in which surface cover-

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ages in diffferent phases of the oscillation cycle were measured (60,71,98,111).These methods were used to study oscillatory reactions before the advent of IR spectroscopy. By stopping the oscillating CO oxidation reaction in different phases of the oscillatory cycle, CO desorption curves were obtained. With this technique it was shown for the first time that the CO coverage on the catalyst surface oscillates counterphase to the reaction rates (98,111). However, IR spectroscopy made these methods almost obsolete, because it allows in situ measurement of coverages and, additionally, reveals information about adsorbate states. There are in principle two different methods of IR spectroscopy: (1) transmission and (2) reflection-absorption IR spectroscopy (RAIRS) arrangements. The transmission setup is usually employed in studies of supported catalysts in which a suitable impregnated support material, such as Aerosils or different aluminas, is pressed into thin disks of about 0.1 mm thick and 10-50 mg catalyst/cm2. Reflectionabsorption IR spectroscopy is used for the investigation of oscillations on metal foils and single crystals. One limitation in studying oscillations on heterogeneous catalysts with IR spectroscopy is that the sampling speed required is relatively high. RAIRS has only been made possible by the advent of Fourier transform IR (FTIR) spectrometers, which allow the collection of up to 100 spectra per second. Other difficulties include the presence of only low concentrations of adsorbed species in the IR beam and superposition of gas-phase bands over the adsorbate bands. Because of these difficulties, IR spectroscopy has been employed for monitoring CO only during the COlO2 reaction (86,87,89-91,107,108,112-114,119,123,125,126,135,138,148, 158,251,253) and the CO/NO reaction (19,123,148) and to a lesser extent for monitoring NO during the CO/NO reaction (148), because these molecules have fairly high extinction coefficients on noble metals. For powerful FTIR spectrometers it has been estimated that as little as 5% of a monolayer of CO can be detected in a reflection-absorption experiment on a catalyst foil (90). It has also been shown that the oxygen coverage of a catalyst surface affects the overall reflectivity in the whole spectral range, and this effect has been exploited to monitor CO and oxygen coverages simultaneously (86,89). This effect, however, is much less pronounced than the CO absorption bands, and it is thus more difficult to infer coverages from these experiments. A somewhat related method that uses IR radiation-but instead measures black-body radiation to yield spatially resolved activity patterns on catalyst disks, ribbons, and wires-was first developed by Delrue et al. (258) and Schmitz et al. (155,262,263) and was later applied by Luss et al. (164,264) and Cordonier et al. (24). The details of this method will be discussed in Section III,B.


67

The methods discussed so far are in principle applicable under all pressure and temperature conditions. Modern surface analysis tools used in the study of clean, well-defined surfaces, however, require high-vacuum conditions, thus limiting their application to reactions that oscillate under these conditions. Currently these include only the CO/O2,the CO/NO, and the NO/H2 reactions. Another possibility is the use of UHV methods on samples that have been introduced into a high-vacuum system after stopping the oscillatory reaction. This, however, violates the in situ measurement requirement. The UHV method most frequently used in the study of oscillations is lowenergy electron diffraction. LEED is utilized for surface analysis and is usually combined with a Kelvin probe for work function measurements in order to monitor oscillations. This setup was extensively applied to the C0/02 reaction on Pt(100) by Ertl et al. (45,46,50,53,66,245) and was also used during the CO/NO reaction on Pt( 100) by Schwartz and Schmidt (141) and Ertl et al. (144). LEED permits the analysis of surface structures by generating the diffraction pattern from an electron beam impinging on singlecrystal surfaces. The long-range order of the underlying substrate [for instance Pt(100), the most frequently used system], as well as that of the adsorbate, is obtained. Originally LEED was a relatively slow method, with the diffraction pattern displayed on a fluorescent screen or on photographic film. However, combination of a LEED setup with a video camera has made much better time resolution possible (53). Vishnevskii and Savchenko utilized X-ray photoelectron spectroscopy (XPS) as a tool for the in situ study of oscillations (71). They used a differentially pumped XPS system to monitor 01s states during oscillations on Pt(ll0) at pressures up to 1 Torr. This allowed them to observe the existence of different oxygen states in different phases of the oscillation cycle. Miscellaneous other UHV methods are of minor importance only. Norton et al. (54) attempted to monitor the hex % 1 X 1 phase transition of Pt( 100) during C O / 0 2oscillations by Rutherford backscattering (RBS), but the results were not very conclusive. RBS, however, might reveal information about the faceting of the Pt( 110) surface during the course of oscillations. Some other UHV methods, such as XPS, have also been applied to obtain the steady-state information necessary for the interpretation of oscillatory results (69).

B. TIMEAND SPATIALLY RESOLVING TECHNIQUES The techniques discussed so far give sufficiently high time resolutions but are spatially averaging. Since the discovery that a catalyst does not necessarily oscillate in synchrony over its entire surface, spatially resolving techniques have become increasingly important. The simplest spatially resolving

68

F. SCHUTH et al.

arrangement is just a combination of nonresolving probes. Wolf et al. (113,148,252,265) combined several thermocouples and area-selected IR spectroscopy to detect wavefronts on catalyst disks during oscillating reactions. Another relatively similar method was employed by Imbihl et al. (59), who applied two Kelvin probes to monitor simultaneously oscillations at different spots on a Pt(ll0) surface. These methods, however, only monitor oscillations at discrete spots on a catalyst and do not allow a spatially continuous analysis. Such an analysis was applied first in the work of Schmitz et al., and later as others used IR emission for measurement of temperature distributions and utilized scanning arrangements in LEED and photoemission experiments. Schmitz et al. (155,262,263)used a high-resolution thermographic imager to analyze temperature-and hence activity distributions-during the H2102 and the C0/02 reactions on metal foils and supported catalyst disks. Figure 4 shows one of the patterns obtained with this method. A similar arrangement was used by Luss (164,264)in his experiments to analyze oscillatory and steadystate patterns during the H d 0 2 reaction, and by Kellow and Wolf (136) to confirm the data obtained earlier by the above-mentioned technique. This setup can be appreciably simplified-at the expense of resolution-if the thermographic imager is substituted by an arrangement of photodiodes. This technique was used by Cordonier et al. (24) during the study of oscillations of methylamine decomposition on noble metal wires.

FIG. 4. Spatial patterns on a Pt foil during the HdOz reaction. Thermograms are in time sequence from left to right beginning with the top row. The temperature scale runs from T < 68°C (1) to T > 122°C (8), with the maximum temperature region numbered on each panel. (From Ref. 155.)


69

The other two continuous spatially resolving methods, scanning LEED and scanning photoemission spectroscopy, were developed by Ertl et al. and can be applied only under HV conditions. In a scanning LEED setup, the system is configured similar to a normal LEED experiment. Two Helmholtz coils, however, are added to allow scanning of the primary electron beam over a certain fraction of the surface (50,53). One hundred data points are recorded within 10 s over a scanning area of 4 X 7 mm. The spatial resolution obtained in such an experiment is limited by the diameter of the incident electron beam. Very recently, scanning photoemission spectroscopy was developed as another spatially resolving tool and was used to study CO oxidation on Pt(100) (57,58) and Pt(ll0) (58,73,74). In such a setup, a UV beam is focused on a single-crystal surface while the total yield of photoelectrons is monitored, thus generating a map of the local work function for a specific area of a Pt(100) or Pt(ll0) single crystal during oscillatory reactions. In the more recent of these studies, lateral resolutions of -0.1 p m and temporal resolutions of -20ms were obtained. The resulting images show remarkable evidence of traveling and standing waves, solitons, rotating spirals, and other spatially inhomogeneous patterns. C. FORCING TECHNIQUES Forcing techniques are somewhat different from experimental techniques that monitor an autonomous oscillator and can be used to elucidate information about oscillating systems. In these experiments, one parameter, typically the concentration of one reactant, is periodically varied and the response of the oscillating system (resonance, subharmonics, or superharmonics) is recorded. Many such experiments have been performed ( I 7,45,63,92,134,206,256,266-269). In addition, forcing techniques in combination with simulations were used for model discrimination (206,267) to show clearly the difference in oscillation mechanisms between Pt(100) and Pt(ll0) (45,63,226) for the CO/Oz reaction, and to discriminate amplified statistical noise from deterministic chaos (104,270). In the latter case, stochastic signals with different amplitudes were applied rather than periodic forcing signals and Liapunov exponents were determined. Extrapolation to a noise amplitude of zero yielded a positive Liapunov exponent, indicating that the oscillations were indeed deterministically chaotic. The topic of forced autonomous oscillators is still, however, in an early stage of development, and future analyses may give more insight when the effects that govern the responses are better understood. Some insight into this area has been gained through studies that analyzed the effects of forcing systems of differential equations that predict oscillations, such as that performed by Cordonier et al. for the Gray-Scott autocatalator (271).

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et al.

There is certainly considerable opportunity for the application of other analysis techniques to the study of oscillations. It can be expected that with the development of more sensitive detectors and wider application of lasers, new analytical tools for the in situ of oscillations will provide greater understanding of these phenomena.

IV. Models of Oscillatory Behavior Although many oscillating reactions have been studied experimentally, as many or perhaps even more models describing the oscillatory behavior have been discussed. At present there is no universal mechanism that explains 0scillations in all heterogeneously catalyzed reactions. Every reaction has to be thoroughly investigated to discover its specific oscillation mechanism. There are, nevertheless, certain classes of models under which several oscillating systems can be considered. This classification, however, is difficult because there are many models and many oscillating reactions, and the merits and drawbacks of each model must be evaluated. Razon and Schmitz (30) restricted their review to the CO/Oz reaction on Pt and classified the models by indicating which elementary steps were included. This procedure resulted in an extensive table discussing (or at least suggesting) the underlying mechanisms. This approach is not feasible in this review because the number of systems considered here is so large that such a table would require numerous pages and would probably increase confusion rather than serve as a guideline for this field of research. We therefore choose a different way to present the various models and their interrelationships (Fig. 5 ) . This procedure was selected in part because our aim was to relate the models discussed to additional experimental evidence beyond mere kinetic studies. We thus distinguish on a first level between general mathematical models and reaction models for particular catalytic reactions. These general mathematical models will be discussed first and they will later be referred to in connection with the impact they have had on the formulation of the more specific experiment-related models. The specific models will be further subdivided into isothermal and nonisothermal models. This distinction is justified because mathematical modeling of a nonisothermal system involves a heat balance in addition to coverage equations (or reactor mass balances), and therefore introduces strong Arrhenius-type nonlinearities into the coverage equations. Nonisothermal processes are much more dependent on the reactor type and the form of the catalyst (supported, wire, foil, or single crystal). Thus these heat balance equations that describe them must take into account the type of catalyst and

MODELS

EXPERIMENTRELATED

ABSTRACT MATHEMATICAL

/ \

SURFACE REACTION

REACTOWREACTION

/ \

ISOTHERMAL

/ I \ ENERGIES

FIG.5 . The classification of oscillation models

NONISOTHERMAL

72

F. SCHUTH et

al.

the type of the reactor, whereas isothermal coverage equations can be formulated on a more general level simply by making assumptions about the processes taking place on the catalyst surface. In Section V, however, it is noted that pressure changes in the reactor gas phase can lead to synchronization of oscillations, thus implying that the reactor type has an influence on all oscillatory reactions. The isothermal models can be divided into elementary-step models wherein interaction of the elementary steps (adsorption, desorption, and reaction) can produce oscillations without additional mechanisms, and models with additional, non-Langmuir-Hinshelwood steps, such as phase transitions and oxidation-reduction cycles. The latter models are usually supported by experimental evidence obtained by the methods discussed in Section 111. In general it should be noted that isothermal models are not strictly isothermal, but would typically display oscillations under nonisothermal conditions as well. In the case of supported catalysts at atmospheric pressure, oscillations are probably never purely isothermal. Usually, however, thermal effects tend to amplify instabilities, so that a model that predicts oscillations for the isothermal case will most probably predict oscillations under nonisothermal conditions. The inherently nonisothermal models that require temperature variations for oscillatory behavior fall into two groups. In both cases, the reaction is blocked and reactivated at different temperatures. The blockage is caused either by a surface blocking or by a bulk-phase transition of the catalyst. In the following sections, these categories and several examples will be discussed in greater detail. Special emphasis will be given to the connection between modeling results, kinetic studies, and additional experimental evidence. This discussion will certainly not be complete but should be representative, given such a large and rapidly developing field. This review covers the most important classes of models that are currently under investigation. The reader will discover in what follows that most of the models were developed for the CO/Oz reaction on noble metals. This is due to the fact that these have been the most intensely investigated systems and the fact that sometimes contradictory results have been published, thus requiring adaptation and refinement of accepted models in order to explain new experimental findings. Further, other reactions are more complex and less understood than CO/OZ, thus researchers have characterized the oscillations in these systems, but often did not attempt to construct detailed models. On the other hand, mechanisms for oscillations in other systems are sometimes almost unanimously accepted, and this leaves no incentive for the development of competing models.


A.

73

ABSTRACT MATHEMATICAL MODELS

When models that describe oscillating reactions are presented, one can distinguish between abstract mechanisms, representing no specific observable reaction, and models for particular reactions that contain additional steps describing surface processes unique to the experimental system under consideration. In some cases, an abstract mechanism might have been developed for a certain reaction, but can be easily generalized to other cases, because no unique, system-specific steps are involved. These abstract mathematical models have become prototypes for classes of oscillation mechanisms often referred to in publications wherein a more detailed model for a certain reaction has been developed. Before beginning a detailed discussion of the models we shall set forth the conventions used throughout this discussion. When reaction steps are presented, we will refer to reactants as A, B, C, . . . and to products as P, Q, R, . . . ; vacant sites are denoted by an asterisk (*). An asterisk directly following a reactant or product indicates an adsorbed species, whereas subscripts represent different forms of one adsorbed species. The differential equation systems for the models discussed will not be given explicitly, but they can typically be derived from the set of reactions given. If additional information is necessary, such as whether gas-phase concentrations are balanced, it will be noted. The abstract models can be divided into two categories, each of which can be further subdivided into three classes (Fig. 5 ) . Some of the models consist of coverage equations only, and these models will be called “surface reaction models.” The remaining models use additional mass and /or heat balance equations that include assumptions about the nature of the reactor in which the catalytic reaction takes place (the reactor could be simply a catalyst pellet). These models will be called “reactor-reaction models.” Some of the models mentioned under the heading “surface reaction models” also incorporate balance equations for the reactor. However, these models need only the coverage equations to predict oscillatory behavior; reactor heat and mass balances are just added to make the models more realistic [e.g., the extension of the Sales-Turner-Maple model (272) given in Aluko and Chang (2731. Such models are therefore included under surface reaction models, which will be discussed first. 1. Surface Reaction Models

Oscillatory surface reaction models were first proposed in the second half of the 1970s and developed in two parallel strains. The buffer-step models originated with the work of Eigenberger (274) and models with coveragedependent activation energies were first proposed by Belyaev et al. (154).

74

F. SCHUTH

et al.

Only later were other models proposed that did not need either of the above-mentioned mechanisms to predict oscillatory behavior (275-277). Eigenberger (274) was the first to introduce a kinetic scheme that contained a buffer step and to discuss the stability and oscillatory behavior of this type of reaction mechanism. His model for the reaction 2A + B-, 2P (CO + f 02-+ COz) can be expressed as A+*=A* 2A*

+ B + 2* -+ 2P + 4* B+*eB*

(6)

Reaction (6) is the buffer step, which describes a slow, reversible adsorption of one reactant. With this model, relaxation type as well as relatively sinusoidal oscillations were predicted for certain parameter values. Eigenberger discusses the nature of several possible buffer steps and generalizes reaction (6) to C

+ k*

$

Ct*

(7)

In the scheme described by reactions (4)-(6), educt B is represented by species C in reaction (7) and k = 1. Species C could also be an adsorbed species, changing from a linearly bonded to a bridge-bonded form, or an inactive oxygen species during oxidation reactions. A generalization of the original Eigenberger model was later analyzed by Lynch and Wanke (278). If one considers reaction (5) not as an Eley-Rideal type of reaction but as a Langmuir-Hinshelwood reaction between adsorbed species A and adsorbed species B, one must replace reaction (5) with the following two equations: B 2A*

+ 2* e B** + B** 2P + 4* --j

(54 (5b)

The Eley-Rideal reaction (5) of the Eigenberger model, however, could also be considered as a limiting case of reactions (5a) and (5b) for low surface concentrations of B** and rapid equilibration of reaction (5a). Hence the model of Lynch and Wanke is a more general formulation of the original Eigenberger scheme. In the analysis of Lynch and Wanke, the predicted periods and amplitudes in the calculated time series and borders of the stability regions in the phase plane differed from those predicted by Eigenberger’s work. Both models, however, predict oscillations for relatively similar parameter values. Other models, though not closely related to the original one proposed by Eigenberger, can nevertheless be thought of as buffer-step models. Takoudis et al. (279) discussed several kinetic schemes, one of which was analyzed in detail:


(8)

A + *=A*

A*

+ B + 2*-t A*

3*

75

+P+Q

s C* + P

c**c+* The first two equations are similar to those of Eigenberger, but with only one adsorbed molecule of species A reacting in Eq. (9) instead of two as in Eq. ( 5 ) . The buffer species, C*, is in this case, however, weakly coupled to the main reaction by A* and P. Nevertheless, according to Eigenberger a weak coupling of the buffer species does not significantly alter his general scheme. The model analyzed by Takoudis et al. predicts a rich phenomenology with single and multiple steady states as well as the coexistence of limit cycles and steady states. Additionally, some of the oxidation/reduction models developed so far can be considered buffer-step models. The most extensive theoretical treatment of such a model was performed by Aluko and Chang (273,280,281), who investigated under nonisothermal conditions the model first proposed by Sales et al. (272).The original isothermal model proposed nominally for the CO/Oz system is able to predict oscillations and can be considered a surface reaction model. The model, described in the general terms of Chang and Aluko, consists of A+*=A*

+ 2* 2B* B* + A* -+ P + 2* Bz

(12) (13)

B* + BI*

+ A* +P + 2*

(16) Species BI* represents a less reactive species that acts as a buffer. The formation and removal of this buffer is slow compared to the main reaction given by Eq. (14). Further application of this model to the CO/Oz system will be discussed in Section IV,B. Chang and Aluko were able to predict the bifurcation and oscillatory behavior of the CO oxidation reaction and exploited the oscillatory data to obtain refined estimates for the reaction parameters of the CO/Oz reaction. The general feasibility of an oscillation model, such as that analyzed by Chang and Aluko, was discussed earlier by Riekert (282) in thermodynamic terms. He argued that oscillations could emerge if the existence of both metal and oxide was thermodynamically permitted during oscillatory conditions, but these species were kinetically unstable. However, no kinetic analysis was performed in this study. Figure 6a illustrates the general scheme for buffer-step models. B,*

F . SCHUTH et at.

76

(b)

(4 buffer

high rate; low buffer conc.; reactant is removed

f

removed

)

buffer is adsorbed

J

low rate; high butter conc.; reactant coverage grows

(4

and reaction of influencing species

high rate; low coverage of inlluencing species: adsorption of influence on rate and consumption of reactant

tow rate: high coverage of influencingspecies; reactants are adsorbed

high rate; high empty site conc.;

high rate; formation of oxide rate I; increasing CO coverage

low reactant conc. rate T; overcompensates increasing empty-site ettect; oxygen coverage reactants adsorb

reactant cow. is high; low adsorption rate low rate; low empty site conc ; high reactant conc.

(4

low rate; removal of oxide

(f)

high rate; high temperature

high rate;

PI 1x1 surface; high 0

2

sticking coeff.

surlace forms 1x1 phase

empty patches reven to hex

\

low rate; pt hex surface; low 02 sticking coetf.

rater; temperature T

J

\

high (low) rate; high temperature;

low (high) rate; low temperature; blocker is removed (formed)

low rate; low temperature

J

high rate; high temperature

I (T);

rate T (I); temp. T

surface becomes 0 covered; temperature falls

catalyst reduced;

low rate; low temperature

FIG.6. General remesentations of heterogeneous oscillatorv mechanisms. (a) Buffer-step model; (b) coverage-dependent activation energy; (c) empty-site model; (d) Sales-TurnerMaple model; (e) Pt(100) phase transition model; (f) Dagonnier model; (g) blocking/ reactivation model; (h) bulk-phase transition model.


77

The other important class of surface reaction models includes those with coverage-dependent activation energies. Such models have been analyzed in great detail, using several different functional forms of these dependencies. The first groups to propose a surface reaction model with coveragedependent activation energies were those of Belyaev and co-workers (154,162). They investigated a reaction scheme for the H2/02 reaction, which can be easily generalized to a system of equations of the following form: A+**A*

A*

B+*eB*

(18)

+ B* + P + 2*

(19)

The activation energy of reaction (19) was assumed to be linearly dependent on the coverage of the catalyst surface. Physically this can be interpreted either as an effect of the adsorbed species, which interact with one another, or as an effect of surface heterogeneity, which would also cause coveragedependent activation energies. Analysis of this model revealed regions where bifurcations and oscillations were predicted, and it thus became one of the first successful mathematical models in this field. A similar model was analyzed by Pikios and Luss (283). They analyzed the same set of reaction steps with the coverage-dependent activation energy interpreted in terms of surface heterogeneity. They derived criteria for the occurrence of oscillations as did Belyaev et al. (154,162). They also found a singular steady state, which became a limit cycle for values of the surface heterogeneity lying above a certain threshold value, and they performed numerical analyses of these oscillatory states. Recently Tambe et al. (284) extended this model and included two different types of adsorption sites for A and B, while permitting the conversion of sites from one type to the other. The authors used the same coverage dependency and the same parameters as Pikios and Luss (283). Introducing the possibility of adsorption on different sites generated a qualitatively new dynamic behavior for the system characterized by a finite amplitude/ infinite period bifurcation that yielded a homoclinic orbit. This new feature was observed when the equilibration between the two types of sites was slow compared to the other reactions. However, if equilibration is fast and the equilibrium constant is assumed to be one, this model is equivalent to the one discussed by Pikos and Luss (283). Several versions of coverage-dependent activation energy models with different stoichiometries and functional forms for the coverage dependencies were analyzed by Ivanov et al. (285) in 1980. They analyzed the following general representation of the original mechanism:

78

F. SCHUTH et

mAl/k*

al.

+ nB,/l* + op + (m + n)*

(22)

with the rate constant for Eq. (22) modified by the following functional forms:

f,(e) = e-'(lW

(23)

Ivanov et al. derived the criteria necessary for the u(i), which were dependent on the different reaction stoichiometries, and they performed numerical simulations. They found behaviors that ranged from sinusoidal oscillations to regular, well-formed relaxation-type oscillations (Fig. 7). Figure 6b shows a schematic representation of models with coveragedependent activation energies. For all the models that assume coveragedependent activation energies, however, one should keep in mind that the coverage dependencies can be very strong. The reaction rate may vary by up to 15 orders of magnitude between coverage zero and coverage one in some cases. There are some oscillatory surface reaction models that contain neither buffer steps nor coverage-dependent activation energies. The first model able to predict oscillatory behavior without these mechanisms was introduced by Takoudis et al. (275,279) and later was analyzed in great detail by McKarnin et al. (286).The model consists of the following three steps, the third of which is strongly autocatalytic:

A*

A+*=A*

(26)

B + * e B *

(27)

+ B* + 2* -+ P + 4*

(28)

The empty-site requirement in Eq. (28) can be physically interpreted in one of two different ways: either the adsorbed A* and B* have to rearrange prior to reaction, or they are bound to more than one adsorption site. For the latter case, the intermediate concentration is low, thus allowing a pseudo-steady-state assumption. Through the application of bifurcation analysis and catastrophe theory this model was found to predict a very rich bifurcation and dynamic behavior. For certain parameter values, sub- and supercritical Hopf bifurcations as well as homoclinic bifurcations were discovered with this simple model. The oscillation cycle predicted by such a model is sketched in Fig. 6c. This model was also used to analyze how white noise would affect the behavior of an oscillatory reaction system

79


x+3 190 0)

.

2

5

0

0

0

2 .I30

.22

x , .26

3

.22

0

2

4

6

0

8

~

1 0 1 2 1 4

t x 10.2

.56u7L-

x .40

.24

0

2

4

6

t

8

1 0 1 2 1 4

I104

FIG. 7. Calculated time series for an isothermal coverage-dependent activation energy model showing both sinusoidal (left) and relaxation (right) behaviors. (From Ref. 285.)

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F. SCHUTH et

al.

(287). The authors found that stationary states could become oscillatory if the correlation time of the noise was of the same order of magnitude as the time constants of the model. Takoudis et al. (275) also proposed a mechanism that was later treated by Fedotov et al. (276). This model also uses two vacant sites in the reaction step, but B reacts via an Eley-Rideal step to form AB*, which then produces P: A + * =A* (29)

+ B Z AB* + 2* -+ P + 3*

A* AB*

(30)

(31) This mechanism also predicts oscillatory behavior. Although Eqs. (26)-(28) and (29)-(31) are quite simple systems, there is an even simpler model that predicts oscillatory behavior. Fichthorn et al. (277) used a Monte Carlo simulation method to model the simplest threestep Langmuir-Hinshelwood mechanism possible: A + * e A *

(32)

B+*+B*

(33)

+ B* -+ P + 2*

(34) Such a model, however, should be stable for all parameters according to the Bendixon theorem (288). Nevertheless, Fichthorn et al. found chaotic oscillations that were sensitive to the desorption probability of one species. This observation, however, is an effect brought about by the finite size of the simulation grid investigated and the lack of synchronization in the system. Increasing the size of the grid on which the simulations were performed led to decreasing oscillation amplitude and would ultimately lead to the disappearance of oscillations on large, perfect surfaces. For a large enough grid, the Monte Carlo simulations would give results similar in character to a description of the system with differential equations. In the latter situation, the Bendixon criterion predicts stability for all cases. Whether these findings indicate that small portions of a surface with different local adsorption properties (perhaps due to defects) can trigger oscillations over the whole surface as discussed by Fichthorn et al. (277) is not clear yet and requires further investigation. There is, however, another system for which Monte Carlo simulations and analyses of differential equations predict oscillations (289). Because this model uses a pressure balance for the reactor, it will be discussed below as a reactor-reaction model. A*

2. Reactor-Reaction Models

Vlachos et al. (289) performed Monte Carlo simulations, modeling a unimolecular surface reaction that had been investigated earlier in the mean-


81

field approximation by Kevrekidis et al. (290). The model describes a unimolecular surface reaction A + * S A * (35) A* -+ P

+Q+*

(36)

and includes the CSTR equation for the reactor pressure in addition to the coverage equation. Kevrekidis discusses several functional forms for the coverage dependencies of the adsorption, desorption, and reaction steps. It was found that such a model can predict oscillations only for cases in which the desorption rate has a coverage dependence of the form e-"'. Interestingly, these cases are those for which multiple thermodynamic phases are also predicted. These same results were obtained with Monte Carlo simulations by Vlachos et al. (289).In addition, they studied the influence of surface diffusion and surface defects and found that these effects do not completely alter the system behavior but can change the size of limit cycles or alter the position of bifurcation points. A mass balance for the reactor was also included in the simple LangmuirHinshelwood model analyzed by Morton and Goodman (291) with constant activation energies: A, + n* nA* (37) 0, oA*

+ m*

+ pB*

-+

rn B*

P

+ (0 + p)*

(38) (39)

Morton and Goodman determined whether oscillations could be predicted for various values of m,n, 0,and p . Their analyses showed that even for the simple Langmuir-Hinshelwood mechanism of the C O / 0 2reaction, oscillations are in principle possible if the reactor mass balance is taken into account. However, the amplitudes of the resultant numerically simulated oscillations were very small. Two models that utilize gas-phase mass balances were proposed by Sheintuch (292).The first, a model for oxidation reactions, was related to an Eigenberger-type buffer-step model. His kinetic scheme differs, however, from the typical buffer scheme: A + * e A * (40)

+ * -+ A** 2A* + A** + A A*

Bz A**

(41 4

+ 2* + 2B* + B* + P + 3*

As reactions (40)-(43) show, the species A* is the buffer in this scheme. In

contrast with the buffers in the mechanisms discussed earlier, in which these

82

F. SCHUTH et

al.

species were formed in a reaction parallel to the main reaction, in this case A* is an intermediate and thus is involved in the reaction mechanism in a different way. Sheintuch discusses this mechanism assuming the existence of both linearly bonded and bridge-bonded forms of A, as did Eigenberger. However, in reactions (40)-(43) the bridge-bonded form is not directly produced by adsorption but rather via the linearly bonded adsorbate. Nevertheless, the Sheintuch model predicts oscillations under certain conditions, and the residence times that yield oscillatory states are in a realistic range for flows past catalytic wires or single-catalyst pellets. The second model discussed by Sheintuch (293), which incorporates a reactor mass balance, is similar to the Sales-Turner-Maple model mentioned in the previous section, but with reactions (14) and (16) changed to B*

+ 2A*-

P

+*

BI* + 2A* (on BI* sites)

-+ P

(14a)

+*

(1 6 4

Reaction (16a) is very slow and is included in the model only to prevent complete blockage of the surface by the B I * species. This model also predicts oscillations and in addition was analyzed for spatial effects. On sufficiently long surfaces, Sheintuch found the existence of symmetry breaking that would in turn result in spatial structures. In a review article on oscillatory reactions (294), Sheintuch discusses the effect of introducing a heat balance for the catalyst rather than a mass balance for the reactor into the differential equation system for a surface reaction with oxidationheduction cycles. Although the coverage equations alone can yield oscillatory behavior, as was the case for the models discussed in the previous section, Sheintuch’s model is discussed in this section because introduction of the heat balance adds qualitatively new features. In this extended system complex, multiple peak behavior and quasiperiodicity was observed as shown in Fig. 8. Sheintuch also investigated the interaction of two oscillators. This work, however, will be treated in detail in Section V, were synchronization and chaos are discussed. Dagonnier et al. (295) also investigated the inclusion of the catalyst temperature into the set of differential equations describing a catalytic reaction. These authors, however, specified an a priori variable, the surface temperature, the meaning of which is not well-defined. The reaction mechanism they present consists of bimolecular reactions as described by Eqs. (44)- (46): (44)

A+*$A* B, A*

+ B* A

--j

AB

+ n* + nB*

+ 2*

+ B* + AB + *

(Langmuir-Hinshelwood) (Eley-Rideal)

(45) (464 (46b)

83


IX

X

B steady

state @upward @dawn

+m rw, -t

FIG. 8. Examples of the types of oscillations arising when a heat balance is included with a simple Langmuir-Hinshelwood differential equation system. Oscillatory time series are shown together with their corresponding projections on the x = 0 plane. (From Ref. 294.)

Stability was investigated for n = 1 and n = 2 in both the LangmuirHinshelwood and the Eley-Rideal cases. The results indicate that the EleyRideal reaction mechanism tends to stabilize the system whereas the Langmuir-Hinshelwood mechanism easily leads to oscillatory behavior. A serious problem with this analysis, however, is that the physical interpretation of the notion of a “surface temperature” in the first few layers of catalyst remains unclear. A much simpler surface reaction model used in conjunction with a heat balance for the entire catalyst unit (wire, foil, or pellet), rather than for a surface layer, was discussed by Ray and Hastings (296).This paper was the first of a series of publications from this research group, culminating in the classic papers of Jensen and Ray (243,297). They analyzed the simplest catalytic reaction conceivable: that having a single adsorption step and a single reaction step, A+*-,A* (47) A*+P+*

(48)

84

F. SCHUTH et

al.

They treated the system much like a CSTR, with the balance for the gasphase concentration substituted by the coverage equation for the catalyst. Ray and Hastings then applied the analytical treatment that they had developed for the CSTR in this same publication. Stability analysis revealed that the critical Lewis numbers for oscillations were in a range that did not allow for oscillations on normal nonporous catalytic surfaces. However, as Jensen and Ray (243) showed, a certain model for catalytic surfaces, the “fuzzy wire” model, with the assumption of a very rough surface with protrusions is able to produce Lewis numbers in the proper range for the occurrence of oscillations. This model, however, included both mass and heat balances as well as coverage equations, thus combining the two classes of reactorreaction models discussed above. The problem of a porous catalyst pellet, which had been addressed in the paper of Ray and Hastings, was later treated extensively by Jensen and Ray (242,297). They used surface coverage equations and mass and heat balances for the whole pellet, all of which, except for the heat balance, were solved for the nonlumped case. The solutions of the resultant partial differential equation set were obtained by collocation techniques. The surface reaction was assumed to be unimolecular and slightly more complex than the mechanism analyzed by Ray and Hastings in that the adsorption step was permitted to be reversible: A + * e A *

(49)

A*+P+*

The coverage equation for A , the crystallite temperature, and the gasphase concentration in the void volume of the support pellet were modeled with partial differential equations. The temperature of the pellet was assumed to be uniform throughout the pellet, and thus description with an ordinary differential equation was sufficient. In-depth analysis of this model yielded rich phenomenology, which, because of its complexity, will be discussed in Section V. The models discussed above were in general developed to study the dynamic behavior of systems often only loosely related to heterogeneous catalysis. The following section describes models, often derived from those previously discussed, which were developed for particular reaction systems.

B. MODELSFOR SPECIFIC EXPERIMENTAL SYSTEMS In this section we will discuss models that were developed to explain oscillations in specific experimental systems. These models are frequently supported by additional experimental evidence for the processes assumed to be occurring on the catalyst surface. The structure of some of the previously


85

discussed abstract mathematical models can often be found in the models discussed in this section. In these cases, reference will be made to the corresponding models of Section IV,A. We will first discuss isothermal models that describe oscillations with elementary-step kinetics for a certain reaction, then the models that incorporate additional experimental evidence and usually non-Langmuir-Hinshelwood surface processes, and finally the somewhat more complex nonisothermal models. Unfortunately the models that follow are not always expressed in terms of differential equations, so that a theoretical investigation is possible only in some cases. 1. Elementary-Step Kinetics Historically the first oscillatory models developed fall into the elementary-step kinetics category. These are the models proposed by the group of Belyaev et al. in a number of experimental and theoretical papers (154,162,167,298,299).This type of model was discussed in Section IV,A with those that utilize coverage-dependent activation energies. The feature that induces instability into the Langmuir-Hinshelwood kinetics is the linear dependence of the surface reaction activation energy on the concentration of the reactants on and underneath the surface. This model was discussed by Belyaev et al. for the HO/Oz reaction on metal catalysts, with the concentration of H and 0 dissolved in the first surface layers influencing the activation energy for water formation. There is indeed experimental evidence--from contact potential difference (CPD) measurements made during the oscillations-that changes in the surface layer are taking place (154,165); however, it has not yet been established where the changes in the surface state have the influence assumed in the proposed models. Another elementary-step model using coverage-dependent activation energies was proposed by Fink et al. for the NO/CO reaction on Pt(100) (144,145).Oscillations were observed in this system in two distinct temperature regions, one associated with the hex 1 X 1 platinum phase change, and the other produced by coverage-dependent desorption activation energies and the autocatalytic generation of empty sites. Models were developed for both regions; however, the phase-change model will be discussed later in Section IV,B,2,b. The reaction mechanism was composed of

co + * e co* NO + * s NO* NO*

+*

---f

N*

+ O*

+ 2* co* + o* + co* 2N* -+ N2

(51) (52)

(53) (54) (55)

86

F. SCHUTH et

al.

Repulsive interactions between NO and CO molecules were assumed and the activation energy for desorption was given by E,,.No,co(O) = EaCt(O)- k . (ONO + Oc0)’. Although oscillations similar to those observed in experiments were predicted by this model, no justification was given for the unusual quadratic dependence on surface coverage for the desorption activation energy, and experiments in which CO and NO have been coadsorbed on Pt(100) have shown small or attractive interactions between the adsorbates (300). Kinetic schemes were also developed for CO oxidation with 0 2 both alone and in the presence of alkenes. Eigenberger-type models were analyzed by Yablonskii et al. (301,302) for the C 0 / 0 2 reaction either with the adsorption of an inert species or with transformation of adsorbed oxygen to a less reactive form. Mukesh et al. discussed elementary-step models for the C0/02/C3H6(207) and the C0/02/C4Hs(204) systems on supported Pt catalysts. Their models consist of steps that describe the adsorption and desorption of each species and the reactions between adsorbed alkene and adsorbed CO with adsorbed oxygen. In addition, alkene and COZadsorption on the support were taken into account. In principle, this model is a variant of the buffer-step models discussed in Section IV,A, with the alkene playing the role of the buffer. Their model fits experimental time series quite well (Fig. 9), but the authors only presented results for one set of conditions. Typically it is difficult to get a satisfactory fit over a significant portion of the parameter space. Nevertheless, given that the model is crude (alkene oxidation is modeled as C3H6** + 90* + 3 c o 2 + 3Hz0), and that the assumption of isothermal conditions fails at high conversions, it is surprising that the experimental time series agrees with model predictions at all. A similar model was proposed by Morton and Goodman (205) for the CO/Oz/C4Hs system. Another isothermal elementary-step model was developed by Takoudis and Nowobilski (303). They modeled the oscillatory reaction between NO and NH3 on Pt with the following reaction equations: NH3

+ * * NH2* + H*

(56)

+ (precursor)

(57)

NO

+ * NO* Nz + HzO + 2* NHI* + NO* H* + NO* ---* HNO* + * 2HNO* + N20 + HzO + 2* HNO* + H* -+ N* + HzO + * 2N* N2 + 2* NO(precursor) -+

(58)

(59)

(60) (61)

(62)

(63)

87


I

a

C

K4/K4=3.1 x 104/0.0471

3-

lb 1

10

20

Time (min)

0

10

20

Time (min)

FIG.9. Experimental (a and b) and simulated (c and d) partial pressure oscillations for the CO/Oz/propylene reaction (a and c), and the CO/Oz/l-butene reaction (b and d), both over Pt. The simulated curves are obtained by an elementary-step model. (From Ref. 207.)

Although this model is able to predict oscillatory behavior, it has one major flaw, because a crucial step in the model is the adsorption of NZ on Pt to yield adsorbed N atoms on the Pt surface. This reaction step is not supported by an existent experimental evidence. This crucial step can also be interpreted as a buffer step as discussed in the general schemes in Section IV,A, and thus omission of this step would probably destroy the oscillatory behavior of the model predictions.

2. Models with Additional Surface Effects a . OxidationlReduction Models. In a feed that consists of only two reactants, strictly there can be no blocker species as in the models discussed previously. If one reactant, however, is adsorbed in both a reactive and a less reactive or nonreactive form, then the conditions for application of buffer-step kinetics are again achieved. This possibility was already pointed out by Eigenberger (274). The models of Sales et al. (23,79,272), Lindstrom and Tsotsis (93), Yeates et al. (56), and Vayenas and co-workers (94,250) are all oxidationireduction models of this type. The structure of the model equations by Turner et al. (272) is similar to the buffer-step models discussed earlier, particularly the model analyzed in general form by Chang and Aluko [Eqs. (12)-16)]. The resultant numerical

88

F . SCHUTH

et al.

simulations are probably the best fit of oscillatory time series currently available (Fig. 10). A particularly attractive feature of this model is that the simulations approximate experimental results at different temperatures. Shobukhov and M. M. Slin'ko investigated the bifurcation behavior predicted by this model not only with respect to temperature variation, but also with respect to variation of oxygen and CO partial pressures; they could reproduce the experimental dependencies of the bifurcation patterns on these variables (304). The principal features of these mechanisms are shown in Fig. 6d. The use of IR spectroscopy and solid electrolyte potentiometry has provided additional experimental evidence for these types of models. Lindstrom and Tsotsis (93,120)report an absorption band around 2120 cm-' for the CO/02 reaction on Pt in addition to the typically observed Pt-CO band around 2070 cm-' . This new band was assigned to CO molecules adsorbed on an oxidized Pt surface. The intensity of this band oscillated counterphase to the reaction rate, but with a much lower amplitude than the 2070-cm-'

l o ) Eiperimenlol

(b)Theorelicol

1.550 K

T:550K

Tz543K

1=543K

T:536K

1~536K

T1529K

1.529 K

1:522K

1.522 K

0

10 Time

20 (rnin)

0

10

20

Time(min1

FIG. 10. A comparison of the experimental and predicted oscillatory COz production time series in the CO oxidation reaction over Pt. The theoretical curves are generated by the Sales-Turner-Maple model. Note the close fit in both period and curve shape over the 522-550 K temperature range. (From Ref. 272).


89

band, which is characteristic of CO adsorbed on a metallic Pt surface. The intensity oscillations of the 2120-cm-' band were explained by the formation of an inactive 0 species that blocks the reaction, as was assumed in the Sales-Turner-Maple model. Vayenas and co-workers (94,189,250), using solid electrolyte potentiometry, found a high oxygen activity on the catalyst surface in the low-reaction-rate state during CO and C2H4 oxidation. They hypothesized that this behavior is caused by the formation of a catalytically inactive PtO, surface layer. Turner and Maple gave additional evidence for the oxidation/reduction model in some of their later work (257). They measured the kinetics of the surface oxidation and reduction of Pt, Ir, and Pd with a gravimetric technique in order to compare experimental data with the values used in their previous theoretical treatment. The measured reduction rates agreed well with the assumed theoretical rates, although the observed oxidation rates were about one order of magnitude lower than the values predicted by the model simulations. However, considering the rate-enhancing effects of impurities such as Si, as suggested by Yeates et al. (56) and Keck and Kasemo (149), the rates are fairly close to the simulation parameters. Although there are problems associated with such oxidation/reduction models-several authors did not observe the 2120-cm-' IR band (91,208,114) and high predicted frequencies in the range of seconds (91) are too fast for an oxidation/reduction reaction-the oxide mechanism is one of the most widely accepted models for the atmospheric pressure oscillations of the CO/OZreaction. Oxidation/reduction cycles have also been used to explain the results of CO oxidation studies on Pd in a zeolite matrix (305). The authors found that the oscillatory behavior depended strongly on whether the pretreatment of the catalyst involved oxidation or reduction conditions. It was argued that the high dispersion of the metal, facilitated by the zeolite, was favorable for fast oxidation/reduction cycles. An oxide model has also been developed for CO/Oz oscillations on the (110) single-crystal plane of Pd (121,248). Oscillations were found in the pressure range between lop3and 1 Torr, and a transition to chaos via a period-doubling sequence was observed. The model assumes the formation of a subsurface oxygen layer that blocks further oxygen adsorption. The surface then becomes covered by CO and the surface reaction is blocked. However, CO slowly reacts with and removes the subsurface oxygen until oxygen adsorption is again possible, yielding a high-reaction-rate state. Under these conditions the subsurface oxide is formed again and the cycle is resumed. This mechanism was mathematically modeled by Bassett and Imbihl (306) in an interesting hybrid between an Eigenberger-type model and one

90

F. SCHUTH

et al.

with coverage-dependent activation energies. The mechanism is described with the following steps: co + * s co* (64)

+ 2* 20* co* + o*+ co2 + 2* o* Osubs 0 2

-3

(65)

The last step is a buffer step, as discussed in the Eigenberger model. In addition, however, the rate of oxygen adsorption is assumed to be dependent on the concentration of subsurface oxygen: rads

= kPo, exp(-u~,,~J

(68)

where a and k are constants. This is similar to a coverage-dependent activation energy for the adsorption step. The model calculations closely reproduced the experimental results, including the bifurcation behavior, the existence region of oscillations, and the direction of hysteresis. Equations (64)-(67) were also modeled by Yamamoto et al. (307) to get qualitatively similar results. Additionally, there is some experimental evidence for the crucial role subsurface oxygen is thought to play during the oscillations. Under conditions in which subsurface oxygen is assumed to be formed, complex LEED patterns evolve, temperature-programmed desorption (TPD) experiments yield oxygen coverages that exceed one monolayer, and two different reactivities with CO are observed (248). Another model that involves subsurface oxygen was developed by Vishnevskii for Pt( 100) (60,70,71). The mathematical formulation, though developed independently, is very similar to that of Bassett and Imbihl, and is also a hybrid between Eigenberger-type models and coverage-dependent activation energy models. This model, however, is in conflict with a surface reconstruction model developed by Ertl et al., which will be discussed later. Experimental support for the subsurface oxygen model comes from TPD and in situ XPS experiments performed by Vishnevskii and Savchenko. They observed a new oxygen state in the oscillation maximum using in situ XPS (60,61) (Fig. 11). This new oxygen peak is not correlated to the 1 X 1 or 1 X 2 structure [the two phases involved in Ertl’s phase transition mode (46)], but rather to a surface state containing subsurface oxygen (62,308). This new oxygen state could also be observed in TPD experiments after stopping the oscillations at a certain point in the cycle, transferring the sample to an adjacent TPD chamber, and then performing desorption experiments (308). Although Vishnevskii and Savchenko observed the Pt( 110) 1 X 1 S 1 X 2 phase transition, they attributed the oscillatory behavior to the effect of subsurface oxygen (309).


t

t

01s

t

PcQ'105/Pa

I

5

91

-

10

Time (min)

FIG. 1 1 . 0 1s XPS spectra taken at three indicated points along the oscillatory cycle in CO oxidation on Pt(ll0). The dotted COz partial pressure curve is from experiments; the solid curve results from simulations. The XPS peak corresponding to subsurface oxygen can be seen at 532 eV when the reaction rate is at a minimum. (From Ref. 61.)

Vayenas et al. (189)developed a model for ethylene oxidation on Pt based on solid-state electrolyte measurements that resembles the Sales-TurnerMaple model described in Section IV,A. However, Vayenas et al. balanced gas-phase concentrations and considered the surface coverages of only two species, namely active and inactive oxygen. Ethylene was assumed to react very rapidly, thus never reaching a significant surface coverage. This model semiquantitatively reproduced the experimentally observed behavior. A relatively similar model has been developed by Kurtanjek (310) as an alternative explanation to the coverage-dependent activation energy models for oscillations of the H 2 / 0 2reaction on Ni. This model is based on experimental studies (165,168)that showed oscillations in CPD measurements as the reaction rate oscillated, indicating formation and removal of an oxygen layer. Kurtanjek's mechanism consists of one differential equation for the oxygen coverage on the reduced portion of the Ni surface and one for the fraction of the surface that is oxidized, and it includes certain assumptions about the mechanisms for oxidation and reduction. H2 and O2 adsorption on the oxidized part of the surface are assumed to be in equilibrium, and thus these coverages are not balanced. Such a model, although using fitted parameters, matches experimental time series of the conversion and experimental CPD oscillations with moderate success.

92

F. SCHUTH

et al.

Rajagopalan et al. (157) investigated the same reaction on a Pd wire held at constant temperature and also explained the observed oscillations with an oxidationheduction model as described above. However, they did not perform model calculations for this mechanism. There are several other oxide models proposed in the literature. These models, however, are not isothermal, and will therefore be discussed in Section IV,B,3. b. Phase Transition Models. A second important class of mechanisms described in this section are the phase transition models, the best understood and most thoroughly investigated of which is the Pt(100) hex 1 X 1 mechanism for CO oxidation on the Pt( 100) surface, first proposed in 1982 by Ertl and co-workers (55). This is probably the most widely accepted and the best experimentally founded model available for an oscillating reaction at present. Figure 6e depicts the mechanism for the oscillations schematically. Initially the clean Pt( 100) surface assumes a nearly hexagonal (hex) 5 X 20 configuration, not unlike the (1 11) surface. Though CO may adsorb on either the hex or the 1 X 1 surface, oxygen adsorption on the hex surface is almost negligible. As CO begins to adsorb on the hex phase, a critical coverage is reached and the 1 X 1 phase begins to form. As the fraction of the surface in the 1 X 1 phase increases, more and more oxygen is adsorbed. Reaction between adsorbed CO and 0 atoms then occurs, producing C 0 2 that immediately desorbs. As the reaction proceeds, the surface coverage decreases and a second critical coverage is reached. The surface then proceeds to revert back to the reconstructed hex phase, and the cycle begins anew. The various steps in this model have all been shown experimentally, The vast difference in the sticking coefficients for 0 2 on the two forms of the surface has been known for a long time (311-314). The mechanism for the surface-phase transition has been studied in detail by Behm et al. (315,316). During the oscillations, this transition has been observed to be coupled with the respective change in the reaction rate (50,53-55,245). Scanning LEED has also been employed to obtain spatial information, for example, the presence of traveling waves of the surface structure transition (50,53),as shown in Fig. 12. This mechanism was modeled with two different techniques, first by Imbihl et al. (52) with a set of coupled differential equations, and later by Moller et al. (46,48)using a cellular automata technique. Experimental data could be fit relatively well (Fig. 13) with the obvious exception of the discrepancy in the time scales in the differential equation model. The model equations, however, are quite complex because the authors tried to model the nucleation mechanism for the phase transition. Recently the model of


93

P t I 100)

TI

(2x21

0

x

'

A

"

20

40

60

80

FIG. 12. Propagation of the Pt(100) surface-phase transition as shown by scanning LEED. The left column shows the intensity of the LEED spots corresponding to the c(2 X 2) structure of a CO-covered 1 X 1 surface; the right column shows the LEED spot intensities representing the hex phase. As CO reacts away, one can see the formation of the hex phase propagating along the surface. (From Ref. 231.)

Imbihl et al. was adapted to include modifications of oxygen adsorption on the 1 X 1 phase caused by defects or irregularities in the CO adlayer (317). Also the dynamics of the surface structure transition were described differently, using a Ginzberg-Landau equation and a lattice model to calculate energies of the different states of the surface. These two modifications resulted in a better fit of the experimental data. A somewhat simpler but mechanistically similar model was developed by b n c h et al. (318). Although presented for supported Pt, the (100) sites

94 94

SCHUTHetetal. al. F.F. SCHUTH

Time (sec)

t

2 .rxv,

. 5. e 0ol

-

10

I I I I '

' I

,

I

1

1

Time (min)

-

FIG. 13. Experimental and simulated time series for CO/02 oscillations on Pt(100). Top: set of curves showing model results with coverages u, (CO on 1 x 1 surface), v, (0on 1 x 1 surface), and b (fraction of sKrLace in the hex phase). Bottom: experimental LEED results showing the intensities of the I 1 beam (corresponding to 0 on 1 X 1 ), the c(2 X 2) (corresponding to CO on 1 X 11, and the hex LEED spots. Oscillating coverages and phase fractions are fit quite well; however, two very different time scales are used. (From Ref. 52.)

were assumed assumed to be the most active, and hence hence integral integral in controlling controlling the were was dynamic behavior of the system. In this case, the surface transition was dynamic modeled as a discontinuous discontinuous hysteresis in the O2 0 2 sticking coefficient at two experimentally determined determined critical critical values. values. This This model model was was able able to to fit fit oscillaoscillaexperimentally tory data for CO/OZ on supported Pt over a wide-parameter range with reatory data for CO/OZon supported Pt over a wide-parameter range with reasonable success. sonable success.

95


800 700

600

1

coo

500

temperature I K I

FIG. 14. Isostere for the 1 X 1 -+hex phase transition of Pt(100) (open circles) together with experimentally observed oscillation conditions for NO/CO on Pt (all other symbols). (From Ref. 91 .)

A phase change scheme similar to those described above was proposed by Schwartz and Schmidt (141,142) on the basis of LEED experiments, and by Schuth and Wicke (91,101) on the basis of IR measurements for the oscillatory CO/NO reaction on Pt(100). The experiments of Schwartz and Schmidt demonstrated that the transition from the high- to the low-reactionrate state was accompanied by a change from the 1 X 1 to the hex phase in LEED patterns. The position of the L-CO band in the IR spectra recorded during oscillations varied between the high- and low-reaction-rate states ... . .. . . . . . . . . . .* witn a relatively nign absorption Dana Delow LUXJ cm present in tne IOWreaction-rate state, which is characteristic of CO on the Pt(100) hex surface. Further proof was provided by Clausius-Clapeyron plot of the conditions for the occurrence of oscillations, which yielded points near the isostere associated with the hex 1 X 1 phase transition (Fig. 14). .

I

1

-A_,-.

-1

3

96

F. SCHUTH

et al.

The role of NO during the CO/NO reaction is analogous to that played by oxygen in the CO/O2 reaction discussed above. Different sticking coefficients have been measured on the two forms of the (100) surface for NO as well (319). A model similar to that of Imbihl et al. (52) for C 0 / 0 2 has been proposed by the same group for NO/CO on Pt(100) (144). The hex 1 X 1 phase transition was modeled in the same manner, and oscillations and multiple steady states similar to experiments were predicted. As discussed in Section IV,B, 1, this model could also display oscillations at lower temperatures, at which the phase transition was not involved. This general model of the hex 1 X 1 phase transition may also be applicable to the recently discovered oscillations in the NO/H2 reaction on Pt(100) (215). The Pt(100) surface is not the only surface that can show oscillations accompanied by a surface reconstruction. The same effect has been observed for Pt(llO), again first by Ertl et al. (64-66,68,72). The corresponding model is relatively similar to the (100) mechanism. The difference in the oxygen sticking coefficient between the reconstructed (1 X 2) surface and the (1 X 1) surface, however, is much less pronounced-they vary only by about a factor of two-and the situation is further complicated by a reversible faceting of the (110) surface that accompanies the oscillations. Oscillations on this surface were suspected to act as a “trigger” for oscillations observed at atmospheric pressure on supported catalysts (91,106).This system has been modeled both with cellular automata (46) and by a phasechange model (320) similar to that of Imbihl et al. (52). Again the phase transition was represented by an expression that changed form dependent on the value of the CO coverage. Hence, the appearance of oscillations is not surprising. A third Pt surface on which oscillations of the CO/Oz reaction have been observed is the (210) plane (75,78).In this case, a surface-phase transition was again proposed as an explanation for the oscillations. The (210) surface was observed to facet into (310) and (1 10) orientations during an induction period, after which oscillations began to occur on the (1 10) facets (78). An alternative model for the oscillatory behavior has been proposed by Ehsasi et al. (76). Another mechanism also belongs to this class of models, although in this case there is no phase transition in the first layer of the catalyst, but rather in the adsorbate. Wicke and Bijcker (125,126) proposed a model in which a CO-induced transition of the 2 X 2 oxygen structure to the more reactive fi X f i l R 3 0 ” structure is the key step in an oscillation cycle of the C 0 / 0 2 reaction on Pd. This was concluded from IR spectra, wherein the characteristic bands for the two different adsorbate structures were identified in the respective phases of the oscillation cycle. Similar results


97

were later obtained by Schiith and Wicke (12.3). This mechanism, however, was not modeled mathematically. A different interpretation of the CO bands was given by Palazov et al. (254).They attributed the different bands to CO adsorbed on the Pd( 111) and Pd( 100) single-crystal planes, and they could not observe the phase coupling of these bands as described above. They also suggest phase transitions as the origin of the oscillations, but they did not work out a detailed mechanism. c . Other Models. The remaining isothermal models cannot easily be classified. In these models, the activity of the catalyst is influenced in various ways. One such model was developed for the Pt(210) surface for C O / 0 2 oscillations as an alternative to a phase transition mechanism (76). In this model, two different adsorption states of oxygen are the key features: one with oxygen strongly bound, with a high sticking coefficient, the other with oxygen weakly bound, with a low sticking coefficient. Mathematical modeling of such a mechanism might lead to an Eigenberger-type formulation, but this task has not yet been performed. A recently developed mechanism for CO/02oscillations on Pt is the carbon model proposed by Sundaresan et al. (26,86,88,89).These authors found a 40% decrease in the number of CO adsorption sites as measured by IR spectroscopy in the ignited state, leading to extinction of the reaction. In the low-reaction-rate state, the CO adsorption capacity recovered within a few minutes and the reaction ignited again. The decrease in adsorption sites could be correlated to a rise in the carbon Auger signal when samples were taken out of the reactor at different points during an oscillation cycle. There was, however, no correlation to the oxygen Auger signal. The oscillations were thus explained by the formation of a blocking C layer and its subsequent removal in the low-reaction-rate state. Model calculations, in principle, used the same formalism as for the Sales-Turner-Maple model (272) and could predict oscillations as well (88). However, there are some problems associated with this model. First, a high CO coverage is typically found in the low-reaction-rate state, implying a low coverage of oxygen. This makes it difficult to imagine that C atoms could be easily removed in the low-reaction-rate state. Second, the source of carbon is not clear. The authors discuss possible sources such as the Pt bulk or hydrocarbon traces in the gas phase, which would find a sufficiently high number of adsorption sites only in the high-reaction-rate state. Bulk carbon should be consumed eventually. However, oscillations can be sustained over several days and it is difficult to accept that the supply from the bulk metal lasts that long. Hydrocarbons can definitely influence oscillatory behavior, as the previously noted studies from Kenney et al. show (204,205,207),but whether they actually induce oscillations remains doubtful because studies with very care-

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fully cleaned gas streams have demonstrated oscillatory behavior as well (321).Nevertheless, the C model, which is supported by some experimental evidence, is a possible explanation for oscillations, at least at atmospheric pressure. Another model in which bulk effects are suspected was discussed by Dress et al. (177,178) for methanol oxidation on supported Pd catalysts. The authors postulate that an initial dehydrogenation of the methanol to CO and Hz occurs. The H, dissolves in the Pd lattice until the concentration necessary for the formation of the Pd-H p phase is reached. On the p phase, CO oxidation is assumed to exhibit a higher rate, and CO is thus removed by reaction with oxygen. The H from the Pd bulk is then oxidized until the Pd-H Q phase is again formed. The stability of a simplified version of this model was analyzed in the same publications. Two other reactions were modeled using the buffer-step model, which originated in Eigenberger’s work, the COiNO reaction on Pt with molecularly adsorbed NO as the buffer species (322),and the CO/Orreaction on Pt with adsorbed COz as the buffer (323).NO adsorption on supported Pt catalysts is known to be a relatively slow step and the coexistence between dissociated and molecularly adsorbed NO under reaction conditions has been observed. No analysis of the NO coverage during oscillations, however, has been performed by Furman (322). In the case of the CO/OZreaction, however, data have been presented showing the marked influence of COZon the reaction rate (218).Freezing out the COz formed during reaction induced the switching of the system from the oscillatory state to the highconversion ignited state, a behavior that could be simulated in the model discussed by Lazman et al. (323). 3 . Nonisothermal Models

The models discussed thus far were essentially isothermal. Nonisothermal models as described below contain an additional degree of complexity that is always related to the characteristics of a particular reactor and the form of a particular catalyst. On the other hand, it is usually easier to construct a nonisothermal model that oscillates, because the exponential term in the Arrhenius law provides a very strong nonlinearity. Nonisothermal models can be classified into two categories: surface blocking/reactivation models and models with additional bulk effects.

a. Surface Blocking/Reactivation. The nonisothermal LangmuirHinshelwood model developed in general form by Dagonnier et al. (295) as described in Section IV,A has also been applied to the special case of CO oxidation by the same authors (324).This model is essentially composed of the same differential equations as in the general mechanism. An oscillation


99

cycle would evolve as depicted in Fig. 6f. This is in principle a blocking/ reactivation model in which oxygen is the blocking species. There are, however, two problems associated with this analysis. First, oxygen adsorption is modeled by the term Tad = k d ( 1 - 0CO - 001 (69) even though this rate should have quadratic dependence on the number of empty sites because of the dissociative adsorption of 0 2 . Second, the authors introduce the concept of a surface temperature that is assumed to deviate from the bulk temperature by a remarkable 100 K . Furthermore, it is questionable whether an Eley-Rideal reaction is a realistic pathway in this catalytic reaction (325). Several other thermokinetic blockingheactivation schemes have been developed, a general representation of which is depicted in Fig. 6g. They were formulated for such different systems as CO/NO on Pd catalysts (ZOZ,Z23) and the methanol to gasoline (MTG) process on ZSM-5 (216) so that thermokinetic blockinglreactivation mechanisms describe the widest range of oscillator types. A stability analysis for such a system was performed by Wicke et al. (98),who modeled the H2/02reaction on Pt catalysts. The reaction was simplified to two differential equations that are easily treated analytically: d0oldt =

kbioc(T)(l

-

00) - k,,,,i,,ti,”(T).eo.f(PH2)

dTldt = r ATA - K ( T - T,)

(70) (71)

The first equation describes the blocked fraction of the catalyst surface and the second describes the catalyst heat balance. The blocking species in this case is oxygen bound as an oxide species. Thus, this model can be considered as a nonisothermal version of the isothermal oxide models discussed above. The experimental support for this model is the variations in CPD measured during the oscillations by other groups (165), indicating the formation and removal of an oxygen layer. The same type of analysis was performed for two other systems, C2H4/H2 on Pt and Pd (223), and CO/NO on Pd (123). In the case of ethylene hydrogenation, partially dehydrogenated species such as (CH), groups are assumed to be the blocking species. However, additional experimental evidence for the crucial role of such a species during the oscillations is not given. For the CO/NO reaction, Nad was found to be the blocking species. At high temperatures the NO dissociation occurs faster than N2 desorption, thus yielding a surface blocked by N atoms. At low temperatures, N2 desorption is faster than NO dissociation, restoring the clean surface. Evidence for this mechanism was found by means of IR spectroscopy during the oscillations, where NCO on the support was used as a probe for the N coverage of the Pd

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surface (123). Numerical simulations for this model were performed using realistic parameter values (101,232), and the simulations described the oscillations of the reaction rate very well considering the crudeness of the model. The model that predicts the oscillations found during the MTG process on ZSM-5 described by Yang et al. (216) was not mathematically formulated. The oscillations are assumed to be caused by exothermic coke formation and subsequent endothermic coke removal in the catalyst bed. Another blockinglreactivation type of oscillatory reaction observed on zeolites is the oxidation of cyclohexane on zeolite KY (200). These oscillations are explained by the accumulation of partly oxygenated intermediates followed by decomposition of these intermediates due to the increased catalyst temperature. Evidence for the accumulation of carbonylated species could be obtained by in situ IR spectroscopy. Peaks for adsorbed cyclohexane and carbonylated species were seen to oscillate in phase with the reaction rate. This reaction system was modeled with five differential equations: two balancing cyclohexane and C02 in the gas phase, two for the coverages of cyclohexane and oxygen -hydrocarbon complexes, and one for the catalyst temperature. With this model it was possible to describe the observed oscillations by fitting the kinetic parameters. An oscillation cycle would start at a low temperature with cyclohexane adsorption and the formation of oxygen-hydrocarbon complexes. As long as the rate of formation was higher than the rate of decomposition, the complexes would accumulate until a certain temperature threshold was reached. If the temperature increased further because of heat generated by further oxidation, the complexes would burn off, leading to a sharp temperature rise. This then would lead to the desorption of all species, allowing the catalyst to cool down and the cycle to begin again. One important oscillating system-namely, the methylamine decomposition on noble metal wires (24,143,227,228)-belongs to this class of thermokinetic blockingheactivation models. This reaction is unique in several ways. It is the only endothermic oscillator (+ 150 kJ/mol), and it is the only unimolecular reaction that displays oscillations caused by surface effects. [The oscillating N2O decomposition, reported by Hugo (5) in 1968, does not oscillate because of the instability of the surface reaction, but rather due to the instability of a CSTR when certain heat and mass transfer conditions exist. Any reaction with similar rate and heat effects would oscillate under such circumstances.] This reaction is also the most vigorous oscillator yet observed and displays frequencies of up to 10 Hz and amplitudes approaching 500 K . Moreover, because the reaction oscillates at temperatures of around 1000 K , the oscillations can actually be observed visually as the metal catalyst heats and cools.


101

The reaction can be modeled by a set of three differential equations (24): d6A/dt = kaAPA(l - 6~ - 6,) - krA6A - kdA6A - kfa6A

(72)

d&/dt = kfB6A - kde&

(73)

where A is methylamine, B is the blocking species, and T is the wire temperature. The wire was assumed to be infinitely long, and heat transfer to the gas phase was approximated by a linear term, T - Tg. Heat loss by transport is much smaller than the reactive heat loss, thus this approximation is allowed even in the radiation regime. The blocker is most likely an adsorbed cyanide species that forms during the reaction and is known to be very stable on Pt [Ead = 250 kJ/mol (227)l.With the exception of the spatial features, which cannot be reproduced by a lumped model, this model semiquantitatively describes the experimentally observed oscillations as shown in Fig. 15. One should note that an endothermic blockingheactivation model is essentially a reversal of the exothermic scheme (indicated in parentheses in Fig. 6g). However, the endothermic mechanism is perhaps more easily envisioned, with a blocking species formed at low temperatures and desorbed at higher temperatures, as opposed to the reverse case, which requires the blocker formed at high temperatures and desorbed or reacted at low temperatures. b. Bulk-Phase Transition Models. The thermokinetic models discussed so far require a blockinglunblocking of the surface as the crucial step of an oscillation cycle. There is a second class of thermokinetic models in which changes of the catalyst bulk are responsible for oscillatory behavior. These bulk changes are exclusively oxidation/reduction cycles, with each phase having a different catalytic activity. Thermokinetic oxidation/reduction models have been proposed primarily for nonnoble metals such as Ni, Fe, and Cu (161,186,195,217,218). The formation of a palladium oxide has also been discussed as a source of oscillatory behavior (130,326), although the Pd in these studies was distributed in a zeolite matrix and was therefore more susceptible to oxidation and reduction. Usually oscillations of this type are described by models of the general form depicted in Fig. 6h. At a high temperature and a high reaction rate, the catalyst begins to oxidize. This causes the temperature and rate to fall, and the system eventually reaches the low-temperature oxidized state. The catalyst then goes through a reduction process that raises the temperature and the reaction rate. The process then repeats to produce oscillations. The reaction systems for which these models have been developed are CO oxidation

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experiment

simulation

0

2

4

8

10

8

10

6

Time (sec)

0) 1.0 0.8

i

b

o*6 0.4

0.2

0.01 0

.

'

2

I

'

4

.

6

.

I

1

Time (sec) FIG. 15. A comparison of experimental and simulated behavior for methylamine decomposition on Pt. (a) Temperature oscillations; (b) predicted coverage oscillations of species a, CH3NH2,and b, a site blocker; (c) simulated bifurcation diagram of temperature vs. current; (d) experimental bifurcation diagram of temperature vs. current. (From Ref. 24.)

103


(c)

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1 Model 0rill.cion Mu.

I/ I

600 2

0rilloticn Min'

.

I

I

4

6

10

8

12

Current (amps)

2

3

4

Current (amps) FIG. 15. (continued)

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on Pd-zeolite (130,326),city gas combustion on Cu (186), propene oxidation on Cu (295,196),H2/02 on Ni (162), and CO/H2 on Fe (217,218).Although the general schemes for these reactions are the same, there are some deviations or additional steps in some cases. In the case of propene oxidation on Cu (295,296),it was assumed that no metal was formed but that the oxidation state of the copper changed between +I and +II. In other words, a redox cycle between CuO and Cu20 was proposed. The oscillations were interpreted in terms of a reversal of the mechanism depicted in Fig. 6h. At a high rate and high temperature, the CuO is reduced due to consumption of O2by propene. However, on CUZOa much less exothermic partial oxidation to acrolein is favored, which leads to decreased temperatures and much less oxygen consumption. The cooling and the lower oxygen consumption lead to a high O2 content in the gas phase, and the Cu20is oxidized to CuO. On CuO, the highly exothermic complete oxidation of propene is the main reaction, hence the catalyst heats up and the cycle starts again. The change between the different states of the catalyst could be observed visually; in fact, the color of the catalyst changed between red-brown (CUZO)and black (CuO) during the oscillations. City gas combustion on Cu is somewhat more complicated than the general mechanism (186).In the high-temperature state, the copper oxide was assumed to adsorb COZstrongly so that the surface was blocked. Thus this model is a hybrid between thermokinetic blockingheactivation models and thermokinetic bulk transition models. This mechanism was modeled mathematically, although the parameters for the simulations were not obtained by independent methods but rather by a curve fit procedure. Another system in this class of mechanisms that was also modeled mathematically is the technically important Fischer -Tropsch process on Fe catalysts investigated by Caldwell (217,228). Oscillation had been previously observed in this system by Tsotsis and co-workers (219).Caldwell reduced the mechanism of the Fischer-Tropsch reaction to three reactions: CO

+ 2Hz -+

-CHZ-

Fe304 + 4CO

+ H20

3Fe

+ 4COz

Fes04 + 4Hz 2 3Fe

+ 4Hz0

%

The catalytically active phase was assumed to be exclusively a-Fe, and Fe304was assumed not to be active for the Fischer-Tropsch reaction. Kinetic parameters for the simulations were obtained independently in separate oxidationh-eduction studies. Balancing the oxidation and reduction rates for the CO/CO, and the H2/H20 systems independently and describing the rate of synthesis in Fischer -Tropsch reactions by a standard expression, Caldwell could predict the oscillations with a simplified model for a tubular reactor fairly well.


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For the remaining bulk-phase transition systems mentioned above, no model calculations or stability analyses have been performed, and the mechanisms are similar to the general scheme. Therefore these models will not be discussed in detail.

V. Synchronization, Spatial Patterns, and Chaos Nearly all models discussed so far have one feature in common: they are not distributed models and can describe only spatially uniform systems. Many of the mathematical models use ordinary differential equations, and the resultant time series are nearly always simply periodic. This approach, however, describes only part of the experimentally observed behavior; there is a great deal of experimental evidence for spatial heterogeneity and chaotic oscillatory behavior in heterogeneously catalyzed systems. At the root of these behaviors is the phenomenon of synchronization of a macroscopic oscillating system. This topic has been discussed occasionally in the literature but has rarely been explicitly treated. In principle there are several stages or hierarchical levels on which oscillations can occur: (1) the single-crystal plane, (2) the catalytically active metal crystallite, (3) the catalyst pellet, (4) arrangements of several pellets in one layer of a flow reactor or a CSTR, and finally ( 5 ) the catalytic packed-bed reactor (327). On each of these levels, different types of oscillations may exist, but to become observable on the next level oscillations on the respective sublevels must be synchronized. For example, if oscillations of the CO/OZ reaction on a Pt( 100) face of a Pt crystallite supported on a pelletized support material in a packed-bed reactor occur, the reaction on the (100) facet as a whole must oscillate in synchrony, other (100) facets of the crystallite have to synchronize, other crystallites in the pellet must couple to the first crystallite, and, finally, all pellets in one layer of the bed must display oscillations in synchrony. If the synchronization on one of these levels fails, different oscillators will superimpose and their effects will cancel. One would then only observe a possible increased level of noise in the measured conversion. On the other hand, if synchronization occurs independently over several regions of a system, then it might exhibit apparently chaotic behavior caused by incomplete coupling. Evidence for the importance of synchronization has been found on all of the five levels discussed above. For single-crystal planes it was demonstrated by Ertl et al. that the surface acts in a cooperative way, thus forming traveling waves (50,53)or regions of different activity (57,58,73,74). Schwartz and Schmidt (141 ) reported the synchronization of oscillations on both sides of a Pt(100) single crystal, and Sander et al. (246) observed the

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synchronization of oscillations on the (loo), (110), and (210) planes of a cylindrical Pt crystal. On the level of a catalyst pellet or a polycrystalline metal catalyst (24,91,99,112,113,123,129,148,328), and for several pellets (100,103,115,184,237,329), evidence for synchrony has been presented by several groups. As was discussed by Schwartz and Schmidt (142), there are in principle three different mechanisms that can act as the synchronizing force. These possible mechanisms are thermal synchronization (either through the catalyst, through its support, or through the gas phase), synchronization through the reactant gas phase (caused by an oscillating rate of consumption on the surface), or a synchronization mechanism on the catalyst surface (either surface diffusion of adsorbed species or a phase transition of the surface structure). All of these mechanisms probably play some role in the synchronization of oscillations. However, which process is dominant depends on the individual system. For instance, in single-crystal studies, where heat effects are minimal, pressure changes or surface processes must be responsible for coordinated oscillations. In a packed-bed reactor, on the other hand, the thermal conductivity of the bed is the major synchronizing force. One additional synchronizer- namely, electronic interactions over the catalyst metal-could also be possible, although its existence has yet to be proved. The possible synchronization mechanisms, together with the systems in which they can be found, are shown schematically in Table IV. The existence of chaotic behavior (22,40,168) and spatial inhomogeneities (53,330,331) in heterogeneous catalysis was first acknowledged in the late 1970s and early 1980s. Simultaneously, theoretical efforts were made to describe the chaotic behavior, because previous models were not capable of generating chaotic responses. The first successful attempt to model this behavior was by Jensen and Ray (242,243) with the pebbly surface model mentioned in Section IV,A. In this section we will first describe some experimental results in this field and then different approaches to modeling TABLE IV Possible Synchronization Mechanisms and Systems in which They Are Likely to Be Important"

Mechanism Single crystallHV Polycrystalline/HV PolycrystallinelTorr range Polycrystalline/atmospheric Supported/atmospheric "

Surface phase transition

+

-

Surface diffusion

Electronic interaction

+ + + +

+ + + +

-

Significant mechanisms are indicated by a "+ ."

-

Pressure

+ +

+ + +

Temperature -

+ / -

+ +

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complex behavior that are classified according to the mechanism of communication in each system.

A. CHAOTIC OSCILLATIONS Chaotic time series have been obtained from a wide variety of experimental systems. Figures 16 and 17 show examples of the irregular time series found in various cases. It is somewhat problematic to assign the term “chaotic” to these reported time series because it was rarely investigated whether the time series were deterministically chaotic in the strict sense of the word. Therefore, when we use the expression chaotic, we are well aware that there is, in many cases, no proof for chaos in the oscillation patterns. The few reports wherein a thorough analysis has been performed (64,104) do though show the existence of deterministically chaotic oscillations. It would be beyond the scope of this review to describe the methods

-

10min

t

FIG. 16. A chaotic front moving through a fixed-bed reactor during CO oxidation over Ptl AI2O3. T,,Tb,T,, and Td are pellet temperatures in the entrance cross section; TI and TZare thermocouple temperatures 24 and 46 mm downstream, respectively. Chaos was characterized by correlation dimension and Liapunov exponent. (From Ref. 100.)

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et al. d

1 I

C

-

$

100

80

-3 40 60

20

b

20

T

..

.

a

50

100

150

Time (sec)

FIG. 17. The transition to chaos (from a to d) observed in the work function during CO oxidation on Pt(ll0) while decreasing CO pressure. Chaos in the upper time series (d) was characterized by the Liapunov exponent, Kolmogorov entropy, and the embedding dimension (From Ref. 68.)

used to analyze chaotic behavior of a system. The reader is therefore referred to other sources specialized in this field (see, e.g., Ref. 332). The first reaction for which complex patterns were explicitly addressed was once again the CO/Oz reaction on Pt ( 2 2 ) . Plichta and Schmitz showed complex time series under nearly isothermal operating conditions. Later, results from Schmitz et al. (333) were interpreted as being deterministically chaotic, and at least four state variables were assumed to be necessary for a description of the system. Complex time series have also been reported for the same reaction on Pt/A120, in a fixed-bed reactor, where the properties of the catalyst bed determined simple or complex behavior (237,329), and on the Pt(ll0) single-crystal plane (68),where a period doubling sequence described by the Feigenbaum scenario led to chaos. In this case, an extensive analysis was performed to prove that the oscillations were truly chaotic. The analysis included examination of power spectra, autocorrelation functions, attractors, correlation integrals, Liapunov exponents, and a lower limit for the Kolmogorov entropy ( 6 4 ) . A period doubling sequence was also observed by Kapicka and Marek (334). They, however, did not vary an external control parameter but rather recorded the evolution of the oscillation pattern with time. They assumed that the catalyst activity, which was changing with time, was the “control parameter”.

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o

b

c

d

e

f

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160O'C 30 VOL% N Y FIG. 18. Evolution of chaotic time series with decreasing oxygen concentration (from a to i) during ammonia oxidation over a Pt wire. (From Ref. 211 .)

A thorough analysis of chaotic oscillations for the NH3/02 reaction over Pt has been performed by Sheintuch and Schmidt (211). Bifurcation diagrams were presented in great detail, as well as phase plane maps and Fourier spectra. Figure 18 shows a series of oscillation traces obtained for various oxygen concentrations. By extracting a next return map from trace g in Fig. 18, evidence for intermittency could be obtained. Complex time series have also been detected for Pd (129,131). In these studies, however, it was shown that the complex structure could be characterized by three or fewer superimposed frequencies (129) and was thus not deterministically chaotic. Chaotic oscillations have been reported for several other systems, namely, NH3/O2 on Pt (40,211,212), H2/02 on Ni (168,169) and Pt (152), CO/NO on Pt and Pd (91,123), C2H4/H2 on Pt ( 3 3 3 , C3Hd02on Pt (194), l-hexene/Oz on Pd (198), and the CH3NH2decomposition over Pt, Rh, and Ir (24). This list, while incomplete, shows that irregular oscillations and often chaos are frequently encountered during heterogeneous catalytic studies. B.

SPATIAL PATTERNS

In the discussion of complex time series, no mention has been made of spatial inhomogeneity, an intrinsically related phenomenon. Several articles, however, have dealt with spatially inhomogeneous oscillating reactions, either in the form of traveling waves or as independently oscillating patches on a catalyst.The latter case can lead to complex time series if the regions are partially coupled. As yet, there has been no definite proof that all chaotic time series are caused by incomplete coupling of differently active regions on a catalyst, but it seems quite probable that this is the predominant chaos-generating mechanism. The pioneering work that examined development of spatial patterns during oscillating reactions was performed by Schmitz et al. (262,263,330). Their IR- thermographic imaging studies clearly demonstrated for the first time that an oscillating heterogeneously catalyzed reaction is not necessarily synchronized over a whole catalyst disk, but that a disk might show several centers of activity. Temperatures on different areas of the catalyst varied by

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up to 90 K during the H2/02reaction on supported Pt catalysts (263). Similar results were later observed by other groups, mainly with supported catalysts, but also using wires and foils. Extensive investigations were performed by Wolf and co-workers (112,114,148,252), who studied the CO/Oz, CO/NO, and C2H4/02reactions on supported catalyst wafers. They used arrangements of several thermocouples and partly shielded disks in IR experiments to prove that different regions on the catalyst surface were oscillating independently and that activity waves were traveling on the disk. This was confirmed by Kellow and Wolf in a recent publication (136) using the thermographic imaging technique used by Schmitz et al. In these measurements, as with the experiments with isolated thermocouples, the catalyst preparation procedures and thus the distribution of the active metal had a strong influence on the synchronization behavior. In the work of Jaeger et al. (131,184,336,337), chaotic conversions and their connection with spatial inhomogeneities were frequently addressed. They describe experiments that show both that coupling and decoupling between different synchronized areas of a catalyst bed lead to chaotic conversions (184) and that within a pellet synchronization is best achieved when the metal loading of the catalyst is high, because this facilitates communication (131). Schuth and Wicke (91,123) reported discrepancies between thermocouple readouts (a local probe) and overall conversions (a global probe) during the oscillating CO/NO reaction on supported Pd and Pt wafers, which also indicates spatial inhomogeneities. On purely metallic catalysts, fewer reports describing spatial patterns have been published. Tsai et al. (328) investigated oscillations of CO oxidation on a Pt wire loop and found by using two thermocouples on opposite sides of the loop that oscillations were not completely synchronized over the entire wire. Lobban and Luss (164) showed by thermographic imaging the existence of traveling waves on a Ni foil during the H2/02 reaction, with fronts traveling at a speed of 1 cm/s. Sheintuch (338,339) performed a detailed analysis of the complex behavior of the controlled wires that were used in the experimental study of Lobban and Luss. He showed that the temperature control of catalytic wires could transform one oscillatory state of the uncontrolled system into several coexisting oscillatory and stable states. The wire underwent a symmetry breaking under these conditions and developed a spatial structure. This was observed for stationary states by Lobban et al. (264), and later for oscillatory states by Lobban and Luss (164) and Philippou et al. (194), as mentioned above. Such behavior was observed experimentally not only for temperature-controlled wires but also for operation in constant voltage and constant current modes. This was also the case during CH3NH2decomposition on a 5-cm-long Pt wire, where traveling waves as well as independently oscillating segments could be observed


111

for certain reactor pressures (340). Proof of the existence of spatially inhomogeneous reaction rates on a single-crystal surface was given by Ertl et al. for Pt(100) (57,58) and Pt(ll0) (58,73,74). They used a novel technique, scanning photoemission spectroscopy, to spatially resolve the work function during oscillations of the CO/02 reaction under HV conditions. The same group also presented the first evidence for traveling oscillation waves on a Pt(100) single crystal (50,53). C . UNDERSTANDING AND MODELING COMPLEX BEHAVIOR Chaotic behavior and synchronization in heterogeneous catalysis are closely related. Partial synchronization can lead to a complex time series, generated by superposition of several periodic oscillators, and can in some cases result in deterministically chaotic behavior. In addition to the fact that macroscopically observable oscillations exist (which demonstrates that synchronization occurs in these systems), a number of experiments show the influence of a synchronizing force on all the hierarchical levels mentioned earlier. Sheintuch (294) analyzed on a general level the problem of communication between two cells. He concluded that if the gas-phase concentration is the autocatalytic variable, then synchronization is attained in all cases. However, if the gas-phase concentration were the nonautocatalytic variable, then this would lead to symmetry breaking and the formation of spatial structures. When surface variables are the model variables, the existence of synchrony is dependent upon the size scale. Only two-variable models were analyzed, and no such strict analysis has been provided for models with two or more surface concentrations, mass balances, or heat balances. There are, however, several studies that focused on a certain system and a certain synchronization mechanism. 1.

Coupling via a Surface Process

For the level of the single-crystal plane, the elegant scanning LEED experiments of Ertl et al. (50,53)showed that a reaction front propagates over the surface, usually starting from a defect such as the edge or corner of the crystal. This system has been successfully modeled by two different approaches. Imbihl et al. (52) developed an extension of the phase-change model mentioned in Section IV,B, which includes a spatial variable in the form of a surface diffusion term. The crystal surface was then modeled by a one-dimensional array of cells coupled by different diffusion algorithms. Integration of this system yielded transformation waves that travel along the surface similar to those found in experiments. The other form of the model was the cellular automata approach by Moller et al. (46,48). This technique also allowed the prediction of traveling wavefronts on the surface of a

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Pt(100) crystal plane. The coupling in this case was caused by a step in the model where an active site is transformed to an inactive site when certain requirements for the neighboring cells were met. A relatively similar model was developed by Gerhardt and Schuster (341 ) to describe the oscillations of CO oxidation on Pt as observed by Jaeger et al. (342). In this case, a change in the activity of a site was initiated when a certain number of neighboring sites were covered by oxygen. This model not only yielded spatial patterns and traveling waves but also predicted chaotic time series as were observed in experiments. The model was not explicitly developed for a single-crystal surface but implicitly assumed a homogeneous surface, which would be best realized by a single-crystal plane. Nonhomogeneous behavior was also predicted by Levine and Zou (343), who began with the oxidation/reduction model of Sales et al. (272) for CO/Oz on polycrystalline Pt and introduced CO surface diffusion. This model predicted one-dimensional, stable traveling waves on the Pt surface. Because no experimental measurements have been performed for traveling waves on polycrystalline catalysts at atmospheric pressures, no comparison to experimental data was possible. if stochastic methods are applied, however, a term representing coupling by a phase transition or surface diffusion does not appear to be necessary to induce chaos in a heterogeneously catalyzed reaction. Fichthorn et al. (277) investigated a simple Langmuir-Hinshelwood mechanism with a Monte Carlo method and found rich, dynamic behavior. Analysis of the time series by the correlation integral method proposed by Grassberger and Procaccia (344) demonstrated that the system was indeed oscillating chaotically. However, although these results are interesting, it is still questionable whether this approach really describes the process leading to chaos in a surface reaction. As mentioned earlier, the Bendixon theorem proves the impossibility of even a limit cycle in the Langmuir-Hinshelwood mechanism if it is described by ordinary differential equations. Hence the results of Fichthorn et al. are most likely caused by an effect of the finite grid size.

2. Coupling via the Gas Phase Coupling across the surface caused by pressure variations in the gas phase is a common and ubiquitous mechanism of synchronization over catalyst regions in a reactor. As with previous mechanisms this coupling can occur at various levels across the catalyst. a. Single-Crystal Level. Wave propagation and coupling through surface processes are not the only mechanisms by which different parts of a single-crystal surface can be synchronized. For the (1 10) plane of Pt, imbihl et al. (59)found by using two Kelvin probes on a single-crystal surface that


113

the coupling in this case proceeded via concentration changes in the gas phase. A comparison of the oscillatory outputs from the two work function measurements and the overall reaction rate as measured by a mass spectrometer ruled out the occurrence of any type of coupling through surface mechanisms. b. The Crystallife Level. The experiments and models discussed above were implicitly or explicitly restricted to coupling over one crystal plane. The next level on which synchronization can develop involves different crystal planes on a single-metal crystallite. The first indication for the existence of such an effect was found by Schwartz and Schmidt (141), who showed that both sides of a Pt( 100) single crystal can oscillate in phase. Because communication by a surface phase transition between both sides was not possible, and because temperature cannot be the synchronizing force because of the low heat of reaction and the large thermal capacity of the single crystal, it was suggested that pressure oscillations in the gas phase of the HV system must be responsible for the synchronization of the two sides of the crystal. The same conclusions were drawn by Ehsasi et al. (247) after investigating the synchronization of two Pd( 110) single crystals that were in the same UHV system. Synchronization between the two crystals proceeded rapidly through the gas phase. Decreasing the degree of coupling by pumping out the gas phase at a higher rate resulted in the appearance of phase-locking, entrainment, and quasiperiodicity in the overall reaction rate. The importance of concentration changes in the gas phase was also observed by the same research group in studies of the Pt(210) plane during CO oxidation (76). Recently, Sander et al. (246) studied a situation very closely approximating a metal crystallite on a support. They used a cylindrical crystal with its axis oriented along the (001) direction, thus exposing (loo), (1 lo), and (210) surfaces. At low pressures, the (100) and (110) faces oscillated relatively independently of each other, with no synchronization between the (100) faces and only weak synchronization between the (1 10) faces due to the small pressure changes in the gas phase. At higher pressures, when the (210) face is in its oscillatory regime, the pressure changes in the gas phase are large enough (25%) to synchronize all crystal planes of the cylindrical crystal. Modeling of these effects has not yet been attempted. Synchronization through the gas phase was also assumed to occur by Tsai et al. (82) for CO oxidation over polycrystalline Pt. At 760 Torr the observed oscillations exhibited complex waveforms. When the pressure was lowered to 10 Torr, communication increased between oscillating patches of the catalyst due to higher gas-phase diffusivity and the oscillatory waveforms become simpler.

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c. The Pellet Level. There has been considerable experimental work on the coupling of crystallites or active regions in a catalyst pellet or wafer. However, only a few attempts have been made to model this behavior. The first models applicable to a catalyst pellet, the “fuzzy wire” and the “pebbly surface” models, were proposed by Jensen and Ray (242,297).They combined the behavior of the metal crystallites on the support (protrusions in the case of the wire) with mass transfer in the particle, while neglecting temperature inhomogeneities over the pellet. The resultant system of ordinary and partial differential equations was solved by a collocation method and yielded a highly complex time series because of mass transfer limitations in the catalyst pellet (Fig. 19). The physical picture of their model is a catalyst pellet on which are arranged active metal particles with a certain size distribution. On the metal, a simplified oxidation reaction took place: A +

*

%A*

(78)

*

(79)

A*-+P+

Four differential equations were used to describe the system: one coverage equation for the crystallites, one heat balance for the crystallites, a mass balance for the pellet, and a heat balance for the pellet, which was lumped over the pellet diameter. A

230 3

2

0

266.0 7

6

4

8

0

D

4

6

8

4 Time (DIML)

6

8

2

E

275.0 274.5 274.0

-

a 0

0

C

238 7

236j

2

234

2321 0

.

, . ,

2

I

4

, 6

1

8

Time (DIML)

FIG. 19. Oscillatory temperature versus time predicted by the “pebbly surface” model under various conditions. (From Ref. 242.)

115


3. Thermal Coupling a . The Pellet Level. The third possible coupling mechanism is synchronization through heat transfer. Schiith et al. (232) lumped the gas phase and introduced coupling via thermal conduction through the support for modeling chaotic traces and spatial inhomogeneities for the COIN0 reaction on catalyst wafers (91,123).The surface of the catalyst wafer was modeled as a 10 X 10 grid with 10 random spots representing active areas of the wafer. Each active spot was described by two differential equations, which were capable of yielding only simple oscillations. Coupling between cells was then described by an additional heat transfer term involving the four nearest neighbors. This system of differential equations was numerically integrated and predicted a wide range of complex behavior, which included deterministic chaos (characterized by continuous Fourier spectra and positive Liapunov exponents), period doublings, quasiperiodicity , periodic oscillations, and stable steady states (Fig. 20). Which behavior occurred depended on the magnitude of the heat transfer coefficient between the catalyst and the gas. It was also shown that one strong oscillator, such as a very active center, could entrain and synchronize the whole disk. The simulation

0.10

-

0.oo.J 0.12

2

.

4 8

.

, 0.14

. . , . 0.16

.

, 0.18

. . , 0.20

. .= , 0.22

FIG.20. Largest Liapunov exponent versus heat transfer coefficient for gas-phase coupling for the NOiCO reaction modeled on a catalyst wafer. The wafer is composed of 100 cells with 10 randomly distributed active cells as shown on the grid. The numbers pointing at various regions indicate the onset of particular periodicities; chaos is observed for 0.15 < K , < 18. (From Ref. 232.) K,

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results qualitatively matched experimental results using realistic parameter values, and the findings can probably be generalized to other reaction systems and catalyst geometries. Experiments that support this idea were reported by Onken and Wolf (99), who investigated disks consisting of one, two, and many active patches.

b. The Packed-Bed Reactor. Several publications deal with the onset of synchrony and chaos in a catalyst bed. The first experiments were carried out by Rathousky and Hlavacek (237), who attributed the chaotic behavior of fixed-bed reactors to the interaction of individual catalyst pellets. This situation was studied in detail experimentally by Onken and Wicke (103,104), first for a small arrangement of catalyst pellets and then for a real adiabatic fixed-bed reactor. In the former case, the thermal contact between three catalyst pellets was varied using the arrangements (a-c) shown in Fig. 21. Oscillations on each pellet were monitored via measurement of the pellet temperature. If the three pellets are independent of each other, each pellet oscillates with its own natural frequency (a). When a weak coupling is introduced, the pellets are synchronized to some extent, but each pellet maintains its own characteristic frequency (b). Increasing the coupling further results in nearly complete synchrony (c). In an arrangement of three pellets, simulating the entrance cross-section of a fixed-bed reactor, a transition from relatively periodic oscillations to chaotic behavior was detected as the feed temperature was varied. At low temperatures it is possible that one pellet oscillates, yielding the periodic time series. When the feed temperature is increased, more and more pellets reach oscillatory states, thus yielding the complex patterns. If these patterns form in the entrance cross-section of a fixed-bed reactor, it is conceivable that they might be amplified and propagate through the bed. Such traveling waves have indeed been detected in fixed-bed reactors (92,105,106,329). The behavior described above could be qualitatively modeled by simulating the bed as a cascade of CSTRs (327). Experiments similar to those of Onken and Wicke were carried out by Kapicka and Marek (115). They investigated the CO/02 reaction in an adiabatic fixed-bed reactor and also observed chaotic time series caused by the coupling of different pellets. The importance of thermal coupling in catalyst beds was also shown by Plath et al. (336). The authors arranged catalyst pellets on support plates made of different materials, such as teflon and stainless steel, and found that the resultant overall oscillatory curves were strongly dependent on the thermal conductivities of the supporting plates. As expected, higher thermal conductivity improved the synchronization, resulting in regular time series.

H

t

10min

o c b

H

10 min

-

t

10min

-t

FIG,21. The influence of coupling strength on oscillations of three Pt/A1203 pellets during the CO/Oz reaction. Top, no coupling; middle, weak coupling; bottom, strong coupling. (From Ref. 103.)

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Summary

Catalytic oscillations occur so commonly that the topic cannot be ignored, yet they are so complex that experimental characterization is seldom satisfactory and theoretical characterization is seldom complete. Therefore we have attempted to discuss experiments and their interpretations together so that readers can attempt to sort through alternate interpretations. However, in spite of many experiments and models, it appears that no oscillatory catalytic systems are as yet understood completely. As more research is done in this field, it becomes clearer that there is probably no single mechanism explaining all oscillatory behavior in heterogeneous catalysis, but rather that every single system has to be studied in detail in order to elucidate a mechanism on the microscopic scale. Oscillations are possible on all levels in a catalytic reactor, from the single-crystal plane to the crystallite to the catalyst pellet to the packed-bed reactor, and each level adds another degree of complexity. Thus it is necessary to isolate the major influences at each level and to separate the characteristics of the oscillations on one level from the effects caused by coupling with other levels. Only when each level is well understood is it possible to fully understand the overall oscillatory behavior. Oscillations in heterogeneous catalysis will therefore remain an intriguing and demanding problem for many years to come. REFERENCES Prigogine, I., Angew. Chem. 90,704 (1978). Zhabotinskii, A . M., Dokl. Akud. Nauk SSSR 157, 392 (1964). Beusch, H., Fieguth, P., and Wicke, E., Chem.-Ing.-Tech. 44, 445 (1972). Wicke, E., Beusch, H., and Fieguth, P., Adv. ACS Symp. Ser. 109, 615 (1972). Hugo, P., Eur. Symp. Chem. React. Eng., 42h, Brussels, 1968, p. 459 (1971). Hugo, P . , Ber. Bunsenges. Phys. Chem. 74, 121 (1970). Hugo, P., and Jakubith, M., Chem.-1ng.-Tech. 44, 383 (1972). Matros, Y. S . , ed., “Studies in Surface Science and Catalysis. Vol. 43: Catalytic Processes Under Unsteady State Conditions.” Elsevier, Amsterdam, 1989. 9. Matros, Y. S . , ed., “Unsteady State Processes in Catalysis.” Utrecht, 1990. 10. Matros, Y. S., Chem. Eng. Sci. 45, 2097 (1990). 11. Hamer, J. W., and Cormack, D. E., Chem. Eng. Sci. 33, 935 (1978). 12. Chiao, L., and Rinker, R. G., Chem. Eng. Commun. 48, 191 (1986). 13. Thullie, J . , and Renken, A . , Chem. Eng. Sci. 46, 1083 (1991). 14. Unni, M. P., Hudgins, R. R., and Silverston, P. L., Can. J . Chem. Eng. 51, 623 (1973). 15. Abdul-Kareem, H. K., Silverston, P. L., and Hudgins, R. R., Chem. Eng. Sci. 35, 2077 (1980). 16. Cutlip, M. B., Hawkins, C. J., Mukesh, D., Morton, W., and Kenney, C. N., Chem. Eng. Commun. 22, 329 (1983). 17. Svensson, P., Jaeger, N. I., and Plath, P. J., J. Phys. Chem. 92, 1882 (1988). 18. Hegedus, L. L . , Chang, C. C., McEwen, D. J., and Sloan, E. M., Ind. Eng. Chem. Fundam. 19, 367 (1980).

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Zeolite-Supported Transition Metal Catalysts* WOLFGANG M. H. SACHTLER AND ZONGCHAO ZHANG lpatieff Laboratory Department of Chemistry Center for Catalysis and Surface Science Northwestern University Evanston, Illinois 60208

1.

introduction

Zeolites have been widely applied in industrial processes as molecular sieves, ion-exchange agents, and heterogeneous catalysts (I-.?). Noncatalytic uses of zeolites of various porekhannel dimensions include selective adsorption in large-scale separation processes. This is outside the scope of the present paper, and the reader is referred to recent reviews [see, e.g., Ruthven ( 4 ) ] .In catalytic applications, zeolites are predominantly used in their acidic form. The most important process in this category is fluidized catalytic cracking, based on rare-earth-exchanged zeolites, mainly X and Y of the faujasite structure with small admixtures of ZSM-5. Another industrial process in this group is catalytic dewaxing using mordenite (5) and ZSM-5 (6). Applications of ZSM-5 and faujasites in synfuel production, isomerization, and other industrially important reactions have been thoroughly reviewed (7-9); applications of zeolites in organic synthesis have been explored more recently (ZO-ZZ). In particular, Ti-substituted zeolites display unique activity in oxidation catalysis (13). Within the vast group of zeolite catalysts, the focus herein is on one subgroup: materials that contain reduced particles of a transition metal or several transition metals dispersed inside zeolite cavities. Present applications and the future potential of these materials in the petroleum and petrochemical industries have been reviewed by Minachev and Isakov (14), Kouwenhoven (15), Uytterhoeven (Z6),Maxwell (I7,18), Delafosse (19), and

* Dedicated to Professor Herman Pines on his 90th birthday. 129 Copyright D 1993 by Academic Press, Inc. All right$ of reproduction in any form reserved.

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Gallezot and Bergeret (20). A majority of transition metal/zeolite catalysts are bifunctional, i.e., strong acid sites are present in the same zeolite. A process based on a zeolite-encaged metal in the absence of acid sites is the dehydrocyclization of small linear alkanes (such as n-hexane) to aromatics. Here the unique selectivity of Pt/KL* and Pt/BaKL is used (21,22).Much research has been devoted to this intriguing system (23-31). More recently, zeolite-supported Cu catalysts have been found to be highly efficient in NO, decomposition (32-35). This finding is of great interest in view of the environmental demand to destroy NO, in exhaust emission from vehicles and in the stack gas of power plants. At the time of writing this text, it is, however, not certain that reduced metal is the important catalytic function of this catalyst. The Cyclar catalyst, which contains Ga in HZSM-5 and is used for the dehydrocyclodimerization of propane or butane to benzene and its homologues, apparently does not belong to the category of catalysts under review herein, because Ga functions in an ionic state (Ga’ and Ga3+)(36). Catalysts used in important petrochemical processes are sometimes misleadingly called zeolites, though they contain nonnegligible amounts of platinum or other transition metals, which render such catalysts bifunctional. We mention hydroisomerization with Pt supported on mordenite ( 3 7 ) , and some variants of hydrocracking, catalyzed by rare-earth-exchanged acidic zeolites of type X or Y mordenite, containing metals such as Pt or Pd (38-40). The traditional model for bifunctional catalysis, introduced by Mills et al. (41), assumes that metal sites mainly provide the hydrogenation and dehydrogenation function, whereas the acid sites catalyze isomerization and cracking reactions. The traditional model further accepts a concomitant role of metals in bifunctional catalysts to hydrogenate coke precursors, thus preventing the deposition of carbonaceous overlayers. In this model it is often tacitly assumed that a small amount of a platinum metal is sufficient to establish the paraffinlolefin ratio in the gas phase. It was, however, shown by Ribeiro et al. (42) that the isomerization activity of Pt/HY and Pt/ H-mordenite catalysts increase almost linearly with the number of surfaceexposed Pt atoms up to Pt contents far in excess of the metal loads of most industrially used Ptlzeolite catalysts. The conventional classification of the reaction network in metal-catalyzed (de)hydrogenation and acid-catalyzed isomerization and cracking steps (43) is certainly an oversimplification of reality. Work of Liberman et al. (44), Gault ( 4 3 , and others (46) revealed

* Metal-loaded

zeolites are expressed in the form MI/M2Z, where M I = reduced metal, = zeolite type. For example, Pt/KL stands for platinum in the channels of an L-zeolite, with K as the charge-compensating cation. Details of zeolite types with their conventional abbreviations can be found in “Zeolite Molecular Sieves” (D. W. Breck; Robert E. Krieger Publishing Company, Malabar, Florida, 1974).

Mz= charge-compensating metal ion, and Z

ZEOLITE-SUPPOWED CATALYSTS

131

that isomerization, ring closure, and ring opening are catalyzed by acid-free metal sites. More recent results also cast doubts on the assumption that metal and acid sites have to be geometrically separated one from another, so that reaction intermediates must shuttle between them. New data suggest that metal particles in zeolites can combine with protons to form one entity (47,48); on such sites (de)hydrogenation and skeletal rearrangement can supposedly take place without intermediate desorption and readsorption (49). As the metal-proton adduct carries a net positive charge, the metal has become “electron deficient” (50). Because these novel concepts were derived from observations on metal/zeolite systems, they provide additional motivation for reviewing the present state of knowledge of that field. Here we primarily focus on transition metals in zeolites, faujasites in particular, though metals on amorphous supports will be mentioned for comparison. Although industrial application is a major incentive for studying transition metal/zeolite systems, a second motivation is their importance for understanding fundamental principles of heterogeneous catalysis. Two key issues should be mentioned: zeolite-imposed stereoselectivity and metal atom reorganization. Stereoselectivity can be imposed by barring entrance into the zeolite of some isomers of the feed while admitting others, or by preventing escape of some products formed in zeolite cages, or by discriminating between transition states of different spatial requirements. In addition, molecules diffusing through narrow zeolite channels tend to become oriented with respect to the channel axis; they will then hit metal clusters in a nonrandom way and one particular configuration of the adsorbed molecule will prevail. Evidence for this has been reported by Moretti and Sachtler ( 5 4 , Tauster and Steger (24), Alvarez and Resasco (52), and Ostgard et al. (53). Secondary rearrangements of an adsorption complex, such as “roll-over,” which are facile in the absence of steric restrictions, will be hindered or even impossible, e.g., in the narrow space between a particle surface and the ceiling of the zeolite cage (54). All of these phenomena impose stereoselectivity to catalytic conversions, reflecting the peculiar geometry of metal particles in defined positions inside a zeolite channel or cage system. Chemisorption-induced metal atom reorganization at low temperatures had been discovered in the 1960s by field ion microscopy in chemisorption studies of nitrogen and carbon monoxide on a tungsten single-crystal tip (55-58). Likewise, surface reconstruction was concluded from research of chemisorption on metal and bimetal films (59-61). This surface reconstruction was observed particularly on “open” crystal faces; the more closely packed faces are often “corroded’ from their edges with the former faces. Surface reconstruction is of particular relevance to alloy surfaces, where it leads to chemisorption-induced segregation of one alloy component to the

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surface (62-66). Surface reconstruction as a direct consequence of chemisorption is of fundamental importance for the understanding of thermodynamics and kinetics of chemisorption and heterogeneous catalysis. The heat of adsorption in general includes endothermic terms due to the partial rupture of metal-metal bonds (67,623). The importance of this phenomenon for catalysis may have been underestimated in the 1970s and 1980s, when much surface science work focused on macroscopic singlecrystal faces of high atomic density, where surface reconstruction is indeed minimal. For the small metal particles present inside zeolite cages, but also on amorphous supports in modern industrial catalysts, the mobility of exposed metal atoms is known to be much higher. With chemisorption on such clusterlike particles, the surface easily reconstructs to acquire the configuration representing minimum free energy of the metal/adsorbate system. Little is known, however, about the dynamics of metal reconstruction of small particles during chemisorption and catalysis. Zeolite-supported metal particles offer favorable conditions for studying such processes because (1) metal particles on zeolites are more uniform than particles on amorphous supports and (2) the zeolite acts as an isolation matrix; bare and ligated clusters, which tend to coalesce to larger particles, can be trapped in zeolite cages at low temperatures and studied while isolated one from another. Recent research on the elementary steps during preparation of zeolitesupported metals has helped to understand the genesis of these particles in considerable detail. For the best studied systems the goal of preparing catalysts by design has been achieved (69-77). An arsenal of modern physical and chemical methods permits characterization of metal particles and acid sites in zeolites. Electron microscopy (EM) and X-ray absorption near-edge structure (XANES) are used to characterize metal particles that are too small for traditional X-ray diffraction (XRD). Extended X-ray absorption fine structure (EXAFS) spectroscopy is unique for the characterization of the size of zeolite-supported metal particles (78-85). Temperature-programmed reduction (TPR) provides relevant information on the extent of reduction and the activation energy of the reduction process. After adsorption of probe molecules such as CO or NO, Fourier transform infrared (FTIR) spectroscopy provides a wealth of information on the electronic and geometric nature of the adsorbing sites. The size of the particle can be derived either from transmission electron microscopy (TEM) micrographs or from the coordination number derived from EXAFS or, within certain limitations, from traditional dispersion measurements. The latter often make use of chemisorption of hydrogen or carbon monoxide, the adsorbed quantity of which can be derived from temperature-programmed desorption (TPD), which also shows the presence of more than one type of adsorbed species.

ZEOLITE-SUPPORTED CATALYSTS

133

For ferromagnetic cobalt particles in zeolite X, spin-echo ferromagnetic resonance has been used to obtain unique structural information (86). In addition, study of the catalytic signature of metalheolite catalysts has been used by the groups of Jacobs (87), Lunsford (88), and Sachtler (47,73,89). Bransted acid protons are identified by their 0-H vibration (90,91) in FTIR or indirectly, by using guest molecules such as pyridine or trimethylphosphine (92,93). An ingenious method to characterize acid sites in zeolites was introduced by Kazansky et al., who showed by diffuse reflection IR spectroscopy that physisorbed HZ clearly discerns different types of acid sites (94). Also, the weak adsorption of CO on Bransted and Lewis acid sites has been used for their identification by FTIR (95). The characterization of the acid sites was achieved also by proton NMR (96). 11.

Preparation of Metal Zeolites

A. GENERAL STRATEGIES Various techniques are available for the introduction of metals into zeolites. Ion exchange and impregnation, e.g., by the “incipient wetness” or “imbibement” techniques, are often used. The former method introduces preferentially the cation into the zeolite, whereas the latter method also incorporates an equivalent number of anions. In both cases, the introduction of ions has to be followed by calcination and reduction steps. For zeolitebased catalysts the ion-exchange method is often preferred (97,98). An alternative route makes use of neutral complexes such as metal carbonyls with zero-valent metal atoms or clusters. Volatile metal precursors are readily introduced into zeolites via vapor-phase adsorption (99,100). Less volatile complexes can be impregnated into zeolites from solution. The use of organometallic precursors is advantageous in several respects. First, the metal catalysts prepared with this method generally exhibit high dispersion, because carbonyl or other ligands can be readily removed at relatively low temperatures. Second, the effect of anions such as chloride, which often exists for samples prepared from salt impregnation or, though to a lesser extent, via ion exchange, is eliminated. A third advantage is that no reoxidation will occur for these samples because no protons are present. An elegant variant of either method is the “ship-in-a-bottle” technique (101),which constructs large complexes in zeolite cages from smaller precursors that can pass through the cage apertures. Research related to this subject will be reviewed in Section VII. The goal of most preparation methods is to obtain reduced metal particles, entrapped inside zeolite cages or channels; their distribution should be homogeneous throughout the zeolite crystals and larger particles at the external surface should be absent. For

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some metals, e.g., Pd, this goal has been fairly approached by using the ion-exchange technique. It appears, however, necessary to optimize each of the steps: the ion-exchange proper, washing, drying, calcination, and reduction. We shall therefore discuss the ion-exchange method in greater detail.

B. ION-EXCHANGE TECHNIQUE For second- and third-row transition metals and rare earth metals, ion exchange is frequently used. The effect of the preparation conditions on the location of ionic precursors, their ligancy, and the size of metal particles after reduction have been the subject of much fundamental research (48,102,103). The pioneering work of Gallezot and co-workers using X-ray diffraction and absorption techniques and electron microscopy has provided the scientific basis for understanding the fundamental chemistry of these processes (20,I04-106). Dynamic techniques, such as temperatureprogrammed oxidation, reduction, and desorption, have been combined with physical methods, including EXAFS, by the group of the present authors for the identification of the elementary processes in the genesis of metal particles (71,72,76,77,107,108). UV-Vis diffuse reflectance spectroscopy and ’29Xe NMR have been successfully employed in the identification of location and coordination of many transition metal ions in zeolites (77,109-111). This spectroscopy has also been used by the research groups of Klier and Schoonheydt in the investigation of the coordination chemistry of metal ions by various ligands and the effects of ligands on the location of the ions in zeolites (112-114). The actual ion-exchange step places ligated metal atoms into the main channel system or, in the case of faujasites, in the supercages. The purpose of the subsequent calcination step is to remove the water and to destroy the ligands, e.g., the ammine groups of Pt (NH&*+ ions. A secondary effect of this calcination is that the metal ions, after being deprived of their ligands, will migrate into smaller zeolite cages, e.g., sodalite cages or hexagonal prisms. There are, however, ways to avoid this. Another side effect of calcination is often that decomposing NH3 groups will reduce the metal ions prematurely. If this “autoreduction” takes place, the metal atoms tend to agglomerate to large particles, which is, of course, undesirable. These particles may become oxidized, after the decomposition of the ammine ligands is completed. Very cautious calcination programs have been proposed by Gallezot et al. in order to minimize autoreduction. Ion exchange of Na+ or K + forms of zeolites normally starts with a dilute aqueous solution of ionic precursors. The ions are solvated or coordinated

ZEOLITE-SUPPOI1TED CATALYSTS

135

by H z 0 molecules in solution. Attention must be paid to the pH of the zeolite slurry when transition metal ions are exchanged in order to avoid hydrolysis of the metal ions prior to their exchange. Multivalent ions such as V3+ and Ga3+exist as free cations only in strongly acidic solutions incompatible with zeolite stability (115). For second- and third-row Group VIII metals neutral ligands such as NH3, which are coordinatively stronger than H20, can be used to prevent hydrolysis in neutral solution. Indeed, Pt (NH&’+ and Pd(NH&’+ are most often used in ion exchange of Pt2+and Pd’+. Amminated trivalent noble metal ions, such as [ R u ( N H ~ ) ~ ] ~ + , [Rh(NH&CI]’+, and [IT(NH~)~C~]’+, are also frequently used. The amminated metal ions persist in a broad pH range. Besides modifying the precipitation constant of metal ions, ammine ligands also affect the distribution of the ions in zeolite cages. A recent study shows that Co2+ions that were introduced into NaY as CO(H~O)~‘+ ions can swiftly migrate into sodalite cages after dehydration at low temperatures. However, when C O ( N H ~ ) ~ ~ + precursors are used, autoreduction leads to monoammine Co2+ ions, which remain in the supercages even at elevated temperatures (77). Another factor to be considered for pH adjustment is that a low pH inhibits the reduction of metal ions due to a high concentration of protons, particularly for less reducible metals. In some cases, controlled hydrolysis at high pH is desired to achieve the reduction of less reducible metal ions, such as Ni2+and Co2+,at a practical temperature range (116,117). In the preparation of bimetal zeolite catalysts, the pH of each individual exchange step must be carefully chosen to meet the designs for the final catalysts. Another parameter of obvious importance is the temperature at which ion exchange is carried out. Although this is often done at ambient temperature, higher temperatures have been used to achieve higher metal loadings (118-120). Another nontrivial parameter is the pH of water used for washing the samples during filtration after ion exchange. This has to be compatible with the pH during ion exchange; otherwise, a new equilibrium will be established. Washing with water at a pH of 7 will, for instance, reintroduce acid sites in neutralized catalysts. Conversely, Ni/NaY that had been ion exchanged at a pH of 7 and then washed with a NaOH solution of a pH of 10.5 leads to hydrolysis of Ni’’ ions inside the supercages of Nay. In agreement with the chemistry suggested by Suzuki, the reduction of such a sample, after calcination, required a much lower temperature than a similar sample that had been washed with doubly distilled water at pH 6 (121). Likewise, an effect of pH on the reduction of Cu2+ was observed (73). Ions of first-row transition metals are notoriously difficult to reduce, in particular when they are located in small zeolite cages, where stabilization of the ions is high. For example, Fe2+ ions are virtually irreducible in zeolites

136


(122-124), although they can be reduced when supported on Si02 or A1203 (125). Various causes have been identified as contributing to the low reducibility of these ions in zeolites. First, the chemistry of the reduction of exchanged ions is different from that of conventional metal precursors, e.g., metal oxides on amorphous supports. Whereas reduction of an oxide with hydrogen yields H 2 0 , hydrogen reduction of a metal ion yields protons. The equilibrium M2+ + HZ S Mo + 2H'

(1)

will be shifted toward the left at high proton concentration. A second difference is the tendency of the transition metal ions Fe2+, Co2+, and Ni2+ on dehydration to occupy preferentially the sodalite cages and hexagonal prisms of faujasite zeolites of high Si/Al ratio (110,126-130). In these small cages, the M2+ ions are favorably coordinated to the framework oxygens. This bonding is generally much stronger for multivalent than for monovalent ions. A third important cause, impeding the reduction of ions in small cages, is that the primary product of this reduction will be an isolated metal atom that strongly interacts with zeolite protons in the small cages, whereas metal-metal bonds can be formed only after migration of this atom out of its cage. In this respect, reduction of ions in supercages is much more favorable because the primary product can be a particle with strong metal-metal bonds. For all these reasons, the activation energy for the reduction of M2+ ions in hexagonal prisms of a faujasite lattice is much higher than for reduction of the same ions in supercages (116-117). In fact, reoxidation of Mo to M2+ under inert atmosphere in zeolite cages has been observed in the presence of protons (208). For platinum this reoxidation is limited to isolated atoms, but for first-row transition metals it has also been documented for small metal particles (131). The location of ions seems to be more critical in zeolite Y than in x , based on an X-ray diffraction study that showed simultaneous disappearance of Ni2+ions from all cages in the course of reduction (132). A method to prevent ion-exchanged first-row ions from escaping to small cages during calcination was proposed by Suzuki et al. (I14,117,133). They used a high pH to induce hydrolysis of metal ions inside supercages prior to calcination. They showed by means of temperature-programmed reduction of Ni2+and Co2+ions in zeolite Y that these ions are very difficult to reduce when they are located in sodalite cages and/or hexagonal prisms. However, the temperature required for their reduction by H2 is much lower if the ionic precursors are hydrolyzed after ion exchange. For hydrolyzed Ni and Co the reduction follows a different chemistry. In this case calcination is likely to result in the formation of hydroxide or oxide particles: M(OH)I -+ MO

+ H20

(2)


137

The reduction of these precursors will produce multiatom particles and water vapor. This reduction process has a much smaller activation energy compared to the reduction of ion precursors, resulting in isolated atoms and protons. A second strategy to prevent escape of transition metal ions into small zeolite cages is blocking of these cages with less reducible and catalytically inactive cations, such as Mn”, Ca2+,or Sr2+.Considerable enhancement of the reduction of Ni2+and Co” has been achieved in this way (79,108,133). Although there are no universally applicable ion-exchange conditions for all metal zeolites, specific procedures can be identified to obtain catalysts by design. The reducibility of metal ions in zeolites largely depends on the reduction potential of the ions and their interaction with zeolites. Ions located at different sites of a zeolite generally have different reducibilities due to different degrees of interaction with zeolites and different degrees of accessibility. Guest molecules, such as ammonia and water, compete with zeolite oxygens as ligands for metal ions; such molecules therefore can affect the location of the ions in zeolites (112). Ammonia also neutralizes the protons that are generated during reduction, thus shifting the reduction equilibrium [Eq. (l)] toward high reduction (134). By neutralizing the protons, ammonia also affects the catalytic activity of metal zeolite catalysts, e.g., in hydrocarbon conversion reactions (47). C. EFFECTS OF CALCINATION CONDITIONS ON M ~ A L DISPERSION Removal of coordinating ligands by careful calcination prior to reduction, is therefore extremely important for metal/zeolite catalysts, because it controls the cation locations and thus the metal particle growth mechanism during subsequent reduction. It has been demonstrated that the ultimate metal dispersion depends on the temperature of the calcination (50,69,71,79,107). An optimum calcination temperature can be defined for obtaining maximum dispersion of metals in zeolites. In the calcination of amminated ion precursors, a low heating rate and a high 0 2 flow are preferred to prevent autoreduction of metal ions by ammonia ligands (106) and subsequent formation of metal oxide particles. The reduction of oxide particles can be discerned from that of metal ions, because the former will be reduced at a much lower temperature than the latter (76). Autoreduction has been found to cause a nonuniform distribution, including large metal particles (135-137). On the other hand, [Ru(NH&I3+/NaY ; dehydration of this particforms bulk Ru oxide on heating in 0 2 therefore, ular system is performed in vucuo or in an inert gas flow prior to reduction (138-140).

138


I

-22

I

I

94

I

I

210

I

I

326

I

I

'

442

Temperature ("C)

FIG. 1. Temperature-programmed oxidation of Pd(NH3)2+/NaY profiles of Ot consumption and NZevolution (71).

ZEOLITE-SUPPOFTED CATALYSTS

139

X-Ray diffraction has been used to identify the locations of noble metal ions in faujasites after subjecting the amminated precursors to a variety of calcination conditions (141-148). These ions remain in the supercages after partial oxidation of the ammine ligands at low temperature, but they migrate into sodalite cages after calcination at high temperature. This has been shown in detail for Pt and Pd ions in zeolite Y. The high negative charge density in the small cavities provides the main driving force for the migration of multivalent transition metal ions from supercages to sodalite cages and hexagonal prisms in faujasites of moderate Al/Si ratio (149). With Pd(NH3)42+/NaY,it was found that tetraammine Pd divalent ions remain intact only up to 150°C (209). Dynamic temperatureprogrammed oxidation (TPO) with on-line mass spectrometry has been used to monitor oxygen consumption and nitrogen evolution in the process of ammine oxidation (71). It is evident from the profiles shown in Fig. 1 that 0 2 consumption and NZformation are essentially each other's mirror images when ammine ligands are oxidized. For complete deammination and concomitant migration from large to small cages, Pt ions require a higher temperature than do Pd ions (146). UV-Vis diffuse reflectance and EXAFS spectroscopic studies of the calcination process provide additional information on the ligancy of Pd2+ ions. As shown in Fig. 2, cis-diammine Pd2+ ions, which absorb at 27.6 kK, are predominant after calcination at 250°C. These ions remain in the supercages because they are unable to traverse the windows of 2.2 A between supercages and sodalite cages. After calcination at a temperature T, above 300"C, the absorption at 20.9 kK corresponds to Pd2+ ions coordinated only by framework oxygens. The migration of naked Pd2+ ions into sodalite cages occurs instantaneously when the last two ammine ligands are oxidized. At 500"C, a fraction of Pdz+ ions is found to occupy hexagonal prisms, as evidenced by the increased A13+ and Si4+second nearest neighbors in the Fourier transforms of the EXAFS data (209). At 275°C cis-diammine Pd" and naked Pd2+ ions coexist (Fig. 2). Monoammine Pd2+ ions appear not to be stable at static conditions. The conclusions derived from the dynamic studies, UV-Vis diffuse reflectance, and EXAFS spectroscopies are in good agreement with earlier XRD results (143,144,148). Although high calcination temperatures are required for the complete removal of ammine ligands, they induce migration of the bare metal ions into small cages, which is usually undesirable, because the activation energy for reducing such isolated metal ions is rather high. In the case of Pt, the temperature required for reducing Pt ions in sodalite cages and moving the Pt atoms to the supercages is so high that large particles are formed (107,146,150,151).Therefore it is desirable to have a strategy that permits

WOLFGANG M . H . SACHTLER AND ZONGCHAO ZHANG

140

1.o 8

e

U

0.5

0.0

1.5

2.0

2.5

Wavenumber (x

1.0

lo'

3.0

3.5

3.0

3.5

Cm")

-

A

8

% 0.5 -

0.0 1.5

2.0

2.5

Wavenumber (x 1 o4 Cm")

FIG. 2. UV-Vis diffuse reflectance spectra of Pd(NH3),*+/NaY after calcination at (a) 150°C, (b) 250°C, (c) 300°C, (d) 4 W C , (e) 500°C, and (f) 275°C (109).

complete ammine destruction, but prevents migration of the noble metal ions into sodalite cages. This goal can be achieved by using ions such as Mg2+ and Ca2+ to block the hexagonal prisms and sodalite cages prior to the introduction of Pd ions (152). This strategy will be further discussed in Section II,D. Blocking of sodalite cages and hexagonal prisms with ions of high charge density is particularly useful for the reduction of less reducible metals, such as Co or Ni. Ca2+,Sr2+,and Mn2+ions have been used as blocking ions for the hidden sites of zeolite Y to force Co2+ and Ni2+ ions in supercages (108,133,153). Other ions, such as Fe2+,La3+,and Ce3+,not only block the


14 1

hidden sites but they also assist in anchoring the reduced metal particles to the cage walls (79,80,154,155). Ba2+ appears to be too large to enter the small cages. Mg2+ ions, which readily hydrolyze, are able to block the sodalite cages but not the hexagonal prisms (152,153). The extent of filling small cages with blocking ions must be near 100%to be effective. Multiple ion-exchange steps are necessary to achieve this. Heating of the sample between successive ion-exchange steps enables the ions to migrate into the small cages. We thus distinguish “exchange” between solution and supercages and “migration,” i.e., place exchange between ions in large and small cages. In this terminology, an exchange step will be as follows: Mee(HzO),Z’

+ NasupNasHY

-+

2Na’

+ Mes(HzO),.supNasHY

(3)

where MeB(H20)2+ is the solvated ion in solution, SUP represents the supercage occupancy of the ions, and SH represents their sodalite and hexagonal prism occupancy. The migration step will then be M e B ( H 2 0 X , ~ ~ ~ N a ~ NasupMea.sHY HY + XHZO +

(4)

In this step, multivalent blocking MeB ions lose their water ligands and migrate into sodalite cages and hexagonal prisms. Na+ ions in the hidden sites are concomitantly forced into supercages. It should be pointed out that the heating rate in the ion migration step should be controlled to prevent hydrolysis of the blocking ions in supercages. For transition metals, such migrations are accompanied by ligand replacement. The chemistry and catalysis of transition metal complexes in zeolites have been recently reviewed by Lunsford (156). In these catalysts, it can be desirable to replace, for example, ammonia by pyridine. Zeolite-encaged cobalt complexes, for example, exhibit high potential for oxygen activation. Although the atmosphere outside the zeolite is oxidizing during calcination, release or decomposition of ammine ligands from ions such as Pt(NH3)42+or Pd(NH3)42+locally creates a reducing atmosphere inside zeolite cavities. This can lead to reduction of transition metal ions during calcination. As a result of this autoreduction, metal particles and protons will be formed. This uncontrolled reduction may lead to undesired particle sizes. Once ammonia and its decomposition products have left the zeolite bed, further calcination will convert metal particles to metal oxide particles. In the case of Pt, the originally divalent ion Pt(NH3)42+will thus first be partially reduced to zerovalent PtO then oxidized to Pt02 with tetravalent Pt4+. The TPR analysis of such samples will obviously lead to spurious amounts of hydrogen consumption. Prevention of autoreduction of Pt is not easy, in particular for zeolites with unidimensional pores such as mordenite or L zeolite. One of the motivations for emphasizing Pd/NaY in this review and in

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the research on which it is based is that for this system autoreduction can be prevented by carrying out calcination under the conditions identified first by Gallezot: a flow of pure O2 of 1 atm and a high flow rate of 2000 ml/min/g, combined with a very slow heating program of 0.5 K/min. Even slight deviations from this, e.g., a heating rate of l K/min instead of 0.5 K/min, lead to visible sample color changes that can be traced back to calcination. D. EFFECTSOF REDUCTION CONDITIONS Reduction of transition metal ions in zeolites is usually carried out in flowing HZ after calcination. Reduction in solution with other agents, such as NaBH4 and hydrazine, has also been reported (157). Some transition metal ions can also be reduced by Na or Cd vapor, NH3, or CO (103,158-161). Hydrogen atoms produced by microwave discharge have been used in the reduction of Ni2+in zeolite X (162). The coproduct of reduction is another ion, i.e., a proton or a metal ion. CO is not an appropriate reductant because it does not readily form ions; it has, however, been suggested that CO might react with lattice oxygen, creating COz and Lewis centers (163). The reducibility of metal ions in zeolites depends on a variety of factors, including location of the ion in certain cages, its accessibility, its coordination with ligands, effects due to other coexisting ions (site blocking or metal anchoring), zeolite structure, proton concentration, Si/Al ratio or zeolite acidity, and metal loading. The temperature required for reduction reflects the number of coordinating ligands and the strength of their bonding to the transition metal ion. In Table I, data are presented for Pd(NH3),'+ ions in NaY with different degrees of deammination owing to previous calcination to different temperatures. The TPR peaks are seen to shift to higher temperatures when deammination is more extensive; oxygen ions of the zeolite cage wall replace these ammine groups as ligands to the Pd2+ ion. These coordinations are identified by diffuse reflectance spectroscopy (DRS). TABLE I Efferts of Deammination of Pd(N&)*'' ~calclnallo" ("C)

100 200 250 500 a

From Ref. 71. From Ref. 109.

TPR Peak ("CY

Ions

Pd Coordinationb


143

Reduction of Pd(NH&2+ in NaY leads to low dispersion ( 7 4 , in part due to the high mobility of the metal precursor through the zeolite channel system and in part due to the absence of “chemical anchors.” Higher H/Pd ratios are obtained for the reduction at identical temperatures of partially deamminated or naked Pd divalent ions, which are much less mobile in the same zeolite. An important factor in obtaining a high degree of metal dispersion is the presence of zeolitic protons, which act as chemical anchors for reduced Pd atoms or particles. In the reduction of only partially deamminated samples, ammonia will be released, which neutralizes acid sites, thus preventing them from becoming chemical anchors (78,152,164). As particle growth is favored at elevated temperatures , the heat released by proton neutralization might also contribute to the lower degree of metal dispersion. The findings on Pd are also valid for other transition metals prepared by exchange of amminated ions into the zeolite, i.e., Pt(NH&*+, [Rh(NH&C1I2+, and [Ir(NH&C1I2+ (102). Ions at different sites have different electrochemical potentials. In general, ions in hexagonal prisms are most difficult to reduce due to their stabilization in small cages of high negative charge density. In addition, such ions are not readily accessible for H2 molecules. As mentioned, blocking of hexagonal prisms or sodalite cages with Ca2+, Mg2+, Sr2+, or Mn2+ ions prior to ion exchange forces transition metal ions, e.g., Ni2+, Co2+, or Pd2+,to stay in supercages where their reduction is easy, as is manifest from a shift to lower temperatures of the TPR peaks (79,108,133,152). For example, bare Pd2+ ions in NaY readily migrate into sodalite cages above 300°C (109).However, when these cages have been filled with Ca2+ or Mg2+ions, the TPR peak of the Pd2+ions is shifted by 70°C to a lower temperature, as illustrated by Fig. 3. Bivalent cations also affect the proton distribution: if small cages have been filled, e.g., with Mg2+ions, the protons that are created during the reduction of a transition metal ion will predominantly stay in the supercages, where they interact with the metal clusters, as will be described in more detail below. Because such adducts can act as very active sites, a catalyst promoter effect of Ca2+and Mg2+on reduced Pd has therefore been attributed, in part, to the enhanced concentration of metal-proton adducts in accessible supercages (165). The reduction of less reducible transition metal ions can be significantly enhanced by readily reducible noble metals. When ion pairs are formed between the less reducible metal ions and the ions of noble metals, both ions enhance each other’s reduction, as was found for the pair Pd-Co (77,166). The mechanism of this phenomenon will be discussed in Section IV. Under conditions in which only one of two elements present in the same zeolite is reduced, chemical interaction of the reduced clusters with ions of the unre-

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I

r 300

Temperature ("C) FIG.3. Temperature-programmed reduction profiles of (a) Pd/NaY and (b) Pd/MgY after calcination at 500°C (152).

duced element has been detected. Because ions interact more strongly with zeolite cage walls than with reduced metal particles, this interaction forms a second example of chemical anchoring. It diminishes sintering of the reduced metal particles at high temperatures. This phenomenon has been established for a number of bimetal/zeolite systems, namely, PtFeNaY, PdFeNaY, and RhCrNaY (79,80,155). The benefits can be mutual: the element that is reduced first often promotes reduction of other coexisting transition metal ions to a lower oxidation state. It was shown that the TPR peak ascribed to the reduction of Fe3+to Fez+ is shifted from 350°C in FeNaY to 150°C in PdFeNaY (155). Fe2+ions strongly interact with Pd particles, as evidenced by the Mossbauer spectra and a CO FTIR spectrum of reduced PdFeNaY that is distinctly different from that of the reduced PdNaY. Remarkably, the interaction of Fe2+ with Pd particles prevents their sintering even at high temperatures. This is shown in Fig. 4; whereas the Fourier transform of the EXAFS data for PdNaY displays a large intensity of the Pd-Pd nearest neighbor bond, that for PdFeNaY mainly shows Pd-Fe bonds, indicating smaller Pd particles. After reduction at 500"C, the average Pd particle in Pd/NaY is larger than Pd13, but in PdFeNaY reduced under identical conditions, the Fe2+-anchored Pd particles are smaller than Pd, . Metal loading and sample preparation conditions are critical for the anchoring effect. This might account for the different observations on PdFeNaY reported by Moller and Bein (83). Zeolites of different structures and pore opening widths exhibit different reduction profiles for the same metal precursors and the same pretreatment

145


Pd/NaY

1

2

3

4

5

6

7

8

7

8

FePd/NaY

1

2

3

4

5

6

RA FIG.4. EXAFS Fourier transform spectra of Pd/NaY and PdFe/NaY after calcination to 500°C and reduction at 500°C (155).

conditions. For example, Pt(NH&*' ions can be readily introduced into zeolite Y of faujasite structure by ion exchange, and homogeneous Pt particle distribution inside zeolite cages can be obtained (79,146). Temperatureprogrammed reduction profiles are typical of bipositive Pt2+ ions in either supercages or sodalite cages (79). With zeolite L, however, similar ionexchange and pretreatment procedures lead to formation of PtO, and some

146

WOLFGANG M . H. SACHTLER AND ZONGCHAO ZHANG

Pt2+ ions that are reduced at rather low temperatures (53). As a consequence, large Pt particles are formed after reduction. On the other hand, incipient wetness impregnation of zeolite L with Pt(NH3)4’+ followed by calcination produces Pt4+ions. Small Pt particles are generated after reduction of the ions at higher temperature. It seems likely that a large concentration gradient is needed to drive Pt(NH&’+ ions into the one-dimensional channels of zeolite L. Impregnation satisfies this condition. Because ion exchange is normally carried out with dilute solutions, the concentration gradient between solution and zeolite is small. Large pore window diameters and three-dimensional accessibility in faujasites may be important for the success of ion exchange with zeolites X and Y. Ni2+ ions are more reducible in L and mordenite zeolites than in Y zeolite (167-169). This dependence of the reducibility of Ni2+ ions on the zeolite structure might be related to the coordination and the accessibility of Ni” for H2. In mordenite and L, the Ni‘+ ions are located in open sites and are less coordinated by framework oxygen. In zeolite Y they are favorably located in hexagonal prisms. The reducibility of transition metal ions also depends on the Si/Al ratio (16,170,171). It is about 2.5 for zeolite Y and 1.2 for zeolite X, although both have the same faujasite structure. Zeolite A, which has supercage and sodalite cage dimensions similar to faujasite, has an Si/Al ratio of unity. The reducibility of Ni2+ions increases with decreasing Si/Al ratio in the order Y < X < A. The same trend was observed for the reducibility of Ni” and Pd2+ ions in untreated zeolite Y (Si/Al = 2.5) and dealuminated Y (%/A1 = 5 ) (172). The ions are less reducible in the dealuminated form. Because the strength of zeolite acidity increases as the Si/Al ratio increases, the distribution of ions in the zeolites of different Si/Al ratios may not be the only factor determining the reducibility. The strongly acidic protons produced during the reduction of metal ions obviously shift the equilibrium in Eq. (1) to the left, thus improving the reducibility of transition metal ions. As mentioned before, neutralization of protons by ammonia has been shown to promote the reduction of Rh+ to Rho (134). The effect of metal loading on the reducibility was examined with PdNaY (71). For samples calcined at 500”C, the TPR peak maximum shifts from 190 to 150°C; the Pd loading increases from 2.0 to 6.7 wt%. This has been attributed to the formation of ion pairs in sodalite cages. The reduction conditions are important for the resultant metal dispersion. TEM and radial electron distribution (RED) evidence shows that reduction of Ir ions in an H2 flow results in much smaller Ir aggregates than reduction under static HZ (173). The location of metal particles after reduction has been examined by a variety of techniques. Transmission electron microscopy provides direct information on particle-size distribution and its homogeneity throughout the zeolite crystallites (20,106,174). X-Ray scattering and EXAFS have also

ZEOLITE- SUPPOKI'ED CATALYSTS

147

frequently been used to determine the average metal particle size in zeolites (20,78,79,106,164,173). Care should be taken in determining metal dispersion from chemisorption of H2 and CO (70), because these molecules have been found to alter the strong interaction of metal particles with the zeolite matrix. For very small clusters, chemisorption of either probe can lead to erroneous values of the particle size (25,78,104,164). This issue will be addressed in more detail in Section 111. The location of reduced Pt particles in Y and L has been chemically probed by the catalytic reduction of methyl viologen (MV") and [Fe(CN),]'- in H2 (175). This method is based on the slow exchange of MV2+ between zeolite and liquid solvent and the absence of [Fe(CN)6l3anions in zeolite cages. As the reduction of the complex is catalyzed by Pt, the rate of the concomitant color change, monitored by UV-Vis spectroscopy, is indicative for Pt particles inside the cage cavities or on the external surface. It has been reported that Pt/L prepared from a Pt(acac);! precursor results exclusively in Pt particles inside the channels of zeolite L, but use of Pt(NH3)42+produces Pt particles both on the external surface and in the channels. With zeolite Y, either precursor generates some external Pt particles. However, it is questionable whether this observation may be generalized, because the Pt/L sample for which no external Pt particles were detected had a Pt loading of only 20 ppm. Scherzer and Fort report that zeolite Y after exchange with Fe2+,Co3+,Ni2+,and Cu2+ions from aqueous solution of their alkaliferrocyanides contained insoluble metal ferrocyanides inside zeolite cavities (176). Reduction with H2 of these cyanides yields monometallic or bimetallic particles. With the same technique, Fe, FeRu, and FeCo particles are produced in zeolite ZSM-5 (177). These results are, however, at variance with the assumption that no alkaliferrocyanide anions are present in the zeolite cages. The mechanism of metal particle formation in zeolite Y during H2 reduction has been reviewed by Homeyer and Sachtler (69,74). The location and particle sizes of noble metal after reduction largely depend on the calcination programs. The oxidation of the ammine ligands in Pt(NH3)42+ and Pd(NH3)42+during calcination gives rise to an abrupt transition, e.g., of the TPR pattern, when the metal ion jumps from the supercage to the sodalite cage. For Pd(NH3)42+in Nay and a slow calcination ramp, this jump occurs between 250 and 300°C. Simultaneously, the color changes from yellow to pink (71,109). For Pt(NH3)42+ions this transition occurs at much higher temperature (20), although the tetrammine complex is destroyed at 250°C (50). Evan after virtually complete oxidation of the ammine ligands at 360"C, a dark grey color indicates the presence of Pt oxide particles in supercages, presumably due to autoreduction and subsequent oxidation. EXAFS shows that the coordination number of Pd particles after reduction with H2 of Pd(NH3)22+in supercages at 200°C is 2.4, corresponding to a nu-

148


clearity of Pd, and P& (164). In the case of platinum, reduction of precursors in supercages at 360°C produces particles with a coordination number of 7.7 (79). The change of color for Pt/NaY from black, after calcination at 360"C, to brownish, after calcination at 550"C, indicates reaction between [PtO], particles and zeolite protons: [PtO],

+ 2nH+ + nPt2+ + nHzO

(5)

where n is a small integer. The Pt2+ ions migrate to sodalite cages (107), where their reduction requires a significantly higher temperature than that of the [PtO], particles or Pt2+in supercages (100 to 400°C vs. -15 to 150°C). Similarly, the TPR of Pd2+ions in supercages peaks at 150"C, whereas that of the same ions in sodalite cages peaks at 190°C. On amorphous supports the reduced metal consists always of multiatomic clusters, although monoatomically dispersed Pt and Pd have been reported in the zeolite Y. A prerequisite for their formation is that the ionic precursors have been located in sodalite cages; reduction has to be carried out under fairly mild conditions (143,164). EXAFS of such samples displays only weak interaction of Pd atoms with framework atoms, because Pd atoms have no heavy atoms among their first neighbors (164). The migration of reduced Pd or Pt atoms from sodalite cages to supercages has to overcome not only the geometric restriction that a 2.75- to 2.76-A atom has to pass through a 2.2-A-wide aperture, but it also has to break the bond due to proton anchoring (152,164). Temperature-programmed desorption measurements suggest that migration of Pd atoms from sodalite cages to supercages is complete only at -350"C, as shown in Fig. 5. Sintering of Pd particles at elevated temperatures will, in general, lower the metal dispersion and hence the ratio of adsorbed hydrogen to metal atoms. This is observed for reduction temperatures, TR, above 350°C. An observed increase of the H/Pd ratio with TRbelow 350°C is attributed to migration of Pd atoms from sodalite cages to supercages. Various causes have been contemplated as to why isolated Pd atoms in a sodalite cage do not chemisorb hydrogen. One possibility is that dissociative adsorption of Hz will be much easier on a Pd, ensemble than on an isolated Pd atom. A second possibility is that the species formed during the "reduction" of an isolated Pd2+ ion by an HZmolecule is of the type [PdH#+. This species is registered in TPR as a reduced species because H2 is consumed; in EXAFS it shows absence of metal-metal bonds, but it is unable to adsorb Hz. When the Pd atom leaves the sodalite cage, the two protons are left behind. In the supercage the Pd atom will either adhere to an existing Pd cluster or act as a nucleus for a new cluster. In either case this process leads to an increase in the H/Pd ratio. The very low H/Pd ratio for zeolites containing most Pd as isolated atoms or as [PdH2I2+complexes obviously does not reflect the true metal dispersion.

149


0.80

0.70

0.60 +I

4 I

0.50

0.40

0.30 150

230

310

390

470

550

Reduction Temperature (“C)

FIG. 5 . HiPd ratio versus reduction temperature for Pd/NaY after calcination to 500°C;

A, 2 wt% Pd; 0 , 1 wt% Pd (72).

Another important manifestation of sodalite-entrapped species containing only one metal atom is observed for platinum. In this case the temperature required for the migration of atoms from sodalite cages to supercages is higher than in the case of Pd. Before this happens, a different process is observed, namely, reoxidation of Pt atoms by the protons in the same cages: Pt’

+ 2H+ -+ Pt2+ + Hi

(6)

This has been observed for Pt/NaY reduced at 400°C (150,151).Heating in an inert atmosphere, i.e., under conditions usually used for studying temperature-programmed desorption, releases H2. This high-temperature “TPD’ peak is shown in Fig. 6 near 450°C.The peak area of this reoxidation peak is larger when the temperature of the preceding calcination was higher, so that more Pt2+ions were “chased” into sodalite cages. This interpretation is confirmed by FTIR: protons in zeolite cages are visible as 0-H vibration bands. These bands are absent in the zeolite containing Pt2+ions only; the band intensity becomes significant during reduction of these ions with H2, but it decreases when protons are consumed in the reoxidation of PtO to Pt2+ (79).

150

WOLFGANG M . H. SACHTLER A N D ZONGCHAO ZHANG

a

0.310

0.230

0.150

0.070

-0.010

0.310 -

I

I

I

I

I

I

I

1

I

I

b

0.230

i?

L P s

-

/

0.150 -

c

\

\

/

0.070 -0.010 -

I

0

\

I \ \ \

, \

I

I

I

I

I

I

I

I

C

0.310

0.230

I

I

-

c

0.150 0.070 -

/

’ I

I

I

I I

-20.0

120.0

260.0

Temperature (“C)

400.0

\

\ \ \

I \ \

540.0

ZEOLITE- SU PPORED CATALYSTS

151

Once Pt atoms are in supercages they agglomerate to Pt particles. At higher temperatures these particles expand over several cages, as suggested by the large coordination number ( 1 1.5) for the Pt particles. Direct observation of these Pt particles by TEM reveals "grapes" filling several contiguous cages (143,178). The reason Pt atoms are less mobile than Pd atoms in migrating out of the sodalite cages is presumably the larger size of Pt in combination with its higher polarizability due to the d electron configuration. This leads to increased anchoring with protons. Several research groups have examined Pt and Pd particles in zeolities X and Y with TEM. The results consistently show that the metal particles obtained from ion-exchanged samples are fairly homogeneously distributed throughout the zeolite crystallite, even when the particles are larger than the supercages (79,143,178). This conclusion is further supported by X-ray photoelectron spectroscopy (XPS) studies (179,180). No accumulation of metal at the external surface is observed when proper preparation conditions have been followed; in fact, a surface depletion of 30% Pd was detected due to the final washing of an ion-exchanged sample (180). Autoreduction of Pt(NH3)42+and Pd(NH3)42+in zeolite X has been studied in detail by Schulz-Ekloff et a / . (137,177,181,182). They found that autoreduction and agglomeration depend on the degree of ion exchange, the extent of dehydration, and the medium used. A high water content in the atmosphere (Ar, He) favors lattice destruction, thus facilitating the growth of Pt particles to an average size of 4 nm. A slow heating rate in dry Ar produces smaller Pt particles of 3 nm. For low exchange levels, autoreduction in vacuo leads to a bimodal size distribution of metal particles, peaking at 1-2 and 3-4 nm, but for higher metal loadings a narrow size distribution was observed around 1-2 nm. The formation of a polynuclear ammine complex precursor, prior to autoreduction proper, has also been suggested (203). Pt particles of supercage size (- 1 nm) were found to prevail after HZ reduction of partially autoreduced samples. Bischoff et al. (182) studied FTIR of CO chemisorbed on Pt and found that CO is dissociatively adsorbed on supercage-sized Pt particles; surface carbon deposits were detected. Temperature-programmed desorption of chemisorbed oxygen (TPDO) from Pt/NaX samples that were prepared by a different autoreduction procedure shows that the TPDO peak maxima shift to higher temperature as the Pt particle size, determined by TEM, decreases (136). The

FIG. 6. Temperature-programmed hydrogen release profiles of Pt/NaY (7.4 wt% Pt) calcined at (a) 360"C, (b) 450°C, and (c) 550°C. Solid lines, after reduction at 550°C; dotted lines, after reduction at 400°C. The peaks at high temperatures for the dotted lines are attributed to reoxidation of Pt atoms by protons (150).

152

.

WOLFGANG M . H SACHTLER AND ZONGCHAO ZHANG

authors attribute the dissociation of CO to the enhanced surface free energy of very small Pt particles. The same authors also report that incomplete autoreduction of Ir(II1) in the chloropentammine Ir(II1) complex in NaX leads to one of two extremes (174): Ir that is autoreduced in vacuo is anchored to unreduced Ir ions and small particles are detected, but autoreduction in Ar produces large Ir particles observed in TEM. Fully reduced Ir particles of high dispersion are obtained in NaX by Hz reduction following careful calcination, in analogy to the behavior of Pt and Pd zeolites, as described above. The work of Bergeret and Gallezot shows that hydrogen reduction after low-temperature calcination (250°C) of Ir(NH3)5C12+in NaY leads to a bimodal distribution of Tr particles ( 300°C

CO’+(Z -O)I(NH,)

(B) TPO of C O ( H ~ O ) ~ ~ + / N ~ Y CO(H~O)~~+Tc

(C) TPO of Pd(NH3)4z’/NaY

-

CO2+(Z-0)6 (predominatesfar Tc > 500°C)

2

200°C

Co2+

Pd(NHd4”

-

Tc = 250°C

J. Pd(NH3)2’+

I

Tc

L

300°C

Pd2+

(D) TPO of [Pd(NH,),’+ + Co(NH3)2+]/NaY“

“Locations of Pd(NH3)d2+with respect to TC are the same as given in (C).

Protons are able to reoxidize selectively, for example, the Cu or Ni atoms in the bimetal particles with Pd or Pt. For CuPd/NaHY it has been suggested that two steps can be discerned (185):

+ Ht -+ Cu’Pd, + 4H2 + nH+ + nCuZt + Pd,, + f n H z

CuOPd,

(9) (10) EXAFS results confirm that the Cu atoms in reduced PtCu/NaHY and PdCu/NaHY are “leached out” of the PtCu and PdCu bimetal particles and migrate to small cages on heating to 500°C in an inert atmosphere (185,202). For PdCu in NaHY, additional evidence for oxidative leaching of Cu from the bimetal particles is obtained from an examination of the hydride decomposition. It is known that Pd metal exothermically forms hydrides (203,204). Two hydride phases are normally distinguished: (1) a - P d hydride, with a heat of solution of 6.0-8.8 kcal/mol (H2), and (2) P-Pd hydride, with a heat of sonCu+Pd,

163

ZEOLITE-SUPPOF3EDCATALYSTS

lution of 9.7 kcal/mol (H2). For bulk Pd both phases can coexist in equilibrium at normal conditions. In zeolites, the p-Pd hydride, which decomposes at about 80”C, is found only after significant sintering, e.g., when the Pd particles are quite large. The a-Pd hydride is observed also with very small Pd clusters in zeolites; it decomposes between 0 and 32”C, depending on the cluster size. This decomposition is easily monitored by TPD. As formation of both hydrides is markedly suppressed by alloying Pd with Cu, detection of the hydride is indicative for “leaching” of Cu from PdCu particles. Indeed the characteristic hydride decomposition peak is absent in Cu-rich bimetallic particles, but it is quite pronounced after extensive leaching of the Cu atoms. The PdsCu particles that are formed in PdCu/NaY after reduction to 5OO0C, and which have been identified by EXAFS, are unable to form a hydride (185). Leaching occurs in two steps: first Cuo is oxidized to Cu+ at 280°C [Eq. (9)]; subsequently Cu’ is oxidized to Cu2+ at about 500°C. The Cu+ ions formed at low temperature are apparently still attached to the Pd particles, but the Cu2+ ions migrate to different ionexchange sites. As a result, subsequent rereduction of Cu’ occurs at a much lower temperature than rereduction of the detached Cu2+ ions. In Fig. 13

I

I

I

I

I

I

1

-100

0

100

200

300

400

500

Temperature (“C)

FIG 13. Temperature-programmed reduction profiles of PdCu/NaY after reoxidation by protons of Cu in PdCu bimetal particles at (A) 350°C and (B) 500°C in an Ar flow. Positive peaks are due to Hzconsumption (rate given in arbitrary units) for the reduction of Cu+ (A) attached to Pd, and for the reduction of preleached Cu2+ ions; negative peaks are due to Hz evolution from hydride decomposition (185).

164


the hydride decomposition peaks in the 5% HdAr gas flow are shown as negative peaks. After leaching at 500"C, the hydride decomposition peak is clearly identified at 17°C in the TPD profile. The composition of the particles after extensive Cu leaching by zeolite protons, as determined by EXAFS, corresponds to P d O . We speculate that this corresponds to Pds particles anchored to the cage wall by a Cu+ ion.

V. Metal Redispersion Redispersion is the reverse of metal sintering. The subject has been briefly reviewed by Gallezot (102). To break the strong metal-metal bond in metal agglomerates, corrosive gaseous molecules or mobile protons in zeolites are necessary. A strategy is to first reoxidize the metal to ions, which migrate into small cages. After this redispersion the ions are reduced under conditions in which formation of small metal particles prevails. Another strategy for metal redispersion makes use of volatile intermediates such as Ni(C0)4 or PdC12. The oxidative leaching of Cu from Pd and Pt bimetal particles that was described in the previous section provides an example of an oxidative redispersion; in this case protons are the oxidizing agents. Oxidative redispersion of Pd and Cu aggregates in zeolite NaHY has been reported (76,143,205).First the metal (M = Pd or Cu) is oxidized to oxide particles:

In the case of Pd/NaY, the reaction given by Eq. (11) predominates below 350°C. The formation of palladium oxide is confirmed with TPR by its characteristic reduction profile near 0°C and the concomitant formation of water. The second step involves reaction of zeolite protons with the metal oxides: (MO),

+ 2nH+ -+ nM2 + nH20

(12)

For relatively small particles of PdO or CuO and high proton concentration in their vicinity, the reaction given by Eq. (12) is virtually complete at 400°C (76,205). The reaction is driven by the formation of water and the migration of the M2+ ions to sites with high electronegative charge density. The ions then have the same oxidation states and locations as they had after original calcination. For Cu/NaY this has been demonstrated by electron

ZEOLITE-SUPPOKTED CATALYSTS

165

spin resonance (ESR), and for Pd/NaY it was shown by TPR and X-ray techniques. For Ag/NaY the stability of the d10 configuration of the Ag+ ions contributes to the thermodynamic driving force. Thus, small Ag clusters inside supercages of zeolite Y can be redispersed to their ionic state in 0 2 at a temperature as low as 90°C (206). For metals such as Pd the TPR profiles reveal the presence of Pd oxide at low temperatures. Subsequent conversion of PdO to Pd2+ions is illustrated by the profiles in Fig. 14. Redispersion, as described so far, will, however, be incomplete when the size of the metal particles exceeds that of the zeolite supercages. Bergeret et al. showed that Pd aggregates of 1 nm are readily reoxidized to Pd2+ions. Larger particles are oxidized to the metal oxide but are not redispersed into ions (143). Similarly, large Cu aggregates are incompletely redispersed (205). This limitation is due to the exhaustion of protons in close vicinity of the oxide particles. Once all protons in this region have been replaced by metal ions, the rate of Eq. (12) becomes limited by the slow site exchange of Pd" ions, in the vicinity of the metal oxide particle, with protons at larger distances from it. Several methods have been published to remedy this situation. One strategy aims at bringing the primary dispersed metal ions back into the supercages, where their mobility is much higher (207). Treatment of partially redispersed Pd/HY with NH3 transforms naked Pd2+ ions in sodalite cages into [Pd(NH3)4]2+ions in supercages:

Simultaneously protons are transformed into NH4+ions: xH+

+ xNH3 -+

xNH4+

(14)

In this situation, ion migration becomes facile

because both species travel easily through the supercage channel system. Thus, mild heat treatment decreases the concentration of [Pd(NH3)4I2+in the vicinity of the undispersed PdO particle, whereas it increases that of the NH4+ ions. Subsequent calcination transforms the latter ions into protons that can attack the PdO particles via the reaction given by Eq. (12). Repetition of this procedure leads to complete redispersion of even highly sintered particles (207). Another strategy for the rejuvenation of metal/zeolite catalysts with very large metal particles achieves mobilization by the formation of volatile

166

WOLFGANG M , H , SACHTLER AND ZONGCHAO ZHANG

metal compounds. Chlorination in the presence of other gaseous molecules such as N2, H 2 0 , CO, and NO has been recommended for the redispersion of noble metal on A1203and SiOz (208,209); this is also applicable for redispersion of large metal particles on the external surface of the zeolites (210,211). Volatile metal chlorides are formed as a first step. For the redispersion of large Pd aggregates, PdC12 is formed, whereas PdCL is unstable. For Pt, the type of Pt chlorides formed has been reported to depend on the

r,,,,

=

200°C

Tho, = 400°C

T b r = 500°C

2 0

100

200

300

Temperature (“C)

FIG. 14. Temperature-programmed reduction profiles of Pd/NaY after reoxidation to different temperatures following initial calcination to 500°C and reduction at 500°C. The peak at about 0°C is attributed to the reduction of Pd oxide; the peak at about 180°C is attributed to the reduction of redispersed PdZ+ions in sodalite cages (76).

167


zeolite and its Si/Al ratio. Initially, large crystals of PtCI2 are formed on silicalite at low temperatures, but chlorination at 350°C for 2 h results in the loss of virtually all Pt from the sample. Reduction following chlorination of large Pt particles on H-ZSM-5 produces some small Pt particles in the zeolite channels. In contrast, no Pt is lost during the chlorination of 2 wt% Pt/L zeolite; Pt particles after reduction are smaller than 1 nm on the zeolite surface and within the channels. The saturation capacity of zeolites of different WAI ratios for Pt after chlorination has been compared by Foger and Jaeger, as shown in Fig. 15 (209). The dependence of the Pt concentra/

I / / f /

/ I

/ /

/

I

I 2

li

I

I

I

4 6 wt% Pt (before treatment)

8

FIG. 15. Effect of AI/Si ratio of various zeolites on their capacity to accommodate PtCI2.., H-L; 0,K-L; H, silinated H-L; A, HZSM-5; and V, silicalite; in 10%C12/N2at 670 K for 4 h (210).

168


tion on the A1 content of different zeolites after chlorination has been attributed to bonding of Pt chloride to the A1 after splitting of the Si-0-A1 linkage in an atmosphere of chlorine or hydrogen chloride. Chlorination of Pt/HL at 400°C for 4 h leads to PtCl this was confirmed by the HdPt ratio from quantitative TPR results and by XPS (209). In the chlorination of large Pt or Pd aggregates in proton-rich zeolites, no interaction of protons with the metal chlorides is apparent below 350"C, but above 400°C HCl has been detected, indicating the occurrence of ion exchange and the reaction given by Eq. (16). PdClr

+ 2Hf -+

Pd2+ + HCl

(16)

Pd redispersion in zeolite Y using this chemistry is achieved via sublimation of Pd,Cl2,, which has a fairly high vapor pressure and is strongly adsorbed in zeolite cages. Subsequent hydrogen reduction of the Pd,Cl2, at temperatures below 35032, where the reaction given by Eq. (16) is too slow, results in highly dispersed Pd particles in supercages and the evolution of HCI. Even physical mixtures of Pd powder with HY have been transformed to highly dispersed Pd inside the zeolite by making use of this chemistry (210). Molecules that are able to coordinate to the metal ions to form complex ions can facilitate the oxidative redispersion of large metal aggregates. Treatment of Pt or Pd catalysts with Clz in the presence of H20, CO, NO, and NH3 or with O2 followed by ammoniation was found to be more efficient than the use of C12 or 0 2 alone (207-209). For Pd particles as large as 2 nm in zeolite Y and for 8-nm Ni particles in mordenite, it has been reported that NO treatment can lead to complete ionic redispersion (212, 213). Combined methods such as X-ray diffraction, ESR, IR, and mass spectroscopies have been used to measure the extent of oxidative redispersion of Pd (212). Magnetic measurements and UV spectroscopy have been used in the study of Ni redispersion (213). The function of NO in the oxidative redispersion of large metal particles in zeolites is twofold; it acts as an oxidizing agent and it coordinates to the metal ions. As discussed in connection with the NH3 treatment (207), coordination with nonanchoring ligands increases the mobility of ions through the zeolite channel system. The high mobility of Pd(NO),*+ thus enables protons to diffuse toward the metal particles, where they will react according to Eq. (12). This seems to be the key for complete redispersion of large metal particles in H-zeolites. The CO molecule is a well-known coordinating ligand for many transition metals. Although CO has been reported to be a reducing agent for metal ions such as Cu2+ and Ag+ in zeolites (214,215), it also induces oxidative redispersion of Rh particles in zeolites at room temperature (104). With FTIR, it has been established that the presence of HzO is crucial for the reduction of Ag+ in zeolite according to the following reaction:

ZEOLITE-SUPPORED CATALYSTS

+ CO -+ Ag+-CO + Ag+ + 3H20 + 2Ag0 + COz + 2H,O+ Ag+

Ag+-CO

169

(17) (18)

The formation of C 0 2 was detected by an IR band at 2345 cm-'. The mechanism of the oxidative redispersion of Rh particles has been studied both for A1203- and zeolite-supported Rh. The hypothesis that dissociative adsorption of CO is responsible (104,197,216) has been ruled out by experimental evidence that surface OH groups on an Al2O3support are consumed as CO interacts with Rh crystallites (217). For Rh in protoncontaining zeolites it was already mentioned that the chemistry given by Eq. (19) has been verified by FTIR and by the quantitative detection of the Rh,

+ 2nCO + nH'

+ nRh'(C0)~

+ iH2

(19)

H2 evolution. A recent XPS study shows the involvement of both protons and CO in the oxidation of Rho to Rh+ (218). The result is in agreement with earlier observations that HzO vapor enhances oxidation of Rho to Rh' on A1203(219), because H 2 0 provides the surface hydroxyl groups and thus the protons needed in Eq. (19). A recent XPS study shows that disintegration of Rh particles is much more pronounced in HY than that in Nay. The Rh in Rh+(C0)2is coordinated to two CO ligands and to three oxygens at the Sn site. This coordination provides the Rh gem-carbonyl cation with a stable 18-electron configuration. Mechanistically, one important role of CO will be to break an Rh atom from a large Rh particle by forming a mobile Rh carbonyl, which subsequently reacts with protons to form the stable Rhf(C0)2 (195). Treatment with a mixture of CO and H20 at the same temperature leads to reduction and agglomeration of Rh+(C0)2 to Rhs(C0)16 (220). Redispersion via mobile carbonyls has also been reported for bimetal/ zeolite catalysts. In a recent study on PdNi/NaY it was found that reduction, after hydrolyzing the ion-exchanged Ni at high pH, produces large PdNi bimetal particles on the external surface of zeolite Y (221). However, after use as a catalyst in CO hydrogenation at elevated pressure, Ni and the Pd were both found redispersed into zeolite cages. Presumably, volatile carbony1 intermediates are instrumental in this redistribution of both metals during the catalytic process. VI.

Metal Particles from Neutral Complexes

Volatile transition metal complexes have been widely used for the production of zeolite-encaged metal particles or organometallic compounds (222,223). The resulting catalysts are active for a wide range of reactions (224) and in some cases are superior to other preparations. For those metals

170


whose ionic precursors are difficult to reduce, severe reduction conditions can be avoided by starting from an appropriate volatile precursor complex. Examples of intensively studied metal carbonyls in zeolites are Fe(CO)5, Fe2(CO)9, Fe3(CO)12,C02(CO)8, and Ni(C0)4. They can be introduced into zeolite cages either via the vapor phase or by impregnation from a paraffin solution (225-231). Most carbonyl complexes remain intact at low temperatures and under CO pressure (232,233). However, at elevated temperature or in vucuo, they decompose by releasing CO ligands. Decomposition of the loaded metal clusters can also be achieved by photochemical action or microware plasma (86,234,235). Fe3(CO)12is more stable than Fe(CO)5 or Fe2(CO)9 during vacuum and thermal treatment. Zeoliteencaged Fe clusters in NaY have been shown to be active for syngas conversion (231). In the presence of zeolite protons, the decomposition of the complexes is sometimes accompanied by oxidation of the metal. When Fe(CO)5 in the supercages of zeolite Y is irradiated with UV light, catalytic activity for butane isomerization is observed (236). This photoinduced activity depends on the alkali ions in the zeolites; it decreases in the order Li+ > Na+ 9 K+ > Rb', Cs+. Studies using diffuse reflectance, UV-Vis, or IR spectroscopies show that coordinatively unsaturated metal atoms are formed when CO is released. It was proposed that the main function of the zeolite is to stabilize the coordinatively unsaturated species by coordination with the lattice oxygen. The isolation of Fe(CO)5 in the cages prevents agglomeration leading to Fe3(CO)12.The mechanism of the photoinduced activation, including the oxidation state and the nuclearity of intermediate carbony1 species, has been studied by a variety of methods, including Mossbauer, IR, and EXAFS spectroscopies (237,238). The decomposition of C O ~ ( C Oin) ~faujasites has been studied in some detail. Low-temperature spin-echo ferromagnetic nuclear resonance spectroscopy shows that very small Co particles are formed in supercages of zeolite NaX by microwave plasma activation at low temperatures (86).In situ far-infrared spectroscopy revealed that adsorbed C O ~ ( C Ointeracts )~ with accessible supercage cations in NaY and COY (239). Carbonyl complexes of different Co nuclearity, such as Co4(C0)12 and Co(C0)4- , are also formed (227,228). In HY the Co atoms are oxidized to Co2+ ions by the zeolite protons. Nickel tetracarbonyl and nickel complexes with mixed carbonyl and alkylphosphine ligands in zeolite Y have been studied with EXAFS, NMR, and FTIR spectroscopies (85,240-242). The loading of Ni(C0)4 in dehydrated zeolite X can be as high as 28 wt%, corresponding to 2.75 Ni(C0)4 molecules per supercages. The IR spectrum of Ni(C0)4 was found to be affected by the cations inside the zeolite supercages. The spectrum of

ZEOLITE-SUPPORTEDCATALYSTS

171

Ni(C0)4 in dealuminated Nay (Si/AI > 400) shows one band at 2046 cm-', similar to that of tetrahedrally coordinated Ni(C0)4 in THF solution. No change of the Ni oxidation state and no loss of CO ligands after adsorption of Ni(C0)4 in alkali zeolite Y are detected with XANES and EXAFS spectroscopies. However, the appearance of four IR bands, which shift when the Ni(C0)4 loading or the alkali cations are varied, indicates an interaction of the type M+----0C-Ni, where M' = Na' or Li'. A reactive Ni(CO), intermediate was proposed to rationalize the observed formation of bridging CO ligands; i.e., it is assumed that this Ni(C0)3 reacts with Ni(C0)4. Also, the oxidation of Ni to Ni2+ is proposed to proceed via Ni(CO),. Thermal treatment of Ni(C0)4 in Nay to 50°C leads to a loss of one CO per Ni. At 100°C the adsorbed Ni(C0)4 is fully decomposed and small Ni particles are formed. A bimodal size distribution was found with one fraction of the Ni particles in the size of the supercage dimensions and another fraction larger than 5 nm. The latter particles are assumed to be located at the external surface of the zeolite crystallites. The reverse process of restoring Ni(C0)4 from the Ni particles and gaseous CO should be facile; however, IR bands typical for coordinated carbonyl ligands were not observed at 25- 100°C and 1 bar of CO (242). Introduction of PPh2CHMe2into Ni(C0)4-loaded NaY led to the formation of Ni(C0)3PPh2CHMe2,as evidenced by XANES, EXAFS, NMR, and IR spectroscopies (85). It appears impossible to reduce Mn or Cr ions in zeolite cages to zerovalent atoms by conventional means. Therefore, Cr(C0)6 and Mn2(CO)lo are preferred precursors in zeolite cages (232,243).In NaX, intact Cr(C0)6 strongly interacts with Na' under CO atmosphere, as evidenced by IR spectroscopy. On heating to 100°C for 30 min, partial decarbonylation takes place and Cr(C0)4 is formed. Its IR spectrum displays four bands at 2042, 1909, 1831, and 1784 cm-'. At 150-200°C in the presence of CO, Cr(CO)3 with trigonal pyramidal symmetry is formed and exhibits two intense bands at 1917 and 1767 cm-'. The process of decarbonylation of Cr carbonyls with Cro can be reversed at low temperatures in the presence of CO. Cr(C0)4 in NaX is inactive for butadiene hydrogenation at room temperature, but Cr(C0)3 shows high activity for this reaction with a remarkable selectivity of >97% toward cis-2-butene. Mo(CO)~and W(CO)6 in zeolite display chemistry similar to that of Cr(C0)6 (244-247). Anchored tricarbonyl species are generally considered to be the active subcarbonyl catalysts for the stereoselective hydrogenations of simple dienes to cis-2-olefins. [HzOs(CO)4]was introduced from the vapor phase into NaN3-treated zeolite Nay. EXAFS and IR spectroscopic results suggest the formation of trisomium carbonyl clusters inside the su-

172


percages (233). The yellow-colored triosmium carbonyl clusters in NaY cannot be extracted by [PPN][Cl] in acetone (248). On the other hand, [H20s(C0)4]is only physically adsorbed on HY, and it can be extracted. Metal carbonyl clusters permit the preparation of bimetal zeolite catalysts that are otherwise difficult to make. A case in point are zeolite-supported PtRe catalysts, with possible applications in catalytic reforming (249-252). In aqueous solution Re does not exist as a free cation at a pH compatible with zeolite stability. Direct ion exchange is therefore not feasible for Re; the same holds for Mo and W. Therefore, volatile carbonyl clusters, such as Re2(CO)lo,are most convenient and readily controllable alternatives for the introduction of metals such as Re into zeolites (75,253,254).It was shown that Re2(CO)lois readily adsorbed in the supercages of the zeolite (75). An interaction of the type Na *.. OC and H OC is suggested. Thermal decomposition of the carbonyl in the presence of hydrogen results in the formation of metal particles via a Re hydridocarbonyl intermediate [HRe(CO)& (n = 1 or 2). Applying this procedure to a Pt/NaY catalyst that is prepared from Pt(NH3)42+by conventional ion exchange with Nay, followed by reduction, provides a convenient path to prepare bimetal PtRe catalysts. The kinetics of the carbonyl decomposition depends markedly on the presence of reduced Pt in the zeolite. When Re2(CO)lois decomposed in metal-free Nay, the agglomeration of Re subcarbonyl appears to be a difficult step. In the presence of Pt clusters, however, which act as nucleation sites, bimetallic PtRe clusters are easily formed. Accordingly, agglomeration of Re subcarbonyl onto existing Pt particlqs is observed at 50°C lower on Pt/NaY than on Nay (75). Use of a bimetal carbonyl cluster as precursor, such as PtRe2(C0)12,has also been studied. Because the volatility of this cluster is very low it has to be applied in THF solution. The catalytic signature of PtRe/NaY is a high selectivity for deep hydrogenolysis of nheptane to CH4 (251). This catalytic criterion shows that the superior formation of PtRe clusters by the ship-in-a-bottle technique, i.e., from Pt + Re2(CO)lo,is distinctly superior to that from the bimetal PtRe carbony1 cluster. Incorporation of CpM(C0)z into zeolite Nay (Cp = cyclopentadiene and M = Rh or Ir) has been investigated with IR and I3CO exchange (255).The complexes are stabilized by interaction with zeolite walls in the supercages of Nay. In acidic zeolites such as H8Na48Y, intrazeolite-protonated [CPM(CO)~H]+is formed (256). Protonation of Pd carbonyl clusters was also observed in Pd/NaHY during hydrogen reduction (189). By independently monitoring the IR bands of the 0-H and C-0 vibrations, the following equilibrium was found to be established: Pd,j(CO),

+ yH’

$

[Hy-Pd~,(CO)2]’’

+ (X

- z)CO,~

(21)

ZEOLITE-SUPPORED CATALYSTS

173

Reversibility is confirmed by the decrease of the 0-H band intensity on removing CO by purging with Ar, and also by the increase of the 0-H band intensity on reexposing the sample to CO.

VII.

Formation of Metal Clusters by the Ship-in-a-Bottle Method

In the above example the ship-in-a-bottle method has been used for the preparation of unligated bimetal PtRe clusters. The technique has more frequently been used for the synthesis of stable intrazeolite organometallic clusters. For example, phthalocyanine complexes of Fe, Co, Ni, Cu, and Zn ions have been synthesized encapsulated in cages of zeolite X and Y (101,257-259). These complexes are too bulky to escape through the zeolite cage windows, but complexes on the external surface of the zeolites can be removed by Soxhlet extraction with pyridine. To fit the approximately 16-A-diameter plane of the phthalocyanine macrocycle into the zeolite supercage with a cage diameter of 13 A and an aperture of 7.5 A, the planar phthalocyanine molecule undergoes a small (< 20")saddle distortion so that its peripheral benzenoid groups can protrude through the four tetrahedrally arranged windows of the supercage. A model is shown in Fig. 16. With iodosobenzene as an oxidant, zeolite-entrapped FePc displays a unique selectivity for the oxidation of alkanes. Shape selectivity is demonstrated by

FIG. 16. Saddle-deformed FePc in the supercage of zeolite Y (259).

174

.


the preferred oxidation of smaller substrate molecules. Regioselectivity is imposed on the oxidation of n-octane and methylcyclohexane; for these molecules oxidation toward the end positions is preferred. Stereoselectivity of this complex catalyst is displayed in the oxidation of methylcyclohexane: the trans C-H bond oxidation products prevail by a factor of two over the cis C-H oxidation products. Unfortunately, diffusion of the oxidation products is the rate-determining step, as evidenced by an increase of the turnover frequency with decreasing concentration of the active complex. As discussed in Section 111, it is well established that Pd13(CO), and Pd6(CO), carbonyl clusters form in zeolite cages as a result of the migration and coalescence of primary carbonyls at room temperature. In zeolites of large window diameter, such as Y, novel palladium trimethylphosphine carbony1 clusters have been synthesized inside the supercages by introducing alkylphosphine ligands (260). The bonding between the P atom of the large trimethylphosphine ligand with Pd is strong, but it is possible subsequently to replace some of the trimethylphosphine ligands by CO molecules. The resulting trimethylphosphine carbonyl clusters have an intense green color. They are stable at 200°C under a CO + H2 gas mixture. Typical FTIR spectra at various levels of CO substitution are shown in Fig. 17. The in situ preparation of Chini clusters in the basic zeolite Y has been achieved by reduction of Pt(NH3)42+with CO at 100°C (261). The intense red-purple color is attributed to the formation of Pt3[(C0)3(p~-CO)3]3~inside zeolite supercages. The IR spectrum of this cluster shows three bands at 2025 (s), 1852 (sh), and 1800 (s) cm-'. C 0 2 was detected after CO reduction. No colored complexes formed in acidic zeolites such as HY, NH4Y, or NaHY. The Pt3[(CO)3(pz-CO)3]32-cluster decomposes after evacuation at 200°C. In Pt/CsNaY, a bright red-orange color, which is apparent after evacuation at 100°C and further heating with CO at 80"C, has been assigned to the cluster Pt3[(C0)3(p2-CO)J]52-,which is supposed to extend through two adjacent supercages. A grey color, indicative of Pt particles, is observed on high-temperature treatment. The Chini clusters cannot be produced if Pt(NH3)42+in Nay or CsNaY is precalcined prior to CO adsorption. CO admission to reduced Pt particles displays only the IR bands of adsorbed CO at 1875 and 2088 cm-I. The samples remain grey. Bimetal carbonyl clusters of RhFe were synthesized inside supercages of NaY by reaction of Rh4(C0)12with [HFe3(CO)II]-/NaYin vucuo at 100°C (262). The pale red-brown RhFe carbonyl clusters in NaY catalyze the hydroformylation of ethylene and propylene with high selectivity to alcohols, whereas this selectivity is lower for supported pale grey Rh carbonyls and pink-purple Fe carbonyls. With Pd phosphine carbonyl clusters in NaY the hydroformylation of propene displays high selectivity for C7 ketones, which are formed by reaction of propene with the primary hydroformylation product, Cd aldehyde (263).

175

ZEOLITE-SUPPORTED CATALYSTS 2.65

2.37

2.09 196

950

1.81

tl

5 .e

1.53

9

1.25

$4

0.97 2042 0.69

0.41

r

0.13 2101

1

2050

1999

1948

1897

1846

1795

Wavenumber (an-')

FIG. 17. CO FTIR spectra of Pd trimethylphosphine carbonyl clusters in supercages of Nay. The PdiNaY was calcined to 500°C and reduced at 250°C before admission of trimethylphosphine. The spectra from bottom to top were recorded after CO exposure to Pro = I bar for 1, 11, 41, 71, 131, and 191 min, respectively (260).

VIII.

Metal-Proton Adducts

Bronsted acidity of zeolite protons is essential for catalytic reactions such as isomerization and cracking and has been studied extensively (15,264). Several characterization methods for acid sites in zeolites have been developed; this subject has been covered in recent reviews (265,266). Pyridine and other basic molecules are often used in IR work as probe molecules for Bronsted and Lewis acid sites (267). Trimethylphosphine has also been used as a probe for the determination of zeolite acidity by IR or NMR (96,268).

176


Pink and purple complexes were recently reported after styrene adsorption in HZSM-5; a colorimetric method for measuring the number of acid sites was suggested (269). Resonance Raman spectroscopy was also applied to determine quantitatively the acid sites in faujasite zeolites after adsorption of a dye molecule, i.e., 4-(phenylazo) diphenylamine (270). Magic angle spinning NMR has been employed for the study of HZSM-5; it was suggested that the enhanced activity for n-hexane cracking in HZSM-5 after mild hydrothermal dealumination is due to the interaction of n-hexane molecules with bridging hydroxyl groups and extraframework A1 species, rather than an enhanced Brensted acidity (271). The discussion here, however, will be focused on the interaction of transition metals with protons in zeolites. Conventionally, bifunctionality is used to describe heterogeneous catalysts that contain both transition metal sites and acid sites (43,272). Catalysis involving both metal and acid sites has often been assumed to proceed via shuttling of reaction intermediates between metal and acid sites (41).

The term “electron deficiency” was introduced by Dalla Betta and Boudart to account for the anomalously high hydrogenation activity of small Pt particles in zeolite Y (50). The electron deficiency was ascribed to an electron transfer from small Pt particles to the zeolite. X-Ray absorption has been applied to measure the Pt Llrlwhite line area as an indication of the electron deficiency because the white line is related to the number of unoccupied electronic states in the 5d and the 6s bands (273). For reduced Pt/NaHY it appeared that the white line area, and hence the electron deficiency of Pt particles, are closely related to the proton concentration of the zeolites. For example, the relative white line areas for Pt/H4RY.Pt/ HNY, and Pt foil are 1.6, 1.2, and 1, respectively. White line areas at the LlllX-ray absorption threshold to determine the d-band occupancy of supported metal catalysts were first reported by Lytle (274). The use of the white line area as an indication for electron deficiency has been questioned by Lewis, who argues that a decrease of the metal particle size will also lead to an increase of the white line area (275). The electronic properties of Pt particles on different supports and in zeolites of different proton concentrations were also probed with the competitive hydrogenation of toluene and benzene (276). It was found that the ratio of the adsorption coefficients of toluene and benzene bTlbe, which can be obtained from a kinetic analysis of the hydrogenation rate data, can be used as a convenient empirical index for the electronic environment of Pt particles. For Pt in zeolite Y, the ratio was found to increase with increasing acidity, as does the electron deficiency. This trend was rationalized by considering that toluene is a stronger electron donor than benzene. Further evidence for a correlation between electron deficiency of reduced Pt or Pd and proton concentration was found in FTIR studies of adsorbed

ZEOLITE-SUPPOKTED CATALYSTS

177

CO. Tri et al. reported that the stretching frequency of CO adsorbed on reduced Pt/NaHY was shifted from 2090 to 2068 cm-’ by neutralization with NaOH (277). A similar shift was found for Pd/NaHY and Pd/CaHY (152,189). The IR band for the linear mode of adsorbed CO on Pd/NaHY is at 2120 cm-I, which is considerably higher than that for CO adsorbed on Pd/SiOz. Neutralization with NH, leads, again, to a red shift. The dependence of the electron deficiency of Pd or Pt particles on the proton concentration of zeolites suggests a direct bond between metal particles and some protons (70). A model is the metal-proton adduct, e.g., [Pd,H,Im+, where m is the number of protons in the adduct. This leaves still two possibilities for the actual structure of this complex. One might assume that the protons in the adduct are either totally detached from the zeolite wall or that they act as bridges of the type Pd, H+-O,, where 0, stands for an oxygen atom in the hexagonal ring of supercages or sodalite cages. The latter model rationalizes both the unusual catalytic behavior and the experimentally observed “anchoring” of Pd particles to the zeolite wall in NaHY; (48,78,104). This evidence does not exclude, however, that some protons are detached from the wall. In Na+-rich Pd/NaHY the number of protons in the hexagonal ring and thus also in the adduct is lower than in Pd/HY and so is the stabilization of small Pd particles. For zeolites with high proton concentrations, a high metal dispersion can be maintained, e.g., at room temperature. The observation that admission of CO leads to spontaneous agglomeration of Pd suggests that protons and CO compete for the same ligand positions on the Pd clusters; as CO forms stronger bonds with the metal it is able to displace the protons from the metal particle. This cutting of the “chemical anchor” induces migration and coalescence of the cluster. Most remarkably, this process is accompanied by an increase of the intensity of the IR absorption band characteristic of the HO, vibration (152). This is additional evidence for the competition of protons and CO ligands for the same Pd sites. Also, after the formation of Pd13carbonyl clusters, it was observed that protons compete with CO for the ligand sites on the Pd cluster. The intensity of the OH band decreases when CO is removed from the carbonyl cluster by Ar purge; it increases again when CO is readsorbed on the cluster. This equilibrium is in good conformity with the adduct model (189). For Pd/NaX, it was found with ESR and XPS that reduction of Pd at low temperature results in Pd+ and small charged clusters Pdyx+ ( y > x ) (278,279). Likewise, monopositive Pd’ was observed with XPS in zeolite Y after mild reduction (180). As shown in Fig. 18, the binding energy of Pd 3dSl2in Pd/NaHY is significantly shifted to higher values with respect to the Pd reference. This shift is, in part, due to the small particle size. However, neutralization of Pd/NaHY at low temperature will leave the particle

178

WOLFGANG M. H , SACHTLER A N D ZONGCHAO ZHANG

size effect

I

electron transer

I

I

I

335.0

336.0 Pd 3d5i2 binding energy, eV

337.0

FIG. 18. Positions of Pd 3dS,*line for (A) bulk Pd, (B) Pd/NaY neutralized with NH3, (C) Pd/NaY neutralized with NaOH, (D) Pd/NaY reduced at 350°C (E) Pd/MgY reduced at 350"C, and (F) Pd/HY reduced at 350°C; width of the bars corresponds to error bar (180).

size constant but lower the binding energy. The difference in binding energies between Pd/NaHY and Pd/NaY is therefore ascribed to the difference between Na' and H+ as ligands to the Pd cluster, with H forming the stronger bond and making the cluster more electron deficient. In accordance with this interpretation, the binding energy on Pd/HY is shifted to an even higher value than in Pd/NaHY. In these samples the electron deficiency appears clearly related to the proton concentration of the supporting zeolite. The results indicate a direct interaction of Pd with protons, in accordance with the metal-proton adduct model. It can be asked whether these shifts are due to a simple equilibrium of the type given by Eq. (22):

If that were the case it would be expected that the binding energy (BE) for Pd, and for [Pd,H]+ have discrete values for a given value of n . In this case two XPS lines would coexist for samples containing both species; these

ZEOLITE-SUPPOmED CATALYSTS

179

lines might merge to a broad line. Remarkably, this is not observed experimentally, but the line width remains constant while the BE shifts from the value of the proton-free to the proton-rich zeolite (48,180). This result indicates that each sample is fairly homogeneous, but adducts [Pd,H,Im+ with different values of m are present in zeolites with different proton concentrations. This conclusion is in line with the geometric concept that a Pd particle of given size inside a zeolite cage will interact with several chargecompensating ions. In HY all these ions are protons, in NaHY some are sodium ions, and in Nay all are sodium ions. The interaction of metal clusters in zeolites with protons has been studied by isotope exchange with D2 (280-283). The presence of metals or metal impurities is essential for this exchange to occur at moderate temperatures; absolutely no exchange between DZand zeolite protons in metal-free HY is detected at room temperature. In the presence of Pt or Pd, however, exchange is fast and includes all protons in the zeolite, which are detectable by their O-H vibration bands in FTIR. Conceptually several stages can be discerned for the exchange with gaseous D2: (1) exchange of deuterium, which is dissociatively adsorbed on the metal, with adsorbed H atoms and with protons in the adduct, (2) exchange between deuterium in the adduct with protons bonded to 0, in the same cage, and (3) further diffusional exchange with more remote protons. Although the first process can be monitored with MS analysis of the gas or by 'H NMR of the solid, the other two stages can be monitored by FTIR. This technique permits also to distinguish protons in supercages from protons in sodalite cages of Y zeolites, as these protons have different positions of the H-0, vibration band. Figure 19 shows how these bands disappear with time, because the H-0, vibration is replaced by the D-0, vibration. It has been found that protons in supercages exchange more rapidly than do those in sodalite cages (284). Figure 19 shows the intensity decrease of the supercage (3650 cm-') and the sodalite (3550 cm-') OH bands and the intensity increase of the supercage (2690 cm-') and sodalite (2630 cm-') optical density (OD) bands. In accordance with earlier findings by Baumgarten et al. (285), with Pt on amorphous supports the exchange rate increases with increasing proton concentration because associated OH groups undergo faster exchange than do isolated OH groups. Therefore, the exchange rate is faster for Pd/HY than for Pd/NaHY (see Fig. 20). It is of relevance that very little isotope exchange is detected, in Fig. 20, for a physical mixture of HY with a neutralized Pd/NaY. As the physical mixture was rereduced after mixing, some solid-state ion exchange between Pd/NaY and HY will have taken place; the Pd-proton adducts formed in this stage will account for the observed low rate of exchange.

.. ..

WOLFGANG M M HH SACHTLER SACHTLER AND AND ZONGCHAO ZONGCHAO ZHANG ZHANG WOLFGANG

n

650cm-' i3650cm-'

I

aa

2690crn-'

b

I( \ 50 i0

I2630cm-'

~--

==4L.-7__-2 50 7 r 2682 2614 2546 2682 2614 Wavenumber (cm-1) (cm-1)

_ _ ( P 1 . -

''

3450 3600 ' 3450 (cm.1) Wavenumber (cm.1)

'

3300

FIG.19. FTIR spectra showing 0-H (a) and 0-D (b) bands in in reduced Pd/NaHY after FIG. exposure to DZfor 0, 2 , 9, 30, 60, and 240 min (284). exposure 1.00

+ A ++ A

0.80

; 1

8

493 K) overestimate the rate of these secondary reactions during FT synthesis. We will also describe how a second catalytic function within transport-limited pellets modifies product molecular weight to a degree not possible by structural and chemical modifications of FT synthesis pellets and chain growth sites. These secondary reactions include hydrogenation, oligomerization, cracking, and carbonylation of olefins, and the hydrogenolysis of linear (C,) paraflIns

(C,,,,C,,-,,,)

paraffins

+ - kc

- - - - (C,)

olefins

I

:k.

- k-.- -

(Cn)

paraffins

v

(Cn+l OH) alcohols

secondary reactions

FIG. 1 . Chain growth and termination and secondary reactions in Fischer-Tropsch synthesis on Co and Ru catalysts.

FISCHER-TROPSCH SYNTHESIS

225

paraffins. These examples will show our increasing ability to tailor catalytic pellets with controlled olefin, CO, and HZconcentrations, thereby introducing key design variables for the control of product distributions formed via polymerization kinetics and secondary reactions during FT synthesis. C. INITIATION OF SURFACE CHAINSBY READSORBED OLEFINS

Secondary reactions that merely reverse one of the termination steps in the overall chain growth sequence do not require a second catalytic function. For example, a-olefins can readsorb and reform the growing chain from which they desorbed (Fig. 1). These reverse reactions occur when termination steps approach equilibrium. a-Olefins, and to a much lesser extent internal olefins, readsorb onto chain growth sites and initiate surface chains that continue to grow and ultimately desorb again as larger hydrocarbons. These readsorption steps are repeated until olefins are removed by diffusion and convection processes from the neighborhood of chain growth sites or until a surface chain terminates as an unreactive paraffin. Termination steps that are far from equilibrium, such as hydrogen addition reactions that lead to n-paraffins, are not easily reversed during FT synthesis because the rate of the reverse reaction is negligible at reaction conditions. Many studies have shown that primary a-olefin products undergo secondary reactions during FT synthesis (15-24). Smith et al. (15) and Craxford (16) observed that the addition of ethylene to CO-Hz reactants increases the formation of oil during Fischer-Tropsch synthesis on Co catalysts. Herington (17 ) first suggested that surface chains terminate by desorption as olefins and paraffins and that olefins can reenter the surface chain growth pathways by reinitiating growing surface chains. Earlier, the presence of olefins in C O / H z feeds was shown to increase product molecular weight in CO hydrogenation (25). More recent work shows that linear (19-22) and cyclic (26,27) olefins act primarily as chain initiators in secondary reactions. Olefins can act as monomers and add as a unit to existing growing chains in oligomerization-type reactions. Readsorbed olefins can also undergo depolymerization by single scissions of terminal methyl groups, a process that merely reverses the methylene addition chain growth step (28, 29). Chain initiation reactions of a-olefins become increasingly selective with increasing CO partial pressures. It appears that high coverages of chemisorbed CO favor the selective reattachment of olefins to chain growth surface sites over the many other secondary reaction pathways available to them. These studies (17-28) also concluded that secondary hydrogenation of a-olefins often competes effectively with readsorption and chain initia-

226

ENRIQUE IGLESIA et

al.

tion. These reactions render olefins unreactive for subsequent readsorption and chain growth; they are also inhibited by high CO partial pressures and become very slow at normal FT synthesis conditions. Recently, we have shown that secondary hydrogenation of a-olefins is also strongly inhibited by water, an indigenous product of the FT synthesis reaction (4,30).As a result, the rate of secondary hydrogenation becomes negligibly small at reasonable CO conversion levels (>5%) and high pressures (>500 kPa) on Ru and Co catalysts (4,30).CO also inhibits olefin hydrogenation reactions, which become very slow as reactant pressure increases. However, many secondary reactions become important again as CO is depleted by strong diffusional restrictions within pellets or by high conversions within the catalyst bed.

D. CHAIN GROWTH AND TERMINATION KINETICS FT synthesis product distributions can be described by equations first developed for polymerization kinetics (31). Flory product distributions (31) an-'

F, = n (1 - a)' (1) where F, is the fraction of the carbon atoms within chains containing n carbon atoms, are obtained when the chain growth probability ( a )is independent of chain size. These product distributions yield a straight line when plotted as In(F,,/n) versus n. This analogy between polymerization and FT synthesis products was first described by Anderson (30, 32), who suggested the relation, Eq. (2), +n = Oi. a"-' (2) which follows directly from Eq. (1) and which relates FT synthesis product distributions to those formed in conventional polymerization processes. An alternate approach that we use here describes FT products in terms of individual chain termination probabilities for each chain size, Pn,

as suggested previously by Herington (17). This approach allows chain growth kinetics to vary with chain size. The total termination probability, &,n becomes a linear combination of the values for the individual termination steps in Fig. l , PT,n

=

Po..

+ PH,n - P r , n

(4)

where the individual termination probabilities are obtained by substituting the mole fraction of each hydrocarbon species in the numerator of Eq. (3).

227


Readsorption steps @, decrease the total termination probability (&,) by reversing the olefin termination step. Readsorption can be enhanced by increasing the olefin reactivity (i.e., &, the ratio of readsorption to propagation rate constants), or by increasing olefin concentrations within pellets and reactors. The net growth probability of chains of size n (a,) is obtained from Eq. (3): a n

= 1/(1

+ Pn>

=

C

+ i / i i=n

i=n+ I

4i

(5)

Clearly, the total chain growth probability is not a linear combination of the chain growth probabilities for each hydrocarbon type within a carbon number. Semilogarithmic plots of Fn/n or 4nversus n cannot be used for individual components (olefins, paraffins, alcohols) within a given carbon number; they are valid only for the entire product within a carbon number and give only the overall chain growth probability ( a ) ,not the individual chain growth probabilities for each hydrocarbon type. The frequent use of such plots to describe products of individual termination steps or to determine individual chain growth probabilities for olefins or paraffins is clearly incorrect, unless each product arises from distinct and independent chain growth sites. E.

PRODUCT FUNCTIONALITY AND MOLECULAR WEIGHT

DISTRIBUTIONS

FT synthesis products seldom follow Flory distributions throughout the entire molecular weight range (4,5,14,18,21,26,33-48).In most cases, chain growth probability increases with chain size, leading to curvature in Flory carbon number distribution plots of FT synthesis products. Often this curvature is slight and occurs over a wide range of hydrocarbon size; consequently, it is difficult to detect deviations from Flory kinetics when measurements are limited to a narrow molecular weight range. Descriptions of product distributions using Eq. (3) and analysis of the entire product slate are often required to detect the carbon number dependence of chain growth and termination pathways (4,5,40,41,44). Pichler et al. (18) first reported deviations from Flory distributions on Co catalysts and suggested that they reflected the higher intrinsic readsorption reactivity of larger olefins, which led to surface chains that continued to reform until they terminated irreversibly to unreactive paraffins. Schulz et d. (42) noted that products begin to leave the reactor with the liquid phase in the same hydrocarbon size range where deviation from Flory kinetics occurred. Later, these authors proposed that the higher solubility of larger

228

ENRIQUE IGLESIA et

al.

olefins in the FT synthesis liquids increased their local concentration and reactor residence time, leading to their apparent higher reactivity (33,34, 42,43). Recently, we have shown that non-Flory distributions cannot arise from the higher solubility of larger olefins because thermodynamic equilibrium between the two phases requires that the fugacity, chemical potential, and kinetic driving force for each component be the same in the two phases (4,5,14,40,41,44).Transport restrictions, however, can lead to higher intrapellet concentrations and residence times of a-olefins, a feature of FT chemistry that accounts for the non-Flory distribution of reaction products and for the increasing paraffin content of larger hydrocarbons (4,5,14,40, 41,44). Non-Flory molecular weight distributions have also been attributed to the presence of several types of active sites with different probabilities for chain growth and for chain termination to olefins and paraffins (45). Two-site models have been used to explain the sharp changes in chain growth probability that occur for intermediate-size hydrocarbons on Fe-based catalysts (46,47). Many of these reports of non-Flory distributions may instead reflect ineffective dispersal of alkali promoters on Fe catalysts or inadequate mass balances and product collection protocols. Recently, we have shown that multisite models alone cannot explain the selectivity changes that occur with increasing chain size, bed residence time, and site density on Ru and Co catalysts (4,5,40,44). We have shown previously that non-Flory distributions often reflect the transport-limited removal of reactive olefins from catalyst pellets on Ru and Co catalysts (4,5,14,40,41,44).This proposal is consistent with the similar effects of bed residence time and of molecular size on chain growth probability and product functionality. It accounts for the observed effects of convective and diffusive rates of reactive olefins and for the non-Flory distribution of highly paraffinic hydrocarbons formed from synthesis gas on Co and Ru catalysts.

EFFECTSIN FISCHER-TROPSCH SYNTHESIS F. TRANSPORT Reactants and products must diffuse through high-molecular-weight liquid hydrocarbons during FT synthesis. The liquid phase may be confined to the mesoporous structure within catalyst pellets or extend to the outer surface and the interstitial spaces between pellets, depending on the reactor design and hydrodynamic properties. In packed-bed reactors, the characteristic diffusion distance equals the radius of the pellets plus the thickness of any liquid boundary layer surrounding them. Intrapellet diffusion usually becomes


229

the controlling resistance in packed-bed reactors. The thickness of the interpellet liquid layer increases dramatically as the reactor configuration changes from packed-bed to trickled-bed and bubble column reactors and diffusion through extrapellet liquids and across the gas-liquid interface becomes the controlling transport resistance. Diffusion of molecules through liquids is much slower than gas-phase transport (49); as a result, primary and secondary reactions will often occur at reactant and product concentrations that differ significantly from those in the interpellet regions within the catalyst bed. Reactants penetrate catalyst pellets slowly and cannot satisfy the kinetic rate of chemical reactions; this lowers the catalyst effectiveness and frequently changes product selectivity. When more than one reactant is involved, their different diffusion rates can lead to marked differences in their relative fugacity within the intrapellet liquid phase. Slow removal of reactive products, on the other hand, can also modify selectivity by enhancing the rate of their secondary reactions. When products are much larger and diffuse more slowly than reactants, as in the case of FT synthesis, intrapellet product fugacity gradients often arise even within small catalyst pellets. The well-recognized strong influence of intrapellet diffusion on FT synthesis rate and selectivity has led to many reports describing the benefits of eggshell catalysts, which contain most of the active sites near the outer pellet surface, in order to control the severity of diffusional restrictions and improve FT synthesis rate and selectivity (50-55). The effects of diffusional restrictions on the activity and selectivity of FT synthesis processes have been widely studied (32,52,56-60). Intrapellet diffusion limitations are common in packed-bed reactors because heat transfer and pressure-drop considerations require the use of relatively large particles. Bubble columns typically use much smaller pellets, and FT synthesis rates and selectivity are more likely to be influenced by the rate of mass transfer across the gas-liquid interface as a gas bubble traverses the reactor (59,61,62). Many reports have shown that larger catalyst pellets lead to lower FT synthesis rates and to lighter products, particularly methane (32,52,56,57-63). Significant effort has been devoted to the measurement of intrinsic kinetics in diffusion-free pellets under differential reactor conditions (32,5236, 57,5Y-65). These kinetic expressions are then combined with descriptions of diffusion through the liquid-filled catalyst pores to produce models that attempt to predict the effect of diffusional inhibition on synthesis rates. Most of these studies have concentrated on reactant diffusion effects and make use of standard effectiveness factor and Thiele modulus treatments (52,56,5Y ,65). These models describe many of the observed experimental trends, but several issues related to the processes that they describe remain

230

ENRIQUE IGLESIA et

al.

unclear. For example, H2 is assumed to be the limiting reactant and FT synthesis kinetics are often described as first order in H2 (56,60). These assumptions are not consistent with available experimental studies, which show that CO is the diffusion-limited reactant under normal FT synthesis conditions on both Co and Ru catalysts. Also, experimental FT synthesis kinetics are more properly described by Langmuir-Hinshelwood expressions, characterized by fractional power dependencies on both CO and H2 pressures (64-67) rather than by first-order kinetics. More recent studies include more complex kinetics and describe a broader range of experimental observations, including, in some cases, predictions of FT synthesis selectivity (4,5,52,64). Two important types of diffusional limitations occur in Fischer-Tropsch synthesis: (1) slow removal of reactive products from pellets and reactors and (2) slow arrival of reactants to catalytic sites. The first regime prevails even within relatively small pellets (-0.1 mm) because reactive a-olefins diffuse slowly through high-molecular-weight liquid hydrocarbons in the catalyst pores. This regime is characterized predominantly by changes in product selectivity; FT synthesis rates equal their kinetic values at the prevailing composition in the gas phase because the faster diffusion of CO and HZmolecules still allows uniform intrapellet fugacities of these reactants. The second regime appears as pellet size or active site density continue to increase because intrinsic reactant consumption rates cannot be satisfied. Consequently, CO and H2 intrapellet fugacities drop significantly. These changes in reactant concentration along the pellet radius also influence product selectivity. Available reaction-transport models describe the second regime (reactant transport), which only requires material balances for CO and HZ. Recently, we reported preliminary results on a transport-reaction model of hydrocarbon synthesis selectivity that describes intraparticle (diffusion) and interparticle (convection) transport processes (4, 5 ) . The model clearly demonstrates how diffusive and convective restrictions dramatically affect the rate of primary and secondary reactions during Fischer-Tropsch synthesis. Here, we use an extended version of this model to illustrate its use in the design of catalyst pellets for the synthesis of various desired products and for the tailoring of product functionality and molecular weight distribution.

111.

Experimental Methods

A. CATALYST SYNTHESIS AND CHARACTERIZATION Detailed synthesis procedures for supported Ru (4) and Co (5, 41) catalysts have been reported previously. These catalysts were prepared by incip-


23 1

ient wetness impregnation of Ti02 (P25, Degussa, 60% rutile), S O 2 (Davison, Grade 62; Shell silica spheres), and y-A1203 (Catapal SB, W. R. Grace Co.) with Ru nitrate (Engelhard Corp.)/acetone mixtures or aqueous solutions of cobalt nitrate (Johnson-Mathey, Puratronic grade). The supports were evenly impregnated with the metal salt solutions, except during the preparation of partially impregnated eggshell catalysts described below. Ru powders were obtained from Johnson-Mathey (Puratronic grade). Samples were sieved to retain desired pellet sizes (80-140 mesh, 0.17 mm average diameter, unless otherwise noted). Cobalt catalysts were calcined in air (373-773 K , 2-4 h) before reduction. All samples were reduced at 723 K for 4 h, passivated with dilute (1%) oxygen, and reduced again at 673 K for 1-3 h before chemisorption and catalytic tests. Eggshell catalysts, in which the metal is exclusively impregnated near the outer surface of support pellets, were prepared by contacting the various supports with Co nitrate melts at 348-363 K for short periods of time (2-5 s) (53,68).These samples were directly reduced in flowing hydrogen without intermediate calcination in order to maximize cobalt dispersion at the high local metal concentrations within the eggshell region. The eggshell thickness was determined by optical and scanning electron microscopy. Metal dispersion in Ru catalysts was determined by titration of preadsorbed oxygen with dihydrogen at 373 K (4,4I,69,70),which provides an accurate measurement of exposed Ru atoms even on TiOz-supported crystallites exhibiting strong metal-support interactions (SMSI) (71). Cobalt dispersions were determined by chemisorption of dihydrogen at 373 K, assuming a 1 : 1 H : Co surface stoichiometry (72). In order to avoid SMSI effects that strongly inhibit hydrogen and CO chemisorption (71), prereduced Co/ Ti02 catalysts were calcined at 573 K for 0.5 h and then rereduced at 523 K in H2 before chemisorption experiments. Chemisorption, X-ray diffraction line breadths, and electron microscopy studies led to similar values for the average metal crystallite size in all samples (4,41). Metal dispersions were used to calculate site densities and FT site-time yields on Co and Ru catalysts. Surface areas and pore-size distributions were measured using dinitrogen physisorption and capillary condensation methods (73). The extent of reduction of Co oxide precursors was determined by temperatureprogrammed reduction in flowing hydrogen and by oxygen uptake measurements at 673 K after reduction, assuming the formation of Co304during oxidation (74).

B. CATALYTIC MEASUREMENTS Catalytic tests were performed in isothermal fixed-bed reactors after samples were rereduced in situ at 673 K for 1-3 h. Steady-state FT synthe-

232

ENRIQUE IGLESIA et

al.

sis rates and selectivities were obtained after 24 h onstream. Reactants and products were analyzed directly (CI-Cls, CO, COZ, N2) or after product collection (CI6+)by gas and liquid chromatography (4,5,41).Carbon selectivities are reported as the percentage of converted CO that appears as a given product. FT synthesis rates are reported as metal-time yields (moles CO converted/gram-atom metal * second) and site-time yields (moles CO converted/gram-atom surface metal * second). Space velocity is defined as the feed flow rate per total (not void) reactor volume. Bed residence time is reported as inverse space velocity. Catalytic data were obtained over a wide range of CO conversion by varying the space velocity at constant temperature and reactant pressure. In many cases, we choose to report rate and selectivity data at integral reactor conditions (45-65% CO conversion) because they favor the synthesis of Cs+ hydrocarbons. All comparisons among catalysts are made at similar levels of conversion, a requirement imposed by the integral operation of the catalytic reactor. Our conclusions, however, remain valid when similar comparisons are made at much lower conversion levels in differential reactors.

c.

KINETIC AND REACTION-TRANSPORT MODELS

I . Surface Chain Growth and Termination The overall stoichiometry of the Fischer-Tropsch reaction is given by n(C0

+ 2H2)

-

-(CH2).-+

nH2O

(6)

for the synthesis of high-molecular-weight hydrocarbons. This reaction describes the step-wise incorporation of CH2 species into a growing chain and the removal of oxygen atoms as water. This chemical equation represents a sequence of elementary steps leading to the formation of hydrocarbons from CO and H2. We use it here to describe global kinetic expressions and mass balances. The hydrogenation of CO and the formation of each individual C, hydrocarbon (e.g., methane) obey Langmuir-Hinshelwood kinetic expressions (32,64,66,67,75):

Table I shows rate constants (L)and pressure orders ( a , b) for CO conversion and methane formation from rate and selectivity data obtained on Co and Ru catalysts over a wide range of pressure (100-2100 kPa) and H2/C0

233


TABLE I CO Hydrogenation and CH, Formation Kinetics on Co and Ru Catalysts

.(

= ki[W'tCOlb) i

1

+K[CO] Catalyst

co ParameteP

co

k, a b

1.96 x 10-7 0.60 0.65

Ru CH4 1.08 X 1 .o 0.05

co 7.5 x lo-'' 1.o 0.6

CH4 4.0

X

lo-''

1.5 0.0

Units: required to give rates in (mole/g-atom surface metal * s) with reaction orders a and b. Obtained on Si02-supported R y and Co catalysts at 473-483 K , 100-3000 kPa, Pa-', estimated to give shift to posiHz/CO = 1-10, l ) , the value of yo determines local H K O ratios and the selectivity of the reactions given by Eqs. (17) and (18). For example, for stoichiometric pellet boundary conditions [HdCO = 2, based on Eq. (17)], values of YO > 1 cause the concentration gradients of CO within pellets to be greater than those of H2. This leads to higher H K O ratios within the catalyst particles than at the pellet surface and to an increase in methane selectivity (methane production/CO consumption). At typical hydrocarbon synthesis conditions (2OO0C, 2100 kPa), yo is about 1.9; consequently, methane selectivity increases as pellets become diffusion-limited in stoichiometric feeds (HdCO = 2.1). The parameter YO depends weakly on temperature; it varies from 1.8 to 2.0 for temperatures between 150 and 250°C (52,53). As in the olefin readsorption model, the Thiele modulus for reactant diffsion (a;) can also be expressed as the product of two components: = +co

. (RZ @ @ M / ~ J

(25)

The second term is identical to the parameter y, described in Eq. (15), where it accounts for the effect of catalyst structural properties on olefin readsorption rates. Not surprisingly, structural catalyst properties that restrict the removal of a-olefins from catalyst pellets also limit the arrival of reactants at catalytic sites. The +co term reflects the diffusivity and reactivity properties of reactant molecules.

IV. Results and Discussion ON A. METALCRYSTALLITE SIZE AND SUPPOR EFFECTS FISCHER-TROPSCH SYNTHESIS RATESON RUCATALYSTS

Specific CO hydrogenation rates (Ru-time yields) are proportional to Ru dispersion (Fig. 2a). Site-time yield values are similar (1.25-1.95 X s-') on all Ru catalysts, including unsupported powders (Table 11). Support and dispersion effects are minor for the supports (SiOz, A1103, TiOz) and dispersion range (0.09-50%) used in our experiments. Site-time yields also do not depend on any strong metal-support interactions (71) introduced by the use of titania as a support for Ru crystallites, perhaps because such SMSI effects are destroyed by contact with water at FT synthesis conditions. Thus, the rate of CO hydrogenation remains proportional to the number of exposed Ru surface atoms, regardless of the size of the crystallite and of the chemical identity of the metal oxide support at least for Ru crystallites larger than about 1.6 nm. Carbon number distributions are qualitatively similar on all Ru catalysts used in our study. Chain termination probabilities (Pn) for small chains

238

ENRIQUE IGLESIA et 0.55

--

Ruthenium-time yield (104s-1) 120

a

t

)

:

0.5

2

0.45

0

.-

0.25

n

e

-

.-0 .-C I

al.

0 0.2

10.6% Ru/Si02

(a,= 0.938, p- = 0.066) 1.2% Ru/Ti02 a, = 0.939, !33 = 0.065)

0.15

e

/

c

'5

L:

0.1

0

-m

c I-0

v-

0.05

n 0

0.2

0.4

Ru Dispersion

0.6

0.8

0

5

10 15 20 Carbon Number (n)

25

30

FIG. 2 . (a) Effects of ruthenium dispersion and support on Fischer-Tropsch synthesis rates (476 K , HdCO = 2.1, 560 kPa, 45-60% CO conversion). A, TiOz; 0 , SiOz; D, AIzO,;o, powder. (b) Support effects on chain termination probability on ruthenium catalysts (476 K, HJCO = 2.1, 560 kPa, 45-60% Co conversion).

( n < 20) depend on the size of Ru crystallites and on the chemical identity of the support (Fig. 2b, Table 11). Previously, we have shown that these changes in Cs+ selectivity and chain termination probability actually reflect differences in the density of Ru sites and in the physical structure of the support pellets, which lead to transport-enhanced readsorption of a-olefins (4,41). Ultimately, termination probabilities for longer chains reach asymptotic values that become independent of support and of metal dispersion (Fig. 2b). Chain termination probabilities initially decrease with increasing chain size (Fig. 2b); product distributions are non-Flory on all catalysts. This reflects an increase in readsorption rate as larger a-olefins become increasingly difficult to remove from liquid-filled catalyst pellets (4,5,14,40,41,44). Large olefins readsorb extensively and leave catalyst pellets predominantly after they form n-paraffins in sequential chain initiation and termination steps. As larger olefins (n > 30) disappear from the products, the chain termination probability reaches a constant value and product distributions become predominantly paraffinic and obey Flory kinetics (Fig. 2b). The asymptotic termination probability (&) reflects the intrinsic probability of

TABLE I1 Elemental Composition. Dispersion, and Catalytic Properties of Ru Catalysts

Support Ti02 (30% anatase) Ti02 (60% anatase) Ti02 (100% anatase) Si02 SiO, Y -A1203 Ru powder

Ru content (wt%)

Ru dispersion"

Ru-time yield (10' s - ' ) ~

Carbon selectivity (%) Site-time yield (lo' s - ~ ) ~ C X 6 C*+b

1.2

0.48

6.8

13.0

3.7

87.3

0.9

0.60

7.5

12.5

5.0

84.0

4.8

0.26

4.2

16.1

3.0

89.0

0.22 0.082 0.25 0.0009

2.5 1.3 3.5 0.018

11.4 15.8 14.0 19.5

4.2 8.1 5.0 2.0

89.1 71.6 88.0 95.0

10.6 1.8 5.1 100

" From hydrogen titration of chemisorbed oxygen. Obtained at 476 K , 560 kPa, H2/C0 = 2.1, 4 5 4 0 % CO conversion; 0.17 mm average pellet diameter.

240

ENRIQUE IGLESIA

et at.

chain termination to paraffins by hydrogen addition to growing alkyl chains, an intrinsic property of Ru surfaces unaffected by crystallite size or support effects (Fig. 2b). As we discuss below, differences in Cs+ selectivity among these Ru catalysts do not require different surface chain growth kinetics; they can be described accurately by olefin readsorption reactions enhanced by varying levels of intrapellet transport restrictions. The structure insensitivity of CO hydrogenation reactions on Ru was previously suggested by Dalla-Betta and co-workers (78). Initial CO hydrogenation turnover rates were independent of Ru crystallite size (1.O-9.0 nm in diameter). Other reports suggested that turnover rates increase with increasing crystallite size over a similar Ru dispersion range (79,80). Methanation turnover rates on Ru decreased from 0.16 to 0.01 s-' as the Ru dispersion increased from nearly zero (Ru powder) to 64% (Ru/A1203)(79). In another study, turnover rates were found to decrease markedly in a Ru dispersion range (0 to 25%) where changes in the surface density of lowcoordination surface atoms are minor and where crystallite size is unlikely to influence surface structure and chemical reactions (83). Kellner and Bell (80) have also reported a decrease in turnover rate with increasing metal dispersion in Ru/AI2O3catalysts at both 0.1 and 1 MPa reactant pressures (H2/C0 = 3, 498 K). Methane turnover rates at 498 K and 1 MPa decreased from 0.003 to 0.0004 s-' as the Ru dispersion increased from 25 to 80% (80). Smith and Everson (84) also reported a 10-fold decrease in turnover rate as the dispersion of 0.5% Ru/A1203eggshell catalysts increased from 16 to 78%. Cs+ selectivity and chain growth probability, however, were not strongly affected by Ru dispersion. These reports suggest that the hydrogenation of CO is sensitive to structural and chemical effects induced by changes in Ru crystallite size or by electronic perturbations induced by their interactions with the metal oxide support. Several other studies and the results that we report here, however, suggest that CO hydrogenation reactions on Ru are structure insensitive (85). For example, CO hydrogenation turnover rates were very similar on Ru(001) and Ru(ll0) single crystals and resemble those on supported Ru catalysts (86),suggesting that methanation pathways and surface kinetics are not strongly influenced by the structural details of the catalytic surface. Vannice and Garten (87) reported that methanation turnover rates on Ru were independent of support identity, but methane selectivity was much lower on Ti02 than on other metal oxide supports. Kikuchi et af. (88) reported that CO hydrogenation turnover rates decreased by less than a factor of two when Ah03 replaced TiO2 as the support for Ru crystallites (7-19% Ru dispersion, 523 K , 100 kPa, H2/C0 = 2), while chain growth probabilities remained unchanged. In contrast, Stoop et al. (89) observed marked rate and selectivity differences among supports; they concluded, however, that these


24 I

changes did not reflect support modifications of intrinsic Ru surface chemistry but instead the effect of C1 anions and other support impurities on the density of active sites and on intrinsic chain growth kinetics (89). Some reported alloy effects are consistent with structure-insensitive pathways requiring small ensembles of Ru surface atoms. For example, alloying Ru crystallites with inactive Cu atoms led only to a modest initial decrease in turnover rate (86). CO hydrogenation turnover frequencies decreased slightly as Au atoms were introduced into Ru catalysts, but activation energies were not influenced by the resulting decrease in ensemble size (90). Similarly, turnover rates remained constant as the fraction of surface Cu on a Ru(001) single crystal increased, also suggesting a small ensemble requirement, probably no larger than isolated Ru atoms (91). These studies suggest that the rate-limiting step in methanation reactions can occur on isolated Ru atoms, albeit with the required assistance of H adatoms that arrive from neighboring Cu atoms or by migration from more distant Ru ensembles. Earlier studies of alloy effects on CO hydrogenation (92, 93), however, concluded that CO hydrogenation steps required much larger Ru ensembles (4-13 Ru atoms), a contradictory observation that suggests a strong sensitivity to surface structure. In catalytic rate measurements, we probe the structural requirements of intermediates and of chemical reactions involved in the rate-limiting steps of a catalytic sequence. Thus, many of the contradicting reports of dispersion and support effects may reflect different operating conditions, which can alter the identity and the structural requirements of the rate-limiting step. Previously reported dispersion and support effects on Ru catalysts were obtained at conditions that favor the formation of light products, especially methane. In our study, we have extended this work to FT synthesis conditions (>500 kPa, 473 K), where Cs+ selectivities exceed 80%. We conclude that the synthesis of higher molecular weight hydrocarbons on Ru is structure insensitive according to the definition given in Refs. 85 and 94, at least on Ru crystallites larger than 1.6 nm. Site-time yields and chain growth kinetics depend only weakly on Ru dispersion and on the identity of the metal oxide support. The structure insensitivity of a reaction as complex as the FT synthesis is somewhat surprising given the number of steps required to build a large hydrocarbon chain from CO and HZ.This structure insensitivity could be an intrinsic property of chain growth pathways on metal surfaces and could reflect rate-limiting steps that are not strongly influenced by local surface structure or by the size of available metal ensembles. Indeed, hydrogenation reactions are generally structure insensitive (85) and synthesis rates appear to be limited by the hydrogenation of surface carbon to form CH, monomers f 7,8,9,26).

ENRIQUE IGLESIA et

242

af.

More probably, the structure insensitivity arises because chain growth occurs on surfaces covered almost completely by CO reactive intermediates, a mechanistic detail reflected in the negative CO pressure order dependence of the Fischer-Tropsch synthesis. As a result, chain growth occurs only on a few surface sites ( 9 3 , generally those that bind CO least strongly in structurally nonuniform but well-covered metal surfaces. Well-covered surfaces hide many of the structural features and specific binding sites initially present on the clean metal surface. In general, catalytic reactions occurring on well-covered surfaces become structure insensitive for similar reasons (96). Similar arguments would suggest that the Fischer-Tropsch synthesis on fully reduced Co surfaces at normal CO partial pressures will also be insensitive to dispersion and support effects. The results in the next section confirm this suggestion.

B. METALCRYSTALLITE SIZE AND SUPPORT EFFECTSON FISCHER-TROPSCH SYNTHESIS RATES-COBALT CATALYSTS Metal-time yields on Co catalysts are also proportional to the dispersion of Co crystallites (10-100 nm in diameter, 0.45-9.5% dispersion) on supported SiOz, A1203, TiOz, and mixed-metal oxides (Fig. 3). Cobalt-time yields increase with increasing dihydrogen uptake on all metal oxide supports. Therefore, site-time yields lie in a narrow range (1.6-3.0 X s-', at 2000 kPa and 473 K ) and do not depend on support or metal disperCobalt-time yield (104s-1)

30

I

25 20 15 10 5

-0

n

0.02

0.04 0.06 0.08 Co Dispersion

0.1

0.12

FIG.3. Effects of cobalt dispersion and support on Fischer-Tropsch synthesis rates (473 K , Hz/CO = 2.1, 2000 kPa, SO-60% CO conversion). A TiOz; 0 , SiOz; H, AIz03; 0, other.


243

sion (Table 111). In this dispersion range (0.45-9.5%), site activity is not strongly affected by metal-support interactions or by crystallite size. Carbon number distributions are similar on all Co catalysts. As on Ru catalysts, termination probabilities decrease with increasing chain size, leading to non-Flory product distributions. The modest effects of support and dispersion on product molecular weight and cs+selectivity (Table 111) reflect differences in readsorption site density and in support pore structure (4,5,14,40,41), which control the contributions of olefin readsorption to chain growth. Carbon number distributions obey Flory kinetics for C30+hyas drocarbons; the chain growth probability reaches a constant value (a"") olefins disappear from the product stream. This constant value reflects the intrinsic probability of chain termination to paraffins by hydrogen addition; it is independent of support and metal dispersion in the crystallite size range studied. Our site-time yields were measured on catalyst granules smaller than 0.2 mm in diameter in order to ensure that CO hydrogenation rates were unaffected by diffusion-limited CO and H2 arrival at catalytic sites. Product removal is still slow on such small pellets, but it affects only the selectivity and not the rate of CO hydrogenation reactions. Our site-time yields were measured on catalysts with more than 95% of the Co atoms in a zero-valent state, in order to avoid complicating factors associated with partially reduced Co surfaces. The initial studies of the structure sensitivity of the FT synthesis on Co catalysts suggested that crystallite size and support strongly affected the rate and selectivity of this reaction (97,98). In one of these studies (97), turnover rates decreased by a factor of seven as Co dispersion decreased from 30 to 15%. These authors concluded that the hydrogenation of CO (at 100 kPa, 473-523 K , H2/C0 = 2) on Co/A1203 is very sensitive to the structure of metal crystallites, even in this narrow dispersion range, where the concentration of surface atoms with unique coordination is not influenced strongly by crystal size (83). This structure sensitivity was attributed to the need for sites that coordinate CO strongly. These authors also proposed that the density of strong binding sites and the nature of reactive CO species depend strongly on cobalt crystallite size. Similarly, they suggested that the higher molecular weight products observed on larger Co particles reflect crystallite size effects on intrinsic surface chain growth kinetics. A later study from the same group (98) also reported marked effects of support and of Co dispersion on the rate and selectivity of the FT synthesis. Turnover rates decreased in the order Ti02 > Si02 > A1203 > C , MgO and also decreased with increasing Co dispersion for each support. However, changes in dispersion and support also modified the reduction properties of Co precursors and frequently led to their incomplete reduction during

TABLE I11 TABLE 111 Elemental Composition, andDispersion, Catalytic Properties of Co Catalystsof Co Catalysts ElementalDispersion, Composition, and Catalytic Properties

Support TiOzd TiOZd TiOzd TiOzd

Ti02" TiOzd SiOz Si02 SiOz SiOz SiOz Si02 Sioz SiOz

Co content (wt%) Support

TiOzd

12.1 11.6 11.8

0.012 11.6 0.022 11.8 0.029 10.5 0.053 12.1 0.030 11.6 0.065 11.8

Si02 Si02 SiOz Si02 Si02 Si02 Sio2 Si02

24.8 23.1 14.0 15.0 13.0 15.0 10.3 32.1

24.80.042 0.032 23.1 0.019 14.0 0.050 15.0 0.063 13.0 0.060 15.0 0.095 10.3 0.0045 32.1

0.012 2.8 0.022 5.0 0.029 7.5 0.053 11.4 0.030 8.9 0.065 15.3 0.042 11.3 0.032 6.7 0.019 3.4 0.050 13.1 0.063 16.7 0.060 14.7 0.095 20.2 0.0045 1.2

0.061 17.2 0.011 10.6

0.061 10.0 0.011 3. I

10.0 16.4 3.1 27-5

5.6 16.4 8.2 27.5

84.0 5.6 83.2 8.2

84.0 83.2

0.036 19.5 0.015 11.2

0.036 0.015

6.9

6.9 19.2 2.4 16.0

7.8 19.2 8.7 16.0

82.5 7.8 8.7 80.2

82.5 80.2

TiOzd TiOZd TiOzd TiOzd TiOzd

11.6 11.8 10.5

1 1 .0% ZrOz/Si02 17.2 11.O% ZrO2/SiO2 Zro.14Tio.~~Oz Zro M T ~810.6 O6 0 ~ A1203 AlzOi

Carbon selectivity Carbon selectivity Co-time yield Co-timeSite-time Co content yield Site-time (104 S - l y yield ~ , ~ (lo3 CHSs - ' ) ~ , ' c 5+ Cobalt (104 s - I(103 ) ~ S K ' )yield CH, cs+ (wt%)dispersion" Cobalt dispersion"

A1203 AlzOs

19.5 11.2

2.4

2.8 23.1 5.0 22.7 7.5 25.9 11.4 21.5 8.9 29.6 15.3 23.5 11.3 27.0 6.7 20.8 3.4 17.5 13.1 26.1 16.7 26.5 14.7 24.5 20.2 20.7 1.2 27.8

8.1 23.1 6.8 22.7 25.9 21.5 7.0 29.6 5.4 23.5

81.5 8.1 84.5 6.8 83.0 82.5 7.0 90.1 5.4

81.5 84.5 83.0 82.5 90.I

4.7 27.0 6.3 20.8 7.5 17.5 7.0 26.1 5.8 26.5 6.4 24.5 5.3 20.7 7.4 27.8

91,O 4.7 85.5 6.3 84.0 7.5 83.5 7.0 89.6 5.8 85.2 6.4 88.5 5.3 83.0 7.4

91.O 85.5 84.0 83.5 89.6 85.2 88.5 83.0

" From Hr chemisorption at 373 Kuptake . a From H2uptake chemisorption at 373 K. K , 2000atkPa, = kPa, 2.1, Hz/CO 50-63%=CO 0.17 mm average pellet Obtained at 473Obtained 473H2/C0 K, 2000 2.1,conversion; 50-63% CO conversion; 0.17 mm diameter. average pellet diameter. ' Based on dispersion from hydrogen chemisorptlon. ' Basedvalues on dispersion values from hydrogen chemisorption. anatase, With 25-40% With 25-40% anatase, balance rutile. balance futile.

'


245

catalyst pretreatment or to partial reoxidation during catalysis; the residual presence of Co oxide species strongly influences CO hydrogenation rates (90-101). As a result, apparent dispersion and support effects may reflect the presence of chemisorbed oxygen and of bulk oxides at varying concentrations during catalysis. Similar phenomena may also account for the reported increase in turnover rate with increasing Co crystallite size (and extent of reduction) on Alz03 surfaces (102) in a dispersion range (0.09-7.6%) in which surface structure remains independent of crystal size (83). Recently, FT synthesis reactions were shown to be independent of metal dispersion on SiOz-supported catalysts with 6-22% cobalt dispersion (103). s-') over the enTurnover rates remained nearly constant (1.8-2.7 X tire dispersion range. Dispersion effects on reaction kinetics and product distributions were not reported. These tests were performed at very low reactant pressures (3 kPa CO, 9 kPa H2), conditions that prevent the formation of higher hydrocarbons and lead to methane with high selectivity and to CO hydrogenation turnover rates 10 times smaller than those obtained at normal FT synthesis conditions and reported here. These low reactant pressures also lead to kinetics that become positive order in CO pressure. Thus, the reported structure insensitivity (103) may agree only coincidentally with the similar conclusions that we reach here on the basis of our results for the synthesis of higher hydrocarbons on Co. Recent reports have confirmed that the synthesis of light hydrocarbons on Co is insensitive to surface structure (104). These authors concluded that the strong structure sensitivity reported previously (97, 98) reflects marked differences in the extent of reduction and in the ease of reoxidation of Co as changes in metal dispersion and support occur. This recent report (104) shows that turnover rates are similar on strained Co overlayers formed on tungsten single crystals of different orientations, in spite of significant differences in the cobalt surface structure. Moreover, turnover rates on these strained Co layers were very similar to those measured on Co crystallites supported on dehydroxylated A1,03, where Co oxide precursors reduced completely to cobalt metal. On these reduced catalysts, turnover rates were independent of cobalt dispersion (5 -37%). These authors concluded that methanation is a facile reaction that proceeds with similar turnover rate and reaction pathways on many transition metal surfaces (Co, Ru, Rh, Ni) of very different geometry and electronic properties (104). Similar turnover rates recently reported on Co single crystals with different exposed crystal planes are consistent with the structure-insensitive nature of CO hydrogenation on Co (105, 106). Turnover rates on crystals with zigzag grooved [Co (1 120)] and close-packed [Co(OOOl)] surfaces differ by less than a factor of two, and CO hydrogenation activation energies are also similar on the two surfaces. The grooved surface favors chain growth

246

ENRIQUE IGLESIA et

al.

(a = 0.36) compared to the close-packed surface (a = 0.20). This difference reflects the lower olefin hydrogenation rates on the grooved surface, which allow readsorption and chain initiation by a-olefins and lead to a lower net rate of chain termination and a higher value of a . Similar rates of CO hydrogenation on these two surfaces are observed even though CO chemisorbs molecularly on Co(OOO1) between 100 and 450 K but dissociates readily above room temperature on Co( 1120) (107-109). Clearly, the effects of local surface structure on CO chemisorption disappear at the high steady-state CO surface coverages that prevail during CO hydrogenation catalysis. Previously reported dispersion effects on Co catalysts were obtained near atmospheric pressure, conditions that favor the formation of light products, especially methane. In our study, we have extended this work to FT synthesis conditions (>500 kPa, 473 K), where Cs+products form with high selectivity (>80%). We conclude that chain growth rates (and as we will show later, selectivity) depend only weakly on Co dispersion and on the identity of the support. This is not entirely unexpected because surface structure does not change markedly in the Co dispersion range studied here (0.45-9.5%) and because changes in electron density should not occur from contact between crystallites in this size range (10-210 nm) and a metal oxide support (94). We cannot rigorously establish, however, the structure insensitivity of this reaction on Co because of the limited dispersion range that we have explored. Instead, we conclude that FT synthesis pathways are not affected by the small structural changes that occur as crystallite size varies between 10 and 210 nm. Consequently, increased FT synthesis rates on Co catalysts will require concomitant increases in metal dispersion. Clearly, previously reported changes in turnover rate or selectivity in similar ranges of crystallite size cannot reflect direct structural effects on surface chemistry. The apparent dispersion and support effects on FT selectivity (Table 111) will be described in a later section. In the following sections, we first describe the reaction pathways and the structural catalyst properties that control the carbon number distribution and the paraffin selectivity during FT synthesis.

c.

BEDRESIDENCE TIMEEFFECTSON CHAIN GROWTH PROBABILITY AND PRODUCT FUNCTIONALITY

Secondary reactions are affected by the residence time of primary products within the catalyst bed while primary reactions are controlled only by the fugacity of reactants (and products if they affect the rate and selectivity of primary reactions pathways). Therefore, studies of the effects of bed residence time and of the presence of reaction products in the H K O feed are

247


ideally suited to examine the relative roles of primary and secondary pathways in the control of FT synthesis selectivity. Methane selectivity decreases and Cs+ selectivity increases with increasing bed residence time and CO conversion on Ru (Fig. 4a) and Co (Fig. 4b) catalysts. Thus, higher residence times of reaction products within the catalyst bed favor chain growth and the formation of higher molecular weight hydrocarbons. Bed residence time was increased by lowering the reactant space velocity at constant reactor temperature and pressure. CO conversion increased almost linearly with increasing bed residence time, suggesting that site-time yields were nearly constant throughout a wide range of CO conversion (5-72%) and bed residence time (2-12 s). Chain growth pathways do not obey Flory kinetics on Ru (Fig. 5a) or Co (Fig. 5b) catalysts. At all conditions, carbon number distribution plots are nonlinear and suggest that the chain growth probability ( a ) increases with increasing hydrocarbon size. Selectivity toward high-molecular-weight hydrocarbons increases as bed residence time increases because of the apparent conversion of light products into heavier ones. Chain growth probabilities become independent of bed residence time and carbon number for Czs+ hydrocarbons (Fig. 5). These findings can be described by an enhancement of secondary readsorption reactions as the residence time and the concentration of reactive a-olefins increase within bed interstices and catalyst pellets. Interpellet removal rates depend on space velocity and give rise to the observed increase in product molecular weight as bed residence time increases (Fig. 4). Intrapellet transport depends on molecular size and leads to the observed increase in chain growth probability with increasing chain size (Fig. 5). 95

95

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8

1 0 1 2 1 4

Bed Residence Time ( 8 )

0

2

4

6

8 1 0 1 2 1 4

Bed Residence Time (s)

FIG.4. Bed residence time effects on methane and Cs+ selectivity on (a) Ru/Ti02 (catalyst: 1.2% Ru, 48% Ru dispersion; 476 K, 560 kPa, H2/C0 = 2.1, 5-60% CO conversion) and (b) Co/Ti02 (catalyst: 11.7 wt% Co, 1.5%Co dispersion, 78 h-' site-time yield at 50% conversion; 473 K, 2000 kPa, H*/CO = 2. I , 9.5-72% CO conversion).

ENRIQUE IGLESIA et al.

248

lo-'

--

10'2

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10-4

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20 30 Carbon Number (n)

10

40

50

0

10

20

30

40

50

Carbon Number (n)

FIG.5. Bed residence time effects on carbon number distributions. (a) Ru/Ti02 (catalyst: 1.2 wt% Ru, 48% Ru dispersion; 476 K , 560 kPa, H2/C0 = 2.1, 5-60% CO conversion). (b) Co/Ti02 (catalyst: 11.7 wt% Co, 1.5% Co dispersion; 473 K , 2000 kPa, H2/CO = 2.1, 9.5-72% CO conversion).

Bed residence time effects on olefin and paraffin selectivity for a given hydrocarbon size also show that chain initiation by readsorbed olefins is the predominant reaction of a-olefins during FT synthesis at high reactant pressure (>500 kPa) on both Co and Ru catalysts. Olefins are selectively consumed in readsorption reactions only to reappear (with a certain probabil4

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Carbon Number (n)

FIG.6. Bed residence time and carbon number effects on the a-olefin to n-paraffin ratio. (a) Ru/Ti02 (catalyst: 1.2 wt% Ru, 48% Ru dispersion; 476 K , 560 kPa, HdCO = 2.1, 5-60% CO conversion). (b) Co/Ti02 (catalyst: 11.7 wt% Co, 1.5% Co dispersion; 473 K , 2000 kPa, H2/C0 = 2.1, 9.5-72% CO conversion).

249


ity) as an unreactive paraffin via a hydrogen addition termination step after several sequential readsorption events. As a result, the paraffin content increases with increasing bed residence time and chain size on Co (Fig. 6a) and Ru (Fig. 6b) catalysts. The paraffinic nature of higher hydrocarbons and of the products of high-conversion FT synthesis has been previously attributed to secondary hydrogenation of a-olefins (23, 33, 34, 48). Such a reaction can occur on a separate catalytic function during FT synthesis, but it is readily poisoned by CO and water (see Sections III,D and 111,F). Also, this proposal cannot explain the concurrent effects of carbon number and bed residence time on chain growth probability because hydrogenation converts olefins to paraffins of equal size but does not affect chain growth. Bed residence time studies also show that secondary reactions selectively extract olefins from low-molecular-weight fractions without a corresponding increase in the selectivity to the corresponding paraffin of equal size (Fig. 7). For example, 1-butene is selectively consumed with increasing bed residence time on Co/Ti02 (Fig. 7a), but n-butane selectivity is essentially independent of bed residence time, suggesting that hydrogenation cannot account for the disappearance of 1-butene. Hydrogenolysis of a-olefins is not responsible for their disappearance from the products because the selectivity to methane and other light paraffins, typical products of hydrogenolysis steps, does not increase with increasing bed residence time. Selectivities to light paraffins are independent of bed residence time, but selectivities to larger paraffins increase because of the combined effect of the readsorption and chain initiation by lighter olefins. The resulting chains continue to grow and ultimately terminate with a low but finite probability as n-paraffins. In-

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Bed Residence Time (s)

FIG.7. Bed residence time effects on olefin and paraffin carbon selectivity. (a) Ru/Ti02; C3 hydrocarbons (catalyst: 1.2 wt% Ru, 48% Ru dispersion; 476 K , 560 kPa, Hz/CO = 2. I , 5-60% CO conversion. (b) Co/TiOz; Cq hydrocarbons (catalyst: 11.7 wt% Co, 1.5% Co dispersion; 473 K , 2000 kPa, H,/CO = 2.1, 9.5-72% CO conversion).

250

ENRIQUE IGLESIA et

al.

ternal butenes are much less affected by bed residence time than are aolefins, suggesting that they readsorb much less readily during FT synthesis. Olefin selectivities also decrease with increasing bed residence time and chain size on Ru catalysts (4,14). For example, propylene selectivity decreases with increasing bed residence time without a corresponding increase in propane selectivity, leading to a net decrease in the fraction of the converted CO that appears as C , molecules (Fig. 7b). Readsorbed olefins initiate chains that continue to grow and ultimately desorb as larger olefins or paraffins. Qualitative trends are similar on all supports and on both Ru and Co catalysts. The selectivity details depend on the support physical structure, on the density of exposed surface metal atoms, and on the intrinsic readsorption properties (&) of Co and Ru surfaces. a-Olefins, n-paraffins, and internal olefins are primary FT synthesis products. They form by the direct termination of surface alkyl chains and are present among synthesis products even at very short bed residence times (24 h onstream).

216

ENRIQUE IGLESIA et

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ultimately become unavailable for FT synthesis reactions. They may, however, become active for olefin hydrogenation, a reaction normally inhibited by higher intrapellet CO concentrations present during kinetic-limited FT synthesis. As a result, FT synthesis products become lighter and contain predominantly paraffins even in the C T C ~product fraction. Reactant transport restrictions that occur at high values of x also lead to a decrease in the FT synthesis activation energy from 125 kJ mol-' to values of 65-75 kJ mol-' (Fig. 23). This behavior reflects primary reactions that become limited by reactant transport. The overall pressure order increases as y, approaches values where CO transport limits overall reaction rates (Fig. 23). This reflects the increased severity of transport restrictions as CO pressures decrease, because transport rates decrease linearly with total pressure but kinetic pellet requirements decrease much less rapidly (cf. Table I). The olefin readsorption and CO transport models suggest the types of catalyst structures and active site distributions that lead to optimum values of CS+selectivity in FT synthesis reactions (Fig. 20). The value of the structural parameter x required to give this optimum selectivity depends on the

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277

relative permeabilities (diffusivity and solubility) of CO and olefins in hydrocarbon liquids contained within catalyst pores. These properties do not depend on pore size (because diffusion occurs through liquids), on intrinsic surface kinetics, or on the chemical identity of the catalyst support. Optimum x values also depend on the relative rates of CO consumption and olefin readsorption, described by &. Therefore, x values required for maximum Cs+ selectivity can be achieved by changing the liquid properties, the readsorption activity of FT synthesis sites, or their geometric distribution within catalyst pellets. Optimum catalysts, in effect, behave as catalytic membrane systems that selectively retain olefin products but often maintain fast transport and equilibrium concentrations of reactants within pellets. The permeability of such catalytic membranes can be modified by changing the width of the catalytic region in a pellet, [e.g., in eggshell catalysts (SO-53)] or the liquid molecular weight [e.g., by extensive recycle (123, 124)]; these changes in x effectively move catalysts along the abscissa in Fig. 20. Changes in permeability do not alter the shape of such curves; changes in transport selectivity (permselectivity), however, can broaden or narrow the range of x values that lead to maximum CS+yields. Selective inhibition of olefin transport moves the left side of the curve in Fig. 20 ( x < 10” m-I), a region where olefin readsorption is the primary modifier of catalytic selectivity. It does not affect catalysts with x values greater than about 1019m-’, where selectivity changes occur as a result of reactant transport restrictions. Similarly, selective enhancement of CO transport without changes in olefin diffusivity moves the right side of such curves farther to the right and also broadens the x region where highest CS+ values are obtained. Such changes in permeation selectivity are likely to occur as the water concentration varies with axial position, and possibly within catalyst pellets, in high-conversion packed-bed reactors. These permeation selectivity changes would require the presence of hydrocarbon liquids that selectively increase the rates of CO or HZdiffusion, but which continue to restrict the exit of hydrocarbons from catalyst pellets. Such effects can, however, be achieved by the use of a size-selective membrane to replace the permeability and selectivity properties of the liquid “membrane” naturally occurring within catalyst pores. Recently, the use of carbon sieve coatings on the outer surface of Fischer-Tropsch catalyst pellets has been reported (125). These coatings appear to retain larger hydrocarbons, which undergo secondary hydrogenolysis reactions within pellets that contain both FT synthesis and hydrogenolysis functions, without introducing CO transport limitations. More subtle effects can be introduced by selectively varying the permeation selectivity phase of CO and HZreactants in the liquid phase. Then, stoichiometric ratios (-2: 1) can be maintained around intrapellet active sites even when feed ratios differ significantly from

278

ENRIQUE IGLESIA et

al.

these stoichiometric values. Porous glass membranes have been used to exploit the different gas-phase diffusivities of H2 and CO in order to alter the HdCO ratio in nonstoichiometric feed mixtures (126). V.

Design of Fischer-Tropsch Catalysts

In a process such as Fischer-Tropsch synthesis, where transport limitations are ubiquitous and often unavoidable even in laboratory reactors, we obtain an unexpected benefit from catalyst design options, unavailable in kinetic-limited pellets, introduced by structural parameters. Site activity and chain growth kinetics appear to be intrinsic properties of surface metal ensembles and depend weakly on site modifications introduced by changes in metal crystaliite size or in metal oxide support (41). Synthesis rate and selectivity changes that take place when Co is alloyed with another metal (I27--129)often reflect increased Co dispersion or reducibility, instead of changes in intrinsic FT surface chemistry (127-129). Thus, the intrinsic behavior of Co catalysts in FT synthesis is difficult to modify by varying impregnation or pretreatment procedures that influence metal dispersion or by choosing different support materials. Consequently, transport restrictions become one of the few selectivity modifiers available for the design of Co catalysts in FT synthesis reactions, a process in which selectivity is additionally constrained by polymerization chain growth kinetics. In what follows, we exploit the inherent coupling between reactive and transport processes in order to design FT catalysts with desired selectivity. We explore structural modifications and reactor conditions that allow the selective introduction and removal of transport restrictions and of additional catalytic functions for secondary reactions of reactive a-olefins. Our discussion includes model simulations to explore optimum designs, new interpretations of reported selectivity improvements, and some of our own experimental attempts in the design of optimum FT synthesis catalysts. Earlier, we identified a challenge in materials synthesis: the preparation of catalysts with optimum values of the structural parameter x. Here, we identify additional challenges but also introduce some flexibility by suggesting alternate and often complementary approaches to achieve higher selectivity to desired FT synthesis products. A. MODIF~CATION OF SURFACE REACTIONS AND OF ISOLATED SITE KINETICS

Olefin hydrogenation reactions control Cs+ selectivity by intercepting reactive a-olefins . Readsorption and chain initiation steps reintroduce olefins into chain growth pathways and reverse the most frequent chain termination

279


step: p -hydrogen abstraction. As a result, readsorption lowers the total chain termination probability and increases product molecular weight during FT synthesis. In contrast, olefin hydrogenation prevents the reversal of the hydrogen abstraction termination step by converting olefins irreversibly to paraffins. In this section, we explore how nlodifications of the intrinsic readsorption activity of FT synthesis sites and of the activity and density of secondary hydrogenation sites influence carbon number distributions during FT synthesis. Our readsorption model shows that carbon number distributions can be accurately described using Flory kinetics as long as olefin readsorption does not occur (pr= 0), because primary chain termination rate constants are independent of chain size (Fig. 24). The resulting constant value of the chain termination probability equals the sum of the intrinsic rates of chain termination to olefins and paraffins (Po + pH).As a result, FT synthesis products become much lighter than those formed on Co catalysts at our reaction conditions (Fig. 24, pr = 1.2), where chain termination probabilities are much lower than Po PH for most hydrocarbon chains. The product distribution for pr = 1.2 corresponds to the intermediate olefin readsorption rates experimentally observed on CoiTiO:! catalysts, where intrapellet transport restrictions limit the rate of removal of larger olefins, enhance their secondary chain initiation reactions, and increase the average chain size of FT synthesis products.

+

10’ I

1

1

I

1

10-2

Carbon Number (n)

FIG. 24. Effect of enhanced a-olefin readsorption rates on carbon number distributions (simulations: experimental/model parameters as in Fig. 16).

280

ENRIQUE IGLESIA et

al.

Carbon number distribution plots also become linear when olefins readsorb very rapidly (large &) or when severe intrapellet transport restrictions (large x ) prevent their removal from catalysts pellets before they convert to paraffins during chain termination (Fig. 24, p, = 100). In this case, chain termination to olefins is totally reversed by fast readsorption, even for light olefins. Chain termination occurs only by hydrogen addition to form paraffins, a step that is not affected by secondary reactions and for which intrinsic kinetics depend only on the nature of the catalytic surface. The product distribution again obeys Flory kinetics, but the constant chain termination probability is given by PH,instead of Po + PH.Clearly, bed and pellet residence times above those required to convert all olefins cannot affect the extent of readsorption or the net chain termination rates and lead to Flory distributions that become independent of bed residence time. The contributions of olefin readsorption and chain initiation to chain growth can also be controlled by introducing a second catalytic function that intercepts olefins and converts them to unreactive paraffins in secondary hydrogenation reactions. Thermodynamic constraints prevent this process from being reversed again; therefore, little benefit can be obtained from the introduction of a dehydrogenation function within FT catalyst pellets. We have introduced into our olefin readsorption model a first-order hydrogenation step ( p s ,Fig. I), which competes with readsorption (pr)for available olefins, in order to simulate the effect of a secondary hydrogenation function. Its contribution as a modifier of FT selectivity depends on the ps/& ratio. Cs+ selectivity decreases as the extent of olefin hydrogenation (Ps/pr)increases (Fig. 25, curves A and B), a catalyst design that is achieved experimentally by increasing the density or the reactivity of hydrogenation sites within FT synthesis pellets. Hydrogenation sites within pellets interfere with chain initiation by olefins much more effectively than do similar sites placed outside such pellets, because hydrogenation steps occur at the higher olefin fugacities, created by intrapellet transport restrictions, from which olefin readsorption steps also benefit (cf. curves B and C in Fig. 25). Intrapellet hydrogenation sites weaken the effect of diffusion-enhanced olefin readsorption and lead to lighter and more paraffinic products that begin to approach Flory distributions. Bed residence time continues to enhance readsorption reactions (curves A and B), except at very high intrapellet values of PI/@, where few olefins exit a pellet unreacted. The presence of hydrogenation sites outside FT synthesis pellets (e.g., as a physical mixture of FT and hydrogenation pellets) is much less effective (Fig. 25, curve C). Such extrapellet sites merely capture (low fugacity) olefins that leave FT synthesis catalyst pellets unreacted and prevent their subsequent readsorption along the catalyst bed. They do not prevent the chain initiation reactions that olefins undergo predominantly within

28 1

FISCHER-TROPSCH SYNTHESIS 100

1

1

-

-

aut

70-

60

102

I

I

10'

100

10'

PJP, Ratio FIG. 25. Effect of secondary hydrogenation of a-olefins on C5+ selectivity (simulations: experimental/model parameters as in Fig. 16). (A) 9.5% CO conversion, intrapellet reaction; (B) 72% CO conversion, intrapellet reaction; (C) 9.5-72% CO conversion, interpellet reaction.

transport-limited pellets. External hydrogenation sites lead to more paraffinic products that retain a non-Flory product distribution. They lessen bed residence effects on FT synthesis selectivity by intercepting olefins before they interact with other FT synthesis pellets, before their convective removal from the catalyst bed. Diffusional restrictions increase the effectiveness of olefin interception sites placed within catalyst pellets. Very high olefin hydrogenation turnover rates or site densities within pellets prevent olefin readsorption and lead to Flory distributions of lighter and more paraffinic hydrocarbons. Identical results can be obtained by introducing a double-bond isomerization function into FT catalyst pellets because internal olefins, like paraffins, are much less reactive than a-olefins in chain initiation reactions. However, light paraffins and internal olefins are not particularly useful end-products in many applications of FT synthesis. Yet, similar concepts can be used to intercept reactive olefins and convert them into more useful products (e.g., alcohols) and to shift the carbon number distribution into a more useful range. In the next section, olefin readsorption model simulations are used to explore these options in the control of FT synthesis selectivity.

B. DIFFUSION-ENHANCED BIFUNCTIONAL CATALYSIS AND THE CONTROL OF CARBON NUMBER DISTRIBUTION AND WODUCT FUNCTIONALITY

Here, we describe the simulated effect of a second catalytic function (hydroformylation or cracking) on product selectivity. Hydroformylation in-

282

ENRIQUE IGLESIA et

al.

volves the reaction of an olefin of size n with CO and Hz, leading to an alcohol of size n + l . Such reactions can also occur during FT synthesis (118-120), often as a primary chain termination step involving CO insertion into surface alkyl groups. Our simulations assume that alcohols formed in these secondary reactions do not react further in chain initiation and growth steps during FT synthesis, an assumption that is unlikely to hold for many catalysts and reaction conditions (130). In effect, we have assumed that hydroformylation reactions intercept olefins and prevent their secondary readsorption and chain initiation reactions, but in contrast with hydrogenation steps, transform these olefins into more useful products. The effectiveness of this secondary reaction (ka, Fig. 1) depends on the ratio of hydroformylation and readsorption rate constants (pa/&).The effect of intrapellet transport restrictions on alcohol selectivity is shown in Fig. 26a for pa/Pr= 0.5. These results were obtained for a catalyst with kinetic parameters identical to those used to describe selectivity data on Co catalysts in Figs. 16, 17, and 19. Not surprisingly, transport-limited pellets favor these secondary hydroformylation reactions and alcohol selectivity increases with increasing values of the Thiele modulus (Fig. 26a, curve A). Clearly, olefin hydroformylation pathways are most effective when they compete locally with readsorption and chain initiation at high intrapellet olefin fugacities within transport-limited FT synthesis pellets. Outside pellets, hydroformylation sites use only those few olefins that exit FT catalyst pellets after extensive readsorption and chain initiation (Fig. 26a, curve B). Hydroformylation reactions on these external sites occur at much lower rates, which simply reflect the lower olefin fugacities in the gas phase; as a result, such extrapellet sites affect FT and alcohol synthesis selectivity only slightly. 0.05

b

4 70

I

l

AlCOhOlS

2

6

0.02

0.01

t 0

0 Thlele Modulus,

41

5 10 15 Carbon number (n)

FIG. 26. Bifunctional Fischer-Tropsch/hydroformylation catalysts-model simulations. (a) Alcohol selectivity vs. Thiele modulus (Po = 0.40, p, = 1.2, BH = 0.058, p. = 0, Co/ TiOn parameters; curve A, intrapellet reaction, Pa& = 0.5; curve B, extrapellet reaction, Pa/& = 1 .0). (b) Carbon selectivity vs. carbon number (intrapellet reaction, Pa/& = 0.5).

283


Secondary hydroformylation sites inhibit chain initiation pathways by converting olefins to less reactive species (alcohols). As a result, they inhibit chain growth and decrease product molecular weight, leading to high yields of C3-C15 alcohols, useful molecules in fuels and petrochemical applications (Fig. 26b). Olefin selectivity decreases markedly with increasing carbon number because of intrapellet conversion of olefins to alcohols (Fig. 26b). Carbon number distributions approach Flory kinetics at higher carbon numbers because paraffins and alcohols, the predominant products in this carbon number range, are assumed to be unreactive FT synthesis products in our simulations. Clearly, the application of this concept requires the use of selective hydroformylation sites that are compatible with FT synthesis catalyst pellets and reaction conditions; it also requires chain growth sites that do not catalyze the subsequent readsorption and chain initiation by these alcohols. Olefin and paraffin cracking have been often proposed as subsequent upgrading steps for FT synthesis products (56, 131). Secondary olefin cracking reactions also inhibit readsorption by intercepting reactive olefins and forming paraffins as one or more of the cracking products ( k c , Fig. 1). In contrast with hydrogenation and hydroformylation reactions, cracking reactions convert olefins into smaller fragments; thus, cracking shifts the carbon number distribution of FT synthesis products. The effect of cracking reactions during FT synthesis (132) has been previously explored using kinetic models of chain growth and secondary reactions that neglected intrapellet transport restrictions. Here, we show how such intrapellet restrictions are essential to exploit the useful cracking activity of this second catalytic function. We use here a simplified kinetic model of cracking reactions in order to illustrate the role of secondary cracking steps on product distribution. More detailed and rigorous models are available but the additional rigor is not essential to describe the concepts that we illustrate here. We assume that sites within transport-limited pellets (all kinetic-transport parameters as in Figs. 16, 17, and 19) catalyze the cracking of olefins with a probability given by

pc = 0,

IZ 5

7;

pc = @(IZ

-

7 ) , IZ

2

8

(27)

The carbon number dependence reflects the relative reactivity of carbon-carbon bonds in cracking reactions (133,134). In this kinetic scheme, cracking does not occur at the three C-C bonds nearest the end of a molecule but occurs with equal probability at all other C-C bonds. Such reactive trends are qualitatively similar to those reported for carbenium iontype cracking reactions on acid catalysts (56,133,134). We express the cracking probability as a ratio (p?/pr)in order to describe the competition between cracking and readsorption pathways for available a-olefins. Also,

284

ENRIQUE IGLESIA

et d.

we assume that both products of cracking events are unreactive in subsequent cracking, readsorption, and chain initiation reactions, an assumption that we can easily relax later as experimental verification of these concepts warrants their further study. Diffusion-enhanced cracking of a-olefins leads to much higher yields of C4-Cs products at the expense of C9+hydrocarbons (Fig. 27a, curves A and B). The combined effects of olefin interception before readsorption and random reversal of chain growth pathways by cracking lead to Clo+product distributions that obey the Flory equation; the chain termination probability for such hydrocarbons resembles those obtained in the absence of readsorption and chain initiation pathways (Fig. 27b, curves A and B). Intrapellet cracking sites lead to high selectivity to Cs-C9 hydrocarbons because cracking kinetics permit reactions of Cs+ products but not of smaller hydrocarbons (Table VI). Thus, cracking reactions alter carbon number distributions and lead to intermediate-size products with selectivities greater than predicted from Flory polymerization kinetics. In contrast with readsorption reactions, which broaden the carbon number distribution of FT synthesis products, cracking reactions narrow such distributions, within the constraints imposed by the random nature of C-C bond cleavage in carbenium ion reactions of large olefins. Cracking of nparaffins can also occur on intrapellet acid sites, but acid-catalyzed paraffin reactions are much slower than those of corresponding olefins of equal size. As in all secondary reactions, cracking sites are used most efficiently when

0

5

10

15

20

CARBON NUMBER (n)

26

30

CARBON NUMBER (n)

FIG.27. Effect of diffusion-enhanced a-olefin cracking catalytic function on carbon number distribution (simulations: experimentalhodel parameters as in Fig. 15, 10%CO conversion). (A) FT synthesis without cracking function; (B) with intrapellet cracking function, /3: = 1.2; (C) with extra pellet cracking function, /3: = 1.2. (a) Carbon selectivity vs. carbon number; (b) Flory plots.

285


TABLE VI Simulated Effects of Intrapellet Cracking Sites on FT Synthesis Selectivity

Carbon number range

FT synthesis

FT synthesis with intrapellet cracking"

CTC4

9.9 15.5 13.3 61.3

30.2 49.6 11.1 9.1

crc9 cI(Tcl6

C17+ a

FT synthesis with extra pellet cracking" 13.7 26.3 17.4 42.6

@ = 1.2, experimental and model parameters as in Fig. 15.

located within transport-limited pellets. The resulting high intrapellet f'ugacities provide the kinetic driving force required to catalyze olefin cracking reactions at FT synthesis conditions. Reactions of a-olefins on extrapellet cracking sites lead to more modest changes in selectivity (Table VI; Fig. 27a, curve C) and preserve the non-Flory nature of the carbon number distribution. This provides an important advantage over extrapellet sites, where cracking usually requires significantly higher temperatures than the FT synthesis reactions (133, 134). Composite materials consisting of mixtures of Fe and H-ZSM-5 components have been widely reported to catalyze concurrent FT synthesis and olefins and paraffin cracking (135-140). These bifunctional processes occur at high temperatures ( > S O K ) and lead to high CH4 and C 0 2 yields. Paraffin hydrogenolysis on metal sites occurs at lower temperatures, especially on Ru catalysts, but it is strongly inhibited by CO during F T synthesis. Transport-enhanced hydrogenolysis can lead to non-Flory distributions, as suggested recently (125), but it will additionally require sites uninhibited by CO; this may be achieved instead when strong reactant transport limitations deplete CO near hydrogenolysis sites placed toward the center of a pellet. In a previous section, however, we showed that even in the absence of CO, Co sites do not catalyze the hydrogenolysis of F T synthesis products at our reaction conditions, perhaps because of the remaining presence of water. Product distribution changes have been previously attributed to solubilityenhanced hydrogenolysis of large paraffins (35).Hydrogenolysis, however, is unlikely to occur even in the absence of CO at the conditions of such studies (Section IV,F). Moreover, we previously showed that solubility cannot affect the kinetic driving force of secondary reactions. Thus, secondary hydrogenolysis reactions are unlikely to be the underlying cause for the reported selectivity changes. We suggest instead that they arise from reactant transport limitations imposed by the liquid phase within the catalyst pores; such limitations can decrease local CO fugacities near catalytic sites and lead to the lighter products reported by these authors (35).

286

ENRIQUE IGLESIA et

al.

Several other secondary reactions of olefins can also be used to modify FT synthesis selectivity. For example, olefins can oligomerize within bifunctional catalyst pellets; they can also incorporate into chain growth sites, rather than just initiate them, at higher olefin concentrations and lower temperatures. Finally, they can be converted into less reactive internal olefins or branched a-olefins and thus kept from reentering chain growth pathways, in much the same way as in secondary hydrogenation steps. These bifunctional pellets share in common an enhancement in bifunctional reaction rates, which arises from the delayed removal of reactive olefins from catalyst pellets. The presence of secondary reaction sites within transportlimited FT synthesis pellets significantly extends our ability to catalyze other selective chemistries at the mild conditions that also favor chain growth during FT synthesis.

c.

CHANGES IN THE LIQUID COMPOSITION WITHIN TRANSPOIC-LIMITED CATALYTIC PELLETS

In Section IV,H, we described how selectivity depends on the transportlimited arrival of reactants and removal of reactive olefins within catalyst pellets. Specifically, we showed how FT synthesis selectivity can be described accurately and entirely by a structural parameter x, regardless of whether olefin removal [Eq. (15)] or reactant arrival [Eq. (25)] is the controlling diffusive resistance. The preceding two sections described how transport-limited removal of olefins from catalyst pellets can enhance the rate of secondary reactions. Here, we show how such transport limitations can be controlled by varying the liquid composition within catalyst pellets. In the discussion that follows, we refer to the experimental and simulation results of Fig. 20, where we showed how Cs+ selectivity depends on the value of the structural parameter (x). Cs+selectivity can be controlled by varying the density of sites within catalyst pellets and the diameter of these pellets. The density of sites determines the reactant requirements and the pellet size controls the required path length of diffusing molecules. Such modifications affect the value of x, causing the performance of these catalysts to move along the curve in Fig. 20. The shape of the selectivity curve, however, depends only on the intrinsic readsorption rate constant (pr)and on the kinetic dependence of chain growth pathways on reactant pressure. Site density and pellet size in FT synthesis catalysts may not be entirely within the control of the catalyst designer. Low site densities may be required in order to bring catalysts limited by CO transport into the optimum region of x values, a modification that unfortunately also lowers reactor productivity. Pellet size may be decreased instead in order to eliminate CO


287

transport restrictions, but small pellets typically lead to high pressure drop in fixed-bed reactors. Thus, we must suggest alternate design strategies that complement these structural catalyst modifications. One such strategy is to lower the molecular weight of the liquid product within catalyst pellets in order to facilitate transport; another is to selectively deposit active sites in specific radial positions within a pellet. The diffusivity of CO, HI, and hydrocarbons within catalyst pores increases as the contained liquid becomes lighter or as it is replaced by a gaseous or supercritical phase. Then, the parameters +, and +co that precede x in Eqs. (15) and (25) decrease. For a given catalyst structure (x), transport limitations become less severe and pellets become increasingly limited by reaction kinetics. The net effect is that selectivity curves, such as the one in Fig. 20, are shifted toward the right without any significant change in their shape. Thus, we can design catalytic materials that avoid CO transport limitations while maintaining high productivity and using the large pellets required in packed-bed reactors. The removal of heavy products from catalyst pores can be achieved in several ways. The average molecular weight of the product can be decreased by continuous recycle of light hydrocarbon fractions that shift the composition of vapor-liquid equilibrium mixtures. The liquid can be totally removed by reactor operation at temperatures well above its boiling point, but such temperatures usually lead to the formation of very light hydrocarbons. More interestingly, it appears that operation of F T synthesis reactors in the presence of supercritical concentrations of light paraffins (e.g., n-hexane) also decreases the average molecular weight of intrapellet liquids (123,124). Such studies report an apparent increase in the rate of CO hydrogenation and in olefin content at supercritical conditions. These authors attribute such effects to the removal of waxy liquids from catalyst pores by the supercritical solvent, a process that decreases the rate of secondary hydrogenation of retained olefins. Supercritical light paraffins can decrease the intrapellet liquid molecular weight simply by shifting the vapor-liquid equilibrium point, rather than by their supercritical behavior. These changes in the vapor-liquid equilibrium point lead to lighter, perhaps even gaseous, products within pellets and to the removal of CO transport limitations. As a result, CO hydrogenation rates and Cs+selectivity increase and the products become increasingly olefinic, as observed experimentally (97,98). Our models are in qualitative agreement with the effect of supercritical conditions of FT synthesis selectivity. A more quantitative analysis of such effects requires more detailed information on the structural and reactive properties of the catalysts used in such studies (123,124) and on the rates of CO, HZ, and hydrocarbon diffusion in supercritical hydrocarbons. We suggest that recycle conditions shift the maximum value of CS+selectivity in

288

ENRIQUE IGLESIA

et at.

Fig. 20 toward higher values of x, without a significant change in the shape of the curve. Therefore, the net effect of supercritical operation will depend on whet her catalysts at normal operating conditions reside on the olefin readsorption or reactant transport parts of the Cs+ selectivity curve (Fig. 20). DISTRIBUTION OF ACTIVE SITES WITHIN D. NONUNIFORM TRANSPOm-LIMITED PELLETS The pellet size term (&) in the defining equation of y, [Eq.(15)] reflects the diffusion path length required to reach catalytic sites. Pellet size and diffusion path lengths are identical only for uniformly impregnated pellets. They can be decoupled by placing catalytic sites in specific regions within catalyst pellets. For example, eggshell catalysts, wherein catalytic sites are placed only in regions near the outer pellet surface, provide short diffusion paths to reactive sites, irrespective of the pellet size required to control pressure drop across the catalyst bed. The benefits of nonuniform activity distributions (site density) or diffusive properties (porosity, tortuosity) within pellets on the rate of catalytic reactions were first suggested theoretically by Kasaoka and Sakata (141). This proposal followed the pioneering experimental work of Maatman and Prater (142). Models of nonuniform catalyst pellets were later extended to more general pellet geometries and activity profiles (143), and applied to specific catalytic reactions, such as SOz and naphthalene oxidation (144-146). Previous experimental and theoretical studies were recently discussed in an excellent review by Lee and Aris (147). Proposed applications in Fischer-Tropsch synthesis catalysis have also been recently reported (50-55,148), but the general concepts have been widely discussed and broadly applied in automotive exhaust and selective hydrogenation catalysis (142,147,149). FT synthesis rates and selectivities on catalyst pellets of varying diameters and active site distributions are shown in Table VII. Evenly impregnated large pellets (2.2 mm in diameter) give lower FT synthesis rates, CS+ selectivities, and activation energies than much smaller pellets. Crushing and sieving large pellets to obtain much smaller particles (0.17 mm) increases synthesis rates, Cs+ selectivity, and activation energy (Table VII), a result consistent with our previous discussion of pellet size effects, with the data in Table V, and with the removal of transport restrictions. Similar improvements in synthesis rates and Cs+ selectivities are observed when Co sites are selectively placed within 0.1-0.2 mm of the outer surface of large Si02 pellets. Eggshell catalysts are more active and selective for CS+ synthesis than evenly impregnated large pellets (Table VII), but both

289


TABLE VII Catalyst Characterization and Fischer-Tropsch Synthesis Data on Co/SiOz Powders and on Even and Eggshell Large Pellets" Catalyst (12.7-13.1 wt% Co/SiOz)

Synthesis data Co site density (106g-atom surface Co m-2) Pellet size (mm) Impregnated pellet region (average eggshell thickness or pellet diameter, mm) Characteristic diffusion distance (mm) Co dispersion (%) y , (10-16m-'c) Site-time yield (h-') C& selectivity (76) C5+ selectivity (%) COz selectivity (%) Activation energy (kJ mol-I) Propylene/propane ratio I-butenelmbutane ratio

Even large pellet

Crushed large pellet

1.67

1.67

2.2 2.2

0.17 0. I7

1.1

0.085 6.3 136 95 5.2 89.6 0.18

6.3 22,000 50 12.1 81.5 0.90 68 0.20 0.11

Eggshell pellet

I .48 (3.2)b 2.2 0.23

0.23 5.5 1910 I05 7.7 88.0 0.50

-

-

2.6 1.7

0.50 0.28

Crushed eggshell pellet 1.48 (3.2)b 0.17 0.17

0.085 5.5 272 140 4.7 90.5 0.16 2.7 1.8

Small pellet 0.38 0.17 0.17

0.085 5.0 95 110 7.0 83.5 0.15 115 2.9 2.7

a Obtained at 473 K , 2000 kPa, H2/C0 = 2.1; 55-65% CO conversion, > 24 h onstream. Within impregnated shell region. y, = L28@/r,;r, in meters, site density (8) in Co atomsim', t as twice the characteristic diffusion distance in meters; calculated from Eq. (15).

cobalt-time yields and Cs+ selectivities increase even further when these samples are also crushed to 0.17-mm pellets. This clearly suggests that synthesis rates on these eggshell catalysts still remain partially limited by the rate of arrival of reactants at catalytic sites. Methane selectivity and olefin content in synthesis products are also influenced by these transport restrictions. Olefin contents are higher on smaller or partially impregnated pellets (Table VII), suggesting that low intrapellet CO concentrations caused by diffusional limitations increase the rate of secondary hydrogenation reactions. In addition, large pellets increase the intrapellet residence time of reactive olefins and increase their probability for secondary hydrogenation and readsorption reactions. Methane selectivity decreases as we remove transport restrictions, whether such changes occur as the result of partial impregnation or of decreasing pellet size (Table VII).

290

ENRIQUE IGLESIA et

al.

Experimental methane and Cs+ selectivities on supported Co catalysts ( S O z , TiO2, A1203)and on Co/Si02 eggshell and even pellets (data in TaThis figure bles 111, V, and VII) are shown in Fig. 28 as a function of also includes data on other materials in which the structural parameter was varied by changing the metal dispersion and the support (Table 111) or the pellet diameter (Table V). CS+selectivity initially increases as transport restrictions worsen with increasing value of (Fig. 28a). This occurs because while reactant transport restrictions are negligible at these low values of a-olefin readsorption, a reaction that increases product molecular weight, is favored by the slow removal of large molecules from catalyst pellets. Methane selectivity decreases over a similar x range (Fig. 28b), also as a result of the favored growth of hydrocarbon chains under mild diffusional restrictions. The dashed curves in Figs. 28a and 28b show the excellent agreement between model simulations and experimental results throughout the entire range of x values. Ultimately, further increases in lead to a decrease in product molecular weight and Cs+ selectivity because structural properties that limit olefin removal rates at low values of also begin to limit the rate of arrival of reactants at catalytic sites as increases (Fig. 28a). At higher values of methane selectivity also reverses its initial trend and begins to increase. These results show that these transport restrictions and the resulting selectivity changes are caused by an increase in x,independent of how the value of this parameter is varied.

x.

x

x,

x

x

x

x,

I

l 95

o

0

y

102

103

"

/--

,,'

-

70 101

I

I

]

0

104

105

107

103

102

x

(m1)

x

104

10s

1016

FIG. 28. Site density, pellet size, and eggshell thickness effects on FT synthesis selectivity (Co catalysts: SiOz, TiOz, and AI2O3supports; 473 K , H2/CO = 2.1, 2000 kPa, 50-62% CO conversion). 0 , Dispersion/support effects; A, pellet size effects; 0, eggshell catalysts (Table VII); A, from Refs. 47 and 48.


29 1

The data in Fig. 28 clearly show that intermediate values of x, which limit olefin removal and enhance secondary readsorption reactions but still permit unrestricted and rapid access of CO and HZto reaction sites, lead to maximum CS+selectivity. They also show that eggshell catalysts allow access to these intermediate values of y, for any pellet size. The design of eggshell pellets with values of between 0.2 and 2.0 X lOI9 m-* leads to high CS+selectivity (Fig. 28a) and maintains catalytic rates and activation energies near their intrinsic kinetic values (Table VII). Recently, we reported the synthesis of eggshell catalysts with sharply defined shell regions and high cobalt dispersion (53,68). The technique exploits the controlled penetration of a molten salt into mesoporous support pellets and maintains small Co crystallites by avoiding calcination pretreatments. This technique provides a simple and rather general method for the synthesis of high metal-loading eggshell catalysts, even when interactions between impregnating species and support are very weak. The preparation of such eggshell materials, however, presents several challenging synthetic tasks in the reproducible formation of thin and uniform catalytic layers. It also requires that we accommodate all metal surface atoms within a very small fraction of the available pellet and reactor volume, thus requiring very high metal dispersions at extremely high local loadings. Yet, materials prepared by these methods lead to remarkable changes in Cs+selectivity and the rigorous modeling of the inherent coupling between reactions and transport guides their synthesis and provides critical flexibility in the design of FT synthesis catalysts. Obviously, transport limitations can always be lessened by the use of less active catalysts, a solution that provides little relief to the design engineer attempting to minimize expensive reactor volume in large-scale FT synthesis plants. Thus, the ability to prepare sharp concentration profiles of active sites near the outer surface of large pellets clearly remains the most attractive solution to transport restrictions that strongly influence FT synthesis selectivity in packed-bed reactors. In contrast with many other applications of eggshell catalysts, it is selectivity rather than activity loss that we seek to prevent in the design of FT synthesis catalysts.

x

E. TRANSPORT EFFECTSAND THE CONTROL OF Fe-BASED FISCHER-TROPSCH SYNTHESIS

SELECTIVITY IN

Many of the chain growth pathways and transport effects described above also occur on Fe-based FT synthesis catalysts. As on Co and Ru catalysts, FT synthesis on Fe often yields non-Flory carbon number distributions of products, where the chain growth probability and the paraffin content increase with hydrocarbon chain size (38-40).These effects were previously

292

ENRIQUE IGLESIA et

ad.

attributed to the presence of several types of chain growth sites (38,39). We have suggested that transport-enhanced olefin readsorption models used to describe FT synthesis products on Co and Ru also account for products formed on Fe catalysts (40). As we have shown previously (40), carbon number distributions do not obey Flory kinetics on Fe catalysts promoted with Cu and K additives (atomic ratios of 100:21.8: 1 for Fe:Cu:K2C03)(Fig. 29a). Catalyst preparation procedures and additional catalytic data were reported previously (40, 58). The curved Flory plot resembles those on Ru (Fig. 5a) and Co (Figs. 5b and 18), but some clear differences exist. For example, CZand C3 hydrocarbons appear to lie on the distribution curve corresponding to the rest of the products (Fig. 29a), in contrast with their appearance below such a curve on Ru and Co. This suggests that intrinsic ethylene readsorption rates on Fe are not significantly higher than those of higher olefins or that the readsorption of the lighter olefin fraction is slow compared to the chain propagation of the corresponding surface chains. On Fe, a-olefins appear among FT synthesis products until much higher carbon numbers than on Co and Ru. Also, a significant fraction of FT synthesis products are alcohols, which can also readsorb and initiate surface chains during FT synthesis (130) and which are present even among highmolecular-weight products. These differences suggest that the reactivity of olefins and alcohols during FT synthesis on Fe is lower than on Co and Ru catalysts; consequently, olefin and alcohol readsorption require stronger diffusional restrictions on Fe catalysts than on Co or Ru. This leads to Flory plots that continue to curve until much higher carbon numbers on Fe, because reactive products remain available for diffusion-enhanced readsorption steps. 0.1,

5.-> 2

0.01 -

..

. ,

0.001 -

0.

*-.

*%.

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,

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'

'

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' b

:A

'

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A

Alcohols

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'PI

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*,.*'.

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.% 0.

0.00001.

,

.

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0.0001

,

,

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,

,

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,

.

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I

,

.

0.0 0

Carbon Number (n)

I

10

?&a :A,

20

[

,

30

,

40

Carbon number (n)

FIG. 29. FT synthesis product distribution on Fe-Cu-K catalysts (ratio 100 : 21.8 : 1, Fe:Cu:K2C03; 1520 kPa, 489 K , H&O = 2.0, 30% CO + Hzconversion). (a) Carbon number distribution; (b) chain termination probability.

50


293

Chain termination probabilities for a-olefins (PO),n-paraffins (pH),and alcohols (Pa)are plotted as a function of carbon number in Fig. 29b. The net rate of chain termination to olefins decreases with increasing carbon number without a proportional increase in the probability of chain termination to paraffins, except perhaps for very light paraffins. This suggests that secondary hydrogenation reactions cannot account for the selective disappearance of olefins as the carbon number increases on Fe catalysts. Chain termination to paraffins decreases slightly for short chains but then remains essentially independent of chain size; the slight decrease in PHwith increasing chain size disappears if we include isoparaffin products in our PH term. Therefore, except for an initial decrease in P H values for short chains, the observed trends are similar to those observed on Co and Ru catalysts. The behavior of alcohols during FT synthesis resemble that of a-olefins, consistent with their observed readsorption when added to H2/C0 reactants (130). Thus, the diffusion-enhanced readsorption of alcohols also contributes to the higher chain growth probability and paraffin content in large hydrocarbons. Chain termination steps leading to alcohols and olefins are partially reversed by their increasing pellet residence time at higher carbon numbers. Similar effects also account for the observed increase in product molecular weight and paraffin content as bed residence time and CO conversion increase. The introduction of an intrapellet double-bond isomerization function inhibits the readsorption of a-olefins by converting them to less reactive internal olefins. Pellets formed by mixing finely ground Fe-Mn-Cu-K catalysts (150) with dealuminated Y-zeolite (Si/Al = 7.0) (151) or with nonacidic SiOz (W. R. Grace, Davison 62 Grade) are compared in Fig. 30. Flory distributions are curved on SiOz-containing pellets (Fig. 30a). The replacement of SiOz with Y-zeolite leads to the disappearance of a-olefins (Fig. 30b) and to their replacement with internal olefins. This effect is accompanied by a decrease in product molecular weight and chain growth probability and to Flory plots that become almost linear (Fig. 30a). These results resemble those obtained by model simulations when a secondary hydrogenation function was introduced within FT synthesis pellets (Fig. 25) or when the effective readsorption probability decreases (Fig. 24). The slight curvature that appears at higher carbon numbers on the zeolite-containing sample reflects transport-enhanced readsorption of less reactive internal olefins. Clearly, the presence and the reactive nature of a-olefin products is essential for the growth of large hydrocarbons and for their non-Flory carbon number distributions also on Fe catalysts. The remarkably similar trends in carbon number distributions and chain termination kinetics on Fe, Co, and Ru catalysts demonstrate the important role of diffusion-enhanced secondary reactions, especially those that reverse

294

ENRIQUE IGLESIA

et d.

Q9 0

10-2

-

0 .

0 O.

"2 0

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.-+

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'

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0

10

20

30

40

50

Carbon Number (n)

Carbon Number (n)

FIG.30. The effect of secondary reactions on FT synthesis selectivity. Composite Fe-Mn/ Y-zeolite and Fe-Mn/Si02 catalysts (501 K , 2500 kPa, H2/C0 = 2.1, 45-47%.5% CO conversion). (a) Carbon number distribution; (b) a-olefin to n-paraffin ratio.


295

chain termination steps, in the control of FT synthesis selectivity. Fe catalysts, however, require more detailed selectivity models than Co and Ru catalysts because FT synthesis products contain significant amounts of internal olefins, isomeric hydrocarbons, and oxygenates. Moreover, secondary reactions on Fe catalysts can be affected by the presence or formation of several oxide and carbide phases during catalysis and by the ineffective dispersal of metal oxides or alkali promoters on their surface. The complexity and nonuniformity of catalytic surfaces and FT synthesis products present modeling and characterization challenges that we shall discuss in future work. VI. Conclusions

Chain growth during the Fischer-Tropsch synthesis is controlled by surface polymerization kinetics that place severe restrictions on our ability to alter the resulting carbon number distribution. Intrinsic chain growth kinetics are not influenced strongly by the identity of the support or by the size of the metal crystallites in supported Co and Ru catalysts. Transport-limited reactant arival and product removal, however, depend on support and metal site density and affect the relative rates of primary and secondary reactions and the FT synthesis selectivity. Diffusion-limited removal of products from catalyst pellets leads to enhanced readsorption and chain initiation by reactive a-olefins. These secondary reactions reverse chain termination steps that form these olefins and lead to heavier products, higher chain growth probabilities, and more paraffinic products. Diffusion-enhanced readsorption of a-olefins accounts for the non-Flory carbon number distributions frequently observed during FT synthesis on Co and Ru catalysts. Diffusion-limited reactant (HdCO) arrival leads instead to lower selectivity to higher hydrocarbons. Consequently, intermediate levels of transport restrictions lead to highest selectivities to Cs+ products. A structural parameter x, containing the pellet diameter, the average pore size, and the density of metal sites within pellets, determines the severity of transport restrictions and the FT synthesis selectivity on supported Ru and Co catalysts. Diffusive and convective transport processes introduce flexibility in the design of catalyst pellets and in the control of FT synthesis selectivity. Transport restrictions lead to the observed effects of pellet size, site density, bed residence time, and hydrocarbon chain size on chain growth probability and olefin content. The restricted removal of reactive olefins also allows the introduction of other intrapellet catalytic functions that convert olefins to other valuable products by exploiting high intrapellet olefin fugacities. Our proposed model also describes the catalytic behavior of more complex Fe-

296

ENRIQUE IGLESIA et

al.

based materials, where several chain termination steps and highly nonuniform and dynamic surfaces introduce additional details into the models required to describe FT synthesis selectivity models. VII.

Nomenclature

Parameter in Eq. (26) Surface area per unit volume of catalyst particle Olefin and paraffin concentration Surface concentration of growing hydrocarbon chain Effective diffusivity of n-olefins and n-praffins Reference diffusivity defined by Eq. (26) Dimensionless effective diffusivity of n-olefins Fraction of carbon atoms contained within chains with n-carbons Hydroformylation rate constant Cracking rate constant Hydrogenation rate constant CO hydrogenation rate constant Hydrogen abstraction rate constant Propagation rate constant Readsorption rate constant Secondary olefin hydrogenation rate constant Adsorption constant Reaction orders for CO hydrogenation reactions Carbon number Avogadro's number Peclet number Total inlet pressure Universal gas constant Particle-averaged rate of CO consumption Particle-averaged rate of H:, consumption CO hydrogenation rate Catalyst pellet radius Mean pore radius Termination rate Propagation rate Dimensionless rate of CO consumption Dimensionless rate of H2 consumption Local surface concentration of growing n-chains Total active site concentration of catalyst Absolute temperature


297

Dimensionless CO concentration Dimensionless inlet CO concentration Dimensionless concentration of growing n-chains Dimensionless H2 concentration Dimensionless inlet H2 concentration Dimensionless n-olefin concentration Dimensionless n-olefin concentration at reactor’s inlet Dimensionless n-paraffin concentration Dimensionless n-paraffin concentration at reactor’s inlet

Greek Symbols Chain growth probability Asymptotic chain growth probability Probability of hydroformylation of a-olefins Probability of cracking of a-olefins Reference cracking probability Probability of chain termination to n-paraffins Probability of chain termination to a-olefins Probability of readsorption of a-olefins Probability of secondary hydrogenation of a-olefins Total chain termination probability Asymptotic chain termination probability H K O effective transport ratio Catalyst site density Dimensionless reactor position Dimensionless pellet position Effectiveness factor Catalyst pellet density Mole fraction of n-hydrocarbons Pellet porosity Thiele modulus of n-hydrocarbons Reference Thiele modulus Structural parameter defined in Eqs. (15) and (25) Diffusivity/reactivityparameter defined in Eq. (15) Diffusivity/reactivityparameter defined in Eq . (25) ACKNOWLEDGMENTS We thank Dr. Rocco A . Fiato and Dr. Chang J. Kim for many discussions about the concepts described in this paper and for their many independent contributions to our understanding of Fischer-Tropsch synthesis chemistry. We also acknowledge the technical assistance of Mrs. Hilda Wroman, Joseph E. Baumgartner, Bruce DeRites, and Craig Womble in the synthesis

298

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at.

and catalytic evaluation of the studies reported here. The patient and expert assistance of Mrs. Patricia A. Kocian in the typing of the many drafts of this manuscript is gratefully acknowledged. REFERENCES 1 . Pichler, H., Adv. Cuful. 4, 271 (1952). 2 . Storch, H. H., Golumbic, N., and Anderson, R. B., “The Fischer-Tropsch and Related Syntheses.” Wiley, New York, 1951; Anderson, R. B., “The Fischer-Tropsch Synthesis.” Wiley, New York, 1984. 3 . Vannice, M. A , , in “Catalysis Science and Technology” (J. R. Anderson and M. Boudart, eds.), p. 139. Springer-Verlag. Berlin, 1982; Vannice, M. A , , Catal. Rev. Sci. Eng. 3, 153 (1976). 4 . Iglesia, E., Reyes, S. C., and Madon, R. J., J . Curd. 129, 238 (1991). 5 . Iglesia, E., Reyes, S. C., and Soled, S. L., in “Computer-Aided Design of Catalysts and Reactors” (C. J. Pereira and R. W. Becker, eds.). Dekker, New York, 1993, in press. 6. Rostrup-Nielsen, J., in “Catalysis Science and Technology” (J. R. Anderson and M. Boudart, eds.), Vol. 5, p.1. Springer-Verlag, Berlin, 1984. 7. Biloen, P., Helle, J. N . , and Sachtler, W. M. H., J . Cutul. 58, 95 (1979). 8. Brady, R. C., and Pettit, R., J. Am. Chem. Soc. 102,6128 (1980). 9 . Brady, R. C., and Pettit, R., J. Am. Chem. SOC. 103, 1287 (1981). 10. Fischer, F., and Tropsch, H., Brennsr.-Chem. 7, 97 (1926). 11. Sachtler, J. W. A,, Kool, J. M., and Ponec, V., J. Cuful. 56, 284 (1979). 12. Horiuti, J., and Polanyi, M., Trans. Furuduy SOC. 30, 1164 (1934). 13. Pines, H., “The Chemistry of Catalytic Hydrocarbon Conversions,” p. 282. Academic Press, New York, 1981. 14. Madon, R. J., Reyes, S. C., and Iglesia, E., J. Phys. Chem. 95, 7795 (1991). 15. Smith, D. R . , Hawk, C. O., and Golden, P. L., J . Am. Chem. SOC. 52, 3221 (1930). 16. Craxford, S. R., Trans. Furuduy SOC. 64, 441 (1939). 17. Herington, E. F. G., Chem. Ind. (N. Y.) 65, 346 (1946). 18. Pichler, H., Schulz, H., and Elstner, M., Brennsr.-Chem. 48, 78 (1967). 19. Kolbel, H. and Ruschenburg, E . , Brennst.-Chem. 35, 161 (1954). 20. Golovina, 0. A., Sakharov, M. M., Roginskii, S. Z., and Dokukina, E. S., Russ. J. Phys. Chem. 33, 471 (1959). 21. Roginskii, S. Z., Proc. Int. Congr. Caral. 3rd, Amsterdam. 1964 p. 939 (1965). 2 2 . Hall, W. K., Kokes, R. J., and Emmett, P. H., J. Am. Chem. Soc. 82, 1027 (1960). 23. Schulz, H., Rao, B. R., and Elstner, M., Erdol Kohle 23, 651 (1970). 2 4 . Novak, S., Madon, R. J., and Suhl, H., J . Catul. 77, 141 (1982). 25. Smith, D. F., Hawk, C. O., and Golden, P. L., J. Am. Chem. SOC. 52, 3221 (1930). 2 6 . Ekerdt, J. G., and Bell, A. T., J . Curul. 58, 170 (1979); 62, 19 (1980). 27. Eidus, Y., Russ. Chem. Rev. 5 , 388 (1967). 28. Kobori, Y., Yamasaki, H., Naito, S.,Onishi, T., and Tamaru, K., J.C.S.Furaday 1 78, 1473 (1982). 29. Percy, L. T., and Walter, R. J., J. Card. 121, 228 (1990). 30. Friedel, R. A., and Anderson, R. B., J. Am. Chem. SOC. 72, 1212 (1950); Anderson, R. B., J. Curd. 55, 114 (1978). 31. Iglesia, E., and Madon, R. J., U. S. Pat. 4,754,092 to Exxon Res. Eng. Co. (1988). 32. Anderson, R. B., in “Catalysis” (P. H. Emmett, ed.), Vol. 4, p. 257. Rheinhold, New York, 1956.


299

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Catalysis by Metal Ions Intercalated in Layer Lattice Silicates YUTAKA MORIKAWA Research Laboratory of Resources Utilization Tokyo Institute of Technology Yokohama 227, Japan

1.

Introduction

Layer lattice silicates, termed phyllosilicates in mineralogy, are commonly found in the Earth's crust. Micas occur as large crystalline forms in rocks and clay minerals are important components of soils. The layer structure consists of the silicate sheets stacked one above the other, sometimes incorporating interlamellar metal ions or hydrated metal ions. Because the stacking is retained weakly by van der Waals forces, hydrogen bonding, and/or electrostatic forces, the stacks are easily cleaved as observed with micas, and have a pronounced capacity for swelling by intercalating polar molecules into the interlayer spaces, as reported by many clay mineralogists ( I ) . The interlayer cations in a swelling layer lattice silicate, usually Na', are exchangeable with any desired cations, even bulky organic or metal chelate cations. These properties play an important role in the fertility of soils and in the stockpiling of water and plant foods, and have therefore been a subject of intensive study in pedology and mineralogy. Acid-treated clay minerals were employed as cracking catalysts in the first commercial process, the Houdry process, widely used in the early petroleum industries to produce high-octane gasoline. The Houdry process catalysts had been discussed extensively by many investigators (2) but were eventually completely replaced by synthetic silica-alumina or zeolite catalysts. Recently, the need for new catalytic materials has revived special interest in the layer lattice silicates because of their ion-exchange properties and their expandable layer structures. The catalysts preparable with swelling layer lattice silicates are classified into four types, as illustrated in Fig. 1. The intercalate of hydrated metal ion (a, Fig. 1) is easily obtained by a simple ion-exchange reaction in an aqueous medium. The intercalate acts as a Br~nstedacid catalyst because 303 Copyright 0 1993 by Academic Press, Inc. All rights of reproduction in any form reserved.

304

YUTAKA MONKAWA

the water molecule coordinated to the metal ion dissociates to give a proton according to the electron withdrawal of the metal ion and the electrostatic field between the negatively charged silicate sheets. Thomas (3) and colleagues conducted a number of reactions using metal ion-exchanged smectites and demonstrated the predominant activity of the intercalates over mineral acids. A cationic metal complex is immobilized between the silicate sheets with negative charges (b, Fig. 1). Pinnavaia and co-workers (4) investigated the catalytic activities of Rh complex intercalates and showed that the catalytic activity of the intercalate is comparable to that of the homogeneous solution containing the same amount of the metal complex when reactant molecules are small enough to diffuse into the interlayers. The selectivity observed with the immobilized catalyst is sometimes different from that in homogeneous systems ( 4 3 ) because of the acidic nature of the interlayer spaces and a steric factor determined by the shape of the reactant molecule and the interlayer spacing between the silicate sheets. A pillared clay (c, Fig. 1) has lately attracted considerable attention as a pore-size adjustable molecular sieve. The intercalate of oligomeric metal hydroxide cations is dehydrated by calcination to give a porous material in

L

Layer lattice -silicate

-

I

F

D

M"+(H~O),

Na+

"u I +t

-

(a) B acid catalyst

-

[M(HZO)m] n+

,J

Calcination

"I. (b) Immobilized metal complex

( c ) Pillared clay

(d) Immobilized metal ion catalyst

FIG. 1. Catalysts preparable with layer lattice silicates.

METAL IONS IN LAYER LAlTICE SILICATES

305

which even-grained metal oxides act as pillars to keep the silicate sheets apart. The alumina-pillared smectite was prepared for the first time by Brindley and Sempels (6)and independently by Lahav et af. (7); extensive studies have since extended the pillar substance to Ni (8),Zr (9),Si (ZO), Fe (ZI),Cr (12), and Ti (13) oxides. Decomposition of the metal complex intercalate in inert gas liberates the coordinated molecules and forms the nude metal ions confined in the interlayer region. The metal ion is surrounded by oxide ions of the silicate sheet, but the coordinative bonding of oxygen is unusually weak considering the geometry of the oxide ions. It is therefore expected that the material (d, Fig. 1) shows the characteristic activity of the metal ion when the silicate sheet is catalytically inactive. Some findings obtained in recent years using a catalyst of this type are introduced herein.

II. Catalytic Materials

Among various 2: 1-type layer lattice silicates, of which the silicate sheet is composed of two tetrahedral silicate layers and one octahedral metal oxide layer, fluorotetrasilicic mica (TSM) shows an unusual property; that is, the silicate sheet of TSM has no acidity and can be said to be catalytically inactive. TSM is a swelling mica with the chemical formula Na [Mg2.,Si4Ol0F2].It was synthesized by Kitajima and Daimon (14)and the layer structure illustrated in Fig. 2 was crystallographically established by Toraya et al. (15). Table I summarizes the acid amounts of silicate minerals determined by temperature-programmed desorption of ammonia (16).Hecis a synthesized mineral and bentorite (Hect), Lil/3[Mg8/,Lil/sSi4010F2], is a naturally occurring clay tonite (Bent), Nal/3[A15/~Mg1/3Si4010(OH)2], mineral. The layer structures of these minerals are similar to that of TSM. Table I demonstrates that acid sites exist more or less on most silicate minerals and even silica gel is frequently used as an inert carrier, but not on this particular mica. The interlayer Na+ ion was exchanged with an aqua or ammine complex ion by a simple ion-exchange method. A dilute solution of metal salt or ammine complex was added to a 0.1-0.3 wt% aqueous sol of TSM with vigorous stirring. The precipitate was readily formed but the stirring was further continued for 12 h to complete the exchange reaction. In most cases, 60% of the interlayer Na' ions was replaced by the addition of slightly larger amounts of the metal ion than the cation-exchange capacity of TSM (calculated from the chemical formula to be 2.54 mEq/g). The precipitate was washed well with distilled water, dried in an oven at lOO"C, and heat treated at 350-400°C in inert gas to obtain a metal ion-exchanged form of TSM

306

YUTAKA MORIKAWA

v

Exchangeable metal ion

FIG. 2. Oxygen network of fluorotetrasilisic mica: 0, oxygen; 0 , fluorine.

(M"+-TSM). The basal spacing of IW-TSM has been determined to be 9.6 A regardless of the interlayer metal ion M"'. The metal ion-exchanged forms of the other layer lattice silicates were prepared in the same manner. 111.

Catalytic Activity of Metal Ion-Exchanged Fluorotetrasilicic Mica

Methanol conversion was adopted as a probe reaction to explore the catalytic activity of M"+-TSM, because methanol is a simple molecule that transforms into easily assignable compounds and is converted into different products through different routes employing different catalysts. Methanol is decomposed into carbon monoxide and hydrogen over metal catalysts [including Nj ( I 7 ) ] ,is dehydrogenated in to formaldehyde or methyl formate over Zn- or Cu-containing catalysts (18), and is dehydrated into dimethyl ether and successively into hydrocarbons over acid catalysts (19). Table I1 briefly summarizes the results of methanol conversion conducted in a pulse reactor using a variety of M"+-TSMs as catalysts (20). Each sam-

307

METAL IONS IN LAYER LAlTICE SILICATES TABLE I Ammonia Desorption and Peak Temperatures Observed in TPD Surface area

Peak temperature

Mineral

(m2/g)

("C)

(mol/g)

(nmol/m2)

TSM Bent Hect Silica gel Zeolite (1 3X)

9.1 38. I 14.2 347 575

-

0 2.3 11 0.54 140

0 60 770 1.6 240

296 292 314 316

Amount of ammonia desorbed

ple was pretreated at 400°C in a stream of helium for 2 h. As seen in Table 11, 20 kinds of I@-TSMs can be classified in five groups. Na+-, Mg2+-, Ca2+-,Mn2+-,Co2+-,Zn2+-,La'+-, and Th4+-TSMsexhibit no activities at reaction temperatures below 400"C, suggesting that these metal ions and the silicate sheets of TSM are inactive Rh3+-, Pd2+-,Ru3+-, Ni2+-, Ce3+-,and Ir3+-TSMs catalyze decomposition to CO 2H2 and subsequently the hy-

+

TABLE I1 Catalytic Activity of Metal Ion-Exchanged Fluorotetrasilicic Mica for the Reaction of Methanol"

Metal ion Rh3+ Pd2+ Ru3+ Ni2+b Ce3 + b

Conversion of CHsOH (%)

Product Major CO, CH4

50.3 72.4 64.I 75.9 21.6 25.7

co co co co co

Cr3+ Fe3+

10.7 6.5 9.8 7.7

CH4 CH4 CH4 CH,

cu2+c

53.6

HCOOCHs

Ti4+

44.8

1r3+b

Sn4+ Ag+

Minor

Na+-, Mg*+-, Ca2+-, Mn2+-, C o 2 + - ,Zn2+-,La3+-, and Th4+-TSMs are inactive. (Reaction temperature, pulse size: unspecified (300°C, 0.2 p l ) . 'Reaction temperature, pulse size: 400°C, 0.2 p i . Reaction temperature, pulse size: 300"C, 1 p I .

308

YUTAKA MORIKAWA

drogenation of carbon monoxide to form hydrocarbons. Although no metal particles are observed in the catalyst samples by means of X-ray diffraction (XRD) analysis, most of the noble metal ions are possibly reduced to zerovalent metals, leaving protons in the interlayer region as countercations of silicate sheets. The catalytic activities of Sn4+-, Ag+-, Cr3+-, and Fe3+TSMs are relatively low and seem rather unusual. The catalysis produces mainly methane, accompanied by no carbon monoxide and no dimethyl ether, which is completely different from both the preceding reaction and acid-catalyzed reactions, in which the formation of dimethyl ether is predominant at such low conversions (Table 11). Supplemental studies with a flow reaction system suggest that methanol is dehydrogenated to methyl formate and decarboxylated subsequently to form CH4 + C 0 2 over Sn4+-and Fe3+-TSMs. It is of particular interest that the catalytic activities of Cu2+and Ti4+-TSMsare considerably high and quite distinct from the others and from each other. Among a variety of M"+-TSMs, only Cu2+-TSMcatalyzes the dehydrogenation to form methyl formate. Ti4+-TSMacts as an acid catalyst, promoting dehydration to form dimethyl ether selectively, although the silicate sheet of TSM has no acid sites. The correlation between the catalytic activity of a metal ion or metal salt and the electronegativity of a metal ion has been discussed for a number of acid-catalyzed reactions (21), including the isomerization of butenes (22). The catalytic activity of metal sulfates for cis-butene isomerization increases with the electronegativity of the metal ion in accordance with the general tendency toward surface acidity induced by metal ions. The catalytic activities of M"+-TSMs for butene isomerization are, however, not subject to such correlation (23), although M"+-TSM is regarded as a metal salt and consists of metal ions and negatively charged silicate sheets. No reaction takes place over Na+-, Mn2+-,and Pb2+-TSMsand the isomerization of 1-butene over Mg2+-,Co2+-,Ti4+-,Ni2+-,and Cu2+-TSMsis accelerated remarkably by the addition of hydrogen into the reaction system. The acceleration effect of hydrogen is observed for the olefin isomerization, which is promoted by metal chalcogenides (24) or metal complexes (25) through the formation of half-hydrogenated species bound to the metal entities. A survey of the reactions of methanol and butene over a series of M"+TSMs leads us to note two common features. First, no activity is observed with some M"+-TSMs, indicating that the silicate sheet of TSM is catalytically inactive. Second, M"+-TSM does not follow the clear activity sequence that is frequently noted in acid catalysts, of which the surface acidity is changed by the inductive effect of metal ions (21, 22). These features strongly suggest that the direct interaction between a substrate and the interlayer metal ion is responsible for the catalysis by M"'-TSM, namely, M"+TSM acts as an immobilized metal ion catalyst.

METAL IONS IN LAYER LATTICE SILICATES

IV.

309

Dehydrogenation of Methanol over Copper Ion-Exchanged TSM

As described above, Cu2+-TSM is completely different from other M"+TSMs in its catalytic activity. We prepared Cu2+-exchanged forms of the silicate minerals listed in Table I by decomposing their Cu(I1) ammine complex-exchanged derivatives at 400°C in a nitrogen stream and investigated the catalytic activities for methanol conversion using a conventional fixedbed flow reactor (26). Table I11 shows comparative results of the reaction. Cu2+-TSM catalyzes the conversion of methanol to methyl formate selectively. The selectivity attains 100% even at the high conversion, close to the thermodynamic limit (55%). As seen in Table 111, the selectivity for methyl formate formation is directly affected by the characteristics of silicate minerals. The acid sites on the silicate sheets of Bent and Hect necessarily catalyze the dehydration of methanol to dimethyl ether, causing the lower selectivity of Cu2+-Bentand Cu2+-Hectfor methyl formate formation. Zeolite is well known as an acid catalyst. The Cu2+-exchanged form of molecular sieve 13X (Cu2+-Zeol)promotes the dehydration of methanol preferentially. The results with Cu+- and CuO-TSMs (27) are also shown in Table 111. Cu+-TSM was obtained similarly to Cu2+-TSM using the Cu(1) ammine complex salt in the ion-exchange reaction. Cuo-TSM, copper supported on TSM, was prepared by calcination and subsequent hydrogen reduction of an oligomeric cupric hydroxide intercalate, which had been obtained by precipitating the hydroxide into the interlayer spaces of TSM after the titration method developed by Yamanaka and Brindley (8). The catalytic activity of Cu+-TSM is so poor as to be less than one-tenth of that of Cu2+-TSM, although the molar content of Cu+ is about twice as high. Because a weak sigTABLE Ill Methanol Conversion over Copper Ion-Exchanged Silicate Minerals" Conversion Catalyst

1%)

Cu2+-TSM Cuz+-Bent Cu2+-Hect Cu2+-Zeal

45 6 34 44

Cu+-TSM Cuo-TSM

4.3 1.6

Selectivity (%) HCOOCH, 100 27 16

0 100

0

HCHO

CH,OCH,

0 5 0 0

0 68 100

0 100

0 0

84

Caltalyst (0.7 g) was exposed to a stream (58 m l h i n ) of a 1.2 mixture of methanol and nitrogen at 240°C.

310

YUTAKA MORIKAWA

nal assigned to interlayer Cu2+ appears on the electron spin resonance (ESR) spectrum of Cu'-TSM, the low but selective activity for methyl formate formation is ascribed to the small amount of Cu2+ contained in Cu+-TSM as an impurity, and Cu+ is inactive for the methanol conversion. Cuo-TSM catalyzes selectively the dehydrogenation of methanol to formaldehyde. The selectivity is again as high as 100%. Because the thermodynamic equilibrium is unfavorable for formaldehyde formation at the reaction temperature, the low conversion does not mean the low activity of Cue-TSM. The value of conversion corresponds to 20% or higher against the equilibrium conversion (-8%). The results demonstrate that the catalytic activity of copper for the methanol conversion is completely different depending on the oxidation state; Cu+ is inactive and Cu2+and Cuo catalyze, respectively, the dehydrogenative coupling to methyl formate and the dehydrogenation to formaldehyde. Figure 3 shows the time course of conversion and selectivity in the methanol conversion at 200°C over Cu2+-TSMand Cu2+-Si02(25). Cu2+SiOz was prepared by a conventional ion-exchange method using a solution of Cu(I1) ammine complex ion and a suspension of silica gel. Both of the catalysts were treated under the same conditions prior to the reaction. The catalytic activity of Cu2+-TSMdoes not change at all, retaining 100% selectivity for methyl formate formation. The stable activity is confirmed in the reaction, which is conducted at 240°C for 30 h. The ESR signal observed

r

100 h

6p

v 3

0

c a c

'

tl

lu

50

o ,0

;

CU~+-TSM

A ,A

;

Cu2+-SiO2

c 0

.3

VI CI

P)

E:

V

0 0

2

4 6 Reaction time ( h )

FIG.3. Change of the catalytic activity of CuZ+-TSMand Cu2'-Si02 with reaction time. Reaction temperature, 200°C; space velocity, 96 ml/min . g for Cu2+-TSM, 301 ml/min . g for Cu2+-Si02;CH,OH concentration, 33%.


31 1

for fresh Cu"-TSM agrees well with that of interlayer Cu2+as assigned by McBride et al. (28),and is perfectly reproduced on the spectrum of the used catalyst. In contrast, the activity of Cu2'-Si02 increases with reaction time. The selectivity for methyl formate formation falls to 80% in 5 h from the initial high value (estimated to be 100% by extrapolation), while carbon monoxide forms in an increasing amount. The observed change in the catalytic activity is speculated to be caused by the reduction of Cu2+during the reaction, because the stepped decrease in Cu2+ signal intensity has been recorded in ESR studies with the Cu2+-Si02sample alternately exposed to methanol and evacuated at the reaction temperature. The high reduction resistivity of Cu2+in the interlayers of TSM likely results from the chemical stability of counteranions, that is, negatively charged silicate sheets, and causes the highly stable catalytic activity of Cu2+-TSM. Methyl formate is expected to be one of the intermediates of the methanol-based industries that produce dimethylformamide, acetic acid, pure hydrogen, carbon monoxide, etc. A number of patented catalysts (18) are therefore composed of copper oxide as a main ingredient and include a variety of metal oxide additives. In our comparative studies, some of these catalysts exhibit fairly good activity but do not exceed Cu2+-TSMin activity and selectivity (26). The results of the reactions over these patented catalysts, which have relatively high performance levels, are illustrated in Fig. 4 together with the result obtained with Cu2+-TSM;the yield of methyl for-

equilibrium conversion

h

*

v

60

a,

e,

m

E 0

rc

40

3

x c e, a,

E k

0

20

'

P

3

a, .d

>

Cu-Zr-Zn-0

I

I

I

312

W T A K A MORIKAWA

mate is plotted against the reaction temperature. CuO Cr203/Si02 appears rather desirable because of the high activity at a lower temperature range. However, the low yield at higher temperatures is a serious problem considering that the thermodynamic equilibrium favors the higher temperature reaction. Cu-Zr-Zn-0 is one of the most selective catalysts but has a lower activity than Cu2+-TSM. The highest yield has been obtained with Cu2+TSM. The selectivity of Cu2+-TSMfor methyl formate formation is so high that it remain at 100%up to the equilibrium conversion (29). This outstanding activity has been realized simply by immobilizing Cu2+in the interlayers of TSM. This suggests that the additives in patented and practical catalysts play a role in stabilizing the oxidation state of Cu2+ under the reducing conditions of the reaction. V.

Reaction of Methanol over Ti4+-TSMand Related Catalysts

A. METHANOL CONVERSION INTO HYDROCARBONS As seen in Table 11, only Ti4+-TSMacts as an acid catalyst to promote the dehydration of methanol to dimethyl ether. This particular acidity seems not to be generated by the inductive effect of metal ions, because Cr3+ and Fe3+ ions-exchanged TSMs, ions of sufficiently high electronegativity, exhibit no acidic characteristics in the methanol conversion. Although the origin of the acidity is not fully understood at this stage of investigation, the acidic properties of the other layer lattice silicates are expected to be modified largely by exchanging the interlayer metal ion with Ti4+. The acidity of proton- or Ti4+-exchanged TSM, Hect, and Bent was measured by TPD of ammonia fully adsorbed on the samples at 200°C (30).The proton-exchanged form of each mineral was prepared by passing the aqueous sol of silicate mineral through a column packed with an acidic ionexchange resin. The sample was treated in a helium stream at 400°C for 2 h prior to the ammonia adsorption. The spectra obtained are illustrated in Fig. 5 , where the results with the original mineral forms are also shown for comparison. It is of interest that no ammonia desorption peak is observed even with proton-exchanged TSM (H+-TSM). A small desorption peak appears with the original form of Hect (Li+-Hect), and the amount of desorbed ammonia does not increase appreciably by changing it into the proton-exchanged form (H+-Hect). In contrast, the large peaks observed in the TPD spectra with Ti4+-TSMand Ti4+-Hect show that considerable numbers of acid sites are generated by exchanging the interlayer Na+ or Li+ ion with

313

ME3AL IONS IN LAYER LATTICE SILICATES

h

2

E

" a10 l C

0 P YI

m

k

L.

a, P

u 0 a,

rr

200

300

400

500

200

300 400 500 Temperature ( "C )

200

300

400

500

600

FIG. 5 . TPD spectra of ammonia adsorbed on various layer lattice silicates.

the Ti4+ ion. The amount of ammonia desorbed from Ti4+-TSMin the temperature range of 200-400°C is estimated to be one-tenth of the cation-exchange capacity of TSM, suggesting that not all of the Ti4+ in Ti4+-TSMis effective in ammonia adsorption. The TPD spectra for the three forms of Bent present complicated curves. The spectra, however, clearly show that the acid site population increases appreciably by changing Na+-Bent into the proton form (H+-Bent) and increases dramatically by changing it into the Ti4+ form (Ti4+-Bent). Table IV summarizes the results of methanol conversion over the catalyst samples employed in the TPD study (30). The reaction was conducted in a conventional fixed-bed flow reactor under the conditions given in the table. The results are in agreement with those of the TPD measurement. Na+- and H+-TSMs are inactive for the methanol conversion, whereas Ti4'-TSM promotes dehydration, converting 50% of the fed methanol into dimethyl ether and a small amount of methane. The negligible activity of Li+-Hect is improved slightly by exchanging the Li+ ion with H+ and dramatically by exchanging Li+ with Ti4+. Na+-Bent is an acidic clay. All of the three Bent catalysts, even Na+-Bent, show higher activity than Ti4+-TSM,and the hydrocarbon yield reflects this difference in catalytic activity. Na+-Bent is sufficiently active to give 60% conversion but has no ability subsequently to dehydrate dimethyl ether into hydrocarbons. The activity of H+-Bent is higher than that of Na+-Bent, but the hydrocarbon yield is as low as 9%. As expected from the results of TPD measurement, the activity of Ti4-Bent is remarkably high and converts 60% of fed methanol into hydrocarbons that are a mixture of methane, CM olefins, and a small amount of C6 hydrocarbons.

314

YUTAKA MORIKAWA TABLE IV Methanol Conversion over Various Laver Lattice Silicates" Conversion of methanol

Yield of hydrocarbons

Catalyst"

Surface area' (mW

(%)

(%I

Na+-TSM H+-TSM Ti4 -TSM

18 71 160

0 0 50.8

0 0 14.2

Li.' -Hect H' -Hect Ti4 -Hect

23 96 166

0 15. I 90.4

0 0.6 38.0

Na+-Bent H+-Bent Ti4-Bent

47 92 24 I

59.8 84.1 93.2

0 8.7 58.6

H' -Ben t* Ti4+-Bent*

260 245

92.5 95.2

5.6 89.5

+

+

* Catalyst, 1 g; total flow rate (15% methanol in nitrogen), 35 ml/min; reaction temperature, 350°C. The asterisks denote acid treatment. ' The value for the original form of each mineral is different from that in Table I, because of a dependency on the particle size of mineral collected and the aggregated form of silicate sheets.

'

We also studied the effect of ion exchange with Ti4+on the catalytic activity of acid-treated Bent (H+-Bent*), sometimes called "activated clay." The results are given in Table IV. H+-Bent* is virtually the same as H+Bent in catalytic activity. However, the catalytic activity of Ti4+-Bent*for methanol conversion to hydrocarbons is much higher than that of Ti4+-Bent. The hydrocarbon yield reaches 90%, and the products, in addition to methane, are primarily olefins lower than Cg . The selectivity for olefin formation is estimated to be 90% or higher based on Cz and C3 hydrocarbon product distribution. Ti4+-Bent* appears to surpass the phosphorus compound-modified zeolite proposed by Kaeding and Butter (31) in selective activity for olefin formation, and has the potential to exceed H-Fe-silicate (32) and Ni-SAPO-34 (33), proposed recently by Inui et al. As seen in Table IV, the specific surface area of each layered mineral increases by ion exchange with H+ and Ti4+.However, the observed increase in catalytic activity is not explainable by the increase in the surface area. Ti4+-Hecthas about twice the surface area of H+-Hect and shows more than six times the activity of H+-Hect. The ion exchange of H+-Bent* with Ti4+ does not change the surface area but increases dramatically the catalytic activity for hydrocarbon production. The XRD pattern of Ti4+-exchanged


315

derivatives as pretreated showed no diffraction lines except for those of layered minerals; there was a shift of the (001) diffraction line from the position observed with its original forms to a lower angle. The shift corresponds to the expansion of basal spacing from 9.6 to 12.1 A, suggesting that nonvolatile entities, possibly Ti-0 species, are deposited in the interlayer region, keeping the silicate sheets apart by 2.5 A. The expanded interlayer spacing is not wide enough for methanol to get in, and is not always observed in the XRD studies with Ti4+-TSM. In addition, methanol conversion over TiO2 produces no hydrocarbons, but only dimethyl ether, even at 100% conversion (34). It is therefore speculated that the enhancement of acidity does not result from the deposition of TiO2. Although extensive studies are needed on the characterization of acid sites on the Ti4+-exchanged layer lattice silicate, ion exchange with Ti4+is very likely widely applicable to the modification of solid acid catalysts. OF METHANOL B . NOVELREACTIONS

-

The catalytic activity of heteroion-exchanged TSM, Ti4+ Zn2+-TSM, is different from the activities of Ti4+- and Zn2+-TSMs. The results of the methanol conversion over the catalysts (35, 36) are summarized in Table V, which includes the data for Ti4-TSM from Table IV. The reaction conditions are the same as given in Table IV. The heteroion-exchange reaction was conducted using a mixed solution of Ti(1V) and Zn(I1) chlorides (mole ratio = 9: 1). The resultant precipitate was washed with distilled water repeatedly and quickly to obtain Ti4+ .Zn2+-TSMand Ti4+ -Zn2+-TSM/CI,respectively. Ti4+-TSM catalyzes the dehydration of methanol to give dimethyl ether and a small amount of hydrocarbons, mainly methane, as described in the preceding section. The catalytic activity of Ti4+ * Zn2+TSM is less than one-sixth as low as that of Ti4+-TSM,although only onetenth of the Ti4+ in Ti4+-TSMhas been replaced with Zn2+,inactive for the TABLE V Methanol Conversion Over Ti4+ andlor Znz+Ion-Exchanged TSMs Conversion

Selectivity (%), ~

Catalyst Ti4+-TSM Zn*+-TSM Ti4+ . Zn2+-TSM Ti4+ . Zn2+-TSM/C1 a

(%I 50.8 0 8.3 60.1

DME

CH4

MF

85.8

14.2b

-

-

-

-

71.3 43.8

25.8 13.9

0.7 17.2

DMM 1.3 0.5

DME, Dimethyl ether; MF, methyl formate; DMM, dimethoxymethane. Including a small amount of Cz and C3 hydrocarbons.

C?+-O 0.9 21.3

COz -

0 3.3

316

YUTAKA MORIKAWA

methanol conversion at the reaction temperature. Ti4+ * Zn2+-TSMcatalyzes the dehydration to dimethyl ether and the dehydrogenation simultaneously to give methyl formate and dimethoxymethane. It is noteworthy that the reaction produces a small amount of various oxygenated compounds, including ketones, esters, and aldehydes, which are collectively included in Table V as C3+-0. Ti4+ Zn2+-TSM/Cl,containing -2 wt% C1- ion as prepared, is much more active than the Cl-free catalyst. The selectivity values show that the presence of C1- accelerates the dehydrogenation of methanol and increases the selectivity for C3+-0 formation by 20%. The product distribution in C3+-0 formed over Ti4+ * Zn*+-TSM/Cl is shown in Table VI, which gives the molecular formulas of unidentified products as determined by mass spectrometry. As seen in Table VI, couples of vinyl and its hydrogenated compounds are formed, and predominant products are branched compounds that have two methyl groups at the a-carbon. Considering the product distribution and the dehydrogenation activity of the catalysts, it is speculated that the oxygenated compounds are formed by the vinylation through aldol condensation with formaldehyde and by the subsequent hydrogenation. Methyl acetate 1 is vinylated with formaldehyde to form methyl acrylate 3. Methyl acrylate is hydrogenated to methyl propionate 4 and subsequently vinylated and hydrogenated to methyl isobutyrate 8. The reaction is analogous to the vinylation of active methyl or methylene groups with methanol, as reported by Ueda et al. (37) using metal ion-containing magnesia catalysts. Although the detailed mechanism of the reaction is not well understood, especially on the first C-C bond formation, the catalyst for G + - O production should have both dehydration and the dehydrogenation activities. Because Cu2+-TSMis excellent catalyst for the dehydrogenation of methanol, Ti4+ * Cu2+-TSM was prepared and tested for the methanol conversion. Against our expectations, the catalytic activity of Ti4+ * Cu2+-TSMwas similar to that of the mechanical mixture of Ti4+-TSM and Cu2+-TSM. TABLE VI Product Distribution in Cs-,-Oxygenated Compounds Formed over Ti4+ . Zn2+-TSM Product

/

(%)

0 1 2.8

(%)

43.1

7

)( 0 2 2.7

8 21.3

eo\ r\fo\ 0

0

0

3 2.9

4 4.4

5 0.8

1.8

9 4.0

3.5

0 6 0.5

1.4

CqH,O, 3.6

10 7.3

317


Methanol conversion to oxygenated compounds has not yet been reported, but has a potential for practical application to produce chemicals from methanol. Further studies are now centered on the characterization of this complicated catalyst to clarify the configuration of bifunctional active sites and to improve selectivity for a desired product. It is predicted that, in the methanol conversion over Ti4+ Zn2+-TSM/C1, methyl vinyl ketone 5 (MVK) and methyl ethyl ketone 6 (MEK) are formed from acetone 2 by the vinylation and the subsequent hydrogenation, respectively, and diisopropyl ketone 10 is the end-product of the successive vinylation. In order to confirm the prediction, the reaction of acetone and methanol was carried out at the same temperature. The reaction in the presence of oxygen was also carried out because the dehydrogenation of methanol to formaldehyde must be a key step in the vinylation of acetone if the prediction is right, and is favored thermodynamically by an oxidizing condition. The results (35,38) are summarized in Table VII, where the values of conversion and selectivity are calculated on the acetone basis. Zn2+TSM showed no activity. MVK and MEK are formed by the reaction over Ti4+ * Zn2+-TSM/Cl in the absence of oxygen and, interestingly, over Ti4+TSM as well. A main by-product is methyl acetate. The column “Others” represents the other by-products, including mesityl oxide, mesitylene, and a small amount of heavier products. By the reaction of acetone without methanol over Ti4+-TSM,mesityl oxide and mesitylene were formed in the ratio of 7:2 and methyl acetate was not formed at all. The results suggest that the condensation of acetone into mesityl oxide and mesitylene takes place over the catalyst. As seen in Table VII, the addition of oxygen into the feed increases both the conversion of acetone and the selectivity for MVK formation. Comparing the yield of each product in the oxygen-added

-

TABLE VII Reaction of Methanol and Acetone Over Ti4’- and Ti4+ . Zn2+-TSMs” Selectivity (%) Catalyst

Conversion of acetone (%)

In the absence of oxygen Ti4+-TSM Ti4’ .Zn2+-TSM/CI In the presence of oxygen Ti4’-TSM Ti4+ .Zn2+-TSM/Cl

MVK

MEK

MeAc

Others

2.3 3.2

40.7 41.8

4.5 5.9

34.5 33.6

20.3 18.7

11.6 7.9

84.5 74.3

0.2 2.8

11.2 17.7

4.1 5.2

MVK, Methyl vinyl ketone; MEK, methyl ethyl ketone; MeAc, methyl acetate. Reaction conditions: reaction temperature, 350°C;feed, methanol:acetone:02 + Nz = 1.5:1.5:O + 20 or 5 + 15 (ml/min).

318

YUTAKA MORIKAWA

system and that in the oxygen-free system, it can be said that the presence of oxygen does not affect the rate of acetone condensation but promotes MVK formation. Ti4+-TSM is more efficient than Ti4+* Zn2+-TSM/C1for the reaction in the presence of oxygen. The selectivity for MVK formation is as high as 85% at about 12% conversion. The value is higher than 55% selectivity at the same conversion level reported by Ueda et al. for the reaction in the absence of oxygen over metal ion-containing magnesia catalysts (37). The catalysis by Ti4+-TSM is not explainable consistently by the mechanism incorporating the dehydrogenation of methanol, because Ti4+-TSMcatalyzes only the dehydration in the absence of oxygen, as mentioned above. Methyl formate is formed in the reaction of methanol and oxygen over Ti4+-TSM, but the selectivity is only 1.3% and the rest is for dimethyl ether formation (33% conversion). However, some evidence for the incorporation of acidcatalyzed aldol condensation has been obtained in supplemental studies. The reaction of acetone and formaldehyde over Ti4+-TSM proceeds at a comparable rate, forming MVK accompanied by a small amount of divinyl ketone. The reaction of acetone and methanol in the presence of oxygen stops completely by the addition of 200 ppm ammonia into the feed, and is not interrupted by the addition of phenol vapor. The results strongly suggest that the aldol condensation of acetone and formaldehyde or formaldehydelike species derived from methanol is facilitated by the acid site of the catalysts and not by a basic site, if any. The vinylation of methyl acetate and methyl propionate took place in the presence of oxygen over Ti4+-TSM (35,39).The results are briefly summarized in Table VIII, where the values of conversion and selectivity are calculated on the ester basis. Because of the acidity of Ti4+-TSMand the presence of water formed by the reaction, the esters fed and produced undergo TABLE VIII Vinylation of Methyl Acetate and Methyl Propionate with Methanol Over T14'-TSM" Selectivity (%) Reactant ester RCH2COOCH3 R=H R = CH,

Conversion of ester

(%I 8.2 7 .O

RCCOOCH-1

RCHCOOCH,

RCHCOOH

RCHCHO

11

I

I

CH2

CH3

CH3

31.4 15.7

1.9 3.2

48.7 73.5

3.2

Others 14.4 4.0

" Reaction conditions: reaction temperature, 350°C; feed, methanol:ester:02:N2= 1.5: 1.5:5: 15 (ml/min).

M E A L IONS IN LAYER LATTICE SILICATES

319

hydrolysis into the corresponding carboxylic acid and methanol. The values for “Others” listed in Table VIII include the selectivity for the hydrolysis of vinylated esters for the most part, 85% of which is occupied by acrylic acid. Methyl acetate is vinylated to methyl acrylate at 35% selectivity and hydrolyzed to acetic acid at 49% selectivity. The selectivity for methyl acrylate is apparently low. However, the selectivity for the vinylation (methyl acrylate and acrylic acid formation) is evaluated to be as high as 92%, considering that acetic acid is equivalent to methyl acetate because they are easily interchangeable on an acid catalyst. Methyl propionate is hydrolyzed predominantly to propionic acid. The selectivity for methyl methacrylate is as low as 16%. However, the selectivity for the vinylation (methyl methacrylate and methacrylic acid formation) is evaluated to be 63%. The selectivity lower than that observed for the reaction of methyl acetate is due to the formation of methyl isobutylate, isobutylic acid, and isobutylaldehyde, which are produced by the successive reactions of methyl methacrylate. It is speculated that methyl propionate and its vinylated product, which are heavier than methyl acetate and its derivatives, adsorb more strongly on the catalyst because of a larger entropy change, and are therefore hydrolyzed or hydrogenated more extensively. Ueda et al. (37) have proposed magnesium oxide catalyst modified with a transition metal ion (M/MgO) for the vinylation of methyl propionate and acetonitrile. Acetonitrile is vinylated to acrylonitrile selectively (94%selectivity at about 10% conversion) over Cr/MgO catalysts at 350°C in the absence of oxygen. The selectivity for the vinylation of methyl propionate over Mn/MgO catalysts is not different from the value obtained with Ti4+TSM in the presence of oxygen. The catalyst system, however, is not effective for the reaction of acetic acid. We conducted the reaction of acetonitrile and methanol over Ti4+-TSM in the presence of oxygen, and found that the vinylation does not take place but the hydrolysis to acetic acid and subsequent esterification with methanol into methyl acetate proceed preferentially. It is likely that Ti4+-TSMis an appropriate catalyst for the vinylation of carbonyl compounds and MiMgO is appropriate for the vinylation of nitriles.

VI. Immobilization of Catalyst Solution Most metal ion-exchanged TSMs swell with water or polar organic solvents. The resultant intercalate stores up both the solvent and the metal ion in the interlayer spaces, and is regarded as a mass of the microvessels filled with the solution of the metal ion. For liquid-phase homogeneous catalyses

320

YUTAKA MORIKAWA

involving a gaseous reactant, an assembly of the microvessels 1-5 p m in diameter X 5 nm thick has the great advantage of a gas-liquid contact area over the conventional reaction vessel, which is equipped with a bubbling system, in addition to the advantage of easy handling. Some attempts have been made to use a cationic metal complex intercalate as a catalyst; these efforts are briefly reviewed in an account of the catalyst type (b) in Fig. 1. The results advocate the possibility of the immobilization of catalyst solution in the interlayer spaces. The catalytic activity of CuZ+-TSMor Cu2+ Pd2+-TSMswelled with water or an organic solvent for the reaction involving molecular oxygen is introduced in the following section.

A. GAS-PHASE OXIDATION OF PROPYLENE OVER Cu2+ * Pd2+-TSM The oxidation of propylene has been chosen as a probe reaction to study the catalytic activity of Cu2+ Pd2+-TSM.The olefin oxidation in an acidic solution of Cu(I1) and Pd(I1) chlorides, well known as the Wacker reaction, is achieved when olefins are selectively oxidized to ketones or aldehydes by hydrated Pd2+,leaving Pd'. The Pdo is oxidized back to Pd2+by 2Cu2+,and the resulting Cu+ is reoxidized by dissolved oxygen. Because the corrosive nature of the catalyst solution is a serious disadvantage for practical use, supported copper-palladium catalysts have been proposed to operate the reaction in a gas flow reactor (40). Cu2+- Pd2+-TSMsof various Cu2+:Pd2+ratios, including the ratios 1:O (Cu2+-TSM)and 0:l (P2+-TSM),were prepared in the same manner as described in the Section I1 using an aqueous solution of Cu(NH3)X12and/or Pd(NH&C12. The degree of ion exchange was 0.4-0.5 for all samples. The catalyst sample was pretreated at 150°C in a stream of a 1:1 mixture of water and nitrogen. The reaction was conducted at the same temperature under an atmospheric pressure feeding 0.036 atm propylene, 0.373 atm oxygen, 0.500 atm water, and nitrogen. Cu2+- Pd2+-, Cu2+-,and Pd2+-TSMsare completely different from each other in catalytic activity. Cu2+-and Pd2+-TSMscatalyze no reaction and the total oxidation of propylene, respectively, whereas Cu2+- Pd2+-TSMcatalyzes the oxidation to form acetone selectively, suggesting that the Wacker type oxidation takes place over the catalyst (41). The results are shown in Fig. 6. The higher initial activity is observed for Cu2+- Pd2+-TSMwith the lower Cu2+:Pd2+ratio, namely the higher Pd2+ loading. This might be explainable by the second order dependency of the reaction rate on Pd2+concentration, observed for the homogeneous system by Vargaftik er al. in the

L -METAL IONS IN LAYER LATTICE SILICATES

-

Cu/Pd

=

4.60

c

Cu/Pd = 1 . 7 1

.

32 1

Cu/Pd = 0 . 9 0 5

Cu/Pd = 0

A +--+-

1

2

3

4

0

1

2

3

4

1

2

3

4

llX=m=% 1

2

3

4

Reaction t i m e ( h )

FIG. 6 . Rate of formation in the oxidation of propylene over Cu2+ . Pd2+-TSM:0, acetone; @,. carbon dioxide; 0, propionaldehyde. Reaction temperature, 250°C; space velocity, 50 ml/min * g; feed composition, C3Hs:O2:H10:N2= 0.036:O. 373:O. 500:0.091.

absence of C1' (42) and by Moiseev et al. at a high concentration of Pd2+ion (43). The catalytic activity declines with the reaction time. The degree of decline is smaller for the catalysts with the higher Cu2+:Pd2+ratio and is negligible for the catalyst prepared at a ratio of Cu2+:Pd2+= 4.60. It is inferred that the initial decline of catalytic activity results from the decrease in the population of Pd2+.At a high concentration of Cu2+,Pd2+reduced to Pdo is supposed to be recovered mostly by the reoxidation with Cuz+,because the equilibrium, 2Cu2+ PdO e 2Cu+ + Pd2+, favors Pd2+ over Pdo. At a low concentration of Cu2+,the population of Pd2+falls with the reaction time toward the population fixed by the equilibrium. Acetone is not formed appreciably with Pd2+-TSMeven in the initial period of the reaction. Because the white color of Pd2+-TSMchanged to pale gray after the pretreatment, most of Pd2+ions in the catalyst sample seem to have been reduced to PdO before the reaction. The constant activity of Cu2+* Pd2+-TSMfor the selective formation of acetone was achieved in a few hours and was maintained at least for 20 h. The results demonstrate that a Pd2+and Cu2+pair in the interlayer region of TSM catalyzes the Wacker reaction as well as that in the aqueous solution. As illustrated in Fig. 7, the interlayer spaces of TSM are considered to be filled with an aqueous solution similar to that used in the Wacker reaction, although the counteranions are now negatively charged silicate sheets instead of C1- ions. Evidence for the presence of water in the interlayer space has been obtained by TPD and XRD studies on the catalyst samples. The TPD spectrum shows two peaks at around 100 and 2OO0C7 corresponding to the desorption of weakly adsorbed water and that of intercalated water, respectively. The latter peak temperature is higher than the reaction tempera-

+

322

YUTAKA MORIKAWA 1/2 0

FIG. 7 . Working mechanism of Cu”

Pd*+-TSM.

ture by 50°C. The basal spacing (cloo,) is calculated from the XRD angle to be 12.3 A for Cu2+* Pd’l-TSM treated at 150°C. The value is larger by 2.7 A than that observed with completely dehydrated TSM, and coincides with the interlayer spacing frequently found for a smectite containing monolayer water between the silicate sheets. As shown in Fig. 7, the oxidation of propylene probably takes place on the exposed Pd2+ ions because the interlayer spacing of 2.7 A is not wide enough for propylene to diffuse into and out of the interlayer spaces. This means only a part of Pd2+ ions facilitate the reaction and suggests that the further enlargement of the interlayer spacing improves markedly the catalytic activity of Cu2+* Pd2+-TSM.

B. LIQUID-PHASE OXIDATION OF 2,6-Dl-tert-BUTYLPHENOL OVER Cu”-TSM

The liquid-phase oxidation of phenols is one of the model systems for biological oxidations and has been investigated extensively using a variety of oxidizing agents and catalysts, including enzymes (44). Schiff base complexes catalyze the oxidation of 2, 6-dialkylphenol by molecular oxygen, giving both diphenoquinone (DPQ) and benzoquinone (BQ) at the ratios determined by the metal complex, Eq. ( l ) , and the oxidation rate is accelerated by the addition of base into the reaction system (45,46). R

R

R

R

(1)

The oxidation of 2, 6-di-tert-butylphenol (DTBP) was carried out at room temperature for 24 h by stirring the mixture of a solvent (50 mi), DTBP (1.00 g = 4.85 mmol), and Cu2+-TSM(500 mg = 0.318 mmol) with bub-

323


bling oxygen. Cu2+-TSM was treated at 400°C for 2 h in a helium stream prior to the reaction. The interlayer spacing of Cu2+-TSMin the solvent was calculated from the doolvalue determined by XRD, subtracting that (9.6 A) of the heat-treated Cu2+-TSM.We examined the catalytic activity of Cr3+-, Mn2+-, Fe3+-,Co2+-,and Ag+-TSMs, in addition to Cu2+-TSM, and found Cu2+-TSMto be most active among them. Table IX shows the results of the reaction with Cu2+-TSM in methanol solvent (41) together with the data obtained with Schiff base complexes by Frostin-Rio et al. (46). The reaction rate and the product distribution in the Schiff base complex systems are very dependent on the solvent and on the base added. With Cu2+-TSM, DTBP was oxidized to DPQ selectively. The addition of base into the Cu2+-TSMsystem accelerates the reaction rate remarkably but does not affect the high selectivity. The rate of DTBP oxidation and the interlayer spacing of Cu2+-TSMvary enormously depending on the solvent employed (41).The results of the reaction and the interlayer spacing in various solvents are summarized in Table X , where the value of selectivity is not listed because DPQ is the only product in every case. The value of interlayer spacing is calculated to be TABLE IX Oxidation of 2,6-Di-tert-butylphenol in Various Catalytic Systems

Catalyst

Solvent

Base

Co(salpn)b Co(sa1pn)' Co(tpP)b Mn( ttp)b Mn(tppIh Cu*+-TSM Cu2 -TSM

CHCI, CHCI, CHCli Toluene CH3CN CH,OH CH3OH

Pyridine Pyridine NaOH NaOH NaOH

+

Reaction time (h) 24 24 9 4 0.5

24 24

Selectivity (%)"

Yield

(%I

BQ

DPQ

30 40

52 80

48 20

100

100

I0

40 10 0 0

0 60

100

19 12

BQ, Benzoquinone; DPQ, diphenoquinone.

* See Ref. 46; salpn, bis(sa1icylidene) propylenediamine; tpp, tetraphenylporphin. TABLE X Interlayer Spacing and Catalytic Activity of &u 2+-TSM in Various Solvents

Solvent

Interlayer spacing (A)

Conversion (%)

CH3 OH CzHsOH CH, CN Benzene Toluene

7.4 I .4 7.5

19 10

0

16 0

0

0

90 100 100

324

YUTAKA MORIKAWA TABLE XI @“ect of Base Added into the Cu2+-TSMCatalyst System ~

~~~

~~

Solvent

Base

Interlayer spacing (A)

Conversion (%)

CHiOH CH30H CH,OH Benzene Benzene Benzene

-

1.4

NaOH Pyridine NaOH Pyridine

19 12 51 0 0 17

7.4 1.5 0 (0 or -3) 5.6

about 7.4 A for Cu2+-TSMsoaked in CH30H, C2H50H,and CH3CN, indicating that Cu2+-TSMswells extensively in the polar solvents. In contrast, CuZ+-TSMdoes not swell at all in nonpolar solvents such as benzene and toluene. As seen in Table X, the oxidation of DTBP proceeds at a considerable rate in the polar solvents and at a negligible rate in the nonpolar solvents. It is likely that Cu2+-TSMacts as an effective catalyst when the interlayer spacing is expanded to a width great enough for DTBP to get into the interlayer spaces. The effects of NaOH or pyridine added into the reaction mixture are shown in Table XI (41). The addition of base into the methanol solvent system does not change the expanded layer structure, and accelerates the reaction rate to give the value of conversion 3.8 or 2.5 times as high as the value obtained with the base-free system. The acceleration effect of NaOH is not observed for the benzene solvent system. The interlayer spacing of Cu2+TSM soaked in the mixture of benzene and NaOH has not been determined because of difficulty in the preparation of XRD sample, but is estimated to be 0 or 3 A, assuming no intercalation or intercalation of the small NaOH molecule, respectively. It is of interest that the addition of pyridine into the benzene solvent system causes both the expansion of the interlayer spacing and the remarkable increase in the reaction rate. The facts suggest that pyridine acts as an accelerator and a pillar to keep the silicate sheets apart, enabling DTBP to diffuse into the interlayer spaces. The results seen in Tables X and XI show that the reaction always proceeds at a considerable rate when the interlayer spacing of Cu2+-TSMis expanded. This means that only a part of the metal ions is distributed on the rim of interlayer space, and the observations give a guide for improvement of the catalytic activity of metal ion-exchanged layer lattice silicates.

-

VII. Conclusions

Fluorotetrasilisic mica is a synthetic swelling layer lattice silicate. Because of the cation-exchange properties and the lack of acid sites on the sili-


325

cate sheet, TSM is a particularly appropriate carrier for use in the study of the intrinsic activities of metal ions in heterogeneous catalysis. The catalytic activities of metal ion-exchanged TSMs for the methanol conversion are individually different, depending on the metal ion. Figure 8 shows the attractive reactions found in this study. It is of particular interest that the catalytic activity of copper varies depending on the oxidation state. Cu+-TSM is inactive, whereas Cu2+-and Cu"-TSMs catalyze selectively the dehydrogenation of methanol to methyl formate and to formaldehyde, respectively. Cu2+-TSMhas potential for practical use because of its high stability and high selectivity for the methly formate formation. Ti4+-TSMacts as an acid catalyst although the original form of TSM has no acidity, suggesting that exchange of interlayer cations with Ti4+ ions enhances the acidity of layer lattice silicates. In fact, Ti4+-exchanged layer lattice silicates are much more active than the original and the proton-exchanged forms for the dehydration of methanol. It is noteworthy that Ti4+-exchanged activated clay shows the selective activity for the methanol conversion into lower olefins. An extension of the studies leads to discovery of novel reactions of methanol. Methanol is converted into a variety of oxygenated compounds over Ti4+ Zn2+-TSM/Cl. Although characterizations of the complicated catalyst and the improvement of selectivity are needed, this particular methanol conversion is most interesting with respect to the reactions shown in Fig. 8. The other three reactions are carbonyl compound vinylations, the products of which are valuable monomers for functional polymers. Cu2+ * Pd2+-TSM catalyzes the oxidation of propylene to acetone selectively in the presence of water vapor. Because Cu2+ Pd2+-TSMswells with water, the interlayer region is considered to be filled with a solution similar to that used in the Wacker reaction. However, the interlayer spacing of Cu2+ * Pd2+-TSM is not wide enough for the reactant molecule to diffuse into the interlayers, and a great number of the Pd2+ ions located at the inside of the interlayers are thought not to contribute in the reaction. For the

-

Ti4+ -Z"2+-TSn/C1

Oxygenated compounds

CHzCHCOCH3

J

Lower olefins

CH2FOOCH3 R

FIG.8 . Reactions of methanol over metal ion-exchanged layer lattice silicates

326

YUTAKA MORIKAWA

oxidation of 2, 6-di-tert-butylphenol, Cu2+-TSMacts as an effective catalyst to produce diphenoquinone selectively when the intercalation of solvent and/or base molecule has expanded the interlayer spacing. The results suggest that the interlayer spaces of layer lattice silicates are able to be utilized as a microvessel of catalyst solution, and all of the interlayer metal ions are effectively incorporated into the reaction when the interlayer spacing is expanded to a width great enough for the reactants to diffuse into the interlayers. REFERENCES 1. Grim, R. E., “Clay Mineralogy,” 2nd Ed., p. 353. McGraw-Hill, New York, 1968. 2. Ryland, L. B., Tamele, M. W., and Wilson, J. N., in “Catalysis” (P. H. Emmett,. ed.), Vol. 7, p. 1. Reinhold, New York, 1960. 3 . Thomas, J. M., in “lntercalation Chemistry” (M. S. Whittingham and A. J. Jacobson,

eds.), p. 55. Academic Press, New York, 1982. 4. Rupert, J. P., Granquist, W. T., and Pinnavaia, J. T., in “Chemistry of Clays and Clay Minerals” (A. C. D. Newman, ed.), p. 297. Wiley, New York, 1987. 5. Miyazaki, T., Tsuboi, A , , Urata, H., Suzuki, H., Morikawa, Y., Moro-oka, Y., and Ikawa, T., Chem. L e t t . p. 793 (1985). 6. Brindley, G. W., and Sempels, R. E., Clay Miner. 12, 229 (1977). 7. Lahav, N., Shani, U., and Shabtai, J., Clays Clay Miner. 26, 107 (1978). 8. Yamanaka, S . , and Brindley, G. W., Clays Clay Miner. 26, 21 (1978). 9. Yamanaka, S . , and Brindley, G. W., Clays Clay Miner. 27, 119 (1979). 10. Endo, T., Mortland, M. M., and Pinnavaia, T. J., Clays Clay Miner. 28, 105 (1980). 11. Yamanaka, S . , Doi, T., and Hattori, M., Muter. Res. Bull. 19, 161 (1984). 12. Pinnavaia, T. J., Tzou, M. S., and Landau, S. D., J . Am. Chem. SOC. 107, 4783 (1985). 13. Yamanaka, S . , Nishihara, T., Hattori, M., and Suzuki, Y., Muter. Chem. f h y s . 17, 87 (1987). 14. Kitajima, K., and Daimon, M., Nippon Kuguku Kuishi p. 991 (1975). 15. Toraya, H., Iwai, S., Marumo, F., Daimon, M., and Kondo, R., Z. Kristullogr. Kristullgeom. Kristullphys. Kristullchem. 44, 42 (1976). 16. Morikawa, Y., Takagi, K., Moro-oka, Y., and Ikawa, T., J.C.S. Chem. Commun. p. 845 (1983). 17. Yasumori, I., and Miyazaki, E., Nippon Kuguku Zusshi 92, 659 (1971). 18. Chono, M., and Yamamoto, T., Shokubui 23, 3 (1981). 19. Chang, C. D., and Silvestri, A. J., J. Curd. 47, 240 (1977). 20. Morikawa, Y., Goto, T., Moro-oka, Y., and Ikawa, T., Chem. L e t t . p. 1667 (1982). 21. Tanaka, K . , Ozaki, A., and Tamaru, K., Shokubui 6, 262 (1964); Tanaka, K . , and Ozaki, A , , Bull. Chem. Soc. Jpn. 40, 1728 (1967); Tanaka, K., and Ozaki, A., J . Cutul. 8, l(1967). 22. Misono, M., Saito, Y., and Yoneda, Y., J. Curd 9, 135 (1967). 23. Morikawa, Y., Yasuda, A , , Moro-oka, Y., and Ikawa, T., Chem. L e t t . p. 1911 (1983). 24. Siegel, S . , J . Cutul. 30, 139 (1973); Tanaka, K . , and Okuhara, T., Cutul. Rev.-Sci. Eng. 15, 249 (1977); J . Curd. 65, 1 (1980). 25. Orchin, M., Adv. Cutul. Relut. Subjects 16, (1966). 26. Morikawa, Y., Takagi, K., Moro-oka, Y., and Ikawa, T., Chem. Lett. p. 1805 (1982); f r o c . l n t . Congr. Cutal., 8th, Berlin 5 , 680 (1984). 27. Takagi, K., Morikawa, Y., and Ikawa, T., Chem. Lett. p. 527 (1985).


327

28. McBride, M. B . , Pinnavaia, T. J., and Mortland, M. M., J. Phys. Chem. 79, 2430 (1975). 29. Morikawa, Y., Liu, Y., Moro-oka, Y., and Ikawa T., Sekiyu Gukkuishi 26, 321 (1983). 30. Morikawa, Y., Wang, F., Moro-oka, Y., and Ikawa, T., Chem. Lett. p. 965 (1983) (in this paper, bentonite is abbreviated to Mont). 31. Kaeding, W. W., and Butter, S. A , , J. Catal. 61, 155 (1980). 32. Inui, T., Matsuda, H., Yamase, O., Nagata, H., Fukuda, K . , Ukawa, T., and Miyamoto, A., J . Cutul. 98, 491 (1986). 33. Inui, T., Patanasry, S., and Matsuda, H., J.C.S. Chem. Commun. p. 205 (1990). 34. Morikawa, Y., Wang, F., Moro-oka, Y., and Ikawa, T., Chem. Lett. p. 1849 (1983). 35. Morikawa, Y., Wang, F., Kurokawa, H., Ueda, W., and Ikawa, T., Proc. Int. Congr. Cutul., 9zh, Calgary, Alberta 1, 292 (1988). 36. Wang, F., Ueda, W., Morikawa, Y., and Ikawa, T., Chem. Lett. 405 (1990). 37. Ueda, W., Yokoyama, T., Moro-oka, Y., and Ikawa, T., J.C.S. Chem. Commun. p. 39 (1984); Ind. Eng. Chem. Prod. Res. Dev. 24, 340 (1985). 38. Wang, F., Ueda, W., Morikawa, Y., and Ikawa, T., Chem. Lett. p. 1991 (1988). 39. Wang, F., Ueda, W., Morikawa, Y., and Ikawa, T., Chem. L e f t . , p. 281 (1989). 40. Fujimoto, K . , Negami, Y., Takahashi, T., and Kunugi, T., Ind. Eng. Chem. Prod. Res. Dev. 11, 303 (1972); Evnin, A. B., Rabo, J. A , , and Kasai, P. H., J . Cutul. 30, 109 (1973); Arai, H . , Yamashiro, T., Kubo, T., and Tominaga, H., Bull. Jpn. Pet. Inst. 18, 39 (1976). 41. Morikawa, Y., Ishikawa, S., Tanaka, H., and Ueda, W., Proc. Int. Con$ Ion Exch. ‘91, Tokyo, 139 (1991). 42. Vargaftik, M. N., Moiseev, 1. I., and Syrkin, Y. K., Dokl. Akud. Nuuk SSSR 147, 399 ( 1962). 43. Moiseev, I . I . , Levanda, 0. G . , and Vargaftik, M. N., J. Am. Chem. SOC.96, 1003 (1968). 44. McDonald, P. T., and Hamilton, G . A , , in “Oxidation in Organic Chemistry” (W. S. Trahanovsky, ed.), Part B, p. 97. Academic Press, New York, 1973. 45. Nishinaga, A., and Tomita, H., J. Mol. Cutul. 7, 179 (1980); Nishinaga, A., Tomita, H., Nishizawa, K., and Matsuura, T., J.C.S. Dalton p. 1504 (1981); Bied-Charreton, C., Frostin-Rio, M., Pujol, D., and Gaudemer, A., J. Mol. Cutul. 16, 335 (1982). 46. Frostin-Rio, M., Pujol, D., Bied-Charreton, C., PerrCe-Fauvet, M., and Gaudemer, A., J.C.S. Perkin I, p. 1971 (1984).



Catalytic Synthesis of ChlorofIuorocarbon Alternatives L. E. MANZER

AND

V. N. M. RAO

Corporate Catalysis Center Central Research and Development Du Pont Company Experimental Station Wilmington, Delaware 19880

1.

Introduction

A. BACKGROUND The photolysis of fully halogenated chlorofluorocarbons (CFCs) in the stratosphere and their role in ozone depletion is of great global concern. In 1971, using electron capture techniques capable of detecting parts per trillion levels, Lovelock (1) showed that the very stable CFCs were accumulating in the atmosphere. The relationship between ozone depletion and CFCs was the subject of a paper in Nature by Molina and Rowland (2). Following this publication in 1974, a review of this relationship in greater detail was published (3).In the following years, a large body of scientific information has been obtained that is consistent with the theory that chlorine atoms, released from long-lived CFCs by the action of short-wavelength solar ultraviolet radiation, initiate the ozone destruction cycle. Equations (1)-(3) represent such a cycle. CCI,F CI

+ UV +0

CIO

3

+0

-4

+ CC12F c10 + c1 + 02 CI

0 2

(1) (2)

(3)

Based on science, there was general agreement that emissions of CFCs to the atmosphere have to be minimized. In September, 1987, the Montreal Protocol, an international United Nations agreement, was signed. This agreement called for periodic review of developing science and a 50% cut in the amount of CFCs produced by the year 1998. In March, 1988, however, 329 Copyright 0 1993 by Academic Press, Inc. All rights of reproduction in any form reserved.

330

L. E. MANZER AND V. N. M. RAO

with the ratification process just beginning, the National Aeronautics and Space Administration’s (NASA) Ozone Trends Panel announced new findings ( 4 ) . Based on these findings the Protocol was further strengthened in June, 1990 and called for a total phaseout by developed countries by the year 2000. Recent analysis of data collected in the stratosphere during the fall of 1991 indicated that ozone depletion was continuing at an accelerated pace and exists over populated regions in both hemispheres during summer months. As a result, Du Pont advanced their timetable to phase out CFCs by year-end 1996, and brominated analogs, by year-end 1994. In addition to the role that CFCs play in the destruction of ozone, they have also been implicated in global warming. Climate change that might occur due to the accumulation of long-lived greenhouse gases in the atmosphere is also a matter of great public concern. Carbon dioxide, which is by far the largest component released to the atmosphere, also contributes to global warming and carries a major share of this potential. Related to the CFCs are the brominated analogs that are made by partial replacement of the chlorine with bromine. These compounds, known by such names as halons, will also be regulated and phased out as part of the Protocol.

B . APPLICATIONS Historically, CFCs have fulfilled a true societal need. Because of their inertness, stability, and low toxicity they have been used in a wide variety of applications. The largest single use of these compounds is in the area of refrigeration. Food chain management and air conditioning, combined, constituted about 30% of the 2.5 billion pounds of CFCs produced in 1988. A close second is in the area of foaming agents, referred to in the industry as blowing agents. The unique thermal conductivity properties of CFCs trapped within foams has been exploited to produce a host of insulating materials with outstanding efficiency. The use of such materials in the housing industry, in refrigerated trucks and railway cars, and in refrigerators and freezers are but a few examples. About 20% of the use in 1988 was in the development of cleaning agents used for cleaning and degreasing of metals and the cleaning of sophisticated medical instrumentation and electronic components, for example. Although used as aerosol propellants in the past, the use of CFCs for such applications in the United States was banned in 1978. Several CFCs find use as chemical intermediates for the synthesis of other fluorine-containing compounds, notably fluorine-containing olefins. Some specialty applications include dielectric fluids, inert liquids, and brominated analogs used as fire extinguishants.

CATALYTIC SYNTHESIS OF CFC ALTERNATIVES

c.

33 1

REQUIREMENTS OF A CFC SUBSTITUTE

An ideal CFC substitute is one that offers identical performance properties of the CFC it replaces with the added stipulation that it does not destroy the ozone layer and has low global warming potential. Low toxicity, a property of CFCs, and shorter atmospheric lifetimes compared to CFCs are also preferred. Summarized, the desired characteristics include short atmospheric lifetimes, low toxicity, ozone compatibility, low global warming potential, thermal and chemical stability, equivalent physical properties, and amenability to realistic processes. A short atmospheric lifetime will ensure that the molecule is degraded in the troposphere, leaving little chance for entry into the stratosphere. For example CHCIR (HCFC-22), a hydrochlorofluorocarbon that is produced commercially and is estimated to have an atmospheric lifetime of less than 20 years, might degrade by a scheme as shown in Fig. 1, leaving very little to enter the stratosphere. To eliminate the possibility of even small quantities entering the stratosphere and starting a chlorine atom-initiated ozone destruction cycle, the ideal replacement would be a hydrofluorocarbon (HFC) containing no chlorine. The presence of hydrogen in the molecule allows for abstraction of this hydrogen by hydroxyl radicals, as shown in Fig. 1, and further degradation in the troposphere. Because CFCs are in such widespread use and realistic HFCs for every need have not yet been identified, HCFCs, when available, have been suggested as bridging molecules to fulfill a short-term need. In contrast to long atmospheric lifetimes, molecules with too short an atmospheric lifetime can contribute to ground level smog, which is also undesirable. In the area of refrigeration, blowing agents, and cleaning agents, the currently used CFCs and their possible replacements are listed in Table I. There are no direct drop-in replacements for the CFCs. Table I1 gives a

02

CF2Cl*

CF&I-0-0.

0.

HCI

H20

FIG.1. Proposed mechanism for HCFC-22 decomposition.

332

L. E. MANZER AND V. N. M. RAO TABLE I Potential CFC Substitutes

Market

Current CFC

Refrigerants

CFC-12 (CFzC12) CFC-115

Blowing agents

CFC- 1 1 (CFCI?)

CFC- 1 14 (CFzCICF2CI) Cleaning agents

CFC-113 (CFzClCFC12)

CFC Alternative HFC- 134a (CF3CFH2); HCFC-22 (CHFZCI) HFC- 125 (CSCF2H); blendslazeotropes HCFC-141b (CH3CFC12); HCFC-123 (CF3CHC12); HCFC-22 (CHF2CI) HCFC- 124 (CF3CHFCl); blends/azeotropes Blends/azeotropes;new compounds

comparison of the physical properties of HFC-134a and HCFC-123 and CFC- 12 and CFC- 11. Minor variations in these properties can result in major equipment changes for some applications. In this regard, the use of blends and azeotropes show promise. In fact, the uses of blends and azeotropes are being patented and evaluated in these applications to minimize performance variations.

DEVELOPMENT OF FLUORINATION CATALYSIS D. EARLY Until recently, large-scale commercial syntheses have focused entirely on fully halogenated derivatives. Two notable exceptions are CHClF2 (HCFC-22) and CH3CHF2 (HFC- 152a). Although there is a considerable body of published literature on the synthesis of fully halogenated molecules, information on catalytic commercial approaches to hydrogen-containing substitutes is still very scant. Based on published literature and patents, one can surmise that most are produced by complex chemical processes whose details are largely undisclosed. The basic step in several of these syntheses is the exchange of a chlorine for fluorine, producing an equimolar amount of HC1. This exchange is quite slow in the absence of a catalyst. In 1891, Swartz (5) showed that liquid-phase contact of a chlorocarbon with SbF3 produced a fluorinated compound. In this case he found that an activated trichloromethyl group was converted to the trifluoromethyl group. The SbF3, which was partially converted to the trichloride or mixed chlorofluorides, was converted back to SbF3 in a separate step by treatment with anhydrous hydrogen fluoride. In subsequent experimentation, Swartz ( 6 ) observed that addition of trace amounts of pentavalent antimony could be added with benefit.

333

CATALYTIC SYNTHESIS OF CFC ALTERNATIVES TABLE I1 Comparison of HFC- 134a with CFC-12 and HCFC-123 with CFC-11 Parameter

HFC- 134a

CFC- 12

HCFC- 123

CFC- 1 1

Molecular weight Boiling point ("C at 1 atm) Freezing point ("C) Critical temperature ("C) Critical pressure (atm) Critical density (g/cm3) Heat of vaporization (cal/g) Viscosity (cps at 25°C) Surface tension (dynedcm at 25°C)

102.02 -26.15 - 101 101.05 40.64 0.508 47.52" 0.204 8.466

120.92 -29.79 - 158 112.0 40.6 0.558 39.476 0.214 9.0

152.9 27.9 - 107 185.0 37.4 0.53 41.6b 0.449 15.626

137.37 23.82

a

-111

198.0 43.5 0.554 43. I b 0.415 18.0

At 0°C. At the boiling point.

It took until 1928, however, for Midgley and Henne (7) to show that CFCs were the "right" class of compounds to be used as refrigerants as opposed to the then prevalent SOz and NH3, which are flammable. Midgley and Henne reacted Cc14 with pentavalent antimony chlorofluoride to produce CC12F2.The nontoxic and nonflammable nature of CClzF2 was dramatically illustrated by Midgley in an American Chemical Society meeting in 1930 (a practice we would defer from today), during which he took a breath of CC12Fzand put a lighted candle out! This demonstration marked the beginnings of the CFC industry as we know it today. Continuous processes were developed wherein a chlorocarbon and HF were fed to a reactor containing antimony pentahalide, usually dissolved in the fluorinated reaction intermediates. Under reaction conditions, pentavalent antimony is somewhat unstable, reverting back to the trivalent state and chlorine. Industry practice is to feed chlorine to oxidize trivalent antimony back to the pentavalent state. In its simplest form the exchange reaction with CC14 can be written as shown in Eqs. (4)and ( 5 ) . Over the years, several improvements to such processes have been made

-

SbC15 + 3HF SbC12F3

+ 2CCI4

+

SbC12F3 + 3HC1

SbCls

+ CC4F + CClzFz

(4) (5)

and reported in the patent literature (8-10). Operating conditions vary widely, with reaction pressures of 0-500 p.s.i.g. and temperatures of about 40-200°C. Operating conditions are usually adjusted in the above ranges depending on the chlorinated compound that is undergoing halogen exchange and the desired end-product. In addition to the liquid-phase processes discussed above, vapor-phase halogen-exchange processes have also been developed. A variety of metal

334


TABLE I11 Catalytic Reactions of Fluorocarbons Halogen exchange: Isomerization: Disproportionation: Elimination: Hydrodehalogenation: Dehydrohalogenation:

-

CF2ClCF2CI + HCI CFzClCFClz + HF CF2CICFC12 ---+ CF~CCII CF2CICFC12 CF2CICCl3+ CFzCICFzCl CFjCHzCI CF2=CHCI + HF CF$XII t H1 -+ CF3CHCI + HCI CF2ClCFClz + Hz CFz=CFCI t HCI

--

-

fluorides, notably chromium fluoride, have been reported to be effective catalysts for this exchange reaction (11-13). Unlike liquid-phase processes in which liquid boil-up provides the needed temperature control, vapor-phase processes have to be controlled carefully. They are also more amenable to the synthesis of compounds having a greater degree of fluorination because the operating temperatures are usually higher.

E. GENERAL APPROACHES TO ALTERNATIVES The need for retention of hydrogen in the molecule has led to the development of a wide variety of catalysts and processes that are selective in this regard. Starting materials for such processes have included hydrocarbons, halohydrocarbons, olefins and haloolefins. The general processes of HF addition to olefins, halogen exchange, isomerization, disproportionation, chlorofluorination, and hydrogenolysis constitute the majority of the approaches. Table I11 lists some representative examples, the key, however, being the development of specific and selective catalysts aimed at the synthesis of a target molecule. Both liquid- and vapor-phase processes have been reported for many of the above molecules.

11.

Catalytic Synthesis of Key CFC Alternatives

A. 1,1,1,2-TETRAFLUOROETHANE 1. From Trichloroethylene

A direct one-step route to 1,l ,1,2,-tetrafluoroethane (HFC- 134a; CF3CH2F)can be written as shown in Eq. (6). Currently, no commercially viable process has been described for this approach involving highconversion single-pass yields. This has to do with equilibrium limitations, as explained below. Because of this, decoupling of the overall process into


+ 4HF

-

335

+ 3HC1

(6) two separate steps has been the subject of several reports. In this approach trichloroethylene (TCE) is allowed to react with anhydrous HF to produce CF3CHKl (HCFC-133a). The first step is the addition of HF to the double bond. This is followed by two successive halogen exchanges to afford HCFC-I33a, as shown in Eq. (7). Both liquid (14) and vapor-phase (15) processes are known for this chemistry. For liquid-phase reactions, Feiring has reported (16,17) the applicability of several catalysts. These include CCIz=CHCI

CC12=CHCI

-

CF3CHZF

CClzFCHzCl

-

CF3CHzCI

(7)

SbXs , BF3, TaFs , NbFs , and MoCls, among others. A drawback of using antimony pentahalide catalysis for this synthesis is the inherent instability of the pentahalide, as mentioned earlier. In such processes both trivalent and pentavalent antimony might be present, but any chlorine added to reoxidize part of the trivalent antimony to the pentavalent state can also chlorinate the hydrogen, thus defeating the very purpose of preserving the hydrogen. Stable pentavalent systems afforded by tantalum or niobium are superior in this regard, as shown by Feiring (16,ZQ. Improvements to the antimony-based system have been suggested by Mitschke and Niederpruem (18), who have used various cocatalysts. These include the addition of salts of Zn, Co, Ni, Fe, Zr, and Pd to the antimony system. In addition, these authors have indicated that preservatives or stabilizers present in TCE should be removed prior to reaction to maintain catalyst performance. Vapor-phase processes for the synthesis of HCFC- 133a have conventionally employed chrome-based systems as catalysts. Use of unsupported chromium oxide (19) and trivalent chrome salts supported on active carbon or alumina has also been applied to this transformation (20). These catalysts are usually pretreated with anhydrous HF at elevated temperatures prior to introduction of the organic feed. The second stage of this two-step process is the conversion of HCFC-133a to HFC-l34a, as shown in Eq. (8). This halogen is difficult to exchange due to equilibrium considerations, as mentioned above. With stoichiometric quantities of reagents and operating CF3CHZCI

+ HF

+

CF3CH2F + HCI

(8)

above 35OoC, only about 3% conversion to the desired HFC-134a is obtained. Forcing conditions and the use of a large excess of HF are required to shift the equilibrium in the desired direction. This leads to costly manufacturing schemes associated with recycle of unreacted HF. Using chromium oxide as catalyst and operating at 35O-40O0C, high equilibrium conversions and selectivity to the desired HFC- 134a have been reported by Bell (21). A major drawback associated with chromium oxide-based systems for this chemistry has been one of catalyst life. These deactivate rather rapidly, es-

336


CF3CH2CI

+

HF W F 3 C Y F

+

HCI

CF,=CHCI 7 I F C ECCII -[coke] HF

FIG.2. Coke formation via HCFC-133a.

pecially at higher operating temperatures. Theories have been proposed as to the mechanism of deactivation. These generally have included the assumption of the loss of HF from either the starting material or product to give olefins, which ultimately lead to coking of the surface. A plausible scheme is shown in Fig. 2. To overcome this deactivation, cofeeds such as chlorine or oxygen have been proposed (22). These two cofeeds, while extending catalyst life, can also contribute to chlorination of the molecule. Chromium oxide is a well-known catalyst for the oxidation of HCI to Clz and H20. The use of oxygen can oxidize the HCl coproduct to chlorine, again leading to chlorination of the molecule. To overcome this, a separate regeneration step to remove carbon deposits can also be used with advantage. However, improved catalysts based on aluminum fluoride or fluorided alumina, which have a longer catalyst life, have been recently reported by Manzer (23). These have a reduced tendency for oxidizing HC1. Although the one-step process described in Eq. (6) has not been outlined in detail, potential variants have been proposed (24,25)claiming high selectivities to the desired HFC-134a. A single catalyst has been discovered that carries out both the hydrofluorination of TCE and halogen exchange of HCFC-133a. Basically, TCE is fed to the catalytic reactor at the rate at which HFC-I34a is removed. HCFC-133a is recycled continuously through the reactor. The recycle of excess HF is thereby minimized.

2. From Tetrachloroethylene A well-studied multistep synthesis of HFC- 134a involves the hydrogenolysis of CF3CCl2F (CFC- 114a) derived from tetrachloroethylene (PCE) through a series of transformations. Possible scenarios are shown in Fig. 3. While circuitous, it might be a method of choice to utilize investment and facilities that might exist for producing CFC-113, an intermediate in this process. Briefly, this approach involves the chlorination of PCE to hexachloroethane, which then undergoes successive halogen exchange reactions to afford fully halogenated species containing three or four fluorines. The


337

FIG.3. PCE routes to HCFC-134a.

disposition of the fluorines and chlorines on the two carbons depends, to a large extent, on the catalytic process used. When using conventional antimony-based catalysts that operate in the liquid phase, the usual products are CF2C1CFCl2 (CFC- 113) and/or CF2CICF2C1 (CFC- 114). Chrome-based caidysts that operate in the vapor phase also produce CFC- 1 13 and/or CFC114 (26). Several versions of the chrome-based systems have been reported, claimiig specific advantages. Vecchio et al. have used salts of iron and chromi,im supported on aluminum fluoride (27). Such systems produce CFC-11.3 and CFC- 114. In making these compounds some isomerization does occuv, leading to the two rearranged species CRCCl3 (CFC- 113a) and CF3CCI2F(CFC-114a). This rearrangement is more predominant in vaporphase processes. This, in itself, is desirable, since it is the rearranged species that is required for further conversion to HFC- 134a. The associated chemistries are shown in Eqs. (9)-( 13).

cc1,=cc12 + c1,

-

CClsCCIs

+ 3HF ------z CFzClCClzF + 3HCl CF2ClCClzF + HF + CFZCICF2CI + HCI CCliCCli

-

CFzClCCliF + CF3CCI-1

(9) (10) (11)

(12)

(13) An exceptional catalyst for isomerization of CFC- 1 13 to CFC- 1 13a is anhydrous aluminum chloride as reported by Miller (28). This isomerization is usually carried out in the liquid phase and under mild conditions. Some disproportionation of CFC- 113a to CFC- 1 14a and CF2CICC13 (CFC- 112a) is also observed on prolonged contact with catalyst. The use of trace quantities of metals such as chromium and manganese is claimed to have a beneficial effect in this process (29). There is an initial activation period CFzClCF2Cl

CF~CCIZF

338


usually associated with isomerization. Partial exchange of the halogens in aluminum chloride to produce an “aluminum chlorofluoride,” which might be the true catalyst, can explain the need for this initial activation. Vaporphase isomerization of CFC- 113 to CFC- 113a has also been reported. A variety of metal salts have been claimed to be effective in this isomerization (30). The actual method of isomerization used might well depend on the synergy it offers to other parts of the manufacturing process. Further reaction of CFC- 113a with HF produces CFC- 1 14a. Again, conventional liquid- and vapor-phase processes are deemed satisfactory. A useful expedient is the isomerization of CFC-114 to CFC-l14a, because CFC114 can be produced in one step, starting from PCE. This is a higher energy process and both liquid- and vapor-phase approaches to this isomerization have been reported (29,31,32). A catalyst that will produce CFC-114a directly is very desirable because extra processing steps can be saved. Vecchio et al. (27) have in fact reported that aluminum fluoride is an effective catalyst for this chemistry. Although the major product is CFC- 114a, some CFC-114 is still produced and the boiling points of the two isomers are such that they cannot be separated by conventional means. The effect of added metals to an aluminum fluoride catalyst is to suppress this isomerization activity. Although the above processes produce CFC- 114a as the major product, disproportionation followed by fluorination of one-half of the product has also been suggested (33).For example, disproportionation of CFC-114 over a variety of metal salt catalysts produces CF3CF2C1(CFC-115) and CFC113a in about equal amounts. Catalyst life also appears to be a shortcoming with this chemistry because the authors recommend cofeeding either O2 or Clz to extend catalyst life. The CFC-113a is then converted to CFC- 114a either in the liquid or vapor phase. 3. Hydrogenolysis of CFC- I14a Replacement of the chlorines in CFC- 114a through catalytic hydrogenolysis produces HFC-l34a, as shown in Eq. (14). Bitner et al. have studied CF~CCIZF+ 2Hz

--+

CF3CHzF

+ 2HCI

(14)

the catalytic hydrogenolysis of carbon-chlorine bonds in molecules containing both chlorine and fluorine (34). Palladium supported on carbon was their preferred catalyst. In the case of CFC-114, they have shown that hydrogenolysis using palladium supported on carbon affords HCSCF2H (HFC-134) as a major product. The facile hydrogenolysis of a trichloropentafluoropropane to the corresponding trihydro analog using palladium on carbon has been reported by


339

Smith (35).Applicability of palladium supported on charcoal or alumina for the hydrogenolysis of CFC- 114a to HFC- I34a has been reported by Darragh (36). In investigating this reaction, Gervasutti et al. (37)have observed that CFC-114a is much more reactive under hydrogenolysis conditions compared to CFC- 114. The dichloromethyl moiety appears to be more reactive under hydrogenolysis conditions. Their studies also indicate that hydrogenolysis of CFC- 114a affords largely HFC- 134a in addition to smaller quantities of CF3CHFCl (HCFC-124) arising by replacement of a single chlorine by hydrogen. This latter molecule also undergoes hydrogenolysis to HFC-134a with palladium catalysis, but at a higher temperature. Quite surprisingly, Kellner et uI. (38)have shown that a deactivated palladium catalyst in the hydrogenolysis reaction can be regenerated by a simple treatment at elevated temperature with a CFC or HCFC in the absence of hydrogen. Since the early 1960s, palladium has been the preferred catalyst for the hydrogenolysis of carbon-chlorine bonds. Recently, catalysts other than palladium have been reported for the hydrogenolysis of CFC-114a. In a series of disclosures, Morikawa and co-workers have reported catalysts that are not based on palladium or those that, when present with palladium, increase catalyst life (39- 42). Group VIII metals have been shown to be catalytically active as well as the carbides of tungsten and molybdenum. Additives to palladium such as iridium, ruthenium, osmium, tungsten, molybdenum, and rhodium have been shown (42) to reduce palladium sintering as a function of length of use under hydrogenolysis conditions. For catalyst systems containing other elements in addition to palladium, Morikawa and co-workers (39)have interpreted performance changes based on chemical adsorption energy as well as geometric factors. In contrast to palladium and other Group VIII metals, Koto et al. have reported an active carbon-catalyzed hydrogenolysis of CFC- 114a (43). The process generally operates at a higher temperature, over 400°C, and appears to produce HCFC- 124 and HFC- 134a as products of sequential reactions, because the former product is produced in higher proportion. 4. Hydrogenolysis of HCFC- 124 Most catalysts discussed in the previous section have also been suggested for the hydrogenolysis of HCFC- 124. Temperatures required are somewhat higher, raising concerns as to palladium sintering. Multimetal systems reported by Morikawa and co-workers (39)are claimed to reduce such sintering . Besides conventional supports such as carbon and alumina, palladium supported on aluminum fluoride has been shown to be an excellent catalyst by Kellner and Rao (44), affording extremely high selectivity to HFC- 134a.

340


5 . From CFC-I I3 An attractive possibility is the conversion of CFC-113 to HFC-134a in the two-step process shown in Eqs. (15) and (16). The formation of a highly reactive olefin in the first step of the process calls for stable catalysts that are not deactivated by polymerization on the surface. The preservation of

-

+ 1.5Hz CF*=CHF + HF

CFzClCClzF

CFZ=CHF

+

+ 3HC1

CF3CHzF

(15) (16)

the olefinic linkage for the second step of the process has also been somewhat of a challenge. Usual hydrogenolysis catalysts might also saturate the double bond, thereby defeating the purpose of the second step. Very recently, lchikawa and co-workers (45) have reported that a palladium catalyst modified by additives such as salts of thallium, tin, copper, or indium can be used to produce mixtures of CFz=CCIF and CFz=CHF. They report that a bismuth-modified palladium affords excellent selectivity to the desired trifluoroethylene. Earlier work by Lerot and co-workers has also shown that modified precious metals are useful in this transformation (46). The stoichiometric reaction of CFC-113 with metallic zinc to produce chlorotrifluoroethylene is a well-known process. Catalytic approaches to this synthesis are also available. Charcoal (47) as well as supported nickel, iron, cobalt, and chromium operating at elevated temperature (48)give chlorotrifluoroethylene yields in excess of 90%. The selective replacement of the vinylic chlorine by hydrogen while still maintaining the olefinic linkage provides still another entry into trifluoroethylene. The second step in the process given by Eq. (16) is catalyzed by chromium-based catalysts. The effect of oxyfluorides of chromium to carry out the addition reaction of HF to trifluoroethylene at mild temperatures has been described by Von Halasz (49). 6. Isomerization of CF2HCF2H Similar to the hydrogenolysis of CFC- 114a, the hydrogenolysis of CFC114 will afford CF2HCF2H(HFC-134). In principle, one can also make mixtures of HFC-134 and HFC-134a starting with a mixture of CFC-114 and CFC- 114a (50). Three groups of investigators have recently shown that isomerization of HFC- 134 to HFC- 134a proceeds over aluminum fluoride (54, over chrome-based or supported chrome systems (52), as well as over chlorofluorides of alumina (53). As a raw material for isomerization, HFC134 can also be made by the hydrogenation of tetrafluoroethylene as reported by Yoneda and Yanase (54). These authors have shown that palladium on carbon is quite effective for this hydrogenation.


34 1

FIG.4. Some potential routes to HCFC-123.

B. 1,1,1-TRIFLUORO-2,2-DICHLOROETHANE Much like HFC- 134a, several catalytic routes to 1, 1,l ,-trifluoro-2,2dichloroethane (HCFC- 123; CF3CHC12)have been studied. Figure 4, although not all-inclusive, shows a generalized schematic of synthetic approaches that have been developed. Similar types of reactions explored for HFC- 134a synthesis also apply here. General literature, however, suggests that fewer catalyst systems have been evaluated. PCE has been the preferred starting material. 1. Methods Based on PCE The most direct synthesis of HCFC-123 involves the initial addition of HF to perclene, which is followed by two halogen exchange reactions. In fact, a by-product in the manufacture of CFC-113 by the antimony halidecatalyzed process is HCFC- 123. Conventional liquid-phase catalysis using antimony halides in the presence of chlorine have the drawback of producing some chlorinated products for reasons mentioned earlier. However, Gumprecht et at. have recently reported that tantalum pentafluoride is an excellent catalyst for this transformation in the liquid phase (55). Instead, one can also use tantalum pentachloride as reported by Rao (56). Tantalum halide catalysis does not have the inherent drawback of instability of the pentavalent state, in contrast to antimony halide catalysis, and excellent product specificity has been shown. Manzer and Rao (57) have shown that highly fluorinated alumina containing added metals such as cobalt salts are an efficient catalyst system for the vapor-phase transformation of PCE to HCFC-123. Catalyst efficacy seems to depend on the extent of fluorination of the alumina. In the course of the reaction, some HCFC-123 tends to react further to give more fluorinated products, such as HCFC-124 and HFC125, as shown in Eq. (17). These end-products can arise by successive exCC12=CC12

+ HF-

CF3CC12H

+ CFzCClFH + CF2CFzH

(17)

L. E. MANZER AND V. N. M.RAO

342

change reactions. Alternative pathways that involve a series of additionelimination reactions are also possible, as shown in Fig. 5 . In contrast to fluorinated aluminas, chromium(II1) oxide as well as chrome supported on alumina are also useful catalyst systems for this chemistry. For example, Gumprecht et al. (58) have claimed that unsupported chromium oxide containing low levels of potassium, and which is produced by the pyrolysis of ammonium dichromate (59), is an excellent catalyst. Chromium oxide is an active catalyst and as such has a tendency to produce overfluorinated products, depending on operating variables. In contrast, Carmello and Mirano (60) report that chrome oxide supported on aluminum fluoride of specific crystalline phases is also a suitable catalyst. Loss of catalytic activity with time in this system is usually restored by treating the catalyst with oxygen or air at elevated temperatures. Presumably carbon deposits or condensed polymeric fluorinated materials that might hinder catalytic activity and selectivity are removed by such treatment. It is interesting, however, that Cuzzato and Maslero (61) have claimed that a chromium salt (for example, CrC13) supported on aluminum fluoride produces an isomer of HCFC-123, namely HCFC-123a (CF*ClCHFCl), at operating temperatures of about 220-280°C. As shown in Fig. 5 , complex mechanistic pathways are involved, and in this instance the reported product

1112

-

- - - - Elimination --* lsomerization

FIG.5. Chemical intermediates involved in PCE fluorinations.


343

distributions seem to indicate the effect of operating temperature being a key variable in determining desired product selectivity. 2. From Chlorination of HCFC- 133a In principle, a TCE-based route to HCFC-123 is the chlorination of HCFC-133a. This intermediate, discussed in Section II,A, can be chlorinated under a variety of conditions. Thermal, photochemical, and catalytic methods are available. The reaction affords successive replacement of chlorines, as shown in Eq. (18), starting from HFC-143a (CCI3CH3),available from the halogen-exchange reaction of CH3CC13.The first step in the above sequence appears the most difficult as determined by early thermal chlorination studies, the progressive steps becoming faster as chlorine substitution CF3CH3

-

CF3CH2C1

-

CF3CHCl2

-

CF3C CI3

(18)

increases. A catalytic route using metal salts has been described (62). It is necessary to operate this process at relatively low conversion in order to minimize the formation of CFC- 1 13a by overchlorination. 3. From CFC-113a

Similar to the hydrogenolysis of CFC- 114a discussed earlier, the treatment of CFC-113a with hydrogen in the presence of suitable catalysts should produce HCFC- 123. Several stoichiometric liquid-phase processes have been reported in recent literature. Kellner et al. (63) have reported that rhenium catalysts are excellent for the substitution of one chlorine to give HCFC-123. The need for a selective catalyst is essential because the reverse of Eq. (18) may occur on very active and nonspecific catalysts due to overhydrogenolysis. Furutaka et al. have reported that a supported platinum catalyst is also useful for this conversion (64). Selectivities to HCFC123 are high at low single-pass conversions of CFC-113a.

c.

1,1,1,2-TETRAFLUOROCHLOROETHANE

1,1,1,2-Tetrafluorochloroethane(HCFC- 124; CF3CHC1F) is a potential replacement in refrigeration and solvent applications, either as a pure compound or as a component of a blend or azeotrope. Several of the methods and catalytic systems discussed in Section II,B have been applied for the synthesis of HCFC- 124. Research focus has been on optimization of operational parameters and also the development of specific catalysts for product maximization. In addition, one can also use HCFC-123 as a starting material or a pentahaloethane instead of PCE. Manzer and Rao (65) have reported that a cobalt salt on alumina, which is fluorided prior to use, is an effective catalyst for the conversion of pentahaloethanes to HCFC- 123 and

344


HCFC- 124, with minimal formation of HFC- 125. A fluorinated olefin such as chlorotrifluoroethylene can also be used as a starting material. Again, chrome-based systems are quite effective for this HF addition (49). Overfluorination results in HFC- 125. In addition to methods involving the use of HF, selective hydrogenolysis of CFC- 114a has also been reported in some detail. As mentioned earlier in Section II,A,2, regarding hydrogenolysis of CFC- 114a using palladiumbased catalysts, the product usually contains some HCFC- 124, the principal product being HFC- 134a. Selective catalytic systems have been developed that replace just a single chlorine, resulting in HCFC-124 as the major product. These include the use of rhenium (63), platinum (64), and activated carbon (43). D. PENTAFLUOROETHANE Catalytic systems available for the synthesis of pentafluoroethane (CF3CHF2;HFC-125) are essentially similar to those reported for the synthesis of HCFC-123 and HCFC-124. If one starts with PCE, the end point after the initial HF addition is HFC-125. Again, as with HCFC-124, a suitable pentahaloethane such as HCFC-123 or even HCFC-124 can also be employed. Active catalysts seem to contain chrome in some form. Pretreatment of such catalysts have been reported to modify activity. For example, Firth and Foll (15) have claimed that if a chromium hydroxide precursor is treated with steam prior to calcination to chromium(II1) oxide, the activity is superior to the one prepared without such a pretreatment. Swamer (66) has disclosed that chromium oxide gel prepared by special techniques involving the addition of alcohol during the preparative procedure is an excellent catalyst for both HF addition as well as halogen exchange. In keeping with the general theme of hydrogenolysis of a suitable substrate, HFC-125 can also be made from the hydrogenolysis of CFC-115 (CRCFzCl). This latter material is made by the chlorofluorination of PCE or through halogen exchange starting from either CFC-l13,-113a,-114,or - 1 14a. The use of conventional palladium catalysts (36),platinum metals, iron group elements (as well as rhenium) (67), and activated carbon (43) has been described. The chemistry presumably follows a path similar to the one observed with HCFC-124.

E. 1,1-DIFLUOROETHANE 1. From Acetylene

A straightforward route to 1,l -difluoroethane (CHXHF,; HFC- 152a) is the addition of HF to acetylene, Eq. (19). Several catalytic routes are avail-


345

able for the synthesis of this molecule. These include both liquid- and vapor-phase processes. A highly selective route to HFC- 152a involves the HCECH

+ 2HF

--+

CH3CHF2

(19)

liquid-phase reaction catalyzed by BF3 as reported by Burke et al. (68) many years ago. In the place of BF3, one can also use FS03H just as effectively, as suggested by Calfee and Bratton (69). The use of SbFs in conjunction with FS03H allows for the reaction to be conducted at as low a ternperature as 25°C (70). Many variants to the above approaches have been reported in early literature. The use of complexes of BR or its complex salts and the use of organoalkane sulfonic acid are examples. These overcome partly the volatility of BF3 as well as hydrolytic instability. Efficient vapor-phase processes for the synthesis of HFC-152a have also been reported. Depending on the nature of the catalyst used, operating temperatures also vary over a wide range, as opposed to liquid-phase processes. If coproduction of vinyl fluoride is the objective, temperatures in excess of 300°C have been employed. In this latter case, the produced HFC-152a simply eliminates HF to afford the olefin. Pauksch et u1. (71) have reported on some highly efficient catalytic systems involving pseudoboehmite and boric acid, which furnish excellent selectivities to HFC- 152a. Optionally, Fez03 has also been employed as part of the catalytic composition. Aluminum fluoride, which has been used as a catalyst in halogen-exchange reactions, also functions as an efficient HF addition catalyst. The product can be shifted to vinyl fluoride by operating at a higher temperature (72). 2. From Vinyl Chloride The higher cost of acetylene has led several investigators to use vinyl chloride, which is available in large quantities. In a liquid-phase process employing SnC14 as a catalyst, Golubev et ul. (73) have reported the reaction of vinyl chloride with anhydrous HF to produce HFC-152a according to Eq. (20). Very good selectivity has been reported. Although this process does produce a mole of HCl by-product, raw material cost savings may well offset this shortcoming. CHZ=CHCl

+ 3HF

-

CH3CHF2 + HCI

(20)

Vapor-phase processes using vinyl fluoride employ a wide variety of catalysts as well. For example, fluorided alumina or aluminum fluoride has been widely used. The conventional chrome-based systems have also been used with success. Carbon-supported VC13 appears to be (74) a selective catalyst for this synthesis.

346


3. From Dichloroethanes Both symmetrical and unsymmetrical dichloroethanes have been used as starting materials for HFC- 152a. In liquid-phase processes, 1,2dichloroethane is converted in good yields to HFC-152a. Catalysts employed include salts of antimony, such as SbCls or SbFs (75). It is interesting that 1,2-dichloroethane produces HFC- 152a under vapor-phase conditions using chromium oxide catalysis (76). It is likely that under reaction conditions an elimination reaction occurs, leading to vinyl chloride, which then undergoes HF addition followed by substitution chemistry. F. 1,l-DICHLORO1-FLUOROETHANE In essence, the catalytic synthesis of 1,l-dichloro-1-fluoroethane (HCFC141b) can be carried out by following one of the two reaction schemes shown in Eqs. (21) and (22). In the first scheme, HF is allowed to react with vinylidene chloride in the presence of a catalyst. Gumprecht (77) reCHz=CC12 CH3CCI3

-

+ HF

+ HF

CHsCClzF

CHjCC12F

+ HCI

(21) (22)

ported that a suitable catalyst for this process is aluminum fluoride. Operating temperatures are mild and selectivity to the desired product is excellent. Swearingen et al. (78) have shown that a modified process wherein part of the product is in condensed form effectively reduces olefinic contamination, which is difficult to remove from the product. Because of the propensity of HCFC-141b to undergo additional halogen exchange, as shown in Eq. (23), extremely active catalysts usually employed for exchange are not very useful to maximize HCFC- 141b formation. For example, a vapor-phase process CH3CClzF

-----*

CH3CClF2

-

CHsCF3

(23)

using chromium oxide, an excellent catalyst for such exchange reactions, produces greater than 90% CH3CF3(79). Liquid-phase processes usually offer greater selectivity to the desired product, as shown by Feiring (17). A moderated SnC14 catalyst operating in the liquid phase has also been reported (80). Starting from CH3CC13,similar catalyst systems have been suggested for the halogen-exchange reaction. In fact, this reaction has been reported to proceed in the absence of any catalyst, but at an elevated temperature. Again, the desired product HCFC- 141b undergoes further halogen exchange, as shown in Eq. (23). Operational variables are controlled to maximize desired product selectivity.


CF2=CF2 + CHCIs CCI,

-

347

CF3CF2CCb

CClF2CFzCHClz

__t

CCIF2CF2CCI3

CYCFzCHj

FIG.6 . Potential routes to HCFC-225s.

G. DICHLOROPENTAFLUOROPROPANES The dichloropentafluoropropanes have received some attention as possible replacements for CFC-113. Figure 6 depicts some synthetic schemes for producing dichloropentafluoropropanes, of which there are several positional isomers. Two such isomers, CF3CF2CHCl2 (225ca) and CClF2CF2CHClF(225cb), have received some attention recently. They are produced by the familiar Prinz reaction as described by Paleta et al. (81).In the presence of a Lewis acid catalyst such as A1C13, tetrafluoroethylene undergoes a condensation reaction with FCHC12 (HCFC-21). Products reported (81) are HCFC-225ca and HCFC-225cb. The direct synthesis of the individual 225 isomers presents a synthetic challenge. H. DIFLUOROMETHANE Recently, difluoromethane (CH2F2;HFC-32) has received some attention as a low-boiling-point refrigerant. It can be made most directly starting from dichloromethane using halogen exchange. In 1937, Henne (82) reported that the reaction of antimony trifluoride with dichloromethane afforded HFC-32. As with earlier work, this is a stoichiometric reaction necessitating reconversion of the antimony halide. Continuous liquid-phase processes using antimony salts in both oxidation states have been reported (83,84).Early literature also contains descriptions of the synthesis using vapor-phase chemistry (85). Improvements, especially to vapor-phase approaches, have been suggested. Takayama et al. have suggested the use of aluminum fluoride or mixtures of chromium fluoride and aluminum fluoride for this transformation (86). A similar approach, but using vaporized aqueous HF, has also been reported (87). To drive the equilibrium in the desired direction, a sandwich-type operation wherein the HCI reacts with

348


sodium fluoride has been described (88). Here dichloromethane and anhydrous HF are passed through alternating layers of chromium oxide and sodium fluoride. Because the HCl is removed from the reaction zone, high conversions of dichloromethane are obtained. 111.

Summary

This review has described the synthetic routes available to several of the HFC and HCFC alternatives to CFCs under consideration for commercial development, as well as some that have already been commercialized. Much of the information has been gleaned from the patent literature. As can be seen, several options to each product are available. Choice of a commercial process is dependent on several factors, including countries in which they are manufactured and raw material availability. There are still many opportunities for innovation, improved catalysts, and above all a fundamental understanding of catalytic mechanisms associated with these interesting new processes. REFERENCES 1 . Lovelock, J. E., Nature (London)230, 379 (1971). Molina, M. J., and Rowland, F. S., Nature (London) 249, 810 (1974). Rowland, F. S . , and Molina, M. J., Rev. Geophys. Space Phys. 13, 1 (1975). Watson, R. T., Prather, M. J., and Kurylo M. J., NASA Ref. Publ. No. 1208 (1988). Swartz, F., Mem. Cour. Acad. R . Belg. 51 (1895). 6. Swartz, F., Bull. Acad. R. Belg. 24, 309 (1892). 7. Midgley, T., and Henne, A. L., Ind. Eng. Chem. 22, 542 (1930). 8. Holt, L. C., and Mattison, E. L., U.S. Pat. 2,005,713 (1935). 9. Benning, A. F., U.S. Pat. 2,450,415 (1948). 10. Benning, A. F., U.S. Pat. 2,478,362 (1949). 1 1 . Miller, C. B., and Calfee, J. D., U.S. Pat. 2,748,177 (1956).

2. 3. 4. 5.

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ADVANCES IN CATALYSIS. VOLUME 39

Molecular Mobility Measurement of Hydrocarbons in Zeolites by NMR Techniques J. CAR0

H . JOBIC

Zentrum f u r Heterogene Katalyse D-I199 Berlin-Adlershof, Germany

Institut de Recherches sur la Catalyse 69626 Villeurbanne, France

M. BULOW

J. KARGER

The BOC Group Murray Hill, New Jersey 07974, U . S . A .

Fachbereich Physik der Universitat Leipzig 0-7010 Leipzig, Germany AND

B. ZIBROWIUS Department of Chemistry University of Manchester Institute of Science and Technology Manchester M601QD, England

1.

introduction

This article provides a review of the most relevant experimental methods to follow molecular translations and/or reorientations of guest molecules in zeolite pores. The benefit of combining these techniques is illustrated by a

35 1 Copyright 0 1993 by Academic Press, Inc. All rights of reproduction in any form reserved.

352

J. CARO et

al.

series of diffusion studies of hydrocarbons adsorbed in zeolite catalysts. Because of the key role of structure-related diffusion processes in shapeselective catalysis and the unique possibilities of the NMR self-diffusion techniques to investigate such processes, the fundamentals of the NMR selfdiffusion techniques are presented in more detail and examples are given of their application to characterize zeolite catalysts. Zeolitic molecular sieves ex hibit remarkable properties that have made them an interesting topic of both fundamental research and industrial application. In shape-selective catalysis, the correspondence between the diameters of typical reactantlproduct molecules and those of the zeolite channels leads to an intense interplay between the host framework and the guest molecule. Therefore, Weisz introduced the term “configurational diffusion” to describe this special kind of molecular mass transport in zeolites ( I ) . Consequently, slight differences in the structure of the guest molecules can lead to a strong variation in their diffusivities. And vice versa, small changes in the zeolite pore structure (e.g., lattice imperfections such as structural defects, stacking faults, and intergrowths of different zeolite types) can significantly modify the diffusion and adsorption patterns. Recent reviews on mass transport in zeolites are given in Refs. 2 and .?. To understand the principles of shape-selective adsorption and catalysis, a detailed knowledge of the microdynamics of the molecules inside the zeolitic pore system is required. Garcia and Weisz ( 4 ) pointed out the relevance of NMR methods to yield a unifying picture for the phenomena, mechanisms, and magnitudes of “diffusivities.” In the past few years, a deeper insight into molecular motions in zeolites has been achieved, especially by combining investigations of molecular translation with studies of reorientation processes on different time scales (5-14). First we will review the relevant NMR techniques and their most recent developments: (1) pulsed-field gradient (PFG) NMR (315-18) for the measurement of translational molecular self-diffusion and (2) ’H (6,7,lO,19,20) and I3C (9,21,22)NMR lineshape analysis as well as ‘ H NMR relaxation analysis (5,23,24)for the study of molecular reorientation. For selected adsorbateladsorbent systems, we demonstrate the benefit of combining these experimental methods to reveal complex molecular transport phenomena. The results of these NMR methods will be compared with those of other experimental techniques, such as quasi-elastic neutron scattering (QENS) (12,13,25-27), the frequency-response technique (28,29) in its singlestep mode (14.30,31), and sophisticated sorption uptake experiments (39,32,33,113) as well as recent molecular dynamics (MD) calculations (34-41).

353

HYDROCARBON MEASUREMENT IN ZEOLITES

II. Basic Principles of NMR Self-Diffusion Studies A. SELF-DIFFUSION MEASUREMENT BY PULSED-FIELD GRADIENT NMR In a magnetic field of strength B, the resonating nuclear spins precess with the Larmor frequency W = -yB (1) about the direction of the magnetic field, where y denotes the gyromagnetic ratio of the resonating nuclei. In NMR diffusion measurements (15-18,23, 24,42,43), this magnetic field consists of a strong time- and spaceindependent component (Bo)and a superimposed field (ABo)that is assumed to increase linearly in the z direction (i.e., ABo = gz). Consequently, the Larmor frequency becomes a function of the space coordinate ( z ) determined by the direction of the field gradient ( g ) . In PFG NMR, the magnetic field gradient is only applied during two short time intervals of identical duration (6) (“field gradient pulses”) and separation ( t ) . By applying an appropriate radio-frequency pulse sequence (a r / 2 - r sequence for generating the “primary” spin echo or a 7r/2-7r/2-7~/2 sequence for generating the “stimulated’ echo), one can observe the transverse nuclear magnetization, i.e., the component of the total magnetization perpendicular to the constant magnetic field, developing under the influence of these two gradient pulses. In order to calculate the vector sum of the total magnetization, one has to determine the difference A$(i) between the average phase and the actual phase for each of the spins after the second gradient pulse. Denoting the z coordinates for the ith spin during the first and second gradient pulses by zI“ and z!’, one obtains from Eq. (1) in the generally considered limiting case of sufficiently short field gradient pulses

A+(d =

I

I

~ d t -

2nd gradient pulse

o dt =

1st gradient

y6g(Zy’ -

zY)

(2)

pulse

Because each spin contributes to the total transverse magnetization via the cosine of the phase difference A+(’),the quantity observed in PFG NMR is given by the relation *(6g, t ) = M J M , ,

=

CoS”)‘@(Zi

-

ZZ)]P(ZI) P(Z2, Z I , t ) dZi dZ2(3)

where M, and M , denote the magnitudes of the total transverse magnetizations with and without field gradients, respectively, p ( z J dzl is the prob-

354

J. CAR0

et a/.

ability of finding a molecule in a position with a z coordinate between z I and z1 + dz, and P (z2, z l , t ) dz2 denotes the conditional probability that a molecule that starts at the time of the first gradient pulse from position z I has reached a position between z2 and z2 + dz2 after the time t. For molecular self-diffusion in a homogeneous medium, one has P ( Z , zl, t) = ( 4 n ~ t ) - I ’exp[-(z ~ - zJ2/4 Dt)

(4)

Equation (4)corresponds to the solution of Fick’s second law in an infinite space for the initial condition c*(z, 0) = S(z - z1). With Eq. (4), Eq. (3) simplifies to *(Sg, t ) = exp{-y262g2Dt}

(5)

or, using the Einstein relation D = (r2(t))/6t, we can write *(Sg, t ) = e ~ p { - y ~ 6 ~ g ~ ( r ~ ( t ) ) / 6 }

(4)

with ( r 2 ( t ) )denoting the mean square displacement during the time interval t . According to Eqs. (4)-(6), the molecular mean square displacements and thus the self-diffusion coefficients may be determined from the slope of a semilogarithmic plot of the PFG NMR signal T versus (Sg)*. The “observation” time of self-diffusion is the separation between the two field gradient pulses, t . Owing to their relatively large gyromagnetic ratio and to their natural abundance of = 1, protons provide very suitable conditions for N M R self-diffusion studies, but I3C (44), 19F ( 4 3 , and 129Xe(46-48) resonances have also been used successfully in recent PFG N M R studies of zeolites. For diffusion in anisotropic systems such as the MFI structure (ZSM-5/ silicalite-I), Eq. ( 5 ) has to be replaced (3,17,42,49)by

-

*(gJ,A) +

= exp[-y2S2

ZEZt]

(7)

with 5 and 2 denoting the diffusion tensor and the field gradient vector, respectively, instead of the scalar quantities D and g. Thus, in macroscopically oriented systems the principal elements of the diffusion tensor may be determined by varying the direction of the field gradient. In powder systems such as beds of zeolite crystallites, one has to integrate over all possible directions and, strictly speaking, the spin-echo attenuation is no longer exponential. For a diffusion tensor of axial symmetry with Dli > D,,Fig. 1 shows the influence of a diffusion anisotropy on the PFG N M R spin-echo attenuation as simulated in numerical calculations (49). The quantitative analysis shows, however, that for principal tensor elements not too different from each other (i.e., for tensor elements within one order of magnitude), the ini-

355


C

Dn,703

-

9

3-

10

5 1

10-

I

,

1

2

, 3

-

,

,

,

,

,

,

,

4

5

6

7

8

9

10

Y*g26*

(!I/

re( units

FIG. 1. Parameter calculations for the PFG NMR spin-echo attenuation due to anisotropic self-diffusion in a diffusion system of axial symmetry with Dll > D I (49).

tial part of the spin-echo attenuation appears to follow Eq. ( 5 ) with a mean Dzz)equal to one-third of the trace of the diffusivity D = i ( D , + D,, diffusion tensor (17,49). It is only this initial part, however, that is considered generally in PFG NMR experiments. Depending on the signal-to-noise ratio of the NMR signal and the magni, range of mean errors typical of self-diffusion meatudes of D ( r 2 ( t ) )the surements performed on zeolite systems is from 10% up to a factor of 2. All self-diffusion measurements reported in this paper have been carried out by means of the homemade NMR pulse spectrometer FEGRIS operating at a proton frequency of 60 MHz at the Department of Physics of the University of Leipzig (16-18).

+

B, SELF-DIFFUSION MEASUREMENTS BY THE NMR TRACER DESORPTION TECHNIQUE In general, for zeolitic self-diffusion at sufficiently high temperatures, the mean molecular displacements outside the crystal are much larger than those inside the zeolites; that is to say, long-range self-diffusion, D L ~ is . , much faster than intracrystalline self-diffusion, Din,,. For observation times comparable with the mean lifetimes of the adsorbed molecules in the individual crystallites, the spin-echo attenuation can be approximated by the superposition of two exponentials of the type of Eq. (6)

356

J. CAW

et al.

where x ( t ) denotes the relative amount of molecules that have left their crystallites during the observation time, t. As a consequence of the initially assumed large difference between the intracrystalline and long-range mobilities, the root mean square (r.m.s.) displacement for long-range selfdiffusion, (r2(r))f!2,is much greater than that for intracrystalline migration, (r2(t)),',/&. Therefore, the quantity x ( t ) can be determined by a plot of In 4' versus (i3g)2.Moreover, by varying the time interval, t , between the two field gradient pulses, the complete time dependence of x (t) is available. This is the same information as that derived from traditional tracer desorption experiments. Therefore, this modification of pulsed-field gradient NMR has been termed tracer desorption (TD) NMR (50) (see also Refs. 3,16,17,51). There are two equivalent ways to achieve a quantitative analysis of the x ( t ) curves as obtained by TD NMR:

1. Comparison of mean intracrystalline lifetimes. For this kind of data evaluation, it is useful to introduce the first statistical moment

MI =

I-, [I

- x(t)I dt

= Tint,

(9)

as a measure of the mean lifetime, Tnntra, of the molecules inside the zeolite crystals. If the molecular exchange between the individual crystals of the sample/pellet is exclusively limited by intracrystalline self-diffusion (i.e., in the absence of surface barriers), for spherical particles of mean square radius (R') the molecular mean lifetime can be calculated from the crystallite size and the value of the intracrystalline self-diffusion coefficient, D,,,, (measured directly by PFG NMR): diff Tintra

=

~ d i f = f 1

(R2)/15Dmtm

(10)

If the calculated value of if if fa is equal to the measured intracrystalline lifetime, T ] . ~ , ~the , rate of molecular exchange between different crystals is controlled by the intracrystalline self-diffusion as the rate-limiting process. Any ~ ~ with T$Lindicates the existence of transport increase of T , in ~comparison resistances different from intracrystalline mass transport. Under the condithus these resistances can only be tions of TD NMR one has D I B D,,,,, brought about by surface barriers. The ratio Tmta/T$k represents, therefore, a direct measure of the influence of surface barriers on molecular transport. 2. Comparison of (effective) self-diffusion coefjcients. To give a clear idea of the adsorption/desorption retardation due to surface barriers, another means to achieve data evaluation from TD NMR experiments is the calculation of an effective self-diffusion coefficient, DeR.From the fractions x ( t ) of those molecules that have left the crystals during different observa-


357

tion times, t , by fitting corresponding solutions of Fick's law for diffusionlimited adsorption/desorption (52) to the x (t) curves, an effective selfdiffusion coefficient, Deff, can be determined. As an example, for x ( t ) = 0.5 (half of the molecules have left/entered the individual crystals of mean radius R during t ) , the relation D b 5 / R 2 = 0.03 (52) provides a rough estimate for the determination of D,B. Note that the sample is in sorption equilibrium; this means that the same amount of molecules that leave the crystals during t also enter them. The value of D,R as derived from TD NMR can be directly compared with Dint-as measured by PFG NMR. In the absence of a transport barrier near the outer crystal surface, the adsorption/desorption behavior is determined exclusively by the intracrystalline self-diffusion, consequently Deff-- Dint=.The presence of any surface barrier as an additional transport hindrance leads to a decisive retardation of the molecular exchange between the individual crystals of the sample. In this case, D,E < Dintra. C. APPLICATION OF PULSED-FIELD GRADIENT NMR TO ZEOLITEPELLETS One of the advantages of PFG NMR is its ability to provide direct information about the entirety of molecular transport phenomena in pelletized adsorbents and catalysts (316). As an example, Table I presents the transport parameters for methane in granulated zeolite NaCaA (16,53). Three quantities that have a key function for the understanding of mass transfer in granules are illustrated in Fig. 2: (i) the coefficients of intracrystalline self-diffusion, Dint=,and of (ii) long-range self-diffusion, as well as (iii) the molecular mean lifetime, n,lra. The coefficient of long-range self-difffusion is approximated by D1.r. =

Pinter

Dinter

(11)

where pinterand Dinterdenote the relative number and the self-diffusion coefficients of the molecules in the void space between the crystals. For sufficiently high molecular concentrations in the gas phase (2lo'* protons/ cm3), NMR provides an experimental access to both of these quantities: Pinter may be determined through the molecular concentration in the adsorbed (ca) and gaseous (cg) phases and the void fraction 7 = K n d (Knter + V,,,,) of the zeolite bed by the relation pinter =

c g Knter

Ca Kntra

+ c g Knter

-

7 7)+ cg7

cg

Ca(1 -

(12)

where cg and ca(l - 7)+ cg7 are directly obtained from the intensities of the NMR signal observed in the gas phase and in the zeolite bed.

I TABLETABLE I Molecular Transport Parameters for Methane in Granulated NaCa Zeolite" MolPrular Transpori Parameters for Meihane in Granulaied NaCu Zmliie" Quantity Quantity

SymbolSymbol

Intracrystalline self-diffusion coefficient Intracrystalline self-diffusion coefficient Molecular mean lifetime a crystal Molecular mean lifetime inside ainside crystal Molecular mean lifetime in a crystal Molecular mean lifetime in a crystal calculated under the assumption calculated under the assumption that that adsorption/desorption is diffusion adsorption/desorption is diffusion limitedlimited (absence of surface barriers) (absence of surface barriers) Long-range self-diffusion coefficient Long-range self-diffusion coefficient DI r of molecules of molecules in the in the RelativeRelative amountamount Pintcr intercrystalline of the granule intercrystalline space ofspace the granule

Mean free path Mean free path Gas-phase self-diffusion coefficient Gas-phase self-diffusion coefficient Molecular mean lifetime in the intercryMolecular mean lifetime in the intercryof the granule stallinestalline space ofspace the granule Mean molecular diffusion Mean molecular diffusion path in path the in the intercrystalline of the granule intercrystalline space ofspace the granule

A

D,,,, T,nm Tdiff ,"Ira

MethodMethod of determination of determination

PFG for NMR for ( r 24( t(R2)IIZ ) ) ' /G 2 (R2)'12 PFG NMR (r2(t))1'2 Tracer desorption Tracer desorption NMR NMR Through D,.,,theand the crystal radius R, Through D,.,,and crystal radius R, Eq. (10) by Eq. by (10)

Value Value Xmz s-I m2 s-' 1.7 X 1.7 loA9 1.1 ms1.1 ms

0.2 ms0.2 ms

PFG for NMR for ( r 25,( tR2)'12 ) ) 1 / 2 R2)1/2 lo-' PFG NMR (r2(t))llZ 2.2 x 2.2 lo-' xm2 s-I m2 s-' Through the concentrations of methane 0.18 0.18 Through the concentrations of methane in the adsorbed gas phases (as in the adsorbed and gasand phases (as determined by NMR) the void determined by NMR) and theand void '1, by Eq. (12) 7,by Fq. (12) fractionfraction A Through the gas-phase concentration Through the gas-phase concentration and and 7.8 nm7.8 nm the collision cross section, Eq. (13) the collision cross section, by Eq.by (13) Dgas PFG NMR PFG applied NMR applied to the gas phase 3.4 x 3.4 xm2s-' m2s-' to the gas phase 240 ps240 p s Through the values P,.,~,,ofPp,.,,, , " , ~=P , " , ~= Through the of values 'T,n,er 1 -and P T,,~,, , and~by T,,, ~,, by ~(14) Eq. ~ (14) 1 - pnnIeC Eq. l,,,,, Through Through the values T ,~and D, ~ ,~ by 42 p m42 p m the values of T ,ofand ~ D, ,~~by Eq. (16) Eq. (16)

DI r

Pl"ler

" Temperature, K; six CH, molecules per cavity. From16 Refs. " Temperature, 293 K; 293 six CH, molecules per cavity. From Refs. and 16 53.and 53.


359

FIG.2. Mass transport parameters in zeolite pellets as determined by PFG NMR and TD NMR (42).

For molecular mean free paths much less than the mean free diameters of the intercrystalline void space in the zeolite bed, D,,,,is controlled by the same mechanism as in the gas phase, with a self-diffusion coefficient D,. Due to the steric confinement, D,,,, is reduced with respect to D, by a tortuosity factor, 7 6 , with values typically of the order of 2-3. The mean free path can be estimated through the relation A = kTaprra2 = 1/V%,mr2

(13)

where k denotes Boltzmann’s constant, T is the absolute temperature, p is the pressure, and m r 2 is the collision cross section. In the example considered in Table I, A is found to be clearly less than the mean free diameters within the zeolite bed, and comparison of D I, PinterDlnter with PlnterDg yields a tortuosity factor of Tb = D,/D,,te, -- 2.7. From the NMR tracer desorption and self-diffusion data (second and third S fa. In the example given, lines of Table I), one obtains the relation 7,ntra intercrystalline molecular exchange is limited, therefore, by transport resistances at the surface of the individual crystals. Combined NMR and highresolution electron microscopy studies (54) suggest that such surface barriers are caused by a layer of reduced permeability rather than by a mere deposit of impenetrable material on the crystal surface, although that must not be the case in general. It becomes obvious from the equilibrium condition pinter/Tinter = pintra/Tintra

(14)

that any enhancement of 7,ntra due to surface resistances leads to an enhance~ . mean distance covered by the molecules in the intercrysment of T , , , ~ ~The talline space before being captured again by a crystal follows from the cor-

360

J. CARO et

al.

responding notation of Einstein’s equation, (liter)

=

(15)

6Dinter~inter

replacing the two quantities on the right-hand side by Eqs. (11) and (14), this relation can be written as (lfnter)

=

6D1.1. Tintralpintra

(16)

where both quantities D I . ~and , Tintraare directly accessible via the NMR experiments. In most cases, the relative number of molecules in the gas phase is negligible in comparison with the amounts adsorbed, so that pint,, pintra and consequently PinIra = 1. According to Table I, the mean diffusion path covered by the molecules in the intercrystalline space before being captured again by a zeolite crystal is one order of magnitude larger than the diameter of the crystals (3.8 pm). A parameter of some technical relevance is the ratio Tintra/Tgranule of the molecular mean lifetimes in the individual crystals and in the granulated particles. With Rgranule denoting the radius of a granulated particle, by analogy to Eq. (10) one has

*

Tgranule

L1

(R&mule)/

so that either of two quantities Tintraand

15Dl.r.

T~~~~~~~ can

(17)

be determined by NMR.

D. ON THE LIMITS OF APPLICATION OF PULSED-FIELD GRADIENT NMR FOR SELF-DIFFUSION MEASUREMENTS IN ZEOLITES The major limitations for an application of PFG NMR to self-diffusion studies can be summarized in four categories:

1. Most self-diffusion studies have been performed at low temperatures and high loadings characteristic of low temperatures. Due to collisional molecule-molecule interactions, often a drastic decrease of the selfdiffusion coefficient at high sorbate concentrations is found (see later, cf. Fig. 10). In the limit, molecular self-diffusion of small molecules in largepore zeolites can be described by diffusion models for liquids (e.g., diffusion of hydrocarbons in zeolite X by the so-called “free-volume model”; see Section V,B.). In contrast, typical catalytic operations work at temperatures above 275 K and, consequently, with sparse populations of molecules per zeolite cage. However, in the case of structure-related molecular selfdiffusion (as found for hydrocarbons in ZSM-5), the systematic study of molecular self-diffusion as a function of temperature and loading (at subcatalytic temperatures and concentrations usually higher than those in catalytic


36 1

reactions) provides a reliable basis for the extrapolation to catalytic conditions. On the other hand, by applying Fourier transform PFG NMR (this technique is described in Section VI,B,2), in recent PFG NMR studies (135) during the conversion of cyclopropane to propene on NaX at 473 K, the self-diffusion coefficients of the individual components involved in this reaction could be measured directly and simultaneously. Furthermore, the development of high-temperature probe heads for PFG NMR (which allowed measurement of the self-diffusion of n-paraffins in zeolite 5A at 625 K) (136) represents, again, substantial experimental progress toward catalytic conditions in PFG NMR. Another approach to determine diffusion coefficients under catalytic conditions can be achieved most effectively by studying the effectiveness factor in catalytic experiments on samples of different crystal size (the intrinsic reaction rate constant and the equilibrium constant must be known) as done by Haag et al. (55) for the cracking of n-hydrocarbons over HZSM-5. 2. The minimum self-diffusion coefficient measurable by PFG NMR depends significantly on the upper limit of the time interval t , which is limited by the damping of the NMR signal due to transverse nuclear magnetic relaxation. In general, t will be of the order of the transverse nuclear magnetic relaxation time, T 2 , or less (17,56). As an example, for benzene in ZSM-5 at 400 K , one has T2 = 0.5 ms and, assuming as characteristic values of t = 1 ms, 6 = 0.5 ms, and g = 10 T m-', only a self-diffusion coefficient D 2 cm2 s-' can be measured. This lower limit is about 3-4 orders of magnitude above the real diffusivity of benzene in ZSM-5 (cf. Fig. 17), i.e., there is only little hope to measure the benzene self-diffusion in ZSM5 by PFG NMR. However, by using partially deuterated compounds, the proton-proton interaction within one molecule can be reduced and as a result T2 can be prolonged, which gives the opportunity to apply larger t and 6 values (24,56).A further possibility to enhance the observation time t involves the application of the stimulated spin echo (3,17). In this case the spin-echo damping is controlled mainly by the longitudinal nuclear magnetic relaxation time, T I ,which may be considerably larger than T2. 3. In Ref. 56 it is shown that small amounts of highly mobile molecules adsorbed on the outer surface or within intracrystalline cracks will, in general, lead to an enhanced damping of the NMR signal that may be interpreted erroneously as a high overall mobility of the adsorbed molecules if the two-phase character of the signal is not taken into consideration. 4. The molecular root mean square displacement, (r2(t))1'2, of the diffusing molecules during the observation time, t , has to be much smaller than the crystal radius, R , in order to guarantee that the measured r.m.s. displacement reflects the undisturbed intracrystalline self-diffusion. Assuming

J. C A R 0 el al.

362

a lower limit o f t = 0.5 ms and a crystal radius R = 1.5 p m , the diffusivities that can be measured amount to D < lop5cm2 s-l. However, this upper limit can be shifted to higher diffusivities if large zeolite crystals are available.

111.

Principles of NMR Techniques to Detect Molecular Reorientations

OF MOLECULAR REORIENTATIONS BY "c A. DETECTION NMR LINESHAPE ANALYSIS

In 13CNMR spectroscopy, deviations from a Lorentzian lineshape, which is usually obtained in liquids, can be caused by a chemical shift anisotropy (CSA). If a CSA is present, the position of the resonance line depends on the relative orientation of the molecule with respect to the direction of the magnetic field applied (21,22). The superposition of the individual resonance lines results in typical lineshape patterns that can be described by two parameters: the chemical shift anisotropy, As, and the asymmetry parameter, q , respectively. In the case of an axially symmetric CSA tensor, i.e., q = 0, the relation between the resonance frequency, w , and the orientation of the molecule is given by w =

wg

[

6i,"

-

1

A6 -j(3 cos26 - 1)

where wo = yBo denotes the Larmor frequency, 6iaois the isotropic chemical shift, and 6 is the angle between the magnetic field and the principal axis of the CSA tensor, thus describing the orientation of the molecule with respect to the magnetic field. Molecular motions can lead to an (at least partial) averaging of the CSA. The degree of averaging is determined by the type of motion as well as by the ratio of the correlation time, T,, and the value of A6 (21). Therefore, several types of molecular motions give rise to characteristic lineshape patterns. For

one has the spectrum of rigid molecules, i.e., molecules of a fixed orientation with respect to the magnetic field. In the case of a fast isotropic reorientation with


363

the CSA is completely removed and one obtains a Lorentzian line with a line width determined by the remaining interactions. In contrast to n-paraffins, which exhibit no or only a slight I3C NMR CSA, aromatics or hydrocarbons with double or triple bonds show a much larger anisotropy. Therefore, benzene (57) and 2-butyne (14) were chosen as suitable probe molecules to study molecular motions by 13C NMR lineshape analysis. Theoretical lineshapes for different motional states of benzene and 2-butyne molecules are depicted in Figs. 3 and 4. The protondecoupled I3C NMR spectra were recorded by means of the homemade NMR spectrometer UDRIS (University of Leipzig) and a BRUKER MSL 400 (Central Institute of Physical Chemistry, Berlin) at frequencies of 22.6 and 100.6 MHz (9,14,57).

h

fast reorientation about F

A 0:

fast aboutreorientation F

fast isotropic reorientations

,

I

I

Z O

200

150

100

6/ppm

-

I

I

50

0

FIG. 3. Theoretical I3C NMR lineshapes of benzene molecules for different types of motions (58).

364

J.

CARO et

a/.

fast isotropic reorientation 200 150 100 50

b/ppm

-

0 -50

FIG.4. Theoretical "C NMR lineshapes of 2-butyne molecules for different types of motions (14). OF MOLECULAR REORIENTATIONS BY 'H B. OBSERVATION NMR LINESHAPE ANALYSIS

The *H NMR lineshape of solids is determined by the quadrupolar interaction, which can be described by two parameters: the quadrupole frequency, W Q , and the asymmetry parameter, 77 (19,20). The parameter W Q is determined by the electric quadrupole moment of the deuteron and the zz component of the electric field gradient at the deuteron site. For deuterons bonded to carbon atoms, the asymmetry parameter is approximately zero and the z axis is along the C-D bond. In this case, the dependence of the resonance frequency, W , from the orientation of the molecule with respect to the magnetic field applied is given by a relation similar to Eq. (18) (19). As long as the correlation time, r,, of a molecular motion is sufficiently large, i.e., TcWQ

%

1

(21)

the lineshape is not influenced by this motion. For faster motions, however, characteristic deviations occur, because-depending on the type of motion-at least a certain part of the quadrupolar interaction is averaged to zero (19,20). Depicted in Fig. 5 , are theoretical 'H NMR lineshape patterns of perdeuterated benzene for various motions. The 2H NMR spectra were


365

fast reorientation about F

H&

fast aboutF reorientation

fast isotropic reorientations

I

I

100

50

0

-50 -100

FIG. 5. Theoretical *H NMR lineshape patterns of perdeuterated benzene molecules undergoing various types of motions (7).

obtained with the quadrupole echo pulse sequence on the previously mentioned spectrometers at frequencies of 13.8 and 61.4 MHz, respectively (7,9).

c. ANALYSIS OF MOLECULAR TRANSLATIONS AND ROTATIONS BY COMBINED 'H NMR RELAXATION AND NMR SELF-DIFFUSION STUDIES The self-diffusion paths of guest molecules in an isotropic host lattice can be described by a sum of successive individual molecular jumps (5,9,25, 26,59,60),

366

J. CARO et

al.

where D is the self-diffusion coefficient and ( r 2 ( t ) )and (Z2) denote, respectively, the long-range mean square displacement during an observation time, t , and the individual mean square jump length after a residence time, 7 , of the jumping molecule at a sorption site, respectively. In addition to QENS, the value of the mean residence time, 7 ,of a jumping molecule can also be estimated from 'H NMR relaxation analysis. The correlation time T~ of the longitudinal proton magnetic relaxation can be approximated by the reciprocal value of the proton magnetic resonance frequency at the temperature of the T I minimum, i.e., T~ -- w0' (5,23,24). The motional process that controls NMR relaxation is usually assumed to be a molecular reorientation together with or without a simultaneous translational motion. Consequently, depending on the dominating mechanism of magnetic interaction, one has T~ 5 T . Therefore, from combined PFG NMR self-diffusion and NMR relaxation studies, according to Eq. (22) a lower limit of molecular mean jump lengths can be given. The longitudinal and transverse ' H NMR relaxation times have been measured using the NMR spectrometer UDRIS. IV.

A.

Principles of Other Mobility Measurements (Comparison with NMR Data)

MOLECULAR TRANSLATIONS AND ROTATIONS BY QUASI-ELASTIC NEUTRON SCATTERING (QENS)

For hydrocarbons in zeolites, only incoherent scattering has to be considered because of the large incoherent cross section of hydrogen. The neutron intensity scattered follows the incoherent scattering law Sinc(Q,w ) , which is related to the self-motion of protons, where 6Q and fiw denote the neutron momentum transfer and the neutron energy transfer, respectively. Assuming for a molecular system the vibrational, rotational, and translational motions as uncoupled, the incoherent scattering law can be expressed as a convolution product of the individual scattering laws for each motion, which can be examined separately: Sinc(6,

0)

= SP:(~, w )

B S \ ; A ( ~w,) B S ; F < ~ w, )

(23)

For the quasi-elastic region ( 5 2 meV) that is of concern here, the vibrational motions affect only the elastic intensity through a Debye-Waller facis the mean square hydrogen amplitude in tor, exp(-Q2(u2)), where (2) the vibrational modes. The scattering law for rotations can be written as a sum of an elastic peak intensity and a quasi-elastic component:

SE(6,

W)

= Ao(&(w)

+ Sqei(6,

0)

(24)


367

The elastic peak intensity is governed by A o ( ~ )which , is kcown astJhe elastic incoherent structure factor (EISF). The variation of Ao(Q) with Q allows determination of the nature of the rotation (e.g., isotropic or uniaxial rotations). The dynapical behavior of protons is reflected by the quasi-elastic component, Sqel(Q,w ) . This term can be expressed as a superposition of Lorentzians whose widths are related to the average time between jumps. The simplest model of translational motions in three dimensions is the continuous random walk diffusion (Fick’s law) giving a Lorentzian-shaped scattering law, SiF(Q,

W)

1 =To2

DQ’

+ (DQ’)’

where D is the self-diffusion coefficient. This model can be applied quite well to simple liquids at small Q values. In zeolites, however, jump diffusion models (61-63) have been found more appropriate. The broadening values obtained for methane diffusion can be fitted to a molecular jump model with a Gaussian distribution of jump lengths (25,62). Thus, the scattering law becomes a Lorentzian,

whose half-width at half maximum (HWHM) 1 Aw(Q) = ;[l

-

1 exp(-Q2(12)/6)] = -[1 - exp(-Q’Dr)] 7

(27)

is determined by both the mean time, r , between two succeeding jumps, and the mean square jump length, (1’). At small Q values, the HWHM becomes equal to AW(Q) = DQ’ and D can be derived. At high Q values, the half-width approaches the mean jump rate, T-’.For systems with a different distribution of jump lengths, the broadening curves were fitted with another jump model (63). The shape of the quasi-elastic peak is still a Lorentzian, but the HWHM is given by (63)

This model was used for ethane and benzene in NaX. The QENS results were obtained at the Institute Laue-Langevin, Grenoble, using the time-of-flight spectrometers IN5 and IN6 (25-27). The time scale on which the motions can be observed by QENS is determined by the elastic energy resolution, which amounted to 20 and 100 p e V for the spectrometers IN5 and IN6, respectively. Based on an appropriate statistical

368

J. CARO et

al.

data evaluation, broadenings smaller than the resolution can be measured so that motions in the range 10-’o-10-’2 s can be followed. Consequently, diffusion coefficients of the order of 10-4-10-6 cm2 s-’ can be measured. Furthermore, by applying the backscattering technique, a better instrument resolution, of the order of lpeV, can be obtained.

B . DIFFUSION COEFFICIENTS FROM SORPTION KINETICS By fitting appropriate solutions of Fick’s diffusion law to the adsorption/ desorption curves, diffusion coefficients can be derived. Depending on the way the experiments are performed, variable boundary conditions have to be taken into account (52). However, as a rough estimate, one can approximate the initial part of the adsorption/desorption curves by the so-called law accounting for the changing boundary conditions (64):

d

where PO,p,, and pmdenote the pressures at the beginning of sorption, at actual time t , and at sorption equilibrium, respectively. A denotes the external surface area and V is the volume of the crystals. Other methods of data evaluation are provided by the method of statistical moments (32)or the full analysis of uptake curves for complex processes by means of the software ZEUS (Zeolite Uptake Simulator) (65). In contrast to all the other techniques considered in this paper, in sorption experiments molecular migration is observed under nonequilibrium sorption conditions. Therefore, instead of self-diffusivities, D, in this case transport diffusivities, D t , are derived. It is generally assumed (see, e.g., Refs. 366) that the “corrected” diffusivities, Do,

D O = D,d [In c ( p ) ] / d[In p ]

(30) with c (p) denoting the sorbate concentration in equilibrium with the sorbate pressure p, should be of the order of the self-diffusivity, D. In the following discussion, sorption experiments will be discussed only in terms of the “corrected” diffusivity, Do. Two experimental methods to follow sorption kinetics have been applied. In the single-step frequency-response method, the frequency-response system at Imperial College, London, has been used. Fast volume perturbations that cause stepwise-like pressure changes have been achieved by activating an electromagnet that has copper bellows attached (3031). In another sorption kinetics technique installed at

369


the Central Institute of Physical Chemistry, Berlin, the pressure changes are initiated by dosing via opening a valve (32). In this mode, adsorption/ desorption processes caused by the pressure jump are followed piezometrically in a constant volume system. V. Combined Application of Different Experimental Techniques to Study Molecular Translations and Rotations of Hydrocarbons in Zeolites

A. COMPARISON OF SELF-DIFFUSION COEFFICIENTS DETERMINED BY PULSED-FIELD GRADIENT NMR AND QUASI-ELASTIC NEUTRON SCATTERING 1 . Methane in ZSM-5

Figure 6 represents typical plots of the spin-echo intensity in PFG NMR experiments. Comparing the slopes of these representations with those of standard liquids, one obtains the mean self-diffusivities, which are found to decrease with increasing sorbate concentration (5,12,16,59,60). It appears from Fig. 6 that within the accuracy of the measurement no deviation from a single exponential decrease may be observed. A comparison of the experimental spin-echo attenuation (Fig. 6) with the results of numerical calcula-

100 80 ffl

60

-h

o

TT

!t

3-

t

2o 10

L

1

2

3

-

4

5

6

7

8

9

1

0

f g 2 6 2 ( t ) / rel-units

FIG.6. PFG NMR spin-echo attenuations due to self-diffusion of CHI in ZSM-5 at 250 K for the loadings (CH4 per u.c.): 0, 4; 0 , 8; and 0 , 12 (49). For calibration, neat water (0) and liquid ammonia (A) are included, with self-diffusion coefficients of 2.3 X lo-' cm2 s-' (67) and 1.5 X cm2 s-' (68).

370

J. CARO et

al.

tions (Fig. 1) shows that for methane in ZSM-5, the principal tensor elements differ from each other by less than a factor of 5-10 (49). This result is in agreement with recent PFG NMR measurements with oriented ZSM-5 crystals, where the mean self-diffusivity of methane averaged over the straight and sinusoidal channels (i.e., in the x , y plane) has been found to be larger by a factor of 3-5 than the self-diffusivity in the c direction, perpendicular to those (42,43).For the self-diffusion of methane in ZSM-S/silicalite-I, the MD simulations were in very satisfactory agreement with the experimental PFG NMR and QENS self-diffusion data (cf. Refs. 34-38). Furthermore, the degree of a mass transport anisotropy in the MFI framework as predicted by MD simulations (34-38) is compatible with the experimental findings (42,43,49). In the QENS experiment (Fig. 7), the broadening of the elastic peak as a function of Q 2 is found to deviate from a straight line, so that a jump diffusion model has to be considered. The mean jump lengths are found to be larger at small loading and high temperatures (12,25,69).For methane in ZSM-5, the mean jump lengths calculated from QENS for small and high loadings at 200 K amount to 1.04 and 0.87 nm, respectively. The selfdiffusion coefficients obtained from PFG NMR and QENS are presented in Table 11. These data yield an activation energy for the self-diffusion of methane in ZSM-5 of 4-5 kJ mol-’ both for PFG NMR and QENS (12,69). Furthermore, Table I1 presents diffusivities determined by sorption uptake/ desorption measurements (single-step frequency-response technique) (30). Table I1 shows that the self-diffusion coefficients measured independently by PFG NMR and QENS are in excellent agreement (12,69).Owing to the different time scales of the two methods, this finding is not trivial. Whereas in QENS, molecular translations are measured up to 6 nm, in PFG NMR the mean molecular displacement may amount to several micrometers. Thus it

.,

0.2

0.L

-

0.6

0.8

02/i3-2

FIG. 7. Elastic peak broadening versus Q’ for CH, in ZSM-5 at 200 K for the loadings (CHJ per u.c.): A, 2; 4;and V, 8 (12.25.69).

37 1


TABLE I1 Self-DiffusionCoef$cients of Methane in ZSM-5Measured by PFG NMR and QENS Compared with Sorption UptakeiDesorption Diffusiivities" At T = 200 K(x

Loading (CR/u.c. 1 1.5 2 2.8 4

8 12

DQENS

DNMR

-

-

2.8

-

-

-

2.5 2.9

3.7 2.8 2.5

-

cmZS K I )

DSVRP 0.6 0.5 -

At T = 250 K ( X DQENS

cm2 s-')

DNMR

5.0

-

5.9

-

-

6.3 5.2 4.5

-

-

NMR and QENS data from Refs. 12, 25, and 69; diffusivity data from Ref. 30.

turns out that the migration over both ranges of displacements is determined by the same mechanism of intracrystalline diffusion. The diffusion data obtained by sorption uptake/desorption are smaller by a factor of 5 than the directly measured self-diffusivities, which is a much better agreement than was stated previously in comparative studies of the same system (70,71), where discrepancies up to 5 orders of magnitude have been obtained. The remaining difference might be attributed to the fact that in sorption experiments a slowing down of uptake rate due to the finite rate of adsorption heat dissipation cannot be entirely excluded. With QENS, both the rotational and translational motions of CH4 can be observed. It was found that the rotational motion of CH4 in ZSM-5 can be described by an isotropic rotational diffusion model with a rotational diffusion time constant, D,(72). The values of D,for CH4 adsorbed at 250 K in ZSM-5 are of the order of 5 X 10'' s-'. In MD simulations at 400 K , D, was found to be of the order of 10l2 s-' (73). This difference is due to the fact that a radius of gyration of 0.15 nm was used in the computer fits of the QENS profiles. This radius is intermediate between a simple rotation model, with 0.11 nm for the distance between the protons and the center of mass of the methane molecule, and the radius of the channel in which the molecule performs oscillations.

2. Ethane in NaX Figure 8 shows a comparison of experimental and calculated energy spectra in the QENS experiments for ethane in NaX at different Q values. Table I11 contains the molecular mobility data determined by QENS. In Fig. 9 the self-diffusion coefficients obtained by QENS (cf. Table 111) are com-

n

;L

u 0 l

I

-0.4

-

-02

0

0.2 0.4 hw/rneV

0.6 0.8

FIG. 8. QENS energy spectra for ethane in NaX at different Q values (253 K , 4.3 C2H6per supercage) (69). TABLE 111 SeEf-DiffusionCoefficients, Mean Residence Times beween Jumps, and Root Mean Square Jump Lengths for Ethane in NaX" Loading (CzHs per supercage)

D cm2 s-')

(x

1.3 4.3 5.8

( x lo-" s)

(12)1'2 (nm)

1.36 1.48 2.22

0.66 0.41 0.44

7

5.4 2.5 L .5

" As determined by QENS at 253 K (69).

N

. E u

A0

2-

n

t

I

I

1

2 3 1 5 6 loading/ C2H6 per supercage

7

FIG.9. Comparison of the self-diffusion coefficients of ethane in NaX at 253 K measured by PFG NMR (0) (59) and QENS (A) (69).

373


pared with the corresponding PFG NMR data (59). As for methane/ZSM-5, agreement is observed in both the absolute values and the concentration dependence. Furthermore, a very similar trend in the decrease of the mean molecular jump lengths with increasing sorbate concentration, as shown in Table I11 for ethane/NaX, has been found by NMR methods for propane/ NaX (5). All the QENS spectra were fitted both by means of the scattering law for isotropic rotation and a jump diffusion model (63) according to Q . (28). The radius of gyration was found to vary between 0.18 nm at high loading and 0.22 nm at low loading. As is the case for methane in ZSM-5, this finding is explained by the fact that the motion does not strictly correspond to a rotation about a fixed center of mass because the position of the center of mass changes during the oscillatory state. The rotational diffusion time constant, D,, was found to be of the order of 10” s-l at 253 K.

B. COMBINED ‘H NMR RELAXATION AND PULSED-FIELD GRADIENT NMR SELF-DIFFUSION STUDIES FOR PROPANE IN ZSM-5 AND NAX For CHJZSM-5 and C2HdNaX, as was previously discussed, the selfdiffusion coefficients decrease monotonically with increasing concentrations. Figure 10 shows that this holds true also for propane in both ZSM-5 and zeolite X. ZSM-5 and zeolite X show similar concentration dependencies of the molecular self-diffusion, so that-at first glance-one might expect them to be caused by identical microdynamic mechanisms. However, according to Eq. (22), the translational mobility can be reduced by both decreasing mean square jump lengths, (1*), and increasing time intervals, T , between succeeding jumps. In fact, the relaxation time measurements (Figs. 1 1 and 12) indicate a decisive difference in the molecular transport mechanisms: for NaX, with increasing sorbate concentration, the temperatures of the TI minima (and associated with it the value of the correlation time, T ~ , of molecular motions) remain unchanged (cf. Fig. l l ) , whereas in ZSM-5 the TI minima are drastically shifted to higher temperatures (cf. Fig. 12), thus indicating a significant decrease of the jump rate. Table IV shows that for ZSM-5 at small and medium pore filling, the mean jump lengths are found to be invariable and amount to 1 nm, which is of the order of the distances between adjacent channel intersections of the MFI framework. Consequently, the concentration dependence of the selfdiffusion coefficients in ZSM-5 has to be explained by structure-related jumps, with jumping rates decreasing at elevated loadings. In contrast, for zeolite NaX the reduction of the molecular translational mobility with increasing sorbate concentration is mainly due to the reduction of the jump lengths, with jump rates practically unaffected by concentration. For hydrocarbons in NaX, the jump lengths are found to be corre-

-

f

2

10-71 L -

2 -

J

lo-*:L

FIG. 10. Self-diffusion coefficients of propane in zeolites NaX

(a),CaA (O), and ZSM-5

(a)at 300 K (5) as a function of sorbate concentration expressed by the number of C3Hs per

24 (Si + Al) atoms (representing one large cavity in the cases of zeolities X and A or one channel intersection with the adjacent pore segments in the case of ZSM-5). Data recently obtained in neutron scattering experiments for ZSM-5 (v) (74) are included.

75 I

temperature/ O C c0 -50 -100 -150 I

r

1

I

A

42-

E".102 +-

I

L-

1011 , 2

,

,

,

,

,

,

,

3

4

5

6

7

8

9

FIG. 11. Longitudinal relaxation times (TJ of propane in NaX for 1.5 (A), 2 (o),4 (o), 5 (01, and 6.5 ( 0 )C3Hg molecules per large cavity measured at a proton resonance frequency of

90 MHz (5).

375

HYDROCARBON MEASUREMENT IN ZEOLITES temperature

75

I

1 2

0 -50

/OC-

-100

I

3

4 5 -IO~K/T

6

-150

L

I

/

7

0

9

FIG. 12. Longitudinal relaxation times ( T I )of propane in ZSM-5 for 1 ( A ) , 1.5 (A),2 (01, 2.5 (m), 3 (o), and 3.5 (0).C3Hx molecules per channel intersection measured at 90 MHz ( 5 ) .

TABLE IV Estimated Vulues for Minimum Meun Jump Lengths in ZSM-5 and NuX" Concentration

NMR

[molecules per 24 (Si + Al) atoms]

Temperature of the TI minimum (K)

0.4 I .o I .5 2.0 2.5 3.0 3.3

150 175 210 270 315 360

I .5 2.0 4.0 5.5 6.5

135 135 I35 135 135

QENS (1 *)I/>

D (X

lo-' cm2 s - ' )

(X

lo-' nm)

PropaneiZSM-5 13 12 13 14 6 5

11.7 11.3 11.7 12.2 8.0 7.3

PropanelNaX 5.8 4.0 2.5 0.6

7.8 6.5 5.1 2.5

0.1

-

(/ 2 ) 1/ 2 (X

10-I nm) 9.2 8.4 6.6

1.o

" Obtained at the temperature of the TI minimum according to Eq. (22) ( 5 ) in comparison with QENS data (74).

376

J. CARO et

al. p e q ~,bid

I b l volume compression- adsorption

2-butyne lsilicalite -I 323 K

Pequ.,reo.: Pstart

l'Lgly I

[a) volume expansion

,I

0 0.1 0 2

"

NS

>

I , , .

-desorption

,

10 20 30"W 100

01

a2

"

D in

30 "goim

t/S

FIG.13. Kinetics of desorption (a) and adsorption (b) of 2-butyne on silicalite-I recorded by the pressure response following the rapid square-wave volume expansion (a) and compression (b). Experimental conditions; 1.5 Torr of 2-butyne at 323 K; 0,without adsorbent; 0 , ZSM-5 present in the system (14).

lated with the molecular free volume, so that the reduction of mobility with increasing concentration may be explained by a modification of the freevolume theory (59).

c.

SORPTION KINETICS AND "cNMR LINESHAPE FOR 2-BUTYNE IN ZSM-5 ANALYSIS

The adsorption and desorption kinetics of 2-butyne on ZSM-5 following square-wave pressure changes are shown in Fig. 13. In contrast to the results obtained for n-butane (14), sorption kinetic curves of 2-butyne on ZSM-5 cannot be described by only a single diffusion coefficient (cf. Fig. 13). After a fast initial sorption process, further adsorption proceeds much more slowly. For 2-butyne the time constant of the succeeding process of comparatively slow adsorption/desorption kinetics is larger by at least one order of magnitude than that for the initial fast sorption process. Applying the V? law to this initial region of sorption kinetics [cf. Eq. (29)], the diffusion coefficient of 2-butyne was found to be larger than 3 X em's-'.


I

1

1

I

I

200 150 100 50 b/PPm

I

I

-

0 -50

377

1

FIG.14. I3C NMR spectrum of 2-butyne sorbed on silicalite-I (1.35 mmollg) at room temperature. The spectrum was obtained by single-pulse excitation (14).

This value is higher than the corresponding value of -0.7-2.5 X cm2s-‘ for n-butane (14) obtained by the same experimental technique.* Our results (14) indicate that the rate of penetration into the molecular sieve lattice is higher for the “stiff’ 2-butyne than for the “flexible” nbutane. Hence, sorption uptake of 2-butyne on ZSM-5 turns out to be a coupled diffusion-rearrangement process. In the initial part, the adsorption/ desorption rates are controlled by intracrystalline diffusion. The final part of the sorption process, leading to a higher sorbate density, is determined by the rate of molecular reorientation of the sorbed molecules finding their optimum sorption arrangement. For ZSM-5, the latter process is especially slow for the relatively stiff 2-butyne but is fast for the flexible n-butane. Similar experimental findings, wherein the sorption kinetic data could not be fitted by a simple diffusion model, have been observed e.g., in the case of xylene isomers (9,78) and picoline (79) on ZSM-5 as well as for benzene on NaX (80) and for p-xylene in microporous gallosilicate (141). Further information on the molecular transport of 2-butyne in ZSM-5 can be derived from l3C NMR lineshape analysis. A comparison of the measured I3C NMR spectrum (Fig. 14) with the theoretical ones (Fig. 4) leads to the following conclusions: 1. The spectrum of sorbed 2-butyne is quite different from the calculated spectrum for molecules fixed in space. This indicates that the directions of

* For comparison, in recent studies the self-diffusion coefficient of n-butane in ZSM-5 was found to be of the order of 10-s-10-6 cm2 s-I: 1.4 X cm2 S K I[two C4HI0per unit cell cm2 s-’ (two C4H10per u.c., 298 K) (1 u.c = 96 SitAL atoms), 353 K] by QENS (73, by PFG NMR (106), and I .7 X lo-’ cm2 S K I (one C4HIoper u.c., 298 K) by M D simulations (77).

378

J. C A R 0

et a / .

the molecular axes are not fixed, i.e., the angle between the axis of the molecule and the magnetic field is changed by molecular motions. Consequently, the motion of 2-butyne molecules is not restricted only to the straight channels of ZSM-5. = 77 ppm, one has to con2. Because there is no Lorentzian line at clude that the sorbed molecules do not perform fast isotropic rotations in the channel intersections, a situation also found for benzene in ZSM-5 (7,9,57,81). Obviously, energetic restrictions prevent such motion, which might have been considered for geometrical reasons. 3. A possible interpretation of the experimental results is a 90" flip motion of the 2-butyne molecules. In the fast limit, such a motion generates a pseudoaxially symmetric CSA tensor with the parameters 611 = 56 2 11 ppm and 6, = 157 2 6 ppm (14). The corresponding lineshape is included in Fig. 4. Considering the ZSM-5 framework, this 90" flip can be a molecular jump from a straight into a sinusoidal pore segment or vice versa. From Eq. (20) it follows that the values of the correlation time must be much less than 10 ps. 4. On the time scale of the NMR experiment, the sorbed 2-butyne molecules behave as a single phase. There are no molecules moving exclusively through the straight channel system and no molecules performing isotropic rotations in the channel intersections.

D. NMR LINESHAPE ANALYSIS, PULSED-FIELD GRADIENT NMR, QUASI-ELASTIC NEUTRON SCATTERING, AND SORPTION KINETICS ON AROMATICS IN ZSM-5 AND NAX 1. Benzene in ZSM-5 I3C NMR lineshapes of benzene molecules sorbed on ZSM-5 are shown in Fig. 15. For high sorbate concentrations and low temperatures, powder spectra are obtained; these spectra are characteristic of an axially symmetric shielding tensor. This lineshape is known to be generated by a fast motion of the benzene molecules about their C-6 axes (cf. Fig. 3). With increasing temperature and decreasing loading, the lineshape changes into a Lorentzian one, i.e., an additional molecular motion becomes fast enough to average out the CSA. For perdeuterated benzene in ZSM-5, the dependencies of the *H NMR lineshape on both temperature and loading are given in Fig. 16. From a comparison of the spectra for high sorbate concentration and low temperature with the theoretical lineshapes for different motional states (cf. Fig. 5), it has to be concluded that the sorbed benzene molecules perform fast reorientations about their C-6 axes, a conclusion identical with the findings from I3C NMR.

-

loading / & H e per

U.C.

6D

3.2

8.0

260 K

E3

e

e

aI

e

0, c

t FIG.15. I3C N M R spectra of benzene sorbed on ZSM-5for various temperatures and concentrations. The spectra at 200 K were taken using signal enhancement by cross-polarization (9,57). 16

-

loading/C6D6 per u c .

56

40

7.6

1"1

200 K

P

5

167K

125 K 0

35

70

-

0

35

70 "W-Wol/;

.d-.,

35 ] / kHz

70

I

,\35

70

FIG. 16. *H N M R spectra of perdeuterated benzene sorbed on ZSM-5for various concentrations and temperatures. The spectra were obtained using 3.5-ps7r/2 pulses and a 60-ps pulse spacing (7,9).

380

J. C A R 0 et

al.

As follows from Fig. 16, a reduction of the sorbate concentration causes an effect similar to that of a rise in temperature. From I3C and 'H NMR lineshape analyses of benzene sorbed on ZSM-5, the following conclusions can be drawn (7,9,57,58):

1. The sorbed benzene molecules perform fast reorientations about their C-6 axes. Even at 125 K , the correlation time of this motion is much shorter than 1 e s . 2. Superimposed on this C-6 reorientation, there are jumps of the benzene molecules between a limited number of sorption sites. These sites allow only distinct orientations of the hexad axis of the benzene molecule with respect to the crystal system. The mean residence time, T ~ between , two succeeding jumps decreases with increasing temperature and decreasing loading. The rotational motion of benzene in ZSM-5 has also been characterized by QENS (81). On the time scale of the QENS experiment (=IO-"s), which is much shorter than in NMR, the spatial distribution of the rotating hydrogen atoms is found to be temperature and loading dependent. At low temperatures (90 K), there is a progressive blocking of the rotational motion, a finding that suggests benzene-benzene interactions. A model of uniaxial rotation in an N-fold cosine potential has been used to interpret the low-temperature patterns. At high temperature (400 K) and low loading (three C6Hdu.c.), the rotational motion converges toward a uniaxial C-6 rotation, but it does not reach the spherical rotational model, i.e., the spatial distribution of the protons remains limited to a circle of a radius R = 0.25 nm and does not reach a sphere. The correlation time obtained at 300 K is 3.2 X lo-'' s, Le., as fast as in the liquid. Simulations of 13CNMR lineshapes have shown that experimental spectra that appear to result from a superposition of two different lines (cf. Fig. 15) can be explained by the above-mentioned molecular jump model. Analogous conclusions were drawn from macroscopic sorption kinetic data (82). From the experimental I3C NMR lineshapes, a mean residence time T~ of 20 and 150 ps for a concentration of six molecules per U.C.at 250 and 200 K , respectively, was derived. Provided that these jumps detected in I3C NMR spectroscopy are accompanied by a translational motion of the molecules, it is possible to derive self-diffusivities D from the mean residence times. Assuming the diffusion path of a migrating molecule as a sum of individual activated jumps, for isotropic systems the relation (1') = 6D7, is valid, where (Z2) denotes the mean square jump length. Following experimental and theoretical studies on the preferential sorption sites of benzene molecules in the MFI framework (83-90), in our estimate the mean distance between adjacent sorption sites is assumed to be 1 nm.

38 1


With the above values for the mean residence time 7j and its activation energy E, = 17 2 4 kJ mol-I, from the I3C NMR data a self-difisivity D = 1-5 X lo-'' cm2 s-' can be predicted for a temperature of 303 K and a loading of six molecules per U.C.(9). It should be noted, however, that this procedure is only justified if the activation energies of the jumping frequency and of the self-diffusion coincide. Furthermore, from 13C NMR a decrease of the molecular mobility with increasing sorbate concentration has to be expected. The concentration patterns of Do obtained from sorption uptake experiments for two ZSM-5 samples are shown in Fig. 17 (9). The data are in satisfactory agreement with the I3C NMR estimates in the absolute values, in the activation energies, and in the concentration dependencies. Moreover, the activation energy data for benzene are found to agree with the results of computer simulations (87).

2 . Benzene in Zeolite X Accounting for the influences of external heat and mass transfer resistances in limiting the sorption rates, in many instances reasonable agreement between diffusion data from sorption experiments and PFG NMR may

lo-*

1

t

a.

10'~

A x

v* v c

A

8

Ll 10'10

N

E

* = I

303K 274 K I

I

1

2

-

I

3

I

4

loading / C s H 6 per u.c

FIG. 17. Concentration dependence of the self-diffusion coefficient DOfor benzene in two ZSM-5 samples. Filled symbols, %/A1 135; crosses, Si/AI > loo0 (9). i=

382

J. CARO

et al.

be obtained (30,3f,91-94). However, there are also numerous welldocumented experimental studies showing differences of up to two orders of magnitude (95). An especially large number of investigations have been devoted to benzene in zeolite NaX. For this system, sorption uptake measurements by different research groups revealed both agreement (10,13,93) and disagreement (95,96) with the PFG NMR self-diffusion data. In Fig. 18 the self-diffusivities obtained by different experimental techniques are compared. It appears that in both the absolute values and the trends in the concentration dependence, the QENS data, the PFG NMR results, and the data derived from sophisticated uptake experiments using the piezometric or single-step frequency-response techniques agree. Nevertheless, disagreement with some sorption results has to be stated. Additional information on the molecular reorientation of benzene in zeolite X has been obtained by QENS and *H NMR lineshape analysis. With neutron scattering, it has been found that the rotational motion of benzene in NaX corresponds to a uniaxial reorientation about the C-6 axis, with jumps of 60". The mean time between successive jumps about the C-6 axis at 458 K was found to be 1 . 3 X s.

1

0

0 1

-

9

2 3 4 5 loading/C6 H6 per supercage

1

-

FIG.18. Self-diffusion coefficients of benzene in NaX at 458 K: PFG NMR, 0 (97) and 0 (92); QENS, A (13);deduced from *H NMR lineshape analysis, 0 (10). Comparison with zerononequilibrium measurements: v, sorption uptake with piezometric control (93); length column method (96); n , frequency-response and single-step frequency-response technique (%). The region of the results of gravimetric measurements with different specimens (92) is indicated by the hatched areas. Asterisked symbols represent data obtained by extrapolation from lower temperatures with an activation energy confirmed by NMR measurements.

+,

383


TABLE V Translational Mobility Data for Benzene in Zeolite NaX" Loading (CbH6per supercage)

(X

0.8 I .3 2.0 a

D cm2 s-') 7

4 1.9

(12)1'2

7

(nm)

( x to-" s)

0.46 0.35 0.24

5.0 5.1 5.0

Obtained at 458 K by QENS (69).

Results obtained by QENS for the translational motion are summarized in Table V. As is the case for light hydrocarbons, the decrease of the selfdiffusion coefficient for benzene in NaX with increasing sorbate concentration is mainly due to reduced mean jump lengths rather than to increasing mean residence times between succeeding jumps. The same conclusions as drawn from QENS data can be drawn from 'H NMR lineshape analysis. The temperature influence on the 2H NMR lineshape of benzene sorbed on NaX is shown in Fig. 19. As in the case of 250 K

222K

200 K

F

182 K

2

c

eaJ

E

167 K

aJ

c

t 125 K

0

35

70

-

-[~w-w,)/~I~I/

kHz

FIG. 19. Temperature dependence of the 2H NMR lineshape for perdeuterated benzene sorbed on NaX (loading: 4 C6D, per supercage) (7).

J. CARO et

384

al.

benzene sorbed on HZSM-5, there is clear evidence (cf. Fig. 5 ) for a fast C-6 reorientation of the sorbed molecules with a correlation time T~ 1 ps. On the other hand, our spectra show that on the time scale of the NMR experiment, the benzene molecules are fixed at their sorption sites. According to Eq. (21), for all motions that change the orientation of the hexad axes of the molecules (e.g., translational jumps between sorption sites), even at T = 200 K, the values of the correlation time, T,, are much larger than the reciprocal value of the quadrupole frequency, w Q , i.e., T~ % 1 ps. This may be caused by the interaction of benzene molecules with sodium ions via their 7~ electrons. In QENS investigations (99), the sorbed benzene molecules were found to interact in Na-mordenite predominantly with the sodium ions, resulting in a uniaxial C-6 reorientation with a correlation time T~ 5 2 ps at 300 K. The correlation time obtained by QENS for the C-6 reorientation of benzene in NaX is of the same order of magnitude. At temperatures above room temperature, Lorentzian 2H NMR lineshapes are observed for perdeuterated benzene in NaX (7). Therefore, in this range of temperature, the translational mobility of benzene in NaX is expected to be significantly higher than in HZSM-5. This expectation was fulfilled indeed in sorption kinetics measurements. The diffusion coefficient Do for benzene on NaX is -2 x lo-' cm2 s-' at 353 K (93),whereas under similar experimental conditions a value of the order of 10-l' cm2 s-' was obtained for benzene on ZSM-5 (100).

*

3. p-Xylene in ZSM-5 The 'H NMR spectrum of p-xylene-d4 (all ring protons are replaced by deuterons) sorbed on ZSM-5 is shown in Fig. 20. From the observed quadrupole splitting of about 140 kHz, it follows that for all molecular mo-

I . . . . I . . . . l . . I . I . . I . l . . I . I . . . . I . . I . I I . . . ,

100

0

- 100

-

[(w-w~)/ZX]/~HZ

FIG.20. *H NMR spectra of p-xylene-d4 sorbed on ZSM-5(7.2 C8D4H6 per u.c., T = 295 K) obtained by composite 7r/2 pulses (9,58).

385


tions at room temperature the correlation times are much larger than 1 ps. Taking into account the pulse interval and the duration of the composite 7r/2 pulses used [40 and 71 ps (9)], the lower limit of the correlation time, T,, of molecular motions that change the orientation of the molecular plane with respect to the magnetic field is estimated to be T~ > 100 ps. A fast rotation about the para axis of the molecule, as deduced from 13CNMR spectra at 310 K (101), can be excluded. Our findings are in accordance with those of a previous 2H NMR study (102). The additional intensity in the center of the spectrum in Fig. 20 can be caused by molecules in another motional state (e.g., molecules outside the pore system performing fast 180” flips) as well as by traces of methyl deuterons as impurities. Assuming for p-xylene a jump mechanism similar to that of benzene, the same procedure can be applied to derive an upper limit for the translational diffusivity. At high sorbate concentrations, the p-xylene molecules in ZSM-5 should be localized in the channel intersections and the pore segments of the sinusoidal channels, based on theoretical (103) and XRD (103,104) studies. With this arrangement of sorption sites and the derived lower limit for the mean residence time (T~2 100 ps), one can predict a diffusion coefficient D I lo-” cm2 s-’ at room temperature and maximum loading, which corresponds well with the uptake data shown in Fig. 21. As in the case for benzene in ZSM-5, the translational motion of p-xylene in ZSM-5 was too slow to be measured on the IN6 neutron spectrometer

A

N

lo-”

i

A

A

A A

1

2 3 looding/C8HI0 per u.C.

.,

L

FIG. 21. Concentration dependence of the diffusivity DO for p-xylene on three different 50; Si/AI = 135; T, Si/AI > 1000 (9).

ZSM-5samples at 363 K: A, Si/AI

i=

386

J.

CARO et

al.

within the temperature range 90-380 K (105). No change of rotational motions was observed for a loading of four molecules per U.C.On the time scale of the neutron experiment (=lo-’’ s), the aromatic ring appears to be immobile and only one of the two methyl groups is rotating. The other methyl group seems to be blocked, most likely due to interactions with the framework. Theoretical and structural studies (103,104) suggest that for a loading of four molecules per u.c., the p-xylene molecules occupy the channel intersections with the methyl groups parallel to the straight channel. In this case, both methyl groups would be equivalent. The neutron results indicate, however, that the molecules located in the channel intersections have to be slightly inclined, leading to different interactions of the two methyl groups of the same molecule with the framework. For concentrations greater than four molecules per u.c., p-xylene becomes highly immobile ( 9 ) .

OF N-HEXANE IN ZSM-5 STUDIED BY E. DIFFUSION PULSED-FIELD GRADIENT NMR, QUASI-ELASTIC NEUTRON AND SELECTIVE SORPTION UPTAKEKINETICS SCATTERING,

For both QENS and PFG NMR spectroscopic methods, the measurement of the translational molecular self-diffusion of n-hexane in ZSM-5 represents the present limit of detection. A value of 4.5 X lop6cm2 s-l is obtained by QENS at 300 K (75); by PFG NMR (106) the value is cm2 s-’ . These self-diffusion coefficients are in reasonable agreement with the data from MD simulations, 1.6 X cm2 S K ’ (107), and with sorption uptake measurements, which indicate that D > 7 X loK6cm2 s-’ (108). However, a difference of three orders of magnitude is observed with the value extrapolated from the “zero-length column” technique (109); five to six orders of magnitude deviation occurs using sorption uptake and gas chromatographic data (110,111). Additional sorption uptake studies on oriented ZSM-5 crystals have been performed (112,113). For anisotropic diffusion systems, this newly developed technique measuring sorption uptake through certain crystal faces, termed “selective sorption uptake kinetics” (113), can be used to determine the tensor components of the diffusion coefficient. For this reason, large ZSM-5 crystals have been aligned in two ways: upright standing crystals, aligned by means of electric fields as shown in Fig. 22, and crystals aligned horizontally in a plane. After embedding the zeolite crystals thus oriented into a gas-tight matrix (e.g., glass, metal, epoxy resin), by careful abrasion selective crystal faces can be opened for selective sorption uptake kinetics. Abrasion of the crystal arrangement shown in Fig. 22 gives the sample shown in Fig. 23. The sorption uptake kinetics on this sample is exclusively controlled by the mobility in the length direction of the ZSM-5 crystal, Dzz.For ZSM-5 crystals hori-

FIG. 22. Large ZSM-5 crystals synthesized by J. Kornatowski (137. f38) and aligned by means of an electric field of strength 2.3 kV cm-I. Crystals are fixed in the upright position by a thin film of an epoxy resin (113).

FIG. 23. Crystal arrangement as shown in Fig. 22 after embedding the crystals into a thermally stable epoxy resin and abrading the top (113). In sorption uptake, this sample was used to measure selectively D,,.

388

J. CARO et

al.

FIG.24. Plane-oriented ZSM-5 crystals, embedded into a copper matrix (Cu deposition by sputtering). The [lo01 and [OlO] faces were selectively opened for adsorption by abrasion. Because of the random orientation of the crystals in the plane, the mean (Dxx 0,,)/2 is determined in sorption kinetics ( I 13).

+

zontally oriented in the plane, as shown in Fig. 24, the tensor component (Dxx+ Dyy)/2becomes accessible. The mean is a result of the random crystal orientation. Using large ZSM-5 crystals (110 X 110 X 310 pm3), the time, in an individual sorption uptake step, to reach 50% of the final amount adsorbed, for the crystal arrangement shown in Fig. 23, is about 15 times larger than for the crystals shown in Fig. 24. At 298 K , the diffusion coefficients amount to D , = 2.4 X lo-’ cm2 s-’ and D,, = 0.8 X lo-’ cm2 s-’ (113);this shows that for n-hexane in ZSM-5 the mass transport in the length direction of the crystal is about three times slower than in the plane perpendicular to this direction. However, these sorption uptake diffusivity values fall between the values from the spectroscopic methods and the other sorption uptake experiments. Possible reasons for these deviations from intracrystalline mass transport could be the internal twinning of the ZSM-5 crystals under study as well as the fact that abrasion could possibly mechanically damage a tiny surface layer (formation of an amorphous surface barrier).

389


VI. Structure-Related Molecular Self-Diffusion in Zeolites by Pulsed-Field Gradient NMR: Influence of Pore Diameter, W A I Ratio, and Concentrations of Internal OH Groups and Cations; Self-Diffusion of Mixtures

A. SINGLE-COMPONENT MEASUREMENTS As Fig. 25 shows, the intracrystalline self-diffusion coefficient of methane in ZSM-5 is between coefficients in zeolites NaCaA and NaX (5,7Z,lZ4,ZZ5,). This order can be interpreted in terms of the minimum apertures of the zeolite channels, which are approximately 0.45,0.55, and 0.75 nm for 5A, ZSM-5,and X-type zeolites. Due to the hydrophobic nature of Z S M - 5 , the mobility of water in ZSM-5 considerably exceeds the mobility in zeolites NaA and NaX. A change in the SiO2/AI2O3ratio of ZSM-5 does not alter the self-diffusion coefficient of methane. On the contrary, for water in ZSM-5 an increase in the self-diffusion coefficients with decreasing A1 concentrations in the framework is indicated.

$

CHblNaX

CHLI ZSM-5 A[

A

CHdNaCaA

I$

+

AA

['A

A

A A

H201ZSM-5

HfllNaX

H201NaA 4

1

-Si02/A1203

10

lo2

I

lo3

FIG.25. Self-diffusion coefficients of methane (open symbols) and water (filled symbols) in zeolites NaCaA, NaX, and ZSM-5 (loading, approximately 1 CH4 or HzO molecule per 24 T atoms, 296 K) (5).

390

J. CARO

et al.

TABLE VI Influence of Hydrothermal Treatment on Sorption and Diffusion Properties of a Silicalite Specimen Si-OH concentration

Sorption capacity

Self-diffusion coefficient ( X lo-' m2 S K I )

(mmol g-') (number g - ' ) 8 . 4 X lozod 4.6 X 10''

c,,,(n-CsHd

Csat(H2O)

D (CHdb

D (H?0)"

1.28 1.38

3.05 1.62

9.8 7.2

1.4 3.1

The silicalite specimen (121) has large concentrations of internal Si-OH groups; as measured by 'H MAS NMR (122,123). all measurements at 300 K . Loading -8CH4 per U . C . ' Loading -12Hr0 per U . C . Parent sample ' After steaming

Due to synthesis conditions, zeolites can contain remarkable concentrations of intracrystalline Si-OH groups (116-121). These groups represent defect sites in the form of nonintact Si-0-Si bonds inside the crystals. Because hydrothermal treatment is a suitable method to recombine zeolitic Si-0-Si bonds, a silicalite sample was held under steam at 1100 K for 5 days (121). In the parent sample, -8% of the framework Si atoms were present as Si-OH. As Table VI shows, during hydrothermal treatment algroups reacted and formed intact most 50% of the former Si-OH Si-0-Si units. As a result, the n-hexane sorption capacity increased by -7%. With water, the opposite effect has been observed. Owing to the increased hydrophobicity of the steamed sample, the former sorption capacity of water decreases by -50%. However, even this reduced amount of sorbed water is approximately three orders of magnitude larger than the amount necessary to produce monolayer coverage on the external surface of the zeolite crystals. With the drastic reduction in the number of silanol groups, the water self-diffusion coefficient increases. In light of these results, the significant scattering of the values of sorption capacity and diffusivity on zeolites as reported in the literature (70,71,92,95) could correlate also with structural defects such as nonintact Si-0-Si bonds. B . PULSED-FIELD GRADIENT NMR MULTICOMPONENT SELF-DIFFUSION During their technical application, molecular sieve catalysts are generally used under the conditions of multicomponent adsorption and diffusion. Selective measurement of the diffusivity of individual components is therefore of both theoretical and practical relevance. The traditional way to perform


39 1

self-diffusion studies in mixtures by PFG NMR is to use deuterated compounds or compounds without any hydrogen, so that the ‘ H NMR signal stems from only one of the mixture compounds (3, 42, 142). Unfortunately, selective diffusion measurement necessitates additional experimental preparations, because for the study of a system containing n components at least n various PFG NMR samples must be prepared, each of them with a different compound in the hydrogen form. A more straightforward possibility of selective self-diffusion measurements is provided by Fourier transform PFG NMR, as was discussed previously. Using this method, the total NMR signal is split up into separate signals due to different NMR chemical shifts. This procedure has been successfully applied to multicomponent liquids, where it was possible to measure simultaneously the diffusivity of up to eight different components (124). In adsorbate-adsorbent systems, however, such experiments are complicated by the reduction of molecular mobility. The linewidths must be small enough to allow resolution of the observed NMR spectra into the spectra of the individual components. 1. NMR SelfDiffusion in Binary Mixtures Using

Deuterated Compounds a. Methanol and Water in HZSM-5. Figure 26 shows the self-diffusion coefficients of water and methanol in a methanollwater mixture absorbed in HZSM-5 for two different total loadings (125). For every mixture composition, the self-diffusion of the proton-containing component was measured selectively; the other mixture component was present in the deuterium form (for ‘ H NMR “invisible”). For a total loading of 35 mg (water plus methanol) per gram of ZSM-5, the mobility of the sorbed water was found to be enhanced with respect to pure liquid water. In contrast, the mobility of the sorbed methanol was slightly reduced. At a total loading of 50 mg per gram of ZSM-5, the self-diffusion coefficients of both adsorbed water and methanol were found to be lower than the values for the free liquids. For nonadsorbed liquid mixtures, a remarkable minimum of the selfdiffusivity was found (Fig. 26). In contrast, no mimima in the self-diffusion coefficients of the sorbed mixtures were obtained. Obviously, the sorption potential and /or the geometric constraints of the intracrystalline channel system prohibit the formation of highly structured methanol/water complexes, which are present in liquid methanol/water mixtures. These bulky complexes have been proposed to cause the mimima found in the diffusion coefficients of the liquid mixtures (126,127). b. C4 Hydrocarbons in the Presence of Water in NaX. Figure 27 shows that the self-diffusion behavior of paraffins and olefins may be influenced significantly by coadsorbed molecules ( 1 15). The self-diffusion coefficient

J. CARO et

392

al.

O

2.0 -

1.5-

n 1.0-

I

I

I

I

I

0

25

50

75

I I 100

mol % methanol

__c

FIG.26. Self-diffusion coefficients of methanol (squares) and water (circles) in their binary mixtures sorbed in HZSM-5 for two total loadings: 35 mg (HzO + CH3OH) g-' (open symbols) and 50 mg g-' (filled symbols) at 300 K (125). Comparison with the self-diffusivity in liquid methanol/water mixtures (126,127): dotted line, D(CH3OH); dashed lines, D(H20).

of n-butane in NaX decreases by up to three orders of magnitude with increasing amounts of coadsorbed water (D20). With a slight increase for small amounts of coadsorbed water, the self-diffusion coefficients of but- 1ene exhibit a different dependence. This effect is due to the specific interaction of but-1-ene with adsorption centers in NaX. These centers are blocked by water, thus leading to an increase in the mobility of the but-lene molecules. The presence of further water molecules leads to a steric hindrance of the but- 1-ene mobility similar to that of n-butane. As shown in Fig. 28, coadsorbed ammonia (ND3) affects the selfdiffusion behavior of n-butane and but- 1-ene in NaX in a manner similar to that seen in the coadsorption of water. In contrast to water and ammonia, the dependence of n-butane and but-1-ene diffusion on the number of coadsorbed C02 molecules is less pronounced.

2. Fourier Transform PFG NMR The first 'H PFG Fourier transform NMR experiments of adsorbed molecules have been carried out with an ethane/ethene mixture adsorbed on

I

-

number of D,O per cavity

in NaX as a function FIG.27. Self-diffusion coefficients of n-butane ( 0 )and but-I-ene (0) of coadsorbed water (0.8 C4 molecules per cavity, 293 K ) (115).

,

I

1

I

___)

I

5

I

I

I

I

I

I

I

I

J

10

number of cosorbed molecules per cavity

FIG. 28. Self-diffision coefficients of n-butane and but-I-ene in NaX (0.8. C, molecules per cavity, 293 K) in the presence of coadsorbed ND, and C02 (115): A, C4HI0+ ND3;0, C4HlU + CO; A , C4H8 ND,; 0, C~HR + COz.

+

J . CARO et

394

al.

ms

/

1 .d

0

/'

l.I-L__LLi_l

1

2

3

4

5

6

&* H/PPm

FIG.29. ' H PFG Fourier transform NMR spectra of an ethane-ethene mixture in NaX (1.5 CzH6plus 1 CzH4per supercage, 293 K ) for increasing values of the width (6) of the field gradient pulses. The pulse separation ( t ) and the field gradient intensity (g) are 4 ms and 2.8 T m-'. The chemical shifts S , , refer to TMS (128).

zeolite NaX (128). This system is especially suitable for such studies, because the spectra of both components consist of only one line and the mobility of both components is sufficiently high to guarantee line narrowing that allows separation of the two spectra. Considering a sorbate concentration of 1.5 molecules ethane and 1 molecule ethene per supercage, it follows from Fig. 29 that the ethene mobility remains the same as in the case of singlecomponent self-diffusion, D = 1.25 X lo-' m2 s-', whereas the selfm2 s-', is found to be reduced by a diffusivity of ethane, D = 4.6 X factor of about 2. Recently, this method has been applied to the in situ observation of the diffusivity of both the reactant and product molecules during the conversion of cyclopropane to propene in NaX catalysts (135).

VII. Location of Diffusion Obstacles Inside the ZSM-5 Framework by Pulsed-Field Gradient NMR For a number of adsorbate-adsorbent systems it has been found that the intracrystalline self-diffusion of a highly mobile component is drastically reduced by the presence of a second, strongly coadsorbed component. By combining these self-diffusion measurements with a computer simulation of


0

-

05

395

1a

number of coadsorbed benzene molecules per 1 / 4 U.C.

FIG. 30. Self-diffusion coefficient of methane adsorbed on ZSM-5 as a function of coadsorbed benzene (3 CH, per u.c., 293 K) (86).

the random walk (of the highly mobile molecule) in a zeolite framework that contains statistically blocked channels and/or channel intersections, information about the location of the strongly adsorbed component (diffusion obstacle) inside the zeolite framework is obtained. A. MOLECULAR SELF-DIFFUSION OF METHANE ADSORBED ON ZSM-5 CONTAINING COADSORBED BENZENE Figure 30 shows the self-diffusion coefficient of methane in ZSM-5 samples that contain different amounts of coadsorbed benzene (86). It can be seen that the benzene molecules significantly reduce the methane mobility. It is evident that the blocking effect of diffusion obstacles such as benzene should depend decisively on their position inside the channel network: (1) for benzene molecules located in pore segments (between the channel intersections), the passage through only these segments is prohibited; (2) a benzene occupation of the channel intersections should lead to a blocking of all four adjacent channel segments.

396

J. CARO et

al.

-

number of coad60rbed benzene molecules per 114 U.C.

1

\

0.01

b=0004

l . I . I I I I . I I

-

0

0.5

1.0

number of dlffusion obstacle8 a8SUmed In computer slmulatlon per 1/4 u.C.

FIG.31. Computer simulation of the random walk in a two-dimensional network with obstacles distributed statistically over the pore segments (dashed lines) and over the channel intersections (solid lines). Different transition probabilities b to pass the barrier are assumed: b = 0.004, 0.020, and 0.120. The experimental results of the methane self-diffusion in ZSM-5containing coadsorbed benzene (0, cf. Fig. 30) are included (86).

Figure 31 shows the results of a computer simulation of the random walk in a two-dimensional channel network with obstacles distributed statistically (1) over the pore segments and (2) over the channel intersections. As expected, the influence of the obstacles distributed over the channel intersections considerably exceeds the influence of the obstacles in the pore segments. The computer calculation shows that the shape of the individual plots of In D versus the amount N of obstacles is typical of the given type of obstacle distribution, i.e., there is no possibility of transferring the In D versus N plots of one case to another by simply changing the transition probability across the obstacles. One has to conclude, therefore, that for loadings 1 indicates an enrichdeposits over the crystals is indicated. Dinrra/Deff ment of phosphorus species near the outer crystal surface. Figure 35 clearly shows that such enrichment occurs with increasing amounts of phorphorus deposits. Unfortunately, no information on the chemical nature of these transport resistances can be obtained from 'H NMR self-diffusion experiments. For this purpose, several MAS NMR methods have been applied successfully (130). Figure 36 shows the 31PMAS NMR spectrum of the HZSM-5 sample containing 2 wt% P. Three signals can be distinguished: (1) the signal at 1 ppm stems from monomeric [P04l3- groups like those in orthophosphoric acid; (2) the signal at -6 ppm is due to the P atoms in pyrophosphoric acid or due to the terminal [PO4I3- groups in polyphosphoric


I

-

0

1

I

2

3

4

5

40 1

I

mass YO P

for methane self-diffusion in HZSM-5 impregnated with H J P O ~ FIG. 35. Ratio Dinlra/Deff (experimental data from Fig. 34) (130).

I

I

20

0

I

-20 c -

1

-40

I

-60

6 1P/ppm

FIG.36. "P MAS NMR spectrum of HZSM-5 with 2% P brought about by H3PO4soaking and subsequent thermal treatment at 823 K (partial sample rehydration before the "P MAS NMR measurement at 293 K) (130).

species; (3) the signal at -30.7 ppm is attributed to aluminum phosphate. The intensity of the latter signal grows with increasing amounts of H3P04 added and indicates a progressive framework dealumination. Consequently, intracrystalline and /or surface deposits of polyphosphates and/or aluminium

402

J. CARO et

al.

phosphates (the presence of which follows from the finding of 27Aland 31P MAS NMR) should cause the transport limitations observed in PFG NMR and TD NMR studies (130).

B. DETECTION OF THE HYDROTHERMAL CRYSTAL OF ZEOLITE A DAMAGE Table VII compares the measured intracrystalline mean lifetimes, T , " , ~of ~, methane and xenon in various zeolites with the values of ~ S f f .calculated ~, according to Eq. (10) from the self-diffusion coefficients and the mean crystal radii. It turns out that in the case of methane, for all three zeolite samples both quantities are in reasonable agreement; for xenon in NaCaA and HZSM-5 the mean residence time is drastically enhanced in comparison with the values expected from the intracrystalline mobility according to Eq. (10). This experimental finding may be correlated with the larger diameter of the xenon atoms (0.49 nm, in comparison with 0.41 nm for methane), making them more sensitive to a reduction in surface permeability of adsorbents with limiting free diameters of this order of magnitude (0.4-0.5 nm for NaCaA, -0.55 nm for HZSM-5). The surface regions of the crystal exhibit a transport resistance to some types of diffusants (e.g., xenon), while other diffusants (methane) may pass without any perceptible restriction; one can therefore conclude that the surface barrier is not brought about by a total blocking of a certain fraction of apertures in the interface, but by a pore narrowing in a surface layer (42,239).

TABLE VII lntrucrystulline Mean Lifetimes of' Methane and Xenon in Zeolites NaCuA, NuX and HZSM- 5" Zeolites

Mean

Methane

crystal diameter (pm)

T,,l,(ms)

T%:,(rns)

Xenon Tdms)

T Lnrra(ms) ~'~'

NaCaA

CaL '

45% 63%

13 20

80%

5 50 20 25

NaX

HZSM-5 "

3*1 4 2 1 0.4 k 0.2 2+ 1 0.4 2 0.2 3 + 1

Obtained at 293 K . From Ref. 42

6 i 2 4 r 2 1 + 0.5 1 + 0.5 2 + 1 4?1

3+-1 4 2 1 0.4 0.2 8 1 2 1.5 C 0.5 11 2 3

*

80 + 20 45 + 10 25 + 8 15 lir 10 5 + 3 > 40


c.

403

O N THE LOCATION OF COKE DEPOSITS IN/ON ZSM-5

CRYSTALS 1.

Coke Deposition by n-Hexane Cracking

The reduction of molecular self-diffusion due to coking can be used to determine the distribution of coke deposits over the crystals. As an example, Fig. 37 shows the characteristic decrease of Dintraand DeRfor HZSM-5 (Si/ A1 = 70, crystal size -14 p m , coked by n-hexane cracking) as a function of coking times. For ZSM-5 samples with a time onstream

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