Advances in Catalysis, Volume 42

Advisory Board

G. ERTL

ADVANCES IN CATALYSIS

BerlidDahlem, Germany

K. TAMARU Tokyo, Japan

V. B. KAZANSKY Moscow, Russia

VOLUME 42 J. M. THOMAS London/Cambridge, England

W. M. H. SACHTLER Evanston, Illinois

P. B. WEISZ State College, Pennsylvania

Advisory Board

G. ERTL BerlidDahlem, Germany

K. TAMARU Tokyo, Japan

V. B. KAZANSKY Moscow, Russia

J. M. THOMAS London/Cambridge, England

W. M. H. SACHTLER Evanston, Illinois

P. B. WEISZ State College, Pennsylvania

ADVANCES IN CATALYSIS VOLUME 42

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

BRUCEGATES University of California Davis, California

WERNER0. HAAG Consultant Lawrenceville, New Jersey

HELMUT KNOZINGER University of Munich Munich, Germany

ACADEMIC PRESS San Diego London Boston New York Sydney Tokyo Toronto

This book is printed on acid-free paper. Copyright 0 1998 by ACADEMIC PRESS 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. The appearance of the code at the bottom of the first page of a chapter in this book indicates the Publisher’s consent that copies of the chapter may be made for personal or internal use of specific clients. This consent is given on the condition, however, that the copier pay the stated per copy fee through the Copyright Clearance Center, Inc. (222 Rosewood Drive, Danvers, Massachusetts 01923), for copying beyond that permitted by Sections 107 or 108 of the US. Copyright Law. This consent does not extend to other kinds of copying, such as copying for general distribution, for advertising or promotional purposes, for creating new collective works, or for resale. Copy fees for pre-1998 chapters are as shown on the title pages. If no fee code appears on the title page, the copy fee is the same as for current chapters. 0360-0564/98 $25.00

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Contents CONTRIBUTORS .................................................................................. PREFACE.........................................................................................

ix xi

The Molecular Basis of Zeolite Catalysis: A Review of Theoretical Simulations SIMONP. BATESAND RUTGERA.

I. 11. 111.

IV.

VAN

SANTEN

Introduction ................................... ........ Diffusion in Zeolites ................................................................ Adsorption in Zeolites .......................................... Bond Activation by Ze ................. References. .......................................................... .

1 3

50 84 107

NMR Studies of Solid Acidity JAMESF. HAWAND TENG Xu I. 11. 111. IV. V.

Introduction ......................................................... The Chemical Shift .................................................................. Computational Chemistry: A Tool for Spectral Interpretatio Sample Preparation Techniques for in Situ NMR. ............. NMR Studies of Solid Acidity Using Pr References. ..................................

120

Vibrational Spectra of Hydrocarbons Adsorbed on Metals Part It: Adsorbed Acyclic Alkynes and Alkanes, Cyclic Hydrocarbons Including Aromatics, and Surface Hydrocarbon Groups Derived from the Decomposition of Alkyl Halides, etc. NORMAN SHEPPARD AND CARLOS DE LA CRUZ I. 11.

Introduction .......................................................................... The Acyclic Alkynes (Acetylenes) ................................................ V

181 183

vi

CONTENTS

111.

IV. V. VI. VII.

VIII. IX. X.

XI. XII.

The Acyclic Alkanes ................._ ................____ 207 Hydrocarbon Surface Species Derived from the Dissociative Adsorption of Halogen- or Nitrogen-Substituted Alkanes or Alkenes.. . .. .. .. . . . . . . . . .. . . . . . . 214 Cycloalkanes ... .. .. .. . . .. . .. .. . . .. .. .. .. .. .. .. .. .. . . . .. . . 229 Cycloalkenes ............................................... 239 General Commen pectra of the Cycloalkanes and Cycloalkenes .. . . .. . . . . . . . . . . . . . . . .. .. .. 244 Aromatic Hydrocarbons . .. .. . . .. .. .. .. .. .. .. .. .. .. . . . .. . . . . .. .. .. . . . . , . , . , . . . . . , . . . 245 Acyclic Alkenes: An Update since Part I .................................... 261 The Reactivity of Surface Species: An Example. Kinetic Aspects of the Interconversion and Hydrogenation of Ethene and Other CzHnSpecies on Platinum Surfaces . .. .. .. .. .. .. .. . . . . . . . . . . . .. . . . . . .. . . . . . . . . . . . . . . . . . . . . . . . . . . . . 212 Looking Ahead: Some Suggested Priorities for Future Research .. .. .. .. .. ... 295 Conclusions. .... .. .. .. ........................... .. ..... .. . . . . . . . ........ .. .......... . 300 References. . . . .. .. .. .. .. .. .. .. .. .. . . . . . . . .. .. . . .. . .. .. .. .. . . . . . . . . . . . . . . . . . . . . . . .. , , , . 301

Application of Combined X-Ray Diffraction and Absorption Techniques for in Situ Catalyst Characterization BJERNES. CLAUSEN, HENRIKT O P S ~ E AN , D RONALD FRAHM 1. 11.

111.

IV. V. VI. VII.

Introducti .... .........._. X-Ray Di SP lyst Characterization . .. .. . ....................... .. .. .... ... . . . . . .. . .... .. .... .. .. ....... In Situ Approaches.. . .................... .............. .. .. .. . . ... .. .. .. .. .. .. .. .... Recent Advances in X-Ray Diffraction and X-Ray Absorption Techniques . . . . . . . . . .. .. .. .. .. .. .. .. . . . .. . . .. .. .. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .. .. . . . . . . . . . . . . . . . . . . . . . . . Combined EXAFS/XRD Methods.. Examples of in Situ Combined EXAFS/XRD Investigations ...... Outlook ............................................................. ...... References ............ .....................................

315 316 321 324 330 332 340 342

Present State of the Art and Future Challenges in the Hydrodesulfurization of Polyaromatic Sulfur Compounds D. DUAYNE WHITEHURST, TAKAAKI ISODA, A N D I. 11. 111.

IV. V. VI. VII. VIII.

ISAO MOCHIDA

Introduction . .. .. .. .. . . .. . . .. .. .. .. . . . . . . . . .. .. .. .. .. .. .. . . .. . . . . . . . . . . . . . . . . . . . . . . . .. Description of Systematic Approach for Describing the ................................. HDS Phenomenon . . . Composition of Sulfur Species in Middle-Distillate Oils Conventional HDS Processes and Catalysts .. ... .. .. .. Computational Aids to Mechanistic Understanding . . . . . . .. .. . . . . .. .. . . . . .. .. .. Limitations in Conventional HDS Processes.. .. .. .. .. . . . . . .. . . . . . . . . . . . . . . . . . . . . Novel Approaches for Deep Desulfurization . .. .. .. .. .. .. . . . . . . . . . . . . . . . . . . . . . . . Concluding Remarks ................... ........................... ...... References. .. . .. .. .. .. .. .. .. .. .. .. .. .. .. .. .. .. .. .. .. . . . .. .. .. ..

345 349 353 366 425 435 455 466 461

CONTENTS

vii

Multiphase Homogeneous Catalysis BIRGIT DRIESSEN-HOLSCHER

............. Introduction .. ..... .. .. .. .. .. .. The Principle of Multiphase Catalysis .. .. .. .. . . .. .. .. .. . . ... .. .. .. ... .. .. ..... .. . Aspects of Mass Transfer in Multiphase Catalysis .... . .. .. .. .. . .... .. . .. .. .. .. . Reactions in Water as the Catalyst Phase.. Multiphase Reactions with Solvents Other Catalyst Phase . .. .. .. .. Industrial Applications Summary and Outlook.............................................................. References. . . . . .. .. .. .. .. .. .. .. . . . .. .. .. . . .. .. .. .. .. .. .. .. . . ... .. .. ..... .. ... .. .. .. .. .

473 474 476 476

INDEX.. . .. .. .. . .. .. .. . .. . . .. .. .. .. .. .. .. .. .. .. ... .. .. .. .. .. .. .. .. .. .. ... . .. . .. . . .. . .. .. .. .. ... .. . .

507

I. 11. 111.

IV. V.

VI. VII.

495 497 501 501

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

SIMON P. BATES,Department of Chemistry, University of Edinburgh, Edinburgh EH9 3JJ, Scotland (1) BJERNES . CLAUSEN,Haldor Topsae Research Laboratories, DK-2800 Lyngby, Denmark (315) CARLOS DE LA CRUZ,Departamento de Quimica, Facultad de Ciencias, L a Universidad del Zulia, Maracaibo, Venezuela (181) BIRGITDRIESSEN-HOLSCHER, Institut f u r Technische Chemie und Petrolchemie der RWTH, 52056 Aachen, Germany (473) RONALD FRAHM, Institut fur Angewandte Physik, Heinrich-Heine-Universitat, 0-40225 Dusseldorf, Germany (315) JAMES F. HAW, The Laboratory for Magnetic Resonance and Molecular Science, Department of Chemistry, Texas A & M University, College Station, Texas 77843 (115) TAKAAKI ISODA, Kyushu University, Fukuoka 812-81, Japan (345) ISAO MOCHIDA, Institute of Advanced Material Study, Kyushu University, Fukuoka 816, Japan (345) NORMAN SHEPPARD, School of Chemical Sciences, University of East Anglia, Norwich N R 4 7TJ, England (181) HENRIK TOPSOE, Haldor Topsae Research Laboratories, DK-2800 Lyngby, Denmark (315) RUTGER A. VAN SANTEN, Department of Inorganic Chemistry and Catalysis, Eindhoven University of Technology, 5600 M B Eindhoven, The Netherlands (1) D. DUAYNE WHITEHURST, Institute of Advanced Material Study, Kyushu University, Fukuoka 816, Japan (345) TENGXu, The Laboratory for Magnetic Resonance and Molecular Science, Department of Chemistry, Texas A & M University, College Station, Texas 77843 ( 115)

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Preface With the present volume, Dan Eley retires as co-editor of Advances in Catalysis, having served the catalysis community in this position for 40 years. Professor Eley’s academic research contributions in catalysis have ranged widely, from fundamental studies of para-hydrogen conversion, which led to establishing the Eley-Rideal mechanism in catalytic kinetics, to work ranging from Friedel-Crafts catalysts, semiconductors, and enzymes to colloidal solutions and the mechanism of adhesive action. The Advances have benefited greatly from his broad experience and knowledge in his many years as co-editor. For this and his diligent, selfless work, he deserves our thanks. We are pleased to have him available for advice in the future, as he will join the ranks of the editorial board. We are also very pleased that Professor Helmut Knozinger has agreed to join us as co-editor, having just completed the enormous task of coediting the five-volume Handbook of Heterogeneous Catalysis published by Verlag Chemie. We have already benefited from his collaboration in the preparation of the present volume; we extend a wholehearted welcome to him. The six chapters of this volume again reflect the diversity of the science that is relevant to catalysis. In the first chapter, Bates and van Santen summarize the theoretical foundations of catalysis in acidic zeolites. Being the most important crystalline materials used as catalysts, zeolites have been the obvious starting point for applications of theory to catalysis by solids and surfaces. Impressive progress has been made in the application of theory to account for transport, sorption, and reaction in zeolites, and the comparisons with experimental results indicate some marked successes as well as opportunities for improving both the theoretical and experimental foundations. Complementing this contribution, Haw and Xu present a detailed assessment of the nature of acidic surface sites (most in zeolites) and their interactions with probe molecules, as assessed in NMR experiments. Their comprehensive approach sheds light on a number of timely issues in acid-base catalysis and demonstrates how successfully NMR spectroscopy has been used recently to understand surface and catalytic phenomena. Sheppard and De La Cruz complete their two-part treatise on the vibrational spectra of hydrocarbons adsorbed on metals; the first part of this xi

xii

PREFACE

comprehensive and critical reassessment of the literature appeared in Volume 41 of Advances. Two of the most incisive physical methods used in catalysis are X-ray diffraction and X-ray absorption spectroscopy, which provide characterizations of relatively large and relatively small structures, respectively. Clausen, Topsae, and Frahm demonstrate how these techniques are used to characterize solid catalysts in the working state and how they complement each other. Whitehurst, Isoda, and Mochida write about catalytic hydrodesulfurization of fossil fuels, one of the important applications of catalysis for environmental protection. They focus on the relatively unreactive substituted dibenzothiophenes, the most difficult to convert organosulfur compounds, which now must be removed if fuels are to meet the stringent emerging standards for sulfur content. On the basis of an in-depth examination of the reaction networks, kinetics, and mechanisms of hydrodesulfurization of these compounds, the authors draw conclusions that are important for catalyst and process design. Homogeneous catalysis by transition metal complexes almost always involves processes in which product-catalyst separation and catalyst recycling are important issues. For years, researchers have worked to find effective ways to isolate metal-complex catalysts in phases separate from those containing the catalyst, usually by anchoring the metal complex to a solid surface. As summarized by Driessen-Holscher, it is now evident that the method that has met with most practical success in this direction involves the use of multiple liquid phases. For example, rhodium complexes with water-soluble sulfonated ligands are used to catalyze alkene hydroformylation, and the aqueous-phase catalyst and the organic products are easily separated as insoluble liquid phases. W. 0. HAAG B. C. GATES

ADVANCES IN CATALYSIS, VOLUME 42

The Molecular Basis of Zeolite Catalysis: A Review of Theoretical Simulations SIMON P. BATES",? AND RUTGER A. VAN SANTEN Department of Inorganic Chemistry and Catalysis Eindhoven University of Technology 5600 M B Eindhoven, The Netherlands

1.

Introduction

There has been a phenomenal growth of interest in theoretical simulations over the past decade. The concomitant advances made in computing power and software development have changed the way that computational chemistry research is undertaken. No longer is it the exclusive realm of specialized theoreticians and supercomputers; rather, computational chemistry is now accessible via user-friendly programs on moderately priced workstations. State-of-the-art calculations on the fastest, massively parallel machines are continually enlarging the scope of what is possible with these methods. These reasons, coupled with the continuing importance of solid acid catalysiswithin the world's petrochemical and petroleum industries, make it timely to review recent work on the theoretical study of zeolite catalysis. If the reaction of a species within the pores of a zeolite is decomposed into its constituent steps, the following elementary processes may be envisaged: i. Transport of reactant to the active sites ii. Sorption of reactant at an active site

* Corresponding author. t Present address: Department of Chemistry, University of Edinburgh, West Mains Road, Edinburgh EH9 355, Scotland. Abbreviations: MD, molecular dynamics; TST, transition state theory; EM, energy minimization; MSD, mean square displacement; PFG-NMR, pulsed field gradient nuclear magnetic resonance; VAF, velocity autocorrelation function; RDF, radial distribution function; MEP, minimum energy p a t h MC, Monte Carlo; GC-MC, grand canonical Monte Carlo; CB-MC, configurational-bias Monte Carlo; MM, molecular mechanics; QM, quantum mechanics; HF, Hartree-Fock; DFT, density functional theory; BSSE, basis set superposition error; DME, dimethyl ether; MTG, methanol to gasoline. 1 Copyright Q 1998 by Academic Press. All rights of reproduction in any form reserved. 0360-056498$25.00

2

SIMON P. BATES AND RUTGER A. VAN SANTEN

iii. Surface reaction and conversion to products iv. Desorption and transport away from active sites Step iv is the complementary process to steps i and ii, involving products rather than reactants. Steps i-iii are the fundamental steps that constitute a reaction, and these broad categories correspond to each of the three main sections in this review. Furthermore, we find that each category is largely characterized by one particular type of simulation method. The first step, transport to an active site, is a dynamic process whereby the reactants enter the pores of the zeolite and move under the influence of the stabilizing topological environment of the pore system. Molecular dynamics calculations, which lead to a solution of the classical equations of motion over time, have been widely applied to this problem. The second stage, sorption at an active site, is a more static process, extensively studied via Monte Carlo simulation techniques. The final category, reaction and transformation, invariably involves breaking and formation of bonds within adsorbate and/or zeolite. Quantum chemical calculations, capable of providing detailed electronic information, have been used almost exclusively to investigate this type of step. By pooling the information obtained by using different calculation methods for each of the individual steps, we obtain a coherent picture of zeolite-catalyzed processes. This kind of information may then be used to help deduce the overall kinetics of a catalytic reaction. In this review, we focus on the information at an atomidmolecular level that is obtainable via the different techniques. The precise methods and techniques used are not extensively discussed; instead we summarize the relevant details and direct the reader toward key references. Nor do we review the potentials that are used in the classical simulations of sorption and diffusion. Derivation and evaluation of these parameters require extensive comparison with detailed spectroscopic data and are beyond the scope of this work. Similarly, the volume of experimental results that may be used in comparison to the calculations is vast. We use representative data taken largely from reviews or books. Throughout the review, there are some trends that become apparent. Irrespective of the nature of the process being modeled-diffusion, sorption, or reaction-the earliest calculations were kept as simple as possible, principally on the grounds of computational expense. Over time, calculations became more and more sophisticated and could therefore be applied not only to study a wider range of systems but also to highlight subtle effects. Examples of this progression in complexity can be seen time and time again in the following pages. Calculations always have represented a balance of available resources and level of computational theory on the

MOLECULAR BASIS OF ZEOLITE CATALYSIS

3

one hand and system complexity on the other. As computing capability advances, one might naively expect this balance to become less of a problem. However, the advances in hardware and software are more than matched by the increased complexity of systems that are being investigated and the level of information that is sought from them. The trade-off of computer time versus system size is just as relevant today as it was 10 years ago. Having stated that, we emphasize that significant advances continue to be made. The following summary shows a progression not only in complexity of calculations but also in the quality and reliability of the predictions made. Throughout the individual sections, we seek to highlight key points or issues that arise from the calculations. Examples of these include the following: Reproducibility: Are results sensitive to slight changes in calculation details, and if so, why? Assumptions: How valid are common assumptions such as maintaining a rigid zeolite framework? Is the imposition of rigidity less serious for some systems than others? Comparability: What can we deduce from a comparison with experimental results characterizing the same system? Are the calculations a realistic model of a real system? Applicability: How applicable are the methods we describe to a wider range of systems? Can they be used for larger molecules or more polar species? The three main sections of this review are organized as follows. The first two, concerned with diffusion and sorption, are structured similarly. The key aspects of the theoretical methods used are described, and then recent literature is reviewed and ordered according to the sorbate molecule. The third section, concerned with activation and reaction, is slightly different. Once again key details of methods are summarized, but then, rather than attempting a complete review of this literature, we focus on two key topics which exemplify the use of this technique. We also highlight the main conceptual advances in this area. The reason for this organization is twofold: these types of quantum chemical calculations are completely reviewed at regular intervals, and the volume of work published merits a complete article for a complete evaluation. II. Diffusion in Zeolites

A. SCOPEOF THISSECTION

This part of the review comprises a brief description of the main principles of the various theoretical methodologies that have been used to investigate

4

SIMON P. BATES AND RUTGER A. VAN SANTEN

diffusion, followed by reports of diffusion simulations, ordered according to diffusing species. Wherever possible, the results of simulations are compared with relevant experimental results to assess the performance of the simulations. We restrict our interest to the simulation of diffusion under equilibrium conditions, i.e., in the absence of any concentration gradients. The random motion of the molecule is then described by a self-diffusivity that can be determined by “micr~scopic”experimental methods. In addition, we make a distinction between the sorption and diffusion of molecules. This section is concerned with the motion of the sorbates alone. The location and conformation of sorbates, together with the energetics of sorption, are considered separately in the next section. The simulation of migration and diffusion of guest molecules within the micropores of zeolites has been most extensively studied by using molecular dynamics (MD) techniques. A wide variety of different systems have been investigated over the past 5 years or so, with diffusing guest species ranging from single atoms such as Xe to hydrocarbons up to six carbon atoms in length and aromatic molecules such as benzene. As the size of the diffusing species increases, the computational cost of the simulations rises dramatically as a result of the long time scales needed to characterize the diffusivity of larger molecules. This cost enforces a lower limit of approximately lo-’’ m2/s on diffusivities that may be practically investigated by MD calculations. Thus, it becomes appropriate to consider alternative methods of simulation. The transition state theory (TST) method of analysis has been used to investigate the dynamics of larger molecules such as benzene in zeolites. In addition to these methods of simulation, a technique based on energy minimization (EM) has also been employed to simulate a “forced” diffusion through a crystalline host. Whereas most effort has been directed toward the zeolite ZSM-5 and its all-silica polymorph silicalite, other frameworks such as zeolite A, faujasite, and ferrierite have also been considered. This theoretical interest mirrors the commercial relevance of these zeolites.

B. THEORETICAL METHODOLOGIES 1. Molecular Dynamics Simulations

In an MD calculation, the classical equations of motion are solved for a system of interacting atoms. Newton’s familiar equation that equates the force on a given atom with the product of its mass and acceleration is integrated with respect to time by a finite-difference method. The resulting coordinates and velocities characterize the trajectory of the system as it evolves over time.

MOLECULAR BASIS OF ZEOLITE CATALYSIS

5

MD calculations may be used not only to gain important insight into the microscopic behavior of the system but also to provide quantitative information at the macroscopic level. Different statistical ensembles may be generated by fixing different combinations of state variables, and, from these, a variety of structural, energetic, and dynamic properties may be calculated. For simulations of diffusion in zeolites by MD methods, it is usual to obtain estimates of the diffusion coefficients, D , from the mean square displacement (MSD) of the sorbate, (?(t)), using the Einstein relationship (I): D = lim-(r2> t+m 6t These calculated intracrystalline diffusion coefficients are particularly appropriate for comparison with those determined from pulsed field gradient (PFG) NMR experiments. Time-independent equilibrium properties such as adsorbate conformations are also readily accessible. The classical nature of the simulations allows a particle’s trajectory to be followed, and from this it is possible to determine all kinds of information, such as how often a particle diffuses through a certain region. The velocity autocorrelation function (VAF) may be used to investigate the possibility of coupling between translational and rotational motions of the sorbed molecules. The VAF is obtained by taking the dot product of the initial velocity with that at time t. It thus contains information about periodic fluctuations in the sorbate’s velocity. The Fourier transform of the VAF yields a frequency spectrum for sorbate motion. By decomposing the total velocity of a sorbate molecule into translational and rotational terms, the coupling of rotational and translational motion can be investigated. This procedure illustrates one of the main strengths of theoretical simulations, namely to predict what is difficult or impossible to determine experimentally. The requirements for a successful MD calculation are quite straightforward: allow the integration of the equations of motion to proceed for a long enough time for the system to explore all the regions of configurational space that are significant; ensure that the integration time step is not so long as to introduce artifactual behavior; select a good initial starting configuration in order to reduce the time spent in equilibration; and choose reliable and tested parameters that determine the strength and extent of interactions between atoms of the system (the force field). Given that computational resources are always limited, the selection of a suitable time step becomes a balance between not introducing artifactual behavior and exploring enough of the host pore system so as to reliably predict longrange diffusivities (2).

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SIMON P. BATES AND RUTGER A. VAN SANTEN

The force field should comprise the following interactions: the intramolecular degrees of freedom of the sorbate, the dynamics of the zeolite lattice, and the intermolecular interactions between zeolite and guest. As with all theoretical calculations, a balance has to be struck between size or complexity of the system under study and the completeness of the theoretical treatment that may be applied to it. This universal constraint has led to the majority of simulations being performed with force fields that neglect one or more of the aforementioned interactions. Most commonly, the dynamics of the zeolite are neglected by keeping it rigid. If the guest molecule is a hydrocarbon, it is often treated as a chain of “united” CH213pseudoatoms [the Ryckaert-Bellmans model ( 3 ) ] .Apart from the obvious reduction in the number of atoms and therefore the number of potential terms to evaluate at each iteration, the united-atom approach also removes highfrequency motion of hydrogen atoms, allowing a relatively long time step to be used. Just how valid all these approximations and simplifications are will be discussed with the theoretical results themselves. Not every parameter used in the various force fields will be described in great detail, as this is not a review concerned with the simulation of aspects of zeolite structure. Instead, we aim to present the essential features of the various parameters that are used and to group force fields into certain families that essentially originate from one of a handful of key references. 2.

Transition State Theory

Transition state theory (TST) ( 4 ) is a well-known method used to calculate the kinetics of infrequent events. The rate constant of the process of interest may be factored into two terms, a TST rate constant based on a knowledge of an equilibrium phase space distribution of the system, and a dynamical correction factor (close to unity) used to correct for errors in the TST rate constant. The correction factor can be evaluated from dynamical information obtained over a short time scale. In principle, this information may all be obtained from a MD simulation, but the real power of the TST method becomes evident when one imagines diffusing sorbates that must pass through a phase-space bottleneck. This results in a prohibitively long MD simulation or a simulation in which the sorbate does not sample a large enough proportion of the phase space to reliably predict the diffusivity. Computational savings with the TST method are estimated to be as high as 2 orders of magnitude compared with a full MD simulation, without significant compromise in the quality of predicted diffusivities. The determination of the diffusivity according to the TST formalism is based on the assumption that the diffusive motion of the sorbate through

MOLECULAR BASIS OF ZEOLITE CATALYSIS

7

the zeolite proceeds via a series of uncorrelated, infrequent hops between potential energy minima inside the zeolite pore system. A sorption state is constructed at each minimum, and a first-order rate constant, kij, is then associated with the transition between a pair of neighboring states, i and j , which are separated by a saddle point. All possible rate constants are determined for all state pairs i and j . Then a Monte Carlo calculation generates random walks of sorbates on the lattice of potential minima, with the rate of hopping given by the TST rate constants. The dynamical correction is then calculated by a MD simulation of “ghost” sorbates thermalized over the surface near state i. This MD calculation need only be run for a short time. The corrected rate constants provide a basic description of hopping between states, but it is necessary to determine self-diffusivitiesfrom these to enable comparison with experimental measurements. In the case of potential minima within a zeolite pore, the lattice of sorption sites is often anisotropic. The probability of a molecule residing in a certain site is dependent on the type of site, and the rate constants, k,, may be different for each ij pair. A Monte Carlo algorithm, based on a first-order description of the hopping process, is usually used to determine the diffusivities. One final point worth mentioning is the location of minima and saddle points within the zeolite pore system. This search is greatly assisted by assuming a rigid lattice and a spherical sorbate, thus restricting the number of degrees of freedom to three. A three-dimensional grid spread over the asymmetric unit of the zeolite is usually used to determine the potential and gradient vector at a series of grid points. Minima are located on the basis of a change in sign of the gradient vector. The a priori determination of potentials is not possible if the zeolite lattice is flexible, which explains the vastly increased computational effort for a flexible lattice calculation of this type. 3.

“Forced” Diffusion from Energy Minimization Calculations

The prediction of diffusional behavior by energy minimization techniques involves the constrained minimization of a sorbate moving in a stepwise fashion between fixed points. What is actually determined is the minimized energy of the guest within the (rigid) host at a number of positions throughout the host lattice, e.g., along the length of a channel. The diffusion coefficients that can be estimated from the slope of the graph of mean square displacement as a function of time are at best order-of-magnitude estimates. The real strength of this method lies within its visualization capabilities; the algorithms for this constrained diffusion were developed as part of MSI’s Catalysis and Sorption Project Software (5). The ability

8

SIMON P. BATES AND RUTGER A. VAN SANTEN

to view the guest's trajectory through the host and to see a spectrum of guest energies as a function of position inside the host is a valuable visual aid.

C. SURVEYOF RESULTS 1. Single Atoms The simplest interaction between host and guest occurs if the guest is a single atom. The guest particle is usually described as a simple LennardJones particle (i.e., a 6-12 Lennard-Jones function is used to describe the guest-host and guest-guest interactions). If there are two interacting atoms, 1 and j , then U(LJ,(r,)

=

4 E . , [ ( ~ l l ~ ~-l l(~l]/rlJ'21. )6

(2)

The values of q1and ullare the well depth and size parameters, respectively, for the two interacting atoms i and j. In the case that one of the interacting atoms is a zeolite atom and the other is a sorbate atom, the cross terms &zeo-sorb and gzeo-sorb are determined from the Lorentz-Berthelot combination rules ( I ) . When polarization interactions are accounted for, such as those between adsorbates and zeolite extra framework cations, Eq. (2) is written in the form

u(IJ)(rl,) = [Alp;

-

B,/r;*I,

(3)

where

A , = 4 ~ , a ~ ; B, = ~ E , C T ? . (4) Here A , and B, are characteristic coefficients of dispersion and repulsion, respectively. [The former can be calculated by the Kirkwood-Muller approach (6, 7) and the latter from the condition of minimum potential energy at a distance equal to the sum of the van der Waals radii.] Interatomic potentials are usually subject to a cutoff at a distance of around 10 A ( I ) . This leads to a somewhat higher potential energy and influences the energetics of the system slightly. Dynamics are less affected by this cutoff radius, as attractive contributions beyond the cutoff radius approximately cancel each other. A single atom adsorbate exhibits only translational degrees of freedom and this relative computational simplicity explains why simulations of the diffusion of noble gas atoms have been reported in the literature for more than 10 years. Another reason why these systems have received so much attention is that the diameters of commonly studied atoms (such as argon) are similar to the molecular diameter of methane, making them simple monoatomic approximations of CH4. Furthermore, 129XeNMR is a well-known technique used to characterize microporous solids (8, 9).

MOLECULAR BASIS OF ZEOLITE CATALYSIS

9

Xenon. The diffusion of xenon in microporous hosts has been investigated extensively by MD methods (10-21). In a series of publications, Yashonath et al. have investigated the diffusive behavior of Xe in zeolite NaY (13, 15, 17, 18, 20). In their calculations, the zeolite had a Si/A1 ratio of 3.0 and the charge-balancing Na cations fully occupied the SI and SII cation sites of the structure, leaving the 12-ring windows that permit entry into and exit from the a-cages free for diffusion of the Xe atom. A concentration of 1 Xe atom per supercage was initially considered at a temperature of 364 K. Host-guest and host-host interactions were described by a 6-12 Lennard-Jones function, with parameters taken from the work of Kiselev and Du (22).Interactions between the guest atom and the Si and A1 atoms of the zeolite were neglected; this is a reasonable approximation because these atoms are surrounded by bulkier, more polarizable oxygens. Rather more significant approximations are that induction effects resulting from the polarizability of zeolite oxygen and xenon atoms were neglected, and the zeolite was held rigid in all these simulations. The calculations were performed in the microcanonical ensemble and allowed to run for approximately 3 ns with an integration step size of 40 fs. The motion of the Xe atom was investigated as it migrated from one supercage to another, via a 12-ringwindow. From an analysis of the trajectories of these window-crossing events, Yashonath et al. were able to obtain a great deal of information concerning the mechanism and energetics of diffusion. They found that the potential that a Xe atom is subjected to as it crosses the 12-ring window depends on the distance of the guest from the window center (13). A potential minimum of -12 kJ/mol was found 1.6 p\ from the center of the 12-ring. The potential barrier to crossing from one cage to another was found to be small (<1 kJ/mol), and the maximum was found just in front of the plane of the 12-ring. Most Xe atoms cross this barrier 1.6 A from the center of the 12-ring. The average time that a Xe atom spends within a given cage was calculated to be 9.9 ps at 376 K (14). The diffusion of Xe in zeolite Y may be considered as being either “surface-mediated” or “centralized”; in the surface-mediated mode, the adsorbate glides along the cage wall and is quasi-two-dimensional. Centralized diffusion dictates that the molecule be near the center of the supercage and is subject to little or no effect from the walls. From an analysis of the distance between the Xe atom and the center of the “parent” supercage (in which Xe resides before the crossing event) and the distance between the Xe atom and the center of the “daughter” supercage (where Xe resides after the crossing event), it was shown that the predominant mode of diffusion is surface-mediated. This is illustrated in Fig. 1. Crossover events are split into surface-mediated or centralized by the minimum distance between the center of the Xe atom and the center of

10

SIMON P. BATES AND RUTGER A. VAN SANTEN

I

FIG.I . Illustration of the predominance of surface-mediated diffusion of Xe in NaY zeolite. The abscissa and ordinate axes give the distance between a Xe atom and the center of a parent and daughter supercage, respectively, during an intercage diffusion event. Reprinted with permission from Ref. 13. Copyright 1991 American Chemical Society.

the supercage it is leaving. Xe atoms that are less than 3 A from the center of the supercage at their closest point are deemed to diffuse through via the centralized mode, and those further from the center follow the surfacemediated route. The value of 3 A is simply half the radius of the supercage, though Yashonath and Santikary (18) stated that perhaps a more appropriate division would be based on the underlying potential energy surface. Simulations representing different temperatures (between 188 and 479 K) and also different Xe concentrations (2 and 3 Xe atoms per supercage) have been reported (l.5,17,18). There is evidence for dimerization at higher loadings, which is fully discussed in the sorption section of this review. The dimerization of Xe is compatible with results for methane (23), reinforcing the similarities between these two sorbates. At higher temperatures than that of the initial studies, the Xe atoms are distributed more evenly as a function of distance from the center of the supercage, but a preference for a site near the walls was still found. This result has important consequences for the overall energy profile of a diffusing Xe atom. The barrier height of cage-to-cage migration (as calculated from the difference in energy in a region of 4 ps either side of the crossover event) was found to decrease with increasing temperature. It was suggested that this is non-Arrhenius behavior (17), in contrast to the temperature dependence that is observed experimentally. However, the observed activation energy comprises two terms: first, the energy needed to free the Xe from its preferred sorption location, and second, the energy barrier associated

MOLECULAR BASIS OF ZEOLITE CATALYSIS

11

with passing through the window. If the activation energy is calculated from the variation in the number of cage-to-cage crossovers with increasing temperature, an Arrhenius dependence is observed and an activation energy value of + 3 kJ/mol is obtained. From the slope of the graph of MSD as a function of time, the diffusion coefficient is calculated to be 0.29 X lo8 m2/s at 188 K and 1.43 X 10' m2/s at 479 K. A residence time of the Xe atom within a cage was estimated to be 15 ps, which is significantly shorter than the experimentally determined residence times quoted for methane molecules within the supercage of zeolite Y. The discrepancy is explained by the lack of accounting for defects and grain boundary effects in the theoretical simulation. At the moment of crossover from one supercage to another, it was predicted at all temperatures that the Xe atom prefers to pass through the 12-ring at a position approximately 1.6 A from the center of the ring. However, the trajectory of the atoms before and after the crossover event was found to be different at different temperatures; at higher temperatures particles can start (and finish) at positions closer to the center of the cage. The energy profiles of the two different diffusion pathways (surface-mediated and centralized) may be interpreted by considering the locations of minima and maxima in the supercage. The global minimum for a Xe atom in a supercage has been shown to be at a site close to the cage wall; the global maximum is at the cage center. The plane of the 12-ring window is a saddle point on the potential energy surface. Thus, it can now be seen that those Xe atoms diffusing via the surface-mediated mechanism will experience a potential barrier at the 12-ringwindow, whereas those diffusing via the centralized regime are subjected to a potential well. The apparent decrease in barrier height as a function of temperature is evident when one considers that the contribution to this centralized diffusion increases with increasing temperature because the Xe atoms are more delocalized, from 18% at 188 K to 45% at 479 K (18). A schematic of energy profile associated with both diffusion pathways is shown in Fig. 2. Xenon has been considered as the diffusing species in simulations of microporous frameworks other than faujasite (20-12,21). Pickett et al. (10) considered the silicalite framework, the all-silica polymorph of ZSM-5. Once again, the framework was assumed to be rigid and a 6-12 LennardJones potential was used to describe the interactions between Xe and zeolite oxygen atoms and interactions between Xe atoms. The potential parameters were slightly different from those used by Yashonath for migration of Xe in NaY zeolite (23).In total, 32 Xe atoms were distributed randomly over 8 unit cells of silicalite at the beginning of the simulations and calculations were made for a run time of 300 ps at temperatures from 77 to 450 K. At 298 K. the diffusion coefficient was calculated to be 1.86 X m2/s. This

12

SIMON P. BATES AND RUTGER A. VAN SANTEN

r

window

cage

windan,

I

I

I

wirdow

cage

1

wincbw

FIG.2. Energy profile as a function of position for the surface-mediated and centralized modes of Xe diffusion in zeolite Y. The upper curve refers to the centralized mode of diffusion, and the lower curve to the surface-mediated mode of diffusion. Reprinted with permission from Ref. 18.Copyright 1993 American Chemical Society.

value is to be compared to 4 X m2/s estimated on the basis of PFGNMR experiments (24, 25). At 77 K the calcufated diffusion coefficient was found to be very small, suggesting that the Xe atoms hardly move at this temperature. From an Arrhenius plot, the activation barrier to diffusion was found to be 5.5 kJ/mol. At higher Xe loadings and a fixed temperature of 298 K, the diffusion coefficients were found to decrease, and the authors speculated that the maximum capacity of Xe in silicalite is around 16 atoms per unit cell. Increased guest-guest interactions and collisions account for a decrease in the calculated diffusion coefficient. The silicalite framework offers the possibility of investigating the anisotropy of diffusion through the straight and sinusoidal channels, which are aligned parallel to the y - and x-axes, respectively. The diffusion coefficients parallel to each crystallographic axis suggest that the motion of Xe atoms is fastest along the straight channels and slowest parallel to the z-axis (as there is no channel along this direction). Diffusion parallel to the z-axis, i.e., with motion from one intersection to another, must involve a tortuous route through a number of different channel segments; hence motion in this direction is slower than in the other two crystallographic directions. The value of the diffusion coefficient of Xe in the straight channel is approximately 3 times that in the sinusoidal channel. It appears that the preexponential term in the Arrhenius equation determines this, as there is no great difference in the activation energies determined for diffusion in each of the three crystallographic directions. The observation of diffusion anisotropy in NMR measurements is complicated when small crystallites are used. It has been estimated (24) that the diffusivities of molecules in the channels of HZSM-5 do not differ by more than 1 order of magnitude, but the uncertainty in the data prevents a definite confirmation of the MD

MOLECULAR BASIS OF ZEOLITE CATALYSIS

13

results. More definitive results are available for methane, as discussed in the next section. June et al. (11) also performed MD calculations to characterize the dynamics of Xe in silicalite. A fixed lattice was assumed with potential parameters close to those used in previous MD studies. The potential between zeolite and guest was determined prior to the calculation over a threedimensional grid spanning the asymmetric unit. From these grid points, the potential at any point in the lattice could be found by interpolation. Temperatures of 200,300, and 400 K were imposed during the simulations, which ran for 1 ns. Diffusion coefficients at different Xe loadings were found to be comparable to those found in the work of Pickett et al. (10) and also to NMR m2/s at 300 K and 4 atoms measurements (24) (approximately 4 X per cell). Again, the predicted diffusion coefficients decreased with increasing Xe loading. This effect was most noticeable for the simulations at 300 and 400 K. The activation energy to diffusion was calculated to be 6.6 kJ/mol, 20% higher than the value reported by Pickett et al. (10). The difference is most probably caused by differences in the length of simulation [ l ns by June et al. (II), 300 ps by Pickett et al. ( l o ) ] .The activation energy for diffusion was found to decrease with Xe concentration, reaching a limiting value of 2.8 kJ/mol at 16 atoms per unit cell. Experimental values of the activation barrier for diffusion tend to be higher than this, as discussed shortly. June et al. (12) used TST as an alternative method to investigate Xe diffusion in silicalite. Interactions between the zeolite oxygen atoms and the Xe atoms were modeled with a 6-12 Lennard-Jones function, with potential parameters similar to those used in previous MD simulations (11). Simulations were performed with both a rigid and a flexible zeolite lattice, and those that included flexibility of the zeolite framework employed a harmonic term to describe the motion of the zeolite atoms, with a force constant and bond length data taken from previous simulations (26)From a calculation to locate the energy minima of Xe in silicalite, three sets of four symmetry-equivalent minima were found, one in each of the two channel systems and one at the intersections. Pathways connect these minima via saddle points and in addition to pathways from a straight or sinusoidal channel minimum to an intersection minimum, there are pathways that are direct transitions between channel segments, circumventing the intersection minima. Simulations of a flexible framework required far more computational effort and the inclusion of flexibility was found to reduce the calculated

14

SIMON P. BATES AND RUTGER A. VAN SANTEN

diffusivity by approximately one-third in the limit of zero coverage; the values are 0.46 X lo-* and 0.33 X lo-* m2/sfor the rigid and flexible lattices, respectively, at a temperature of 100 K. At 200 K, the value was calculated (for a rigid lattice only) to be 1.1 X lo-* m2/s. Previous MD simulations by the same authors (11) lead to a value of 1.5 X lo-’ m2/s for the same conditions. Experimental values from NMR measurements may be extrapolated to zero coverage to give a value of approximately 1 x m2/s at 298 K. Both the diagonal elements of the diffusion tensor and the activation energy for diffusion determined from an Arrhenius plot give values similar to those found in the MD simulations (11).Values of the diagonal elements of the diffusion tensor (in units of 10’ m2/s) are 0.10, 0.12, and 0.017 for the x-, y - , and z-components, respectively, at 150 K. The value of the activation energy for diffusion is 5.2 kJ/mol [5.5 kJ/mol from MD calculations (11)].Both these activation energies are significantly lower than experimental values determined from NMR measurements. The discrepancy seems unlikely to be explained on the grounds of calculation deficiencies alone. The assumption that Xe atoms at low loadings diffuse as monomers rather than dimers has been tested. Santikary et al. (15)found, on the basis of MD simulations of Xe in NaY zeolite, that approximately 15% of Xe atoms diffuse as dimers at a concentration of 1 Xe atom per supercage. From a plot of the R D F of the Xe atoms, June et al. (12) found that approximately 10%were dimerized. This is an important point, as Xe atoms diffusing as dimers rather than monomers present a very different potential energy hypersurface. (This surface was used to locate the position of the minima in the TST method.) Xe diffusion through cloverite has also been simulated (21). Cloverite is a recently synthesized gallophosphate comprising a-cages that are joined by “rpa” cages (27). This arrangement prFduces a principal channel containing large cubic supercages of 29- t o 30-A diameter, accessed via clover-ieafshaped windows of 20 T atoms (T = G a or P) of approximately 6-A diameter. In the calculations, Xe was confined to this channel, and one unit cell of cloverite was simulated with 25 sorbate atoms for 500 ps with a time step of 20 fs. Once again the framework was held rigid. Interaction parameters were from Kiselev and D u (22),and the parameters characterizing the interaction of Xe with the OH groups that protrude into the cloverleaf window were taken from simulations of alcohols. The results are similar to those found for zeolite Y .At low temperatures, Xe appears to be localized at the adsorption site, with an increase in temperature yielding increased delocalization. Diffusion coefficients were calculated from the slope of the graph of MSD versus time, and the values are 0.25 X 10’ m2/s at 397 K and 0.74 X 10’ m2/s at 716 K.

MOLECULAR BASIS OF ZEOLITE CATALYSIS

15

Argon. Calculations complementary to those of Xe in zeolite Y have been performed by Yashonath et al. (20, 28, 29) for the diffusion of Ar in NaCaA. Zeolites Y and A have some structural similarities; each consists of large cages, called a-cages or supercages, and smaller ones, known as 6-cages or sodalite cages. The a-cages have an approximate diameter of 12 A and are tetrahedrally interconnected in zeolite Y and octahedrally interconnected in zeolite A. The cages in zeolite A are linked via eightmembered rings ca. 4 A in diameter, far smaller than the 12-ring windows in zeolite Y. The calculation parameters were very similar to those used in the investigations of Xe diffusion in zeolite Y, except that a larger array of 8 unit cells was used. The Si/Al ratio was 2.0, such that the Na and Ca counterions occupied sites close to the 6-rings (i.e., not in the 8-ring windows that allow access into the a-cages). This seemingly similar system produces some rather surprising results. As for Xe in Nay, the Ar atom was found to be localized near the cage walls at lower temperatures, with the delocalization increasing with temperature. The diffusion coefficients, determined from the slopes of the MSD plots as a function of time, vary between 0.30 X 10' mis at 139 K and 1.08 X 10' m2/s at 585 K. At (close to) room temperature (261 K), a value of 0.65 X 10' m2/s was found. The need for long simulations to obtain good estimates of the diffusion coefficients is demonstrated by these calculations. Earlier simulations by the same authors using the same potential parameters (28) employed a run time of 200 ps for 8 argon atoms, whereas the latter simulations (29) ran for 600 ps and comprised 64 argon atoms. The predicted diffusion coefficients differed by about 15%. An Arrhenius plot of the diffusion coefficient (determined from the longer simulations) as a function of inverse temperature gives an activation energy for diffusion of 1.8 kJ/mol. This is lower than the value found for Xe in zeolite Y (It?),and the reason for the difference is explained in the following paragraphs. The Ar atom is estimated to reside for about 30 ps inside any given cage before diffusing into a neighboring one. The rate of cage-to-cage crossovers for Ar in zeolite A was found to be higher than that of Xe in zeolite Y, even though the ratio of the pore opening to the sorbate size is smaller for Ar in zeolite A than for Xe in zeolite Y [the values are 1.32 and 1.95, respectively (30)].Furthermore, the potential barrier to diffusion-either the surface-mediated or the centralized mode-was found to be negative for Ar diffusing through the 8-rings of zeolite A (Fig. 3). This situation is very different from that representing Xe in zeolite Y, whereby the surface-mediated diffusion is characterized by a potential barrier to cage-to-cage crossings,consistent with the lower activation energy to diffusion of Ar in zeolite A. As was found for Xe, the surface-mediated

16

SIMON P. BATES AND RUTGER A. VAN SANTEN

window

cage

windcm

I

window

cage

winduw

FIG.3. Energy profile as a function of position for the surface-mediated and centralized modes of diffusion of argon in zeolite A. The upper curve refers to the centralized mode of diffusion, and the lower curve to the surface-mediated mode of diffusion. Reprinted with permission from Ref. 29. Copyright 1993 American Chemical Society.

diffusion dominates at low temperature, but as the temperature increases the diffusion mechanism becomes principally centralized. At 600 K, the surface-mediated diffusion comprises only 33% of the total. On the basis of purely geometrical considerations, intuition would suggest that the diffusion of Ar through the smaller 8-rings of zeolite A would be slower than the diffusion of Xe through the 12-rings in zeolite Y. The fact that the reverse trend was found led the authors to suggest that the interactions between host and guest are critical in determining the diffusion resistance of a particular molecular sieve (30). This inference is supported by supplementary calculations whereby the dispersive interactions were omitted from the Lennard-Jones potential function while the repulsive interactions were retained. For the Xe/zeolite Y system, the number of cageto-cage crossovers in these supplementary simulations was approximately equal to the number found when dispersive interactions were present. However, for Ar in zeolite A, no cage-to-cage crossovers were found at all when the dispersive interactions were not present. The authors asserted that the interaction energy of the sorbate as it crosses from one cage to another decreases as the size of the window that connects the cages increases. It was speculated that very large pore materials (VPI-5, cloverite, etc.) will lead to more unfavorable guest-host interactions at the cage windows (30). The most pronounced effect is expected if the sorbate is large and reasonably polarizable, e.g., Xe. The authors also claimed that this effect is not restricted to monoatomic sorbates, and MD calculations of methane diffusion (31) (discussed in the next section) support this claim. This seemingly anomalous result, called the “ring effect,” has been investigated in more detail (20, 32). With a potential expression as simple as that for a single Lennard-Jones particle as the sorbate, it is possible to vary systematically the size (uguest-guest) or interaction parameters, as

MOLECULAR BASIS OF ZEOLITE CATALYSIS

17

was done by Yashonath and Santikary (20) for both zeolites Y and A. The size parameter was varied at a fixed interaction parameter for both zeolites Y and A, and the mass of the sorbate molecule was held constant. Both zeolites showed maxima in the rate of intercage diffusion when the size of the imaginary sorbing molecule approached the size of the ring opening. Furthermore, it was shown that the shapes of the potential energy curves shown in Figs. 2 and 3 were a consequence of the sorbate size and not of the particular zeolite for which they were originally determined. This result raises the important question of whether this anomalous behavior can be shown to be a general feature of MD calculations of diffusion. The authors contended that the feature is general (32). The fact that the maximum in the cage-to-cage crossover rate (or diffusion coefficient) occurs at the same to size of window provides convincing evivalue of the ratio of gguest-guest dence for this contention. We return to this question shortly with details of simulations in other host frameworks. In principle it is possible to verify these findings experimentally. Diffusion measurements characterizing single-crystal samples, free from defects and grain boundaries, should be able to demonstrate that smaller guests do not always diffuse faster. However, there are a number of problems in trying to find suitable experimental results with which to corroborate the claim of the theoretical simulations. One of these problems is the rather ideal distribution of cations that is assumed in the simulations of zeolites A and Y (such that no windows are blocked). An illustration is provided by the early work of Ruthven and Derrah (33). They measured the diffusion of Ar through pellets of zeolite 4A (zeolite A with cations blocking the 8-ring windows) by gravimetric uptake, finding a diffusion coefficient of 3 X m2/s. This value is many orders of magnitude smaller than that calculated from the simulations of Yashonath and Santikary (28,29),and the difference is attributed to the presence of the window-blocking cations. However, intracrystalline diffusion coefficients determined from chromatographic and sorption rate measurements for Ar in zeolite 5A (24) are in reasonable agreement with the simulated values of Santikary (0.2 X lo8 compared with 0.65 X lo8 m2/s, respectively, at around 298 K). A valid question raised about the diffusion anomaly observed by Yashonath et al. is that it may be nothing more than an artifact of the assumed rigidity of the zeolite lattice. The imposition of a rigid zeolite lattice is clearly one of the most serious assumptions in all the calculations described so far, particularly when the size of the adsorbate molecule approaches the size of the framework window. With a rigid lattice, not only are the motions of these atoms ignored, but any energy exchange between guest and host and any dynamic coupling of sorbate and framework vibrations are also ignored. In recognition of this issue, Santikary and Yashonath (34) per-

18

SIMON P. BATES AND RUTGER A. VAN SANTEN

formed calculations for a Lennard-Jones particle diffusing in zeolite A in which framework flexibility is included. A harmonic force field of bondstretching terms was used to account for zeolite-zeolite interactions, as proposed by Demontis et al. (26). This was stated by the authors to “satisfactorily reproduce the main features of aluminosilicate frameworks” and has previously been used in a MD simulation of the structure and dynamics of NaY zeolite (16). The other calculation details were similar to those of previous investigations of this zeolite (20, 29). Results of calculations at various values of ugguest-guest suggest the persistence of the ring effect. However, the peak heights in the diffusion coefficient as a function of uguest-guest were found to be reduced when the flexibility of the framework was introduced. Furthermore, the 8-ring window was found to expand slightly when the atom passes through it, but only when the diameter of the atom is comparable to that of the window. Thus, the influence of the sorbate-zeolite interactions on the framework-however small-can be investigated properly only if the host is allowed to relax. The rate of intercage diffusion was found to increase when the framework was allowed to relax. This result can be rationalized by a change in the mechanism of intercage diffusion. The flexible framework calculations show a significant increase in the proportion of diffusion propagated via the centralized rather than the surface-mediated route; this result was inferred from a change in the potential energy surfaces of surface-mediated and centralized diffusion. (As the potential energy near the surface of the cage increases, the surface-mediated route is disfavored, whereas no such change is observed in the potential near the cage center.) The increase in the rate of intercage diffusion was estimated to be 1.520%. The changes observed upon allowing relaxation of the framework (employing even a modest description of the zeolite) suggest that certain details of the ring effect are artifacts due to calculation approximations, although the overall effect was still found. The existence of the ring effect has been investigated for the diffusion of Lennard-Jones particles in other frameworks (35, 36). Yashonath and Bandyopadhyay (35) presented results for the silicalite framework. In contrast to the frameworks based on a-cages that have displayed the ring effect, silicalite has a totally different pore structure comprising two slightly elliptical intersecting channels, one straight and the other zigzagged. Calculations were performed with both rigid and flexible lattices. Because of increased computational demands of allowing the framework to relax, flexible lattice simulations were done with a smaller system, together with a shorter time step relative to the rigid framework simulations. The adsorbate atoms were distributed evenly throughout the channels at the start of each simulation, and a nominal temperature of 300 K was selected. The size

MOLECULAR BASIS OF ZEOLITE CATALYSIS

19

parameter of the adsorbate, uguest-guest, was varied between 1.5 and 4.5 A. Once again, as in the case of zeolites Y and A, a peak in the diffusion coefficient was found as the diameter of the guest approached the diameter of the channel. This feature persisted when framework flexibility was included. The V P I J framework has also been investigated to determine whether it produces the ring effect for a diffusing Lennard-Jones particle (36). VPI-5 is an aluminophosphate molecular sieve, the channel system of which comprises a straight 18-ring pore parallel to the z-axis, with a free diameter of 13 x 21 A. The simulations were done with an approximate loading of one sorbate (of mass 40 amu) per channel. A time step of 5 fs and a temperature of 620 K were used. A peak in the diffusion coefficient as the sorbate diameter approached that of the window was once again found (13.2 A was used as the window diameter as it is the dimension of the minor axis of the elliptical channel). The results not only show that the diffusion anomaly appears to be independent of the host zeolite but also show that it appears to be independent of the topological features of the host micropore. The concept of the ring effect has recently been applied to the diffusion of a binary mixture of Lennard-Jones particles in zeolite NaY (37). The first particle was varied in size while the second was held fixed. The results suggest that when the diameter of the larger sorbate is close to that of the 12-ring window in zeolite Nay, the larger sorbate will diffuse faster than the smaller one. 2. Water Leherte et al. (38-40) reported a series of MD calculations for water molecules adsorbed in the ferrierite framework. The ferrierite lattice is modeled with a %/A1 ratio of 8, with the siting for aluminum atoms taken from ab initio calculations. (The T4 site is preferentially substituted.) Four unit cells were included and the lattice and intramolecular water parameters kept rigid. Three different water concentrations were considered: 23, 33, and 41 molecules in 4 unit cells. The interaction parameters for the water molecules were taken from nonempirical configuration interaction calculations for water dimers (41) that have been shown to give good agreement between experimental radial distribution functions and simulations at low sorbate densities. The potential terms for the water-ferrierite interaction consisted of repulsion, dispersion, and electrostatic terms. The first two of these terms are the components of the 6-12 Lennard-Jones function, and the electrostatic term accounts for long-range contributions and is evaluated by an Ewald summation. The

20

SIMON P. BATES AND RUTGER A. VAN SANTEN

importance of the long-range electrostatic term has been emphasized in previous MD studies by the same authors (38, 39). Partial atomic charges for all atoms involved in the electrostatic contribution were determined from STO-3G quantum chemical calculations, with the assumption that the zeolite is only partially ionic. All simulations were performed for 26 ps, after an equilibration period of 3 ps. For this length of simulation, the water molecules do not diffuse randomly but oscillate around equilibrium positions (as shown by the nonlinearity of the MSD versus time curves). Thus it is rather difficult to obtain precise diffusion coefficients from the simulations. An order of magnitude value of m2/s was quoted, similar to experimental measurements of water diffusion in other hosts (24). The concentration dependence of the predicted diffusion coefficient goes through a maximum value at the intermediate concentration of 33 water molecules in 4 ferrierite cells. This result is characteristic of the dependence predicted for small polar molecules interacting with zeolite acid sites (42).The initial increase in the diffusion coefficient observed for a concentration increase from 23 to 33 water molecules is explained by adsorbed molecules strongly interacting with acid sites, leaving others free. As the concentration increases further, a decrease in diffusion coefficient is observed as a result of mutual intermolecular interaction. A comparison of the estimated diffusion coefficient with those found experimentally for liquid water and ice suggests that the state of the adsorbed water molecules is characterized by values closer to those of the liquid state than the solid state. 3. Methane

Considerable effort has been directed toward simulation of the diffusive behavior of methane, particularly in silicalite. Predicted diffusion coefficients and activation energies, together with a representative selection of those determined by experimental methods, are collected in Table I. A thorough study of the concentration and temperature dependence of the diffusion of methane in NaY zeolite has been presented by Yashonath et al. (43, 44). Zeolite Y was modeled with a Si/A1 ratio of 3.0, which gave a cation distribution equivalent to that used to investigate the diffusion of Xe in NaY zeolite (13,15,17,18,20), as discussed earlier. The interactions between methane molecules were modeled using the modified RMK potential of Meinander and Tabisz (45);the parameters were determined by fitting to solid- and gas-phase properties of methane. The potential comprises a dispersive part, which acts only over the carbon atoms and is a sum of r6, 6 , and r1' terms, and a repulsive term, which acts over all atoms and is based on an exponential term. The interactions between guest and the

MOLECULAR BASIS OF ZEOLITE CATALYSIS

21

(rigid) zeolite host atoms were modeled with a 6-12 Lennard-Jones function, with the parameters taken from a previous Monte Carlo study of the adsorption of methane in NaY zeolite (46), as it appeared to yield good results. One cell of faujasite was used in the simulations, with periodic boundary conditions. The Coulomb contribution to the interaction was estimated by assigning charges from MO calculations and using an Ewald summation technique. It was found that this contribution was approximately 20% of the total interaction. Temperatures between 50 and 300 K were simulated with a fixed methane concentration of 6 molecules per supercage, and loadings between 2 and 8 molecules per cage were simulated at a fixed temperature of 300 K. Simulations were allowed to run for 25-45 ps with a time step of 1 fs. At higher simulation temperatures (220 and 300 K), the adsorbates appear to exhibit fluidlike behavior; evidence for this includes a lack of definite peaks in the RDFs between host and guest atoms, and this is supported by neutron time-of-flight measurements characterizing methane in faujasites (47). At lower temperatures, the methane molecules become increasingly likely to be found at or near the preferred site of adsorption close to the wall of the cage. The RDFs between the hydrogens of methane and the zeolite atoms indicate that the influence of the zeolite on the guest is considerable and the rotation and translational motion of the methane molecule is restricted at these low temperatures. These results are in accord with proportions of trapped and free methane molecules predicted from neutron scattering experiments (48). The estimated diffusion coefficients range from 0.13 X lo-' m2/s at 50 K to approximately 2 X lo-' m2/s at 300 K, but the error in the lowtemperature values is higher because of the short simulation time. Comparison of the simulated values at 300 K with those determined by NMR spectroscopy (42) (1.5 X lo-' m2/s at 300 K) and neutron scattering experiments (48) (0.9 X lo-' m2/s at 250 K) characterizing a sample of NaX zeolite at a concentration of 6 methane molecules per supercage shows reasonable agreement. However, the uncertainties in the low-temperature simulations are exemplified by a calculation of the activation energy for diffusion; it is approximately 6 times smaller than the value of 6.3 kJ/mol determined from a neutron scattering study (48). The calculated power spectra as a function of decreasing temperature reveal the center-of-mass vibrations as the molecule becomes increasingly trapped around the minimum energy site. The three frequencies are 36, 53, and 83 cm-' at 50 K. The mechanism of cage-to-cage diffusion was investigated by plotting the distance between the center of mass of the methane molecule and the center of the parent supercage versus the distance to the center of the daughter supercage (as has been described for Xe diffusion in NaY zeolite).

TABLE I Diffusion Coeficient ( D ) and Activation Energy (Eo) Data for the Diffusion of Methane in Various Zeolites Determined from Simulations and Experimental Methods" Zeolite type

Reference

D (10"

X

m2is)

E, (kJ/mol)

NaY Simulation

Experiment All Si A Simulation

-

?lsc

300 300

6lsc

Infinite dilution 6lsc

0.9

250

6lsc

0.25 0.08

360 300 300 300

1isc disc

0.08 -

0.049

Experiment

Number of CH4

293

4 1.5

0.35

NaCaA Simulation

Temperature (K)

0.03 0.05

- ve -

7

hlsc hlsc

-

?/sc

173 173

6lsc

300

hlsc 1isc

Framework rigid rigid rigid

flexible rigid flexible rigid rigid rigid

-

-

1.60 0.25 1.05 0.75 0.65 1.17 0.21

300 200 300 300 300 446 167

81uc Infinite dilution 8Iuc 12luc 1-3luc 81uc 81uc 0.51~~ 12luc 121uc 41uc 4luc 161uc 41uc 41uc 81uc 1luc 8Iuc 121uc 121uc 121uc

-

-

4-12luc

0.5 -

0.75 0.66 -1 0.74 0.62 0.36 0.83 0.30 1.61 0.90 0.10

200 300 300 300 300 300 300 446 167 300 300 300

’“sc” denotes supercage and “uc” denotes unit cell. There are 8 supercages to a unit cell of Y and A.

rigid rigid rigid flexible rigid rigid rigid flexible flexible flexible flexible rigid rigid rigid rigid -

-

-

24

SIMON P. BATES AND RUTGER A. VAN SANTEN

It was found that the diffusion of methane from one cage to another takes place by the surface-mediated route, whereby the methane molecule skates along the inner surface of the cages. At 300 K, the residence times of a methane molecule within a given cage were found to be of the order of a few picoseconds. This is substantially shorter than the residence time of Xe (14), notwithstanding their similar sizes. The reason for this difference lies in the very different masses of the two sorbates. Bandyopadhyay and Yashonath (31), in an extension of their work on MD studies of noble gas diffusion, presented MD results for methane diffusion in NaY and NaCaA zeolites. The zeolite models were the same as those used in the noble gas simulations (13, 15, 17, 18, 20, 28, 29) and the zeolite lattice was held rigid. The methane molecule was approximated as a single interaction center and the guest-host potential parameters were calculated from data of Bezus et al. (49) (for the dispersive term) and by setting the force on a pair of atoms equal to zero at the sum of their van der Waals radii (for the repulsive term). Simulations were run for 600 ps with a time step of 10 fs. Calculation results at 177 K led to the prediction that methane diffuses from one a-cage to another at a rate of 1.9 X 10" per sorbate per second in NaY zeolite and 24.3 X lo1' per sorbate per second in NaCaA zeolite. Thus, this is another example of a guest diffusing faster through a host with smaller pore windows (the so-called "ring effect"), a phenomenon that had been observed previously only for noble gas atoms. The faster cage-to-cage diffusion in zeolite A persists at higher temperatures, although the differences become substantially less. This result is explained by an increase in the proportion of centralized diffusion, as compared to surface-mediated diffusion, as the temperature increases. At low temperatures, the surface-mediated route dominates, and this route presents a potential barrier in the case of NaY zeolite and a well in the case of NaCaA zeolite. As the temperature increases and the centralized route becomes more important, the rate of cage-to-cage crossings becomes approximately equal since both zeolites present a potential well to centralized diffusion. From the temperature dependence of the rate of cage-tocage crossovers, the activation energy for diffusion may be determined. This is found to be negative for NaCaA zeolite and positive for NaY zeolite. [The value of 6.4 kJ/mol for NaY zeolite is almost identical to that from a neutron scattering study (48).] Other simulations of the diffusion of methane in zeolite A have been performed by Cohen de Lara et al. (50), who reported calculations for a single methane molecule in an a-cage of zeolite A. They used a 7-cage array as a model for the zeolite, with cations fully occupying the SI sites, half-filling the SII sites, and occupying only 1112th of the SIII sites. Ionic

MOLECULAR BASIS OF ZEOLITE CATALYSIS

25

charges were assigned to the crystal atoms (which led to an overestimation of the mean field by approximately a factor of 2). The methane sorbate was modeled as a single interaction center, thus precluding any information on the rotational behavior of methane inside zeolite A. The authors assumed that on the time scale of their simulation (375 ps), the methane molecule would be enclosed within the cage and have a negligible probability of jumping into a neighboring cage. This assumption is based upon proton spin-spin relaxation measurements (51) that predict a residence time of 68 ps at 300 K. [We note that this is far longer than the rate of cageto-cage crossovers determined from the calculations of Bandyopadhyay and Yashonath (31) would imply.] The intermolecular interactions are accounted for by a 6-12 Lennard-Jones function, with an electronic contribution to account for the effect of the crystal field on permanent and induced moments of the methane molecule. The trajectory and radial distribution of the methane molecule were predicted from the calculations. There are many ways the trajectory of a particle may be quantified; the authors used the number of times the methane molecule passes through the volume element bounded by r and Sr (where Sr = 0.1 and r has a value between 0 and 5 A). At high temperatures, it was found that the methane molecule was delocalized over the whole cage. As the temperature was decreased, the cage center was never visited, and the molecule sweeps around the walls in the region of the SIII cation, finally locating itself between an SIII and SII cation site. Fritzsche et al. (52-56) published a series of papers considering aspects of the diffusion of methane in zeolite A and its cation-free analogue, ZK4. The first of these considered the thermal equilibrium of the diffusing species (52). In calculations employing a fixed zeolite lattice, the thermal equilibrium of the diffusing molecules was maintained somewhat artificially by renormalizing the particle velocities after each time step. At the beginning of the decade, MD calculations that allowed energy exchange between host and guest and at the same time were performed for a long enough time to ensure that long-range diffusive behavior was simulated were beyond the reach of all but the longest simulations on the fastest machines. Fritzsche et al. sought to investigate how much the thermal equilibrium of the sorbates was maintained by their own mutual interactions. Their simulations suggested that the diffusing molecules are able to thermalize their own kinetic energy to a Boltzmann distribution at any point within the zeolite lattice, regardless of the different potential. So, to a first approximation at least, the energy exchange between the zeolite and the guest can be neglected. Subsequent calculations (53) for methane in ZK4 took account of the applicability of the Einstein relationship to determine the diffusion coefficient, by examining the probability density of finding a particle at a given

A

26

SIMON P. BATES AND RUTGER A. VAN SANTEN

point at a given time. This was done by deriving the first four moments of the distribution curve of molecular displacements; at such a time that the diffusion equation is valid, all moments are compatible with the expression for the probability density. From MD calculations for variable concentrations of methane molecules (which were treated as spherical particles), it was found that after 30-60 ps of simulation time, the diffusion equation was satisfied. By comparison, for NaCaA zeolite this time can be as long as 30 ns (55). Details of the diffusion in cation-free zeolite A as a function of concentration and intermolecular potential parameters were presented in later papers (54, 55).Two different sets of potential parameters were used; the first (set A) was determined from spectroscopic data, whereas the second (set B) had a slightly larger value of CTCH,-O and a smaller value of E C ~ , -than ~ set A. Qualitatively, set A may be thought of as simulating slightly larger windows than set B. Isopotential surfaces, viewed through the center of an a-cage, determined from both sets of parameters showed important similarities. The center of the cage was found to be a region of high potential energy and therefore likely to be free of guests at all but the highest concentrations. Calculations with set A predicted that the 8-ring windows present a potential minimum in the plane of the window, whereas those with set B predicted a smaller minimum near the inlet of the window, with the plane of the window corresponding to a saddle point. Simulations were performed with models comprising up to 343 a-cages, with a methane loading of 1-7 molecules per cage. Runs of 0.75 ns were performed with a step size of 5 fs. The residence times of a methane molecule were investigated by monitoring the passage of particles through a plane perpendicular to an 8-ring window and expressing the probability density of finding a molecule in a given cage as a function of time. A fraction of the methane molecules that cross into a neighboring cage jump directly back into the cage from which they originate. This contribution to the probability density as a function of time was found to increase with increasing loading and decreasing temperature; it seems reasonable that this behavior is brought about by methane-methane collisions. However, a given methane molecule will pass through a window more frequently at a higher loading, even though it is more likely that the increased intermolecular collisions knock the molecule back t o where it came from. Thus, the intracavity collisions help the molecule to find its way out of the cage but are also responsible for knocking it back more frequently. The residence times of a methane molecule in a cage are determined to be ca. 5-10 ps, i.e., when the probability density approaches zero. [This value is in good agreement with the rate of cage-to-cage crossovers calculated by Bandyopadhyay and Yashonath (32).]

MOLECULAR BASIS OF ZEOLITE CATALYSIS

27

The dependence of the diffusion coefficients on loading was found to depend on the parameter set used. Set A (larger window size) predicts a decrease as loading increases, whereas set B predicts the trend to reverse. The experimentally measured concentration dependence of the diffusion coefficient of methane in NaCaA zeolite is an increase at higher loadings (57). However, a direct comparison is not straightforward because of the presence of cations in NaCaA zeolite. Nonetheless, the calculations do not E ~ ~ , that - ~ is , largely show that it appears to be the variation in sCH,-O, responsible for the behavior of the diffusion coefficient. Indeed, from these calculations it would appear that it is the selection of the potential parameters that governs the predicted behavior as a function of concentration, rather than other assumptions such as that of a rigid lattice or no contributions from Si atoms. The logical next step in this sequence of calculations was to perform simulations that included the extraframework cations Na and Ca in the zeolite A structure (56).The cations were situated near the hexagonal faces of the sodalite units in zeolite A such that the geometrical hindrance to methane molecules passing through the 8-ring windows was small. Interactions between the methane molecule and the cations were derived from the data of Ruthven and Derrah (58),and the same two sets of parameters were used ( A and B). Long simulations were performed, 25 ns with a time step of 10 fs. It was found that the diffusivities of methane in NaCaA zeolite were far lower than for those found for ZK4; the difference was found to be as large as 2 orders of magnitude, depending on the parameter set used and concentration. The results from parameter set A (larger effective windows) were found to be consistently better than those from set B, as judged by the comparison to PFG-NMR measurements (57). A t loadings greater than 5 methane molecules per cage, the results from set A were found to be in good agreement with those from NMR measurements; for example, the values at 7 molecules per cage are 5.52 X lo-'' and 5 X lo-'' m2/s for the simulated and experimental values, respectively. The overall trend in the calculated values predicts an increase in the diffusion coefficients as a function of concentration, in agreement with NMR measurements (57). The poor statistics from some of the simulations, particularly those of set B, are explained in terms of the inadequate length of even these very long simulations as a result of the slow diffusion in the presence of chargebalancing cations. However, the length of the simulation may not be the only factor that leads to uncertainties in the calculated diffusion coefficients; the approximation of a rigid framework acts so as to reduce calculated diffusion coefficients, as described in the following paragraphs. This point was highlighted by Demontis and Suffritti (59) for a simulation

28

SIMON P. BATES AND RUTGER A. VAN SANTEN

of methane in cation-free zeolite A, employing both fixed and flexible frameworks. The force field parameters used to account for the motion of the zeolite atoms were those used to reproduce the main structural features of zeolite A (26) and for purposes of this calculation the methane molecule was considered to be a single spherical particle. Calculations were performed for a maximum of 6 ns in the low loading limit of 1 methane molecule per a-cage. The diffusion coefficients calculated from a simulation employing a flexible framework were all between 5 and 10 times larger than those calculated from fixed lattice simulations. A comparison between flexible framework results and NMR measurements (57) illustrated the influence of the cations in the experimental sample; calculated diffusion coefficients from the cationfree (flexible) framework were approximately 5 times higher than the experimental results. The increase in diffusion coefficient as a function of loading found in experimental studies was reproduced by the simulations. Three different diffusive regimes were highlighted in this work; the first is the quasi-free motion of the particle on the picosecond time scale, when the influence of the vibrations of the framework are not yet established. The intracavity diffusion lasts for tens of picoseconds and is dependent on the loading. The long-range diffusion is controlled by the 8-ring windows and is the origin of the large differences in calculated diffusion coefficient for the fixed and mobile frameworks. The breathing of the 8-ring windows, of an amplitude up to 0.3 A, is critical in enhancing the long-range diffusion. Considerable research has been directed toward the simulation of the diffusive properties of methane within the pores of silicalite ( f f , 2 3 ,60-69). June et al. ( I f ) presented variable temperature and concentration MD calculations. (Calculation details were the same as for their simulation of Xe in silicalite, described earlier, with methane treated as a rigid polyatomic with six degrees of freedom.) At a temperature of 200 K, the diffusion coefficient was found to be largely independent of concentration, but a marked decrease with increasing concentration was observed at higher temperatures (in contrast to the calculations of den Ouden et al. (60),which showed the diffusion coefficient of methane at 300 K to be approximately 1.2 X m2/s,independent of loading). The values of the diffusion coefficient calculated by June et al. (11) are in good agreement with those from NMR measurements (70, 71). A t 200 K, the simulations overestimate the diffusion coefficient by a factor of 2, and at 300 K theoretical and experimental values are coincident. The activation energy for diffusion was calculated to be 5.6 kJ/mol in the limit of infinite dilution, which is close to the value found for NaZSM-5 (4.7 kJ/mol) (72). The anisotropy of the diffusion process is most pronounced in the limit of infinite dilution in the absence of intermolecular collisions that scatter

MOLECULAR BASIS OF ZEOLITE CATALYSIS

29

adsorbate molecules. These collisions act to reduce the larger components of the diffusion coefficient (parallel to the y- and x-axes, corresponding to motion down the straight and sinusoidal channels, respectively). Diffusion down the straight channel was predicted to be only slightly faster than down the sinusoidal channel, in contrast to the marked preference exhibited by Xe for the straight channel (10). Estimates of the rotational diffusivity may be made from MD calculations by fitting an exponential function to Legendre polynomials that express the decorrelation of a unit vector that is fixed in the methane coordinate frame (11).The rotational diffusivity was found to increase with concentration (as a result of sorbate-sorbate collisions which act to decorrelate the molecular orientation). The values are of the same order as those for liquid methane and are 2 orders of magnitude larger than those found by Jobic et QZ. (73) from a quasi-elastic neutron scattering study of methane in NaZSM-5. Similar rigid lattice simulations have been reported by Goodbody et al. (62). Methane was once again considered as a single spherical particle. The interatomic potential parameter Eguest-host was fitted to reproduce the experimentally measured Henry’s law coefficient for methane in silicalite. Simulation lengths were dependent on the concentration of methane molecules in the 8 unit cell simulation box; times ranged from 5 ns in the limit of low concentration (1 molecule per unit cell) to 600 ps at a concentration of 16 molecules per cell. All simulations were for room temperature. The calculated diffusion coefficients in all three Cartesian directions showed a smooth decrease with increased loading. The averaged value at 8 molecules per unit cell was identical to that of June et al. (11) (0.75 X m2/s). Diffusion was predicted to be strongly anisotropic; diffusion parallel to the y-axis (corresponding to motion along the straight channel) was predicted to be twice as fast as along the sinusoidal channel. At low loading, transport through the straight channel was found to be nearly an order of magnitude faster than motion parallel to the z-axis. These ratios are comparable to those of Pickett et aZ. (10) for the simulation of Xe in silicalite, but they are in direct disagreement with those of June et al. ( 1 1 ) described earlier. Fixed framework-flexible adsorbate calculations were also reported by Dumont and Bougeard (68, 69). The diffusion coefficient calculated from a 42-ps calculation of 4 molecules per unit cell at 300 K was 1.60 X m2/s. Once again the anisotropy of the diffusive process was calculated to be strong; transport through the straight channel was found to be twice as fast as through the sinusoidal channel and an order of magnitude faster than motion parallel to the z-axis.

30

SIMON P. BATES AND RUTGER A. VAN SANTEN

The simulation length may be critical in determining the validity of a direct comparison between calculated and experimental diffusion coefficients, as noted by Nowak et al. (63).They studied the diffusion of methane in three all-silica zeolite frameworks: silicalite, mordenite, and EU-1. (Both mordenite and EU-1 may be thought of as having a unidimensional pore system, with side pockets perpendicular to the main channel direction.) The diffusion coefficient of methane in silicalite is in reasonable agreement with previous simulated and experimental values, 0.62 X m2/s at 298 K and a loading of 2 molecules per unit cell. It was found to be by far the largest of the three values; the reason for this is that the time scale of the simulation is sufficiently short for diffusion or trapping of the methane molecules by the side pockets of mordenite and EU-1 to contribute significantly to the overall diffusion coefficient. This trapping reduces the rate of movement along the principal channel direction. In a very long simulation, or on the time scale of a PFG-NMR experiment, the molecular motion in the directions of the side pockets would be averaged out, and the resulting diffusion coefficient would be nearer the value found in shorter simulations for diffusion along the principal channel axis. Nowak et al. found these values to be similar for all three zeolites and concluded that a long simulation or experimental measurement would predict approximately equal diffusivities for all three structures. Nicholas et al. (67)presented a thorough discussion of methane diffusion in silicalite. Test calculations were initially performed to ascertain the best of four sets of potential parameters (assessed by their ability to reproduce heats of adsorption and diffusivities).The MM2 parameters (74)were found to be the best. A large simulation box of 27 unit cells was used, into which a large number of methane molecules (up to 216) were placed, in independent sets so as to obtain accurate statistical averages of physical properties. The zeolite lattice and methane molecules were assumed to be rigid, but in contrast to all previous studies, intermolecular interaction contributions from the silicon atoms of the lattice were included. Simulations were run for 60 ps using a step size of 1.5 fs. An electrostatic term was included in the potential energy summation, and the atomic partial charges for the zeolitic part of this term were taken from a study of the structure of all-silica sodalite (75). The charges for methane were determined as those that reproduced the electric field around the molecule. Variable concentrations, from infinite dilution to 16 methane molecules per unit cell, were considered at a temperature of 300 K. The calculated diffusion coefficients from this investigation are in good agreement with other theoretical and experimental values. Values showed a decrease with increasing methane concentration, from 0.9 X lo-* m2/s at 4 molecules per

MOLECULAR BASIS OF ZEOLITE CATALYSIS

31

unit cell to 0.1 X m2/s at 16 molecules per unit cell. The activation energy for diffusion was found to be 4.3 kJ/mol. The mechanism of molecular diffusion in zeolites has often been described in terms of a jump model, whereby molecules reside in energetically favorable sites until they acquire enough energy to overcome potential barriers. Car0 et al. (71) proposed such a model, suggesting that methane diffuses through silicalite by jumping between intersections. From the MSD plots of individual molecules, diffusion of methane in the simulations of Nicholas et al. does seem to be somewhat jumplike, especially at higher methane loadings. However, individual molecules were demonstrated to show inhomogeneous diffusive behavior. This result strengthens the authors' arguments for using a large number of sorbate molecules to obtain good statistics. By decomposing the velocities of sorbates into translational and rotational components, the possibility of coupling in the sorbate motion can be investigated. The frequency spectra, obtained by Fourier transforming the VAFs, indicated little or no coupling. Demontis et al. presented a detailed investigation of the diffusion of methane in silicalite (23,61,65).Their calculations took account of a flexible zeolite framework, with potential parameters taken from work successfully reproducing the main structural features of zeolite A (26). The methane molecules were simulated as spherical particles, with methane-zeolite interactions taken from the work of Ruthven and Derrah (33).The loading was fixed at 12 methane molecules per unit cell (3 per channel intersection) and, at the start of a simulation, the methane molecules were distributed over positions occupied by the CH3and CH2 groups of the organic template used to synthesize ZSM-5. The step size was 1 fs, and simulations were allowed to proceed for 200 ps at room temperature. The authors admitted that their force field for the flexible zeolite lattice was somewhat crude, yet reasonable, being an efficient compromise to the problem of long simulation times associated with slow diffusion behavior. In fact, the force field used in this work costs only approximately 50% more computer time than that used when the framework is kept rigid. The force field was demonstrated in a later work (61) to reproduce sorbate-induced symmetry changes adequately only at low temperatures. At high temperatures, the effects are overestimated, with the possibility that this may affect the diffusion mechanism. The calculated diffusion coefficient is in remarkably good agreement with a value found from NMR measurements at the same temperature and methane loading; the simulations predicted a value of 0.66 X lo-' m2/s, and the experiments gave (0.65 2 0.1) X lo-' m2/s (76). The uncertainty of the theoretical result is difficult to gauge. When the framework is fixed,

32

SIMON P. BATES AND RUTGER A. VAN SANTEN

the calculated values of the diffusion coefficient are approximately 20% smaller, in agreement with the same authors’ findings for methane in dealuminated zeolite A (59). The unidimensional diffusion coefficients along each of the three Cartesian axes were also calculated. In the oscillating framework, the ratios of the values were found to be approximately 3 : 12 : 1 for the x-, y-, and z-directions, respectively. In the rigid framework, the values (on the same scale) were found to be in the ratio 4 : 8 : 1. The results indicate a large increase in the diffusion coefficient parallel to the y-axis (the direction of the straight channel). It was argued that a possible explanation for this enhancement is that the methane molecules, which the R D F shows to favor the centers of the channels, essentially float unhindered along the straight channel. In the sinusoidal channel, the oscillating zeolite framework acts to tighten up the diameter of the channels, thus slowing diffusion. The moving framework may be thought of as acting like a thermal bath on the sorbed molecules, and this is reflected in a far wider spread of temperatures of the methane molecules than in the case of the fixed framework. There are experimental results that show the anisotropic nature of diffusion of methane in silicalite (24, 77). From a stochastic jump model of the diffusion process, Karger et al. (24) found that the ratio of the rate of diffusion in the direction of the two channel systems should not be less than 4.4 times that in the orthogonal direction:

Zibrowius et al. (77) used an NMR spin-echo attenuation technique to estimate the ratio of the diffusion tensor element related to motion along the y-axis (the straight channel) and the average of the other two elements. For a methane loading of 8 molecules per unit cell, the ratio was estimated to be less than 5 at room temperature:

The MD values of Demontis et al. give ratios somewhat larger than these for the flexible framework and values between 3 and 4.5 for the fixed framework. The ratios determined from this work are substantially larger than those from the fixed framework simulations of June et al. ( 2 1 ) and Goodbody et al. (62). It has been stated (23) that the temperature dependence of the calculated diffusion coefficients can illuminate discrepancies between calculated and experimental values, even when results at a given temperature are in good

MOLECULAR BASIS OF ZEOLITE CATALYSIS

33

agreement. Thus, the temperature dependence of the calculated diffusion coefficient is an important relationship to investigate. Demontis et al. (65) performed MD calculations characterizing diffusion at four temperatures between 446 and 167 K, using the same calculation parameters as in their previous simulations (23). The differences in the calculated diffusion coefficients for the fixed and flexible framework models were small at all the temperatures considered. These results are in marked contrast to the results of simulation of methane in a cation-free zeolite A framework performed by the same authors (59),whereby a difference of up to 1order of magnitude was seen. At low temperatures, the predicted diffusion coefficients tended to be larger than those found from NMR measurements, whereas at temperatures greater than 350 K they tended to be smaller. The smallest discrepancy between experimental and theoretical values was found for diffusion at 300 K (the values are almost coincident), while the largest-an underestimation by the calculations of 30%-was found for diffusion at 446 K, the highest temperature considered. The ratio of the unidimensional diffusion coefficients derived by Zibrowius et al. (77) [Eq. (6)] was reasonably well reproduced at all temperatures. That of Karger et al. (24) [Eq. ( 5 ) ] was consistently higher than the lower limit of 4.4. The behavior of the diffusion coefficient as a function of reciprocal temperature showed characteristic Arrhenius behavior, with an activation energy (2 kJ/mol) somewhat lower than that found experimentally (4 kJ/mol) (24). Unidimensional activation energies, corresponding to the barrier to diffusion along a particular Cartesian direction, were also calculated, confirming the conclusion of Pickett et al. (10)from a simulation of Xe in silicalite; preexponential factors were found to be largely responsible for the different diffusivities, as there is little difference in the unidimensional activation energies. The values for fixed and flexible frameworks are remarkably similar. Other flexible framework calculations of methane diffusion in silicalite have been performed by Catlow et al. (64, 66). A more rigorous potential was used to simulate the motion of the zeolite lattice, developed by Vessal et al. (78),whose parameters were derived by fitting to reproduce the static structural and elastic properties of a-quartz. The guest molecule interactions were taken from the work of Kiselev et al. (79), with methane treated as a flexible poiyatomic molecule. Concentrations of 1 and 2 methane molecules per 2 unit cells were considered. Simulations were done with a time step of 1 fs and ran for 120 ps. The slope of the MSD curve as a function of temperature was not constant, indicating that the methane molecules were trapped in a region of the lattice (for a time estimated by Catlow et al. as 5 ps). The diffusion coefficient, calculated as the slope of the MSD graph, is in satisfactory

34

SIMON P. BATES AND RUTGER A. VAN SANTEN

agreement with the NMR value of Car0 et al. (71).The diffusion coefficient was predicted to increase by a factor of 3 with a concentration increase from 1 to 2 methane molecules per cell, in contrast to what was observed experimentally and inferred from other simulations. The methane molecules were predicted to favor the center of the straight channel, but at higher temperatures, motion through the sinusoidal channels was also indicated. Calculations at higher methane loadings (66) (4 molecules per unit cell) led to diffusion coefficients that were somewhat larger than those found experimentally and to an activation energy for diffusion of 3.6 kJ/mol. 4.

Other Hydrocarbons

We now consider the diffusion of other hydrocarbons. Most calculations have been performed for n-alkanes, up to and including n-hexane, but alkenes and alkynes have also been considered. Calculations involving larger molecules have been mainly restricted to silicalite. Nowak et al. (63) presented a comparative study of the diffusivities of rigid models of methane, ethane, and propane in silicalite. (The details of the calculation are reported in the preceding section.) The calculated diffusion coefficients decreased as the length of the carbon chain increased, and the effect was found to be far more pronounced for ethane than propane. The calculated diffusivities, in units of 10' m2/s, were 0.62, 0.47, and 0.41 for methane, ethane, and propane, respectively. The ethane value is in satisfactory agreement with PFG-NMR measurements [0.38 (71),0.3 (80),0.4 (42)] for silicalite. The value for propane, however, was calculated to be almost an order of magnitude larger than the NMR results of Briscoe et al. (80). [The agreement with the value of Car0 et al. (71) is better, but still an overestimation.] The origin of this overestimation is believed to be the rigid model adopted for the sorbates, and this supposition is supported by values characterizing the rate of diffusion parallel to each Cartesian axis. Ethane and propane have almost equivalent diffusion coefficients in the straight and sinusoidal channels, and all these are lower than the corresponding values for methane. Motion from one intersection to another is predicted to be almost negligible for propane. As already discussed, this diffusion involves a tortuous route through neighboring straight and sinusoidal channel segments. This negligible rate of diffusion indicates that the molecule is virtually unable to move from one channel segment to another to effect an overall translation. The inflexibility of the carbon backbone is cited as a possible explanation for this slow diffusion; however, the rather short length of the simulations is not considered as a possible cause for the discrepancy. The potential

MOLECULAR BASIS OF ZEOLITE CATALYSIS

35

parameters used in the calculations are not thought to have contributed to this error; they were shown to reproduce the heats of adsorption of the alkanes reasonably well. Dumont and Bougeard (68,69)reported MD calculations of the diffusion of n-alkanes up to propane as well as ethene and ethyne in silicalite.Thirteen independent sets of 4 molecules per unit cell were considered, to bolster the statistics of the results. The framework was held rigid, but the hydrocarbon molecules were flexible. The internal coordinates that were allowed to vary were as follows: bond stretching, planar angular deformation, linear bending (ethyne), out-of-plane bending (ethene), and bond torsion. The potential parameters governing intermolecular interactions were optimized to reproduce infrared spectra (68). The molecules were predicted to travel through the center of the channel, avoiding the channel intersections. This result is in agreement with the similar calculations of Demontis er al. (23) and June et al. (11).While the centers of mass of molecules translate around the channel centers, libration also occurs to permit a closer approach of some C and H atoms to the zeolite walls (68). The calculated diffusion coefficients for the n-alkanes, estimated to be accurate to within 15%,are in good agreement with experimental results. The values for ethane and propane, in units of 10' m2/s, are 0.59 and 0.19, respectively; the value for ethane is slightly overestimated and that for propane slightly underestimated. In contrast to the findings of Nowak et al. (63), the differences between the unidimensional diffusion coefficients in the straight and sinusoidal channels for ethane and propane are more distinct. Both values are approximately 3 times larger for ethane. The calculation of the ratios of monodimensional diffusion coefficients (24, 77) gives rather high values, due to the short simulation time (42 ps) and therefore poor statistics for motion between intersections. The authors stated that the influence of molecular flexibility on the calculated diffusional data is not obvious and also suggest that the influence of allowing framework flexibility is small. [The latter statement is based on MD results of Demontis ef al. (23) obtained for methane]. A useful comparison may be made between bent, flat, and linear molecules by considering the diffusion coefficients for ethane, ethene, and ethyne. In the sinusoidal channel, ethane diffuses the slowest, ethene approximately 30% faster, and ethyne 3 times as fast. In the straight channel and parallel to the z-axis, ethene and ethyne both diffuse approximately 3 times faster than ethane. These ratios are consistent with the relative crosssectional area of the three C2 hydrocarbons. The value of the diffusion coefficient for ethene in silicalite (1.29 x lo-' m2/s) is an order of magnitude larger than that from a similar MD study

36

SIMON P. BATES AND RUTGER A. VAN SANTEN

undertaken by Catlow et al. (64) (0.15 2 0.5 X lo-* m2/s). The latter work employed a far lower concentration of ethene molecules. Catlow et al. suggest that the simulations point to a preference for diffusion through the sinusoidal channel, rather than the straight channel. Nicholas et al, (67) have performed MD calculations of propane in silicalite in which the propane molecule is given complete flexibility. The calculations, which have been detailed previously for methane diffusion, employed a large simulation box with multiple sets of adsorbates to ensure good statistics. The framework was kept fixed and data were collected over a 40-ps run. The results predict diffusion coefficients in very good agreement with the values of Car0 et al. (71).The calculated values for a concentration of 4 and 12 propane molecules per silicalite unit cell are 0.12 and 0.005 X m2/s, respectively. These values for propane are far lower than those of Nowak et al. (63); the reason for this is that Nicholas et al. used flexible adsorbate molecules, whereas Nowak et al. used rigid ones. Bell et al. (81) presented forced diffusion calculations of butene isomers in the zeolite DAF-1. DAF-1 (82) is a MeALPO comprising two different channel systems, both bounded by 12-rings. The first of these is unidimensional with periodic supercages, while the other is three-dimensional and linked by double 10-rings. The two channel systems are linked together by small 8-ring pores. It is a particularly useful catalyst for the isomerization of but-1-ene to isobutylene (83);its activity and selectivity are greater than those of ferrierite, theta-1, or ZSM-5. Diffusion calculations were performed by using the MSI ( 5 ) constrained diffusion software. Force field parameters were taken from the work of Hagler et al. (84)for the all-silica DAF-1. Coulombic effects were neglected, the host was held rigid, and the guest alkene was allowed to flex freely. From the calculation, it did not appear possible for any of the four butene isomers (but-1-ene, isobutylene, cis- and trans-but-2-ene) to pass through the 8-ring windows that connect the two channels. Thus, the two channel systems are isolated for these sorbates. The limitation of a flexible framework was believed to be partly responsible for this result. Flexible lattice calculations may give indications of diffusion via the %rings, although probably with a high barrier. Favorable sorption sites for the butene isomers were found to be the double 10-rings of the three-dimensional channel system. Thus, diffusion was investigated between adjacent 10-rings. Results showed that the diffusion was an activated process; the lowest barrier was 17.5 kJ/mol (but-lene) and the largest 22.5 kJ/mol (cis-but-2-ene). The authors concluded that all of the three-dimensional channel system is accessible by the four butene isomers, since the diffusion barriers are small enough to be overcome at ambient temperatures.

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As the size of the sorbates increases, it becomes increasingly likely that the sorbates have to be treated as a chain of united atoms, according to the model of Ryckaert and Bellmans ( 3 ) . This is indeed the case for the MD simulations of n-butane and of n-hexane discussed shortly (62,85, 86). Goodbody et al. (62) presented simulation results for four butane molecules in a unit cell of silicalite at 298 K. The bond lengths and angles of the butane molecules were fixed at 1.53 A and 109.45",respectively. All interactions were modeled by using a Lennard-Jones potential; intramolecular interactions were accounted for by using the dihedral potential of Ryckaert and Bellmans ( 3 ) .The calculated diffusion coefficients showed the characteristic anisotropy; diffusion in the straight channel was found to be twice as fast as in the sinusoidal channel and almost an order of magnitude faster than motion parallel to the z-axis. The orientationally averaged value of the diffusion coefficient was found to be 0.17 X lo-* m2/s, approximately an order of magnitude larger than that determined from PFG-NMR measurements (24). June et al. (85) presented united-atom calculations for butane and for hexane in silicalite, whereby the bond and dihedral angles of the alkanes were allowed to vary. In addition, the calculation of hexane took account of an additional intramolecular Lennard-Jones potential for nonbonded atoms more than three bonds apart (which prevents the alkane crossing over itself). The interaction parameters for the alkane molecules were taken from Ryckaert and Bellmans ( 3 ) , and those governing the interaction of the alkanes with the zeolite from a previous study of the low-occupancy sorption of alkanes in silicalite (87). Variable loadings of alkanes were considered; from 1to 8 molecules per unit cell were considered, and calculations were allowed to run for 500 ps for diffusion at 300 K. The diffusivities show a monotonic decrease with increasing alkane loading. The value for n-butane at an intermediate loading of 4 molecules per m2/s. PFG-NMR measurements predict unit cell is (0.24 2 0.08) X m2/s(24),with membrane measurements values no larger than 0.015 X up to 2 orders of magnitude lower. (Alkanes as long as butane and pentane are the longest for which PFG-NMR techniques may be applied.) Similar magnitudes of discrepancy were found between the calculated and literature values for hexane. The MD simulations predicted a value of (0.1 C 0.08) X m2/s, and experimental measurements were of the order of m2/s. It is not easy to quantify just what proportion of these discrepancies may be attributed to calculation assumptions or a biased comparison of theoretical and experimental data. Examples of effects that may lead to an unequal comparison are crystal defects, other physical processes in sorption uptake measurements, etc. The assumption of a rigid lattice has been shown to be relatively unimportant for calculations of methane diffusion (23);the

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same is not necessarily true for a larger molecule with more degrees of freedom. The potential parameters are assumed to give a reasonable description of interatomic interactions, as they have been found to reproduce Henry’s law coefficients well. The strong anisotropy of diffusion was predicted for both alkanes. The ratio suggested by Karger (24) based on a stochastic jump model [Eq. ( 5 ) ] is predicted to be above the lower limit of 4.4 for all loadings, with the exception of 8 molecules per unit cell for butane. This loading is close to the saturation loading for butane, and it is unlikely that the diffusion regime is well characterized by a jump model. The ratio of the diffusion coefficients in the straight and sinusoidal channels was found to be higher for hexane than for butane (at a loading of 4 molecules per unit cell). This result reflects the larger torsional barriers that must be overcome for hexane to diffuse through the sinusoidal channels. The microscopic motion of the molecules may be investigated in detail with the velocity autocorrelation function (VAF) and its Fourier transform (FT). Findings agree with previous simulations for Xe and methane (11); the low-frequency component of the FT of the VAF corresponds to motion through the channel segments between potential barriers at the intersections. The high-frequency component corresponds to a rattling motion perpendicular to the channel axis, on a length scale of 1-2 A. It was suggested by the authors that very long simulations (as long as 1 ns) need to be performed for molecules such as n-hexane. The reason for this is the observation of two long-time-scale dynamical processes in the simulations which need to be probed to predict properly the long-range diffusional behavior. These are the interchange of molecules between straight and sinusoidal channels and conformational isomerization of sorbates about nonterminal carbon-carbon bonds. The interchange of molecules between channels was examined by investigating the decay of a normalized end-toend vector correlation function of the sorbate molecule. The vector changes significantly only when sorbates pass from one channel segment to another, due to the rotational constraints imposed by the pore walls. Given that the channels are orthogonal and populated approximately equally, the average value of the correlation function of the end-to-end vector is expected to approach zero at long times. Time constants of the decay of correlation of this vector were estimated to be of the order of hundreds of picoseconds and were found to be as long as 500 ps for hexane at a loading of 4 molecules per unit cell. The suggestion of 1 ns as a simulation length arises from the need to sample at least two relaxation times for interchange between channel segments to gather adequate statistics. The rate of conformational isomerization, an integral part in the process of larger molecules such as

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hexane moving from one channel segment to another, was also shown to have a time constant of the order of 100 ps. Hernandez and Catlow (86) recently reported an investigation of n-butane and of n-hexane diffusion in silicalite; the work is similar to that of June et al. (85). Many calculation details were the same as in the earlier work, including the assumption of identical Lennard-Jones coefficients of intermolecular dispersion and repulsion. Simulations were performed at different loadings for butane, namely, 2, 4, 5.3, and 8 molecules per unit cell. In addition, simulations were performed at a constant loading and variable temperature (200, 300, and 400 K) for both butane and hexane. These calculations were performed for 1000 ps, twice the length of those of June et al. The zeolite framework was held rigid. The effect of these longer simulations is evident from a comparison of the calculated diffusion coefficients for butane as a function of loading. The values up to 4 molecules per unit cell are in good agreement with each other, as are the ratios of the components of the anisotropic diffusion tensor. However, at a loading of 8 molecules per unit cell, the value of Hernandez and Catlow (0.10 X lo-* m2/s) is twice that of June et al. This loading corresponds to the slowest diffusion rate, as sorbates are slowed by frequent intermolecular collisions.The larger diffusion coefficient, corresponding to the longer simulation time, is probably the more reliable. (This suggestion is supported by the recommended length of simulations made by June et al.) The variation of the temperature appears to produce a greater variation in the calculated diffusivities than the loading. For both butane and hexane the diffusivity is predicted to increase by a factor of 4 as the temperature is raised from 200 to 400 K. Butane diffuses somewhat faster than hexane, although the difference is never greater than a factor of 2 at any of the three temperatures considered. The ratio of the unidimensional diffusion coefficients defined by Eq. (5) increases with temperature. At 200 K, both butane and hexane yield values close to the lower limit suggested by Karger (24), whereas at 400 K, the values are 15 and 7 for butane and hexane, respectively. The Arrhenius activation energy barrier to diffusion was calculated to be 4.4 kJ/mol for butane and 5.3 kJ/mol for hexane; it appears to be relatively insensitive to changes in chain length. This insensitivity, and indeed the actual values, are in good agreement with neutron scattering measurements of Jobic et al. (88) for a similar loading (5 kJ/mol). PFGNMR measurements (24) predict a value of 8.1 kJlmol for a sample loaded with 8 molecules per unit cell. The decay of the end-to-end vector correlation function was used to gain insight into the mechanism of diffusion. As in the simulations of June et

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al., the z-component of the end-to-end vector decorrelate very rapidly, because the pore structure prohibits the alignment of molecules parallel to this direction. The z-component of the end-to-end vector is constrained to take on values close to zero, and the rattling of molecules within the pores leads to interchange of small positive and negative values, thus accelerating decorrelation. For the other two components, the decay was found to be far slower, and at 200 K the decay time was found to be of the order of hundreds of picoseconds for butane and thousands of picoseconds for hexane. The rate of decorrelation was found to increase with increasing temperature, but even at 400 K the decay time (estimated by fitting an exponential function to the decorrelation curve) was found to be 235 ps for hexane in the sinusoidal channel. Decorrelation is always faster for molecules in the straight channel, indicating the quicker motion along this section of the pore. The mechanism of diffusion was found to be a series of hops; molecules stay in one energy minimum until sufficient energy is acquired to overcome the potential barrier posed by the intersections. At intermediate loadings, the jump lengths were found to be of the order of 10 A, roughly the distance between channel segments across intersections. This result is in agreement with the average jump length found from neutron scattering experiments (88).At lower loadings, longer jump lengths were predicted by the calculations. This result is probably a consequence of the fixed framework approximation; at lower loadings there are fewer intermolecular collisions, and molecules cannot transfer their energy to the lattice. Supplementary flexible framework calculations corroborate this result; jump distances were found not to exceed 10 A. A rigid lattice may lead to prediction of an unrealistic diffusion mechanism, but it does not preclude good agreement with experimentally measured diffusion coefficients.

5. Benzene and Other Molecules The adsorption and diffusion properties of benzene are of immense interest in zeolite research; aromatics play important roles in a number of zeolite-catalyzed processes. Theoretical simulations of benzene diffusion first began to be published in the late 1980s. The first studies evaluated and minimized the potential energy of a molecule such as benzene within the channels, a method less computationally demanding than the MD simulations that followed. Most recent studies have used the TST formalism. Possibly the earliest theoretical study of diffusion of aromatics in zeolites was published in 1987 by Nowak et al. (89), who considered diffusion of benzene and toluene in the pores of silicalite and theta-1. Theta-1 (90)has a unidimensional medium-sized pore opening bounded by 10-rings. In this study, only the straight channel of silicalite was considered, making the

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pore topologies only slightly different. Potential maps were evaluated for two degrees of freedom of the benzene molecule, translation down the main pore, and rotation about the channel axis. The interatomic potentials were taken from Kiselev et al. (79) and both the zeolite and guest were held rigid during the simulations. Both zeolites were considered as all-silica structures, and contributions from the Si atoms were neglected. Activation energies for diffusion were calculated from differences in energy on the potential maps. The barrier to diffusion of benzene and of toluene in silicalite was estimated to be 22-45 kJ/mol. The diffusion pathway in theta-1 was predicted to be far less hindered. In 1989, an article was published that simulated the diffusion of benzene in silicalite, theta-1, and EU-1 (91). The method of a forced diffusion path through the pores was used and energy minimization calculations of the flexible guest were performed at regular intervals (every 0.1 A, for a distance of 10 A). The intermolecular potentials between the benzene molecule and the rigid framework were taken from the work of Kiselev et al. (79). The three zeolites have different pore structures: theta-1 and silicalite were in length described earlier; EU-1 has large 12-ring side pockets ca. 8 that branch off the main elliptical pore that has dimensions of 5.8 X 4.1 A (80). The energy profiles of the diffusing benzene molecules as a function of distance from the starting point were used to estimate the barrier to diffusion. This was calculated as the difference between the maximum and minimum energies obtained during the course of the calculations. For theta-1, diffusion was predicted to be facile, with the barrier calculated to be only slightly higher than thermal energy at room temperature. In silicalite, the barrier height was estimated to be approximately 25 kJ/mol, which is in good agreement with the activation energy values found from experimental studies (21-28 kJ/mol) (24).In EU-1, there was a possibility of trapping diffusing molecules in the side pockets. The side pockets are large enough to accommodate benzene without obstructing the passage of other molecules through the main pore. From two separate simulations, the authors found that the barrier to diffusion into the side pockets from the straight channel was smaller than that for diffusion along the main channel. Thus, it appears that trapping in the side pockets is likely. The activation energy for diffusion along the main channel was calculated to be 28 kJ/mol. Similar energy minimization calculations were reported for benzene and p-xylene in silicalite (92). Diffusion coefficients were estimated from minimum energy paths through the pore. The value for benzene, 27.6 kJ/mol, is in good agreement with that of Pickett etal. (91). For the bulkierp-xylene molecule, the activation barrier was predicted to be slightly lower (23 kJ/

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SIMON P. BATES AND RUTGER A. VAN SANTEN

mol). Geometrical constraints prevent rotation of the aromatic ring in the channel segments, whereas at the intersections, even xylene can rotate. Schroder and Sauer (93) calculated the minimum energy path of a single benzene molecule through silicalite and applied TST to estimate the diffusivities. The parameters of Kiselev et al. (79) were used to describe the interaction of a rigid benzene molecule with a rigid, cut-out portion of the silicalite lattice. Energetically favored sorption sites were determined by an optimization of all six translational and rotational degrees of freedom of the benzene. A large number of minimum energy structures were found, with the global minimum in the straight channel midway between two intersections. Assuming that a minimum in the straight channel was separated from a minimum in the intersection by a single barrier and neglecting diffusion in the sinusoidal channel, the authors calculated diffusion coefficients. These values were of the order of 10-11-10-13m2/s, somewhat higher than experimental values (24) determined by various methods (approximately m2/s). In fact, diffusion through the sinusoidal channels was found to be slower than diffusion through the straight channels but not negligible, as will be discussed later, and it could partly account for the rather high calculated values. Schroder and Sauer pointed out that a large number of different minimum energy paths exist, and it is rather difficult to follow a particular path exactly. Depending on the algorithm used for following the path, the benzene molecule may be predicted to undergo sudden unphysical reorientation as it switches from one path to another. Demontis et al. (94) reported an early MD study of the sorption and mobility of benzene in zeolite N a y . The zeolite was modeled with a %/A1 ratio of 3.0, as in previous calculations for Xe and methane. The zeolite and benzene molecules were treated as rigid. The authors supported the assumption of a rigid zeolite lattice by quoting structural studies (95), in which the cell parameter of NaY zeolite was found to contract little upon uptake of benzene. It is, however, more than possible that the lattice undergoes substantial deformation without an overall change in volume; quantum chemical calculations (96) have shown that the Si-0-Si bending potential is very soft. When these calculations were performed, the assumption of a rigid lattice was more a matter of computational necessity than it is today. Benzene-benzene interactions were modeled with a Buckingham potential that was shown to yield reasonable predictions of the properties of liquid and solid benzene. Benzene-zeolite interactions were modeled by a short-range Lennard-Jones term and a long-range electrostatic term. In total, 16 benzene molecules were simulated in a unit cell of zeolite Y , corresponding to a concentration of 2 molecules per supercage. Calculations ran for 24 ps (after an initial 24-ps equilibration time) for diffusion at 300 K.

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It was found from experimental studies (95) that there are two preferred adsorption sites for benzene in zeolite Y. The first has the benzene molecule lying atop an SII cation site (center of a 6-ring). The distance from the benzene plane to the Na cation is approximately 2.7 A. The second position for benzene is in the 12-ring window to the supercage. The MD simulations were initiated with benzene exclusively atop SII cations. As the calculation progressed, the benzene molecules migrated to occupy window sites, in agreement with neutron diffraction studies of sorption location as a function of loading. The frequency spectrum, obtained by Fourier transforming the VAF, provided information on the ease of motion of the benzene molecule around the cation site. Two peaks were observed, the lower (20 cm-l) being assigned to motion parallel to the cage surface, the higher (110 cm-') to motion perpendicular to the cage surface. It is clear that the nature of the pore system results in motion that is highly anisotropic. m2/s.Experimental The diffusion coefficient was estimated to be 4 X values for benzene in faujasites range from 10-l' to m2/s, depending on the measurement technique (24, 97). PFG-NMR measurements are the closest to the MD value, which was admitted by the authors to be a crude estimate (mainly on grounds of a short simulation and inflexible molecules). The simulation time was too short to permit a calculation of the residence times of the benzene at either the cation or the window site or inside a particular cage. The cage residence times were estimated to be at least an order of magnitude longer than those for methane in NaY zeolite (43). Bull et al. (97) reported a systematic 2H-NMR and MD study of siliceous faujasite. MD calculations were performed for 1 molecule of benzene adsorbed in a single unit cell of faujasite. Full framework flexibility was incorporated, using potential parameters from MSI's cff91 force field (5). Simulations were performed for diffusion at 298,350,400, and 450 K, using a time step of 1 fs for a 25-ps calculation run (following 5 ps of equilibration). The agreement between the MD and 2H-NMR diffusion coefficients is remarkably good, given the rather short simulation times imposed by the limited computational resources. The values are (2.0 ? 0.3) X m2/s for the MD simulations at 300 K and (4.5 ? 3.3) X lo-' m2/s for the 2HNMR experiments. The activation energy barrier to diffusion was calculated to be approximately 29 kJ/mol, nearly 3 times that derived from the NMR measurements. Molecular trajectories demonstrated the increasing localization of benzene at lower temperatures. At 400 K, the benzene molecule was shown to exhibit cage-to-cage diffusion on the 25-ps time scale of the simulations. Only at this temperature did the calculated diffusion coefficient correspond to intercage motion; all other values describe intracage migration.

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On the basis of their 'H-NMR measurements for siliceous Y and NaY zeolites, the authors estimated the simulation time needed to observe cageto-cage migration for an MD calculation for NaY zeolite. Their estimate of 200 ns is still beyond even the longest MD simulations on the most powerful machines. It is unsurprising, therefore, that more recent theoretical investigations into diffusion of benzene have used the more appropriate TST approach. Klein et al. (98) performed molecular mechanics calculations to investigate the diffusion of a group of aromatic compounds in NaY zeolite (benzene, toluene, the xylenes, mesitylene, aniline, rn-nitroaniline, and rn-dinitrobenzene). In their calculations, a molecular mechanics force field was used to determine the PES of 1 aromatic molecule in a unit cell of zeolite Y. Initially, guest-host energies were calculated on a coarse grid over the asymmetric unit. Possible adsorption sites were further refined by performing calculations at these positions with a very fine grid and 5000 trial orientations of the guest molecule. Passage from one minimum energy sorption site to another characterizes the minimum energy path (MEP) of the guest molecule in the limit of infinite dilution at 0 K. The difference in energy between the initial and final states is a measure of the activation barrier of the motion. The zeolite was modeled with a Si/Al ratio of 3; both the lattice and guest molecules were held rigid and the total interaction energy was calculated as a sum of Lennard-Jones and Coulombic terms. Lennard-Jones parameters were taken from Kiselev and D u (99) and Demontis et al. (94). Partial charges in the Coulombic interaction were taken from Demontis et al. (94)for benzene and from ab initiu calculations. The calculations led to predictions of adsorption sites for the nonpolar compounds that are in good agreement with those determined experimentally. The cation site is preferred over the window site. The activation barrier for movement between two cation sites was calculated to be 30 kJ/ mol and that for movement between a cation and a window site 43 kJ/mol. Experimental measurements of activation barriers to diffusion of benzene in faujasites are between 17 and 27 kJ/mol (24).The calculations provide strong support for the mechanism of surface-mediated diffusion for all guest molecules in the limit of infinite dilution and 0 K. The MEPs show that molecules slide along the wall of the supercage, with the plane of the aromatic ring almost parallel to the pore wall. The calculations of Klein et al. predicted a larger activation barrier for the diffusion of rn-xylene than for diffusion of u- and p-xylene. Thus, rn-xylene may be expected to diffuse more slowly than u- or p-xylene, which is inconsistent with the diffusion coefficients and activation energies determined experimentally (24,100).Of course, one of the main factors that precludes a truly meaningful comparison is that the calculations simulate the

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limit of infinite dilution and 0 K. The real dynamical behavior of the system can differ greatly from the MEP prediction. Auerbach et al. (101) used a variant of the TST model of diffusion to characterize the motion of benzene in NaY zeolite. The computational efficiency of this method, as already discussed for the diffusion of Xe in NaY zeolite (12), means that long-time-scale motions such as intercage jumps can be investigated. Auerbach et al. used a zeolite-hydrocarbon potential energy surface that they recently developed themselves. A Si/A1 ratio of 3.0 was assumed and the potential parameters were fitted to reproduce crystallographic and thermodynamic data for the benzene-NaY zeolite system. The functional form of the potential was similar to all others, including a Lennard-Jones function to describe the short-range interactions and a Coulombic repulsion term calculated by Ewald summation. Four distinct hopping events were considered in the calculations, which correspond to the movement of the benzene molecules between the cation and window sites of minimum energy in Nay, i.e., cation to window (C-W), C-C, W-C, and W-W. The associated rate constants of these processes were used to calculate the activation barrier to each hopping process and the Arrhenius prefactor. The MEP of benzene molecules was followed by a constrained optimization method that drags benzene from its initial site of minimum energy, through the transition state, to the final state. The results show that the motion of the benzene molecules between two cation sites within the same supercage can proceed in two different ways, illustrated in Fig. 4. The first (consistent with the MEP) is cartwheel type motion, whereby the molecule remains roughly orthogonal to intermediate

FIG. 4. Illustration of the cartwheel (left) and skateboard (right) intercage migration between two cation sites of benzene in NaY zeolite. Reprinted with permission from Ref. 101. Copyright 1995 American Chemical Society.

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4-rings. The second, denoted skateboard motion, occurs as the plane of the benzene molecule remains approximately parallel to the intermediate 4-rings. These two pathways characterize intracage motion. To examine intercage motion, the C-W hop was investigated, illustrated in Fig. 5. The MEP of cation-to-window motion also displays cartwheel behavior with an activation energy of 41 kJ/mol. This result is in contrast to the findings of Klein et al., who predicted the MEP for cation-to-window motion proceeds via a skateboard motion (98).The final hop considered, W-W, only becomes important at higher loadings when all the cation sites are occupied. Leaving a window site is more facile than leaving a cation site; this is reflected in the lower activation energies for hopping from a window site compared with those for hopping from a cation site. The process of C-W hopping is considered to be rate-limiting in intercage diffusion. The overall activation energy to diffusion was calculated as 41 kJ/mol, the same as the hopping activation energy between cation and window sites. The reason for this agreement is not simply that this hopping activation energy was the largest observed; this is the only jump that propagates longrange motion. Even if the activation energy for cation-to-cation hops was greater than 41 kJ/mol, the overall activation barrier to diffusion would still be 41 kJ/mol because cation-to-cation hops do not contribute to diffusion. This value is in good agreement with that of Klein et al., but both are overestimations of experimental activation barriers (24). Although the potential parameters used in the simulations may need some refinement, a reasonable comparison has been achieved. The same authors also applied MM calculations to compare the behavior of benzene in NaY and NaX zeolites (102).NaX zeolite was modeled with

FIG.5. Illustration of the motion of benzene in NaY zeolite between a cation and a window sorption site. Reprinted with permission from Ref. 101. Copyright 1995 American Chemical Society.

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a Si/Al ratio of 1; the Na I and I1 sites are completely filled, with 32 remaining Na cations distributed throughout Na 111 sites, generally found in the proximity of the 12-ring window. [The Na 111 cation distribution of Hseu (103) was assumed.] From the calculations, three minimum energy positions were located: the cation and window site described earlier and an additional structure whereby benzene interacts with the Na 111 cation, at a distance of 2.8 A away inside the cage. The binding energies indicate that NaX zeolite provides a homogeneous environment for benzene (in contrast to what was found for NaY zeolite). Stabilization energies were found to be similar to those found for the cation site in NaY zeolite (approximately 70 kJ/mol). Subtle differences were found to exist; for example, in the case of the cation site, the plane of the benzene molecule is no longer predicted to be parallel to the supercage 6-ring. Repulsive interactions with the Na 111 cation distort the plane of the benzene, leading to a destabilization in NaX zeolite of 7 kJ/ mol relative to the value for NaY zeolite. In the case of the window site, favorable interactions with the Na 111 cation mean that the interaction in NaX zeolite is 14 kJ/mol more stable than in NaY zeolite. Thus, more window sites are populated in NaX than in NaY zeolite. This siting information was used to construct pathways by which the benzene molecule hops between neighboring sites. The presence of the Na I11 ions yields seven distinct hopping paths. It was found that hopping between SIII and SII sites and between SIII and W sites was characterized by approximately the same activation energy, 15 kJ/mol. This result suggests that the time scales of inter- and intracage orientational randomization will be similar. In NaY zeolite, hopping activation energies were found such that orientational randomization was controlled by intracage motion. An NMR spin-relaxation measurement will reflect this and will therefore not be expected to agree with a PFG-NMR experiment. Conversely, in the case of NaX zeolite, inter- and intracage motions are expected to contribute to orientational randomization, and so the two NMR experiments would be expected to agree. This expectation was demonstrated by the 2H-NMR results of Burmeister et al. (104) and the PFG-NMR measurements of Germanus et al. (105). Both values (22.5 and 20 kJlmo1, respectively) are in satisfactory agreement with the value of 15 kJ/mol found from the simulations. In addition, the authors presented NMR correlation times that predict a value of the activation energy in NaX zeolite of 14 kJ/mol, in excellent agreement with the calculated value. The predicted enhancement of benzene mobility in NaX zeolite was explained by using a simple model, which accounts for both the enhancement effect and its order of magnitude. The addition of a new sorption site (the SIII site) in NaX zeolite that is of approximately the same strength

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acts to increase the mobility of the benzene molecules. The attractive interactions from the additional SIII sites overlap with those from the SII sites, making them act as “stepping stones” to facilitate increased benzene mobility in NaX zeolite. Snurr et al. (106) performed a TST study of the dynamics of benzene in silicalite. The choice of the molecular sieve and sorbate makes this system quite a challenge. Silicalite is electrically neutral; hence one predicts a greater number of preferred binding sites for the nucleophilic benzene in silicalite than in the cationic faujasites. In addition, benzene has six degrees of freedom, so each minimum energy structure that corresponds to a preferred sorption site represents a local minimum in six-dimensional phase space (assuming a fixed lattice). This expectation was borne out by the location of 27 unique minima and 100 unique transition states (that connect two minima) for benzene in the asymmetric unit of silicalite. The minima were grouped into microstates and macrostates. The benzene molecule was found to move frequently between microstates within a given macrostate; hence each macrostate characterizes the phase space around a local minimum. Transitions between macrostates were found to be approximately 1 order of magnitude less frequent, as a result of the relatively high energy barriers that interconnect them. Microstates within a given macrostate may be assumed to be in equilibrium with each other. Macrostates were found in both channels and in the intersections, where the benzene molecules were predicted to have the highest occupation probability. The rate constants for hops between macrostates were found to be strongly temperature dependent, in contrast to the hops within a macrostate. These latter movements invariably correspond to slight rotations or displacements of the benzene molecule, such as rotation about the c6 axis. The facile nature of the rotations, even at temperatures as low as 100 K, has been indicated by neutron scattering studies (107).Rotational barriers around the c6 axis have been calculated by VignC-Maeder and Jobic (108). Using a potential model that included polarization and electrostatic as well as dispersive and repulsive terms, they found a barrier of 2 kcal/mol for a benzene molecule occupying a minimum energy position in the channel intersection. [They also calculated minimum energy paths for diffusion between potential minima and found an activation barrier to motion along the straight channel of 33 kJ/mol, slightly higher than that found from MD calculations (91).] The predicted diffusion behavior of benzene showed the familiar anisotropy observed for other molecules in silicalite, although evolving over a far longer time period. The calculated diffusion coefficients at temperatures between 200 and 500 K varied by 6 orders of magnitude. The orientationally averaged value at 300 K is 1.1 X m2/s, approximately 1-2 orders of

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magnitude smaller than experimental results (24). Diffusion in the sinusoidal channel is not negligible; at 300 K it was found to be a factor of 3 slower than motion through the straight channel. The activation barrier to diffusion was found to be 37 kJ/mol and the preexponential factor was found to be 2.54 X lO-'O m2/s.The activation energy appears to be overestimated by the calculations and the preexponential term underestimated. The authors suggested that the principal source of error in their calculations arose from the rigid framework assumption. This suggestion was supported by citation of calculations of diffusion through polymers (109).When the polymer network was given flexibility, diffusivities within a factor of 2 of experimental values were predicted, whereas rigid network calculations predicted underestimations of 3-5 orders of magnitude. Finally, for completeness we briefly mention a TST study of SF6 in silicalite (12);the calculation is complementary to that for Xe diffusing through the same material. Limited experimental results exist for comparison; Karger and Ruthven (24) cited a diffusion coefficient from PFG-NMR measurements of 0.4 X lo-' m2/s for NaX zeolite at 300 K with a loading of 1.2 molecules per cavity. The TST-calculated diffusion coefficients (in the limit of zero coverage at 300 K) differ by at least an order of magnitude from the experimental value. Notwithstanding the different frameworks and the impossibility of a direct comparison, it is possible that there are inaccuracies in the SF6 model potential and also that the assumption of a rigid lattice is seriously flawed for a sorbate molecule as large as SF6. D. SUMMARY It is still not easy to draw general conclusions regarding the assumptions made in the simulation of diffusion in zeolite pores. Different simulations of the same system have been shown in the preceding pages to lead to different conclusions. The calculation requirements to ensure a good comparison between simulations and true diffusion are still very much dependent on the system under consideration, although researchers find themselves in a far better position these days than 5 years ago. Computational advances have progressed so rapidly that a great many of the assumptions that were introduced by necessity (on the grounds of limited resources) at the beginning of the decade may now simply be avoided. Flexible framework calculations are the norm and tried and tested potential parameter sets are now available. Furthermore, software has been developed to the point that many of the seminal works of the late eighties are now short, routine bench tests. It is nonetheless instructive to have a sense for which calculation parameters are most likely to affect the quality of the predicted results. We suggest

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that these are the simulation length, the size of system considered, the force field parameters, and the inclusion of framework flexibility. The first of these applies to all calculation systems, irrespective of their nature. To obtain reliable estimates of long-range diffusive behavior, MD calculations must be allowed to run for a sufficiently long time. As sorbates become more complex, these simulation times progress from the order of picoseconds to the order of nanoseconds. Modeling very slow diffusive processes may require the application of more computationally expedient methods. The reliability with which diffusive behavior is predicted has been shown to be affected by the size of the simulation box and the number of adsorbate molecules, as the diffusive behavior of any two adsorbates may be radically different. When new force fields are applied to simulate diffusive behavior, it is essential that they already be thoroughly tested and evaluated for their ability to reproduce experimental spectroscopic results. The flexibility of the framework is of maximum importance when the sorbate and pore window sizes become commensurate. The outlook for theoretical simulations of diffusion in zeolites is certainly encouraging. Order of magnitude agreement between experimental and theoretical results constituted success at the inception of these calculations, but now methods have progressed and parameters refined to the point whereby theoretical methods can rival experimental methods in accuracy and cost-effectiveness.

111.

Adsorption in Zeolites

A. SCOPEOF THISSECTION

The study of sorption of guests within zeolite hosts is complementary to the study of diffusion in zeolites. Having discussed the pathways and trajectories of molecules through micropores, we now consider the favored sorption locations, conformations of sorbates, and sorption energetics. Indeed, so close are the two subjects that they are frequently considered within the same paper. It is therefore unsurprising that the MD and TST methods used to characterize diffusion processes are also used to simulate sorption. In the theoretical methodologies section that follows, these methods are not mentioned further as they were summarized in the preceding section. Monte Carlo methods are discussed in detail, including a recently developed technique to simulate the location and adsorption of longer chain molecules than would normally be possible by using conventional methods. Furthermore, we present the methodology of a combined MD/Monte Carlo/EM tech-

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nique, as implemented in the MSI molecular modeling software. We then proceed with a discussion of the results of the sorption studies, ordered according to sorbate molecule. METHODOLOGIES B. THEORETICAL 1. Monte Carlo Simulations One of the main limitations of using MD methods to investigate the sorption energetics and locations of molecules within zeolite pores becomes evident when the trajectory of the sorbate molecule evolves rather slowly with time. This can occur when the sorbate is large or the trajectory passes through a bottleneck in the structure, such as a narrow pore window. For a given simulation length, the guest species might not visit all relevant areas of the host pore structure. Monte Carlo methods avoid this inherent difficulty by not following the natural path of the sorbate through the host; instead, the guest molecule is moved to random points within the host. At low temperatures, the location of a molecule within a zeolite is confined to the most favored sorption site. A t higher temperatures, location becomes a probability distribution and energy an integral over many sites. The numerical evaluation of the integrals is hindered by the fact that the total potential energy is often highly repulsive (e.g., when a guest molecule occupies a site in the zeolite close to or overlapping with the walls). These configurations contribute little to the total energy, and so to overcome this a method of importance sampling was developed by Metropolis et al. (110). The method of importance sampling confines the exploration of configurational space to regions of significant probability. In general, a particle is selected and displaced in a particular direction. In the case of a molecule, there is the possibility of displacement and rotation of the molecule about a fixed axis. The direction and degree of movement are selected at random. is accepted if it is more favorable The energy of the new trial state, Etrial, or if than the previous, Einitial,

E trial . -E. .. initial exp(-

kT

< rand(O,l),

(7)

where rand(0,l) is a random number between 0 and 1, T is the absolute temperature, and k is Boltzmann’s constant. States are generated that sample the equilibrium states of the system according to a Boltzmann probability distribution. At low temperatures, the probability of accepting highenergy states is small, and at high temperatures it is correspondingly higher. The determination of equilibrium properties reduces to a numerical average over the states sampled.

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Monte Carlo simulations are performed within a statistical ensemble. In the canonical ensemble (with the number of molecules, volume, and temperature fixed), the average value of a thermodynamic quantity, ( T ( x ) ) , as a function of the states of system, x, is given by

where P ( x ) is the Boltzmann-weighted probability. In the grand canonical ensemble (pressure, volume, and temperature constant), the number of molecules is varied so as to preserve the stipulated values of constant parameters. This type of simulation allows prediction of the equilibrium between two phases, such as a sorbate and zeolite. In addition to the displacement steps in the simulation, there is a (random) choice between an attempted deletion or creation of a particle. Separate calculations at different pressures can be collected to predict the adsorption isotherms of guest species in zeolite hosts. The parameter determined from simulations that can be most readily compared to an experimentally observable quantity is the heat of sorption. The mean total internal energy determined by the simulations, (U), can be equated to the isosteric heat of adsorption, Q,, in the limit of low occupancy, as follows:

where R is the gas constant and T the absolute temperature. Both isosteric heats of adsorption and internal energies are often quoted in research papers. 2.

Configurational-Bias Monte Curlo Methods

Conventional Monte Carlo methods pose their own set of problems as the guest molecule increases in size. The trial movements of the guest in the zeolite become increasingly unlikely to satisfy the acceptance criteria, as there is a far higher probability that part of the guest molecule will overlap unfavorably with the zeolite wall. For a typical zeolite structure, the probability of successfully moving a methane molecule to a location in the pore system where none of the atoms are in contact with the pore wall is approximately (111). For ethane, this probability is only and for longer alkanes it becomes vanishingly small, resulting in acceptance of almost none of the trial moves. The configurational-bias Monte Carlo method (CB-MC) (112)was developed to overcome these sorts of problems. Instead of a random insertion into the zeolite host, the guest molecule is “grown” atom by atom within the host in a way that avoids unfavorable overlap with the zeolite atoms.

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This growth procedure introduces a bias in the conformation of the molecule which can be exactly removed by adjusting the acceptance rules (111,113). These types of simulation are particularly well suited to chain molecules and the results we discuss in the following section are all for normal alkanes. Simulations are performed in cycles, each of which consists of a randomly selected attempt to either translate, rotate, regrow part, or regrow ail of the molecule. The latter two choices exploit the efficiency of the CB-MC method. In addition, the process of completely regrowing a chain inside the zeolite leads to information with which to calculate the Henry’s law coefficient, which may be compared with experimentally measured values. At regular intervals throughout the calculations, snapshots of the energetics and conformation of the sorbed molecules may be taken for postcalculation analysis.

3. Combined Monte Carlo/Energy Minimization Methods This combined method consists of three basic calculation steps. The first involves a short, high-temperature MD simulation of the guest in the gas phase. Snapshots of the guest are saved at regular intervals during this sampling of flexibility. The second stage involves Monte Carlo docking of these random MD snapshots inside the host structure. The MC method generates successive docked structures entirely at random, in contrast to the Metropolis importance sampling method. The docked structures are selected for the final stage of the calculation if the total energy lies below a threshold value. Structures that are selected are subject to energy minimization of first the guest, and, if desired, a portion of the host lattice as well. The energy minimization step encourages relaxation of the system to the nearest minimum energy site, and all of these minimum energy sites within the whole lattice may be found if the number of randomly docked structures that are selected is statistically large enough. C. SURVEY OF RESULTS

1. Single Atoms The use of statistical calculations of configuration integrals to determine thermodynamic adsorption characteristics of zeolites dates back to the late 1970s (49).Kiselev and Du (22) reported calculations based on atom-atom potentials for Ar, Kr, and Xe sorbed in NaX, Nay, and KX zeolites. Their calculations, which included an electrostatic contribution, predicted changes in internal energy in excellent agreement with those determined experimentally. The largest deviation between calculated and experimental values, for any of the sorbates in any of the hosts, was a little over 1 kJ/mol.

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In Faujnsites. The sorption of single-atom adsorbates in the cavities of zeolites has been reported by Woods et al. (114, 115), first, for model Lennard-Jones particles in model spherical cavities and then Xe and methane in the supercages of faujasite. In the latter case, the potential parameters of Kiselev and D u (22) were used, parameterized to fit experimental isosteric heats of adsorption at zero coverage. Grand canonical Monte Carlo calculations (GC-MC) were performed, and the calculated adsorption isotherms were in good agreement with experimental results. The distribution of cage occupancies of Xe in NaY zeolite was found to be broad and roughly symmetrical about the most probable occupancy, except at very low loadings. For an average loading of 2 atoms per cage, as many as 6 atoms were found in any given cage. The Xe atoms were found to occupy six minimum energy sites within the cage in the limit of low occupancy. At higher loadings it was predicted that the center of the cavity would become populated. Calculations for the X and Y forms of faujasite demonstrated the “holes” in the structure of the adsorbed fluid induced by the cations. The isosteric heat of adsorption was predicted to increase steadily as a function of loading, up to approximately 60 atoms per unit cell, at which value a sharp decrease was observed. The initial increase was attributed to attractive interactions between adsorbed particles, and the decrease to the supercages becoming full. The heat of adsorption in the limit of zero coverage shows a decrease with increasing temperature. [This is in contrast to the results obtained from calculations on model spherical cavities (114).] A great deal of information about the location of single-atom adsorbates such as Xe has been obtained from M D studies that were primarily concerned with the diffusion of these species (13, 15, 17, 18, 20, 28, 29). In general, the sorbates are localized at low temperatures in a minimum energy location close to the inner surface of the zeolite pore. A t higher temperatures, the delocalization of the sorbates increases, and regions of the void volume further away from the walls become populated. The structural similarities between zeolites Y and A mean that this description of sorbate location applies to Ar in NaCaA zeolite (20, 28, 29). The location of the Xe in NaY zeolite has been explained in terms of the potential energy surface inside the supercage (13).The minimum energy location close to the cage wall corresponds to the global potential energy minimum, and the cage center represents the global maximum. The isosteric heat of adsorption of Xe in NaY zeolite at 376 K was calculated to be -15.5 kJ/mol (14) from MD simulations. The majority of the discrepancy between this and experimental values (approximately 20 kJ/mol) (22) arises as a result of the omission of the polarization interaction. The interaction distribution function is essentially unimodal with a peak at - 12 kJ/mol and a small shoulder around 15 kJ/mol. The shoulder arises

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55

from Xe atoms localized near the adsorption site. Thus, a spherically averaged potential would seem to be a reasonable assumption for the adsorption of Xe in NaY zeolite, as there do not seem to be two or more different energetic regions within the cage. (This is in contrast to the situation for methane and benzene, as discussed shortly.) The distribution of occupancies demonstrates that there may be as many as five Xe atoms in a cage at any given time, starting from an initial loading of 1 atom per cage. For Xe concentrations greater than 1 Xe per cage, the position of minimum energy close to the cage walls is less populated and the Xe atoms adsorb closer to the cage center. There is evidence for dimerization from the radial distribution function (RDF) of the Xe-Xe interactions for even 1Xe per cage, but this dimer population decreases with increasing temperature. The dimerized Xe atoms were found to reside approximately 4.6 A apart (again from the RDF), and it was estimated from the intensity of the peaks in the Xe-Xe energy distribution function that 15% of the Xe atoms are paired at any one time when the average concentration is 1 Xe per cage. There is evidence for higher order clusters, even for this lowest concentration of Xe atoms, in agreement with the findings of Demontis et al. for methane (23). When there are 2 Xe atoms per supercage on average, there are few isolated atoms and when there are 3 Xe atoms per cage on average, most exist as dimers or trimers. Hence, a progression toward liquid behavior is observed. From a calculation of the power spectrum (the inverse Fourier transform of the VAF), the Xe-Xe dimer stretching frequency was estimated to be 16 cm-'. In Silicalite. Different values of the change in internal energy upon Xe adsorption in silicalite have been calculated from MD simulations (10, 11). Pickett et al. (10) quoted a value of -27.4 kJ/mol at 250 K and 4 Xe atoms per unit cell. This is slightly higher than the value of June et al. (11) of -21.5 kJ/mol under similar conditions. Experimental uptake rate measurements give an intermediate value of -24.5 kJ/mol. A decrease in the internal energy of adsorption is predicted as the loading is increased because of additional interatomic interactions. June et al. (11)found that increasing the temperature led to increased internal energy of adsorption. From the single-particle density distribution function, the distribution of Xe over the silicalite cell was determined. At infinite dilution, sorbates sample a large fraction of the essentially energetically homogeneous pore structure. At a concentration of 16 atoms per unit cell, there is clear evidence for structure in the sorbate phase as the sorbates confine themselves to smaller volume regions of the channels. The Xe atoms prefer the sinusoidal channels and avoid the unfavorable intersections. This result is supported by the findings of a TST study by the same authors (12). The occupation probability of a given state was found by normalization of the configuration integral by the

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sum of those for all accessible states. This calculation yielded occupation probabilities at 150 K of 0.57, 0.41, and 0.01 for the sinusoidal, straight, and intersection regions. At higher temperatures, the occupancy probability is governed by pore region volumes as the zeolite adsorption sites become more energetically homogeneous. Hope et al. (116)presented a combined volumetric sorption and theoretical study of the sorption of Kr in silicalite. The theoretical calculation was based on a potential model related to that of Sanders er al. (117), which includes electrostatic terms and a simple bond-bending formalism for the portion of the framework (120 atoms) that is allowed to relax during the simulations. In contrast to the potential developed by Sanders et al., these calculations employed hard, unpolarizable oxygen ions. Polarizability was, however, included in the description of the Kr atoms. Intermolecular potential terms accounting for the interaction of Kr atoms with the zeolite oxygen atoms were derived from fitting experimental results characterizing the interatomic potentials of rare gas mixtures. In contrast to the situation for hydrocarbons, there are few direct empirical data to aid parameterization, but the use of Ne-Kr potentials is reasonable, because Ne is isoelectronic with 02-. Three distinct sorption locations were identified by both experimental and theoretical investigations. The most favorable site is located in the intersection, the least favorable site is located in the pore opening of the straight channel, and an intermediate site is located in the pore opening of the sinusoidal channel. The internal energy changes upon sorption calculated from the simulations are 70% larger than those found experimentally; the calculated and experimentally observed energy differences between sites are nearly comparable to each other. The deviation in absolute values is principally a consequence of the uncertainty in the calculation of the K r - 0 interatomic potential parameters and the fact that the calculation method (EM) is based on the assumption of a temperature of 0 K. The effect of sorbate-sorbate interactions was examined in a separate calculation with two Kr atoms, represented by shell models. The sorbate-sorbate interaction energy was estimated to be 1.5 kJ/mol, with the two Kr atoms separated by a distance of 4.3 This distance is similar to the interatomic separation of Xe atoms in NaY zeolite (15, 17). In Zeolite A. An extensive series of papers concerned with the sorption location and isotherms of Xe in zeolite A have been published (118-122). The locations of sorbates and their structures were investigated by using Metropolis Monte Carlo simulations of zeolite A models (118, 119). Initially, an idealized truncated cuboctahedron was used, with Si and A1 atoms occupying vertices and 0 atoms occupying the midpoints of line segments (118). Subsequent calculations were based on the positions of atoms in

A.

MOLECULAR BASIS OF ZEOLITE CATALYSIS

57

an a-cage determined by X-ray diffraction (119). The potential function included terms to account for repulsive, dispersive, polarization, and sorbate-sorbate interactions. Two Si/AI ratios were considered, 2.0 (termed the “cation-poor’’ model), in which eight cations occupy the centers of the six rings of the a-cage, and 1.18 (termed the “cation-rich’’ model), where six additional cations occupy the 8-ring windows. In all simulations, a “hardwall” potential was used in the plane of the 8-ring windows to simulate repulsion of Xe atoms in neighboring cages. The Xe atoms were not predicted to traverse the cage freely during the simulations; instead, specific adsorption sites were located, as found for other zeolites. At these points, the dispersion interaction with both the cation and the framework is maximized. In the case of the cation-rich cage, the sites that are occupied as the loading is increased from 1 to 12 atoms per cage are at vertices in front of the four-membered rings. The sites constitute the vertices of a cuboctahedron. Although the first eight of these sites are found to be energetically equivalent, as indicated by a level region in the differential potential energy graph, they are not entropically equivalent. The nearest-neighbor separation of these sites is 3.9 approximately 0.5 A less than the van der Waals diameter of Xe. Thus, occupancy of one site will affect the probability of occupancy of a neighboring site. At loadings higher than 12 atoms per cage, the cage center becomes occupied, and beyond that, a highly distorted arrangement of sorbates is predicted. Insertion of a 14th Xe atom is just energetically favorable. In the case of the cation-poor model, the structure of the model resembles that of CaA zeolite, though in the simulations the cations are given a +1 charge. Thus only changes due to structure may be detected rather than polarization. The potential surface indicates that a Xe atom inserted into the cage would have a choice of six potential minima, arranged at the vertices of an octahedron. These sites lie in front of the six 8-ring windows and have been predicted to be favorable by other workers (123).The cationpoor cage is less crowded, and these six adsorption minima consequently are independent of each other. Under these conditions, a Langmuir model of adsorption is justified. As the loading is increased beyond 6 atoms per cage, the octahedral sites are altered and new sites are added. At a maximum loading of 19 atoms per cage, all octahedral, cuboctahedral, and central positions are occupied. A decreased temperature of simulation results in adsorbates that are more tightly localized around the minimum energy positions, as illustrated in Fig. 6. At low loadings, the heats of adsorption in the cation-poor cage are greater than those in the cation-rich cage. This result seems surprising at first, since the cation-rich cage is more polar, but this energy benefit is offset by the limitation of adsorption to less favorable sites.

A,

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SIMON P. BATES AND RUTGER A. VAN SANTEN

FIG.^. Densityprofilesalongaplanebisecting the cation-poorwcageofzeolite Aat aloading of 6 Xe atoms per cage. Increased localization about the minimum energy positions is observed at lower temperatures: left, 300 K; right, 100 K (from 119, with permission of Taylor & Francis).

GC-MC calculations have also been reported for the Xe/zeolite A system (120-122). Complementary to their canonical calculations, van Tassel et al. (120) used the cation-rich and cation-poor models just described, together with a “cation-saturated’’ model with a Si/Al ratio of 1.0. In this last model, an additional type I11 cation was placed in front of a 4-ring. In addition to calculations for a single a-cage, a network of 8 (2 X 2 X 2) interconnected cages was also investigated. The entropy for each loading was calculated by extraction from the grand canonical ensemble average entropy as a function of chemical potential. The simulated isotherms represent loadings up to saturation, which require pressures of up to lo4 atm. They are in good agreement with experiment. The cation-saturated simulation predicts an isotherm in good agreement with that of Chmelka et al. (124) for NaA zeolite. In addition, the distribution of cage occupancies is very similar, with deviations only found at extremely high pressures. The cation-poor isotherm is similar to those of Barrer et al. (125) and Jameson et al. (126) for CaA zeolite. The plateaus that exist in the simulated isotherms are in accord with the findings of the canonical ensemble calculations reported earlier. The pressure required to admit more than 6 Xe atoms per cage into the cation-saturated model was found to be prohibitively high; similarly, such a pressure was found for a loading greater than 8 Xe atoms per cage in the cation-poor model. The calculated entropy decreased smoothly as a function of loading, differing little from that calculated from the canonical simulations. In the cation-saturated model, the type I11 cation was found to be located (approximately) in one of the 12 sorption minima. This result was confirmed by

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the decrease in cage capacity from 13 to 12. As the cage is loaded with progressively more Xe atoms, they locate near the type 111 cation. Thus, the presence of the type 111 cation removes the energetic homogeneity of the adsorption sites. Jameson et al. (121, 122) presented a detailed GC-MC study of the locations and structure of Xe in NaA zeolite. Their model of NaA zeolite corresponds to the cation-saturated model of Van Tassel et al., i.e., with a single Na type 111 cation located above a 4-ring. The Xe-Xe potential was represented by a Maitland-Smith function, rather than a 6-12 LennardJones function, to enable accurate prediction of 129Xechemical shifts. The interaction parameters for Xe with zeolite 0 and Na atoms were taken from Woods and Rowlinson (215).Without explicitly considering the structure of the adsorbed Xe atoms, the authors found the cage occupancy distributions to agree with those determined experimently at all loadings up to 8 per cage. Eleven minimum energy sites were found for Xe sorption, in agreement with the results for the cation-saturated model of van Tassel er al. (120). These were not energetically equivalent, because of the presence of the type 111 cation, but instead were spread over an energy range of 0.2 kJ/ mol. Furthermore, the lowest energy transition states that lay on paths between adjacent minima were also located, by a method of “walking uphill” from a potential energy minimum (127).These pathways were found to run along the walls of the a-cage, in agreement with the work of other authors. Most lowest energy transition states were only a few kJ/mol higher than the minima they bridged. By contrast, the barrier to location of a Xe atom at the cage center was calculated to be nearly 14 kJ/mol. The distribution of Xe atoms over the minimum energy sites was found to be similar for loadings up to 8 atoms per cage, with higher values producing Xe atoms more tightly localized near the minimum energy site. The exact positions of the Xe atoms were found to depend on the loading. For example, two Xe atoms were found to adsorb slightly farther apart than the distance between minimum energy sites. A t higher loadings, Xe-Xe separations were found that are slightly lower than the distance between next-nearest-neighbor sites of minimum energy. The exact structure of the sorbed molecules appears to reflect a balance between the stabilization achieved with the framework and the favorable or unfavorable Xe-Xe interactions. More than one unique distribution of sorbates over the minimum energy sites was found. In the case of 1 Xe per cage, there were 11 different distributions over each of the 11minimum energy sites. For higher loadings, up to as many as 25 different sorbate configurations were found. Each of these involves a slightly different arrangement of the Xe atoms around the minimum energy sites, and thus a slightly different energy. The

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range of energies for different configurations is typically 1-2 kJ/mol. This picture conflicts with that of an ordered adsorbed fluid in the cages at higher loadings, described earlier. van Tassel et al. (128) also presented a GC-MC study of adsorption mixtures within zeolite A, using the GC-MC method with the cation-rich and cation-poor models. Initial simulations to predict single-component isotherms showed a clear correlation with adsorbate size and polarizability. Larger adsorbates (such as Xe or methane) were found to occupy welldefined polyhedral sites, whereas smaller ones (such as Ar) were found to fill the pore in a less regular fashion. The isotherms of the latter sorbates exhibit no obvious plateau regions. The isotherms of binary mixtures are somewhat more complicated. In the case of the Xe/Ar mixture, Xe, being larger and more polarizable, adsorbs preferentially at low pressures. At higher pressures, the smaller, less-polarizable Ar is able to displace the larger molecule. Xe is restricted to adsorption sites near the four-membered rings, whereas Ar is able to squeeze into crevices and thus more efficiently fill the pore volume. This type of filling is entropically favored as the Ar atoms have more freedom than Xe. The packing advantage of Ar at higher pressures is therefore a function of both its smaller size and weaker energetic interaction with the pore. A binary mixture of methane and Xe does not show the same packing behavior. Methane is closer in size to Xe, and so both molecules compete for the same sites. Xe dominates at low pressures, because of a greater stabilization by the pore walls. Methane, however, does adsorb preferentially at high pressures, as it can saturate the sites without encountering unfavorable sorbate-sorbate repulsions. The form of the isotherms of the mixtures is largely independent of the cation distribution within the cage, i.e., whether the cation-poor or cationrich model is used. This result is somewhat surprising, especially in view of the different adsorbate structures predicted by single-component isotherms (118-120). Only nonpolar adsorbates were considered in this study and the insensitivity to cation arrangement may well change if one component possesses a permanent dipole. These simulations were based on simple spherical molecules, but the competition for pore space as it depends on size, shape, and polarizability may be extended to other adsorbates. Indeed, Santilli et al. (129) observed experimentally that a branched hydrocarbon adsorbs in preference to a linear one at low loading. Smaller single-component, single-atom adsorbates, such as Ar, in zeolite A have also been investigated (20, 28, 29, 130). Kono and Takasaka (130) calculated the sorption characteristics using classical statistical mechanics methods. The particular MC method that they used is applicable at sorbate concentrations higher than infinite dilution. They found that a London

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model of dispersive forces and a point charge model to account for the electrostatics gave a good approximation of the interaction potential. The adsorption isotherm and isosteric heat of adsorption were found to be in good agreement with experiment, but the coincidence is explained by an overestimation of the polarization energy by the point charge distribution. Santikary and Yashonath (20,28,29) used MD simulations to investigate the sorption of Ar in NaCaA zeolite. For a simulation temperature of 393 K, the guest-host interaction energy was found to be -12.4 kJfmol. The sorbate-zeolite RDFs show well-defined peaks, indicating that the sorbate has a preference for certain regions of the cage. The most intense interaction appears to be between Ar atoms and the extraframework calcium ions. The guest-host interaction energy distribution shows a broad band between -15 and -10 kJ/mol and a small shoulder around -18 kJ/ mol. The latter is indicative of the Ar atoms in minimum energy positions within the framework, and the former broad peak is indicative of more mobile guest molecules. The contrast to the essentially unimodal distribution of energies for Xe in NaY zeolite is most probably a consequence of a difference in framework structure or cation distribution. The guest-guest dimerization energy distribution function shows that there is a small Ar-Ar dimer population with an interaction strength of approximately 1 kJ/mol. Hardly any trimers or higher order clusters were found. In Zeolite rho. Vernov et al. (131,132)reported MD studies of Xe sorption within the a-cages of zeolite rho. Two different models were used for zeolite rho: an all-silica model and one with 0.2 Cs' atoms per unit cell, located in the double 8-ring windows that interconnect the a-cages. Xe-0 parameters were derived by fitting to the experimental Henry's law coefficients, resulting in a potential maximum at the center of the cage, a local minimum close to the walls, and the most favorable sorption site located in the double 8-ring windows. With these potential parameters, it is not possible for the Xe atom to diffuse through a 6-ring window. The Cs'-Xe pair potential was estimated from the Kirkwood-Muller approach (6, 7).Concentrations of up to 9 Xe atoms per cage were investigated. The presence of the Cs cation was found to have little effect other than to reduce the available packing volume for Xe. Xe was found to be initially adsorbed at window sites with an interaction energy of 6.5 kcal/mol. As the loading increased, the interior wall sites became occupied (approximately 4 kcal/mol). At the highest loading of 9 Xe atoms per cage, the cage center was occupied by one Xe atom (1.4 kcal/mol). The intermolecular repulsions between neighboring Xe atoms meant that they were slightly farther apart than their equilibrium minimum energy separation. The maximum loading is in agreement with that found experimentally (133). It was found that the Xe atoms presented a rather

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ordered body-centered cubic structure, executing primarily vibrational motions about their equilibrium positions. Subsequent calculations employing smaller sorbates (Ar and Kr) indicated a similar ordered cubic structure at high loadings (132).The maximum loading of Kr was 15 atoms per cage, and that for Ar was 21 per cage. 2. Water The calculations of Leherte et ul. (38-40), characterizing the sorption and diffusion of water within an aluminosilicate ferrierite, provide energetic and distribution information for the sorbed water molecules. The heat of adsorption was calculated to be 64.71 kJ/mol for a loading of approximately 6 molecules per unit cell. This value increased slightly to 65.25 kJ/mol for a loading of 10 molecules per unit cell, with both values being slightly higher than the unpublished experimental result quoted by Leherte et al. of 59.9 kJ/mol. The presence of aluminum, and hence charge-balancing protons, has a significant effect in these calculations on the distribution and orientation of the sorbate molecules. The water molecules were found to avoid central positions within the larger 10-ring pore; in the smaller 8-ring pore, they occupy both sides of the aluminum sites. The R D F between water and framework atoms demonstrates a preferential orientation of the hydrogen atoms of water toward the pore center. A t a concentration of 4 molecules per unit cell, 8- and 10-ring channels were found to be equally occupied; at 10 molecules per unit cell, a slight preference for the 10-ring channel was found (as it is larger). 3. Methane In Faujusites. Bezus et al. (49) reported in 1978 statistical calculations on the low-coverage adsorption thermodynamics of methane in NaX zeolite (Si/Al = 1.48). As for single-atom adsorbates described earlier, the agreement between their calculated values and a range of experimental values was excellent. Allowing for different orientations of the molecule, they calculated a value of 17.9 kJ/mol for the isosteric heat of adsorption at 323 K. Experimental values available for comparison at that time (134-136) ranged from 17.6 to 18.8 kJ/mol. Treating the methane molecule as a hardsphere particle, with a radius of 2 A, resulted in a far lower heat of adsorption (12.6 kJ/mol). Further calculations (99) yielded heats of adsorption of 19.8 and 18.1 kJ/mol for methane in NaX and NaY zeolites, respectively. Yashonath et al. (46)used a Metropolis Monte Carlo method to simulate the infinite-dilution adsorption of methane in NaY zeolite. The lattice had a %/A1 ratio of 3.0 and was treated as rigid, whereas methane was modeled

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with six degrees of freedom. The interactions between the framework and methane included a short-range Lennard-Jones term and a long-range Coulombic term. The charges for the latter sum were derived from MNDO calculations. Isosteric heats of adsorption were calculated between 10 and 298 K. Results were in close agreement with experimental and previous calculated values. Energy minimization with respect to the position and orientation of the methane molecule led to a site near the central fourmembered ring containing the hexagonal rings of the supercage (Fig. 7). At this site, the methane molecule is approximately 5 A from two SII cation sites and has an energy of -22.1 kJ/mol. There are six of these sites related by symmetry in one unit cell. At low temperatures, the methane was found to be localized at or close to this minimum energy position. At 175 K, the methane molecule was found to have gained enough energy to hop from one minimum energy site to another, and the high-energy shoulder in the energy distribution function represents the regions occupied during the

FIG.7. Predicted minimum energy sites of methane in faujasite from Monte Carlo calculations. The methane molecules are shown as large dark circles within the faujasite cavity. Reprinted with permission from Nature, Ref. 46. Copyright 1988 Macmillan Magazines Limited.

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process of hopping. The energy distribution function at 200 K is clearly bimodal and, at 298 K, the methane was found to be so spread out over the cage volume as to hardly spend any time in the minimum energy position. These findings were corroborated by subsequent, more detailed, MD calculations that focus on the diffusive behavior of methane in NaY zeolite (31, 44). These predictions are in excellent agreement with results of neutron diffraction and infrared studies on methane in NaCaA zeolite (137-1 39). The location of the adsorbate close to the supercage wall in faujasite frameworks has been found to be similar for methane and single-atom adsorbates such as Xe. Calculation of the adsorption isotherms of these adsorbates helps to reveal differences in adsorbate structure as a function of loading. GC-MC simulations by Woods and Rowlinson (115) of 1 methane per supercage of NaY zeolite predicted adsorption isotherms at a temperature of 198 K that were in good agreement with experimental results over a pressure range of lo4 atm (140). At 87 K, adsorption takes place at extremely low pressures, with a marked step in the isotherm occurring around an occupancy of 48 molecules per cell, when all six of the minimum energy positions in each supercage become filled. At low average occupancies and temperatures around 300 K, the structure and location of the adsorbed methane molecules were found to be similar to that described earlier. Differences emerge at higher occupancies: additional methane molecules were found in the windows of the supercages. This additional packing volume is available because the methane molecules can interact favorably with each other, thus increasing the sorption capacity, notwithstanding a hard-sphere radius similar to that of Xe. At even the highest loading, sites at the center of the supercage were not predicted to be occupied. Isosteric heats of adsorption as a function of occupancy at 300 K were found to closely resemble experimental results (140).As in the case of Xe, the heat of adsorption at zero coverage was found to decrease with increasing temperature. This concentration and temperature dependence of the calculated heat of adsorption was also predicted by June et al. (11) (i.e., decreased heat of adsorption as the temperature is increased and the loading decreased). Their value of the isosteric heat of adsorption computed in the limit of infinite dilution at 300 K (17.5 kJ/mol) is slightly lower than the range of experimentally measured values of 18-20 kJ/mol (134-136, 140-1 42). In Zeolite A. The location and distribution of methane within the acages of zeolite A were found to be very similar to those found in zeolite Y, consistent with the structural similarities of these two materials. Cohen de Lara et al. (50) complemented their extensive experimental work with MD calculations of a single methane molecule adsorbed in zeolite A. At

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high temperatures, the methane molecule was once again found to be distributed over the whole cage volume. As the temperature decreased, the cage center became empty and the molecule moved close to the cage wall in the region of an SIII cation. As the temperature decreased still further, the molecule became trapped between SIII and SII cation sites. The calculated heat of adsorption (determined from the average potential energy when the methane molecule is delocalized) was calculated to be 16.7 kJ/mol, in satisfactory agreement with the experimental value of 20 kJ/mol(143). At very low temperatures, three well-defined vibrations were predicted which correspond to the external vibrations in the adsorption site. The calculated values are somewhat too high in comparison with the experimental studies, but this is most probably a result of the assignment of ionic charge to the lattice atoms. The three bands move to lower frequencies as the temperature is increased and eventually flatten out. Only at relatively high temperatures does a new frequency appear, and this is attributed to the free rotation of the delocalized molecule in the cavity. In Mordenite. Smit and den Ouden (60, 144) reported a Monte Carlo investigation of methane adsorbed in mordenite of varying Si/A1 ratios. In their calculations, both the zeolite and sorbate were held rigid, infinite dilution was assumed, and sorbate-zeolite interaction parameters were taken from Kiselev et al. (79). Electronic neutrality of the zeolite framework was preserved by compensating the trivalent aluminum exactly with sodium cations, located in experimentally determined crystallographic locations. The value of the heat of adsorption of methane in mordenite at an %/A1 ratio of 11 was calculated to be 21.4 -+ 2.6 kJ/mol. An experimental value of 23.0 kJ/mol was quoted by the authors, demonstrating the quality of the simulations. If the heat of adsorption is calculated as a function of Si/AI ratio, a marked decrease is observed (approximately 30%) around %/A1 = 6-7. This result is explained by considering the location of the sorbed molecules as a function of Si/Al ratio. In the case of WAl = 11, a significant proportion of the methane molecules were found to reside in the side pockets that branch off the main channel. These molecules enjoy a more favorable stabilization than those within the main channel. When the WA1 ratio decreased to 5 , the extra Na cations blocked off the side pockets; the methane molecules were then all located in the main channel and the heat of adsorption was reduced (Fig. 8). If such a decrease in the heat of adsorption could be seen experimentally, it would not be expected to be quite so pronounced. The use of periodic boundary conditions in the simulations overemphasizes the regularity of positions at which cations are to be found. As soon as a cation is placed on a site, the boundary conditions replicate it at equivalent points within the cell. In real zeolites, it is expected that Na cations would be distributed

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FIG.8. Distribution of methane molecules in mordenite at 300 K. The left figure shows a framework with a Si/AI ratio of 11, and the right, a framework with a %/A1 ratio of 5. Each spot represents the center of mass of a methane molecule. Population of the side pockets is clearly seen at low Al concentrations. Reprinted with permission from Ref. 144. Copyright 1988 American Chemical Society.

over several nonequivalent sites. Nonetheless, the simulations demonstrate the importance of being able to visualize where the sorbates reside. This is a point we return to later, when discussing longer alkanes. In Silicalite. A variety of papers are concerned with sorption of methane in the all-silica p e n t a d , silicalite. June et al. (87) used a Metropolis Monte Carlo method and MC integration of configuration integrals to determine low-occupancy sorption information for methane. The predicted heat of adsorption (18 kJ/mol) is within the range of experimental values (18-21 kJ/ rnol) (145-150), as is the Henry’s law coefficient as a function of temperature (141, 142). Furthermore, the center of mass distribution for methane in silicalite at 400 K shows that the molecule is delocalized over most of the total pore volume (Fig. 9). Even in the case of such a small sorbate, the channel intersections are unfavorable locations. In an MD study of methane sorption and diffusion in silicalite, Nicholas et al. (67) identified favorable sites for sorption. From the MD calculations, the time-averaged position of the center of mass of the methane molecule was plotted. Energy minimization calculations were then performed, locating the methane molecule at positions where the MD calculations predicted they spent the most time. Each channel intersection region was found to contain two sites that are minima for methane-zeolite interactions. These two sites are separated by a translation parallel to the straight channel

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Z

FIG.9. Center of mass distribution of methane in silicalite at 400 K. The light grid illustrates the pore volume, and the dark shows delocalization over most of the cell. Reprinted with permission from Ref. 87. Copyright 1990 American Chemical Society.

A,

axis of approximately 0.5 so they cannot be occupied simultaneously. However, a single sorbed methane molecule can rapidly jump between them. The straight channel segments also provide two minimum energy sites, which bind methane more strongly than those sites in the intersections. The two sites are 2.5 A apart, again preventing simultaneous occupation by two methane molecules. Each sinusoidal channel segment was found to provide one minimum energy position, where methane experiences essentially the same interaction energy as in the straight channel segments. Thus, on a strictly energetic basis, these results predict no preference for sorption in either the straight or the sinusoidal channels, although the intersections are somewhat less favored. These minimum energy locations remain at higher loadings, until the crowding of sorbates takes place. However, the distribution of sorbates is somewhat more complicated than that predicted solely by energetic considerations. If distribution depended

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solely on channel volume, then, according to the calculations of Nicholas et af. (67), the sinusoidal channel should be slightly favored over the intersections and the straight channel (which comprises approximately 40% of the total pore volume). The anisotropic nature of diffusion and the existence of favorable sorbate-sorbate interactions at certain concentrations further complicate matters. Nonetheless, the fraction of time spent by the adsorbate in each pore region was determined as a function of concentration (from infinite dilution to 16 molecules per unit cell). The results show that at all concentrations there is a clear preference for sorption in the channels rather than the intersections. Higher concentrations show convergence toward the fractional occupancies predicted simply by the volume of pore region. The sinusoidal channel is consistently preferred over the straight channel, irrespective of energetic equivalence for methane sorption. Titiloye et al. (1.51)used energy minimization techniques to investigate the sorption of methane and longer alkanes in silicalite and ZSM-5. A portion of the zeolite lattice, approximately 240 atoms immediately surrounding the sorbate molecule, was allowed to flex. The potential parameters for the zeolite atoms, which included the shell model to account for oxygen polarizability, were taken from previous structural simulations (152). Formal charges were assigned to the zeolite atoms for the evaluation of long-range electrostatic contributions. The charges on the sorbate atoms were assigned on the basis of ab initio calculation results that employed the 3-21G basis set. The potential used t o describe the interaction of the framework with the adsorbate was that of Bezus et al. (1.53). The siting of aluminum in ZSM-5 was taken from the study of methanol sorption in ZSM-5 by Vetrivel et al. (154). The heat of adsorption of methane within silicalite was calculated to be 22.7 kJ/mol, which is slightly higher than the range of experimentally measured values. This is to be expected as the energy minimization technique assumes a temperature of 0 K and locates the absolute minimum energy rather than a distribution of values obtained with MD or MC. Starting from the center of the straight channel, the methane molecule was found to migrate into the intersection. In the model used for ZSM-5, the Si atom substituted for an A1 is located in the channel intersection (154). When the methane molecule is introduced into this framework, it is once again found to migrate to the intersection and reside close to the chargebalancing proton. The flexible framework acts to increase the heat of adsorption of methane by approximately 1kJimo1 in both silicalite and ZSM-5. This increase is more significant for longer alkanes, as discussed later. The presence of clusters of adsorbates has already been discussed for single-atom adsorbates, and Demontis and coworkers addressed this issue for methane (65). From the methane-methane RDF, clear evidence for

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clustering was found. Many dimers were observed as well as clusters of higher order in a linear or branched configuration near the channel intersections. It was estimated that only one-third of the adsorbates were free or isolated any one time, a far lower proportion than that estimated for singleatom adsorbates (approximately 90%). The mean number of dimers at room temperature is two from the ensemble of 24 molecules, with dimer lifetimes predicted to be as long as 1 ps. Trimer and higher n-mer lifetimes were predicted to be far shorter than this. A fixed framework model predicts that cluster lifetimes are longer than those of molecules in the flexible framework, and this effect is most evident at low temperatures. It can be accounted for in terms of energy exchange of the clusters with the vibrating framework walls, which discourage clustering. Thus, the imposition of a rigid framework leads to an “enhanced clustering effect” but other properties, such as the diffusion coefficient and activation energy, remain largely unaffected. Adsorption isotherms of methane in silicalite have also been predicted in a number of calculation studies (62, 155, 156). Goodbody et al. (62) predicted a heat of adsorption of 18 kJ/mol and simulated the adsorption isotherm up to 650 bar. From the adsorption isotherm, they found that the sinusoidal pore volume contains more methane molecules at all pressures. Snurr et al. (155) performed GC-MC and MD simulations over a wide range of occupancies at several temperatures. The intermolecular zeolitemethane potential parameters were taken from previous MD studies (11, 87) and the methane-methane parameters from MD simulations were adjusted to fit experimental results for liquid methane (157). Electrostatic contributions were neglected on account of the all-silica framework, and methane was represented by a rigid, five-center model. GC-MC simulations were performed for the system at 200 and 300 K. At 200 K, the saturation occupancy was found to be in good agreement with that obtained by Yamazaki et al. (146, 147), although the amount of sorption in the pressure range 0.05-0.3 atm was somewhat underestimated. At 300 K, the GC-MC low-pressure isotherm (up to 1 atm) lies well within the range of values determined experimentally (146,147,158).(Experimental variations may arise because of differences in samples, such as lattice imperfections, which are not accounted for in the simulations.) The highpressure isotherm (up to 55 atm) may be compared to the experimental isotherm of Ding et al. (159),determined for NaZSM-5 zeolite. The simulated isotherm predicts a higher sorption capacity at pressures exceeding 10 atm, possibly because of the presence of cations in the experimental sample that reduce available void volume. As the loading of sorbate within the zeolite increased, the interaction between zeolite and sorbate was found to become slightly less favorable. At low occupancies, the methane mole-

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cules occupy the most favorable lattice sorption sites, whereas at higher loadings they are forced into some of the less favored regions. This is more than compensated for by the increased methane-methane interaction energy, resulting in an overall more favorable interaction as the loading increases. MD simulations, complete with “ghost” particle insertions (160, 261), may be used to obtain static and dynamic information. (These particle insertions were performed after the MD runs and do not affect the calculations; they merely probe the insertion of particles into the system.) The MD simulations performed by Snurr et al. (155) were slightly more expensive than the GC-MC calculations, but they produced similar isotherms and also yielded important information about the structure of the adsorbed fluid. The methane molecules appeared to behave like an ordered fluid at all concentrations, although the structure does change. This change reflects the changing importance of sorbate-sorbate and zeolite-sorbate interactions as a function of loading. Smit (156) also presented adsorption isotherms for methane in silicalite, calculated by standard Monte Carlo methods. The methane molecule was modeled as a single united atom with methane-methane interactions described by the potential parameters of Verlet and Weis (162),which give a reasonable prediction of the vapor-liquid curve of methane. The zeolitemethane interaction parameters were optimized to reproduce both the experimental heat of adsorption and Henry’s law coefficient. Comparison of these values and other parameter sets that have been used in conjunction with a united-atom model of methane reveals wide variations. The size parameter (u)may vary by 25%, while the interaction parameter ( c ) can vary by almost a factor of 2. However, virtually all parameter sets predict values of the heat of adsorption within the experimental range of 18-21 kJ/mol. The deviation from the experimentally determined Henry’s law coefficient [ l O S mg/(g atm)] is far larger, up to a factor of 5 . By calculating the Henry’s law coefficients predicted by parameter sets used in other simulations, Smit was able to show that a reasonable estimation of the heat of adsorption does not guarantee an accurate prediction of the adsorption isotherm. The simulation of the adsorption isotherm up to a pressure of lo4 kPa is in very good agreement with that of Goodbody et al. (62). 4.

Other Hydrocarbons

A large proportion of the work described in this section refers to nalkane sorbates; however, other molecules such as alkenes and alcohols have also been considered. Most interest has been focused on the silicalite/ ZSM-5 system.

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Titiloye et al. (151) extended their EM calculations for methane in silicalite, ZSM-5, and all-silica faujasite to cover linear alkanes up to and including octane. A region of the lattice and the adsorbate molecule were allowed to relax. In general, the calculated heats of adsorption are all too high, as is expected when EM techniques are used. The overestimation generally increases as a function of carbon number, up to approximately 10 kJ/mol for octane in silicalite. Fixed framework simulations gave lower heats of adsorption, and the effect of flexibility was found to increase with chain length. When heats of adsorption in silicalite and ZSM-5 are compared, the values for methane and ethane are similar and the ZSM-5 values are approximately 10 kJ/mol larger. This same trend and the 10 kJ/ mol energy difference are also evident in experimental heat of adsorption measurements. For both silicalite and ZSM-5, the alkane molecules were found to favor the straight channels over the sinusoidal channels and to avoid the channel intersections strictly. The location of sites of minimum energy within silicalite and Z S M J was found to be considerably easier than in Na faujasites. (In Na faujasites, there were a number of preferred sites in the vicinity of charge-balancing cations.) Preferential sorption in the sinusoidal channels was confirmed by Nicholas et al. (67) in an MD study of methane and propane adsorption. This preference was most noticeable at infinite dilution; at a loading of 12 molecules per unit cell the distribution of molecules over the channels was found to be close to that expected from the relative volumes of the channel segments. The propane molecules were predicted to spend more time in the intersections than the straight channel at infinite dilution. This result is rationalized by considering the slow motion of the molecules and the conformational changes necessary to move from one channel type to another via an intersection. The distribution of propane backbone bond angles was predicted to be similar to that of gas-phase propane, indicating the rather minor effect of the zeolite on the internal coordinates of propane. June et al. investigated the sorption and spatial distribution of butane and three hexane isomers within the pores of silicalite, using a Metropolis MC method (87) and MD simulations (85).Perturbations of conformation as a result of confinement within the pore were also reported. Heats of adsorption and Henry’s law coefficients were found to be in good agreement with experimental values for butane (48-51 kJ/mol) (142,148,150,163-165) and n-hexane (70-71 kJ/mol) (163,166, 167). The heats of sorption of the other two hexane isomers, 2- and 3-methylpentane, were predicted to be 5 kJ/mol lower than that of n-hexane. Changes in the conformations of the sorbed alkane molecules were found to be far more pronounced for hexane than for butane when compared with their gas-phase conformations. In the case of butane, only slightly

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more trans bonds were predicted for the molecules adsorbed within the pores of silicalite. In the case of hexane, however, this increase was predicted to be 20%. More than 70% of the C-C bonds of hexane inside the pores of silicalite were predicted to be in the trans conformation. Thus, the assumption of rigid sorbate geometry for anything but the simplest of molecules is not realistic. The MD simulation results (8.5) subdivide the proportions of trans and gauche bonds into pore regions, for a loading of 4 molecules per unit cell. The molecules in the straight and sinusoidal channels are responsible for the observed increase in trans bonds for butane, this time compared to C-C bonds in the liquid-phase alkane. (The increase in the proportion of trans bonds was predicted to be far larger from MD simulations than that from MC results, described earlier.) In the case of hexane, the conformations of molecules in the intersections were also found to be dramatically different from those in the liquid phase. The linear alkanes were found to exhibit a clear preference for sorption within the channels, as opposed to the intersections. Such a location maximizes their dispersive interactions with the zeolite wall, without them encountering unfavorable repulsions. n-Hexane prefers the slightly narrower sinusoidal channels. Simulations for lower temperatures, when the spatial extent of the sorbate is reduced and it “freezes” into its minimum energy positions, showed minima in the sinusoidal channels. From the MD simulations, the distribution of butane within the channels was found to be 50% in the sinusoidal channel and 45% in the straight channel. These figures are almost independent of loading over a range of 1-8 molecules per unit cell. The proportions of hexane in the straight and sinusoidal channels were found to be almost equal and close to 50% (i.e., hardly any hexane molecules are found in the intersections). The two branched methylpentane isomers have a larger cross-sectional diameter, approximately 5.3 A, than the linear alkanes, 4.5 A. Although the simulations do not predict that the branched isomers will be excluded from the channels by virtue of their size, they do show a definite bias toward the channel intersections. The configuration-bias Monte Carlo (CB-MC) technique (112) has also been extensively applied to characterize the sorption of alkanes, principally in silicalite (111, 1.56, 168-171) but also in other zeolites (172-174). Smit and Siepmann (111, 168) presented a thorough study of the energetics, location, and conformations of alkanes from n-butane to n-dodecane in silicalite at room temperature. A loading of infinite dilution was simulated, based on a united-atom model of the alkanes and a zeolite simulation box of 16 unit cells. Potential parameters were very similar to those used in the MD study of June et al. (8.5). As expected, the static properties (heat of adsorption, Henry’s law coefficient) determined from the CB-MC simulations are therefore in close agreement with the values of June et al. The

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simulation time, however, was several orders of magnitude less for the CB-MC calculations. The calculated heats of adsorption were predicted to be in good agreement with selected values taken from the wealth of experimental measurements characterizing the alkane/silicalite system ( I 75). For alkanes shorter than n-octane, the simulations marginally underestimate the heats of adsorption, which was found to increase by 11 kJ/mol per CH2 group. For longer alkanes, a change of slope was observed and the incremental heat of adsorption was found to be 13 kJ/mol. The reason for this change in slope may be understood by examining the locations of the alkane molecules. Snapshots of the alkane conformation were saved periodically throughout the simulations for postcalculation analysis. By decomposing the end-toend vector of the alkane into components parallel to the three crystallographic axes, it was possible to determine the orientations of the molecule. This, together with the location of the center of the alkane backbone, allowed conformational snapshots to be divided into those in either the straight or sinusoidal channels or in the intersections. This distribution over the channels and intersections was found to change as a function of carbon number, as shown in Fig. 10. For the shorter alkanes (SC,), the probability of occupying the straight and sinusoidal channel segments was found to be approximately equal, and both are strongly favored over occupation of the intersections. The straight channel is preferred over the sinusoidal for longer alkanes. Nearly 80% of the n-decane molecules were found to be located within the straight channels. Thus, silicalite is predicted to become more unidimensional with increasing length of hydrocarbon, as a greater propor-

0 zig-zag Ostraigth

0.80

minter.

O'*Ot

0.00

0 0

3

5

u

0

7

a 9

1;; 1

1

1

3

Nc

FIG.10. Probabilities of finding an n-alkane in each of the three pore regions of silicalite as a function of carbon number, N,. Reprinted with permission from Ref. 111. Copyright 1994 American Chemical Society.

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SIMON P. BATES A N D RUTGER A. VAN SANTEN

tion of the hydrocarbon molecules are found in the straight channel. This description of the siting of the alkanes conflicts with those found from the simulations of Titiloye et al. (151),who concluded that the intersections are favored locations for all chains. There is correspondence between these results and that from the MD study of June et al. (85) for butane and hexane and the work of Nicholas et al. (67). Conclusions drawn from experimental studies offer similar inconsistencies with respect to siting. In most cases, a preferential adsorption location is predicted, but these preferences have been found to be different in different investigations. The heat of adsorption of a given alkane is not the same in the straight channels as in the sinusoidal channels, because of the different channel topologies. The increasingly preferred siting in the straight channels as the carbon number increases is responsible for the observed change in slope of the graph of heat of adsorption as a function of carbon number. The probability of finding a molecule within the channel intersection decreases from 0.1 (butane) to approximately 0.05 (dodecane). Conformational analysis of the alkane snapshots shows that for those alkanes sorbed in the straight channels, the percentage of all-trans conformers decreases with increasing carbon number up to heptane and then increases. This pattern clearly demonstrates the conformational changes induced by confinement within the zeolite pores. In the liquid state, the percentage of all-trans conformers decreases with increasing chain length. For those molecules in the sinusoidal channels, the proportion of all-trans alkanes decreases steadily with increasing carbon number. In the intersections, it is not possible to accommodate a molecule longer than heptane in the all-trans conformation, and most longer molecules adopt a coiled conformation. The efficiency of the CB-MC technique has been used by Maginn et at. (169), who considered the low-occupancy thermodynamics of sorption of alkanes as long as C25 in silicalite. The locations of such long molecules are no longer correctly predicted by considering the end-to-end vector and the chain midpoint. To overcome this problem, a “coarse-graining” technique was used to describe both the adsorbate and the zeolite, allowing for accurate microscopic characterization. The siting of short-chain alkanes (SC,) was predicted to be similar to that found by Smit and Siepmann (111, 168), notwithstanding different potential parameters and methods of characterizing conformations. A slight preference for butane sorbing in the sinusoidal channel was found, but no real preference was found for hexane. In contrast to the results of Smit and Siepmann, there is negligible probability of an alkane coiling exclusively in the intersection. As the temperature is increased to 800 K, the most probable coarse-grained conformations are depopulated at the expense of

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less favorable, more twisted ones. This prediction is in agreement with the delocalization of many sorbates found at elevated temperatures. Medium-length chains are considered to range from C8 to CI4. Octane appears to explore all regions of the pore volume. At low temperatures, there is a preference for the linear conformation (which spans two straight channel segments and an intersection). At a temperature of 650 K, a bent conformation of hexane spanning a straight channel, an intersection, and a sinusoidal channel segment was found to be most populated. At room temperature, longer chains increasingly favor the straight channels, and this preference increases with chain length. In this temperature region, silicalite does indeed appear to become more unidimensional, as stated earlier. In many ways, the high-temperature behavior is of more relevance since the low diffusivities of such long molecules preclude the catalytic use of such materials at low temperatures. It was shown that the preference for straight channel siting disappears at moderate to high temperatures, as higher entropy conformations become populated. The alkanes still explore the whole of the void volume; thus silicalite is still three-dimensional under the catalytically relevant conditions of elevated temperatures. The trend continues up to the longest alkanes investigated, and these chains can now populate a wide range of coarse-grained conformations at elevated temperatures as a consequence of their length. Macroscopic thermodynamic properties, such as Henry’s law coefficients and heats of adsorption, were also calculated. When experimental results were available for comparison, the agreement was good. The Henry’s law coefficients were compared to a correlation proposed by Hufton and Danner (150) and with the simulations of Smit and Siepmann. An approximately linear increase as a function of chain length was observed. Heats of adsorption at 300 and 650 K were found to be in similar good agreement. The heat of adsorption was shown to be almost independent of temperature for chains up to octane; however, for longer molecules the decrease with increasing temperature was found to be distinct, as a result of the randomization of alkane conformations. The agreement with the correlation of Hufton and Danner (150) is better at high temperature for the longer chains, whereby localization effects disappear. The CB-MC method has been used to simulate the adsorption isotherms of various alkanes in silicalite (170, 171). Using potential parameters that were fitted to obtain good agreement with experimental Henry’s law coefficients, Smit and Maesen (170,171) have simulated the adsorption isotherms of straight-chain alkanes in silicalite. Good agreement was obtained for ethane and propane in comparison with the different type-I curves measured experimentally. The overall agreement with experimental isotherms was found to be satisfactory with hexane and heptane, and a kink is seen

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in the isotherms at a pressure corresponding to approximately 50% of maximum loading. This kink is far more pronounced for heptane than it is for hexane, and it has been observed experimentally (176-178). The information provided by the simulations on a molecular level sheds light on a phenomenon that had hitherto been unexplained. Snapshots of the locations of some hexane molecules at 50% loading demonstrate that they are approximately the same length as a period of the sinusoidal channel and that they explore all regions of the pore volume. At this loading they are relatively free and molecules can occupy the channel intersections as well as the straight and sinusoidal channel segments. At near-saturation loadings, nearly all hexane molecules were found to be in the sinusoidal channels, no longer being afforded the same range of motion evident at lower loadings. The kinks in the isotherms arise from the fact that the hexane molecules have to be “frozen” into this distribution, resulting in a loss of entropy. This “freezing” will only happen if the pressure of sorbate is high enough to compensate for the loss of entropy. The presence of the step in the adsorption isotherms was found to become more pronounced at elevated temperatures and to occur at progressively higher temperatures with increasing chain length. The entropy contribution is more important at higher temperatures and a higher entropy value is necessary to freeze the alkane molecules. This predicted behavior may explain why one roomtemperature experimental measurement failed to show a kink ( I 78), whereas one at a slightly higher temperature (176) did. The fact that the lengths of only hexane and heptane molecules are commensurate with one period of a sinusoidal channel explains why kinks are not observed in the isotherms of shorter or longer molecules. If they are longer, they cannot fit exactly into the sinusoidal channel, and one or both ends extend into an intersection. Nothing is gained by a collective freezing in the zigzag channels. If the molecules are shorter, part of a second molecule can occupy the same channel segment, resulting in a different type of packing. The presence of this type of phase transition has subsequently been confirmed by experimental temperature-programmed desorption measurements (179). Bates et al. (172-174) considered the energetics, locations, and conformations of alkanes ranging from butane to decane in a variety of different all-silica zeolites. Calculations similar to those described already were performed for alkanes in mordenite, zeolite rho, faujasite, ferrierite, and zeolite A. A linear increase in the calculated heat of adsorption with increasing carbon number was found for all zeolites. Less experimental information is available to compare with the calculated heats of adsorption, and thus the performance of the technique and parameters cannot be subjected to quite the same scrutiny as the results for silicalite (111).Nonetheless, where

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comparisons could be made, the simulation results were found to be in good agreement with experiment. In the case of faujasite, the calculated value of the incremental heat of adsorption per CH2 group was found to be almost identical to that determined experimentally for a sample of ultrastable Y zeolite (175), a material with only a few highly active acidic sites. This similarity is rationalized by considering the interaction with an acidic site to be a perturbation to the all-silica case. Most of the molecule is subjected to an all-silica environment, with only one CHZl3group interacting with the acidic site. The energetics data are presented in terms of heat of adsorption as a function of average zeolite pore diameter. Average pore diameter is applicable to those zeolites with elliptical pore openings, and the pore dimensions employed are those usually used to characterize the zeolite. The heat of adsorption as a function of pore diameter was predicted to exhibit a maximum around 5 A for all the alkanes studied, as shown on Fig. 11. The optimum heat of adsorption of straight-chain alkanes appears to be achieved by a pore with dimensions close to that of the 10-ring channel in ferrierite.

0

i 3

4

5

6

7

8

average pore diameter (A] FIG. 11. Heat of adsorption as a function of pore diameter for alkanes ranging from butane to decane as found by CB-MC calculations. Reprinted with permission from Ref. 172. Copyright 1996 American Chemical Society.

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The decrease in the heat of adsorption as the pore size is increased beyond this size is not surprising; dispersive interactions with the zeolite pore decrease. The behavior at lower pore dimensions is explained by considering the location of the sorbed molecules. In the cases of zeolites rho and A, the alkanes were found to adopt highly coiled conformations in the centers of the a-cages that form these structures. The alkanes thus are located in pores with a larger diameter than that usually used to characterize the zeolite (namely, the diameter of the windows that connect the cages). If the heat of adsorption as a function of pore diameter is replotted to reflect the locations of the sorbed molecules, a more straightforward inverse relationship is obtained. The locations and conformations of the alkanes in the other zeolites were also considered (173).In mordenite, butane was predicted to adsorb relatively unconstrained, being able to orient itself either parallel to the main pore direction or parallel to the longer elliptical pore direction. Longer alkanes align exclusively parallel to the main pore. In ferrierite, butane is distributed over the two channel systems, with ca. one-third favoring the smaller 8-ring channel over the larger 10-ring channel. Pentane and longer alkanes were found only in the 10-ring channel, being forced to adopt an all-trans conformation because of the narrowness of the pore channel (5.4 X 4.2 The large void volume within the supercages of faujasite results in conformations that are similar to those found in the free, gasphase alkanes. Freeman et al. (180) investigated the sorption of isomeric butenes in silicalite, using a combined MD/MC/EM technique (5). This is a system of considerable catalytic interest since the isomerization of but-1-ene to give isobutylene and subsequent reaction with methanol produces methyl rer-t-butyl ether, an effective replacement for lead in automotive fuel (181). The combined calculation method is not dependent on any assumption about the initial distribution of the sorbates at the start of the E M calculation. Results show that isobutylene is the least strongly bound of the compounds, although by only a few kJ/mol. The authors stated that this result is in accordance with the selectivity of silicalite/ZSM-5 for production of the isobutylene isomer. But-1-ene was predicted to adsorb the most strongly, followed by trans- and cis-but-Zene. In all cases, the minimum energy location of the isomers was found to be at the channel intersections. The technique, while not as theoretically rigorous as some, does give insight into systems that are beyond the limits of techniques such as MD. Vetrivel et al. (154) applied similar EM methods to investigate the sorption of methanol and ethene in silicalite and ZSM-5. Calculation parameters were taken from an earlier structural study (152), and a portion of the silicalite lattice comprising 120 atoms was allowed to relax. The sorbate

A).

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molecules were held rigid, at geometries determined from ab initio calculations. A crucial question is the location of the aluminum atoms within the lattice; the authors chose the T2 site, stating that there was no clear evidence for preferential siting. Methanol was found to adsorb in the 10-ring of silicalite where it opens into the sinusoidal pore, with an adsorption energy of approximately 90 kJ/mol. In ZSM-5, it was found to be located close to the acidic OH group in which the proton bridges T2 and T8. The adsorption energy was predicted to be lower, approximately 61 kJ/mol. The authors stated that more reliable estimates of the adsorption energy may be found by using quantum chemical calculations. Indeed, there is a vast body of literature on the modeling of methanol interacting with a zeolite cluster based on quantum chemical methods, and we return to this subject in the next part of the review. Ethene was also found to adsorb less strongly in ZSM-5 than silicalite. In the aluminosilicate model, a large displacement away from the starting point (the center of the 10-ring channel) was observed. In the all-silica material, little displacement was found. Shubin et al. (182) reported a study of the isomers of butanol in silicalite and ZSM-5. Calculations were done with the cvff force field (183), augmented with a shell model description of oxygen polarizability (152). The locations of the aluminum atoms-over a variety of T sites-were taken from the calculations of Schroder et al. (184). The charges for butanol were taken from ab initio calculations. Calculation results were found to be very sensitive to the charges apportioned to atoms. Two charge sets were used: formal charges (Si4+,02-) on the zeolite atoms and the ab initio butanol charges, scaled by 1.0 and 0.6. [This has the same effect as choosing lattice in accordance with the force field of van Beest charges of Si2.4+and 01.2-, el al. (185).] The calculation method was basically the combined MD/ MC/EM method, with a number of small alterations. In calculations for ZSM-5, the number of silicalite-like structures (where the alcohol interacts with just the siliceous framework, not the OH group) was minimized by making the OH group a more preferred site for sorption. This was done by scaling the interaction parameter between zeolite H and butanol 0 by 1000. A second minimization was then performed with normal interactions as a starting point for flexible lattice minimizations, by relaxing spherical regions around the adsorption site of 8 A (Model I) or 9.8 A (Model 11). Sorption energies in silicalite were found to be spread over a relatively small range, 18 and 9 kJ/mol, for full- and reduced-charge models, respectively. In each case, the preferential binding order was the same; 1-butanol adsorbs the most strongly, then 2-butano1, and finally butyl alcohol. The locations of adsorption were also found to be similar for all isomers; 1-butanol prefers to adsorb in the sinusoidal channels, whereas 2-butanol

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prefers the straight channel. The other two isomers showed different preferences, depending on which of the two different charge models was used. This information about the energetics and locations would suggest that at ambient temperatures the molecules would be distributed over a range of different sites. In ZSM-5, the preferred locations of butanol are the AlOSi bridges. In these calculations, the six most energetically favored acidic OH group locations determined from the calculations of Schroder et al. (184) were used, and all but two of these direct the proton into the sinusoidal channel. Once again, a rather narrow range of adsorption energies was found. Which isomer adsorbs most strongly depends on the charge model used; in the case of the full charges, it is l-butanol but for the reduced-charge model it is 2-butanol. Both models predict that the same sinusoidal channel acidic OH group is favored for l-butanol, 2-butanol, and butyl alcohol. The heat of adsorption was found to be greater by 3-15 kJ/mol in ZSM-5 for all isomers except 2-butanol; experiments suggested that this preference in ZSM-5 is 20-25 kJ/mol (186). The discrepancies are most likely a consequence of the assumption of a fixed zeolite lattice and inadequate intermolecular potential parameters.

5. Aromatics and Other Compounds Early studies of the favored sorption sites of larger cyclic and aromatic molecules date back to the mid-1980s. Wright et al. (187) predicted the location of pyridine inside the channels of gallozeolite L by systematically varying the distance between the N atom and the charge-balancing K cation and also the an$e of tilt of the molecule. The predicted position is found to be within 0.2 A of that determined by neutron powder diffraction analysis. Pickett et al. (188) used similar potentials (79) to investigate the location of p-xylene in silicalite. They found the global minimum energy position within the straight channel system, with partial rotation about an axis parallel to the straight channel activated by approximately 9 kJfmol. Other minima were located elsewhere in the channels, higher in energy by only a few kJ/mol. The discrepancies between the global minimum location and that predicted from X-ray powder diffraction (at the intersection) are explained by the limitations of the theory to 0 K and by the relatively small barriers between minima at ambient temperature. In these two studies the limitation of a rigid zeolite and adsorbate molecule was imposed. The preferred adsorption sites for benzene in zeolites such as faujasites have been thoroughly characterized experimentally (95, 97). The two preferred sites are atop an SII cation and (at higher loadings) in the 12-ring

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window of the supercage. This siting has been predicted by MD simulations (94,97),MM calculations (98), and the TST formalism (101). (These studies were discussed in the section of this review concerned with diffusion.) These two sites for benzene adsorption were also predicted by Uytterhoeven et al. (189) on the basis of EEM calculations (190, 191). Different faujasites were modeled: US Y, Y (Si/Al = 3), and X (WAl = 1);zeolite X was modeled with both monovalent and divalent cations; both the benzene and crystal geometries were held rigid. A single benzene molecule was found to be positioned in the supercage in a plane perpendicular to the body diagonal of the cubic unit cell. The interaction energy profiles calculated as a function of position along the threefold axis clearly show the existence of the two favored sorption sites in all the zeolites. In NaY and NaX zeolites, the type I1 cation site is preferred over the window site by approximately 30 and 20 kJ/mol, respectively. In the case of US Y zeolite, the trend is reversed, and the window site is more favorable, although the energy differences between the two sites are of the order of only a few kJ/ mol. It is clear that the concentration of framework aluminum-and hence extraframework cations-is critical in determining both the preferred site and the strength of adsorption. As expected, the cation site in NaY zeolite binds benzene more strongly than that in NaX zeolite. At the cation site, binding occurs by the formation of a .rr-complex between the cation and the aromatic ring. At the window site, it is the interaction between hydrogen atoms of benzene and framework oxygen atoms that stabilizes the complex. If the size of the cation in the zeolite X is increased from that of Na to that of K, the window site becomes preferred, and this preference is enhanced by further increases in cation size to those of Rb and Cs. Similarly, increasing the cation size in NaY zeolite destabilizes the binding at the cation site, although in this case it is still more favorable than at the window site. If the formal charge on the cation is increased to +2, the cation site is preferred as a result of the formation of very strong complexes with the benzene molecule. Thus, the influence of aluminum content on the adsorption energy cannot be examined without considering the nature of the charge-balancing cation. The charge distribution over the framework partially explains some of the observed energetic trends, but not all. It was shown that the only property to correlate consistently with the trend in interaction energies is the mean electric field gradient on the C-H bonds of benzene. The complexity of modeling the adsorption of benzene in silicalite has already been discussed in the section concerned with diffusion. A TST study by Snurr et al. (106) led to the identification of 27 unique sorption minima in the asymmetric unit. Given this result, it is unsurprising that there have been relatively few simulation studies of this system. However,

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those that have been published have relied on simulation methods that allow orders of magnitude improvement in efficiency (192, 193), while still allowing predictions that are in good agreement with experiment. Snurr et al. (192) used biased GC-MC simulations to predict isotherms, isosteric heats of adsorption, and locations of benzene and p-xylene at various concentrations. The suitability of a bias method is clear, at low coverages to prevent trial insertions overlapping with the zeolite walls and at high coverages to prevent overlap with other sorbate molecules. (Slightly different bias schemes were used for the two extremes of concentrations.) Interactions between sorbates and zeolites-both of which were considered to be rigid-were modeled with parameters from the literature (79, 87). Electrostatic interactions were included to account for the quadrupole moment of the sorbates. Sorbate-sorbate interaction parameters were taken from Shi and Bartell (194) for benzene and from Jorgensen et al. (195) for p-xylene. Simulations were performed for two forms of the silicalite structure; one was the as-synthesized orthorhonibic form (196) (hereafter designated ORTHO) and the other was the form with P212121symmetry (hereafter designated PARA) observed for silicalite at high loadings of p-xylene (197). The experimental isotherm for p-xylene in silicalite shows a discontinuity (or step) at a loading of about 4 molecules per unit cell. This has been interpreted as a phase change from the ORTHO to the PARA form. The simulated isotherms are in good agreement with this observation; a compound isotherm-the ORTHO form for up to 4 molecules per cell and the PARA form for more-is similar to that measured experimentally. The general trend seen in experimental isotherms of benzene in silicalite is that at high temperatures they follow a Langmuir shape and level off at 4 molecules per unit cell. A t low temperatures, the isotherms display a kink or step around 4 molecules per unit cell and then level off to a plateau around 8 molecules per unit cell. The slow diffusion of benzene in silicalite, and thus the slow equilibration for adsorption measurements, means that there are considerable discrepancies among experimental isotherms (198201). The simulated isotherm for the ORTHO structure does not agree at all well with this general picture; no steps were predicted and the saturation loading was found to be too high. Given the structural similarities between benzene and p-xylene, the authors suggested that the same sort of sorbateinduced phase transition occurs in the benzene-silicalite system. A composite isotherm does indeed show better correspondence with the general trend from experimental investigations. At a high benzene loading (8 molecules per unit cell), well-defined adsorption sites exist in both channels and intersections. If the isotherm for the ORTHO structure is decomposed to reflect where benzene molecules

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adsorb, it is shown that the channel intersections are filled initially. At a loading of 4 molecules per unit cell, when all the intersections are occupied, the straight and sinusoidal channels fill almost simultaneously until a maximum loading is reached. The picture is somewhat different for the PARA structure. The sinusoidal channels and the intersections fill simultaneously, with almost no adsorption in the straight channels. The results show that the subtle changes in the position of the zeolite atoms from one structure to another are sufficient to induce completely different siting of the benzene molecules within the pore network. If the benzene-silicalite system does show a phase transition, the siting at low loadings corresponds to that seen from the ORTHO simulations, and that at high loadings corresponds to that from the PARA simulations. Similar sitings were deduced for p-xylene. There is a lack of uniform agreement regarding siting in the experimental literature; neutron diffraction studies (202) suggest that the sorbate locations at high loading are the intersections and the straight channels, whereas NMR measurements (203) were interpreted to indicate preferential filling of the channels before the intersections. Sorbate-sorbate interactions were found to be negligible in the ORTHO regime of the isotherm; molecules occupy intersections and are too far away from each other to interact. At higher loadings, in the PARA region of the isotherm, the sinusoidal channels become filled, and the significant sorbate-sorbate interactions between neighboring benzene molecules result in an increase of the isosteric heat of adsorption. The predictions of these atomistic simulations have been confirmed by the same authors on the basis of a lattice model simulation (193).Features of the atomistic simulations, such as the steps in the isotherms, were also indicated by the lattice model, and more than an order of magnitude saving in computer time is realized. This increased efficiency originated from the adoption of a hierarchical simulation strategy. Short simulations based on the atomistic model were used to obtain parameters that were then fed into a coarse-grained lattice model. Only at high loadings of benzene were appreciable differences found between the results of the lattice and atomistic models. The suitability of a lattice model to reproduce the results of atomistic simulations rests on the fact that well-defined adsorption sites exist for benzene within silicalite. The technique would not be applicable to long-chain alkanes, for example.

D. SUMMARY As was the case for diffusion calculations, tremendous advances have been made recently in the simulation of the sorption locations, energetics, and conformations of adsorbates within zeolites. As far as the prediction

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of the energetics of sorption is concerned, a consensus has been reached regarding the theoretical techniques. It is unusual nowadays to find a calculated heat of adsorption that does not lie within the range of values determined experimentally. In the preceding discussions, we have demonstrated that the consensus appears relatively insensitive to small differences in, for example, potential parameters and model representations. Nonetheless, there are certain basic requirements that must be satisfied to guarantee the reliability of the results. One such requirement is the inclusion of adsorbate flexibility for medium- and long-chain alkanes. The broad agreement between calculated and experimental values of adsorption energies does not extend to the favored sites for sorption. Prediction of the locations for sorption can lead to very different results depending on the calculation methods. It is not always possible to determine whether this inconsistency is a consequence of the imposed calculation variables (potential parameters etc.) or of the methodologies themselves. Furthermore, the uncertainty in the theoretically calculated results finds a parallel in the experimental studies, as the debate concerning siting continues here as well. In general, plenty of experimental results are available to allow comparisons with calculated adsorption energies, although in some cases the systems being simulated cannot be replicated in an experimental laboratory. Experimental information regarding siting is rather more scarce. Notwithstanding the differences in the predicted location of sorbates, a great deal of valuable information has emerged from these theoretical studies. Together with the dynamic information resulting from diffusion studies, they have provided atomic-level detail of many of the fundamental steps that take place in zeolite-catalyzed reactions. It now remains for us to address the question of activation and reaction of sorbates; this is the subject of the next section.

IV. Bond Activation by Zeolites

A. SCOPEOF THISSECTION The previous two sections of this review deal with classical simulation methods. A description of the activation of adsorbates by acidic sites, together with any bond breaking or bond formation that may take place, is the realm of quantum mechanical (QM) simulations. These types of calculations are particularly well-suited to zeolite-adsorbate systems when the cluster approximation is used. The active acidic site in the zeolite is modeled by a molecular cluster, formed by cutting out a small portion of

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the zeolite lattice and terminating open valences with hydrogen atoms. The size of the cluster is chosen so that the system can be modeled quantum chemically and detailed electronic structure information about the acidic site and the sorbate molecule may be obtained. The method has clear shortcomings; the long-range effects from the rest of the crystal are ignored and any subtle effects that are influenced by pore topology are not accessible to this type of calculation. Furthermore, the clusters chosen are (almost always) completely general; i.e., they apply to no specific zeolite structure. However, aspects of a catalytic reaction which are only dependent on the local properties can often be treated very well by using a cluster approach. The applicability of the cluster method originates from the covalent nature of the chemical bonds in a zeolite; 90% of the bonding can be ascribed to covalent interactions. The great strength of the method is demonstrated by the volume of material that has been published on these clusters over the past 10 years or so. The detailed information that can be obtained for very modest clusters is in good agreement with what is known about zeolite structures from experimental studies. The applicability of the cluster approximation to zeolite chemistry has been extensively reviewed recently (204). In addition to cluster calculations, we report on recent advances made in the application of density functional theory applied to periodic lattices. So great is the volume of work on zeolite cluster calculations, both with and without adsorbate molecules, that to attempt to present a comprehensive review would result in a doubling of the length of this article. Instead, we select certain systems to highlight. These are the activation and reaction of methanol on a Brensted acid site and the activation of the C-H and C-C bonds in alkanes. These reactions are central to the industrial uses of zeolites as catalysts in the petroleum and petrochemical industries. Cracking, isomerization, and alkylation of hydrocarbons all proceed via proton transfer or carbocation formation. Unsurprisingly, these areas have attracted immense theoretical interest for a number of years, but there is still considerable debate in the literature. Interactions of zeolites with other molecules, such as CO (205-207) or CH3CN (208,209), are not considered here, although these too have attracted wide attention. We proceed with a short description of the quantum chemical methods and models that are used in these simulations before addressing each of the selected topics in turn.

B. THEORETICAL METHODOLOGIES AND MODELS At the time of their inception, cluster calculations of adsorbate-zeolite systems were largely treated by using a semiempirical method. In the mid-

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to late-1980sa general shift was seen toward the Hartree-Fock (HF) formalism, which employs no empirical parameters. The size and sophistication of the systems considered grew very rapidly, facilitated by advances in computing power. Electron correlation was incorporated into calculations by use of Moller-Plesset perturbation theory (210), first as a postoptimization correction and then later self-consistently. From the beginning of this decade, a second general shift was seen toward first-principles calculations that employed density functional theory (DFT), as these scale less expensively with the size and complexity of the system under consideration. Implementation of gradient correction terms to the exchange-correlation energy have made DFT calculations cost-effective competitors to highly correlated HF methods. Practically all of the calculation studies we discuss here were performed within the HF or DFT formalism, and most employed acid site cluster models that may contain anywhere between one and five Si and A1 atoms. Basis sets used to represent the electrons of the system were usually of double-zeta quality or higher; i.e., each filled orbital of an atom has been represented by two separate exponential functions. In addition, extra functions have been added-so-called polarization functions-to represent orbitals that are empty. These allow the orbital more flexibility and result in better theoretical predictions. In this review, we focus on cluster models of Bronsted acid sites, bridging hydroxyl groups that result from the incorporation of trivalent aluminum atoms into the siliceous framework during synthesis. These sites are by no means the only active sites within zeolites, but they are among the best characterized. Stable adsorption complexes are characterized by local minima on the potential energy hypersurface. The reaction pathway between two stable minima is determined by computation of a transition state structure, a saddle point on the potential energy hypersurface, characterized by a single imaginary vibrational mode. The Cartesian displacements of atoms that participate in this vibration characterize movements of these atoms along the reaction coordinate between sorption complexes. The choice of cluster model size is critical. It is essential that the cluster model be neutral and not subjected to optimization constraints. Both these restrictions have been shown to lead to artifactual behavior. Small clusters cannot be used to investigate concentration dependence, and if this dependence is to be considered, a larger model must be used. Similarly, a cluster should not be so small that it artificially constrains the spatial extent of the adsorbate complex or transition state. The acidity of the cluster-quantified by the deprotonation energy-is found to change as a function of cluster size. The deprotonation energy of a 3T atom cluster terminated with hydro-

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gen atoms yields a deprotonation energy close to that found for high-silica acidic zeolites (approximately 1250 kJimo1). Smaller clusters generally give higher deprotonation energies, indicating that they are less strongly acidic. Clusters that overestimate deprotonation energies also lead to uncertainties in the activation barriers to catalytic reactions, according to the BronstedPolanyi principle. (The effect of this has been estimated from calculations and is discussed later.) In addition to changes in acid strength as a function of cluster size, we discuss in the following sections how it is possible to vary the acid strength of an individual cluster by manipulation of terminal Si-H bonds. Of course, certain features of overall kinetics are inaccessible via a cluster model method, such as the influence of pore structure on reactivity. The cluster model method cannot integrate reaction rates with concepts such as shape selectivity, and an alternative method of probing overall kinetics is needed. This has recently been illustrated by a study of the kinetics of the hydroisomerization of hexane catalyzed by Pt-loaded acidic mordenite and ZSM-5 (221). The intrinsic acidities of the two catalysts were the same, and differences in catalyst performance were shown to be completely understood on the basis of differences in the heat of adsorption of hexene, an intermediate in the isomerization reaction. Heats of adsorption are strongly dependent on the zeolite pore diameter, as shown earlier in this review (Fig. 11). In addition to the cluster calculations, we report details of recent firstprinciples calculations based on the density functional formalism. These calculations employ periodic boundary conditions to allow investigation of the entire zeolite lattice, and therefore the use of a plane-wave basis set is applicable. This has a number of advantages, most notably that the absence of atom-centered basis functions results in no basis set superposition error (BSSE) (212),which arises as a result of the finite nature of atom-centered basis sets. Nonlocal, or gradient, corrections are applicable also, just as they are in the cluster calculations. AND REACTION OF METHANOL C. ACTIVATION

1. Proton Transfer

The theoretical modeling of the activation and reaction of methanol by Bronsted acid sites within zeolites has attracted a wide interest. This is in part a consequence of the industrial importance of the interaction-as the first step of the conversion of methanol into gasoline in the MTG process (213). However, a great deal of the theoretical interest has arisen because of the possibility of proton transfer from the zeolite lattice to methanol. An early investigation was that of Vetrivel et al. (214), who employed ab

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initio HF and lattice simulation methods to analyze the adsorption within zeolite pores. They predicted two possible orientations of methanol; the first showed a hydrogen bond between the methanol oxygen and the zeolite proton. The second had the methyl group oriented toward the zeolite framework, facilitating hydrogen abstraction via the lattice oxygens. It was later shown by Gale et al. (215) that this second geometry arose from an imperfect charge balance within the embedded cluster used by Vetrivel et al., resulting in the interaction of a methanol molecule with a lattice bearing a charge of -4. The work pointed to the need to ensure charge neutrality of cluster models. The semiempirical calculations of Gale et al. (215),based on a 20T-atom cluster, led to the identification of an adsorbed structure in which the hydroxyl group of the methanol hydrogen bonds with the zeolite proton. The question of whether the methanol molecule is adsorbed in a physisorbed or chemisorbed state was addressed by a number of authors (216222). The former state involves hydrogen bonding between the oxygen of methanol and a proton in an acidic hydroxyl, together with a similar interaction between the proton of methanol and a basic lattice oxygen of the zeolite. The latter interaction consists of the interaction of a protonated methoxonium ion and a negatively charged deprotonated framework. It was unclear for some time whether both the neutral physisorbed complex and the ion-paired complex were local minima on the potential energy surface. If both are minima, an equilibrium between the two states may be assumed. If only one is a minimum-a single potential well-this may be either the ion-paired complex (which implies a barrierless proton transfer) or the neutral complex (which implies that proton transfer does not take place). Early calculations that used symmetry constraints to shorten computation times led to predictions that both the neutral and ion-paired complexes were local minima (216, 217). Proton transfer was slightly favored (216) or disfavored (217),depending on the details of the calculation method. Gale et al. (218) inferred that the ion-paired structure was actually a saddle point on the potential energy surface, as they were unable to locate it as a local minimum. Other authors (219-221) demonstrated this unambiguously by analysis of the matrix of second derivatives of the ion-paired complex. The structure corresponds to a transition state for the exchange of the Bronsted and methanol protons. The nature of the ion pair complex is clearly shown in Fig. 12. The Cartesian atomic displacements of the imaginary stretching frequency that characterizes the transition state are added to the middle diagram of the figure. Experimental investigations of the interaction of methanol with a Bronsted acid site have failed to provide unambiguous evidence as to the

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v

b

n

C

FIG. 12. Minima and transition states on the reaction path of hydrogen exchange for methanol on a zeolite cluster. The upper and lower diagrams shown the equivalent neutral complexes, and the middle figure illustrates the transition state. Reprinted with permission from Ref. 221. Copyright 1995 American Chemical Society.

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nature of the interaction (223-231). The results from quantum chemical simulations are frequently compared to the experimental results from, for example, infrared measurements, and those involving methanol are no exception. Interpretation of the spectra (and the suggested correspondence between calculated and experimental bands) was often complicated, but a coherent explanation was offered by Pelmenschikov er al. (232). Strongly hydrogen-bonded complexes display a characteristic triplet of bands in the regions 2900-2800, 2550-2400, and 1740-1610 cm-l. In addition, the methanol spectrum exhibits another band at 3570-3489 cm-' . Upon interaction with methanol, the OH stretch of the acidic zeolite hydroxyl group is shifted to lower wavenumbers by the hydrogen-bonded interaction and is significantly broadened. The in-plane and out-of-plane zeolite OH deformations are shifted to higher wavenumbers, and the first overtones of these bands overlap with the broad OH stretch band. The characteristic triplet of bands is formed by creation of two transmission windows arising from Fermi resonance between the OH stretching and bending modes. These OH band splittings have recently been confirmed by Meijer et al. (208), who predicted the anharmonic infrared frequencies of acetonitrile adsorbed on a zeolite cluster. Haase and Sauer (220)have shown that the approximate positions of the in-plane and out-of-plane bending fundamentals of the zeolite O H group complexed with methanol are well reproduced by MP2 calculations, although poorly by standard HF methods. (The need for a self-consistent correlation or gradient correction in the H F or DFT calculation, respectively, has frequently been demonstrated in a number of these calculation studies.) The final band at 3570-3489 cm-' was suggested to arise from a different surface species (220), such as SiOH and SiOCH3 groups formed by dissociative adsorption of methanol on siloxane bridges (233, 234). Recently, first-principles simulations based on density functional theory via the molecular dynamics approach of Car and Parrinello (235)have been applied to study the problem of methanol within the pores of a zeolite (236-238). Nusterer et al. (236) performed calculations characterizing a methanol molecule trapped within sodalite at 400 K over a time period of 11 ps. The whole periodic lattice was included in the calculations. They found two structures for a single methanol molecule adsorbed on a Bronsted acid site. The first of these was found to be similar to the neutral complex found from cluster calculations (Fig. 12). The second involved a hydrogen bond between the methanol hydrogen and a zeolite oxygen that is a nextnearest neighbor of the Bronsted acid hydroxyl group. In other words, the complex between the zeolite and methanol constitutes an eight-membered ring, rather than the six-membered ring shown in Fig. 12. This latter structure is predicted to be 10 kJ/mol more stable and had not been suggested

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91

for the methanol-zeolite system before. The result is surprising since the basicity of an oxygen next to an aluminum should be larger than that of an oxygen bridging two silicon atoms. Hence the six-membered rather than the eight-membered ring structure should form a stronger complex. The question of methanol protonation was revisited by Shah et al. (237, 238),who used first-principles calculations to study the adsorption of methanol in chabazite and sodalite. The computational demands of this technique are such that only the most symmetrical zeolite lattices are accessible at present, but this limitation is sure to change in the future. Pseudopotentials were used to model the core electrons, verified by reproduction of the lattice parameter of a-quartz and the gas-phase geometry of methanol. In chabazite, methanol was found to be adsorbed in the 8-ring channel of the structure. The optimized structure corresponds to the ion-paired complex, previously designated as a saddle point on the basis of cluster calculations. No stable minimum was found corresponding to the neutral complex. Shah et al. (237)concluded that any barrier to protonation is more than compensated for by the electrostatic potential within the 8-ring. As a comparison, similar calculations were performed for methanol within sodalite, a structure in which the channel aperture is made up of 6rings unlike the 8-rings of chabazite (239).Steric hindrance prevents methanol from lying in the 6-rings, so it is forced into the more open cages of the structure. Methanol was predicted to exist physisorbed as a neutral molecule, in contrast to what was found for chabazite. The authors concluded that the structure of the zeolite is crucial in determining the nature of the adsorbed complex. An interesting point is that the binding energy of the neutral complex in sodalite is not much smaller than that found for chemisorption in chabazite. Thus, it appears that the energy penalty for deprotonating the zeolite lattice is offset by the strong interactions found for the ion-paired complex. We note, however, that the energy differences between protonated and hydrogen-bonded methanol are of the order of only a few kJ/mol. The infancy of these first-principles methods as applied to periodic zeolite lattices means that further detailed work is necessary, particularly in the area of verification of the ability of the pseudopotential to reproduce dynamic as well as static structural properties. However, the results found with these methods demonstrate that the debate concerning the modeling of the activation of methanol within a zeolite is far from concluded. The proton transfer to methanol as a reaction in its own right is, however, of relatively little interest. It does not govern the pathway or energetics of reactions such as dehydration to give dimethyl ether (DME). These are governed instead by the individual transition states that lead to the products, as we discuss in the next section.

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2. Dehydration The initial activation of methanol by a Bransted acid proton is merely the first of many stages in the MTG process (213).Many mechanisms have been proposed on the basis of experimental studies; of these the surface methoxonium ion pathway of Hutchings and Hunter (240) has received a partial experimental justification. The first stage of this mechanism involves the dehydration of adsorbed methanol to give water and a surface methoxy species. This reaction has been investigated theoretically by a number of authors (221, 222, 241). Blaszkowski et al. (221) demonstrated that the methanol molecule is capable of adsorbing in a physisorbed state in two different modes, the end-on mode, shown in the first part of Fig. 12, and a side-on mode, shown in Fig. 13a. In this side-on mode, a C-H bond of the methanol CH3 group is directed toward the zeolitic basic oxygen site, while the acidic zeolite proton retains its strong hydrogen bond with the methanol oxygen. The authors used TST ( 4 ) to determine the equilibrium constants for the two modes of adsorption from the computed adsorption energies. The equilibrium constant for the side-on mode is a factor of lo6 smaller than that for the end-on mode at 300 K. Thus, nearly all methanol molecules adsorb in an end-on manner, but the dehydration reaction necessitates conversion to the side-on form. The transition state leading to the dissociative complex of surface methoxy and water is shown in Fig. 13, as are the final adsorbed products. The transition state is product-like, indicating that the reaction barrier appears to be dominated by the breaking of the methanol C-0 bond. The overall thermodynamics of dehydration were predicted to be between 2 kJ/mol endothermic and 5 kJ/mol exothermic, depending on the precise method and cluster model used. With respect to the side-on physisorbed complex, the activation energy was calculated to be 185 kJ/mol. A similar reaction path has been described by Zicovich-Wilson eral. (222), who used single-point MP2 energy corrections applied to HF-optimized geometries. The corresponding calculated activation energy (approximately 220 kJ/mol) is somewhat higher than the DFT value of Blaszkowski and van Santen (221). It was stated earlier that DFT methods have a tendency to underestimate activation barriers somewhat (242). Initial DFT calculations of Sinclair and Catlow (241), employing a 3T-atom model cluster, suggested a value of 230 kJ/mol. The transition state leading to surface methoxy formation (Fig. 13) is reported to contain a planar CH3 carbenium ion fragment (221, 222). It is now accepted (204) that carbenium ion fragments do not exist freely in zeolites but must be bonded to the lattice oxygen atoms. They have been

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93

a

O

H

b

FIG.13. Dehydration of methanol on a BrGnsted acid site: (a) shows the side-on complex, (b) the transition state, and ( c ) the dissociative complex of surface methoxy and water. Reprinted with permission from Ref. 221. Copyright 1995 American Chemical Society.

observed experimentally by NMR spectroscopy (243).This carbenium fragment may be considered as an SN2-like intermediate species, with the oxygen atoms from the water fragment and the basic framework site acting as nucleophiles. SN2 reactions are generally characterized by a linear arrangement of the central species and arriving and leaving groups. The transition state leading to methoxy formation clearly shows a deviation

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SIMON P. BATES AND RUTGER A. VAN SANTEN

from this ideal value and therefore a certain amount of strain. ZicovichWilson et aZ. (222) found a bond angle of 124",whereas that of Blaszkowski et al. (221) is 148". Two different approaches have been proposed to eliminate this strain. The first is the use of a different acid site model cluster, as suggested by Zicovich-Wilson et al. (222). The basic oxygen site is no longer connected to an aluminum atom, but is in a next-nearest-neighbor shell. The authors demonstrated that this structure, which may be visualized as a pocket of the catalyst's inner surface, generates a transition state in which the strain is significantly reduced. (The angle is 168".) This results in a reduction of the activation energy by ca. 50 kJ/mol. A competing feature to this reduction in energy is the penalty incurred as a result of complex formation involving a next-nearest-neighbor atom, which is less basic than the nearest-neighbor oxygen. The use of this cluster results in overall reaction thermodynamics that are strongly endothermic (120 kJ/mol). The second method adopted to relieve the strain around the transition state leading to the formation of a surface methoxy species involves a second methanol molecule, as in the HF calculations of Sinclair and Catlow (241). These calculations were based on some constraints that prevented the results from being fully quantitative; 2T-atom clusters were used to model the acidic site and the 3-21G basis set was used for optimizations, which sometimes led to unphysical bonding effects, which forced the fixing of dihedral angles. The energies were calculated at the MP2 level with a larger basis set. Nonetheless, the method was considered by the authors to be semiquantitative, and calculation constraints were thought to be unlikely to affect mechanistic features. In addition, one advantage of the two different 2T-atom clusters employed is that effects associated with different proton affinities of basic oxygen sites may be investigated. The role of the second methanol molecule is to accept the zeolitic proton, while transferring its own hydroxyl proton to the other methanol molecule, which is then dehydrated. This results in a transition state (on both zeolite clusters) that contains a bond angle of around 170" for the SN2 participant fragments. The activation barrier to dehydration was then calculated to be lower, between 130 and 160 kJ/mol, depending on the nature of the acid site model used. This range of values was substantially lower than those found when just one methanol molecule was adsorbed (221, 222) (vide supra). The overall energetic profile of the reaction was found to be highly dependent on the specific acidic properties of neighboring lattice oxygen sites. Increasing either the basicity of the framework basic oxygen or the acidity of the zeolitic proton leads to a decrease in the activation barrier.

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3. Dimethyl Ether Formation After initial dehydration of methanol, the conversion to olefins, aliphatics, and aromatics up to Clo can proceed via the formation of dimethyl ether (DME). The reaction path to DME has been characterized by Blaszkowski et al. (244, 245). Different mechanisms for formation of DME have been proposed. Bandiera and Naccache (246) proposed that two methanol molecules adsorb simultaneously, forming two surface species ([CH,OH,]+ and [CH30]-) which then condense to give DME and water. Alternatively, Kubelkova et al. (227) predicted initial formation of a surface methoxy species via protonated methanol CH,OH,+. Subsequent reaction of the surface methoxy groups with another adsorbed methanol molecule produces DME. Both reaction pathways were investigated by using DFT calculations with self-consistent gradient corrections, without imposition of any symmetry constraints. It transpires that the route involving simultaneous adsorption of two methanol molecules as a first step presents an activation energy (with respect to the initial adsorbed complex) that is 70 kJ/mol lower than the route involving a surface methoxy. Thus, it would appear that DME is formed by an associative mechanism, without formation of intermediate methoxy. Calculated reaction rate constants (245) support this argument; the rate constant for the associative mechanism route is a factor of 10" larger at room temperature. Even at the elevated temperature of 700 K, the associative pathway rate constant is still larger by a factor of 1000. The calculated reaction profile agrees with the general idea of the mechanism proposed by Bandiera and Naccache (246),although the details are slightly different. The reaction profile determined from the QM calculations is shown in Fig. 14. Specifically, the first intermediate species proposed by Bandiera and Naccache, [CH30H2]+,has previously been characterized as a transition state rather than a stable intermediate, and the second ({CH,O]-) is not found. Two methanol molecules initially adsorb with an interaction energy of 65 kJ/mol per molecule (i.e., 130 kJimol in total). This value is reassuringly lower than the value found by the same authors for adsorption of a single molecule (73 kJ/mol) (221). The adsorption is followed by a rotation of one of the methyl groups of methanol (the one on the right in Fig. 14) to enable interaction with the hydroxyl group of the other methanol. Calculation of reaction rate constants (245)shows that at reasonable temperatures for DME formation (400 K), for every 7 million pairs of methanol molecules that exist in the as-adsorbed state (PII-ads1 in Fig. 14), only one pair exists in the rotated state. The transition state that subsequently leads to formation of adsorbed DME and water exhibits little strain on the SN2-likespecies;

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SIMON P. BATES AND RUTGER A. VAN SANTEN

t ;

74

ads1

I

s3

FIG.14. Reaction energy diagram for the formation of DME from methanol, proceeding by an associative path without intermediate methoxy formation. Energies are in kJ/mol; ads refers to an adsorbed complex, ts to a transition state. Reprinted with permission from Ref. 244. Copyright 1996 American Chemical Society.

the angle between intact methanol oxygen, carbenium carbon, and water oxygen is close to 180". The activation energy with respect to the reactant molecules is very low, only 15 kJ/mol. The activation energy with respect to the two adsorbed methanol molecules is 145 kJ/mol; this is somewhat higher than the experimental barrier found for mordenite, 80 kJ/mol (246). In a subsequent study, the same authors have considered another pathway that leads to DME formation (245). It is a hybrid of the associative path and that involving methoxy described in the preceding text. Initially, two molecules of methanol are adsorbed, and then a surface methoxy species is produced prior to DME formation. The reaction energy diagram for this hybrid path is shown in Fig. 15. The obvious difference in the mode of interaction of virtually all species, compared with Fig. 14, is the threefold nature of interactions. The model of the zeolite cluster is reduced to a 1-T model on grounds of computational expense (the authors suggested that the effect of cluster size is far less important than that of including nonlocal corrections self-consistently). The reaction begins as in Fig. 14 with the adsorption of two methanol molecules. The differences in computed interaction energies are a result of the use of different cluster models, each of which implies a different acid strength.

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FIG.15. Reaction energy diagram for the formation of DME from methanol, proceeding by initial adsorption of two methanol molecules, followed by surface methoxy formation. Energies are in kJ/mol; ads refers to an adsorbed complex, ts to a transition state. Reprinted with permission from Ref. 245. Copyright 1997 American Chemical Society.

The transition state that leads to the formation of the surface methoxy species is similar to that of Sinclair and Catlow (241) insofar as it exhibits little strain around the planar CH3 component. However, the activation energy for methoxide formation (referenced to the initial sorption complex of two methanol molecules) is lower in the calculations of Blaszkowski et al. (245).The value is 160 kJ/mol, compared with the value of Sinclair and Catlow (241) of 180-190 kJ/mol. The reason for this difference lies in the mode of interaction of the methanol molecules. The twofold interaction modeled by Sinclair and Catlow results in a costly rotation to allow formation of the transition state. Such a rotation is not required in the case of a threefold interaction. Two adsorption complexes were found for the interaction of the neutral water and methanol molecules with the methoxy-zeolite lattice, as shown in Fig. 15.The water molecule essentially is a spectator during the formation of DME, although it does stabilize the DME once it is formed. The largest activation energy of the elementary steps is that for methoxy formation (160 kJ/mol). This is similar to the barrier to DME formation via the

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associative mechanism (described earlier), and the conclusion is that either pathway may be responsible for the formation of DME. Calculated reaction rate constants support this conclusion. After discussing the dehydration of methanol and formation of DME, we are able to illustrate a number of key theoretical concepts. The first is that carbocation fragments are found in transition states, rather than as stable intermediates. Furthermore, the nature of these species is different from what is predicted from gas-phase studies, experimental or theoretical. The cluster, i.e., the zeolite, controls the stabilization of this carbocationic fragment. Second, we see that each different reaction requires a different transition state, rather than formation of a transition state that can be converted in a number of possible reactions. (This latter view received some support as a result of different processes possessing very similar activation barriers.) D.

ACTIVATION OF HYDROCARBONS

1. C-H Bonds

The activation of the C-H bonds in methane and the hydrogen exchange and dehydrogenation reactions that arise from it are the simplest prototype of a heterogeneous catalytic reaction involving a hydrocarbon and an acidic zeolite. Such a prototype reaction is of great importance, as the conversion of methane into fuel or desirable chemical products is a major goal of catalytic science and technology. Here we consider both the hydrogen exchange and dehydrogenation of methane on zeolite cluster models. We also describe how the model for hydrogen exchange may be tailored to account for differences in proton affinity of neighboring oxygen atoms and reproduce experimental exchange rates for different zeolites. In liquid superacidic media, the hydrogen-deuterium exchange reaction of alkanes is believed to involve carbonium ions (247),which are pentavalent, positively charged carbon atoms. The possibility of the existence of these species as stable intermediates in the hydrogen exchange of methane catalyzed by a zeolite has been investigated by a number of authors using cluster calculations (248-250). The methyl carbonium ion (CH;) was not found to be a stable intermediate, but instead a transition state for hydrogen exchange. This transition state is shown in Fig. 16,with the arrows representing displacement vectors along the reaction coordinate. Figure 16 shows that the reaction coordinate for exchange involves two zeolite oxygen atoms, one which acts as a proton donor (acid) and the other as a proton acceptor (base). Both exchangeable hydrogen atoms lie approximately halfway between the carbon atom and their respective oxygens. Comparison of literature geometries indicates that the precise

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99

@ A' Si C

0

O H

FIG.16. Reaction coordinate for hydrogen exchange between methane and an acidic zeolite cluster. Arrows represent displacement vectors along the reaction coordinate. Reprinted with permission from Ref. 248. Copyright 1994 American Chemical Society.

distances are somewhat dependent on the cluster model and calculation method used. The same is true of a charge analysis of the transition state complex. Blaszkowski et al. (248) and Kramer et al. (250) asserted that the interaction in the transition state is more covalent that ionic. Evleth et al. (249) asserted that the CH3 fragment is actually anionic in character, but they did mention that this may be partly as a consequence of the particular cluster and calculation method they chose. (Furthermore, the calculation and analysis of partial atomic charges on cluster fragments are subject to question and can only be used in an approximate manner.) Similar calculations characterizing the hydrogen exchange of ethane with an acidic zeolite cluster have also been performed (251), and a similar activation energy was obtained. Different calculations have predicted the activation barrier to hydrogen exchange between 125 and 167 kJ/mol: the DFT calculations of Blaszkowski et al. (248)yielded a value of 125-130 kJ/mol; the HF calculations of Evleth et al. (249) and Kramer et al. (250) predicted 140-167 and 150 2 20 kJ/ mol, respectively. All these values are in good agreement with what has been found experimentally (120-140 kJ/mol) (250,252).Kramer et al. (250) showed how the general 3T-atom cluster used in their calculations may be manipulated to vary the proton affinity of oxygen atoms participating in the exchange reaction. The proton affinity differences allow the cluster to mimic the behavior of two structurally dissimilar, compositionally equivalent samples of faujasite and ZSM-5. This is achieved by variation of the Si-H bond lengths at the edges of the cluster. In accordance with the bond order conservation principle, shortening the Si-H bonds will weaken neighboring bonds, strengthen next-nearest-neighboring bonds, etc. The authors found that the activation barrier increases in proportion to the difference in proton affinity induced between donor and acceptor oxygens (with a constant of proportionality 0.7). Thus, proton affinity differences will always act to reduce the rate of hydrogen exchange, which can be

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SIMON P. BATES AND RUTGER A. VAN SANTEN

calculated from activation barriers. This result suggests that the proton affinity differences are controlling the effective acid strength of the system and provide a way to simulate the behavior of different zeolites. Classical force field calculations led to estimates of the variations in strength of acidic sites within faujasite and MFI to be as large as 60 kJ/ mol (253).It was found (250) that variation of terminal Si-H bonds in the 3T-atom molecular cluster between 1.3 and 1.7 reproduced such a variation in proton affinity of oxygens. Using the proton affinity results calculated for ZSM-5 and faujasite from the classical simulations (250),the reductions in exchange rates were deduced for all pathways as results of cluster calculations at geometries that reproduce the proton affinity differences. For faujasites, there is only one crystallographically unique T atom in the asymmetric unit, and some of the exchange pathways are forbidden on steric grounds. For ZSM-5, the larger asymmetric unit (12 T atoms) makes the problem more complicated. The exchange rate was found to vary by up to 2 orders of magnitude over the T sites of ZSM-5. Assuming a random distribution of aluminum over all T sites in ZSM-5, the overall exchange rate at 700 K was found to be a factor of 20 lower than for the molecular cluster with equilibrium Si-H bond lengths. The equivalent reduction for faujasite is a factor of 80. The reduction factors vary with temperature, and the predicted rates as a function of temperature were found to be in very good agreement with those determined experimentally by the same authors for faujasite and ZSM-5 samples. Dehydrogenation of methane on a zeolite cluster has also been proposed to proceed via interaction of a CH; fragment with the deprotonated zeolite lattice. DFT calculations performed with a 3T-atom cluster (248) and HF calculations with a 1T-atom cluster (254) gave very similar results. The calculated transition state determined from the DFT calculations (248) that leads to dehydrogenation is shown in Fig. 17. In contrast to what was found for the hydrogen exchange reaction, the basic oxygen site stabilizes the CH3+ fragment of the reaction complex,

A

FIG.17. Reaction coordinate for dehydrogenation of methane on an acidic zeolite cluster. Arrows represent displacement vectors along the reaction coordinate. Reprinted with permission from Ref. 248. Copyright 1994 American Chemical Society.

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eventually leading to production of a surface methoxy species. The fourth C-H bond of methane is severely elongated, as is the zeolite hydroxyl bond. The displacement vectors in Fig. 17 clearly show the coalescence of the two hydrogen atoms to produce molecular hydrogen. Following on from the earlier discussion of strain in the SN2-likecomponents of a transition state, we note that the HF-calculated transition state (2.54)is far more strained than that shown in Fig. 17. This is not surprising; the DFT calculations employ a larger basis set and cluster model and are therefore considered to be more rigorous. The activation barrier to dehydrogenation [ca. 340-350 kJ/mol from the DFT calculations (248)]was found to be ca. 220 kJ/mol higher than that characterizing the hydrogen exchange reaction. Thus, hydrogen exchange is predicted to be a more facile process than dehydrogenation, as expected. Experimentally determined activation energies for dehydrogenation of C4 hydrocarbons are available for comparison (255-257). Values for isobutane are quoted as 216-225 kJ/mol on ultrastable Y zeolite (257) and 289-298 kJ/mol on HZSM-5 (255), whereas for n-butane a range of 156-165 kJ/ mol was reported (256). (All these values are true activation energies suitable for comparison with the theoretical results. They are the apparent experimental activation energy plus the heat of adsorption of the hydrocarbon on the zeolite.) First, all the experimental values are lower than the equivalent value for methane as the dehydrogenation of isobutane and of n-butane generates a tertiary and a secondary carbenium ion, respectively. These are both far more stable than the primary cation produced by methane dehydrogenation. On the basis of a DET calculation for the gas-phase species (248),a free tertiary carbenium ion was predicted to be 360 kJ/mol more stable than a free primary ion. The difference found between zeolitecatalyzed methane and isobutane dehydrogenation activation energies is far less than 360 kJ/mol and reflects the stabilization of the carbenium ion species as a transition state by the zeolite. Similar dehydrogenation calculations have been performed for ethane (251, 254), and the activation barrier was found to be ca. 50 kJ/mol lower than that for methane, on account of the secondary carbenium ion produced. 2. C-CBonds In considering the activation of C-C bonds, we limit ourselves to discussing the activation of ethene and the cracking of ethane, butane, and isobutane. Some of the earliest calculations characterizing the activation of olefins were performed by Kazansky and Senchenya (258-260) and Pelmenschikov et al. (261). In these calculations, it was shown that ethene could interact with a zeolite proton to form either a n- or a-bonded complex

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..

2.94

,,&

A

(4

i 207

1.79

:'

1.55

/h

C

(b)

(C)

FIG.18. Interaction of ethene with a zeolite proton. (a) shows the a-bonded complex, (b) the transition state, and (c) the u-bonded complex (260).

(Fig. 18). Initially, the n-bonded complex is formed, stabilized by approximately 7 kcal/mol (on the basis of calculations for a 1T-atom cluster with the 3-21G basis set). At low temperatures, the heat of adsorption of ethene corresponds to n-complex formation. Upon complexation, the ethene geometry remains largely unchanged, while the zeolite 0-H bond is lengthened slightly. The proton is transferred to the ethene in the transition state structure, and the positive charge on the other end of the ethene molecule is stabilized by the negative charge on the lattice. This ionic interaction between the carbenium ion and the deprotonated lattice proceeds to a covalently bonded species-the u-bonded complex-and generates a surface ethoxy fragment, similar to the methoxy fragment described earlier. This complex is stabilized by approximately 11 kcal/mol, referenced to the free cluster and ethene. The existence of surface alkoxy species has recently been confirmed by 13C-NMR spectroscopy (243). Notwithstanding the limitations of these calculations in terms of cluster model and basis set size, they clearly illustrate a number of key points. These include the existence of carbocations as transition states rather than intermediates and the bifunctional nature of a BrGnsted acid site with neighboring basic sites. The adsorption of saturated C2 hydrocarbons facilitates investigation of the cracking of ethane to give methane and an adsorbed methoxy species. Blaszkowski et al. (251)and Kazansky et al. (262)investigated the protolytic cracking of ethane with a zeolite cluster model. The former DFT calculations identified two different reaction pathways that led to cracking. In the first, ethane cracks directly; in the second, an additional rotation of the adsorbed ethane from staggered to eclipsed is necessary prior to cracking. This latter pathway closely resembles that found by Kazansky et al. (262)

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MOLECULAR BASIS OF ZEOLITE CATALYSIS

from their HF calculations. The reaction coordinate and optimized geometry of the transition state obtained in the direct cracking path of Blaszkowski et al. (251) are shown in Fig. 19. The transition state shows that the zeolitic proton has been transferred to one side of the ethane molecule; the C-C bond is almost broken, and the other side of the ethane molecule resembles a methyl cation. The reaction coordinate indicates subsequent formation of a surface methoxy species and adsorbed methane. The true activation barrier to cracking was computed to be 316-321 kJ/ mol, depending on the cluster model used (251). There was no energetic

a

b

1.913

FIG.19. Reaction coordinate (a) and optimized geometry (b) for the transition state that leads to cracking of ethane in the presence of a zeolite. Distances are in angstroms and angles in degrees. Reprinted with permission from Ref. 251. Copyright 1996 American Chemical Society.

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SIMON P. BATES AND RUTGER A. VAN SANTEN

difference between the two different cracking pathways of Blaszkowski et al. (251). The HF calculations of Kazansky et al. (262) predict a value of 390 kJ/mol, overestimated because of the assumption of a small cluster model and basis set. The small difference (20-30 kJ/mol) between the DFT-calculated value and the experimental value of Stefanadis et al. (255) for the cracking of isobutane catalyzed by ZSM-5 is fortuitous, as the secondary ion formed from isobutane cracking is expected to be more stable. The cracking reaction is expected to compete with dehydrogenation, which was calculated to have a similar activation barrier by the same authors (251). Larger alkane molecules, such as propane and butane, present a greater computational challenge and also afford the possibility of a number of different cracking paths. The cracking of n-butane in the presence of a zeolite cluster model was investigated by Collins and O’Malley (263, 264), who used 2T- and 3T-atom models and a combination of semiempirical and DFT methods. These authors found that protonation of the alkane molecules occurs at or near the center of the C-C bond being protonated. The resulting carbonium ion is formed directly over the aluminum atom, at significant distances from the zeolite surface. In other words, the acid-base functionality demonstrated in so many activations of C-H and C-C bonds was not observed. Furthermore, collapse of the transition state toward products produces an alkene and an alkane, rather than a surface alkoxy species, as suggested by Kazansky and Senchenya (258-260). The authors used the results from AM1 and DFT calculations to support their conclusions. In the AM1 case, the carbonium ion was predicted to be a highenergy, yet stable, intermediate. In the DFT calculations, it was found to correspond to a transition state. While the pathway via alkane and alkene formation is certainly plausible and corresponds to the observed catalytic reaction products at low conversions, further quantum chemical calculations are needed to resolve this point. The alkoxide formation suggested by Kazansky and Senchenya (258-260) is said to be specific to ethane protonation, as a consequence of its lack of P-protons. However, the recent calculations of Kazansky et al. (265, 266) and Viruela-Martin et al. (267) characterizing the cracking of isobutane and propene (267) in the presence of zeolites all refute this claim. The calculations of Kazansky et al. (265, 266), with a 1T-atom cluster and H F calculation methods, predict that the protolytic cracking of isobutane proceeds in a similar manner to that of ethane (described earlier). The zeolite proton attacks a C-C bond, resulting in the formation of a transition state that consists of a C4Hllcarbonium ion species stabilized by the negatively charged lattice. Other methyl hydrogens form hydrogen bonds with neighboring basic oxygen sites. The reaction proceeds via abstraction of methane

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to yield a surface secondary-propoxide species. The reaction is calculated to be almost thermoneutral, with an activation energy of 240 kJ/mol. This is far less than the value the same authors obtained for the protolytic cracking of ethane in the presence of the same cluster (262).The difference of 150 kJ was shown to arise in part as a consequence of the smaller basis set used in the earlier work (approximately 50 kJ/mol) but mostly as a consequence of the enhanced stability of the secondary-propoxide species relative to that of the methoxy species. The MPZcorrected HF calculation results of Viruela-Martin et al. (267) for a 3T-atom cluster yield a similar picture, with slightly different energy values. Kazansky et al. (265,266) also described similar reaction profiles to show the presence of carbonium ion-like transition states and adsorbed alkoxy products. These are the protonation of isobutylene, the dehydrogenation of isobutane, and hydride transfer from isobutane to surface tert-butoxide. The theoretical activation energies of the cracking, dehydrogenation, and hydride transfer reactions were all found to be somewhat higher than experimental measurements for isobutane activation (257).[The deviations, with corrected experimental values for the heat of adsorption of isobutane (150), were calculated to be approximately +20, +65, and not more than + 16 kJ/mol, respectively.] Kazansky et al. (265)recently proposed a method to correct these overestimated activation energies for the effect of using a finite cluster to model the zeolite. It has already been stated in the methodologies subsection of this part of the review that the cluster deprotonation energy, and hence acid strength, varies with size. Extrapolation of the deprotonation energy as a function of cluster size to a value representative of an average zeolite has been made and a value of 1236 kJ/mol obtained (268). This lowered deprotonation energy manifests itself in a weaker, more strongly acidic OH group. According to the BrmstedPolanyi principle, this lowers the activation barriers to a reaction. Kazansky et al. (265) proposed corrections to the barriers to cracking, dehydrogenation, and hydride transfer on the basis of this extrapolated value of average zeolite acidity. The reductions in activation barriers range from 25 kJ/ mol (for dehydrogenation) to 31 kJ/mol (for hydride transfer). Thus, the comparison between experimental activation energies and calculated values corrected for cluster size effects is significantly improved. In the case of protolytic cracking of isobutane, the corrected calculated value is within 4% of the experimental value. This quantification of the effect of cluster size and, more importantly, its effect on activation barriers allow far more meaningful comparisons between theoretical and experimental results. As these types of calculations are extended further, for larger hydrocarbons, not only does the complexity of individual calculations increase but so does the number of possible reaction paths. At present, calculations for

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larger hydrocarbons, e.g., hexane, are restricted to a semiempirical treatment (264).While this is a useful method to visualize adsorption geometries, the energies can at best be considered semiquantitative. E. SUMMARY

The use of zeolite clusters in quantum chemical calculations has now progressed to quite a sophisticated level. Elementary steps of reaction mechanisms can now be characterized and the results used to distinguish which steps are the most plausible. Computational power is such that clusters and methods can avoid obvious pitfalls (too small a cluster, basis set, etc.). Several key concepts that have arisen from theoretical studies are illustrated in the preceding discussion. These include the following: carbocations exist as parts of transition state structures, rather than as stable intermediates, and their stabilization is controlled by the zeolite lattice. The transition states are very different from the ground states to either side of them, and each different reaction has been shown to proceed via a different transition state. The nature of the methanol-zeolite interaction has been shown to be sensitive to a number of parameters and as such has proved to be a good benchmark for judging the reliability of quantum chemical methods. Not only are there a number of possible modes whereby one and two molecules interact with an acidic site (245), the barrier to proton transfer is small and sensitive to calculation details. Recent first-principles simulations (236-238) suggest that the nature of adsorbed methanol may be sensitive to the topology of the zeolite pore. The activation and reaction of methane, ethane, and isobutane have been characterized by using reliable methods and models, and realistic activation energies for catalytic reactions have been obtained. As with the summaries of the other sections, we mention a number of calculation parameters or variables that have been demonstrated to be of critical importance for accurate prediction of aspects of the interactions. Symmetry constraints on the clusters have been shown to introduce artifactual behavior. Corrections to account for the correlation of electrons have become essential in a calculation, and they must be incorporated selfconsistently rather than as postoptimization corrections. Basis sets need to have the flexibility afforded by double- or triple-zeta functionality and polarization functions to reproduce known parameters most accurately. The choice of the model cluster and its size affect the acid strength, and the cluster must be large enough not to spatially constrain reactants or transition states. The choice of cluster is invariably governed by the available resources, but a small cluster can still perform well. Indeed, some of the

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most elegant descriptions of the reaction pathways to DME formation have been performed with a cluster containing only 1 T atom (245). The need to use such small clusters inevitably comes as a result of the costly and time-consuming practice of locating transition states along reaction pathways. These transition states fill in the mechanistic details between the thermodynamic information provided by sorption minima. They have been shown in this section to provide invaluable information about the nature of a reaction coordinate between reactants and products. REFERENCES 1. Allen, M. P., and Tildesley, D. J., “Computer Simulation of Liquids.” Clarendon Press, Oxford, 1987. 2. El Amrani, S., and Kolb, M., J. Chem. Phys. 98, 1509 (1993). 3. Ryckaert, J. P., and Bellmans, A., Faraday Discuss. Chem. Soc. 66, 95 (1978). 4. Gilbert, R. G., and Smith, S. C., “Theory of Unimolecular and Recombination Reactions.” Blackwell, Oxford, 1990. 5. Biosym Technologies/Molecular Simulations Inc., San Diego, CA. 6. Kirkwood, J. G., Phys. Z. 33, 57 (1932). 7. Muller, A,, Proc. R. Soc. London, Ser. A 154,624 (1936). 8. Fraissard, J., and Ito, T., Zeolites 8, 3.50 (1988). 9. Barrie, P. J., and Klinowski, J., Prog. N M R Spectrosc. 24, 91 (1992). 10. Pickett, S. D., Nowak, A. K., Thomas, J. M., Peterson, B. K., Swift, J. F. P., Cheetham, A. K., den Ouden, C. J. J., Smit, B., and Post, M. F. M., J. Phys. Chem. 94,1233 (1990). 11. June, R. L., Bell, A. T., Theodorou, D. N., J. Phys. Chem. 94, 8232 (1990). 12. June, R. L., Bell, A. T., Theodorou, D. N., J. Phys. Chem. 95,8866 (1991). 13. Yashonath, S., J. Phys. Chem. 95,5877 (1991). 14. Yashonath, S., Chem. Phys. Lett. 177, 54 (1991). 15. Santikary, P., Yashonath, S., Ananthakrishna, G., J. Phys. Chem. 96, 10469 (1992). 16. Schrirnpf, G., Schelenkrich, M., Brickmann, J., Bopp, P., J. Phys. Chem. 96,7404 (1992). 17. Yashonath, S., and Santikary, P., Phys. Rev. B 45, 10095 (1992). 18. Yashonath, S., and Santikary, P., J. Phys. Chem. 97, 3849 (1993). 19. Santikary, P., Yashonath, S., J. Solid State Chem. 106, 184 (1993). 20. Yashonath, S., and Santikary, P., J. Phys. Chem. 98,6368 (1994). 21. Bandyopadhyay, S., and Yashonath, S., J. Solid State Chem. 111,151 (1994). 22. Kiselev, A. V., and Du, P. Q., J. Chem. Soc., Faraday Trans. 2 77, 1 (1981). 23. Demontis, P., Fois, E. S., Suffritti, G. B., Quartieri, S., 1.Phys. Chem. 94, 4329 (1990). 24. Karger, J., and Ruthven, D. M., “Diffusion in Zeolites and Other Microporous Solids.” Wiley, New York, 1992. 25. Heink, W., Karger, J., Pfeifer, H., Stallmach, F., J. Am. Chem. SOC. 112,2175 (1990). 26. Demontis, P., Suffritti, G. B., Quartieri, S., Fois, E. S., Gamba, A., J. Phys. Chem. 92, 867 (1988). 27. Eaterrnann, M., McCusker, L. B., Baerlocher, C., Merrouche, A,, Kessler, H., Nature (London) 352, 320 (1991). 28. Santikary, P., and Yashonath, S., J. Chem. Soc., Faraday Trans. 88, 1063 (1992). 29. Yashonath, S., and Santikary, P., J. Phys. Chem. 97, 13778 (1993). 30. Yashonath, S., and Santikary, P., Mol. Phys. 78, 1 (1993). 31. Bandyopadhyay, S., and Yashonath, S., Chem. Phys. Lett. 223,363 (1994). 32. Yashonath, S., and Santikary, P., J. Chem. Phys. 100, 4013 (1994).

108

SIMON P. BATES AND RUTGER A. VAN SANTEN

Ruthven, D. M., and Derrah, R. I., J. Chem. Soc., Faraday Trans. 1 71,2031 (1975). Santikary, P., and Yashonath, S . , J. Phys. Chem. 98, 9252 (1994). Yashonath, S., and Bandyopadhyay, S., Chem. Phys. Lett. 228,284 (1994). Bandyopadhyay, S . , and Yashonath, S . , J. Chem. Phys. 99,4286 (1995). Chitra, R., and Yashonath, S . , Chem. Phys. Left. 234, 16 (1995). Leherte, L., Andre, J.-M., Vercauteren, D. P., Derouane, E. G., J. Mol. Catal. 54, 426 (1989). 39. Leherte, L., Andre, J.-M., Derouane, E. G., Vercauteren, D. P., Comput. Chem. 15, 273 (1991). 40. Leherte, L., Andre, J.-M., Derouane, E. G., Vercauteren, D. P., J. Chem. Soc., Faraday Trans. 87, 1959 (1991). 41. Matsuoka, O., Clementi, E., Yoshirnine, M., J. Chem. Phys. 64, 1351 (1976). 42. Karger, J., and Pfeifer, H., Zeolites 7, 90 (1987). 43. Yashonath, S., Demontis, P., Klein, M. L., Chem. Phys. Lett. 153, 551 (1988). 44. Yashonath, S . , Demontis, P., Klein, M. L., J. Chem. Phys. 95,5881 (1991). 45. Meinander, N., and Tabisz, G. C., J. Chem. Phys. 79,416 (1983). 46. Yashonath, S . , Thomas, J. M., Nowak, A. K., and Cheetham, A. K., Nature ( L o n d o n ) 331, 601 (1988). 47. Stockmeyer, R., Zeolites 4, 81 (1984). 48. Jobic, H., Bee, M., Kearley, G. J., J. Phys. Chem. 98, 4660 (1994). 49. Bezus, A. G., Kiselev, A. V., Lopatkin, A. A,, Du, P. Q., J. Chem. Soc., Faraday Trans. 2 74,367 (1978). 50. Cohen de Lara, E., Kahn, R., Goulay, A. M., J. Chem. Phys. 90, 7482 (1990). 51. Alloneau, J. M., and Volino, F., Zeolites 6, 431 (1986). 52. Fritzsche, S . , Haberlandt, R., Karger, J., Pfeifer, H., and Wolfsberg, M., Chem. Phys. Lett. 171, 109 (1990). 53. Fritzsche, S . , Haberlandt, R., Karger, J., Pfeifer, H., Heinzinger, K., Chem. Phys. Left 198,283 (1992). 54. Fritzsche, S . , Haberlandt, R., Karger, J., Pfeifer, H., and Heinzinger, K., Chem. Phys. 174, 229 (1993). 55. Fritzsche, S . , Haberlandt, R., Karger, J., Pfeifer, H., and Waldherr-Teschner, M., in “Proceedings of the 10th International Zeolite Conference, Garmisch-Partenkirchen, 1994” (J. Weitkamp, H . G. Karge, H. Pfeifer, and W. Holderich, eds.), Part C, pp. 2139-2146. Elsevier, Amsterdam, 1994. 56. Fritzsche, S . , Haberlandt, R., Karger, J., Pfeifer, H., and Heinzinger, K., Chem. Phys. Lett. 242, 361 (1995). 57. Heink, W., Karger, J., Pfeifer, H., Salverda, P., Datema, K. P., Nowak, A., J. Chem. Soc., Faraday Trans. 88, 515 (1992). 58. Ruthven, D. M., and Derrah, R. I., J. Chem. Soc., Faraday Trans. I 68, 2332 (1972). 59. Demontis, P., and Suffritti, G. B., in “Proceedings of the 10th International Zeolite Conference, Garmisch-Partenkirchen, 1994” (J. Weitkamp, H. G. Karge, H. Pfeifer, and W. Holderich, eds.), Part C, pp. 2107-2113. Elsevier, Amsterdam, 1994. 60. den Ouden, C. J. J., Smit, B., Wielers, A. F. H., Jackson, R. A., Nowak, A. K., Mol. Simul. 4, 121 (1989). 61. Demontis, P., Suffritti, G. B., Quartieri, S., Gamba, A., Fois, E. S., J. Chem. SOC.,Faraday Trans. 87, 1657 (1991). 62. Goodbody, S . J., Watanabe, K., MacGowan, D., Walton, J. P. R. B., and Quirke, N., J. Chem. Soc., Faraday Trans. 87, 1951 (1991). 63. Nowak, A. K., den Ouden, C. J. J., Picket!, S . D., Smit, B., Cheetharn, A. K., Post, M. F. M., and Thomas, J. M., J. Phys. Chem. 95, 848 (1991).

33. 34. 35. 36. 37. 38.

MOLECULAR BASIS OF ZEOLITE CATALYSIS

109

64. Catlow, C. R. A., Freeman, C. M., Vessal, B., Tomlinson, S. M., and Leslie, M., J. Chem. Soc., Faraday Trans. 87, 1947 (1991). 65. Demontis, P., Suffritti, G. B., Fois, E. S., Quartieri, S., J. Phys. Chem. 96, 1482 (1992). 66. Kawano, M., Vessel, B., Catlow, C. R. A., J. Chem. Soc., Chem. Commum. 12,879 (1992). 67. Nicholas, J. B., Trouw, F. R., Mertz, J. E., Iton, L. E., Hopfinger, A. J., J. Phys. Chem. 97,4149 (1993). 68. Dumont, D., and Bougeard, D., in “Proceedings of the 10th International Zeolite Conference, Garmisch-Partenkirchen, 1994” (J. Weitkamp, H. G. Karge, H. Pfeifer, and W. Holderich, eds.), Part C, pp. 2131-2138. Elsevier, Amsterdam, 1994. 69. Dumont, D., and Bougeard, D., Zeolites 15, 650 (1995). 70. Jobic, H., Bee, M., Caro, J., Karger, J., in “Zeolites for the Nineties, Recent Research Reports, International Zeolite Conference, Amsterdam, 1989” (J. C. Jansen, L. Moscou, and M. F. M. Post, eds.), pp. 305-306. Elsevier, Amsterdam, 1990. 71. Caro, J., Biilow, M., Schirmer, W., Karger, J., Heink, W., Pfeifer, H., and Zdanov, S., J. Chem. Soc., Faraday Trans. 81, 2541 (1985). 72. Jobic, H., Bee, M., Caro, J., Biilow, M., and Karger, J., J. Chem. Soc., Faraday Trans. I 85,4201 (1989). 73. Jobic, H., Bee, M., Kearley, G. J., Zeolites 9, 312 (1989). 74. Burkert, U., Allinger, N. L., “Molecular Mechanics.” Am. Chem. Soc., Washington, DC, 1982. 75. Nicholas, J. B., Hopfinger, A. J., Trouw, F. R., Iton, L. E., J. Am. Chem. SOC.113, 4792 (1991). 76. Caro, J., HoEvar, S., Karger, J., Riekert, L., Zeolites 6, 213 (1986). 77. Zibrowius, B., Caro, J., Karger, J., 2. Phys. Chem. (Leipzig) 269, 1101 (1988). 78. Vessal, B., Leslie, M., Catlow, C. R. A,, Mol. Simul. 3, 123 (1989). 79. Kiselev, A. V., Lopatkin, A. A., Shulga, A. A., Zeolites 5, 261 (1985). 80. Briscoe, M. A,, Johnson, D. W., Kokotailo, G. T., McCusker, L. B., and Shannon, M. D., Zeolites 8, 74 (1988). 81. Bell, R. G., Lewis, D. W., Voigt, P., Freeman, C. M., Thomas, J. M., Catlow, C. R. A., in “Proceedings of the 10th International Zeolite Conference, Garmisch-Partenkirchen, 1994” (J. Weitkamp, H. G. Karge, H. Pfeifer, and W. Holderich, eds.), Part C, pp. 2075-2082. Elsevier, Amsterdam, 1994. 82. Wright, P. A,, Jones, R. H., Natarajan, S., Bell, R. G., Chen, J., Hursthouse, M. B., and Thomas, J. M., J. Chern. Soc., Chem. Commun., 633 (1993). 83. Natarajan, S., Wright, P. A., and Thomas, J. M., J. Chem. Soc., Chem. Commun., 1861 (1993). 84. Hagler, A. T., Lifson, S., Dauber, P., J. Am. Chem. SOC. 101, 5122 (1979). 85. June, R. L., Bell, A. T., Theodorou, D. N., J. Phys. Chem. 96,1051 (1992). 86. Hernandez, E., Catlow, C. R. A., Proc. R. Soc. London, Ser. A 448,143 (1995). 87. June, R. L., Bell, A. T., Theodorou, D. N., J. Phys. Chem. 94, 1508 (1990). 88. Jobic, H., Bee, M., Caro, J., in “Proceedings of the 9th International Zeolite Conference” (R. von Balmoos, J. B. Higgins, and M. M. J. Treacy, eds.), pp. 121-136. ButterworthHeinemann, Boston, 1993. 89. Nowak, A. K., Cheetham, A. K., Pickett, S. D., Ramdas, S., Mol. Simul. 1,67 (1987). 90. Barri, S. A. I., Smith, G. W., White, D., Young, D., Nature (London) 312, 533 (1984). 91. Pickett, S. D., Nowak, A. K., Thomas, J. M., and Cheetham, A. K., Zeolites 9,123 (1989). 92. Talu, O., Mol. Simul. 8, 119 (1991). 93. Schroder, K.-P., and Sauer, J., Z . Phys. Chem. (Leipzig)271, 289 (1990). 94. Demontis, P., Yashonath, S., Klein, M. L., J. Phys. Chem. 93, 5016 (1989). 95. Fitch, A. N., Jobic, H., Renouprez, A,, J. Phys. Chem. 90, 1311 (1986).

110

SIMON P. BATES AND RUTGER A. VAN SANTEN

96. Dwyer, J., and O’Malley, P. .I. Zeolires , 8, 317 (1988). 97. Bull, L. M., Henson, N. J., Cheetham, A. K., Newsam, J. M., and Heyes, S. J., 1. Phys. Chem. 97,11776 (1993). 98. Klein, H., Kirschhock, C., and Fuess, H., J. Phys. Chem. 98, 12345 (1994). 99. Kiselev, A. V., and Du, P. Q., J. Chem. Soc., Faraday Trans. 2 77, 17 (1981). 100. Germanus, A,, Karger, J., Pfeifer, H., Samulevic, N. N., and Zdanov, S. P., Zeolites 5, 91 (1985). 101. Auerbach, S. M., Henson, N. J., Cheetharn, A. K., and Metiu, H. I., J. Phys. Chem. 99, 10600 (1995). 102. Auerbach, S. M., Bull, L. M., Henson, N. J., Metiu, H. I., and Cheetham, A. K., J. Phys. Chem. 100, 5923 (1996). 103. Hseu, T., Ph.D. Thesis, University of Washington, Seattle (1972). 104. Burmeister, R., Schwarz, H., and Boddenberg, B., Ber. Bunsenges. Phys. Chem. 93, 1309 (1989). 105. Germanus, A., Karger, J., Pfeifer, H., Samulevic, N. N., and Zdanov, S. P., Zeolites 5, 91 (1985). 106. Snurr, R. Q., Bell, A. T., and Theodorou, A. T., J. Phys. Chem. 98,11948 (1994). 107. Jobic, H., Bee, M., and Dianoux, A. J., J. Chem. Soc., Furaday Trans. 185,2525 (1989). 108. VignC-Maeder, F., and Jobic, H., Chem. Phys. Lett. 169, 31 (1990). 109. Gusev, A. A,, and Surer, U. W., J. Chem. Phys. 99, 2228 (1993). 110. Metropolis, N., Rosenbluth, A. W., Rosenbluth, M. N., Teller, A. H.. and Teller. E., J. Chem. Phys. 21, 1087 (1953). I / / . Smit, B., and Siepmann, J. I., J. Phys. Chem. 98, 8442 (1994). 112. Frenkel, D., in “Proceedings of the Euroconference on Computer Simulation in Condensed Matter Physics and Chemistry, Como 1995” (K. Binder and G. Ciccotti. eds.), Chapter 7. Italian Physical Society, Bolugna, 1995. 113. Smit, B., Mol. Phys. 85, 153 (1995). 114. Woods, G. B., Panagiotopoulos, A. Z., and Rowlinson, J . S., Mol. Phys. 63, 49 (1988). 115. Woods, G. B., and Rowlinson, J. S., J. Chem. SOC., Faruday Trans. 2 85,165 (1989). 116. Hope, A. T. J., Leng, C. A., and Catlow, C. R. A., Proc. R. Soc. London Ser. A 424, 57 (1989). 117. Sanders, M. J., Leslie, M., and Catlow, C. R. A,, J . Chem. Soc., Chem. Commun., 1273 (1984). 118. van Tassel, P. R., Davis, H. T., and McCormick, A. V., Mol. Phys. 73, 1107 (1991). 119. van Tassel, P. R., Davis, H. T., and McCormick, A. V., Mol. Phys. 76, 411 (1992). 120. van Tassel, P. R., Davis, H. T., and McCormick, A. V.. J. Chem. Phys. 98, 8919 (1993). 121. Jameson, C. J., Jameson, A. K., Baello. B. I., and Lim, H.-M.. J. Chem. Phys. 100, 5965 (1994). 122. Jameson, C. J., Jameson, A. K., Lim, H.-M., and Baello, B. I., J. Chem. Phys. 100, 5977 (1994). 123. Derrah, R. I., and Ruthven, D. M., Can. J . Chem. 53,996 (1975). 124. Chmelka, B. F., Raferty, D., McCormick, A. V., de Menorval, L. C.. Levine. R. D., and Pines, A,, Phys. Rev. Lett. 66, 580 (1991). 125. Barrer, R. M., Papadopoulos, R., and Ramsey, J. D. F., Proc. R. Soc. London, Ser. A 326, 331 (1972). 126. Jameson, C. J., Jameson, A. K., Gerald, R., and de Dios, A. C.. J. Chem. Phys. 96, 1690 (1992). 127. Cerjan, C. J., and Miller, W. H., J. Chem. Phys. 75, 2800 (1981). 128. Van Tassel, P. R., Davis, H. T., and McCormick, A. V., Langmitir 10, 1257 (1994). 129. Santilli, D. S., Harris, T. V., and Zones, S. I . , Microporous Materials 1, 329 (1993).

MOLECULAR BASIS O F ZEOLITE CATALYSIS

111

130. Kono, H., and Takasaka, A,, J. Phys. Chem. 91,4044 (1987). 131. Vernov, A. V., Steele, W. A., and Abrams, L., J. Phys. Chem. 97, 7660 (1993). 132. Loriso, A,, Bojan, A. J., Vernov, A. V., and Steele, W. A,, J. Phys. Chem. 97,7665 (1993). 133. Tsaio, C., Kauffman, J. S., Corbin, D. R., Abrams, L., Carroll, E. E., and Dybowski, C., J. Phys. Chem. 95, 5586 (1991). 134. Barrer, R. M., and Sutherland, J. W., Proc. R. Sac. London, Ser. A 237,439 (1956). 135. Habgood, H. W., Can. J. Chem. 12, 2340 (1964). 136. Neddendriep, R. J., J. Colloid Interface Sci. 28, 293 (1968). 137. Cohen de Lara, E., and Nguyen Tan, T., J. Phys. Chem. SO, 1917 (1976). 138. &hen de Lara, E., and Kahn, R., J. Phys. (Paris) 42,1029 (1981). 139. Kahn, R., Cohen de Lara, E., and Moller, K. D., J. Chem. Phys. 83,2653 (1985). 140. Chkhaidze, E. V., Fomkin, A. A., Serpinskii, V. V., and Tsitsishvili, G. V., lzv. Akad. Nauk SSSR,Ser. Chim., 974 (1985). 141. Chiang, A. S., Dixon, A. G., and Ma, Y. H., Chem. Eng. Sci. 39, 1451 (1984). 142. Chiang, A. S., Dixon, A. G., and Ma, Y. H., Chem. Eng. Sci. 39, 1461 (1984). 143. Yucel, H., and Ruthven, D. R., J. Chem. SOC.,Faraday Trans. 2 76, 60 (1980). 144. Smit, B., and den Ouden, C. J. J., J. Phys. Chem. 92,7169 (1988). 145. Hyun, S. H., and Danner, R. P., AZChE Symp. Ser. 78, 19 (1982). 146. Yamazaki, T., Watanuki, I., Ozawa, S., and Ogino, Y., Nippon Kagaku Kaishi 8, 1535 (1987). 147. Yarnazaki, T., Watanuki, I., Ozawa, S., and Ogino, Y., Langmuir 4, 433 (1988). 148. Abdul-Rehman, H. B., Hasanain, M. A,, and Loughlin, K. F., Ind. Eng. Chem. Res. 29, 1525 (1990). 149. Rees, L. V. C., Briickner, P., and Hampson, J., Gas Sep. Purif: 5, 67 (1991). 150. Hufton, J. R., and Danner, R. P., AZChE J. 39, 954 (1993). 151. Titiloye, J. O., Parker, S. C., Stone, F. S., and Catlow, C. R. A,, J. Phys. Chem. 95, 4038 (1991). 152. Jackson, R. A., and Catlow, C. R. A., Mol. Simul. 1,207 (1988). 153. Bezus, A. G., Kocirik, M., Kiselev, A. V., Lopatkin, A. A,, and Vasilyeva, E. A,, Zeolites 6, 101 (1986). 154. Vetrivel, R., Catlow, C. R. A., and Colbourn, E. A,, J. Chem. SOC.,Faraday Trans. 2 85,497 (1989). 155. Snurr, R. Q., June, R. L., Bell, A. T., and Theodorou, A. T., Mol. Simul. 8,73 (1991). 156. Smit, B., J. Phys. Chem. 99, 5597 (1995). 157. Murad, S., and Gubbins, K. E., A C S Symp. Ser. 86, 62 (1978). 158. Papp, H., Hinsen, W., Do, N. T., and Baerns, M., Thermochim. Acta 82, 137 (1984). 159. Ding, T., Ozawa, S., and Ogino, Y., Zhejiang Daxue Xuebao, Ziran Kexueban 22, 124 (1988). 160. Widom, B., J. Chem. Phys. 39, 2808 (1963). 161. Widom, B., J. Phys. Chem. 86, 869 (1982). 162. Verlet, L., and Weis, J. J., Mol. Phys. 24, 1013 (1972). 163. Stach, H., Thamm, H., Fiedler, K., and Schirmer, W., in “Proceedings of the 6th International Zeolite Conference” (D. Olsen and A. Bisio, eds.), pp. 210-216. Buttenvorth, Guildford, 1984. 164. Thamm, H., Stach, H., Schirmer, W., and Fahlke, B., 2. Phys. Chem. (leipzig) 263, 461 (1982). 165. Thamm, H., Stach, H., and Fiebig, W., Zeolites 3, 95 (1983). 166. Stach, H., Lohse, U., Thamm, H., and Schirmer, W., Zeolites 6, 74 (1986). 167. Flanigen, E. M., Bennett, J. M., Grose, R. W., Cohen, J. P., Patton, R. L., Kirchener, R. M., and Smith, J. V., Nature (London) 271, 512 (1978).

112

SIMON P. BATES AND RUTGER A. VAN SANTEN

Smit, B., and Siepmann, J. I., Science 264, 1118 (1994). Maginn, E. J., Bell, A. T., and Theodorou, D. N., J. Phys. Chem. 99,2057 (1995). Smit, B., J. Phys. Chem. 99,5597 (1995). Smit, B., and Maesen, T. L. M., Nature (London) 374,42 (1995). Bates, S. P., van Well, W. J. M., van Santen, R. A., and Smit, B., J. A m . Chem. SOC. 118, 6753 (1996). 173. Bates, S. P., van Well, W. J. M., van Santen, R. A., and Smit, B., J. Phys. Chem. 100(44), 17573 (1996). 174. Bates, S. P., van Well, W. J. M., van Santen, R. A,, and Smit, B., Mol. Simul. (in press) (1997). 175. Stach, H., Fiedler, K., and Janchen, J., Pure Appl. Chem. 65,2193 (1993). 176. Richards, R. E., and Rees, L. V. C., Langmuir 3,335 (1987). 177. Dubinin, M. M., Rakhmatkariev, G. U., and Isirikyan, A. A,, Izv. Akad. Nauk SSSR, Ser. Khim. 10,2333 (1989). 178. Lohse, U., and Fahlke, B., Chem. Tech. (Leipzig) 35, 350 (1983). 179. van Well, W. J. M., Wolthuizen, J. P., Smit, B., van Hooff, J. H . C., and van Santen, R. A,, Angew. Chem. 107, 2765 (1995). 180. Freeman, C. M., Catlow, C. R. A,, Thomas, J. M., and Brode, S., Chem. Phys. Lett. 186, 137 (1991). 181. Ecklund, E. E., and Mills, G. A. Chem. Tech. (Heidelberg) 19, 549 (1989). 182. Shubin, A. A,, Catlow, C. R. A,, Thomas, J. M., and Zamaraev, K. I., Proc. R. SOC. London, Ser. A 446,411 (1994). 183. Dauber-Osguthorpe, P., Roberts, V. A,, Osguthorpe, D. J., Wolff, J., Genest, M., and Hagler, A. T., Proteins: Struct., Funct., Genet. 4, 31 (1988). 184. Schroder, K.-P., Sauer, J., Leslie, M., and Catlow, C. R. A., Zeolites 12, 20 (1989). 185. van Beest, B. W. H., Kramer, G. J., and van Santen, R. A,, Phys. Rev. Lett. 64,1955 (1990). 186. Pimental, G. C., and McClellan, A. L., “The Hydrogen Bond.” Freeman, San Francisco, 1960. 187. Wright, P. A,, Thomas, J. M., Cheetham, A. K., and Nowak, A. K., Nature (London) 318,611 (1985). 188. Pickett, S. D., Nowak, A. K., Cheetham, A. K., and Thomas, J. M., Mol. Sirnul. 2, 353 (1989). 189. Uytterhoeven, L., Dompas, D., and Mortier, W. J., J. Chem. SOC.,Faraday Trans. 88, 2753 (1992). 190. Mortier, W. J., Ghosh, S. K., and Shankar, S., J. Am. Chem. Soc. 108,4315 (1986). 191. The EEM method does not strictly belong in a section concerned with classical simulations. It is a method based on density functional theory that allows proper consideration of long range effects and parameters that are calibrated to non-empirical charges. Given the subject of this reference (benzene) it was included here. 192. Snurr, R. Q., Bell, A. T., and Theodorou, A. T., J. Phys. Chem. 97, 13742 (1993). 193. Snurr, R. Q., Bell, A. T., and Theodorou, A. T., J. Phys. Chem. 98,5111 (1994). 194. Shi, X., and Bartell, L. S., J. Phys. Chem. 92, 5667 (1988). 195. Jorgensen, W. L., Laird, E. R., Nguyen, T. B., and Tirado-Rives, J., J. Comput. Chem. 14,206 (1993). 196. van Koningsveld, H., van Bekkum, H., and Jansen, J. C., Acta Crystallogr, Sect. 3:Struct. Sci. B43, 127 (1987). 197. Fyfe, C. A,, Feng, Y . ,Grondey, H., and Kokotailo, G. T.,J. Chem. Soc., Chem. Commun., 1224 (1990). 198. Talu, O., Guo, C.-J., and Hayhurst, D. T., J. Phys. Chem. 93, 7294 (1993). 199. Thamm, H., Zeolites 7,341 (1987). 168. 169. 170. 171. 172.

MOLECULAR BASIS OF ZEOLITE CATALYSIS 200. 201. 202. 203. 204. 205. 206. 207. 208. 209. 210. 211.

212. 213. 214. 215. 216.

217. 218. 219. 220. 221. 222. 223. 224. 225. 226. 227. 228. 229. 230. 231. 232. 233. 234. 235. 236.

113

Wu, P., Debebe, A,, and Ma, Y. H., Zeolites 3, 118 (1983). Tsikoyiannis, J. G., and Wei, J., Chem. Eng. Sci. 46,255 (1991). Taylor, J. C., Zeolites 7, 311 (1987). Portsmouth, R. L., and Gladden, L. F., J. Chem. Soc., Chem. Commun., 512 (1992). van Santen, R. A., and Kramer, G. J., Chem. Rev. 95(3), 637 (1995). O’Malley, P. J., and Dwyer, J., Chem. Phys. Lett. 143, 97 (1988). Garrone, E., Kazansky, V. B., Kustov, L. M., Sauer, J., Senchenya, I. N., and Ugliengo, P., J. Phys. Chem. 96, 1040 (1992). Bates, S. P., and Dwyer, J., J. Phys. Chem. 97,5897 (1993). Meijer, E. L., van Santen, R. A,, and Jansen, A. P. J., J. Phys. Chem. 100,9282 (1996). Haw, J. F., Haw, M. B., Alvaro-Swaisgood, A. E., Munson, E. J., Lin, Z., Beck, L. W., and Howard, T., J. Am. Chem. SOC.116,7308 (1994). M~iller,C., and Plesset, M. S., Phys. Rev. 46,618 (1934). van de Runstraat, A., Stobbelaar, P. J., van Grondelle, J., Anderson, B. G., van IJzendoorn, L. J., and van Santen, R. A., in “Proceedings of the 11th International Zeolite Conference, Seoul, 1996” (M. Chon, S.-K. Ihm, and Y. S. Uh, eds.), Part B, pp. 1253-1260. Elsevier, Amsterdam, 1997. Sauer, J., Chem. Rev. 89,199 (1989). Meisel, S. L., McCullogh, J. P., Lechthaler, C. H., and Weisz, P. B., Chem. Tech. (Heidelberg) 6,86 (1976). Vetrivel, R., Catlow, C. R. A,, and Colbourn, E. A., J. Phys. Chem. 93,4594 (1989). Gale, J. D., Catlow, C. R. A,, and Cheetham, A. K., J. Chem. SOC.,Chem. Commun. 3, 178 (1991). Sauer, J., Kolmel, C., Haase, F., and Alrichs, R., in “Proceedings of the 9th International Zeolite Conference, 1992” (R. von Ballmoos, J. B. Higging, and M. M. J. Treacy, eds.), p. 679. Butterworth-Heinernann, Boston, 1993. Bates, S., and Dwyer, J., J. Mol. Struct. (Theochem.) 306, 57 (1994). Gale, J. D., Catlow, C. R. A,, and Carruthers, J. R., Chem. Phys. Lett. 216, 155 (1993). Haase, F., and Sauer, J., J. Phys. Chem. 98, 3083 (1994). Haase, F., and Sauer, J., J. Am. Chem. SOC. 117, 3780 (1995). Blaszkowski, S. R., Nascimento, M. A. C., and van Santen, R. A., J. Phys. Chem. 99, 11728 (1995). Zicovich-Wilson, C. M., Viruela, P., and Corma, A., J. Phys. Chem. 99, 13224 (1995). Anderson, M. W., and Klinowski, J., Nature (London) 339, 200 (1989). Anderson, M. W., and Klinowski, J., J. Am. Chem. SOC.112, 10 (1990). Anderson, M. W., and Klinowski, J., J. Chem. Soc., Chem. Commun. 918 (1990). Anderson, M. W., Barrie, P. J., and Klinowski, J., J. Phys. Chem. 95, 235 (1991). Kubelkova, L., Novakova, J., and Nedomava, K., J. Catul. lZ4,441 (1990). Mirth, G., Lercher, J., Anderson, M. W., and Klinowski, J., J. Chem. SOC.,Faraday Trans. 86, 3039 (1990). Bronnimann, C . E., and Maciel, G. E., J. Am. Chem. SOC.108,7154 (1986). Munson, E. J., Kheir, A. A,, Lazo, N. D., and Haw, J. F., J. Phys. Chem. %, 7740 (1992). Bosacek, V., J. Phys. Chem. 97, 10732 (1993). Pelmenschikov, A. G., van Santen, R. A., Janchen, J., and Meijer, E. L., J. Phys. Chem. 97, 11071 (1993). Pelmenschikov, A. G., Morosi, G., and Gamba, A,, J. Phys. Chem. %, 2241 (1992). Pelmenschikov, A. G., Morosi, G., Garnba, A., Zecchina, A,, Bordiga, S., and Paukshtis, E. A., J. Phys. Chem. 97, 11979 (1993). Car, R., and Parrinello, M., Phys. Rev. Lett. 55, 2471 (1985). Nusterer, E., Blochl, P. E., and Schwarz, K., Angew. Chem., lnt. Ed. Engl. 35,175 (1996).

114

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237. Shah, R., Payne, M. C., Lee, M.-H., and Gale, J. D., Science 271, 1395 (1996). 238. Shah, R., Gale, J. D., and Payne, M. C., J. Phys. Chem. 100, 11688 (1996). 239. The calculation is somewhat artificial since the acidic form of sodalite is not known experimentally. However, the soldalite cage building unit is found within other materials such as zeolite A. 240. Hutchings, G. J., and Hunter, R., Catal. Today 6, 279 (1990). 241. Sinclair, P. E., and Catlow, C. R. A,, J. Chem. SOC.,Faraday Trans. 92, 2099 (1996). 242. Ziegler, T., Chem. Rev. 91, 651 (1991). 243. Oliver, F. G., Munson, E. J., and Haw, J. F., J. Am. Chem. SOC. 96, 8106 (1992). 244. Blaszkowski, S. R., Nascimento, M. A. C., and van Santen, R. A,, J. Am. Chem. Soc. 118, 5152 (1996). 24.5. Blaszkowski, S. R., Nascimento, M. A. C., and van Santen, R. A,, J . Phys. Chem. B 101, 2292 (1997). 246. Bandiera, J., and Naccache, C., Appl. Catal. 69, 139 (1991). 247. Olah, G. A,, Prakash, G. K., Williams, R. E., Field, L. D., and Wade, K., “Hypercarbon Chemistry.” Wiley, New York, 1987. 248. Blaszkowski, S. R., Jansen, A. P. J., Nascimento, M. A. C., and van Santen, R. A,, J. Phys. Chem. 98, 12938 (1994). 249. Evleth, E. M., Kassab, E., and Sierra, L. R., J. Phys. Chem. 98, 1421 (1994). 250. Kramer, G. J., van Santen, R. A., Emeis, C. A., and Nowak, A. K., Nature (London) 363,529 (1993). 251. Blaszkowski, S. R., Nascimento, M. A. C., and van Santen, R. A., J. Phys. Chem. 100, 3463 (1996). 252. Dalla Betta, R. A,, and Boudart, M., J . Chem. SOC.,Faraday Trans. I 72, 1723 (1978). 253. Kramer, G. J., and van Santen, R. A., J. Am. Chem. SOC.115, 2887 (1993). 254. Kazansky, V. B., Frash, M. V., and van Santen, R. A., Catul. Lett. 28, 211 (1994). 255. Stefanadis, C., Gates, B. C., and Haag, W. O., J. Mol. Catul. 67, 363 (1991). 256. Lercher, J. A,, van Santen, R. A,, and Vinek, H., Catal. Lett. 27, 91 (1994). 257. Corma, A,, Miguel, P. J., and OrchillCs, A. V., J. Catal. 145, 171 (1994). 258. Senchenya, I. N., and Kazansky, V. B., Kinet. Catal. 28, 566 (1987). 259. Kazansky, V. B., and Senchenya, I. N., J. Catal. 119, 108 (1989). 260. Senchenya, I. N., and Kazansky, V. B., Catal. Lett. 8, 317 (1991). 261. Pelmenschikov, A. G., Zhanpeisov, N. U., Paukshtis, E. A., Malyshave, L. V., Zhdimirov, G. M., and Zamaraev, K. I., Dokl. Akud. Nuuk SSSR 393,915 (1987). 262. Kazansky, V. B., Senchenya, I. N., Frash, M. V., and van Santen, R. A,, Catal. Lett. 27, 345 (1994). 263. Collins, S. J., and O’Malley, P. J., Chem. Phys. Lett. 246, 555 (1995). 264. Collins, S. J., and OMalley, P. J., J. Catal. 153, 94 (1995). 265. Kazansky, V. B., Frash, M. V., and van Santen, R. A,, in “Proceedings of the 11th International Congress on Catalysis-40th Anniversary” (J. W. Hightower, ef al., eds.), pp. 1233-1242. Elsevier, Amsterdam, 1996. 266. Kazansky, V. B., Frash, M. V., and van Santen, R. A,, Appl. Catal. 146, 1 (1996). 267. Viruela-Martin, P., Zicovich-Wilson, C. M., and Corma, A,, J. Phys. Chem. 97, 13713 (1993). 268. Brand, H. V., Curtiss, L. A,, and Iton, L. E., J. Phys. Chem. 97, 12773 (1993).

ADVANCES IN CATALYSIS, VOLUME 42

NMR Studies of Solid Acidity JAMES F. H A W AND TENG XU The Laboratory for Magnetic Resonance and Molecular Science Department of Chemistry Texas A &M University College Station, Texas 77843

1.

Introduction

The prospects of obtaining a detailed molecular-level understanding of heterogeneous catalysts would appear to be best for solid acids ( I ) . Catalysis by solid acids often involves an appreciable concentration of reasonably uniform sites and restricted roles for defect structures. Furthermore, the great number of reaction studies, physical property measurements, and spectroscopic studies of solid acids provides a background (2-4) for the design and evaluation of further experimental and theoretical work. NMR spectroscopy has contributed to recent progress in the understanding of solid acidity. With the benefit of hindsight, this was inevitable. Nuclear magnetic resonance in condensed media (as opposed to beam experiments) was demonstrated in 1946 independently by groups led by Bloch ( 5 ) and Purcell (6). Commercial NMR instrumentation for chemical studies first became available in the mid-1950s and was common in the early 1960s. Physical organic and inorganic chemists quickly applied *H and 19F NMR to a variety of fundamental problems in solution chemistry, including chemistry in strongly acidic media. More or less simultaneously, semiempirical and ab initio methods were maturing to the point at which the conclusions of spectroscopic investigations of simple reactive species could be rationalized or sometimes predicted by suitable computations (7). Some of the most important reaction intermediates in organic chemistry are the carbocations. Neglecting some heteroatom-stabilized cations, most carbocations are divided into two groups: trivalent carbenium ions and fivecoordinate or higher coordinate carbonium ions. The parent carbenium ion is CH; , and the parent carbonium ion is CHS . Carbonium ions have been proposed as reactive intermediates in superacid-catalyzed reactions; however, they have never been directly observed in condensed media. In contrast, a variety of carbenium ions have already been prepared in superacidic media and been characterized by various physical methods, mainly I3C NMR spectroscopy (8). 115 Copyrlght 0 1998 by Academic Press All rights of reproduction in any form reserved 0360-0564198 $25 M)

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JAMES F. HAW AND TENG XU

The role of carbenium ions as reactive intermediates was not recognized until the 1920s, although there was evidence for the existence of carbenium ions (actually triarylcarbenium ions) as early as in 1901 (9).Meerwein (10) was the first to propose a cationic intermediate to explain the rearrangement of camphene hydrochloride to isobornyl chloride. In the 1930s the concept of cationic intermediates was extended by Ingold, by Hughes and co-workers ( I I ) , and by Whitemore (12).Heteroatom-substituted carbocations such as the acetylium ion 1 were prepared using boron trifluoride in the 1940s (I.?), but the direct and unambiguous observation and spectroscopic characterization of alkyl carbenium ions had to wait until the 1960s.

+ -

CH;-C=O

1

A 2

+

n 3

Olah and his associates unexpectedly observed the tert-butyl cation 2 by ‘H NMR spectroscopy while studying the decarbonylation reactions of pivaloyl complexes with strong Lewis acids such as SbF5 (14, 15). Olah’s work was initially met with skepticism by many of the established leaders of the physical organic community. Direct NMR observation of 13C(even for enriched samples) was technically very difficult in the mid-l960s, but it was possible in some cases to measure 13C chemical shifts indirectly by observing the effect of low-power irradiation at the 13Cfrequency on protons scalar coupled to the carbon of interest. The indirect (16) and later direct observation (17) of the 13Cspectrum of the tert-butyl cation in SbFS/ SOzprovided indisputable evidence of the existence of free carbenium ions as stable species in superacid solution media. The evidence for the tert-butyl cation was so strong because the 13Cshift measured (335 ppm downfield of TMS) (17) was far larger than any previously seen for neutral carbon in any type of bonding, even with heteroatoms. It took little time for Olah and his associates to recognize the extremely strong acidity and low nucleophilicity of SbF5, and a general method for preparing carbenium ions in SbF5 and related superacid solutions was soon developed. In short order a number of other alkyl cations, including the isopropyl cation 3, were observed (8).Olah’s contribution to carbocation chemistry was recognized by the 1994 Nobel prize in chemistry (18). The typical acids used in solution studies of carbenium ions are, by thermodynamically rigorous methodology, stronger than 100%sulfuric acid. Table I (19, 20) lists some of the common superacids as well as their approximate Hammett acidities. NMR studies of the products from organic

NMR STUDIES OF SOLID ACIDITY

117

TABLE I Hammett Acidity Values (Ho) for Selected Superacidsnsb (100% H r S 0 4 , the Threshold of Superacidity, Is Also Included) Superacid

Hammett Acidity (Ho)

100% H2S04 HClO4 CIS03H CF3S03H 1: 1 Oleum (H2SO4ISO3) 1:1 Magic acid (HSO,F/SbFS)C 1: 1 HF/SbFSd

-12.0 -13.0 -13.8 -14.1 -14.4 -22.8 ca. -30

a All the data were taken from Olah et al. (8) unless otherwise noted. The ratios in the table are in moles. From Gold et al. (19). From Olah et al. (8) and Brouwer and Van Doorn (20).

’

and inorganic precursors in superacid solutions soon provided a wealth of information about the structures, dynamics, and reactivities of electrophilic species with atypical coordination numbers or oxidation states. Because these species, especially three-coordinate carbenium ions, were often drawn in formal descriptions of reaction mechanisms in more conventional solutions, there was an inadvertent invitation to take quite literally the existence of free cations as true intermediates in a wide range of reactions and media. Of course, one can also observe cations and cation radicals in the gas phase where solvation is either not present or restricted to clusters with a few solvent molecules. But does the existence of free carbenium ions in certain solution media (the superacids) necessarily imply even a transient existence for these same cations in a wide range of media in which reactions proceed by formally similar routes? Physical organic chemists have long recognized that part of the power of the carbenium ion formalism lies in its modification to accommodate various degrees of solvation, ion pairing, and other deviations from the free-ion limit and that the detailed reaction coordinate for a process may contain, for example, a true intermediate in one medium but a transition state similar to that intermediate in a different medium. When a sophomore organic chemistry student goes to the blackboard and predicts that p-see-butyltoluene is a plausible product of the reaction of 1-butene and toluene in acid solution but that rn-n-butyltoluene is not plausible, that student has demonstrated knowledge of a formal system that will usually predict or rationalize a wide variety of organic transformations. Yet when one studies this trail of chalk marks, can it be claimed that

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it is a complete and accurate description of reality? In the early and mid parts of the 20th century, many of the details of organic reaction mechanism were not testable by either experimental or theoretical methods; now, at the end of this century, the tools are becoming available. One can with relative ease and confidence carry out first-principles calculations for small molecules that yield bond distances, angles, energies, and spectroscopic observables that more closely approximate a true description of nature than the intuition of experts. Using density functional theory or in some cases MP2 calculations, one may often treat plausible models of Bransted sites in zeolites. The object of catalysis is the acceleration of some but usually not all reactions by which a chemical system may more closely reach equilibrium. Thus, some reactions that proceed at a negligible rate without the catalyst, due to an appreciable activation barrier, proceed faster if a catalyst lowers the barrier, and selectivity is achieved by lowering the barriers to some but not all channels in a discriminating manner. (Selectivity of course also arises by differences in rates of diffusion and factors that may be cast in the preexponential term.) In the past, when one drew the mechanism of, for example, the oligomerization or rearrangement of an olefin in a zeolite, one was engaging in a purely formal exercise (much as the sophomore organic student) and was “graded” purely on conformance to accepted models. For example, if one pages through leading catalysis journals, one will see many examples that explicitly describe carbenium ion chemistry involving even primary cations. This practice is, at the very least, an oversimplification. The distinction between free-ion intermediates and more subtle chemistry is more than one of semantics. New catalyst systems are often designed to better express the property believed to be responsible for the activity of existing catalysts. If an existing zeolite must be a superacid that makes primary carbenium ion intermediates at will, then an even better catalyst might be an even stronger acid that makes such ions in an even greater concentration. Yet, if these ideas are wrong (as evidenced by the experience of some industrial laboratories that weaker acids are often better catalysts), then the rational design of better catalysts may be delayed. Indeed, if a catalyst works by lowering the activation energy of a reaction, will it do so by means of high-energy intermediates accessed through even higher energy transition states? Furthermore, rational syntheses of all sorts of pharmaceuticals and other useful products are being jilted for empirical methods in the youthful guise of combinatorial chemistry. Although there is some evidence that templating effects in zeolite systems can be understood and even controlled, the rapid progress in the discovery of new zeolite

NMR STUDIES OF SOLID ACIDITY

119

solid acids still has a very empirical component. Could this be in part due to confusion about the mechanism of reactions on solid acids? The aforementioned progress in NMR spectroscopy (and other experimental methods as well) in combination with computational chemistry has reached a stage in which an understanding of the most general features of organic reactions on solid acids may reasonably be expected in several years. This does not yet quite exist; this report is written in a time at which the sophisticated application of NMR and computational quantum chemistry to solid acids is becoming widespread, and specialists in various areas are suddenly having to evaluate evidence from other specialties. A frequent if not central question in solid acid studies is, “How can one best measure (or calculate) acid strength?” Although some of the probe molecule work described herein is interpreted in statements about relative acid strength, the authors believe that there has been too much emphasis on acid strength. The half-reaction concept of a zeolite Bronsted site dissociating to give a free proton and a conjugate base site ignores the specific interactions of the base molecule (or reactant) with the site. Acid dissociation constants in solution would appear to be defined without regard to the specific base that might be involved in a subsequent reaction step, but the idea works because the constants are defined for the reaction with a common base (the solvent). Nevertheless, workers will continue to debate various measurements and interpretations of solid acid strength, and to the extent that NMR measurements will be used in this regard, they can at least be interpreted with full appreciation of the caveats and oversimplifications in their common application. Indeed, the response of a given probe molecule is acceptable as a purely operational definition of “acidity,” provided one does not expect too much from such a figure of merit. To quote a good friend and colleague from his review (21) of a work by another good friend and collaborator, “There have been those who have wondered, often to themselves, and occasionally out loud, just where these various attempts to be precise and quantitative about the nature of acidity and basicity become too quixotic to be valuable. It is not impossible that the venerable sport of jousting at windmills is being practiced by the more zealous defenders of the various acid and base ‘religions’.’’ There has been too much theology associated with zeolite acidity, and the difference between what we know from direct evidence and what we infer should be kept in mind. The intent of this report is thus twofold: We first review those less familiar aspects of NMR spectroscopy, including its combination with computational quantum chemistry, and then we attempt to critically evaluate the application of NMR to solid acidity. We restrict ourselves primarily to aluminosili-

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cate zeolites and other solid acid systems (especially metal halides) against which the properties of zeolites may be meaningfully contrasted.

II.

The Chemical Shift

The NMR observable most commonly exploited in studies of solid acidity is the chemical shift. While some NMR observables (e.g., dipolar couplings) lend themselves to a more or less direct quantitative evaluation, the chemical shift must be interpreted. Changes in the I3C or I5N isotropic shifts of adsorbates are observed upon complexation with Bronsted sites, and the same is true of the 'H shift of the Bronsted site, but one is hard pressed to interpret such changes quantitatively in terms of a detailed structure of the adsorption complex or even the extent of proton transfer. Problems in interpretation also arise for the 'H shifts of Bronsted sites in the absence of adsorbates. One commonly sees claims of a correlation between the 'H shift of the uncomplexed Bronsted site and its relative acid strength. For example, two 'H shifts are frequently quoted for zeolite HZSM-5 (MFI), 2.0 pprn for external silanols and 4.3 ppm for the Bronsted site (22). For the manifestly more weakly acidic boron analog of this material, the latter shift is only ca. 2.2 ppm (23). One can find plenty of caveats that arise with these tempting correlations. Figure 1 shows 'H MAS NMR spectra of activated (dehydrated) HZSM-5 acquired over a range of temperatures (24). At lower temperatures, a shoulder on the downfield side of the 4.3 ppm signal resolves itself into a previously unexpected signal at 6.9 ppm. Should one claim that this site is much more acidic than the 4.3 pprn site? Such a claim would run counter to a body of evidence that the distribution of acid site strength on HZSM-5 is narrow. The problem becomes even more apparent when one considers how the 'H shift of the Bronsted site might be affected by interaction with adsorbates. A trivial example is the ammonium form of ZSM-5 for which the proton shift is ca. 6.5 ppm (25), yet the acidity is markedly less than the proton form of the zeolite. Figure 2 motivates a more subtle example of the perils of oversimplifying the relationship between chemical shift and acid-base properties. Acetylene and ethylene form n complexes with zeolite Bransted sites of approximately equal strength, yet they induce very different shifts for the Bronsted sites (26). These adsorbates, as well as aromatic rings, carbonyl groups, and other systems with multiple bonds, are characterized by an anisotropy (orientation dependence) in their magnetic susceptibilities. For acetylene, the sign of the magnetic anisotropy is such as to provide a downfield (deshielding) contribution in addition to that due to complexation. The magnetic anisotropic effect is reversed for ethylene and

NMR STUDIES OF SOLID ACIDITY

121

423 K 313 K

- JA

323K

,,

l h

A

A

296 K

273K

A

223K

r

"

1'

25 20

J fi

"

' I'

"

' I ' "'I""I"'

15 10

5

0

'T" "

I

"

''1

-5 -10 -15 PPm

FIG. 1. 300-MHz 'H MAS spectra of zeolite HZSM-5 over a range of temperatures. In addition to the well-known Bronsted sites (4.3 ppm) and external silanols (2.0 pprn), a broad shoulder at 296 K sharpened to a third peak at 6.9 ppm when the sample temperature was reduced to 123 K. The spinning speed was 3.5 kHz. (Reprinted with permission from Beck ef al. (24). Copyright 1995 American Chemical Society.)

FIG.2. A qualitative pictorial drawing showing the effect of the neighboring group magnetic anisotropy on the 'H chemical shifts of the Bronsted acid that is hydrogen bonded to acetylene and ethylene. The more shielded regions induced by the magnetic anisotropies of triple and double bonds are indicated with +u in the diagram. Upon forming hydrogen bonds with acetylene and ethylene, the 'H chemical shift of the Bronsted sites in HZSM-5 was shifted from 4.3 ppm to 7.3 and 6.3 ppm, respectively. This example highlights the perils of oversimplifying the relationship between chemical shifts and acid strength. (Reprinted with permission from White ef al. (26). Copyright 1992 American Chemical Society.)

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JAMES F. HAW AND TENG XU

it partially offsets the shift due to the electronic effect of complex formation. Thus, these two molecules bind to the Bronsted site very similarly, but an oversimplified interpretation of the 'H spectra would suggest little similarity. Since many NMR studies of solid acids entail the observation of a chemical shift change in a probe molecule, reactant, intermediate, or product upon complexation with an acid site, there is an opportunity to fundamentally impact the application of NMR to solid acids by making better use of this information. We therefore review the salient parts of the chemical shift in some detail. Since some of the more visible controversies in NMR studies of solid acids regard carbenium ions and related electrophilic species, this treatment will use such ions as examples wherever possible. It is well known that electrons in atoms and molecules can shield the nucleus in part from the applied magnetic field and so alter its resonance frequency, Y =

ly/27TJBo(l- o),

(1)

where v is the resonance frequency of the nucleus in the external magnetic field B o , y is the magnetogyric ratio of the nucleus, and 0- is the shielding constant. The chemical shift is related to the shielding constant by the appropriate reference to a chemical shift standard (for 13Cthis standard is tetramethylsilane, or TMS) as shown in Eq. (2). 8 = VTMS

- umnp~e

(2)

The nuclear shielding is more precisely the second derivative of energy with respect to magnetic field ( B ) and nuclear magnetic moment ( p ) (27).

where a, /3 = x, y, z. The components oi, form a second-rank tensor that is nonsymmetric (i.e., gij # ajJ.

This nonsymmetric second-rank tensor can be decomposed to a symmetric (i.e., 0-[,= u,J and an antisymmetric tensor through a symmetrization process (28). = gsymmetric +

Oant~symmetrlc

(5)

usymmetnc

=

r r

aantisymmetric =

(a,

( a , + @,V (Czx

123

NMR STUDIES OF SOLID ACIDITY

+ ayx)/2

uyy

+ @A12

(cry

+ @yz)/2

(a,

-(Gy- a y x ) / 2

0

-(ux7 - a,x)/2

-(ayz

-

(axr

+ a7x)/2

(ayr

+ U7Y)/2

Czz

ayM

(vxz - U , x P

(UY7 -

a7,)/2

0

-

a7,)/2

1

(7)

The antisymmetric tensor is generally not observable in NMR experiments and is therefore ignored. The symmetric tensor is now diagonalized by a suitable coordinate transformation to orient into the principal axis system (PAS). After diagonalization there are still six independent parameters, the three principal components of the tensor and three Euler angles that specify the PAS in the molecular frame. PAS

a,

aPAS

=

[:

0

:J 0

PAS

7

(8)

The absolute shielding of a nucleus is referenced to that of the bare nucleus. Although absolute shieldings are convenient for theoretical calculations, experimental measurements on bare nuclei are usually impractical. Thus shielding tensors are converted to chemical shift tensors to facilitate comparison with experimental measurements. The conversion from absolute shielding to chemical shift can be made for each component of the shielding tensor to give the chemical shift tensor SPAS.

According to convention, if the three diagonal elements in Eq. (9) are arranged in a decreasing numerical order, these are called S,, , S2,, and &3 (Sl, > h2> S33), respectively. The isotropic chemical shift Sisois defined as the average of the tensor trace: Siso= 4 Tr

PAS = +(al1+ & + &).

(10)

The chemical shift anisotropies (CSA) and the asymmetry factor (q)are defined in Eqs. (11)-(14) following the conventions of Duncan (29), Harris (30),and Fyfe (31).

124

JAMES F. HAW AND TENG XU

Figure 3 shows several examples of 13Cchemical shift tensors of carbenium ions oriented in the molecular frame as indicated. If the molecule or ion has a rotational axis of C3or higher order that passes through the nucleus in question, symmetry demands that the two components perpendicular to the C3or higher order rotational axes be identical; in such cases the identical components are designated 8,. The unique component that lies along the 633

4

633

611

Isopropyl Cation

Propionylium Ion O

H C

0

622

633

Benzoylium Ion

Benzenium Ion

FIG. 3. Chemical shift tensor orientations for the isopropyl cation, propionylium ion, benzoylium ion, and benzenium ion obtained from MP2 chemical shift calculations.

NMR STUDIES OF SOLID ACIDITY

125

rotational axis is designated 41. One obvious consequence is that 7 must be zero. Rapid motion may also induce time-averaged symmetry. For example, at 87 K the aromatic carbons of hexamethylbenzene show an asymmetrical chemical shift tensor (alI= 232 ppm, & = 151 ppm, and & = 24 ppm), while at 298 K the fast rotation of the molecule along its C6axis makes the chemical shift tensor symmetrical (6, = 189 ppm, 41 = 21 ppm) (32). The isotropic chemical shift is a single number derived from the six components of the symmetric tensor. Given that the isotropic averaging process reduces the three principal components to their average and discards the orientation of the PAS, apparent correlations between Si,, and properties related to solid acidity should be closely examined. The chemical shift tensor may be determined from a series of measurements on a suitable single crystal as a function of orientation. In this manner Waugh and co-workers measured the principal components and orientations of the chemical shift tensors for the aromatic carbons of durene (33). Grant and co-workers have performed similar measurements for a variety of compounds (34).Even when a large single crystal is not available, the orientation of the chemical shift tensor may sometimes be established if the labeled nucleus also has a well-defined dipolar coupling (35).Unfortunately, the latter must be fit to a model including the shift and dipolar tensors and their relative orientations; if the simulations show strong correlation of some parameters, a unique specification of the chemical shift tensor and dipolar coupling can be elusive (36). For polycrystalline samples, one is limited to measurement of only the principal components of the chemical shift tensor. For a single resonance these are afforded by the cusps or inflection points in the powder pattern (vide infra). Rapid, isotropic tumbling in solution averages the anisotropic (orientation-dependent) character of most nuclear spin interactions. NMR spectra of adsorbates can range from solution-like for weakly physisorbed species to solid-like for chemisorbed species with little internal motion. Thus, a general approach to surface NMR requires the capabilities of a modern solid-state NMR instrument to effect line narrowing and to deal with other consequences of long correlation times. Solids probes are generally less compact than solution probes, motivating the choice of a “widebore” magnet. A wider bore size is essential for reliable variable-temperature operation over a wide range, and as this review will demonstrate repeatedly, variable-temperature operation is essential for NMR studies of the diverse phenomena associated with solid acidity. The most frequently applied line-narrowing technique is magic angle spinning (MAS), which requires that the sample be rotated 1,000to >20,000 times a second while tilted at the “magic” angle. Rotation about the “magic” angle (54.7”, the angle between the body diagonal of a cube and

126

JAMES F. HAW AND TENG XU

any one of its faces) transforms the x, y, and I faces of a cube into each other. Several nuclear spin interactions are second-rank tensors and affect resonance frequencies to first order; the most important interactions are chemical shift anisotropy (CS) and dipolar coupling (DD). The resonance frequency (v) from first-order static perturbation theory for spin 1/2 nuclei is shown in the following equation (37): v

= vcs,iso + ( ~ C S , Z +

VDD,~)~'~(COS

(15)

where p is the magic angle and P2(c0s p) is the second-order Legkndre polynomials of cos 0. The chemical shift has isotropic (orientation-independent) components and anisotropic components, while the dipolar coupling tensor (D) has only anisotropic components (Tr D = 0). The rotation of the sample about the magic angle averages P2(cosp) to zero, thus removing the anisotropic components. In contrast to the spin 1/2nuclei, the resonance frequencies for quadrupolar nuclei (IZl > 1/2) are affected by both firstorder and second-order quadrupolar contributions. For an integer spin, the first-order broadening is the dominant contribution to the linewidth. However, for the central transition (- 1/2* + 1/2) of half-integer quadrupolar nuclei, the first-order quadrupolar broadening disappears because of the Hamiltonian's quadratic dependence on the angular momentum (38). Consequently the second-order contributions dominate the spectral broadening of the central transition. A perturbation treatment of the secondorder quadrupolar interaction gives the following equation for the central transition of half-integer quadrupolar nuclei (37): VQ

= vQ,iso

+ vQ,2p2(cOs6 ) + vQ,4p4(c0s p),

(16)

where P,(cos p) is the fourth-order Legbndre polynomial. Inspection of the functions PZ(cosp ) and P4(cos 0 ) reveals that there is no single angle at which the sample can be rotated so as to average both P,(cos p) and P4(cos p) to zero. To mechanically average the second-order quadrupolar interactions, double rotation (DOR) or dynamic angle spinning (DAS) is necessary (37). Frydman and co-workers recently demonstrated that second-order quadrupolar interactions can be removed for half-integer quadrupolar nuclei by the use of multiple-quantum MAS NMR (39).Quadrupolar nuclei are not a primary focus of this review, but Fig. 4 clearly demonstrates the use of 2H NMR to study the dynamics of zeolites. While many chemical investigations of solids seek to emulate solution investigations by averaging all orientation-dependent properties, this approach may be shortsighted in that it necessarily reduces the information conveyed by the full anisotropic tensor quantities. 2D MAS experimental techniques originally developed by Bax, Szeverenyi, and Maciel (40) and refined by Grant and co-workers (41) can provide isotropic chemical shifts

NMR STUDIES OF SOLID ACIDITY

127

CSX non-spinning

HY non-spinning 3.0 kHz

1

HZSM-5 non-spinning

3.0 !iHz f~""""""""'""'"l''''I''''1

50

40 30 20 10

0 -10 -20 -30 -40 -50 kH2

FIG. 4. 45-MHz 'H spectra of CD3I on zeolites CsX, HY, and HZSM-5. On CsX the adsorbate forms a framework-bound CD3group. Methyl iodide is much less reactive on acidic zeolites. It tumbles isotropically in H Y and shows restricted motion in HZSM-5. (Courtesy of Larry W. Beck.)

on one axis and resolved powder patterns on the orthogonal axis, but these experiments require prolonged acquisition times. Alternatively, one may perform slow-speed magic angle spinning to obtain peaks at the isotropic shifts and many sidebands at Si,, r+_ nv,l x,where u, is the rotational speed of the rotor, and VL the Larmor frequency. Numerical analysis of the sideband intensities by the method of Herzfeld and Berger (42) provides the principal components, all,&, and a, for each resolved isotropic shift. This method is often appropriate for one to four distinct isotropic shifts, it eliminates or attenuates weak dipolar couplings in species enriched in 13C, and it is time-efficient. Figure 5 reports two I3C spectra of a sample prepared by adsorbing acetyl chloride-l-13C onto TaCl, powder (43-45). The major product of this reaction is the acetylium ion. The upper spectrum is a powder pattern obtained without sample rotation; one notes a characteristic

128

JAMES F. HAW AND TENG XU 8, = 269 ppm I

,

6, = -75 ppm

PPm FIG.5. 75.4-MHz I3C cross-polarization spectra of the acetylium ion 1 on TaCI5 acquired at 298 K. The nonspinning spectrum shows a broad and axially symmetric powder pattern. The MAS spectrum shows that ca. 80% of the starting material (a~etyl-I-’~C chloride) was = 153 ppm), and the rest formed the donor-acceptor ionized to form the acetylium ion complex with TaCls (a, = 189 pprn). The nonspinning spectrum requires 20 times more scans to acquire than the MAS spectrum. The principal components of the I3C shift tensor of 1 were measured from both spectra, and the results are in very good agreement.

axially symmetric chemical shift tensor, and the two unique principal components (8, and 4,) can be measured from the inflection points. The lower spectrum was obtained over a much shorter time using slow-speed MAS because of line narrowing. The principal components of the chemical shift tensor are preserved in the relative sideband intensities. Furthermore, the MAS spectrum reveals the presence of a minority species that was not resolved in the static spectrum. We have prepared a number of acylium ions on metal halide powders and measured the principal components of their chemical shift tensors (43-45). Most of these cations have isotropic I3C shifts of 154 t 1 ppm. Often insensitivity to substituents results from opposite and offsetting variations in the principal components. The acetylium ion has an axially symmetric chemical shift tensor because of its C, rotation axis. When the symmetry is reduced from C,, to C,, or lower, a nonzero value may be observed. The sensitivity of chemical shift tensors to symmetry is a powerful means of probing molecular structure and temperature-dependent molecular dynamics. Multiple orders of spinning sidebands may offend those who seek “solution-like’’ NMR spectra of solids, but discarding most of the information inherent in the chemical shift is a considerable concession to aesthetics.

NMR STUDIES OF SOLID ACIDITY

129

We emphasize that attempts to correlate chemical shift and properties related to acidity should consider, whenever possible, the principal components of the chemical shift tensor and not simply their average.

111.

Computational Chemistry: A Tool for Spectral Interpretation

SHIFTAND MOLECULAR STRUCTURE A. THE CHEMICAL How does one relate chemical shift to structure? At first this may seem too obvious a question; one assigns the spectrum like a textbook problem and draws a picture of the “correct results.” Unfortunately, the recent literature of NMR studies on zeolites and other catalysts contains a few spurious reports that proved to be due to major errors in spectral interpretation, and this evidence highlights the sometimes deceptive characteristics of NMR spectra of surface species. Several features are common to many of the misassignments that reached publication. First, the results were interpreted in terms of a preconceived notion of the chemistry. In some cases, the chemical shifts were clearly not quite right, but this was rationalized. Second, the critical intermediate or product gave a very simple NMR spectrum (one or two peaks) and was part of a mixture of other products. Third, appropriate control experiments requiring an additional stable isotope chemical or a different isotopomer of the same reactant were not performed. Lastly, supporting evidence from spectroscopic techniques other than NMR was not obtained. If the species in question can be desorbed without further reaction, one can obtain mass spectra and even delineate label incorporation. Figure 6 demonstrates such an example of the use of GC-MS to identify reaction products and to support spectral assignments in NMR studies. Even if one’s skill (and luck) is so great as to never misassign a spectrum, the resulting cartoon structure of an adsorbate interacting with an acid site does not contain the same level of information as, for example, an X-ray single-crystal structure. What one would really like from the interpretation of an NMR study would be quantitative information about molecular structure, but unlike the case of a suitable set of diffraction data, one cannot invert chemical shift data to yield molecular structure. Measurements of dipolar couplings can oftentimes specify one or a few distances, but the relationship between chemical shift tensors and structure does not admit to a straightforward interpretation. Fortunately, these can be related using computational chemistry. The method proceeds as follows: a series of spectroscopic results are first collected using appropriate pulse sequences for spectral assignments and control experiments. The spectral features are then assigned and critical

130

JAMES F. HAW AND TENG XU

Linear Dimer

Cyclic Dimer

cis- and trans-isomers of cyclic dimer

I

I

1i.O

14.5

linear dimer

14.5

14.0

4

-

I

Time(min)

MassKharge

FIG. 6. An example showing the use of gas chromatography and mass spectrometry (GC-MS) for identifying reaction products on zeolites. Styrene was first reacted for ca. 30 min on activated zeolite HY at 298 K in a sealed glass tube; the sample was extracted using toluene as solvent, and the extracts were then analyzed with GC-MS. The total ion chromatogram of the extracts (a) shows three major peaks eluting at 13.83, 13.93, and 14.31 rnin, respectively. The peak at 14.31 min was readily identified as the linear dimer. The peaks at 13.83 and 13.93 min show equal ion intensity and nearly identical mass spectra (the mass spectrum of the 13.93-min peak is shown in (b)), and these were assigned to the cis and trans isomers of the cyclic dimer.

NMR STUDIES OF SOLID ACIDITY

131

experiments are designed in an attempt to falsify one’s own conclusions. This is the stage where one’s appreciation of the scientific method is most important. Now the model of the structure of the chemisorbed species is visualized using whatever graphical tools are available; this could range from a crude sketch to a three-dimensional computer model incorporating hundreds of framework atoms at their presumed crystallographic locations. At this point, part of the structure may even be optimized using empirical or semiempirical methods. An elaborate visualization strategy might be appropriate when there are important questions about steric constraints that cannot easily be addressed using more exact quantum mechanical treatments. No matter how nice the picture may look at this point, its validity is yet to be tested. The geometry must then be optimized using a reliable level of theory, and this may require some severe approximations. There are many considerations that critically affect the accuracy of a quantum mechanical calculation, the most obvious of which are the degree to which the atomic coordinates of the theoretical model relate to the experimentally observed compound, the flexibility of the basis set, and the extent to which electron correlation is accounted for (46). For chemisorption studies in zeolites, cluster models of appropriate size have been used to elucidate a wide range of zeolite chemistry (47-49). In some cases, long-range electrostatic interactions may need to be included in the Hamiltonian, and recent advances in methodology make the full periodic treatment of crystalline catalytic systems possible (50). Since charge separation is important in many chemisorption processes, large basis sets treating polarization and incorporating diffuse functions are often required. In addition, electron correlation is critical for the treatment of transition state structures (vide infra), hydrogen bonding, and chemical reactions. The perturbation theory of Maller and Plesset is the most straightforward way of improving a Hartree-Fock calculation; such treatments truncated at second order (MP2) usually provide much of the correlation effect at a reasonable computational cost (46). As an alternate means of including electron correlation, density functional theory (DFT) methods have been demonstrated to give results comparable to high-level post-Hartree-Fock calculations with considerably less computational cost (51-53). Recent studies have validated the DFT approach in the computation of geometries, vibrational frequencies, energetics, and other molecular properties. More important to catalytic applications, DFT has also been shown to give good results for reaction barriers and transition state geometries in proton transfer reactions (54). Figure 7 shows DFT calculations (BLYP/DNP) of the geometries of acetone and mesityl oxide on a cluster model of the Bransted site in zeolite

Hydrogen Q Carbon Oxygen Aluminum Silicon

Q

?3

Q

FIG.7. Energy-minimized structures of acetone and mesityl oxide adsorption complexes on a cluster model of HZSM-5 using DFT calculation. Note that in the case of the acetone complex, the proton remains bonded to the bridging oxygen, while in the case of the mesityl oxide complex, the proton is more fully transferred to the ketone. (Reprinted with permission from Haw er al. ( I ) . Copyright 1996 American Chemical Society.)

NMR STUDIES OF SOLID ACIDITY

133

H Z S M J (I).A number of papers have supported the use of such a cluster model for calculations (2, 47, 55). The calculated deprotonation energy of the cluster used in Fig. 7 is 280 kcal/mol, close to the experimental deprotonation energy (56) of 284 kcal/mol for the most acidic protons in zeolite HZSM-5. For comparison, the calculated proton affinities of acetone and mesityl oxide are 198.8 and 217.3 kcal/mol, respectively. A close inspection of Fig. 7 reveals that when the zeolite Bronsted site complexes with acetone, the proton remains bonded to the bridging oxygen, although the bond is slightly elongated. In the case of the zeolite complex with mesityl oxide, the proton is more fully transferred to the base molecule. In addition to the energy-minimized structure, an appropriate calculation also gives an estimate of the energy of the system. Since it is possible to determine local minima and miss the global minimum, one needs to look carefully for other geometries. If the energy calculation is performed with requisite accuracy, it is possible to compare it with any available experimental data on energetics and assess it with one’s intuition. Figure 7 suggests a number of NMR observables that should be diagnostic of, for example, the extent of proton transfer in complexes, and these observables can be approached computationally as well in an effort to validate the computed structure as a model of the physical system. One such observable is clearly the deuterium quadrupole coupling constant. Comparison of the two structures in Fig. 7 suggests that they may have significant differences in electric field gradient tensors at the nucleus of the Brgnsted site. Although ’H NMR has been applied to acid sites in uncomplexed zeolites by Root (57), its potential for first-principles studies of adsorption on acid sites is clearly underexploited. Since the discussion here focuses on the assignment of chemical shifts for species adsorbed to acid sites, chemical shift calculations are the next step. B. THEORETICAL CHEMICAL SHIFTCALCULATIONS It has become possible to carry out computations of the chemical shift tensor with sufficient accuracy that meaningful comparison with experiment should be expected (58). Of course, other second-order properties (dipole polarizabilities, harmonic vibrational frequencies, etc.) can also be calculated by straightforward ab initio methods. Calculations of the chemical shift have long been complicated by the problem of the gauge origin of the magnetic vector potential. For finite basis sets, calculated chemical shifts vary with changes in the Cartesian coordinates of the molecule. Several methods have been developed for dealing with the gauge problem, including Ditchfield’sGIAO method (gauge-including atomic orbitals) (59), Kutzelniggs’ and Schindler’s IGLO method (individual gauge for localized

134

JAMES F. HAW AND TENG XU

orbitals) (60-62), and Hansens' and Bouman's LORG approach (localized molecular orbitals with local origins) (63). In 1990, Wolinski, Hinton, and Pulay reformulated the GIAO approach in terms of analytic derivative techniques, making NMR calculations much more tractable (64). More recently, Gauss presented the GIAO-MP2 NMR equations (65), allowing the assessment of the effect of electron correlation on shieldings-a very significant contribution for carbenium ions and molecules with multiple bonds (66). GIAO methods have been further extended to include correlation effects at the fourth order (MP4) and coupled-cluster (CCSD) levels (67-69). MP4 and CCSD are more exacting methods of treating electron correlation and are required in the most strongly correlated cases (19F in F2, "0 in ozone, etc.). There has been much recent progress in the application of density functional theory (DFT) to the calculation of shift tensors, and several methods are presently available. The sum-over-state (SOS) DFT method developed by Malkin et al. (70) does not explicitly include the current density, but it has been parametrized to improve numerical accuracy. Ziegler and coworkers have described a GIAO-DFT method (71) that is available as part of the Amsterdam density functional package (72). An alternate method developed by Cheeseman and co-workers (73) is implemented in Gaussian 94 (74). For other details on chemical shift calculations and their application to chemical problems, the reader is directed to a symposium volume (58) and papers by Jameson and co-workers (75), Oldfield and co-workers (76), Grant and co-workers (77), Ellis and co-workers (78), and Farrar and coworkers (79). We have carried out a series of structure and shift calculations pertinent to problems in solid acidity ( I , 44, 80-84). These calculations range from obtaining structures, energies, and chemical shift tensors of isolated carbenium ions and other electrophilic species to attempting the same calculations for analogous species in contact with a cluster sufficiently large to model a specific acid site in a zeolite. In the simplest case, one is completely neglecting the medium, thus excluding effects such as incomplete proton transfer, ion pairing, more extensive coordination or complexation, and any long-range electrostatic effects. The computational cost of chemical shift calculations is strongly dependent on the number of basis functions used to approximate molecular orbitals and the manner and extent to which electron correlation is treated. In the case of isolated carbenium ions, one can do a full optimization at MP2/6-311G* or better and do single-point energy calculations at MP4 and perform vibrational calculations to obtain the zero-point energy (ZPE) corrections. With fewer atoms or in the presence of appreciable symmetry, one can go much higher in theory. For

NMR STUDIES OF SOLID ACIDITY

135

example, in preliminary calculations we have optimized the isopropyl cation (83) at the CCSD level. Although there have been many published reports of chemical shift calculations neglecting electron correlation, it is well established that electron correlation is important for carbenium ions, triple bonds, and other systems sensitive to correlation. Currently, Gauss' MP2 method (distributed as part of ACES 11) (85) is the most extensively validated methodology for carbenium ion chemical shift calculations. Treatment of electron correlation was absolutely essential to compute the 13Cchemical shift tensors for the acylium ions, reflecting not only the charge but also the triple bond. We optimized the geometries of the acylium ions and a number of neutral model compounds at MP2/6-311G* and calculated the chemical shift tensors using a variety of basis sets and treatments of correlation. Table I1 (86-88) compares computed and experimental results for methylacetylene, acetonitrile, CO, and the acetylium ion (44). Neglect of correlation (RHF) gives very large errors for the species in Table 11; in fact, the double-J basis set 6-31G* outperformed the triple-J set 6-311G* at the RHF level, a sign that the apparent agreement at lower level is fortuitous. When correlation is included at the MP2 level, the agreement with experiment is typically quite good and often impressive, as the data in Table I1 as well as analogous computations on six other acylium ions (44) demonstrate. In chemical shift calculations for acylium ions, it was not necessary to model the ionic lattice to obtain accurate values. These ions have tetravalent carbons with no formally empty orbitals, as verified by natural bond orbital calculations (89). Shift calculations for simple carbenium ions with formally empty orbitals may require treatment of the medium. We prepared the isopropyl cation by the adsorption of 2-br0mopropane-2-'~C onto frozen SbF5at 223 K and obtained a I3C CP/MAS spectrum at 83 K (83).Analysis of the spinning sidebands yielded experimental values of all = 497 ppm, &2 = 385 ppm, and & = 77 ppm. The isotropic I3C shift, 320 ppm, is within 1ppm of the value in magic acid solution (17).Other NMR evidence includes dipolar dephasing experiments and observation at higher temperature of a scalar doublet ('Jc-H = 165 Hz) for the cation center. Schleyer and co-workers performed MP2/6-311G(d,p) calculations of various geometries of the isopropyl cation, including single-point MP4 energies and zero-point energies, and predicted that the point group of the lowest energy configuration is C, (90). Our MP2/6-311f +G** optimizations confirm the important features of Schleyer's findings, and our C, structure is shown in Fig. 8. We calculated the chemical shift tensor of the isolated cation using the GIAO-MP2 implementation of Gauss (65),with tzp (C) and dz (H) basis

136

JAMES F. HAW AND TENG XU TABLE I1 Comparison of the Experimentally Measured "C Isotropic Chemical Shifts and Principal Component Data for Selected Molecules with Theoretically Calculated Values Using Various Levels of Theory

CH3C= CH

CH3CZ CH

CH,C=N

CEO

CH,CrO+

RHF/6-31G* RHF/6-311G* RHF/tzp/dz MPZ/tzp/dz Experiment a RHF/6-31G* RHF/6-311G* RHF/tzp/dz MP2/tzp/dz Experiment RHF/6-31G* RHF/6-311G* RHF/tzp/dz MPZ/tzp/dz Experiment * RHF/6-31G* RHF/6-311G* RHFltzpldz MPZ/tzp/dz Experiment RHF/6-31G* RHF/6-311G* RHF/tzp/dz MPZItzpldz Experiment

70 86 88 82 80 67 79 81 65 69 116 133 135 118 117 200 224 227 183 181 157 177 178 152 152

156 180 182 166 166 131 149 152 136 140 223 250 252 219 216 337 373 377 313 303 276 307 309 268 268

- 102 -102 - 102 86 - 93 -61 -61 -60 -71 - 74 -99 - 99 - 99 -83 - 78 - 74 -74 - 74 - 76 - 62 -83 - 83 - 83 -81 -80 ~

From Beeler et al. (86). From Ripmeester et al. (87). From Gibson and Scott (88).

sets, as well as using the DFT method at B3LYP/tzp/dz. The calculated isotropic shifts were in every case 8-10 ppm larger than those from experiment. The moderate discrepancy in the isotropic shift actually belies larger and partially offsetting discrepancies in the principal components. Distortions in the geometry of the cation did not improve agreement, but modeling the medium using a simple ion-pair model did. FHF- was used at MP2, and we were able to go as high as SbF; with DFT. The importance of modeling the medium has previously been suggested by other shift calculations for simple cations (91); it would be interesting to demonstrate examples of this for carbenium ions in zeolites. In favorable

NMR STUDIES O F SOLID ACIDITY

137

633

FIG.8. (a) MP2/6-311++G** optimized geometry of the isopropyl cation with selected bond distances in angstroms. The orientation of the chemical shift tensor for the cation is indicated. (b) The isopropyl cation complexed with FHF-. (c) Isopropyl cation complexed with SbF,. The computed chemical shift tensor of the isolated cation using the MP2/6311+ +G** geometry shows large discrepancy with experimental values, especially with the 611and &, components. The inclusion of anions ((b) FHF- and (c) SbFi) to model the medium effect improves the agreement between theory and experiment. (Reprinted with permission from Nicholas er al. (83). Copyright 1996 American Chemical Society.)

cases, it may be possible to evaluate the location and geometry of the cation based on calculations modeling perturbations by the lattice.

IV. Sample Preparation Techniques for in Situ NMR

Since a large portion of the in situ NMR work we will discuss has been performed by our group, we will give a brief description of our sample

138

JAMES F. HAW AND TENG XU

preparation techniques for NMR studies (92). Many solid acids are extremely moisture sensitive, and a trace of water will spoil the sample. Furthermore, many samples must be prepared at cryogenic temperatures (typically 153-233 K, and sometimes as low as 77 K) to avoid reaction of adsorbates before acquisition of NMR spectra. Regardless of the preparation and handling requirements, the sample must end up inside an MAS rotor with a sample compartment that is ca. 1 cm in length by 4-6 mm in diameter. The sample has to be transferred to the NMR probe without warming and must be well balanced for the rotor to spin at 1,000 to >20,000 times per second in the NMR probe. Since 1988 we have developed more than a half-dozen devices to accommodate the challenging sample preparation requirements, and we coined the acronym “CAVERN” to represent such devices (93). The first CAVERN device was introduced in 1989 in an in situ study of the reactions of propene on acidic zeolites (93) and soon became obsolete. A much improved version was described in Munson et al. (94), and Fig. 9 shows a diagram of such a device. This CAVERN device permits the rotor to be sealed or unsealed in the vacuum line to facilitate multiple reagent adsorption stages. Volatile adsorbates are introduced into the CAVERN from the vacuum line, but nonvolatile adsorbates can also be introduced by crushing a small glass bulb above the catalyst bed during the sealing step. These steps can be performed at cryogenic temperatures if necessary for subambient studies of very reactive species. One of the drawbacks of this CAVERN device is the occurrence of a nonuniform distribution of reactant on catalysts because adsorption occurs on a “deep bed” of catalyst packed in a MAS rotor. To overcome this problem, we developed several shallow-bed CAVERN devices (95), and Fig. 10 shows a version of one such design. A thin layer of catalyst is supported on a glass trapdoor, and the device is evacuated. A furnace is clamped in place so that the catalyst can be activated if necessary. The catalyst is cooled with a cryogen bath, and a controlled amount of adsorbate is introduced from the vacuum line. The trapdoor is raised, the loaded catalyst falls into the MAS rotor, and the seal is driven into place. Finally the cold, sealed rotor is manually transferred into the cold MAS probe. The added advantages of the shallow-bed CAVERN is that all manipulations can be carried out without using a glovebox in any step. An ultra-shallow-bed CAVERN device (Fig. 11) was developed to accommodate occasions where an extremely homogeneous distribution of adsorbates is required. Here the same amount of catalyst is loaded onto a much larger surface. After catalyst activation, reactants are introduced onto an ultrathin layer of catalyst bed ( c 0 . 5 mm) to achieve a homogeneous distribution of adsorbates on catalyst.

NMR STUDIES OF SOLID ACIDITY

1819 Ball Joint-

139

9

FIG.9. Diagram of the CAVERN apparatus used for both capping rotors following adsorption of reactants onto catalysts and the capping and uncapping of rotors to allow the sequential adsorptions of reactants onto catalysts. The diagram on the right shows an expanded view of the optional uncapping assembly.

V.

NMR Studies of Solid Acidity Using Probe Molecules

For the purposes of this review, we include probe molecules that can be either directly adsorbed or formed in situ. Examples of the latter case are carbenium ions and related electrophilic species. We will also consider several important heteroatom-substituted carbenium ions and heteroatom analogs of carbenium ions. Acylium ions are the intermediates in FriedelCrafts acylation reactions (96). The most simple, stable acylium ion is the acetylium ion, 1, and others are formally derived by replacing the methyl group with other R groups. Oxonium ions, formed by alkylation of an ether, resemble carbenium ions but are in fact onium ions in terms of their structures. Their stabilization requires strongly acidic media, and like carbenium ions, oxonium ions have been proposed as intermediates in a

140

JAMES F. HAW AND TENG XU

-Mechanism for Opening Trap Door and Driving Seal into Rotor

3 mm Stainless Steel Rod

I

-

35/25 Ball & Socket Joint

Catalvst

FIG. 0. Diagram of a trapdoor CAVERN allowing catalyst activation and shallow-be adsorptions at low temperature. The inset shows the expanded view of the trapdoor assembly.

variety of zeolite-catalyzed reactions. Analogous sulfonium and selenonium ions are also of interest. It is important to note that while a few of these species clearly form on zeolites, many carbenium ions can only be formed on much stroger solid acids. For example, we demonstrate that a free isopropyl cation does not form on zeolites at low temperature, even as a transient intermediate. As examples of probe molecules directly introduced onto solid acids, we consider ketones and aldehydes, amines and other nitrogen-containing compounds, phosphines, and molecules that form multiple hydrogen bonds. A. CYCLOPENTENYL CATIONS In 1987, Zardkoohi, Haw, and Lunsford (97) reported I3C MAS NMR spectra of zeolite HY treated with propene-2-I3C at room temperature. A

141

NMR STUDIES OF SOLID ACIDITY

To Vacuum line

t

t

Catalyst Bed FIG.11. Diagram of an ultra-shallow-bed CAVERN allowing catalyst activation and ultrashallow-bed adsorptions.

peak was observed at ca. 250 ppm, and this was interpreted in the context of the prevailing view that zeolites, as solid superacids, readily produce carbenium ions by protonation of olefins. The chemical shift of the isopropyl cation in SbF5 solution is ca. 321 pprn (I?, and the large discrepancy relative to the shift in the zeolite was interpreted with a model in which the cation coordinated to the conjugate base site of the zeolite to give a sort of alkoxyl species with substantial positive charge on the carbon. Propene on HY was, therefore, selected for the first in situ variabletemperature study using the CAVERN method. These experiments were carried out in early 1988 and published in 1989 (93). The central features of the CAVERN experiments were that the propene was introduced into the zeolite at cryogenic temperature and the sample was manipulated so that spectral acquisition could commence with an unreacted sample. Additional spectra were then acquired as the sample was slowly raised to room temperature. Detailed experiments of this sort were carried out for propene-2I3C and p r ~ p e n e - l - ' ~and C less extensive experiments were performed for pr0pene-3-'~C. These experiments showed, among other things, that the 250 ppm peak was formed coincident with a second peak at ca. 156 ppm and the relative intensities of these peaks were 2 : 1. A careful study of the literature of carbenium ion chemistry in sulfuric acid and superacid solution media suggested the assignment of these resonances (250 and 156 ppm) to alkyl-substituted cyclopentenyl cations similar to 4. The late Herman Pines (98) described the process of conjunct polymerization by which olefins reacted in sulfuric acid to give a complex mixture of

142

JAMES F. HAW AND TENG XU

4

5

6

products including cyclopentenyl cations. Den0 and Pittman (99,100) made careful measurements of the extent of formation of various cations from suitable precursors and concluded that the acid strength of cations like 5 was equivalent to that of ca. 37% HzS04.The remarkable stability of alkylsubstituted cyclopentenyl cations is due in part to the five-membered ring, which fixes the planarity of the n system and maximizes resonance stabilization. This effect can be appreciated by comparing 5 with 6 and 7.6 has an acid strength of ca. 50% H2S0, and 7 ca. 73% H2S04(99,100). This stability order is reflected in a wide range of solution reactions; for example, cyclohexanol derivatives react to form cyclopentenyl cations in 96% sulfuric acid (99).

I

8

Cyclopentenyl cations with structures similar to 5 were reported in a 1992 account of a high-temperature study of ethylene oligomers on HZSM-5 (101). Cyclopentenyl cations similar to 8 were prepared in HZSM-5 by more rational routes based on cyclic precursors (102). The principal components of the 13C shift tensors of the 1,3 carbons of 8 (S,,,= 250 ppm) were estimated to be Sll = 375 ppm, & = 290 ppm, and S,, = 86 ppm. The error bars on the principal components in these ions could be appreciable as a result of broad lines and low signal-to-noise. Figure 12 compares 13CMAS spectra of oligomerization products, including cyclopentenyl cations, prepared by adsorbing propene-2-13C on HY and c y ~ l o p e n t e n e - ' ~(random C~ label) on HZSM-5. In addition to reports from our group, the formation of cyclopentenyl cations has also been observed in NMR studies by others (103).

143

NMR STUDIES OF SOLID ACIDITY

A

PPm FIG.12. 75.4-MHz 13CCP/MAS spectra showing the formation of cyclopentenyl cations 8 from cycl~pentene-'~C~ (random) on zeolite HZSM-5 and 4 from pr0pene-2-'~Con zeolite HY.

B. THE ALLYL CATIONCONTROVERSY The allyl cation (9) is the simplest member of the class of resonancestabilized cations that includes the alkyl-substituted cyclopentenyl cations. But one could also say that the carbenium ion (CH;) is the simplest member of a class of cations that includes the trityl cation. In each case, 10 or so orders of magnitude of acidity separate the primitive member from its more elaborate derivatives.

9

10

The allyl cation has never been characterized as a persistent species in solution. If prepared, it would be the smallest carbenium ion universally accepted to have been formed in condensed media (this title has for many years been held by the isopropyl cation). No allyl cation derivatives, e.g., 10, were observed in a 1990 report of a CAVERN study of butadiene on HY and HZSM-5 (104).The same year, Hutchings reported a flow reactor study of allyl alcohol on HZSM-5 (105) and Gorte (106) reported a TPD study of allyl alcohol on the same zeolite. Hutchings and co-workers found that allyl alcohol had two reaction paths, one in which propanal was formed and another that formed hydrocarbons. The latter route was proposed to proceed through an allyl cation intermediate, but no claim for its persistence or spectroscopic observation was made or implied by Hutchings. Gorte

144

JAMES F. HAW AND TENG XU

interpreted his TPD results as evidence of an “allylic species” formed from dehydration on a Bronsted site. In early 1993, Haw and co-workers (107) reported in situ studies of allyl alcoh01-1-’~Con HZSM-5 and CsHX. No persistent carbenium ions were observed, but 1,3label exchange was observed for the alcohol on the weakly acidic zeolite. We interpreted this as support for a transient allyl cation, The low stability of this cation was invoked to explain the failure to observe this species as a persistent species. Downfield signals observed in that study were attributed to the formation of propanal. Later in 1993,Biaglow, Gorte, and White (BGW) (108) reported similar studies conducted at different loadings and assigned a downfield resonance (variously reported at 216 and 218 ppm by BGW) to the allyl cation in HZSM-5. In 1994,Buzek et al. (109) reported that the allyl cation could be prepared from a variety of halide precursors, e.g., allyl chloride or cyclopropyl bromide, on SbF, at cryogenic temperature, based on the infrared spectrum of the products. Those workers challenged BGW’s claim of the persistent allyl cation based on the discrepancy between the isotropic I3C shift in the zeolite and that calculated at MP2/6-31G*. This was one of the first examples of the use of chemical shift calculations to interpret (and in this case challenge) an NMR study of a species on a solid acid. A few months later, Farcasiu (110) independently disputed the allyl cation claim of BGW based on several lines of reasoning, including previously published evidence that the zeolite was not a sufficiently strong acid to stabilize this species as well as logical inconsistencies in BGW’s report. We (111) then reopened the allyl alcohol investigation (Fig. 13) and demonstrated unambiguously that BGW’s claim of the allyl cation and other unspecified secondary cations was based on a misinterpretation of the NMR spectrum of propanal and its oligomerization products. After acceptance of our manuscript, BGW submitted a manuscript (112) revising their interpretation, so the allyl cation controversy can be said to be concluded. Had the allyl cation formed in measurable concentration as a persistent species in the zeolite, this would be implied remarkable acid strength for zeolites, as some workers would still maintain. Instead, the experimental and theoretical rebuttals of the allyl cation claim highlight a paradigm shift regarding zeolite acid strength and function,

C. INDANYL CATIONS The clearest and most direct experimental evidence from zeolite studies for the existence of a free carbenium ion intermediate obtained by any means is summarized in Fig. 14 (113). We followed the dimerization of styrene to form cis and trans isomers in a series of low-temperature MAS NMR experiments. Identification of the dimeric products was further

145

NMR STUDIES OF SOLID ACIDITY

216

I

b

11"'1'"'1"'1'""'"'1''"1~'''1 250 200 150 100 50 0

300

-50

PPm FIG. 13. 90.5-MHz I3C CP/MAS spectra of allyl-I-13C alcohol (spectra a and b) and c and d) on zeolite HZSM-5. A11 the spectra were acquired at ambient p r ~ p a n a l - l - ~(spectra ~C temperature: (a) after heating for 5 h at 323 K; (b) after heating at 393 K for 0.5 h; (c) prior to heating; (d) after heating at 353 K for 0.5 h. The downfield spectral features in spectra c and d, e.g., the isotropic shifts and the number of resonances, are consistent with those derived from allyl-l-13C alcohol on HZSM-5 (spectra a and b), thus providing unambiguous evidence that the disputed resonance at 216 ppm is pr~panal-l-'~C. See Xu et af. ( I l l ) for a more detailed assignment of the resonances.

confirmed by GC-MS analysis following solvent extraction as has been discussed previously. We did not observe the styryl cation 11 as a persistent species in HZSM-5 or HY, although the observation of a small amount of ethylbenzene could imply a transient styryl species (not necessarily a free carbenium ion intermediate). Upon furtherheating of the di-

+

I

.p 0

11

12

13

146

JAMES F. HAW AND TENG XU Styrene-a-13C 298 K

523 K

FIG.14. 90.4-MHz 13C MAS spectra of styrene-a-13C reacting on zeolite HZSM-5. The methylindanyl cation 12 (251 ppm), formed through the cracking of the cyclic dirner (cf. Fig. 6) followed by intramolecular hydride transfer, was converted to naphthalene at 523 K. This is the clearest example of a free carbenium ion as a reaction intermediate on a zeolite.

meric products in HZSM-5, these species cracked to form the methylindanyl cation 12 in appreciable yield. 12 was indefinitely stable in the zeolite at room temperature, but upon further heating, it was converted to naphthalene. The methylindanyl cation, like the related alkyl-substituted cyclopentyl cations, is a surprisingly stable and well-characterized species in acidic solutions. The acid strength of 12 is comparable to 78% H2S04 (114).

We also obtained NMR spectra of the phenylindanyl cation 13 in the large-pore zeolite HY, and a small amount of cation 14 was formed on HZSM-5 by dimerization of a-methylstyrene. The dimethylphenyl carbenium ion 15 was not persistent on any zeolite we examined. This is not surprising if one reads the solution acid literature. 15 cannot be observed in 100% H2S04; stabilizing this cation requires 30% oleum (S03/H2S04) or other superacids (115). HZSM-5 is not a superacid. The observation of the much less stable styryl cation 11 was hailed as a triumph of superacid solution chemistry (116). If the styryl cation, with the phenyl group provid-

NMR STUDIES OF SOLID ACIDITY

147

ing resonance stabilization, is not persistent in HZSM-5, is it reasonable to assume zeolite reaction mechanisms in which even a transient existence is claimed for the isopropyl cation? Not at all, as we will show.

15

14

D. THETRITYL CATION Maciel (117) described the formation of the trityl cation 16 on silicaalumina in a 1984 symposium; in retrospect it is surprising that no one followed up on this work until much later. 16 easily forms from triphenylcarbinol and other precursors in solutions of modest acid strength. The early observation of such an “easy” cation had the unintended effect of suggesting that “real” carbenium ions would not so easily be detected in NMR studies of solid acids. Over a decade elapsed before we characterized the cut-butyl cation on AlC13 powder (43).

& 0

0

16

In 1995, Maciel and co-workers (218) synthesized the trityl cation in the supercages of zeolite HY by a clever application of Friedel-Crafts chemistry-13CC14 was reacted with an excess of benzene (Fig. 15). Maciel and co-workers carried out a number of spectroscopic and chemical manipulations that unambiguously demonstrated that the product was the trityl cation and that the cation was in the zeolite. A b initio calculations at various levels of theory predict that the point group of isolated 16 is D3 rather than D3h. It is interesting to speculate about the extent to which the zeolite environment might force the degree of twist away from the “gas-phase” equilibrium value.

148

JAMES F. HAW AND TENG XU

FIG.15. "Ship-in-a-bottle" synthesis of the trityl cation 16 from l3CCl4and benzene inside the HY supercage. (Reprinted with permission from Tao and Maciel (118).Copyright 1996 American Chemical Society.)

6 1 = 282 ppm l = 55 ppm

120 K after

20jppm

207 ppm

207 ppm

L

I@ I

I 8

5 8 PPm

FIG.16. 50.1-MHz 13CMAS spectra of ben~aldehyde-a-'~Cand benzene reacting on zeolite HY. The spectrum acquired at 120 K after the sample was heated at 448 K clearly shows an isotropic chemical shift at 207 ppm, consistent with the chemical shift of the trityl cation. Furthermore, the Herzfeld-Berger analysis of the sideband intensities reveals an axially symmetric tensor, thus providing unambiguous evidence for the trityl cation 16.

149

NMR STUDIES OF SOLID ACIDITY

We prepared the trityl cation by heating benzaldehyde-a-13C and unlabeled benzene in HY in situ in an NMR probe (Fig. 16) (119).This synthetic strategy was motivated by a 1995 literature report. Olah discussed the reactions of benzaldehyde with an excess of benzene in a variety of liquid acids at room temperature (120). In excess superacids such as CF3S03H, benzaldehyde was converted almost quantitatively to triphenylmethane. Yet, there was no reaction in 100%sulfuric acid. On the basis of experiment and theory, Olah proposed a diprotonated benzaldehyde as a superelectrophilic intermediate. In our studies of benzaldehyde and benzene on zeolite HY, there was no reaction at room temperature, but at 458 K the reactants were converted to appreciable yields of the trityl cation and diphenylmethane without 13Clabel scrambling. The observed reaction products in zeolites at high temperatures are reminiscent of reactions in superacid solutions at temperatures perhaps 400 K lower. The significance of the benzaldehyde study is that it provides one example (out of many) of a reaction requiring superacidity at moderate temperature but merely strong acidity at high temperature. The principal components of the trityl cation in zeolite HY are 6, = 282 ppm and 4,= 55 ppm. It is instructive to tabulate all of the 13Cprincipal component data measured for free carbenium ions in zeolites as well as for a few carbenium ions characterized in other solid acid media (Table 111). The zeolitic species, in addition to the trityl cation (119), are the substituted cyclopentenyl cation 8 (102),the phenylindanyl cation 13, and the methylindanyl cation 12 (113). Values for the tert-butyl cation 2 and methylcyclopentyl cation 17 (prepared on metal halides) (43, 45) are included for comparison. Note that the ordering of isotropic chemical shifts is reasonably consistent with one’s intuition from resonance structures; i.e., the more delocalized the positive charge, the smaller the isotropic shift. This effect is even more apparent in the magnitudes of the CSA. Since TABLE 111 Summary of Chemical Shift Parameters for the Trityl Cation 16, Phenylindanyl Cation 13, Cyclopentenyl Cation 8, Methylindanyl Cation 12, tert-butyl Cation 2, and Methylcyclopentyl Cation I7 Carbenium ion

S,, (PP4

16 13 8

282 310 375 359 468 484

12

2 17

282 286 290 320 467 459

55 68 86 76 54 51

228 230 246 263 414 420

0.00 0.16 0.52 0.22 0.00 0.09

207 221 250 251 330 331

150

JAMES F. HAW AND TENG XU

twisting of aromatic rings alters resonance overlap, the CSA value of the trityl cation might be strongly sensitive to the twist angle. Furthermore, note the r] values; the trityl cation assignment is supported by the zero value of r] which is required by a C, or higher symmetry axis passing through the nucleus in question. Neither the substituted cyclopentenyl cation nor the indanyl cations have these symmetry properties, and their r] values are nonzero. Comparison of 2 vs 17 shows that in favorable cases even subtle breaking of symmetry can be reflected in the chemical shift tensor. As demonstrated previously, r] can be very close to zero with some deviation from symmetry, but r] must be zero for C3and higher order rotational axes.

17

E. ARENIUM IONS We prepared the benzenium (18), toluenium (19), and ethylbenzenium (20) ions on solid HBr/AIBr3 (direct protonation) or AlBr3 (alkylation with the alkyl bromide) (45,84).The principal components of the chemical shift tensors of the ring carbons were measured and found to be in reasonable agreement with those calculated at GIAO-MP2/6-311G*. The benzenium cation exhibited temperature-dependent dynamics reflected in the averaging of both the CSA and the isotropic shifts (Fig. 17). A t low temperature, one observes isotropic peaks and sidebands for all four nonequivalent carbons. Rapid 1,2-hydrogen shifts and anisotropic ring motion at higher temperature result in a single exchange-averaged resonance with an axially symmetric tensor.

6

Q

i

+:

H

H H 18

19

20

There are three different isomers of the toluenium cation: ortho, meta, and para. MP2/6-311G* calculations predict that the para isomer predomi-

151

NMR STUDIES OF SOLID ACIDITY

6,,=206 298 K, CP

400

350

Q

300

250

200

150

100

50

0

-50

H

p

QHH

%

H H

ti

ti

6 l2..

FIG.17. 75.4-MHz 13CMAS spectra showing the dynamics of the benzenium ion 18. The spectrum at 77 K shows a static benzenium ion with four isotropic shifts. The other features in that spectrum are spinning sidebands. The benzenium ion 18 shows a single isotropic peak (147 ppm) and an axially symmetric tensor at 298 K because of rapid intramolecular hydrogen shifts.

nates at low temperature, and a mixture of the ortho and para isomers would be found near room temperature. In accord with this prediction, 13C MAS spectra of the toluenium cation prepared from benzene and methyl bromide on A1Br3 were strongly temperature dependent (Fig. 18). Note in particular the temperature dependence of the signals of the two carbons that change hybridization in the ortho-para isomerization reaction. This is a clear-cut example of chemical exchange with a temperature-dependent equilibrium constant. There is no evidence of protonation of simple aromatic rings in zeolites to form appreciable equilibrium concentrations of arenium ions. Figure 19 shows 13CMAS spectra of benzene in zeolite HY as a function of temperature (81).No ring protonation is reflected in the 13Cshift, although temperature-dependent ring motion is observed. DFT calculations at the BLYP/

152

JAMES F. HAW AND TENG XU

243 K

273 K

A

1.

I

A ,

1 ’ ~ 1 1 1 1 1 1 1 1 1 1 1 ~ 1 1 ’ 1 ’ 1 1 250 200 150 100 50

para

0

ortho

FIG.18. 75.4-MHz 13CBloch decay MAS spectra showing the dynamics of the toluenium ion. The cation was synthesized by reacting br~momethane-’~C with benzene-I3C6on AIBri at 233 K. The spectrum at 213 K shows all the peaks for the toluenium ion at 32 (methyl), 50 (C-4), 178 (C-3), 139 (C-2), and 201 pprn (C-1). The peak at 129 pprn was the unreacted benzene-13C6.A t 243 K, the peaks were much sharper, and the 138 and 50 ppm peaks were NMR “invisible.” A t 273 K, the spectrum shows two extra peaks at 128 and 73 ppm. All these spectral features are rationalized by the chemical exchange between the para and ortho isomers

DNP level found no evidence for a stable benzenium cation in contact with a cluster modeling the zeolite conjugate base site. We were able to locate a transition state for benzene H/D exchange as shown in Fig. 20, which is similar to the transition state for methane H/D exchange on zeolites (222). These transition states clearly show that hydrogen exchange is a concerted process. One recent study that used H/D exchange to probe the mechanism of toluene disproportionation proposed a mechanism involving arenium ions (Fig. 21) (222). This mechanism is attractive in many respects, and it is probably much closer to the truth than an alternate mechanism involving methylation of the zeolite. Nevertheless, it would be interesting to deter-

153

NMR STUDIES OF SOLID ACIDITY

298 K

rll

350

I I

I11III

300

1 1 I I

250

I

I I I I

200

I11III

150

I I I I

100

I11

50

1 1 1 1 1 I I

I

-50

0

PPm FIG.19. 90.4-MHz I3CMAS spectra of benzene-% on zeolite HY, showing the temperature-dependent dynamics of benzene inside zeolite HY. Note that benzene is not protonated by zeolite HY.

Hydrogen

8

Carbon

lr!

Aluminum

3

1

Silicon

L

3

W

FIG.20. Optimized BLYPlDNP transition state for benzene HID exchange on a zeolite cluster model. (Reprinted with permission from Beck et al. (81). Copyright 1996 American Chemical Society.)

154

JAMES F. HAW AND TENG XU

t

I

FIG.21. A reaction mechanism proposed by Xiong et al. to rationalize the toluene disproportion reaction on zeolites. The benzyl cations and benzenium ions were proposed as reactive intermediates for this zeolite-catalyzed, high-temperature reaction. (Reprinted with permission from Xiong et al. (122).Copyright 1995 American Chemical Society.)

mine whether the species depicted in Fig. 21 are free carbenium ion intermediates or transition states with stabilization afforded by the zeolite lattice.

IONS F. ALKYLCARBENIUM Alkyl carbenium ions have no resonance or neighboring group stabilization due to heteroatoms. However, with the exception of the simplest member (CH;), they do tend to have significant stabilization due to neighboring cr bonds. As shown in Fig. 8, the isopropyl cation has C, rather than Czvsymmetry due to hyperconjugation-one C-H bond is bent up on one methyl group and one down on the other (83). This effect can be even more apparent for primary cations; for example, the ethyl cation has a C, structure with a bridging proton involved in a three-center, two-electron bond (Fig. 22) (123). The ethyl cation has never been characterized in condensed media of any sort. Our attempt to form the ethyl cation from ethyl bromide in frozen SbF, resulted in the bromonium ion (C2H5)2Br+. In situ NMR studies have shown that ethylene is relatively unreactive on acidic zeolites and oligomerizes in the vicinity of room temperature to afford more or less linear species without persistent carbenium ions (101). Free primary cations are frequently drawn in the catalysis literature as “intermediates” in zeolite-catalyzed reactions; this practice has little experi-

NMR STUDIES OF SOLID ACIDITY

155

FIG.22. CCSD (tzp, spherical) optimized structures for the bridged and “classical” forms of the ethyl cation used in NMR calculations. (Reprinted with permission from Ajith Perera et al. (123). Copyright 1995 American Chemical Society.)

mental or theoretical justification and should be recognized as an extraordinary claim in dire need of extraordinary justification. The isopropyl cation is the simplest secondary carbenium ion. When prepared in frozen SbFS, I3C label scrambling goes to completion at ca. 253 K (45).This is the result of a well-known process involving a protonated cyclopropane intermediate or transition state (124). The barrier for this scrambling process is low. We have observed complete scrambling of the I3C label for isopropyl bromide on A1Br3 (Fig. 23), a system that does not produce the isopropyl cation as a persistent species (45). The observation of label scrambling provides compelling evidence that the dominant species on the AlBr3 catalyst, the donor-acceptor complex 21, is in equilibrium with a small amount of isopropyl cation formed as a transient species (Scheme 1). Our study of ~ropene-l-’~C, -2-13C, and -3-13C reacting on zeolite HZSM-5 clearly shows that the isopropyl cation is not formed in measurable concentration as a persistent species (45). Furthermore, there is no label scrambling of the 2 position, although 1,3-label scrambling is facile on the zeolite. This strongly argues against a free isopropyl cation-even as a transient intermediate! At low temperature, the equilibrium structure of propene is a 7~complex 22 with the Bronsted site. This mode of coordination

156

JAMES F. HAW AND TENG XU

a: 2-Bromopropane-2-13C on AIBr3 at 298 K

b: 2-Bromopropane-2-13C on SbF5 at 233 K

c 2-Bromopropane-2-13C on SbF5 at 253 K

r 1 1 1 1 i 1 1 1 1 i 1 1 1 1 i 1 1 1 1 l 1 1 1 1 l 1 1 l 1 i ’ 1 1 1 i 1 1 1 1 i 1 1 1 1 1

450 400 350

300

250

200

150

100

50

0

PPm FIG.23. 75.4-MHz I3CMAS spectra of 2-br0mopropane-2-’~Creacting on Lewis superacids AIBr, and SbFS. Spectrum a shows the formation of an adsorption complex (89 ppm for CZ and 29 ppm for the methyl carbons) with AlBr, at 298 K. Note that I3C label scrambling from Cz to Cl is complete in the adsorption complex as indicated by the 2: 1 intensity ratio of the 29 pprn to 89 ppm peaks. Spectrum b shows a completely I3C scrambled isopropyl cation (320 and 52 ppm) and a partially 13Cscrambled adsorption complex with SbF5 (92 and 26 ppm) at 233 K. (c) Upon raising the temperature, the adsorption complex was completely converted to the isopropyl cation.

is apparently general; it has also been reported for ethylene (26),acetylene (26), isobutylene (125), and benzene (26, 81). 1,3-Label scrambling in propene occurs through the framework-bound alkoxyl intermediate 23, which is seen at 89 ppm in studies of propene-2I3C on HZSM-5 (45). Analogous alkoxyl species have been reported on acidic zeolites from the reactions of certain alcohols and acetylene. In the latter case, structure 24 was proposed as the product obtained by heating acetylene on HZSM-5 (126). Evidence for this assignment included the formation of acetaldehyde as a hydrolysis product. Acetylene also reacts

SCHEME 1

157

NMR STUDIES OF SOLID ACIDITY

22

23

24

25

on metal oxides, including MgO, but in this case the product seems to be the acetylide species 25 (127). The authors would not be overly surprised if the zeolite species was also acetylide-perhaps formed on extraframework material-but the zeolite-acetylene system has not been reexamined in light of the metal oxide studies. Alkoxyl species form very readily from the reaction of alkyl halides on alkali, alkaline earth, transition metal, and lanthanide exchanged zeolites (128, 129). The more basic the zeolite, the more readily the reaction proceeds. Alkyl halides have been used to generate methoxyl, ethoxyl, isopropoxyl, and rert-butoxyl species on metal-exchanged zeolites. The mechanistic significance of alkoxyl species in zeolite acid catalysis is not in general clear; in some reactions they may be true intermediates, and in others mere spectators. Figure 24 reports 13CMAS spectra of the tert-butyl cation (43) and the methylcyclopentyl cation 17 (45)on the solid metal halides AIC13 and A1Br3; the asymmetry parameters, CSAs, and isotropic shifts (Table 111) are unambiguous for the species indicated. Repeated attempts in various laboratories to observe the tert-butyl cation as a persistent species in a zeolite have thus far been unsuccessful. DetaiIed theoretical work wiII be required to determine whether or not the tert-butyl cations are local minima (i.e., true intermediates) on typical reaction pathways in zeolites. The ease with which these cations form in true superacids (liquid or solid) should be contrasted with the history of negative observations in zeolites. G. CARBONIUM IONS Sommer (130, 1 3 0 ~ and ) Hall (131) have independently described the low-temperature H/D exchange of isobutane on zeolites. The traditional mechanism involves a five-coordinate carbonium ion intermediate; yet no exchange occurred for the methine position, and this is inconsistent with a carbonium ion. This surprising result was explained by Sommer with a reaction sequence beginning with hydride abstraction by an unknown route

158

JAMES F. HAW AND TENG XU

-

I ""I 1 ' " I " '1.P ' I 1 ' ' 1 I " " I " ' 1 I 550 500 450 400 350 300 250 200 150 100 50

0

PPm

273 K

550

41

332

450

350

1 250

150

50 PPm

2

-50

FIG.24. 75.4-MHz 13C MAS spectra showing the formation of the rert-butyl cation and the methylcyclopentyl cation on AICL and AIBr, . The methylcyclopentyl cation was synthesized by an intermolecular hydride transfer reaction as shown in the figure.

to generate either a framework alkoxyl or carbenium ion in equilibrium with the olefin (Fig. 25). Deuteration then occurs following the Markovnikov rule, and hydride abstraction from another isobutane provides the methyl-deuterated product and chain propagation. Do these results also suggest that five-coordinate carbonium ions are not essential to explain alkane cracking? The evidence is mixed. Kazansky and van Santen (132) reported low-level calculations and found a metastable carbonium ion (CH,-H-CH,t) formed from ethane and a zeolite BrGnsted site, but this species was so high in energy that it did not appear to be thermally accessible. More extensive work by van Santen (133) shows, however, that the transition states leading from this species do not relate to ethane cracking! Blaszkowski, Nascimento, and van Santen (134) found other transition states for ethane cracking (Fig. 26) that are similar to carbenium ions albeit with stabilization from the lattice.

159

NMR STUDIES OF SOLID ACIDITY - HD

RHD+

superacid - D+ R"H

-

AH

RH R= Acid-Base bifunctional

-H, AD

FIG.25. Two possible pathways proposed by Sommer and co-workers explaining the observed H/D exchange of the alkanes. Pathway 1: the carbenium ion RA (nondeuterated) is formed by protolytic cleavage of a C-H bond (via a carbonium ion intermediate). Pathway 2: the olefin R =is formed by acid-base bifunctional dehydrogenation. (Reprinted with permission . 1995 American Chemical Society.) from Sommer ef al. ( 1 3 0 ~ )Copyright

TS 1

TS 2

FIG.26. Geometries of the two transition states for zeolite-catalyzed cracking reactions of ethane obtained from a density functional study on a zeolite cluster model. The distances are in angstroms and angles in degrees. (Reprinted with permission from Blaszkowski et al. (134). Copyright 1996 American Chemical Society.)

160

JAMES F. HAW AND TENG XU

It is clear that suggestions of carbonium ion “intermediates” in zeolite catalysis should be viewed with caution pending extension and confirmation of the work by Sommer, van Santen, Kazansky, and others.

IONS H. ACYLIUM Acylium ions can be formed in superacid solutions from carboxylic acids and acyl halides (8). They are among the best characterized carbenium ions, and single-crystal X-ray structures of a number of them have been determined as BF, , SbF;, or TaC1; salts (135-139). Solid-state NMR characterization of these species on A1Br3 and other solid superacids was described earlier in this review.

. ‘ . -) , , \ \

26

27

Acylium ions are intermediates in the Koch-Haaf reaction (140), in which an alcohol (or the equivalent) reacts with CO in the presence of a strong acid to form a carboxylic acid. Stepanov et al. (141) observed the Koch-Haaf reaction of tert-butyl alcohol and CO on Bronsted sites in HZSM-5. Although the expected carboxylic acid product formed, the acylium ion intermediate was not observed. Our efforts to prepare persistent acylium ions in zeolites by any of several routes have all been unsuccessful. However, the Koch-Haaf reaction can be used to make these ions on true solid superacids. For example, we reacted 2-chlor0-2-rnethylpropane-2-’~C and I3CO on AlC13. Formation of the acylium ion 26 occurred quantitatively, and the carbon-carbon bond was verified from the observation of (through-space) homonuclear dipolar coupling, which is reintroduced under the rotational resonance condition (44, 45). We measured the principal components of the 13Cchemical shift tensors for a number of acylium ions, formed either directly from the acyl chloride or by in situ synthesis. Almost all of these cations have isotropic I3Cshifts of 154 2 1 ppm; yet the principal components show appreciable variations that relate to symmetry and structure. For example, the measured asymmetry parameters (7)of 1 and 26 are both zero, as expected for point group C3“,but the benzoyl cation 27 has a nonzero asymmetry parameter, reflecting its lower Czvsymmetry. The CSA magnitude of 27 is significantly smaller than that of 1 and 26, as can be rationalized by theoretical predictions of resonance hybrids.

161

NMR STUDIES OF SOLID ACIDITY

I. CHALCOGENONIUM IONS Trialkyl cations of oxygen, sulfur, and selenenium (and less commonly tellurium) are onium ions by virtue of their having one extra valence, although these three-coordinate ions invite comparison with carbenium ions. These ions (e.g., 28, 29, and 30) are not planar (as opposed to the analogous carbenium ions), but rather are pyramidal. These species have a well-characterized solution acid chemistry; concentrated sulfuric acid is required for observation of the trimethyloxonium cation, but the sulfonium and selenonium ions require less strongly acidic solutions.

28

30

29

The trimethyloxonium cation was proposed to be an intermediate in one of the proposed mechanisms of carbon-carbon bond formation in MTG chemistry on HZSM-5 (Scheme 2) (142). Chang and co-workers demonstrated that this cation could be ion exchanged into HZSM-5 from cold aqueous (CH&O+BFi (143). Munson and Haw (144) demonstrated that

H

I

-.

p,.

Si

CH30CH3, CH30H

"**

'?+ CH3

Basic Site

.

CH3OCH3 or CH30H

SCHEME2

CH3

CHI-

I

162

JAMES F. HAW AND TENG XU

the adsorption of excess dimethyl ether on HZSM-5 produces 27 and a stoichiometric amount of methanol (Scheme 3). The analogous cations 29 and 30 readily formed in HZSM-5, HY, and even the much weaker acid HX by the reaction of methanol with dimethyl sulfide or dimethyl selenide (145). Formation of the trimethylselenonium ion in zeolite HZSM-5 was established unambiguously by the observation the 77Se resonance (22 ppm downfield of dimethyl selenide). Extensive evaluation of the reactivity of methanol in the presence of trimethylchalcogenonium ions found no evidence of a rate enhancement in MTG chemistry, and this mechanism (Scheme 2) was rejected.

J. ALDEHYDES AND KETONES NMR has been extensively applied to carbonyl compounds in acidic zeolites and other solid acids. The unshared pairs of electrons on the oxygen can interact with either Bronsted or Lewis sites, and aldol condensation reactions are commonly observed. Acetone was first studied on a zeolite by Boshcek and co-workers (146) followed by Haw and co-workers (147) and later by Gorte and co-workers (148). The conclusion of an earlier acetone paper of Gorte and co-workers (149) was that acetone forms a static complex on the Bronsted site of HZSM-5 at room temperature, but this claim was later revised (150)upon the realization that molecular motion in the complex is not halted except at appreciably lower temperatures. Munson and Haw (151)reported the first in situ NMR study of acetaldehyde in a zeolite. Figure 27 shows 13C spectra of this species reacting on HZSM-5 in the presence of water to form crotonaldehyde with high selectivity (an example of aldol condensation). We later reported a very detailed study of the aldol reactions of acetone and cyclopentanone on various zeolites (Scheme 4) (147). Dimerization of acetone followed by dehydration gives mesityl oxide (31), and the 13C isotropic shifts of this conjugated ketone are strongly dependent on state of protonation. Farcasiu and Ghenciu (152, 153) have reported extensive measurements of the 13Cshifts of 31

SCHEME 3

163

NMR STUDIES OF SOLID ACIDITY

c 10 min at 353K

I

0

30 rnin

I

20 min at

393 K

I

20 min at

1

250

1

1

200

1

1

150

1

1

100

1

1

50

1

1

0

1

1 -50

PPm FIG.27. 50.1-MHz I3C MAS NMR spectra of a~etaldehyde-I,2-'~C~ on HZSM-5 that had been saturated with water. Crotonaldehyde (199,160,135, and 19 ppm) was produced selectively at 353-393 K. (Reprinted with permission from Munson and Haw (151). Copyright 1993 VCH Verlagsgesellschaft.)

in various solution acids and proposed that the 13Cchemical shift difference between the a and 6 carbons of mesityl oxide be used as a measure of acid strength. We interpreted the mesityl oxide shifts in the context of the Farcasiu proposal. Literal comparison of the zeolite and sulfuric acid shifts (Table IV) (154) suggests that the acid strength of common zeolites does not exceed 70% H,SO,. The aldol reaction chemistry of propanal (111) was also elucidated in the course of resolving the ally1 cation controversy. Measurements of the static 13C line shape or sideband intensities of acetone on many solid acids at room temperature underestimate the chemical shift anisotropy due to motion, but the principal components of the chemical shift tensor can be accurately measured at reduced temperature. Table V reports these data for acetone on a wide variety of Bronsted and Lewis acids (43, 45); note that the largest contribution to the isotropic shift is all.The shift induced by A1C13 and other Lewis acids is rationalized by

164

JAMES F. HAW AND TENG XU

diacetone alcohol

isophorone

mesityl oxide (31)

phorone

trindane SCHEME 4. Summary of the reaction sequences of acetone and cyclopentanone on acidic zeolites observed by in situ NMR

32

165

NMR STUDIES OF SOLID ACIDITY TABLE IV I3C Isotropic Chemical Shifts of Mesityl Oxide in Various Media at 298 K 13Cchemical shift’ Media

Hob

100.0% H2S04 70.2% H2S04 HZSM-5 HZSM-5 HY 62.5% HzS0.q HY CIZCHCOOH HX NaX csx CH3COOH CSY CSZSM-5 CDC13

-12.0 -5.96

Acetone loading (equiv)

0.25 2.0 0.1 -4.90 1.2 -0.75 0.50 0.50 0.50 0.0

1.o 2.0

Carbonyl

C,

C,

210 211 211 210 210 211 210 205 207 206 205 201 205 204 198

122 122 122 122

203 191 190 188 188 183 175 163 162 162 162 158 157 157 155

122 123 124 123 123 124

124

Reference

Reported in ppm relative to TMS. Hammett acidity value reported in Farcasiu and Ghenciu (152). Initial loading as a fraction of the number of Bronsted acid sites. Measurements were performed after ca. 20% conversions to mesityl oxide.

the resonance structure 32.Shifts due to Lewis acids or Lewis acid-enhanced Bronsted complexation in zeolites have been observed in dealuminated HY (255) and Beta (256).

K. NITROGEN-CONTAINING COMPOUNDS We consider amines, imines, nitriles, and nitro compounds. Ammonia itself is discussed in a section on species with behavior dominated by multiple hydrogen-bonding interactions. Ellis and co-workers published a 13CMAS study of ethylamine on solid acids in 1981 (157).Maciel and Haw (158, 159) published NMR studies of pyridine as a probe molecule on solid acids in 1983. We have recently begun to reexamine the I5N spectrum of pyridine on zeolites and other solid acids (160). At low temperatures pyridine is remarkably sensitive to the kinds of acid sites present. Figure 28 shows 15N spectra of pyridine adsorbed on HY samples before and after dealumination. Dealumination in this case seems to make four kinds of Lewis sites distinguishable by NMR of adsorbed pyridine, suggesting pyridine as a good candidate for

166

JAMES F. HAW AND TENG XU

TABLE V Summary of Chemical Shift Parameters for the Donor-Acceptor Complex between Acetone and a Variety of Lewis and Brgnsted Acids Donor-acceptor complex

Temp (K)

4,"

41

&2

&3

(PP4

(PP4

(PP4

(PP4

CSA (PP4

CH3COCH3 CF3CH(OH)OCF3 HZSM-5 HZSM-5 HZSM-5' HZSM-5 HZSM-5 30% oleum

87 83 298 153 125 93 298 83

208 221 223 223 223 223 224 246

279 309 291 310 312 314 285 392

265 276 271 276 269 266 280

79 78 115 89 82 86 121 66

193 215 162 20 1 211 206 155 270

MgC12 ZnC12 TaCl, AH3 SCTf3 AIBr3 AICI3 TaF, SbF5

83 83 113 83 113 83 193 83 83

221 230 237 238 239 243 245 248 250

319 335 363 375 370 396 387 392 404

267 277 264 266 262 265 256 270 280

77 80 84 73 86 69 93 81 67

216 225 230 248 230 261 228 250 274

263

From Biaglow et al. (150). From Biaglow et al. (149).

studying the structures of solid acids. A number of proposals have been put forward for the structures of Lewis sites in partially dealuminated zeolites. If some of these can be successfully modeled using existing or emergent methodologies, it will be possible to test whether the model predicts the experimental I5N data as well as other observables such as the "A1 quadrupole coupling constant. Grey and Vega (161) have demonstrated the use of the TRAPDOOR experiment for measuring the 27Alquadrupole coupling constant in zeolites and applied it to studies of trimethylamine in zeolites. In other spectroscopic work, Fripiat and co-workers (162) used REDOR and various other NMR methods to characterize acid sites in zeolites treated with ammonia. Ernst and Pfeiffer (163) have reported a I3C MAS NMR study of the reactions of methanol and ammonia to make methylamines in zeolite HZSM-5. We reported NMR evidence of the synthesis of cyclopentylamine from cyclopentanol and ammonia on zeolite CsX (102). BosAcek and co-workers (146) reported spectroscopic evidence that acetone reacts with ammonia on zeolite HZSM-5 (MFI) to form the corre-

167

NMR STUDIES OF SOLID ACIDITY

h I 1

200

1 1 1 1

100

0

I l l 1

1 1 1 1

1 1 1 1

1 1 1 1

1 1 1 1

1 1 1 1

-100 -200 -300 -400 -500 -600

PPm FIG.28. 36-MHz 15N CP/MAS spectra of pyridine-15N on zeolite HY. The experimental conditions were all the same for (a) and (b), except that sample b was extensively dealuminated by increasing the activation temperature to 550°C (400°C for sample a). Both spectra were acquired at 77 K to prevent chemical exchanges on the NMR time scale. (a) The single resonance at -176 ppm as well as its associated sidebands indicates protonation of pyridine by the Br@nsted sites. (b) In addition to the protonated pyridine, four additional resonances at -68, -88, -116, and -140 ppm are also seen, indicating complexation of pyridine with different extraframework Lewis sites.

sponding imine. We later studied the reactions of benzaldehyde with ammonia in zeolite HY and found that a surprisingly stable gem-amino-hydroxy tetrahedral intermediate was formed (164). This tetrahedral intermediate dehydrated upon mild heating to form the imine (Scheme 5). Figure 29 shows 'H + 13Cand 'H -+I5N cross-polarization and 'H -+"N + 13Cand 'H + I3C -+ 15N double-cross-polarization spectra of the intermediate in the zeolite. Figure 30 shows a qualitative picture for the reactants, intermediate, and products in the zeolite. The multiple hydrogen bonds rationalize the unexpected stability of the tetrahedral intermediate and the imine product in a zeolite. Nitromethane was proposed as a probe molecule of basic sites on catalyst surfaces. The nitro form is in equilibrium with a small amount of the aci tautomer. Kheir and Haw (165) demonstrated that the aci tautomer was deprotonated on MgO and CaO to form chemisorbed aci anion (Scheme OH

d

H 0

Nq

NH

c0

SCHEME 5

0

168

JAMES F. HAW AND TENG XU

l ' ' ' ~ l ' ' ' ' l ' ' ' ' l ' ' ' ' l ' ' ' ' l ' ' ' ' i

300

250

200

150

100

50

0 PPm

I ' " ' I " " I " ' ' I ' ' " l " ~ ' l ' ' ' ' l " " l " ' ' i

0

-50

-100

-150

-200

-250

-300

-350

-400 PPm

FIG.29. 13C and 15N MAS spectra of ben~aldehyde-a-'~C reacting with ammonia-"N in zeolite HY to form an imine by way of a tetrahedral intermediate. Spectral editing by means of "double cross polarization" affords unambiguous assignment of the intermediate's spectra. The imine 15Nhas a short TIPand is absent from the 'H + I3C + ''N spectrum. (Reprinted with permission from Haw ei al. ( I ) . Copyright 1996 by the American Chemical Society.)

6). This process has been partially modeled theoretically (166). A small amount of aci anion was also observed on basic CsX zeolite (167). Many nitro-substituted aromatic compounds are Hammett indicators. Although the Hammett formalism cannot be rigorously applied to solids, the degree of proton transfer from a zeolite to any given base molecule can provide insight into the nature of acidity. "F and "N NMR were = - 12.4), p - fluoroaniline recently applied to p-fluoronitrobenzene (HBH+ (HBH+ = ca. +2.4), p-nitrotoluene (HBH+ = -11.4), and p-nitroaniline ( H B H t = +0.99) in zeolites HY and HZSM-5 (82). The 19FMAS spectra in Fig. 31 show that p-fluoronitrobenzene is not protonated in the zeolites even though a structurally similar molecule that is a much stronger base @-fluoroaniline) is protonated. Figure 32 shows energy-minimized geometries for these probe molecules on zeolite clusters; it is clear that p-fluoroaniline is not protonated and that proton transfer to p-fluoroaniline is assisted by multidentate coordination to the zeolite lattice, as demonstrated previously for ammonia (168, 169). Acetonitrile is the most frequently studied nitrile compound in zeolites.

169

NMR STUDIES OF SOLID ACIDITY

B

A R'

C FIG.30. A qualitative picture showing the reactants, the tetrahedral gem-amino-hydroxy intermediates, and the products of reactions of acetone and benzaldehyde with ammonia on zeolite HY. (Reprinted with permission from Xu et al. (164). Copyright 1995 American Chemical Society.)

Figure 33 shows plots of the nitrile carbon shift as a function of temperature for acetonitrile (<1equiv of loading) on various acidic zeolites (80).These data suggest larger shifts at moderate or high temperatures for zeolites that are commonly regarded as more acidic. There is no generally accepted mechanism for these remarkable temperature-dependent shifts. Haw et al. (80)proposed a model involving bending and proton transfer based on restricted Hartree-Fock (RHF) calculations. van Santen and co-workers

Aci anion SCHEME 6

170

JAMES F. HAW AND TENG XU

1.0 M HC1 0.1 M HCI

'

':CDC13 E$)&.D3

PPm

HY

I""I"' 0

-150

-200

FIG.31. Representative 188-MHz I9FNMR spectra ofp-fluoroaniline (top) andp-fluoronitrobenzene (bottom) obtained in zeolites HY and HZSM-5. Spinning sidebands are denoted by asterisks. Spectra were acquired (several thousand scans) using magic angle spinning (4rnrn rotors), cross polarization (2 ms), and proton dipolar coupling. (Reprinted with permission from Nicholas et al. (82). Copyright 1995 American Chemical Society.)

(170) have reported more detailed theoretical calculations of the acetonitrile-zeolite potential surface and explained features of the vibrational spectrum. Further theoretical work on acetonitrile to better explain the NMR behavior would be valuable. Munson (271) carried out experimental work on HCN in zeolites in parallel with the acetonitrile experiments. The results were consistent with oligomerization. Although this process by itself is less interesting from a catalytic standpoint, the oligomerization of HCN and its reaction with formaldehyde remain of interest in the study of prebiotic syntheses of monomers that may have combined to form the first self-replicating systems (172).

L. PHOSPHINES

The use of 31PNMR of trimethylphosphine to study acid sites in zeolites was first reported by Lunsford and co-workers (273,174). The 31Pisotropic

NMR STUDIES O F SOLID ACIDITY

171

FIG. 32. (a) Selected internal coordinates and partial charges for p-fluoronitrobenzene adsorbed on the zeolite model. p-Fluoronitrobenzene is weakly hydrogen bonded to the zeolite, but not protonated. Values in parentheses are those of an isolated, neutral p-fluoronitrobenzene molecule. (b) Optimized geometry for p-fluoroaniline adsorbed onto the zeolite. In this case, we started with the proton on the zeolite; the optimization resulted in protonation of the adsorbate. (Reprinted with permission from Nicholas et al. (82). Copyright 1995 American Chemical Society.)

172

JAMES F. HAW AND TENG XU

180 170 -160 --

a

2 4.;a

150 -140 --

-

HZSM-5

+W Steamed --.)Beta €I-

+W - 2 . 5 5 ~wth.)

I HY-19 ( a ~

A H-Mordenite

d

110 250

350

450

550

Temperature / K FIG.33. Summary of temperature-dependent 13C isotropic chemical shifts for the nitrile carbon of acetonitrile on various zeolites. Loadings were in the range of 0.3-0.9 equiv. (Reprinted with permission from Haw et al. (80). Copyright 1995 American Chemical Society.)

shift of adsorbed trimethylphosphine is very different for complexes with BrQnsted and Lewis sites, and the resonance of the former shows wellresolved 'H scalar coupling in the absence of proton decoupling (Fig. 34). The observation of scalar coupling ('Jp-n = 550 Hz) is unambiguous evidence for a complex with at least some covalent bonding between 'H and 31Pthat is static on a time scale greater than the reciprocal of the coupling constant. Fripiat and co-workers ( I 75) have also applied trimethylphosphine to zeolites. Maciel and co-workers (176) have described the use of trialkylphosphines on solid acids.

THATBOTHDONATE AND ACCEPT HYDROGEN BONDS M. ADSORBATES Some of the more controversial experimental and theoretical work on adsorbates in zeolites concerns simple molecules such as ammonia (168, 169,177, 178),methanol (55,178-184), and water (178,185-191), which by virtue of their great facilities for hydrogen bonding have very complicated potential energy surfaces. The ammonia system is probably best understood, at least for the case of a single ammonia molecule interacting with a Brmsted site. It is now agreed that proton transfer to the ammonia is energetically unfavorable if the product is either free NH,' or ammonium

173

NMR STUDIES OF SOLID ACIDITY

P 200 160 120 80

40

0

-40 -80 -120 -160 -200 -240

ppm

FIG.34. 31P MAS spectra of trimethylphosphine on zeolite HY. Both proton-decoupled (a) and nondecoupled (b) spectra are shown. The nondecoupled spectrum clearly shows the scalar coupling of '3P-'H ('.Ip-~= 550 Hz), providing unambiguous evidence for the protonation of trimethylphosphine by the Brmsted site. (Reprinted with permission from Rothwell et al. (173). Copyright 1984 American Chemical Society.)

d$b

H

8 8

I I I

J

8

J b

a

I

I

\ \ \ 8

b

FIG. 35. Models showing the interaction of NH3 with zeolite Bronsted acidic sites: (a) bidentate-coordinated NH4 and (b) tridentate-coordination of NHb . (Reprinted with permission from Teunissen et nl. (168). Copyright 1993 American Chemical Society.)

174

JAMES F. HAW AND TENG XU

singly hydrogen bonded to the zeolite lattice. Consideration of doubly and triply hydrogen-bonding species (268, 269) shows that proton transfer does occur. One can get a sense of this from the models shown in Fig. 35. As this review is written, there is disagreement regarding the detailed potential surfaces for methanol and water on zeolites. Some of the reports on these systems found evidence for proton transfer, even for one or two adsorbate molecules, while proton transfer has been predicted only for larger adsorbate clusters or not at all in other studies. It would be inappropriate to prejudge the outcome of these ongoing debates. ACKNOWLEDGMENTS This work was supported throughout by the National Science Foundation and more recently by the Basic Energy Sciences program of the U.S. Department of Energy. We thank Dr. John B. Nicholas at Pacific Northwestern Laboratory for his many helpful suggestions. The contributions of former and present graduate students and postdocs are gratefully acknowledged. REFERENCES 1. Haw, J. F., Nicholas, J. B., Xu, T., Beck, L. W.. and Ferguson, D. B.. Acc. Cherrz. Rex

2. 3. 4. 5. 6. 7.

8. 9. 10. 11.

12. 13. 14.

29, 259 (1996). van Santen, R. A., and Kramer, G. J., Chem. Rev. 95,637 (1995). Corma, A,, Chem. Rev. 95, 559 (1995). Farneth, W. E., and Gorte, R. J., Chem. Rev. 95, 615 (1995). Bloch, F., Hansen, W. W., and Packard, M., Phys. Rev. 69, 127 (1946). Purcell, E. M., Torrey, H. C., and Pound, R. V., Phys. Rev. 69, 37 (1946). See, for example, Saika, A., and Slichter, C. P., J. Chem. Phys. 22, 26 (1954): Pople, Chem. . Phys. 37,53 (1962): Mol. Phys. I, 301 J. A,. Discuss. Furuday Soc. 34,7 (1962); .I (1964): Stevens, R. M., Pitzer, R. M., and Lipscomb, W. N., J. Chem. Phys. 38, 550 (1963); Jarneson, C. J., and Gutowsky, H. S., ibid. 40, 1714 (1964); Karplus. M., and Pople, J. A., ibid. 38, 2803 (1963); Yonezawa, T., Morishima, I.. and Kato, H.. B d . Chem. SOC.Jpn. 39, 1398 (1966); Alger, T. D., Grant, D . M., and Paul. E. G.. J. A m . Chem. SOC. 88, 5397 (1966); Cheney, B. V., and Grant, D. M.. ibid. 89, 5319 (1967); Grant, D. M., and Cheney, B. V., ibid., p. 5315; Velenik, A.. and Lynden-Bell. R. M., Mol. Phys. 19, 371 (1970); Ellis, P. D., Maciel, G. E.. and McIver. J. W., J. A m . Chem. SOC.94,4069 (1972). Olah, G . A,,Surya Prakash, G. K., and Sommer, J., “Superacids.“ Wiley, New York. 1985. Norris, J. F., Am. Chem. 1. 25, 117 (1901); Kehrniann. E., and Wentzel. F.. Ber. Dfsch. Chem. Ges. 34, 3815 (1901). Meerwein, H., and van Emster, Ber. Dtsch. Chem. Ges. 55, 2500 (1922). Ingold, C. K., “Structure and Mechanism in Organic Chemistry.” Cornell University Press, Ithaca, NY, 1953, and references therein. Whiternore, F. C., J. A m . Chem. Soc. 54,3274 (1932); Chem. Eng. News 26, 668 (1948). Seel, F. Z., Anorg. A&. Chem. 250,331 (1943);Meerwein, H., in “Houben-Wey Methoden der Organischen Chemie,” (E. Mueller et al., eds.), 4th ed. Thieme. Stuttgart. 1952. Olah, G. A., Tolgyesi, W. S., Kuhn, S. J . , Moffatt, M. E.. Bastien, I. J.. and Baker, E. B., J. Am. Chem. SOC. 85,1328 (1963).

NMR STUDIES OF SOLID ACIDITY

175

15. Olah, G. A,, Kuhn, S. J., Tolgyesi, W. S., and Baker, E. B., J. Am. Chem. Soc. 84, 2733 (1962). 16. Olah, G. A,, Baker, E. B., Evans, J. C., Tolgyesi, W. S., McIntyre, J. S., and Bastien, I. J., J. Am. Chem. Soc. 86, 1360 (1964). 17. Olah, G. A., and Donocan, D. J., J. Am. Chem. SOC. 99,5026 (1977). 18. Olah, G. A,, Angew. Chern., Int. Ed. Engl. 34, 1393 (1995). 19. Gold, V., Laali, K., Morries, K. P., and Zdunek, L. Z., J. Chem. Soc., Chem. Commun., 769 (1981). 20. Brouwer, D. M., and Van Doorn, J. A,, Reel. Trav. Chim. Pays-Bas 89,553 (1970); 91, 895 (1972). 21. Davenport, D. A,, Chem. Intell. 451 (1996). 22. Freude, D., and Klinowski, J., J. Chem. Soc., Chem. Commun. 1411 (1988); Brunner, E., J. Chern. Soc., Faraday Trans. 86, 3957 (1990); Hunger, H., Freude, D., Pfeifer, H., and Schieger, W., Chem. Phys. Lett. 167,21 (1990); Hunger, M., Freude, D., and Pfeifer, H., J. Chem. Soc., Faraday Trans. 87, 657 (1991). 23. Datka, J., Kolodziejski, W., and Klinowski, J., Catal. Lett. 24, 265 (1994). 24. Beck, L. W., White, J. L., and Haw, J. F., J. Am. Chem. Soc. 116, 9657 (1994). 25. White, J. L., Ph.D. Dissertation, Texas A&M University, College Station (1992). 26. White, J. L., Beck, L. W., and Haw, J. F., J. Am. Chem. Soc. 114, 6182 (1992). 27. Webb, G. A,, in “Nuclear Magnetic Shielding and Molecular Structure” (J. A. Tossell, ed.), NATO AS1 Ser. Kluwer Academic Publishers, Boston, 1993. 28. Anet, F. A. L., and O’Leary, D. J., Concepts Magn. Reson. 3, 193 (1991). 29. Duncan, M. T., “A Compilation of Chemical Shift Anisotropies.” The Farragut Press, Chicago, 1990. 30. Harris, R. K., “Nuclear Magnetic Resonance Spectroscopy,” p. 247. Wiley, New York, 1989. 31. Fyfe, C. A,, “Solid State NMR for Chemists,” p. 157. C.F.C. Press, Ontario, 1983. 32. Pines, A., Gibby, M. G., and Waugh, J. S., Chem. Phys. Lett. 15, 373 (1972). 33. Pausak, S., Pines, A., and Waugh, J. S., J. Chem. Phys. 59, 591 (1973). 34. See, for example, Soderquist, A., Hughes, C. D., Horton, W. J., Facelli, J. C., and Grant, D. M., J. Am. Chem. Soc. 114, 2826 (1992); Facelli, J. C., and Grant, D. M., Nature (London) 365,325 (1993). 35. Wu, G., Lumsden, M. D., Ossenkamp, G. C., Eichele, K., and Wasylishen, R. E., J. Phys. Chern. 99, 15806 (1995); Wang, J., and Ellis, P. D., J. Am. Chem. Soc. 113,9675 (1991). 36. Ellis, P., Presentation at Pacifichem, Honolulu, HI, December (1995). 37. Wooten, E. W., Mueller, K. T., and Pines, A,, Ace. Chem. Res. 25, 209 (1992). 38. Abragam, A,, “Principles of Nuclear Magnetism.” Oxford University Press, New York, 1961. 39. Frydman, L., and Hanvood, J. S., J. Am. Chem. Soc. 117, 5367 (1995); Medek, A., Harwood, J. S., and Frydman, L., ibid., p. 12779. 40. Bax, A,, Szeverenyi, N. M., and Maciel, G. E., J. Magn. Reson. 52, 147 (1993). 41. Hu, J. Z., Alderman, D. W., Ye, C., Pugmire, R. J., and Grant, D. M., J. Magn. Reson. 105, 82 (1993); Hu, J. Z . , Wang. W., Liu, F., Solum, M. S., Alderman, D. W., Pugmire, R. J., and Grant, D. M., ibid. 113,210 (1995). 42. Herzfeld, J., and Berger, A. E., J. Chem. Phys. 73, 6021 (1980). 43. Xu, T., Torres, P. D., Beck, L. W., and Haw, J. F., J. Am. Chern. SOC. 117,8027 (1995). 44. Xu, T., Barich, D. H., Torres, P. D., Nicholas, J. B., and Haw, J. F., J. Am. Chem. SOC. 119,396 (1996). 45. Xu, T., Ph.D. Dissertation, Texas A&M University, College Station (1996).

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46. Hehre, W. J., Radom, L., Schleyer, P. von R., and Pople, J. A., “Ab Initio Molecular Orbital Theory.” Wiley, New York, 1986. 47. Sauer, J., Chem. Rev. 89, 199 (1989). 48. Nicholas, J. B., Winans, R. E., Harrison, R. J., Iton, L. E., Curtiss, L. A,, and Hopfinger, A. J., J. Phys. Chem. 96, 10247 (1992). 49. Nicholas, J. B., Winans, R. E., Harrison, R. J., Iton, L. E., Curtiss, L. A,, and Hopfinger, A. J., J. Phys. Chem. 96, 7958 (1992). 50. Nicholas, J. B., and Hess, A. C., J. Am. Chem. SOC. 116, 5428 (1994). 51. Parr, R. G., and Yang, W., “Density-Functional Theory of Atoms and Molecules.” Oxford University, New York, 1989. 52. Ziegler, T., Chem. Rev. 91, 651 (1991). 53. Stave, M. S., and Nicholas, J. B., J. Phys. Chem. 97, 9630 (1993). 54. Baker, J., Muir, M., and Andzelm, J., J. Chem. Phys. 102, 2063 (1995). 55. Pelmenschikov, A. G., Morosi, G., Gamba, A,, and Coluccia, S., J. Phys. Chem. 99, 15018 (1995). 56. Datka, J., Boczar, M., and Rymarowics, P., J. Catul. 114, 368 (1988). 57. Kobe, J. M., Gluszak, T. J., Dumesic, J. A,,and Root, T. W., J. Phys. Chem. 99,5485 (1995). 58. Tossell, J. A., ed., “Nuclear Magnetic Shieldings and Molecular Structure,” NATO AS1 Ser. Kluwer Academic Publishers, Boston, 1993. 59. Ditchfield, R., Mol. Phys. 27,789 (1974). 60. Kutzelnigg, W., Fleischer, U., and Schindler, M., in “NMR Basic Principles and Progress” (P. Diehl, E. Fluck, H. Giinther, R. Kosfeld, and J. Seelig, eds.), p. 165. Springer, Berlin, 1991. 61. Kutzelnigg, W., J. Mol. Strucc. (Theochem.) 202, 11 (1989). 62. Kutzelnigg, W., Isr. J. Chem. 19, 193 (1980). 63. Hansen, A. E., and Bouman, T. D., J. Chem. Phys. 82,5035 (1985). 64. Wolinski, K., Hinton, J. F., and Pulay, P., J. Am. Chem. SOC. 112, 8251 (1990). 65. Gauss, J., J. Chem. Phys. 99, 3629 (1993). 66. Sieber, S., Schleyer, P. von R., and Gauss, J., J. Am. Chem. SOC. 115,6987 (1993). 67. Gauss, J., Chem. Phys. Lett. 229, 198 (1994). 68. Gauss, J., and Stanton, J. F., J. Chem. Phys. 102, 251 (1995). 69. Gauss, J., and Stanton, J. F., J. Chem. Phys. 103,3561 (1995). 70. Malkin, V. G., Malkina, 0. L., Casida, M. E., and Salahub, D. R., J. Am. Chem. Soc. 116, 5898 (1994). 71. Schreckenbach, G., and Ziegler, T., J. Phys. Chem. 99, 606 (1995). 72. “Amsterdam Density Functional (ADF), Users Guide, Release 1.1.” Department of Theoretical Chemistry, Free University, Amsterdam, The Netherlands, 1994. 73. Cheeseman, J. R., Trucks, G. W., Keith, T. A,, and Frisch, M. J., J. Chem. Phys. 104, 5497 (1996). 74. Frisch, M. J., Trucks, G. W., Schlegel, H. B., Gill, P. M. W., Johnson, B. G., Robb, M. A,, Cheeseman, J. R., Keith, T., Peterson, G. A,, Montgomery, J. A,, Raghavachari, K., Al-Laham, M. A,, Zakrzewski, V. G., Ortiz, J. V., Foresman, J. B., Cioslowski, J., Stefanov, B. B., Nanayakkara, A,, Challacombe, M., Peng, C. Y., Ayala, P. Y., Chen, W., Wong, M. W., Andres, J. L., Replogle, E. S., Gomperts, R., Martin, R. L., Fox, D. J., Binkley, J. S., Defrees, D. J., Baker, J., Stewart, J. P., Head-Gordon, M., Gonzalez, C., and Pople, J. A,, “Gaussian 94, Revision B.l.” Gaussian, Inc., Pittsburgh, PA, 1995. 75. Jameson, C. J., and de Dios, A. C . , J. Chem. Phys. 97, 417 (1992); 98, 2208 (1993). 76. de Dios, A. C., Pearson, J. G., and Oldfield, E., Science 260,1491 (1993); Le, H., Pearson, J. G., de Dios, A. C., and Oldfield, E., J. Am. Chem. SOC. 117, 3800 (1995). 77. Orendt, A. M., Sethi, N. K., Facelli, J. C., Horton, W. J., Pugmire, R. J., and Grant,

NMR STUDIES OF SOLID ACIDITY

177

D. M., J. Am. Chem. SOC. 114, 2832 (1992); Soderquist, A., Hughes, C. D., Horton, W. J., Facelli, J. C., and Grant, D. M., ibid.,p. 2826. 78. Ellis, P. D., Odom, J. D., Lipton, A. S., and Gulick, J. M., J. Am. Chem. SOC. 115, 755 (1993). 79. Bohmann, J., and Farrar, T. C., J. Phys. Chem. 100,2646 (1996). 80. Haw, J. F., Hall, M. B., Alvarado-Swaisgood, A. E., Munson, E. J., Lin, Z . , Beck, L. W., and Howard, T., J. Am. Chem. SOC. 116, 7308 (1994). 81. Beck, L. W., Xu, T., Nicholas, J. B., and Haw, J. F.,J. Am. Chem. SOC. 117,11594 (1995). 82. Nicholas, J. B., Haw, J. F., Beck, L. W., Krawietz, T. R., and Ferguson, D. B., J. Am. Chem. SOC. 117, 12350 (1995). 83. Nicholas, J. B., Xu, T., Barich, D. H., Torres, P. D., and Haw, J. F., J. Am. Chem. SOC. 118,4202 (1996). 84. X u , T., Barich, D. H., Torres, P. D., and Haw, J. F., J. Am. Chem. SOC. 119,406 (1997). 85. ”ACES 11,” Quantum Theory Project. Dept. of Chemistry and Physics, University of Florida, Gainesville. 86. Beeler, J. B., Orendt, A. M., Grant, D. M., Cutts, P. W., Michl, J., Zilm, K. W., Downing, J. W., Facelli, J. C., Schindler, M. S., and Kutzelnigg, W., J. Am. Chem. SOC. 106, 7672 (1984). 87. Ripmeester, J . A., Tse, J. S., and Davidson, D. W., Chem. Phys. Lett. 86, 428 (1982). 88. Gibson, A. A. V., and Scott, T. A,, J. M a p . Reson. 27, 29 (1977). 89. Reed, A. E., Weinstock, R. B., and Weinhold, F., J. Chem. Phys. 83,735 (1984). 90. Koch, W., Schleyer, P. von R., Buzek, P., and Liu, B., Croat. Chem. Acta 65,655 (1992). 91. Olsson, L., Ottosson, C., and Cremer, D., J. Am. Chem. SOC. 117,7460 (1995). 92. X u , T., and Haw, J. F., in “Topics in Catalysis” (J. J. Fripiat and J. Dumesic, eds.), p. 109 (1997). 93. Haw, J. F., Richardson, B. R., Oshiro, I. S., Lazo, N. D., and Speed, J. A., J. Am. Chem. SOC. 111,2052 (1989). 94. Munson, E. J., Ferguson, D. B., Kheir, A. A,, and Haw, J. F., J. Catal. 136, 504 (1992). 95. Munson, E. J., Murray, D. K., and Haw, J. F., J. Cafal. 141,733 (1993). 96. Olah, G. A,, ed., “Friedel-Crafts and Related Reactions.” Wiley, New York, 1964. 97. Zardkoohi, M., Haw, J. F., and Lunsford, J. H., J. Am. Chem. SOC.109,5278 (1987). 98. Pines, H., “The Chemistry of Catalytic Hydrocarbon Conversions,” Chapter 1. Academic Press, New York, 1981. 99. Deno, N. C., in “Carbonium Ions” (G. A. Olah and P. von R. Schleyer, eds.), Vol. 2, Chapter 18. Wiley-Interscience, New York, 1970. 100. Deno, N. C., and Pittman, C. U., J. Am. Chem. SOC. 86,1744 (1964). 101. Oliver, F. G., Munson, E. J., and Haw, J. F., J. Phys. Chem. 96, 8106 (1992). 102. X u , T., and Haw, J. F., J. Am. Chem. SOC. 116,7753 (1994). 103. Stepanov, A. G., Sidelnkow, V. N., and Zamaraev, K. I., Chem. Eur. J. 2, 157 (1996); Derouane, E. G., Gilson, J. P., and Nagy, J. B., Zeolite 2,42 (1982); Lange, J.-P., Gutsze, A,, Allgeier, J., and Karge, H. G., Appl. Catal. 45, 345 (1988). 104. Lazo, N. D., White, J. L., Munson, E. J., Lambregts, M., and Haw, J. F., J. Am. Chem. SOC. la,4050 (1990). 105. Hutchings, G. J., Lee, D. F., and Williams, C. D., J. Chem. SOC., Chem. Commun., 1475 (1990). 106. Pereira, C., Kokotailo, G. T., Gorte, R. J., and Farneth, W. E., 1. Phys. Chem. 90, 2063 (1990). 107. Munson, E . J., Xu, T., and Haw, J. F., J. Chem. Soc., Chem. Commun., 75 (1993). 108. Biaglow, A. I., Gorte, R. J., and White, D., J. Chem. SOC., Chem. Commun., 1164 (1993).

178

JAMES F. HAW AND TENG XU

109. Buzek, P., Schleyer, P. von R., Vancik, H., Mihalic, Z., and Gauss, J., Angew. Chem.,

Int. Ed. Engl. 33, 448 (1994). 110. Farcasiu, D., J. Chem. Soc., Chem. Commun., 1801 (1994). 111. Xu, T., Zhang, J., Munson, E. J., and Haw, J. F., J. Chem. Soc., Chem. Commun.

2733 (1994). 112. Biaglow, A. I., Sepa, J., Gorte, R. J., and White, D., J. Catal. 154, 208 (1995). 113. Xu, T., and Haw, J. F., J. Am. Chem. Soc. 116, 10188 (1994). 114. Deno, N. C., Groves, P. T., Jaruzelski, J. J., and Lugasch, M. N., J. Am. Chem. Soc. 82, 4719 (1960). 115. Olah, G. A., Pittman, C. U., Waack, R., and Doran, M.,J. Am. Chem. Soc. 88,1488 (1966). 116. Olah, G. A,, Porter, R. D., and Kelly, D. P., J. Am. Chem. Soc. 93, 464 (1971). I f 7. Maciel, G. E., in “Heterogeneous Catalysis” (B. L. Shapiro, ed.), pp. 349-381. Texas A&M University Press, College Station, 1984. 118. Tao, T., and Maciel, G. E., J. Am. Chem. Soc. 117, 12889 (1995). 119. Zhang, J., and Haw, J. F., to be submitted. Z20. Olah, G. A., Rasul, G., York, C., and Surya-Prakash, G. K., J. Am. Chem. Soc. 117, 11211 (1995). 121. Kramer, G. J., van Santen, R. A., Emeis, C. A,, and Nowak, A. K., Nature (London) 363,529 (1993). 122. Xiong, Y., Rodewald, P. G., and Chang, C. D., J. Am. Chem. Soc. 117,9427 (1995). 123. Ajith Perera, S., Bartlett, R. J., and Schleyer, P. von R.,J. Am. Chem. Soc. 117,8476 (1995). 124. Olah, G. A,, and White, A. M., J. Am. Chem. Soc. 91, 5801 (1969); Lee, C. C., and Woodcock, D. J., ibid. 92,5992 (1970); Karabatsos, G. J., Zioudrou, C., and Meyerson, S., ibid., p. 5996. 125. Lazo, N. D., Richardson, B. R., Schettler, P. D., White, J. L., Munson, E. J., and Haw, J. F., J. Phys. Chem. 95, 9420 (1991). 126. Lambregts, M. J., Munson, E. J., Kheir, A. A,, and Haw, J. F., J. Am. Chem. Soc. 114, 6875 (1992). 127. Kheir, A. A., and Haw, J. F., to be submitted. 128. Murray, D. K., Chang, J.-W., and Haw, J. F., J. Am. Chem. Soc. 115,4731 (1993). 129. Murray, D. K., Howard, T., Goguen, P. W., Krawietz, T. R., and Haw, J. F., J. Am. Chem. Soc. 116, 6354 (1994). 130. Sommer, J., Hachoumy, M., Garin, F., and Barthomeuf, D., J. Am. Chem. Soc. 116, 5491 (1994). 130a. Sommer, J., Hachoumy, M., Garin, F., Barthomeuf, D., and Vedrine, J., J. Am. Chem. Soc. 117, 1135 (1995). 131. Engelhardt, J., and Hall, K., J. Catal. 151, 1 (1995). 132. Kazansky, V. B., Senchenya, I. N., Frash, M., and van Santen, R. A,, Cutul. Lefr. 27, 345 (1994). 133. van Santen, R., Presentation at Texas A&M University, November (1995). 134. Blaszkowski, S. R., Nascimento, M. A. C., and van Santen, R. A,, J. Phys. Chem. 100, 3463 (1996). 135. Chevrier, B., Le Carpentier, J.-M., and Weiss, R., J. Am. Chem. Soc. 94, 5718 (1972). 136. Chevrier, P. B., Le Carpentier, J.-M., and Weiss, R., Acta Crystallogr. 28, 2673 (1972). 137. Le Carpentier, J.-M., and Weiss, R., Acta Crystallogr. 28, 1430 (1972). 138. Le Carpentier, J.-M., and Weiss, R., Acta CrysruNogr. 28, 1421 (1972). 139. Fink, F. H., Turner, J. A., and Payne, D. A,, Jr., J. Am. Chem. Soc. 88, 1574 (1966). 140. Koch, H., and Haaf, W., Justus Liebigs Ann. Chem. 618,251 (1958). 141. Stepanov, A. G., Luzgin, M. V., Romannikov, V. N., and Zamaraev, K. I., J. Am. Chem. Soc. 117,3615 (1995).

NMR STUDIES O F SOLID ACIDITY

179

142. Chang, C. D., Cutul. Rev. 25, l(1983). 143. Hellring, S. D., Schmitt, K. D., and Chang, C. D., J. Chem. SOC., Chem. Comrnun.,

1320 (1987). 144. Munson, E. J., and Haw, J. F., J. Am. Chern. SOC. 113, 6303 (1991). 145. Munson, E. J., Kheir, A. A., and Haw, J. F., J. Phys. Chem. 97,7321 (1993). 146. Bosicek, V., Kubelkovi, L., and Novakovi, J., in “Catalysis and Adsorption by Zeolites” (G. Ohlmann, ed.), pp. 337-346. Elsevier Science, Amsterdam, 1991. 147. Xu, T., Munson, E. J., and Haw, J. F., J. Am. Chem. SOC. 116, 1962 (1994). 148. Biaglow, A. I., Sepa, J., Gorte, R. J., and White, D., J. Curul. 151, 373 (1995). 149. Biaglow, A. I., Gorte, R. J., and White, D., J. Phys. Chem. 97,7135 (1993). 150. Biaglow, A. I., Gorte, R. J., Kokotailo, G. T., and White, D., J. Cural. 148, 779 (1994). 151. Munson, E. J., and Haw, J. F., Angew. Chem., Int. Ed. Engl. 105, 643 (1993). 152. Farcasiu, D., and Ghenciu, A,, J. Catal. 134,118 (1992). 153. Farcasiu, D., and Ghenciu, A,, J. Am. Chem. SOC. 115, 10901 (1993). 154. Breitmeier, E., Haas, G., and Voelter, W., in “Atlas of Carbon-13 NMR Data,” Heyden, London, 1979. 155. Biaglow, A. I., Gorte, R. J., and White, D., J. Cuful. 150, 221 (1994). 156. Beck, L. W., and Haw, J. F., J. Phys. Chem. 99, 1076 (1995). 157. Dawson, W. H., Kaiser, S. W., Ellis, P. D., and Inners, R. R., J. Am. Chem. SOC. 103, 6780 (1981). 158. Maciel, G. E., Haw, J. F., Chuang, I. S., Hawkins, B. L., Early, T. E., McKay, D. R., and Petrakis, L., J. Am. Chem. SOC. 105, 5529 (1983). 159. Haw, J. F., Chuang, I.-S., Hawkins, B. L., and Maciel, G. E., J. Am. Chem. SOC. 105, 7206 (1983). 160. Torres, P. D., Xu, T., and Haw, J. F., unpublished results. 161. Grey, C. P., and Vega, A. J., J. Am. Chem. SOC. 117, 8232 (1995). 162. Blumfeld, A. L., Coster, D., and Fripiat, J. J., J. Phys. Chem. 99, 15181 (1995). 163. Ernst, H., and Pfeifer, H., J. Catul. 136, 202 (1992). 164. Xu, T., Zhang, J., and Haw, J. F., J. Am. Chem. SOC.117,3171 (1995). 165. Kheir, A. A,, and Haw, J. F., J. Am. Chem. SOC. 116, 817 (1994). 166. Allouche, A,, J. Phys. Chem. 100, 1820 (1996). 167. Kheir. A. A,, and Haw, J. F., unpublished results. 168. Teunissen, E. H., van Santen, R. A,, Jansen, A. P. J., and van Duijneveldt, F. B., J. Phys. Chern. 97,203 (1993). 169. Teunissen, E. H., van Duijeveldt, F. B., and van Santen, R. A., J. Phys. Chem. 96, 366 (1992). 170. Pelmenschikov, A. G., van Santen, R. A,, Janchen, J., and Meijer, E., .I. Phys. Chem. 97,11071 (1993). 171. Munson, E. J., and Haw, J. F., unpublished results. 172. Mason, S. F., “Chemical Evolution.” Oxford University, New York, 1991. 173. Rothwell, W. P., Shen, W., and Lunsford, J. H., J. Am. Chem. SOC.106, 2452 (1984). 174. Lunsford, J. H., Rothwell, W. P., and Shen, W., J. Am. Chem. SOC.107, 1540 (1985). 175. Bendada, A,, DeRose, E. F., and Fripiat, J. J., J. Phys. Chem. 98,3838 (1994). 176. Baltusis, L., Frye, J. S., and Maciel, G. E., J. Am. Chem. SOC.109, 40 (1987). 177. Jacobs. W. P. J. H., de Haan, J. W., van de Ven, L. J. M., and van Santen, R. A., J. Phys. Chem. 97, 10394 (1993). 178. Sauer, J., Ugliengo, P., Garrone, E., and Saunders, V. R., Chem. Rev. 94, 2095 (1994); Haase, F., and Sauer, J., J. Phys. Chem. 98, 3083 (1994); Sauer, J., Kolmel, C . , Haase, F., and Ahlrichs, R., in “Proceedings of the 9th International Zeolite Conference”

180

179. 180. 181. 182.

183. 184. 185. 186.

187. 188. 189. 190. 191.

JAMES F. HAW AND TENG XU

(R. von Ballamos, J. B. Higgins, and M. M. J. Treacy, eds.), p. 679. Butterworth, London, 1993. Shah, R., Payne, M. C., Lee, M.-H., and Gale, J. D., Science 271, 1395 (1996). Blaszkowski, S. R., and van Santen, R. A., J. Phys. Chem. 99, 11728 (1995). Haase, F., and Sauer, J., J. Am. Chem. SOC.117,3780 (1995). Gale, J. D., Catlow, C. R. A., and Carruthers, J. R., Chem. Phys. Lett. 216, 155 (1993); Vetrivel, R., Catlow, C. R. A,, and Colbourn, E. A,, J. Phys. Chem. 93, 4594 (1989); Gale, J. D., Catlow, C. R. A,, and Cheetham, A. K., J. Chem. Soc., Chem. Cornrnun., 178 (1991). Bates, S., and Dwyer, J., J. Mol. Struct. 306, 57 (1994). Nusterer, E., Blochl, P. E., and Schwarz, K., Angew. Chem., Int. Ed. Engl. 35,175 (1996). Smith, L., Cheetham, A. K., Morris, R. E., Marchese, L., Thomas, J. M., Wright, P. A,, and Chen, J., Science 271, 799 (1996). Sauer, J., Science 271,774 (1996); Krossner, M., and Sauer, J., J. Phys. Chem. 100,6199 (1996); Sauer, J., Horn, H., Haser, M., and Ahlrichs, R., Chem. Phys. Lett. 173,26 (1990). Zygmunt, S. A., Curtiss, L. A., Iton, L. E., and Erhardt, M. K., J. Phys. Chem. 100, 6663 (1996). Parker, L. M., Bibby, D. M., and Burns, G. R., Zeolites 13, 107 (1993). Jobic, H., Czjzek, M., and van Santen, R. A., J. Phys. Chem. 96, 1540 (1992). Batamack, P., Morin, C. D., Fraissard, J., and Freude, D., J. Phys. Chem. 95,3790 (1991). Jentys, A., Warecka, G., Derewinski, M., and Lercher, J. A., J . Phys. Chem. 93, 4837 (1989).

ADVANCES IN CATALYSIS, VOLUME 42

Vibrational Spectra of Hydrocarbons Adsorbed on Metals Part II. Adsorbed Acyclic Alkynes and Alkanes, Cyclic Hydrocarbons Including Aromatics, and Surface Hydrocarbon Groups Derived from the Decomposition of Alkyl Halides, etc. NORMAN SHEPPARD School of Chemical Sciences University of East Anglia Norwich NR4 7T4 England

AND CARLOS DE LA CRUZ Departamento de Quimica Facultad de Ciencias La Universidad del Zulia Maracaibo, Venezuela

1.

Introduction

The development of vibrational electron energy loss spectroscopy [VEELS; otherwise known in the literature as high-resolution energy loss spectroscopy (HREELS)] in the 1970sand of reflection-absorption infrared spectroscopy (RAIRS; sometimes given the alternative acronym IRRAS) in the 1980s led for the first time to the possibility of obtaining vibrational spectra of chemisorbed molecules, one species at a time, on single-crystal surface planes of metals with known and regular atomic arrangements. Abbreviations: AES, Auger electron spectroscopy; amu, atomic mass unit; ARUPES, angleresolved ultraviolet photoelectron spectroscopy; ASED-MO, atomic superposition electron delocalization-molecular orbital; bcc, body-centered cubic (lattice); BPTDS, bismuth postdosing thermal desorption spectroscopy; ESDIAD, electron-stimulated desorption ion angular distribution; fcc, face-centered cubic (lattice); FTIR, Fourier-transform infrared; GC, gas chromatography; hcp, hexagonal close-packed (lattice); INS, inelastic neutron scattering; IUPAC, International Union of Pure and Applied Chemistry; L, the Langmuir (1 X Torr . s);LEED, low-energy electron diffraction; LITD, laser-induced thermal desorption; MS, mass spectrometry; MSSR, metal-surface selection rule; NEXAFS, near-edge X-ray absorption

181 Copyright 0 1998 by Academic Press All rights of reproduction in any form reserved. 0360-0564198 $25 00

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Allied with the diffraction methods, such as low-energy electron diffraction (LEED) and photoelectron diffraction (PED), which can also be applied in single-crystal research, these advances have led to much better interpretations of the vibrational spectra of chemisorbed hydrocarbons in terms of the structures of the surface species. The new results have in turn led to the possibility of reassessing more reliably earlier interpretations of the infrared or Raman spectra of adsorbed hydrocarbons on the finely divided metal samples (usually oxide supported) that are more closely related to working solid catalysts. Such spectra are more complicated because of the occurrence of a variety of different adsorption sites on the metal particles, with the consequence that the observed pattern of absorption bands frequently arises from overlapping spectra from several different surface species. After an introduction to the principles involved in such spectral analysis, in Part I of this review article ( I ) , we attempted a comprehensive and critical reassessment of the literature on the vibrational spectra of adsorbed acyclic alkenes, going back to the pioneering studies by Eischens and Pliskin on Ni/Si02 in 1956-1958 (2,3).The single-crystal work on adsorbed hydrocarbons up to 1988 had already been reviewed by one of us ( 4 ) . In Part I, we updated the single-crystal results for the acyclic alkenes and used them for assistance in reassessing the many published spectra of such hydrocarbons adsorbed on finely divided catalysts. In the figures of Part I, we illustrated many of the significant spectra in the literature of finely divided metals. The spectra of the cyclic alkenes have been left for Part I1 because of their ready surface conversion into aromatic hydrocarbons, to be considered herein. In general, in Part I1 we apply the same pattern of analysis to the numerous published vibrational spectra derived from the adsorption of alkynes, alkanes, and aromatic hydrocarbons. In addition, we summarize recently obtained spectroscopic results characterizing hydrocarbon species obtained by thermal, photochemical, or electron-bombardment dissociation of halogen- or nitrogen-substituted alkanes on single-crystal metal surfaces. The hydrocarbon surface species so obtained are normally as anticipated from the replacement of the heteroatoms by surface metal atoms. The

fine structure; NMR, nuclear magnetic resonance; NRA, nuclear reaction analysis; PED, photoelectron diffraction; RAIRS, reflection-absorption infrared spectroscopy: SER(S), surface-enhanced Raman (spectroscopy); SFG, sum-frequency generation; SIMS, secondary-ion mass spectrometry; STM, scanning tunneling microscopy; T-NEXAFS, transient near-edge X-ray absorption fine structure; TPD, temperature-programmed desorption; UPES, ultraviolet (valence level) photoelectron spectroscopy; UHV, ultrahigh vacuum; VEEL(S), vibrational electron energy loss (spectroscopy); XPES, X-ray (core level) photoelectron spectroscopy.

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spectra provide additional aid toward the interpretation of the spectra originating from adsorbed hydrocarbons. Finally, we summarize, again with emphasis on finely divided metals, recent vibrational spectroscopic kinetic investigations of hydrocarbon surface reactions on platinum and relate these to possible mechanisms of the transformations themselves.

II. The Acyclic Alkynes (Acetylenes) A. INTRODUCTORY COMMENTS The literature of the vibrational spectra of adsorbed alkynes (acetylene and alkyl-substituted acetylenes) is very much in favor of single-crystal studies, with fewer reported investigations of adsorption on oxide-supported metal catalysts. Fewer studies still have been made of the particulate metals under the more advantageous experimental conditions for spectral interpretation, namely, at low temperatures and on alumina as the support. (The latter has a wide transmittance range down to ca. 1100 cm-l.) A similar number of different single-crystal metal surfaces have been studied for ethyne as for ethene adsorption. We shall review in more detail the lowtemperature work which usually leads to HCCH nondissociatively adsorbed surface structures. Only salient features will be discussed for higher temperature ethyne adsorption that often leads to dissociative chemisorption. Many of the latter species are those already identified in Part I from the decomposition of adsorbed ethene.

B. ETHYNE (ACETYLENE) 1. Low-Temperature Single-Crystal Work

a. Spectroscopic Classification. Ethyne has been adsorbed on the following single-crystal planes: Ni(ll1) (5-15); Ni(100) (16-18); Ni(ll0) (19, 20);Ni[5 (111) x (?lo)] (7); Pd(ll1) (21-27); Pd(100) (23,28,29);Pd(11O) (20, 30, 31);Pt(ll1) (32-35); Pt(100) (36, 37);Rh(ll1) (38, 39); Ru(0001) (40-42); Re(0001) (43);Ir(ll1) (44,45);Fe(ll1) (46);Fe(ll0) (47);Fe(100) (48);W(111) (49); W(100) (50);W(110) (51, 52); Cu(ll1) (53);Cu(100) (54, 55); Cu(ll0) (56);Ag(ll0) (57, 58). Table I gives a summary of the types of spectra obtained. In the earlier single-crystal review (4,the low-temperature ethyne spectra were classified in terms of the three types indicated in Table I. Types A and A' spectra have in common strong features from vCC modes (as identified with the help of spectra from adsorbed C2H2and C2D2)in the 1300- to 1050-cm-l region together with vCH absorptions between 2950

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TABLE I Spectral Types,‘ Temperatures, and References for the Adsorption of Ethyne (Acetylene) on Single-Crystal Metal Surfaces at Low Temperatures Surface Metal

(111)

Fe (bcc)b Ni (fcc) c u (fcc) Ru (hCP)

A (120 K) (46)’ A (150 K) (8) A (110 K) (20, 53) low 0, A (130 K) (41)f high 0, B (150 K) (40, 4I)f B (77 K) ( 3 9 ) g B/B‘ (150 K) (21, 22, 25-27)

Rh (fcc) Pd (fcc) Ag (fee) Ir (fcc) Pt (fcc)

-

B (180 K) (44, 45) B (150 K) (32-35)

(110) A (120 K) (47)“ B (80, 120 K) (19, 20) A/B (280 K) (56) Not applicable

(100) A (170 K) (16, 17)‘ A’ (140 K) (51) Not applicable

-

high 0, B (90 K) (20,30)‘” low 0, B’ (90 K) (30) weak i~ (100 K) (56)

-

A ‘ (210 K) (28) -

-

-

A’? (120 K) (36)

a The significance of the type designations, A , A‘, B, and B’, are discussed in the text (also see Sheppard ( 4 ) for a visual display of the spectra). (bcc), body-centered cubic; (fcc), facecentered cubic; (hcp), hexagonal close packed. < Exhibits a “soft” vCH mode. Two closely related species. ‘Somewhat different spectra were reported in Dinardo et al. (16). f T h e term 0 denotes surface coverage. g When ethyne is coadsorbed with CO, this spectrum has a type A’ profile.

”

and ca. 2850 cm-’. Type A‘ spectra differ from those of type A in having features of notably greater intensity in the 970- to 850-cm-’ region, from SCH or yCH modes (8, in-plane, and y , out-of-plane, with respect to the plane of the nonlinear HCCH species). Type B spectra have weak vCC absorptions in the 1400- to 1 1 4 0 - ~ m -region, ~ vCH absorptions close to 3000 cm-’, and very prominent yCH absorptions between 780 and 650 cm-’. The positions of the vCC and vCH modes imply that the type B spectra arise from surface species in which the carbon hybridization is rather less changed toward sp3,from the sp value of the originating ethyne, than is the case for the species giving the A or A’ spectra. In an early paper, Demuth and Ibach (6) suggested that the type A spectrum of ethyne on Ni(ll1) may correspond to a surface species involving four metal atoms by adsorption across the central M-M bond of the (111)unit cell of the face-centered cubic (fcc) crystal. This involves interactions with the two somewhat different threefold sites on either side of the central M-M bond. This model implies that the plane of the HCCH group is perpendicular to the (111)face so that, as found experimentally, the yCH

VIBRATIONAL SPECTRA OF HYDROCARBONS

185

out-of-plane mode is forbidden according to the metal-surface selection rule (MSSR, see Part I). Recently, this proposed structure has been confirmed directly by photoelectron diffraction for Ni(ll1) (59) and also for Cu(ll1) (60),both of which give type A spectra. The type B spectrum, by contrast, shows a strong yCH feature, implying that the plane of the HCCH group is at a substantial direction away from the surface normal. Following a suggestion by Ibach and Lehwald (34) in relation to the spectrum on Pt(lll), there is a strong consensus that on such planes this is likely to correspond to a di-a/a structure, with a a-bond from each carbon atom to metal and a a-bond to a third close-packed metal atom. This configuration finds support in the similarity between the type B spectral pattern and the infrared spectrum of a metal cluster compound with such a structure (61). As yet there seems to have been no direct photoelectron diffraction confirmation of this structure, although such a structure is consistent with low-energy electron diffraction (LEED) results for ethyne on the Rh(ll1) crystal plane (39), which also gives a type B spectrum. It remains to discuss the origin of the extra strong feature between ca. 850 and 970 cm-l in the type A' spectra relative to those of type A. On one particular surface, Ni(lll), one laboratory recorded a clear type A spectrum at 150 K (8) whereas another reported a type A' spectrum at 240 K in which the additional strength of the ca. 870-cm-' feature was the only significant difference (12). This change could be interpreted as from the growth of a second species at the somewhat higher temperature. Pd(100) shows the strength of the 850- to 970-cm-' feature most markedly (24,28), and a more prominent than usual feature also occurs in the same region in the type B spectrum on Pd(ll1). It has been shown by temperatureprogrammed desorption (TPD) that this metal shows a particularly ready self-hydrogenation reaction of ethyne to give ethene (62). As a-complexes from adsorbed ethene absorb most strongly in the 900-cm-l region, such an additional species could provide a possible candidate for an explanation of this feature. However, one would expect such a species to give a dipoleactivated feature, strongest on-specular, but this is not the case. On the contrary, there is clear evidence in many cases that this feature in the VEEL spectra is impact-allowed. For example, DiNardo et al. (16,17) have shown in the case of Ni(100) that a feature at ca. 940 cm-' is much stronger off-specular but that it also varies in intensity down to zero on-specular when the plane of incidence changes from the (100) to the (110) direction within the surface plane. Both of these features denote impact scattering, and the variable intensity of the same band with the direction of the plane of incidence denotes the operation of on-specular impact selection rules (Part I, Section 1V.B). There are other examples, e.g., Ni(ll1) itself (8, 23)

186

NORMAN SHEPPARD AND CARLOS DE LA CRUZ

and Fe(ll0) (47),whereby a small degree of off-specular orientation leads to a very rapid rise in intensity of the feature near 870 cm-l relative to the vCC band at ca. 1220 cm-l. In such cases, an imprecise orientation of the detector with respect to the specular direction, or even the use of a spectrometer with an unusually wide collection angle, could also lead to enhancement of the ca. 900-cm-' feature. Clearly, several spectra that exhibit strong type A' features, e.g., those of ethyne on Ni(11l) or Pd(100), would merit reinvestigation, with careful attention being paid both to the identification of the specular direction and to the direction of the plane of incidence of the electron beam. If such experiments fail to account for type A/type A' variations, the additionalspecies hypothesis could be tested by remeasuring systems giving type A' spectra at markedly lower temperatures, when dissociative adsorption or self-hydrogenation would not be expected, and perhaps by deliberate postadsorption of ethene and/or perdeuterioethene. b. Spectroscopic Trends as a Function of the Metal and of the Crystal Face. Table I summarizes the general pattern of spectra obtained from ethyne adsorption at low temperatures as a function of metal and of crystal face. Both type A and type B spectra occur on (111) faces of the facecentered cubic (fcc) metals and, because the close-packed (111) plane has usually been the first to be investigated, this has provided the model for correlating spectral pattern with surface species, as described earlier. Type A spectra are obtained on the (111) surfaces of the earlier groups VIIA and VIII (IUPAC groups 7-10) transition metals, and type B on those of the later ones. Ruthenium seems to be an intermediate case, where a type A spectrum is obtained at low coverage and a type B one at high coverage. With respect to type B spectra obtained on (111) surfaces, it should be noted that benzene can be formed by the trimerization of ethyne; for example, after ethyne adsorption it is desorbed from Pd(1ll) at temperatures exceeding 250 K (26). If present, benzene would be expected to give a strong band close to that normally attributed to the yCH mode of the di-uin species. Gates and Kesmodel(22, 24) had attributed a strong VEEL feature from ethyne adsorbed on Pd(ll1) near 680 cm-' to the latter mode. Marchon later showed that this could increase greatly in intensity with coverage at 153 K and attributed the increase to the coadsorption of benzene (25).However, Timbell et al. (26) attributed this phenomenon to a coveragedependent negative-ion resonance of the di-u/n species. The results of these experiments emphasize that the kinetic energy of the electron beam is another variable that can affect the appearance of VEEL spectra. Turning next to the other flat plane, (loo), we see that all the spectra are of type A or A', with similar spectra reported for Ni(100) (one azimuthal

VIBRATIONAL SPECTRA OF HYDROCARBONS

187

direction), Pd( loo), and Fe(100). Indeed the d i - d r species, which requires a close-packed triangular site and is correlated with type B spectra on (111) planes, would not be expected since such sites are not present on an unreconstructed (100) plane. It should be noted, however, that a near-edge X-ray absorption fine-structure (NEXAFS) study, while agreeing that a C-C bond (elongated to 0.141 nm with respect to 0.121 nm of free ethyne) occurs parallel to the Cu(100) surface, has been interpreted in terms of a three-metal-atom site based on the diagonal of the square unit cell of the bare (100) metal surface (63). As against this, the lack of strong r C H features in the VEEL spectrum (55) does suggest an orientation of the HCCH plane near to perpendicular to the surface. In the work of Dinardo et al. cited earlier (16, I 7 ) , it was concluded that the spectrum of ethyne on Ni(1OO) exhibited a plane of symmetry along the diagonal of the square unit cell. This would be consistent with an HCCH species across the diagonal with its plane perpendicular to the surface. The bond and angle dimensions would be reasonable for such a species. The expectation from the low wavenumbers of the vCC modes observed forethyne adsorbed on (110) faces [on all but Ag(ll0) (57)] is that nonlinear arrangements of three or four metal atoms are involved in the adsorption sites. These cannot be supplied by atoms along the ridges of the (110) surface, but locally flat three- or four-metal-atom sites can occur involving facets within the troughs of the (110) surface. However, species formed by adsorption on these surfaces will often have different orientations with respect to the surface normal compared with similar species on (111) sites. As a result, the relative intensities of the bands associated with vCC and yCH modes are no longer adequate identifiers of such adsorbed species on these nonflat surfaces. For example, the use of a triangular site to form a three-metal-atom species within a trough could lead to quite different orientations of the HCCH plane, and hence of the strength of the yCH feature, depending on whether the a-bonds are formed to ridge metal atoms or to those at the bottom of the troughs. However, the strengths or weaknesses of the yCH modes, which give dipole changes essentially perpendicular to the HCCH plane, can still be used to tell qualitatively whether the latter plane is more nearly parallel or more nearly perpendicular to the metal face. On this more limited criterion, the spectra obtained from Ni(ll0) (19,20) and Pd(ll0) (20,30) are dominated at higher coverage (5 L or greater dosage) by closely similar species with HCCH planes oriented more nearly parallel than perpendicular to the surface. The weakness and breadth of the vCC absorptions may possibly also indicate a threemetal-atom site, but this is less certain. On Pd(ll0) another species has been indicated by weak bands at 1560 (HCCH) and 1520 cm-' (DCCD) (30). A quite different spectrum, obtained at low coverage ((0.3 L) on

188

NORMAN SHEPPARD AND CARLOS D E LA CRUZ

Pd(llO), on the other hand, implies a species with the HCCH plane essentially perpendicular to the (110) surface. The latter spectrum has also a prominent “soft” vCH mode at 2820 cm-l, indicating an interaction between one of the CH bonds and a surface metal atom. The suggested structure-an HCCH plane perpendicular to an M-M bond of the ridge but tilted down toward the bottom of the trough (30)-is consistent with these spectroscopic criteria. A recent combined LEED, UPES, NEXAFS, and theoretical study of the ethyne/Ni(llO) system (64) suggested a structure with the C-C bond parallel to the direction of the surface-atom rows and bonded to two metal atoms on adjacent ridges. This structure implies an HCCH plane perpendicular to the surface and is therefore unfortunately at variance with the spectroscopic conclusions, at least at near-saturation coverage. Experimental spectroscopic work at lower coverage, analogous to that carried out on Pd(ll0) (30) would be advantageous. The spectrum of ethyne on Cu(ll0) at 280 K differs from those of ethyne on Ni(ll0) and Pd(ll0) at low temperatures in showing additional absorption bands and strong and well-defined vCC and vCH absorptions (56). Both of the latter bands are broad and weak for the other two metals. In the case of Fe(llO), which has the different body-centered cubic (bcc) structure, the (110) plane is nearest to close-packed so that the observed type A spectrum at 120 K may have its more usual significance as indicating a four-metal-atom site not very different from that on fcc (111) planes. Adsorption on Ag(ll0) at 100 K gives a spectrum much less strongly perturbed relative to the spectrum of the free ethyne molecule than any of the others. This clearly denotes relatively weak n-bonding to the surface (57, 58), in marked contrast to the copper case. Lehwald and Ibach (7) investigated ethyne adsorption at 150 K on a stepped Ni surface [5 (111) X (?lo)]. They found a unique spectrum, with strong bands at 2220 and 350 cm-’, starting from either C2H2or C2D2.This is clearly from a C2 species with essentially a triple bond, probably located along the inside of the steps. By the remarkably low temperature of 180 K, the C2 species decomposed further to give surface carbon (520 cm-’) and H atoms (690 cm-I). The saturation of a Pd(ll0) surface with Cs leads to a surface reconstruction of the missing-row form. This provides larger (111) facets for the adsorption of ethyne. At low coverage at 90 K, again a spectrum is found implying an HCCH plane perpendicular to the surface. This showed different band positions compared with clean Pd(llO), and there was no sign of a “soft” vCH absorption (31). At higher coverages di-cr-bonded vinylidene and physical adsorption were considered to occur, with the latter being

VIBRATIONAL SPECTRA OF HYDROCARBONS

189

driven off at 135 K. At 235 K the surface vinylidene was largely desorbed as ethene. 2. Ambient and Higher Temperature Single-Crystal Work Table I1 is a collection of information about the dissociatively adsorbed species thought to be present at higher temperatures, usually 300 K or higher, on the various metal surfaces after ethyne adsorption at low temperatures. A d i d r C=CH2 structure was earlier proposed by Gates and Kesmodel to occur on Pd(ll1) at 250 K (22). A similar spectrum occurs on Pt(ll1) at 350 K (34). Ultraviolet photoelectron spectroscopy (UPES) combined with NEXAFS has recently confirmed this structure on Pd(ll1) (65).The C-C bond, of length 0.146 nm, was found to be oriented at about 50" with respect to the surface normal. The spectra show that the C=CH2 species converts to ethylidyne (CCH3) by self-hydrogenation on Pd(ll1) and Pt(ll1) (22, 34) at 300 and 420 K, respectively. It probably also coexists with a HCCH d i - d r (type B spectrum) species on Rh(ll1) at 320 K (38) and on Pt(ll1) at 370 K (35).A d i d r C=CH2 structure has also been suggested on Fe(100) at 393 K (48). Ethylidyne occurs on the triangular threefold sites on fcc (111) or hexagonal close-packed (hcp) (0001) faces and is formed at lower temperatures on Pd(ll1) and Pt(ll1) in the presence of coadsorbed hydrogen. Its spectral signature also occurs on Ru(0001) at 330 K and on Ir(ll1) at 300 K. Ni(ll1) is exceptional in not giving spectroscopic evidence for the ethylidyne species derived from adsorbed ethyne (or from ethene, I ) . Table I1 shows that at temperatures between 450 and 550 K all the close-packed surfaces investigated, and additionally Cu(lOO), Pd( loo), and Fe(llO), give similar spectra. These have latterly been attibuted to a surface species that we designate a(CCH). The spectrum has bands at ca. 3000 cm-' (m), vCH; 1360-1300 cm-' (m, bd), vCC; and 870-730 cm-' (s), 6CH or yCH. We designate this species cr(CCH) to distinguish it from others given the same formula but with notably different spectra that we discuss later. A particularly good example of an a(CCH) spectrum was observed for the species on Pd(ll1) (24),whereby it was clearly distinguished from a spectrum from adsorbed benzene, which also has a strong absorption (at 730 cm-')in the low-wavenumber region. Another example is provided by the species on Ni(ll1) at 550 K (14). The description a(CCH) is probably an oversimplification. The same type of spectrum is also obtained from the thermal dissociation at similar temperatures of initially adsorbed G H , species. Recent scanning tunneling

TABLE 11 Suggested Structural Interpretations of Higher-Temperature (>300 K ) Spectra of Ethyne Adsorbed on Different Metal Surfaces

Fe (bcc) Ni (fcc)

a(CCH) (375 K) (46)",' .1 Carbon species (500 K) A' (300 K) (12)

CH? (550 K) (47)c

-

-

-

I

c u (fcc)

a(CCH) (450 K) (7) -

-

a(CCH)

+ ?CCH2 (375 K) (54) J.

Ru (hcp)g

high 8,d B

+ CCH3 (330 K) (41)< 1

Rh (fcc)

a(CCH) (550 K) (54)

a(CCH) (450 K) B + CCH2 (320 K) (39)f

Not applicable

Not applicable

-

-

P(CCH) (>ZOO K) (30)

?CCHZ + B (300 K) (28)

I

Pd (fcc)

CH and/or a(CCH) (470 K) (39) CCH2 (250 K) (22)

I CCH? + CCHz (300 K) I

Ir (fcc)

a(CCH) (500 K) (24) a(CCH) + CCH3 (300 K) (44)

1

?cr(CCH) (450 K)

1

1

a(CCH) (450 K) (24, 28)'

Carbon (600 K) -

-

-

-

1

Pt (fcc)

a(CCH) (500 K) (44) B + CCHP (350 K) (33)

1

CCH? (420 K ) (34, 35) " a (CCH), P(CCH), and y (CCH) denote species with decreasing perturbation (increasing CCH angles) from linear MCCH. See further discussion of a(CCH) in Sections II.B.2 and 1V.D. 'The original authors suggested CH. ' An off-specular spectrum. "The term 8 denotes surface coverage. ' Also see Jakob et al. (41).f With coadsorbed CO, the spectrum is clearly that of ethylidyne. 8 Close-packed (0001) surface.

VIBRATIONAL SPECTRA OF HYDROCARBONS

191

microscopy (STM) images of the thermal evolution of ethene on Pt(ll1) suggest that the species of this composition is in fact a polymerized material, one carbon atom thick, that transforms into surface graphite at yet higher temperatures (66).The strong bands in the 870- to 730-cm-' region could be out-of-plane modes (yCH) of peripheral CH bonds of a growing aromatic system. One, two, or three adjacent aromatic CH groups give absorptions in this region, the value moving to higher values with reduced numbers of adjacent CH groups. This type of spectrum is also associated with saturation coverage of ethyne on Ni(ll1) at 420 K and thicker polymerized layers at 550 K (14). See a further discussion of such spectra in Section 1V.D. Considering the intermediate temperature range of 160-400 K, Yoshinobu et al. (30) assigned a spectrum with a strong band at 925 cm-', a broad weak band at ca. 1290 cm-l, and a medium vCH band at ca. 3000 cm-' to a CCH species formed by scission of the "soft" C-H bond of initially adsorbed ethyne on Pd(ll0). At low coverage, this was replaced by a spectrum attributed to surface carbon at 500 K; at higher coverage, an intermediate spectrum similar to that of a(CCH) was observed before carbon formation at 600 K. We designate the first-formed CCH species with a strong band at 925 cm-l as P(CCH). The same spectrum occurs after decomposition of ethene on Pd(ll0) [see Part I, ref 1861. Ethyne on Ag(ll0) with preadsorbed oxygen at 170 K gave a spectrum with absorptions at 690 and 3250 cm-', reasonably attributed to a tilted AgC= CH species without n-bonding to the surface (67). We designate this y(CCH). The CCH species, if such they be, are presumed to be increasingly nonlinear in the sequence y + P + a. At an early stage of the investigation of hydrocarbon adsorption by VEELS, spectra were measured on the ( l l l ) , (110), and (100) faces of tungsten at ambient temperatures (49-52). The first dose of ethyne in each case gave spectra that were attributed to dissociation to give carbon and hydrogen. At higher coverages, spectra with vCH features appeared on surfaces that were already carbon-contaminated. The spectra are not easy to interpret, but the vCC modes in the 1200- to llOO-cm-' region imply a carbon hybridization closer to sp3 than to sp'. 3. Spectra of Ethyne on Finely Divided Metals The results reported for ethyne adsorbed on finely divided metals are rather fragmentary. We therefore review the results metal by metal and attempt an overview at the end. We start with the group VIIIC (IUPAC group 10) metals Ni, Pd, and Pt, which show wide catalytic activity. The available wavenumber ranges are limited to down to 1300 cm-' on SiOz supports or to 1100 cm-' on A1203.

192

NORMAN SHEPPARD AND CARLOS DE LA CRUZ

a. Nickel. Spectra from ethyne adsorbed on Ni/Si02 were obtained during the pioneering period by Eischens and Pliskin (2, 3 ) (Fig. 1A) and by Little et al. (68) near room temperature. However, we commence our discussion by considering the low-temperature and detailed spectra obtained recently by Lapinski and Ekerdt (69) with NilAl2O3 and Ni/Si02. The spectrum, measured at 204 K after adsorption at 178 K, is illustrated in Fig. 1D. The authors convincingly assigned absorptions at 2877, 1333, and 1122 cm-I to the presence of the ethylidyne species, the formation of which might have been assisted by the presence of some residual adsorbed hydrogen on the metal particles after reduction. They also suggested that strong features at 292812913 and 1260 cm-' correspond rather well to the type A/A' spectra from ethyne adsorbed on Ni(ll1) and Ni(100). Other prominent features at ca. 2960 and 1460 cm-l indicate the presence of CH3 groups in alkyl rather than alkylidyne species. Also weak absorptions at 3015 ( v = CH), 1630, and 1585 cm-I (vC=C) indicate that probably minor

3000

2900 2800

1600 1500 1400

1300

1200

cm-'

FIG.1. Infrared spectra of ethyne adsorbed on Ni: (A) Ni/SiOz, H-covered; (B) Ni/SiOz, (-) H-covered, (---) H-depleted; (C) Ni/SiOz; (D) Ni/AI2O3,204 K. [(A) from Ref. 3; (B) from Ref. 7 0 ( C ) from Ref. 71; (D) reprinted with permission from Ref. 69. Copyright 1990 American Chemical Society.]

VIBRATIONAL SPECTRA OF HYDROCARBONS

193

amounts of unsaturated ethylene species, such as the r-complex, were also present. Earlier room-temperature spectra of species formed from ethyne adsorbed on Ni/Si02 obtained by Sheppard and Ward (70) (Fig. 1B) and by Erkelens and Wosten (71) (Fig. 1C) are similar to each other. They, and the low-resolution spectrum reported by Eischens and Pliskin (Fig. lA), which also showed absorption bands in the SCH2/6CH3region, imply the presence of alkyl groups (2960 cm-', vCH3 as; 2930 cm-', vCH2 as; ca. 2870 cm-', bd, vCH3 s + vCH2 s; 1460 cm-', SCH2/SCH3as; and 1380 cm-', SCH3 s). In addition, a weaker absorption at 3025 cm-l denotes the presence of v= CH groups of some form of unsaturated species. Erkelens and Wosten (71) additionally observed a weak absorption at 1685 cm-l. Similar absorptions have been observed for species on Pt/A1203and Co/ Si02, and their possible chemical significance is discussed in the section on cobalt. The addition of hydrogen led to the elimination or strong weakening of the 3025 (and 1685) cm-' bands and to a very substantial growth in intensity of the alkyl absorptions. There is general agreement, from Eischens and Pliskin onward, that the spectrum after hydrogenation has a strong resemblance to that expected from an n-butyl group. This implies that, as was found to be the case with ethene (Part I, Section VI.B.3), dimerization had occurred before or during hydrogenation. Eischens and Pliskin suggested that their spectrum obtained before hydrogenation could be representative of an ethyl group, but the later work done at higher resolution suggests that it is more likely that a mixture of alkyl species was present. By gas chromatography, Lapinski and Ekerdt showed that at temperatures exceeding 200 K, methane, ethane, and ethene occurred in the gas phase together with ethyne. Room-temperature Raman wavenumbers for ethyne (C2H2and C2D2) on Ni/Si02have been recorded by Krasser et al. (72), but without illustration of the spectra. The higher wavenumber listed bands at 2988, 2910, 1204, 1044, 864, and 806 cm-l show some promising coincidences, particularly the 1204-cm-' feature from C'C, with the bands in the VEEL spectrum of ethyne on Ni(ll1) (8). In a few cases, adsorption on particulate nickel has been studied other than in the form of the conventional oxide-supported metal catalysts. Nash and De Sieno (73) exploded nickel wires in a rare-gas atmosphere to give Ni particles of ca. 20-nm diameter. Results were reported (but not illustrated) characterizing adsorption of ethyne; they are similar to those found by Eischens and Pliskin. Ito et al. (74) evaporated Ni films onto quartz glass. Upon adsorption of C2H2or C2D2,they observed pairs of absorptions that were assumed to

194

NORMAN SHEPPARD AND CARLOS DE LA CRUZ

arise from two different, but related, adsorption sites. For C2H2 the absorptions occurred at 3250/3150 ( vCH), 1845/1800 ( VCC), and 770/ 850 cm-' (SCH). Corresponding wavenumbers of absorptions of ethyne in the gas phase that, taking into account the MSSR, would be expected for a molecule adsorbed parallel to the surface are 3274, 1974, and 730 cm-'. The uCH absorption of the adsorbed species at lower wavenumber, implying stronger adsorption, increased in intensity at higher temperatures. However, both sets of wavenumbers imply relatively limited perturbations by bonding to the surface, probably only through n--bonding. No such HCCH species seem to have been detected for species on single-crystal or oxide-supported nickel surfaces. Additional weaker absorptions at 2900 and 2750 cm-' for species on the evaporated nickel film imply the presence of some sp3-hybridized species also. The second of these bands is unusually low in wavenumber even for an alkyl grouping. This probably implies that the C-H bond in question has an agostic C-H...M interaction with a metal atom of the surface, of a type independently observed for cyclohexane adsorbed on Pt(ll1) or on Ni(ll1) and for ethene adsorbed on Ni(ll1) (75). Both these absorptions characterizing the species on the evaporated film may correspond to those from the di-a CZH4 species on N i ( l l l ) , which, with the lower precision of VEELS, are reported to occur at 2940 and 2690 cm-'. Bobrov et al. have, remarkably, recorded the Raman spectrum of ethyne adsorbed at 80 K on evaporated films of Ni, showing shifted u=CH and u C = C bands at 3300 and 1860 cm-', respectively. They also recorded bands at 3080, 3030, and 1525 cm-l, which they assigned to n--bonded ethene (76). b. Palladium. The earliest investigation of ethyne adsorbed on finely divided Pd supported on porous silica glass by Little et al. (68; Fig. 2A) showed absorption bands at ca. 3090 and 3030 cm-'. These are clearly evidence of unsaturated alkenyl species with C = CH2 and C =C H groups, the former possibly indicative of physically adsorbed ethene. An improved spectrum reported by Clark (77),who also used a sample with porous glass as the support, is shown in Fig. 2B. It shows additional weak bands at ca. 2970, 2930,and 2870 cm-' in the alkyl region. A very similar spectrum was obtained by Dunken et al. (78; Fig. 2C) with Pd/Si02 in a pressed disk starting from finely powdered silica. Clark repeated his work with hydrogencovered Pd, and this enhanced the intensity of the absorptions in the alkyl region. A later study by Beebe et al. (79; Fig. 2D) of Pd/AIZO3clearly identified the presence of ethylidyne species at room temperature, as anticipated from single-crystal results; as only a partial spectrum was provided, it was not made explicit whether other species coexisted with ethylidyne.

VIBRATIONAL SPECTRA OF HYDROCARBONS

3100

3000 2900 2800 crn-

1

1350

195

1250

FIG. 2. Infrared spectra of ethyne adsorbed on Pd: (A) Pd/SiOz; (B) Pd/SiOz; (C) Pd/Si02; (D) Pd/AlzOs. [(A) from Ref. 6 8 (B) from Ref. 77; ( C ) from Ref. 78; (D) from Ref. 79.1

What is common to all these results is that the adsorption of ethyne on Pd (as on Ni) leads to gradual self-hydrogenation to give first alkenyl and then to give alkyl-type surface species. For chemical balance these must coexist with hydrogen-deficient carbonaceous species. There is general agreement that the addition of H2 to the initially adsorbed species on Pd leads, over a period of hours, to a large increase in intensity, giving absorption bands from n-alkyl species with four or more carbon atoms. The intensity increase is expected for the hydrogenation of the carbonaceous species. The nature of the products shows that oligomerization of the initial C2 units had occurred. Clark suggested that this might have occurred before the addition of hydrogen, as considerably more molecules of ethyne than of ethene were adsorbed on the same catalyst. Several authors commented that after ethyne adsorption it is much more difficult to restore Pd catalysts to their original clean state by hydrogen treatment; i.e., the oligomerized surface species are difficult to remove.

196

NORMAN SHEPPARD AND CARLOS DE LA CRUZ

Sokolova et al. (80) reported two absorptions, at 3050 and 2960 cm-l, for species on Pd/A1203at temperatures between 223 and 373 K. Using inelastic electron tunneling spectroscopy, Bayman et al. (81) obtained complex spectra of species formed from C2H2 and C2D2 adsorbed on Pd/ A1203. These show the presence of a mixture of sp2- and sp'-hybridized hydrocarbons in addition to bands indicative of undissociated ethyne. Parker et al. (82)reported a remarkably strong Raman spectrum of CzH2 (and also G D 2 and 13C2H2)adsorbed on Pd/A1203and also on Rh/A1203 (Fig. 3A). They suggested that this is evidence of the di4ii.i species which give the type B VEEL or FTIR spectra. However, Patterson and Weaver (83) observed a closely similar spectrum for ethyne adsorbed on a gold electrode (Fig. 3B) and convincingly interpreted it in terms of the resonance Raman spectrum of all-trans-polyacetylene (84, 85). It is possible that the very high selectivity of resonance Raman spectra for conjugated polyenes enabled the identification of the initially polymerized species that give rise on hydrogenation to lengthy polymethylene chains. However, it should be borne in mind that such ethyne polymerization has also been shown to occur on metal oxide surfaces alone (Ti02; 85). For ethyne adsorbed on an evaporated Pd film, Ito et al. (74) reported infrared spectra similar to those described earlier for species formed from ethyne on Ni, with the modified wavenumbers of 3350/3220, 1880/1850, and 8351755 cm-' for the pairs of perturbed ethyne species and 2870 and 2750 cm-' for the saturated species. On wire-exploded Pd, Nash and De Sieno (73) reported a spectrum similar to that observed by Dunken et al. on Pd/Si02.

3000

2000

1000

crn - 1 FIG.3. Raman spectra of ethyne (in polymerized form as a long-chain poiyacetylene) on (A) Rh/AI2O3 and (B) a gold electrode. [(A) Reprinted with permission from Ref. 94. Copyright 1985 American Chemical Society; (B) reprinted with permission from Ref. 83. Copyright 1985 American Chemical Society.]

VIBRATIONAL SPECTRA OF HYDROCARBONS

197

c. Platinum. Room-temperature spectra of species formed from ethyne on Pt/porous glass reported by Clark (77; Fig. 4A) and on Pt/Si02 reported by Ward (70, 86; Fig. 4B) and Prentice (87; Fig. 4C) are very similar and once again denote the presence of a mixture of unsaturated and saturated hydrocarbon surface species. Once again, hydrogenation led to n-alkyl groups corresponding to a mean value of about six for the number of carbon atoms. Randhava and Rehmet (88) reported a single absorption at 1690 cm-I on Pt/AI2O3,similar to one described earlier for Ni/Si02. Its possible interpretation is discussed in the section on colbalt. Sziligyi (89; Fig. 4D) reported a similar band at 1700 cm-l on Pt/Si02 together with additional absorptions that resemble those obtained (without a 170O-cm-' feature) from adsorbed but-1-ene (Part I; Section VI.D.2) and which are again consistent with the occurrence of surface dimerization.

0

-

0 m

s

"

O E C N ", 0. 1

1

I

I

I

I

I

I

I

-

0 Q c

I

I

3100

3000

I

f

2900 2800

I

1700

'

A Pt

1600 1500 1400 1300

FIG.4. Infrared spectra of ethyne adsorbed of Pt: (A) Pt/SiOz; (B) Pt/SiOz; (C) Pt/SiOz; (D) Pt/Si02. [(A) from Ref. 77; (B) from Ref. 70; (C) from Ref. 87; (D) reprinted from Ref. 89, Vibr. Spectrosc. 2, T. Szilagyi, p. 29. Copyright 1991 with kind permission of Elsevier Science-NL, Sara Burgerhartstraat 25, 1055 KV Amsterdam, The Netherlands.]

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NORMAN SHEPPARD AND CARLOS D E LA CRUZ

d. Rhodium. Room-temperature spectra of species formed from ethyne adsorbed on Rh/Si02were recorded by Kavtaradze and Sokolova (90) and, in more detail, by Pearce et al. (91,92).The latter spectrum is illustrated in Fig. 5A. The presence of ethylidyne is clearly denoted by the absorptions at 1342 and (partial absorption) 2880 cm-l. The other alkyl absorptions are similar to those obtained after the adsorption of but-1-ene on Pt/ Si02 (93), and so it is no surprise that hydrogenation leads to n-butyl surface species. Parker et al. (94) obtained strong Raman spectra from ethyne adsorbed on Rh/A1203 (Fig. 3A) of the same form as those for Pd/A1203.Ethyne adsorbed on a Rh3'-modified silver sol gives rise to a strong SERS feature at ca. 1915 cm-', which gradually decays to give another strong and broad

I

I

I

I

I

1

cm-'

FIG.5. Infrared spectra of ethyne adsorbed on several metals: (A) Rh/SiOz ; (B) Co/Si02; (C) Cu/SiOz; (D) infrared spectrum of propyne (methylacetylene) adsorbed on PtlSiO,. [(A) from Ref. 91; (B) reprinted with permission from Ref. 96. Copyright 1974 American Chemical Society; (C) from Ref. 68; (D) from Refs. 77 and 111.1

VIBRATIONAL SPECTRA OF HYDROCARBONS

199

band at ca. 15.50 cm-' (95).These are attributed to species with C = C and C = C bonds, respectively. A more detailed discussion of such spectra on silver is given in Section g. e. Cobalt. A single spectrum of species formed from ethyne adsorbed at room temperature on Co/Si02 was published by Blyholder and Wyatt (96) (Fig. 5B). In addition to alkenyl and alkyl vCH absorptions, the latter supplemented by a band at 1450 cm-l, there is at 1690 cm-' a strong version of the analogous absorptions mentioned earlier in connection with adsorption on Ni/Si02 and Pt/A1203.This absorption seems to occur sporadically, which suggests that it is associated with some special condition of catalyst preparation. In this case, with Co/Si02, it occurred also after the adsorption of ethene but only when the catalyst had been evacuated at high temperature (673 K) after the reduction of the salt to give the metal. The same conditions applied for ethyne adsorption. Catalysts which are evacuated at high temperature and cooled in vacuum stand a greater chance of contamination by oxygen. The possibility should therefore be borne in mind that the ca. 1690-cm-' absorptions might represent relatively small amounts of organic carbonyl groups produced by oxidation. Compared with the absorptions of hydrocarbon groupings, organic vCO absorptions are intrinsically very strong. On the other hand, Sziliigyi (89) reported the same spectrum for species on Pt/Si02, whether or not the catalyst had been evacuated at high temperatures. If the 1690-cm-' absorptions are associated with a hydrocarbon surface species, it is, as pointed out by Erkelens and Wosten ( 7 I ) ,very unlikely that they arise simply from metal-substituted alkenyl groups. Metal substitution for hydrogen normally lowers vC= C. HC=CH \ / The latter authors suggested the cyclic M structure as a possibility. Szilhgyi (89) suggested an H2C= C = M species on Pt/A1203.Model compounds of these types [structures 12 and 19 of Figs. 4 and 5 and Table IV of Part I ( I ) ] do give absorptions between 1700 and 1600 cm-l. f. Copper. Copper on porous silica glass (68) gave spectra from adsorbed ethyne that are very similar to those described for Pd (Fig. SC). Once again, alkyl groups form by hydrogenation, although very slowly. g. Silver. Ethyne, as CZH,, C2D2,and I3CzH2,when adsorbed on coldevaporated silver (97),gives surface-enhanced Raman spectra (SERS) indicative of the presence of two species (Fig. 6A). One of these, with a lowering of vCC by 40 cm-' in comparison with gas-phase ethyne at 1993 cm-', exhibits all the infrared- and Raman-active wavenumbers of a weakly

200

NORMAN SHEPPARD AND CARLOS D E LA CRUZ

3000

2000

1000 crn - 1

FIG.6. SER spectra of hydrocarbons adsorbed on cold-deposited silver: (A) ethyne; (B) ethene; (C) ethane; (D) benzene. [(A) reprinted from Ref. 97, J. Electron Spectrosc. Relat. Phenorn. 29, I. Pockrand, C. Pettenkofer, and A. Otto, p. 409. Copyright 1983 with kind permission of Elsevier Science-NL, Sara Burgerhartstraat 25, 1055 KV Amsterdam, The Netherlands; (B) and ( C ) reprinted from Ref. 98, Sur$ Sci. 188, C. Pettenkofer, I. Mrozek, T. Bornemann, and A. Otto, p. 519. Copyright 1987 with kind permission of Elsevier ScienceNL, Sara Burgerhartstraat 25, 1055 KV Amsterdam, The Netherlands; (D) reprinted from Ref. 99, J. Electron Spectrosc. Relat. Phenom. 54/55, J. Mrozek and A. Otto, p. 895. Copyright 1990 with kind permission of Elsevier Science-NL, Sara Burgerhartstraat 25,1055 K V Amsterdam, The Netherlands.]

perturbed HCCH molecule. It desorbs at 145 K. The second species (with vCC lowered by 86 cm-' and with two X = C H bands, at 756 and 635 cm-') is more strongly adsorbed and is probably also from intact HCCH molecules adsorbed on a different site. SER spectra from ethane, ethene (98),and benzene (99) also on cold-deposited silver are shown for comparison in Figs. 6B-D. In general, the close relationships of the band positions and relative intensities of many of the SER bands that are common to those in the Raman spectra of the parent hydrocarbons imply that the adsorbed species are weakly perturbed. However, in all cases additional bands occur due to the reduction of symmetry of the hydrocarbon compo-

VIBRATIONAL SPECTRA OF HYDROCARBONS

201

nent within the adsorption complex. For the unsaturated hydrocarbons the attachment at the surface is of a n-bonding nature, and the CH stretching bands are abnormally very weak. An SER spectrum from ethyne on colloidal silver particles isolated in a solid argon matrix gave absorptions closely similar to those of the aforementioned more weakly bound species (100). SER spectra from ethyne adsorbed on silver electrodes show multiple bands in the v C = C region at ca. 2130,1985, and 1810 cm-' plus, at more negative potentials, a strong and broad feature at ca. 1550 cm-l. The latter seems to grow at the expense of that at 1810 cm-l (101). The higher wavenumber bands were tentatively attributed to polymerized acetylide species in which the o-bond to one metal atom is supplemented by a nbond to another. There does, however, appear to be a good correlation between the ca. 1985- and 1550-cm-' features, and it is remarkable that closely similar spectra are obtained from ethyne and alkyl-substituted alkynes, including 2-alkynes. These features suggest the formation of oligomerized acetylene species, less well-defined compared with a clear alltrans spectrum earlier obtained from ethyne on a gold electrode (83).The position of the strong ca. 1550-cm-l feature suggests a chain of about four vC= C bonds, possibly with some showing cis configurations (85). The ca. 1810-cm-' absorption may represent a monomeric surface species with strong v-bonding to the metal, which is converted into a polyene at more negative potentials. Bobrov et al. (76) reported Raman spectroscopic data for ethyne adsorbed on evaporated Ag at 80 K. h. Gold. Spectra similar to those obtained on silver were observed from ethyne adsorption on gold and gold/silver electrodes (101). Longerterm bubbling of ethyne through the solution gave well-defined spectra from linear all-trans-polyenes (83). 4. General Comments on the Spectra of Ethyne Adsorbed on Finely Divided Metal Catalysts

The results observed for ethyne adsorbed on single-crystal surfaces are not easy to relate to those observed for ethyne on the finely divided metals. This is because the more readily interpreted spectra of the species on the flat surfaces occur at low temperature, whereas, with one exception, the spectra of the species on the metal particles have been obtained at room temperature. There is, however, one common finding, namely, that ethylidyne is frequently found in single-crystal spectra near 300 K and that this species has been found at room temperature on oxide-supported catalysts of Ni, Pd, and Rh.

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NORMAN SHEPPARD AND CARLOS DE LA CRUZ

All the spectra from initially adsorbed ethyne on oxide-supported metals provide evidence for self-hydrogenation on the surface to give alkenyl- and alkyl-type surface species. Ethylidyne formation is just one example of this. Additionally, the spectra of hydrogenated species all show alkyl chains of length four carbon atoms or more, implying that the carbonaceous residues left by self-hydrogenation tend to polymerize. The Raman spectra obtained by Hexter ei al. for species on Rh/AI2O3and Pd/A1203provide clear evidence for polyene formation, suggesting the existence of oligomerized carbon-rich product before the addition of hydrogen. There are always vCH absorptions from =CH2or =CH groups in evidence before hydrogenation, which could be associated with the polyenes. Unfortunately, the skeletal modes of the latter, which are strong in the Raman spectra, are forbidden in the infrared spectrum of a trans single chain and at best are likely to be weak for the chain adsorbed on the metal surface. Of the coinage metals, finely divided copper shows evidence of strong rehybridization of adsorbed ethyne, such as occurs with transition metals Ni, Pd, Pt, and Rh, which have unfilled d orbitals. On the other hand, SER spectra from ethyne on cold-deposited silver at low temperatures show only moderately perturbed n-bonded species. There is SERS evidence of polymerization to long-chain polyenes on silver and gold electrodes at room temperature. Room-temperature spectra of species on Ni and Pd evaporated films that were also deposited at room temperature showed spectra from a mixture of ethyne-like species, probably n-bonded, and alkyl species with “soft” vCH modes indicating agostic C-H...M interactions. C. HIGHERALKYNES

1. Spectra of Higher Alkynes on Metal Single Crystals To date, only propyne (methylacetylene) and but-2-yne (dimethylacetylene) seem to have been studied as adsorbates. Nondissociative adsorption at low temperature is supported by the experimental results in all cases. We first discuss the results obtained from but-2-yne, as the adsorbed species are likely to be more symmetrical and hence, with the use of the MSSR, more effective for structure elucidation. But-2-yne has been adsorbed on Pt(ll1) and studied by VEELS at 320500 K (102) and by RAIRS at room temperature (103). It has also been investigated on Ni(ll1) by VEELS at 80-300 K (104) and on Cu(ll1) by RAIRS at 150 K (105, 106). Three possible structures for the adsorbed species have been considered (1-111). In the first one, I, the triple bond of the parent butyne opens up to give di-a-bonding to two surface metal atoms; in the second, 11, an additional n-bond is formed to a third metal

VIBRATIONAL SPECTRA OF HYDROCARBONS

CH 3

M i HC,\

4

/ c=c,\

- 'M

203

3

M

atom (for ethyne itself, this is the structure assigned to type B spectra on flat metal planes); and in the third, 111, there is bonding to four metal atoms (as has been assigned to type A species from adsorbed ethyne). In cases I and 111, the symmetry of the adsorption complex is C2,, and the MSSRSCH3 as, allowed modes would be vCH3 as, vCH3 s, vC=C/vC"C, SCH3s, CH3rocking, vCC s, vCM s, and SC-C C s. A distinction between these two structures should be possible through the position of the vC=CI vC'C mode, probably between 1600 and 1500 cm-' for I and <1500 cm-' for 111. The latter situation would also apply to 11, which has its symmetry reduced to C, so that one each of an additional vCH3 as, 6CH3 as, CH3 rocking, and an out-of-plane skeletal deformation would in principle be allowed; however, the first two of these might not be resolved from the same types of vibration allowed in the C,, case. Avery and Sheppard (202) were the first to study the VEEL spectrum of but-2-yne adsorbed on Pt(ll1). From the absence of a vC=C mode at wavenumbers greater than 1500 cm-I, they discounted structure I and assumed that the vC-C mode occurred in the 1450- to 1350-cm-l region already occupied by strong SCH3as and SCH3s absorptions. They preferred structure I1 on the grounds that ethyne adsorption on the same crystal plane gives type B spectra. In a more recent RAIRS study of the same system, Pudney (103) gave more precise positions for the absorption bands down to 800 cm-l and was able to resolve the absorption in the vCH3 region into three components. The probable vibrational assignments are as follows: 2930 (w), vCH3 as; 2886 (s), vCH3 s; 1433 (m), SCH3 as; 1354 (ms), vC'C or SCH3 s (see following discussion); and 1040 cm-' (m), CH3 rocking. Fricke et al. (104) obtained a VEEL spectrum from but-2-yne on Ni(ll1) that is closely similar to that on Pt(ll1). However, they found additional weak features at 1580 and 840 cm-l. They assigned the former to the vC =C mode; the latter could represent the vCC s mode of a C2, structure. From a consideration of LEED data, they preferred the structural assignment 111, which is of the type of structure associated with the type A spectrum from ethyne of Ni(ll1). However, 1580 cm-' seems to be a very high value

204

NORMAN SHEPPARD AND CARLOS DE LA CRUZ

for vC'C of 111, as for adsorbed ethyne the value is as low as 1200 cm-'. Furthermore, the 1580-cm-' feature is sufficiently weak to be possibly assigned to an overtone or combination. The RAIRS study of but-2-yne adsorbed on Cu(ll1) by Chesters and McCash (105) shows a similar spectrum in the vCH3 region to that on Pt(lll), but only a single absorption of appreciable intensity occurs in the vC'CISCH~ region, at 1392 cm-l. As this only shifts to 1356 cm-l in the spectrum from the perdeuteriobutyne, this is essentially the vC'C mode. These results show that, in contrast to the Pt(ll1) case, both 6CH3 modes absorb very weakly on Cu(ll1). The results do, however, raise the question whether the 1354-cm-l band in the Pt(ll1) case is from vC"C or from SCH3s. Either the former is very weak on Pt or possibly there is a resonance between vC'C and SCH3s such that one of these coupled modes absorbs very strongly and the other only weakly. Clearly, for both Ni(ll1) and Pt(lll), it is important to identify the vC=C/vC'C and 6CH3 modes unambiguously by study of the adsorption of the perdeuterio- and partially deuteriobutynes. For none of the spectra obtained to date on the (111) surface is there sufficient complexity to give preference to the C, structure I1 over the C2,structures of I and 111,but the anticipated additional features may be weak. However, for ethyne itself the comparative intensities of the vC-C modes (strong in type A spectra from I11 and broad and weak in type B spectra from 11) provide a valid criterion for distinction between I1 and 111, and this may prove to be applicable also to the spectra of the methyl-substituted ethynes. The adsorption of propyne at 1.50 K on Cu(ll1) has been studied by RAIRS (105, 106), as has the adsorption of propyne on Cu(ll0) at 90 K (106a), on Ni(ll1) at 110 K (106a), and on Pt(ll1) at 100 K (107); the adsorption on Rh(ll1) at 240 K has been investigated by VEELS (108). Compared with the spectra of but-2-yne on Cu(ll1) and on Pt(lll), at low coverage extra absorptions were found in the vCH region of the propyne spectra, at 2855 and 2845 cm-', respectively; analogous absorptions occurred on Ni(ll1) at 2869 cm-I and on Cu(ll0) at 2751 cm-'.The 2855-cm-' band shifted to 2157 cm-l in the spectrum of CH3CCD on Cu(lll), so it is clearly indicative of the stretching vibration of the lone CH group. CD3CCH spectra on Ni(ll1) and Cu(ll0) also confirmed the remaining 2869- and 2751-cm-' bands in the vCH region to be from this group. The major change in the vCH frequency from its value at 3305 cm-' in the spectrum of the propyne itself clearly signals a hybridization change at the carbon atom from sp to between sp2 and sp3 and leads to an expectation of a vC'C rather than a vC=C absorption in the surface complex. A prominent feature at 1361 cm-l characterizing the species on Cu(lll), which moved to 1353 cm-' in the spectrum of CH3CCD,was assigned to vC' C,

VIBRATIONAL SPECTRA OF HYDROCARBONS

205

and once again, the SCH3 modes were so weak as to be unobservable. However, bands at 1415 and 1356 cm-' characterizing the species on Pt(ll1) could be assigned to SCH3 as and SCH3 s, although the latter might once again alternatively be assigned to vC'C. Spectra of CD3CCHor CD3CCD would resolve the latter ambiguity. However, the close similarity of the spectra of propyne and butyne on Pt(ll1) in this region would seem to favor the SCH3 s assignment, with the implication that the vC'C absorption is very weak for the species on this metal. The SCH3 s interpretation of absorptions between 1360 and 1350 cm-' has recently been confirmed by RAIRS studies of propyne and CD3CCH on Ni(ll1) and on Cu(ll0) (1064 109). Higher coverage spectra of propyne on Cu(ll1) and on Cu(ll0) showed multilayer features. The intensity of the vCH3 as absorption relative to that of vCH3 s is greater for the species on Pt(ll1) than for those on Cu(ll1) for both propyne and butyne adsorption. This would be expected for a type I1 structure in which the four-carbon plane is inclined at an angle to the surface, compared with for a type 111 structure in which the C, plane is perpendicular to the surface. By analogy with the case of ethyne, adsorption of propyne on Ni(ll1) and on Cu(ll1) is likely to occur with structure 111. It should be noted that for CH3 groups attached to double bonds, as in the spectrum of tetramethylethene (110), the vCH3 s infrared absorptions are normally stronger than those of vCH3 as, in contrast to the opposite situation which applies to CH3 in saturated alkyl groups. On Rh(ll1) at 240 K, propyne and its isomer propadiene (allene) have very similar VEEL spectra (108),which, like the RAIR spectrum of propyne on Pt(lll), show prominent features in both the vCH3 and 6CH3 regions. Warming the propyne on Rh(ll1) to 310 K gives a very different spectrum, which is clearly indicative of ethylidyne formed by C-C bond cleavage. Results of deuterium-substitution experiments suggest that it is the bond that was originally C = C that is broken on the surface (108). VEEL spectra have also been presented for but-2-yne adsorption on Pt(ll1) at 400 and 500 K (102).They reveal complex dissociative changes shown by sp2-hybridized vCH absorptions of two species with strong absorptions at 790 and 720 cm-', respectively.

2. Spectra of Higher Alkynes on Finely Divided Metals Surprisingly, the literature seems to record only a single infrared study of propyne adsorption on a metal oxide-supported metal, Pt/SiO,, where the silica support was in the form of porous glass (77,111). The spectrum at room temperature, obtained only in the vCH region, showed absorptions at 3015, 2970, ca. 2930, and ca. 2880 cm-' (Fig. 5D). These denote the

206

NORMAN SHEPPARD AND CARLOS DE LA CRUZ

presence of a mixture of alkenyl and alkyl groups. A very large increase in the intensity occurred on the addition of hydrogen to give a methylrich spectrum. Originally, this spectrum was attributed to the presence of surface-bonded 2-propyl groups. However, it was later shown that analogous spectra obtained by the hydrogenation of surface species from other adsorbed alkenes arose from the physical adsorption of the alkane products, in this case propane, on the silica support (112). A recent study of the hydrogenation of propyne on Pd/ZrOz showed residual absorptions from surface species after evacuation. These were attributed to di-o and propylidyne adsorbed species, but no spectra were illustrated (223). SER spectra of the alkynes but-2-yne, pent-1-yne, and pent-2-yne on a gold electrode have been measured (83).At -0.4 V the v C r C modes are lowered by about 110 cm-' from the values for the free alkynes. For the alk-2-ynes the shift increases by a further 10 cm-' upon a change of the potential to +0.4 V. The spectrum of but-2-yne has two principal component bands, at 2134 and 2109 cm-I in the vC=C region, and also shows a downward shift for the vCH3 modes. Multicomponent V C S C bands were also observed in the SER spectra of phenylacetylene adsorbed on copper, silver (114), and gold (83) electrodes. The principal components characterizing the species on Ag and Au -93 and -125 cm-I), respectively. are at ca. 2017 and 1985 cm-' (Av The higher wavenumber band is displaced by the adsorption of chloride ions when the potential was changed from -0.6 to 0.0 V. Four vC=C components were observed for the species on the copper electrode, centered at ca. 1900 cm-I (Av -210 cm-'). It is considered that all these spectra represent alkyne molecules adsorbed on the metal electrodes via the nelectrons of the C z C groups (a -+d and d + n* electron transfers). In the case of phenylacetylene on Ag, it was considered that the V E CH band was also detected (114), although this conclusion has been questioned (83).

-

-

D. PROPADIENE (ALLENE, H2C=C=CH2) RAIR spectra have been reported for propadiene adsorbed on Cu(llO), N i ( l l l ) , and on an evaporated Ag film (124a), and VEEL spectra have been obtained on Rh(ll1) (108). Closely similar spectra on Cu(l10) at 90 K and on the Ag film at 130 K were convincingly interpreted in terms of nondissociative adsorption, with the C= C= C skeleton and the plane of one of the C= CH2groups slightly tilted from parallelism with the surface (symmetry Cs).The bonding to the surface is probably of a weak 7~ nature. On Ni(ll1) at 130 K, a different type of sectrum was obtained with some features that intensified at 295 K, indicating that isomerization had occurred to give adsorbed propyne. The remaining features in the spectrum at 130 K

VIBRATIONAL SPECTRA OF HYDROCARBONS

207

were considered to arise from an intermediate in the isomerization process, possibly from nondissociative adsorption with one of the C = C groups converted to di-c bonding to the surface. The VEEL spectrum of propadiene on Rh(ll1) at 77 K also changed at 240 K to a spectrum identified as being from an adsorbed propyne species.

111.

The Acyclic Alkanes A. INTRODUCTION

Relatively few vibrational spectroscopic investigations have been reported for acyclic alkanes, whether adsorbed on single-crystal or on finely divided metal surfaces. The spectra of the cyclic alkanes are more conveniently discussed later (Section V.A) because of their relationships to the spectra of aromatic species into which they are readily converted by metal catalysts. More experimental difficulties have to be overcome to observe chemisorbed species with the alkanes than with the alkenes or alkynes. The study of adsorption on metal single-crystal surfaces often starts from multilayer adsorption at low temperatures followed by warming and evacuation to remove the multilayers and reveal the chemisorbed monolayer. However, in the case of the low-molecular-weight alkanes under ultrahigh-vacuum (UHV) conditions, frequently the physically adsorbed layers are completely removed before a sufficiently high temperature is reached to overcome the activation energy for chemisorption. Also, direct adsorption at higher temperatures can similarly be hampered by low dissociation probabilities. However, Ceyer and her group (115,116) have overcome these difficulties in the case of methane on Ni(ll1) by causing a beam of high-velocity molecules to collide with the surface, so that the kinetic energy dissipated at the collision overcomes the energy of activation. Alternatively, a physically adsorbed layer of methane can be bombarded by a beam of inert-gas atoms of high kinetic energy (“chemistry with a hammer”) (116).

B. METHANE

1. Single-Crystal Work The species first to be expected from the chemisorption of methane on a metal surface is a surface-attached methyl group. If, as seems probable, the threefold axis of the CH3 group is perpendicular to the metal surface (C3” overall symmetry of the surface complex), the expected completely symmetrical modes are vCH3 s, SCH3s, and vCM (M = metal). According

208

NORMAN SHEPPARD AND CARLOS DE LA CRUZ

to the metal-surface selection rule (MSSR), these will be the only ones allowed in a RAIR spectrum and should be dipolar-dominant in on-specular VEELS. The additional modes of vCH3us, SCH3 us, and CH3 rock should additionally become observable in off-specular VEELS. Just such a spectrum was obtained by Lee et al. (117,118) using VEELS after the impact of methane molecules with a normal-incidence energy of >70 kJ mol-' on a Ni(ll1) surface held at 140 K. More recently, these spectra have been analyzed at higher resolution by Yang et ul. (119). The assignments are as follows: ca. 2730 (w, sh), vCH3 as; 2655 (ms), uCH3 s; 2610 (m), 2SCH3 as; 2430 (w), 2SCH3 s; 1320 (ms), SCH3 as; 1220 (vs), 8CH3 s; 965 (w), CH3 rock; 485 (ms), CH3 torsion; and 385 cm-I (s), vCNi s. The intensities relate to the on-specular spectrum. Off-specular, the 1220- and 385-cm-' bands were relatively much weakened as expected for completely symmetrical modes. It was deduced that the CH3 axis is perpendicular to the metal surface and that a threefold hollow site is occupied. A similar on-specular spectrum was obtained by the bombardment of a physically adsorbed methane layer on Ni(ll1) by a beam of highenergy A r atoms (115). A high-energy impact experiment involving CD4 gave the on-specular VEEL spectrum of 2030 (w), vCH3 as; 19.55 (ms), vCD3 s; 1925 (m, sh), 2SCD3 as; 1815 (w), 2SCD3 s; 978 (m), SCD3 as; 915 (vs), 6CD3 s; 730 (w), CD3 rock; 375 (w), CD3 torsion; and 365 cm-' (ms), vCNi s (118). The bands at 915 and 365 cm-' were much weakend offspecular. A weak feature at 2200 cm-' was attributed to vCD2 s from surface-adsorbed CHDz (118, 120). Collectively, the number of vibrational modes observed on- and offspecular is just what is expected for a C3"surface methyl species. However, on comparison with expectations based on model compounds (CH3)4M (M = Sn and Pb, Table 111, Part I) or (CD3)4M, the vCH3 us, vCH3 s, and vCD3 s modes, observed at 2730, 2655, and 1955 cm-I, respectively, have very low values. Yang et al. (119) more recently attributed this to back-donation from metal d orbitals to antibonding orbitals of the CH3/ CD3 groups. The 2655-cm-' band in the VEEL spectrum is not very broad (ca. 60 cm-') compared with the "soft-mode" absorptions resulting from agostic hydrogen-bonding-like C-H ..-M interactions. In their first published spectrum (117), the authors also showed a band at 2940 cm-', which they attributed to a C-H bond free of hydrogen bonding. Hence at that time they suggested that the CH3 group was tilted toward a surface metal atom on one side. However, the latter absorption has been missing (seemingly without explanatory comment) in subsequent papers of the series, while the vCH3 as mode was identified at 2730 cm-' in the later off-specular work. It was reported that the surface methyl group on Ni(ll1) is stable up to 150 K but thereafter decomposes to give surface

VIBRATIONAL SPECTRA OF HYDROCARBONS

209

CH groups at 220 K, adsorbed ethyne (HCCH) at 320 K, and adsorbed benzene at ca. 400 K (120, 121). Intermediate CH2 groups were not observed. A similar experiment by Oakes et al. (122) involving the impact of highkinetic-energy methane on Pt(ll1) at 150 K by RAIRS gave vCH3 s as a single broad band at the much more normal position of 2885 cm-'. In that case, ethylidyne, CH3C, was observed by heating to 300 K or by highenergy impact at 400 K. This was possibly formed via the dissociation of CH3 to CH2 followed by dimerization to ethene and then dissociation to ethylidyne (Part I, Section 1V.A.l.e). However, neither of the suggested CH2 or ethene surface intermediates was detected spectroscopically. More ) shown that adsorption of methane recently, Yoshinobu et al. ( 1 2 2 ~have on Pt(ll1) at 25 K gives a first undissociated layer with the v3 and v4 modes lowered in wavenumbers by 22 and 11 cm-', respectively, from the gasphase values. Irradiation of this layer with an ArF laser gave the CH3/CD3 species with absorptions at 288212098 cm-'. Camplin et al. (123) reported that at 25 K methane is nondissociatively adsorbed on Cu(100) but with displacements of the triply degenerate vCH2 and SCH2 modes to lower wavenumbers by 21 and 8 cm-', respectively. Wu and Goodman (124) studied the reaction of methane with singlecrystal surfaces of Ru(0001) and Ru(1120) over a range of temperatures from 330 K upward. A spectrum obtained for the species on Ru(llT0) at 330 K was very reasonably attributed to an ethylidyne species necessarily tilted with respect to the surface on Ru3 sites. At the higher temperature of 500 K, a different spectral profile contained a feature at 1620 cm-' denoting vC=C, leading to a common spectrum at 600 K reasonably attributable to a surface vinylidene species (3010, v= CH2s; 1620, vC= C; 1410, S=CH2; 1165, =CH2 rock in-plane). An absorption at 830 cm-' grew in intensity with temperature relative to the other bands but decreased again at temperatures exceeding 600 K. Together with a weaker companion at 3010 cm-', it was assigned to surface methylidyne, CH. When the surface was Ru(0001), the latter spectrum dominated from 500 K, when surface species were first observed, and was accompanied between 500 and 600 K by a weaker version of the spectrum attributed to vinylidene. After the 830/301O-cm-' spectrum was removed at 800 K, the surface was found to be covered with a graphitic carbon layer, and once again the possibility is that this spectrum arises from single aromatic CH groups incorporated in the growing graphitic layer (see the discussion on the origin of the a(CCH) spectrum in Section II.B.2 and further discussion of the topic in Section 1V.D). Other methods have since been devised for obtaining spectra from surface methyl groups, from the dissociation of methyl halides or (CH3)2N2on

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NORMAN SHEPPARD AND CARLOS D E LA CRUZ

metal surfaces, and we return to this aspect of the literature and to further discussion of the CHJNi(111) spectrum in Section 1V.B. 2. Finely Divided Metals Infrared spectra, together with magnetization measurements, were made at temperatures between 298 and 373 K for a Ni/Si02 catalyst in the presence of gas-phase methane (125). No infrared absorption bands of surface species were obtained, although the magnetization results suggested that throughout this temperature range methane was adsorbed in the highly dissociative forms of Ni3C and NiH. In view of the magnetization results for the Ni/Si02 catalyst, it seems that further FTIR studies of the adsorption of methane on supported catalysts, starting at subambient temperatures, should be worthwhile. Only a weak Raman spectrum of physically adsorbed methane, against a strongly fluctuating background, was reported when methane was in contact with a polycrystalline Co catalyst at ambient temperature (126). C. ETHANE 1. Single-Crystal Work Studies of adsorption of ethane, CzH6 or C2D6, on the single-crystal surfaces Cu(llO), C u ( l l l ) , and Pt(ll1) [(127),(128),and (129,130),respectively] at low temperatures by RAIRS have only shown absorptions from undissociated adsorbed molecules. In each case, there was strong evidence that at the monolayer adsorption level the C-C bond is parallel to the metal surface. As expected for such an orientation, when the MSSR is taken into account, the vCH3 as and SCH3 as modes gave strong absorptions. In the cases of the high-resolution RAIR spectra obtained with Cu(ll1) and Pt( 111) by Chesters et al. (128,130),the orientation is confirmed by detailed analysis of the spectra in the vCH3 region. 2. Finely Divided Metals

At an early stage, Eischens and Pliskin (3) reported that no absorption bands were observed when ethane was in contact with “hydrogen-covered” Ni/Si02, but when the catalyst was “hydrogen-free,” essentially the same absorptions were obtained as from the adsorption of ethene (Fig. 13, Part I). The latter spectrum implied the presence of both unsaturated and saturated hydrocarbon groups. Somewhat similar ethane/ethene spectral relationships for adsorbed species have been reported for ethane adsorbed on Pt/Si02near room temperature (131,132).The spectrum obtained by D e La Cruz and Sheppard (132)

VIBRATIONAL SPECTRA OF HYDROCARBONS

A

3700

'

I

3000 2900

I

I

211

1

1600 1500 1400 1300

2800 cm-'

FIG.7. Infrared spectra of (A) ethane on Pt/SiOz and (B) propane on Pt/Si02. [(A) from Ref. 132; (B) reprinted from Ref. 140, Spectrochim. Acfa 46A, G. Shahid and N. Sheppard, p. 999. Copyright 1990 with kind permission of Elsevier Science-NL, Sara Burgerhartstraat 25, 1055 KV Amsterdam, The Netherlands.]

for ethane adsorbed on Pt/Si02 at 294 K is illustrated in Fig. 7A. There are prominent features at 2881 and 1348 cm-' characterizing ethylidyne, other bands at 2915 and ca. 1415 cm-', previously assigned to the di-uethene species, and another weak but broad absorption at ca. 2950 cm-', clearly evidence of the vCH3as mode of a methyl-containing surface species. The spectrum differs from that obtained for ethene adsorbed on the same catalyst (Section VI.B.1.c and Fig. 7, Part I) in the enhanced strength of the ca. 2950-cm-' absorption and the virtual absence of the absorptions characteristic of the n-complex near 3020 and 1500 cm-l. The broad and weak absorption near 2950 cm-l could be evidence of the methyl group of a surface ethyl species, but it is more likely that its breadth denotes the presence of a small amount of ethane. This could be formed on standing after the initial evacuation of the gas phase by a gradual rehydrogenation of unsaturated surface species by surface hydrogen. The spectrum of Fig. 7A was little changed by heating in a closed cell to 413 K, but when the temperature reached 513 K, methane had appeared in the gas phase, denoting C-C bond scission, initially at the expense of ethylidyne absorptions. When the temperature reached 573 K, all the surface absorptions had disappeared in the accessible spectral region (4000-1300 cm-') and gas-phase methane absorptions were further enhanced. In an early study of ethane adsorbed on Pt supported on porous silica glass, Clark (77) found no infrared bands on initial adsorption at room temperature, 453, or 513 K, but on hydrogenation observed absorptions near 2935 and 2870 cm-', which could be attributed to surface alkyl species.

212

NORMAN SHEPPARD AND CARLOS DE LA CRUZ

An infrared study of the interaction of ethane with a Pd/A1203catalyst did not reveal any absorption bands of adsorbed species at (seemingly) 373 K but strong bands indicative of oxygenated (carboxylate) surface species were obtained at temperatures between 373 and 478 K when ethane plus air was in contact with the catalyst (133). Ito et af.(74) obtained a room-temperature spectrum of ethane adsorbed on an evaporated film of palladium. Absorptions were observed at 2850 and 2700 cm-’ characteristic of an alkyl species. The “soft” vCH 2700cm-l band implies the presence of agostic hydrogen-bond-type interactions of C-H bonds with surface metal atoms. SER spectra have been obtained for undissociated ethane adsorbed on cold-deposited silver films (134-137). On adsorption at 22 K (98) (Fig. 6C), a spectrum was obtained showing both Raman- and infrared-active bands, the latter indicating the modes expected to be present under the MSSR if the C-C bond is parallel to the surface. The vCH3as and vCH3s absorptions at 2916 and 2858 cm-l were 45 and 22 cm-l, respectively, lower than the corresponding absorptions in the multilayer spectrum, indicating substantial physical interaction with the metal surface. A weak absorption at 2720 cm-’ was assigned to 26CH3 s. Similar spectra were reported for species on In, K, Cu, and Au.

D. HIGHER ACYCLICALKANES 1. Single-Crystal Work Two papers by Chesters et al. (130, 138) relate to the adsorption of the linear alkanes (CnH2n+2r n = 3, 4, 5, and 6) and to the highly branched neopentane (2,3-dimethylpropane) as monolayers on Pt(ll1) at temperatures near 95 K. At this low temperature, the n-alkanes are expected to be predominantly in the trans (planar zigzag) conformations. These seem most likely to be stable on a metal surface given that “hydrogen-bondlike” interactions are known to occur between C-H bonds and surface metal atoms in favorable cases such as the adsorption of cyclohexane (see Section V.A.l). For the aforementioned n-alkanes, the spectra obtained for the species on Pt(ll1) at low coverage appear to refer to monolayer coverage or less. Under these conditions, the almost complete lack of absorptions in the RAIR spectra in the vCH3/vCH2 region below 2900 cm-’ that are characteristic of vCH3 s or vCHp s modes provide a firm indication that the planar zigzag carbon skeleton lies parallel to the Pt surface. Strong features in the 2950- to 2925-cm-’ region, and others between 1450 and 1440 cm-’, are readily assigned to vCH3 as and 6CH3 as, respectively. For each of these vibrations, one component gives a vibrational

VIBRATIONAL SPECTRA OF HYDROCARBONS

213

dipole perpendicular to the carbon chain and the metal surface and is hence as modes are 20-30 cm-’ lower in the allowed by the MSSR. The v C H ~ spectrum of the surface species than in the spectrum of the free alkane or of the multilayers. Bands in the 2910- to 2900-cm-’ region, which increase in intensity with the length of the alkyl chain, are assigned to the MSSRallowed vCHz as modes, again at somewhat lower wavenumbers relative to those of the free alkanes. The investigation of the spectra of the n-alkanes on Pt(ll1) was undertaken to determine whether any broad and low-wavenumber absorptions indicative of “soft” vCH modes could be detected as in the cyclohexane case. Such an absorption may have been present as a broad tail in the 2800-cm-l region of the spectra of n-pentane, n-hexane, and neopentane. However, such bands with wide half-widths give weak maxima and are difficult to detect under the high-resolution conditions of RAIRS. Such “soft-mode” phenomena show up more readily in VEELS. It should be noted, however, that for geometrical reasons there are unlikely to be linear C-H...M interactions characterizing the adsorbed alkanes such as occur in the cyclohexane case. In the neopentane case such a situation would be possible if three of the four CH3 groups interacted with the surface, each via one C-H...M group. However, four less-directed such interactions would be possible if only two methyl groups were each in contact with two surface metal atoms. Recently, n-butane and isobutane have been studied by RAIRS when adsorbed at low temperature on Ag(ll0) (138a). They gave undissociated monolayers which exhibited intermolecular attraction. The translational vibration frequencies of n-C,Hzn+2(n = 6,8,10) against Cu(ll0) surfaces at temperatures near 150K have been measured by helium atom scattering (139). 2. Finely Divided Metals At an early stage, Clark (77) attempted without success to obtain spectra of propane, n-butane, and 2-methylpropane (isobutane) adsorbed on Pt/ SiOs in the range 433-463 K. However, more recently, good spectra of adsorbed propane have been obtained at room temperature for Pt/SiO, by Shahid and Sheppard (140) (Fig. 7B). The dominant absorption bands were of the dissociatively adsorbed species propylidyne. Compared with the case of adsorbed propene, there was less sign of di-c and no sign of 8complexes. Heating to 473 K led to the removal of absorptions indicative of propylidyne and to the generation of a small amount of ethylidyne together with broad and weak vCH absorptions, probably evidence of the superposition of spectra of a number of alkyl species.

214

NORMAN SHEPPARD AND CARLOS DE LA CRUZ

Brogan et al. (141) reported that n-butane does not react at ambient temperature with Pt/A1203or Pt/A1203/Ce02catalysts. CONCLUSIONS REGARDING ACYCLIC ALKANES E. GENERAL Low-temperature spectra of species formed from the adsorption of alkanes on single-crystal metal surfaces have all indicated the presence of undissociated molecules physically adsorbed flat on the surface. The wavenumber lowering of the vCH3/vCH2 absorptions indicates appreciable adsorptive interactions with surface metal atoms. Room-temperature adsorption on finely divided metal catalysts has been shown in several cases (Ni/Si02 and Pt/Si02) to give rise to dissociative adsorption as alkylidynes and other products. It is therefore very clear that, contrary to often-expressed views, C-H bonds within the alkanes can readily be broken by interaction with metal catalyst surfaces. Methane is a very important feedstock and, although this may be the most resistant to chemisorption, there is clearly much further of interest to be discovered in this area involving interactions of light alkanes with different metals. IV. Hydrocarbon Surface Species Derived from the Dissociative Adsorption of Halogen- or Nitrogen-Substituted Alkanes or Alkenes

A. SCOPEAND STRATEGY During the past decade a very considerable literature has developed concerning the generation and reactivity of alkyl and alkylidene groups adsorbed on metal single-crystal surfaces produced via the photochemical or thermal decomposition of adsorbed alkyl halides or nitrogen-substituted alkanes. In this review, we concentrate on publications which exhibit VEEL or RAIR spectra of the hydrocarbon groupings that can be used as reference spectra for the identification of such species in spectra of species derived from hydrocarbon chemisorption. Reviews of such work cover the kinetic as well as spectroscopic aspects of this area of research (142-144). The following subheadings in Section IV relate to the adsorbed species themselves rather than, as in the rest of Parts I and 11, to the adsorbed molecuies from which the spectra are derived. The structure of the section is different here because the surface hydrocarbon groups of interest are often derived from a number of different adsorbed halides or nitrogencontaining adsorbates. METHYLGROUPS AND THEIRDECOMPOSITION PRODUCTS B. SURFACE A particularly large amount of spectroscopic data is available for adsorbed methyl groups on different metal surfaces as generated from the

VIBRATIONAL SPECTRA OF HYDROCARBONS

215

decomposition (thermal or photolytic) of a number of heteroatom-substituted alkenes. The results are collected in Table I11 (117,118,122,145-162). For comparison, we have also incorporated in this table the data obtained from the high-energy-impact experiments of methane on Ni(ll1) and Pt(l1l) (115, 117, 118, 122), discussed earlier in Section 1II.B. Most of the experimental results were obtained by VEELS, but in a few cases much more accurate band positions were obtained by RAIRS in the ca. 3000cm-‘ vCH3 region, where this method has high resolution and its highest sensitivity. It is in just this region that the VEEL spectra give ambiguous band positions because the broad VEEL resonances can be blends of several different but unresolved vCH3bands. Furthermore, it is in this high-energyloss region that, even for on-specular experiments, impact scattering leads to features that compete in intensity with those generated by the dipole mechanism. It is noteworthy, for example, that the VEEL spectra of surface CH3 groups indicate a range of band positions on the same crystal face when several experiments have been carried out involving different initial adsorbates. In contrast, all the RAIR spectra give essentially identical band positions for the vCH3 s mode on Pt(lll), irrespective of whether the surface species is derived from CH31,(CH3)2Zn,or (CH3)2N2as a precursor or from impact experiments involving CH4 (119, 122, 147, 154, 156). The same conclusion was drawn from closely similar VEEL spectra for CH3 groups on Cu(lll), whether the initial adsorbate was CH31 (which gives coadsorbed I atoms after decomposition) or (CH3)2N2(whereby N2is driven off during methyl formation so that there is no residual coadsorbate) (160). Intensities in general are more sensitive to changes in chemical environments than band positions, and this is of relevance when attempts are made to deduce the symmetry of an adsorption site by using allowed/forbidden criteria based on the MSSR. A methyl adsorbate would have C3vsymmetry if, in isolation, it were bonded to an on-top or threefold hollow site on a (111) surface or if it were part of a regular array of adsorbed molecules based on such sites. If the methyl group were adsorbed on a twofold bridged site, the symmetry would be approximately C2, (C, or C1, depending on the orientation if the methyl group does not rotate about the CH3-tosurface linkage). If, on the other hand, the methyl group were bonded to a C3vsite of the clean surface but had a halogen atom as a neighbor on one side but not the other (or, to a lesser extent, if it had other CH3 neighbors in incomplete irregular arrays), then the true symmetry would at best be C,. In MSSR terms, the former C3”case would lead to only the A I modes, with vibrational dipoles perpendicular to the surface, to be allowed in RAIR or to dominate on-specular VEEL spectra. The doubly degenerate E modes would then occur principally off-specular via the impact mechanism. In the latter case, C,, one each of the E modes would also be formally allowed in the on-specular VEEL and RAIR spectra.

TABLE 111 Vibrational Wavenunibers (cm-') and Intensities for Surface Methyl Groups on Metal Single Crystals Generated by Thermal or Photolyric Decomposition of Heteroatom-Substituted Alkanes

Surface Pt(ll1)

Reactant

CH31 CHII CH31 CH31 CH3Br CH3Br CH3Cl (CH3)3Bi (CH&Zn (CH&NZ CH4 Ru(0001) CHJ CH4 Ni(1ll) Cu(l11) CH31 (CH&N2 CHIBr Cu(100) CH31

Method of Type of decomposition spectrum Thermal Thermal Thermal Thermal Photolysis (Xe) Photolysis (Hg) Photolysis (Hg) Thermal Thermal Thermal Impact Thermal Impact Thermal Thermal Photolysis ( X e ) Thermal

VEEL VEEL RAIR RAIR VEEL VEEL VEEL VEEL RAIR RAIR RAIR VEEL VEEL VEEL VEEL VEEL VEEL

vCH3 us

-

vCH3 s 2925 (s) 2925 (s) 2879 (s)

26CH3 as 2775 (w) -

2755 (w)

ca. 2860 (m) 2950 (s) 2790 (w) 2971 (s) 2918 (s) 2882 2885 (m) 2750 (w) 2885 2910 (ms) 2730 (w, sh) 2655 (ms) 2610 (m) 2950 (m)' 2830 (m) 2945 (m) 2790 (s) 2910 (ms) 2781 (ms) 2910 (s) 2760 (s) -

SCH3 as

6CH3 s

1425 (mw) 1165 (s) 1405 (w) 1180 (s) [ 1244?] 1363 ( m ) 1153 (ms) 1410 (s) 1180 (s) 1390 (s) 1174 ( s )

1335 (m) 1320 (ms) 1370 (ms). 1380 (s) 1386 (mw) 1430 (s)

pCH,

vCM

Reference

790 (m) 520 (w) 770 (ms) 545 (w)

145 146 I47 148, 149

734 (w) 492 ( s ) 820 (s) 495 (s) 784 (s) 516 (vs) [765 (s)l

152 I53 156 154, 155

-

1180 (w)? -

1190 (ms) 1220 (vs) 1190 (ms) 1190 (vs) 1185 ( s ) 1150 (s)

150 151

965 (w) 830 (w)" 890 (m) 854 (w)

ca. 470 (s) 385 (s) 345 (s) 355 (s) [480 (bd)] 370 (s)

122 157 117-119 158, 159 160 161 162

Note: Features observed only in off-specular VEELS are in italics. Doubtful assignments are in square brackets. vs, very strong; s, strong; ms, medium strong; m, medium; mw, medium weak; w, weak; bd, broad. Low-coverage spectra only.

217

VIBRATIONAL SPECTRA OF HYDROCARBONS

In a considerable number of cases both sets of modes have been observed in on-specular VEEL spectra, and the deduction has been made that the symmetry of the surface complex is C, (or less) (145,146,151,152,160,162). The question remains whether this implies a twofold bridged adsorption site or a neighbor-induced asymmetry within an essentially C3,site, as already described. However, there are examples of species on Pt(ll1) (150), Ni(ll1) (117), and Cu(ll1) (161) surfaces for which MSSR as applied to VEEL spectra clearly indicates C3, symmetry of the surface complex, without significant differences in the other frequencies as observed off-specular. These favorable cases may arise from particularly regular arrays of adsorbed species, the presence of which could very profitably be confirmed by LEED. We deduce that the CH3 adsorption sites are intrinsically C,, as far as the bare surface is concerned, i.e., on-top or threefold hollow in nature with the threefold axis of the CH3 group perpendicular to the surface. We turn next to the information provided by the band position data. Table 111 shows that the surface methyl group can be characterized by bands in six regions; three of them are always observed in on-specular VEELS (and in principle in RAIRS) and three virtually always in offspecular and sometimes also on-specular VEELS, as discussed earlier. Their values exhibit a metal dependence, as shown in Table IV, where average values are quoted, except that in the case of vCH3 s on Pt(lll), we have made use of a particularly accurate and consistent RAIR value. It is seen that the SCH3 as and SCH3 s modes have relatively narrow wavenumber spreads between metals, and therefore these are particularly useful characteristic modes for identifying this surface species. vCH3 s and vC-M show particularly wide wavenumber variabilities. The values of vCH3 s and of YC-M for Pt and Ru are reasonably typical for such modes in metal coordination compounds involving methyl ligands (Table 111,Part I; also 163), and there seems to be no reason to doubt that these CH3 TABLE IV The Observed Metal Dependence of the Wavenumbers (cm-') Characteristic of Surface-Bound Methyl Groups Metal

Ni Cu" Ru Pt

uCH3as

KH3s

6CH3as

GCH3s

pCH3

uCM

2730 (ms) 2926 (m)

2655 (m) 2790 (ms) 2910 (ms) 2885 (ms)

1320 (m) 1390 (ms) 1335 (m) 1400 (m)

1220 (s) 1180 (s) 1190 (ms) 1170 (s)

965 (w) 860 (mw)

385 (ms) 360 (s) 470 (s) 515 (ms)

-

-

780 (m)

Note: s, strong; ms, medium strong; m, medium; mw, medium weak; w, weak. " A spectrum of CH3 groups on Cu/SiOz has recently been obtained (V. H. Grassian, personal communication), with absorptions at 2913 (rn,sh) and 2805 cm-' (w, bd).

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NORMAN SHEPPARD AND CARLOS DE LA CRUZ

groups are directly bonded to metal atoms in on-top sites. The species on Ni and Cu have notably lower wavenumbers for both the vCH3 s and vC-M modes, implying the occurrence of weaker C-H and C-M bonds. These two metals have very similar values for the carbon-to-surface modes, but in the vCH3 s region the value for Ni(1ll) at 2655 cm-l is very low, whereas the value of 2790 cm-’ for Cu(ll1) is intermediate with respect to the values near 2900 cm-’ for Pt(ll1) and Ru(0001) (Table 111). As has been suggested on experimental (159) and theoretical (164-167) grounds, it seems probable that for Ni(ll1) and Cu(lll), adsorption occurs on the C3”hollow site, probably with another metal atom in the next layer. There is experimental evidence that twofold bridged CH3 or CH2 species as in Mg(CH3)2(168), Mg (CD3CH2)(169),and A12(CH3)h(170) do give lower than normal vCH3/vCH2modes near 2800 cm-l. It has been suggested that in the Cu case the weakening of the C-H bonds is caused by back-donation from the d orbitals of metal atoms into nonbonding A l and antibonding E electronic orbitals of the CH3 group (156,164,165).The alternative possibility of donation from bonding orbitals of CH3 to vacant metal d orbitals could apply to the Ni case, as the d shell of Ni is not full. The even lower vCH3 s wavenumber in the Ni case might hence relate to synergetic (s, p ) + d and d + (s, p ) * electron transfers. Another possible explanation for low wavenumber values for vCH modes is a hydrogen-bond-like end-on C-H...M interaction such as is well known to occur for cyclohexane on Pt(ll1) (75,138). It is difficult to see how such CH3/M interactions could occur for a C3”surface symmetry, and such a structure is not supported by theory (164, 166, 167). There is also a clear interest in increasing and widening the number of metals (and individual surfaces) to be studied for methyl formation, preferably with investigations by both VEELS and RAIRS. The use of methyl radicals derived from the gas-phase decomposition of (CH3)2N2 ( I 71) would now seem to be the preferred route so as to avoid complications from the presence of other coadsorbed species. It would be of particular interest to generate methyl groups on Ni(ll1) by this alternative to the high-energy-impact route. Photoelectron diffraction is a recently developed experimental technique that would seem to be capable of distinguishing between the threefold hollow site proposed for CH3on Ni( 111) and Cu( 111) and the on-top site proposed for Pt(ll1) and Ru(0001). Little work has yet been published with finely divided catalysts, but Rask6 et al. (172) investigated the interaction of CH3C1 and CH31 with a Pd/Si02catalyst by FTIR. New absorptions were observed at temperatures greater than 260 K for CH3C1and at temperatures greater than 213 K for CH31.However, it seems probable that these are indicative of the presence

VIBRATIONAL SPECTRA OF HYDROCARBONS

219

of adsorbed alkyl chains rather than CH3groups. The addition of potassium with CH3C1probably led to longer, CHz-dominated, alkyl chains.

C. SURFACEMETHYLENE (CH,) GROUPS Attempts have been made to obtain VEEL spectra of CHz on singlecrystal planes by the decomposition of ketene, CHzCO, on Fe(ll0) (173), Ru(0001) (174, 1 7 3 , and Pt(ll1) (176); of diazomethane, CHzNN, on Fe(ll0) (173) and Ru(0001) (177, 178); of diazirine, cyclic CH2N2, on W(100) (179); of CICHzI on Pt(ll1) (180); and of CHz12on Rh(ll1) (181) and Mo(l10) (182). From these, the most promising spectroscopic results, seemingly from a single hydrocarbon species, have been obtained for the reaction of ketene or diazomethane on Fe(ll0) (173) after adsorption at low temperatures followed by heating to temperatures exceeding 450 K to remove coadsorbed CO. Even simpler spectra, as expected for CHz (see next paragraph), have been obtained for ketene on polycrystalline iron at 300 K. Somewhat more complex spectra, thought to represent coadsorption of CHz and CH, have also been obtained for Fe(ll1) after a period of Fischer-Tropsch reaction of CO and H2 (for which iron is a well-known catalyst) at 300 K or after hydrogenation of carbon on Fe(ll0) at the same temperature (183). A CH2 group bonded to two metal atoms with the plane of the CHz group perpendicular to the surface would be expected (according to the MSSR) to give bands only from vCH2 s, SCH2 scissors, and vCMz s modes for on-specular VEELS. The prominent features of CHzCO on Fe(ll0) at 2970, 1430, and 640 cm-' have been assigned to these modes; with diazomethane the corresponding values are 2970,1420, and 640 cm-' (173). Additional weak bands at 1020, 930, and 790 cm-' were attributed to CH2 wagging, twisting, and rocking modes, respectively. The VEEL band positions agree rather well with the mean values obtained for the various modes of CH2M2groups, obtained from the infrared spectra of a series of organometallic cluster compounds of Os, Ru, and Fe, which have been reported as follows: ca. 2985, vCH2 as; 2940, vCH2 s; 1410, SCH2 scissors; 970, CH2 wagging; -890, CH2 twisting; 800, CHz rocking; and 655 cm-', s (184-186). vCM~ A VEEL spectrum obtained by thermal decomposition of CHzIz on Mo(ll0) (182) has also been assigned to the CH2 group. It has bands at 2920, 1320, 930, 780, 590, and 330 cm-l. The intensity pattern implies a non-Czvtilted species, but this would not be unexpected on a (110) plane. Zhou et al. attributed absorptions at 2880 and 1440 cm-I to CH2 groups formed by the decomposition of CICHzI on Pt(ll1) (180). The spectroscopic results characterizing the adsorption of ketene, diazo-

220

NORMAN SHEPPARD AND CARLOS DE LA CRUZ

methane, or diazirine on the surfaces of other metals are more complex. Ketene adsorbed on Ru(0001) at low coverage at 105 K has been suggested to give absorption indicative of surface CH2 (274, 275), but no prominent band occurs for 6CH2 near 1420 cm-', and the authors commented that an additional prominent feature at 1295 cm-' is considerably different from the expected position for this mode. Hills et al. (278) have given brief reference to bands at 2965,1155, and 785 cm-' resulting from the decomposition of H2CNN on Ru(0001). These might possibly correspond to a H2C=Ru species rather than to H2CRu2 (187). The isomeric H2C=M species in the form of the molecule H2C=Fe (187) gives, at wavenumbers less than 1500 cm-', the somewhat different band positions of 1122 (CH2 rocking), 700 (CH2wagging), and 624 cm-l (vC=Fe), without a prominent 6CH2 scissors feature. Similar VEEL data reported by Solymosi and KlivCnyi (282) for CH2 resulting from the decomposition of CH212on Rh(ll1) occur at 2940,1190,780, and 650 cm-'. These add to the possibility that a surface H2C= M species exists, but, until more spectroscopic details regarding band intensities are available, the possibility has to be considered that the 1 1 9 0 - ~ m -absorption ~ is indicative of di-a-CH2CH2. Monim et al. (179) assigned bands at 2950 and 1440 cm-' to methylene on W(100), but these occur together with strong bands from other adsorbed species. Ketene adsorbed on Pt(l1l) gave complex spectra which have not been attributed to the presence of the methylene group (176). Zaera (148, 249) postulated the presence of surface CD2 groups (2040 cm-', vCH2 s) resulting from the decomposition of CD3 groups on Pt(ll1). Solymosi and Rask6 (288) were the first to study the decomposition of CH2C12on an oxide-supported metal, Pd/Si02, by infrared spectroscopy. There had been previous nonvibrational (TPD, XPES, UPES, etc.) evidence for CH2 groups formed from the decomposition of methylene dihalides on metal surfaces (see references cited in 288). The authors assigned bands at 2984 (s) and 2907 (m) observed at 243 K to CH2groups, with this species decomposing to give another species at 293 K, with absorptions at 2922 (s) and 2857 cm-' (m). They identified the latter as ethylidyne, but the relative intensities of the absorptions more strongly suggest a polymethylene chain. Unfortunately, no data are yet available in the 6CH2 region. More VEELS or RAIRS investigations of the decomposition of methylene halides on metal single-crystal surfaces and infrared work on aluminasupported metals would be welcome.

D. SURFACE METHYLIDYNE (CH) GROUPS A CH group equally bonded to three metal atoms on an fcc (111) surface should, accordingly to the MSSR, have an allowed vCH band but a forbid-

VIBRATIONAL SPECTRA OF HYDROCARBONS

221

den doubly degenerate 6CH one. Infrared spectra of the model compounds (HC)Co3(C0)9 (189) and ( H C ) R U ~ ( C O ) (190) ~ H ~ gave these bands at 30411 850 and 2988/894 cm-', respectively. The high wavenumber of the vCH mode implies a carbon hybridization of ca. sp2, but the vCH absorption is weak, and so such a species might be difficult to find on a metal surface. No work appears yet to have been published on such a surface CH group generated from the decomposition of a haloform, HCX3 (X = C1, Br, or I), on a metal surface. However, in many cases hydrocarbons, such as ethene or ethyne, give VEEL spectra of decomposition products generated at high temperatures (>400-500 K) on fcc (111) or hcp (0001) faces which have distinct bands at ca. 3000 (m) and 760-850 cm-' (s). These have been assigned to CH groups occurring before the thermolysis of the last C-H bonds, and the relevant experimental results have been collected by Hills et al. (278).If this structural assignment is accepted, it is necessary for the C-H bond to be quasi-parallel rather than perpendicular to the metal surface, as discussed by Demuth and Ibach (191),with essentially sp2 @-bondingto two metal atoms and r-bonding to a third, or sp a-bonding to one and n-bonding to two metal atoms. There are, however, often weak broad bands present in the 1200- to 1300-cm-' region that could be poorly defined vCC absorptions indicative of alternative a(CCH) structures, probably of a polymeric form. In Section II.B.2 we speculated whether the a(CCH) species could be aromatic in nature with residual C-H bonds on a growing graphitic structure. Since this was written, a similar idea has been proposed by Yang et al. (219)in their paper on the decomposition of surface CH3groups on Ni(ll1). The decomposition occurs via CH (220 K), adsorbed HCCH (320 K), and, at high coverage, adsorbed benzene (395 K), and then a polymerized C,H, polyaromatic species with y < x. The authors attributed a simple but different spectrum to CH (CD) produced in the first stage of CH3 decomposition, with features at 2970 (2200), vCHIVCD; 1275 (930), SCH/SCD; and 650 (630) cm-', vCNi3 s. They assigned this species to a C3, hollow site, the GCH1SCD mode being allowed by impact scattering as shown by offspecular data. The unexpected feature of this assignment is the high value of the SCD/GCD mode compared with the value of ca. 875 cm-' found for the model compounds. E. SURFACE ETHYL GROUPS

Much less attention has been given to obtaining vibrational spectra of surface ethyl compared with those of surface methyl groups. Two research groups have attributed VEEL spectra to surface ethyl groups formed by the photolysis of ethyl chloride on Pt(ll1) at 160 K (292)and by the thermal

222

NORMAN SHEPPARD AND CARLOS DE LA CRUZ

decomposition of ethyl iodide on Cu(ll1) at 180 K (159). The spectra obtained are indeed readily assigned in terms of the vibrational modes of surface ethyl groups, as set out in Table V. In the Cu(ll1) case, parallel experiments involving the thermal decomposition of CH3CD21 and CD3CH21helped in the assignment of the CH3 and particularly the CH2 modes. As was found in the latter case, the directly surface-attached CH2 group gives a lower wavenumber vCH2 s mode at 2745 cm-l compared with 2918 cm-’ for this species on Pt(ll1). By analogy with the methyl case, it would be expected that an isolated ethyl group should have a plane of symmetry incorporating the CCM skeleton that is perpendicular to the surface. According to the MSSR, only the A’ modes would be active; the out-of-plane A” modes would be forbidden. With the further expectation that the C-M bond will be approximately perpendicular to the surface, it would be anticipated that the vCH3 as, 6CH3 as, and CH3 rocking modes should be prominent. All these expectations are fulfilled in the assignment given in Table V for the ethyl group on Pt(ll1). One band characterizing the species on Cu(ll1) seems best assigned to an A” CH2-rocking mode, possibly arising from symmetry distortion caused by adjacent iodine adsorbates. In an isolated or regular-array situation, this is the mode most likely to be informative about the presence or absence of a plane of symmetry in the surface complex. Lin et al. (162, 193) also obtained a similar spectrum by decomposition of ethyl bromide on Cu(lOO), but with poorer resolution in the 900-cm-’ region. This spectrum does, however, show a much lower wavenumber VCCUmode at 370 cm-I than that for vCPt at 484 cm-’. This difference between the two metals is similar to that found from the spectra of the adsorbed methyl species. Another attempt to obtain ethyl spectra from the thermal decomposition of triethylbismuth on Pt(1ll) was frustrated, probably through a more comprehensive decomposition reaction (153). A RAIR spectrum attributed to C2D5on Pt(ll1) formed from the thermal decomposition of ethyl iodide at 220-240 K gave just two absorptions, at 2182 and 2170 cm-’, in the vCD31vCD2 region (194). These authors gave reasons for being reluctant to attribute somewhat different low-coverage RAIR spectra of ethyl iodide adsorbed at 100 K to surface ethyl groups. However, on the basis of the band positions (Table V), this would have been a reasonable assignment. Once again, further single-crystal VEELS and RAIRS work would be very advantageous for the ethyl group on other metals, specifically Ni. LEED evidence for the presence of regular arrays would enable reliable symmetry assignments to be made, and PED work could probably distinguish between on-top and threefold hollow sites.

TABLE V Assignments of the Vibrational Spectra of Ethyl Groups Adsorbed on Metal Surfaces Metal surface Mode and symmetry" uCH2 as; A" vCH3 as; A ' , A" V C HS;~ A' vCH, S; A' 6CH3 as; A', A" 6CH2; A' 6CH3 S; A' CHZwag; A' CH2 twist; A" ~ C H ~ I V CA' C; pCH3; A" v C C / ~ C HA' ~; CH2 rock; A" vCM; A'

Cu(ll1) (159)* Cu(100) (162, 193)b Cu/Si02 (296)' 2935 (s) 2745 (m)

2900 (s) 2730 (m)

1430 (ms)

1420 (ms)

1140 (ms)

1140 (m)

2967 (s) 2927 (s) 2873 (m) 1453 (m) 1379 (w)

-

-

ca. 1000 (sh)

-

-

-855 (s, bd)

950 ( s ) 710 (m) -

1

Pt(ll1) (192)b Pt/SiO, (195)"" 2918 (s, bd)

1450 (sh) 1430 (ms) 1376 (ms) 1173 (ms) -

1022 (s) -

Pt/A1203 (195)"."

Pt/Si02 (197)'

2959 (m) 2923 (m) 2872 (w) (1446)

2962 (m) 2925 (m) 2878 (w) (1440)

2957 (ms) 2939 (ms) 2870 (m) 1468 (mw)

(1382)

(1381) -

1383 (ms)

I

I

T

941 (s) ?720 (sh) 370 (s)

-

484 (s)

'

a A' and A" denote in-plane and out-of-plane modes, respectively, assuming the presence of a plane of symmetry (C$). VEELS. ' Infrared transmission. Band positions in brackets denote overlap with absorptions of ethyl chloride. NA: this spectral region is not available because of strong absorption by the oxide support. s, strong; ms, medium strong; m, medium; mw, medium weak; w, weak; sh, shoulder; bd, broad; v.bd, very broad.

224

NORMAN SHEPPARD AND CARLOS DE LA CRUZ

Recently, McGee et al, (195) thermally decomposed ethyl chloride on finely divided Pt/SiO, and Pt/A1,03 catalysts. The FTIR spectrum was limited to wavenumbers >1100 cm-' because of absorption by the A1203 support, but it provides much higher resolution than VEELS, particularly in the vCH31vCH2region near 3000 cm-l. Absorption bands near 2960, 2925, and 2875 cm-' were assigned to surface ethyl groups, as listed in Table V. Driessen and Grassian (196) also reported spectra of ethyl groups on Cu/Si02 formed from thermal decomposition of ethyl chloride. The spectra are very similar to those obtained with Pt/SiO, and imply predominant bonding to single surface metal atoms rather than to three or four atoms in hollow sites, as has been deduced from VEELS work on Cu(ll1) and Cu(100). Table V also shows spectroscopic data, derived from a limited degree of hydrogenation of the adsorbed species formed from ethene adsorbed on Pt/Si02 (197),that have also been attributed to the ethyl group. There is good agreement with the results observed for species formed from thermolysis on this metal. We discuss in Section X information about the decomposition of ethyl groups obtained from the thermal evolution of such spectra.

F. SURFACE ETHYLIDENE GROUPS An attempt by Janssens and Zaera (198) to obtain a spectrum of ethylidene (CH3CH) formed from the thermal decomposition of 1,l-diiodoethane on Pt(ll1) at temperatures exceeding 150 K was unsuccessful. There was immediate decomposition to give ethylidyne at low coverage and at high coverages, to give di-cT adsorbed ethene. These findings do, however, support ethylidene as a probable intermediate in the formation of ethylidyne from ethene. They also raise the interesting possibility that at high coverages the di-cT-etheneto ethylidyne transformation might occur by a 1,2-H internal bond shift.

G.

SURFACE

1-PROPYL AND 2-PROPYL GROUPS

Vibrational spectra assigned to surface 1-propyl groups following the decomposition of 1-propyliodide or bromide have been obtained as follows, where the surface used, the halide precursor, and the decomposition temperatures are all indicated: A1(100), iodide, 310 K (199);Cu(lll), bromide, 180 K (200);Cu(lll), iodide, 180 K (201);Cu(lOO), bromide, 185 K (193); and Cu(llO), iodide, 165 K (202).The first four cases were investigated by VEELS, and the fifth one by RAIRS. The band positions observed, the band intensities, and their probable vibrational assignments are listed in

225

VIBRATIONAL SPECTRA OF HYDROCARBONS

Table VI. Recently, 1-propyl and 2-propyl groups on Pt(lll), obtained by the decomposition of the adsorbed iodides, have been studied in the vCH31 CH2region using sum-frequency generation (SFG) (202a) (see Section IX for a description of this technique). Unless there is a specific attraction between the metal surface and the end methyl group of the 1-propyl chain, the surface can be considered to be a very large substituent which would cause the trans rather than the gauche conformational isomer to be preferred. Given the expectation that the C-M bond will be approximately perpendicular to the surface, this implies that the axis of the terminal methyl group will be oriented likewise. Hence, according to the MSSR, the vCH3 s and 6CH3 s modes should be particularly prominent. This is very notably the case for the RAIRS study of adsorbates on Cu(ll0) and the SFG study on Pt(ll1). VEEL spectra are, as noted before, less reliable in the vCH3 region because of poor TABLE VI Proposed Partial Assignments of the Vibrational Spectra from trans-I-Propyl Groups Adsorbed on Metal Surfaces Metal surface Mode vCH3 as

Al(100)"

PCH3

vcc

vCM

I 1

2945 (s, sh)

CU( 111)'

I

2950 (s, bd)

2935 (s, bd)

2740 (m) -1450 (sh)

2730 (ms) 1445 (s)

1385 (ms)

1385 (s)

1360 (vs)

1160 (w) 1000 (mw) 804 (s, bd)

1165 (mw) 1055 (mw) 880 ( s )

1150 (m) 1010 (w) -800 ( s )

605 (s)

720? -

v C H s~ v C H ~ Ms 6CH3 as 6CH2 6CH3 s CH2 wag CH2M wag

Cu(11l)b

I

t

-600 (s, v.bd)

1

Cu(l10)d

Cu(100)'

I

2954 (w) {2926(s)f 2865 (s)

2925 (s, bd)

I

2690 (m) 1457 (w)

1371 (vs) 1228 (w) 1124 (w)

T

}l400 (s, bd)

1135 (w) ca. 1020 (m)

T

1

i*'

910 (v.bd) ~~

C3H71,310 K, VEELS (199). C3H7Br, 195 K, VEELS (162, 200). C3H71, 180 K , VEELS (201). C3H71,165 K , RAIRS (202). C3H7Br, VEELS (193). On the basis of the wavenumber this could be from CH2as, but this assignment contradicts the finding that the 1-propyl group has the trans conformation determined on the basis of the observation of strong bands of vCH3 s and SCH3 s. g N A region not available. a

226

NORMAN SHEPPARD AND CARLOS DE LA CRUZ

resolution and a tendency for impact-excited modes to give measurable features on-specular. However, in all four cases, the band near 1375 cm-l, associated with SCH3 s, is very prominent. In the Cu(ll1) case, low-wavenumber bands, indicative of the vCH2 s modes of the CH2M group, are present, as has already been discussed for the methyl and ethyl cases. It can be concluded that the assignment of these spectra to the frans form of the 1-propyl group is very satisfactory. The RAIR spectrum of 2-propyl, formed by the thermal decomposition of 2-propyl iodide on Cu(ll0) at 170 K, is a surprising one (202).It consists of prominent uCH bands at 2893,2880, 2843, and 2822 cm-’, with a weak feature at 2972 cm-’. In the low-wavenumber region, the spectrum is dominated by a strong band at 1110 cm-’, with broad and weak absorptions at perhaps 1420 and 1360 cm-l. A closely similar v C H ~ I V C H spectrum was obtained on Pt(ll1) using SFG. If the 2-propyl surface species has its C-M bond perpendicular to the surface, the modes with dipole changes approximately perpendicular to the surface would be vCH3 as, SCH3 as, and CH3 rocking modes plus vC-M. One of the bands at wavenumbers less than 2900 cm-I could be evidence of the lone vCH mode, but the others are typical of vCH3 s (not vCH3 as) plus overtones brought up by Fermi resonance with this. Only the very weak band at 2972 cm-’ can correspond to vCH3 as. According to the MSSR, assuming C, symmetry for the complex, there is in fact only one completely symmetrical ( A ’ ) vCH3 s mode but two 6CH3 as modes. Equal numbers of further vCH3 s and vCH3 as modes are of symmetry A” and therefore MSSR forbidden. The spectrum as observed would be more consistent with a situation in which the plane of the C3 skeleton is approximately perpendicular to the surface, as in M2C(CH3)2.A lower resolution spectrum obtained by VEELS by decomposition of 2-propyl bromide on Cu(100) at 195 K shows similar features (193). This structural ambiguity requires further investigation. In particular, the 2-propyl species should be generated on other metal surfaces and studied spectroscopically. H. SURFACE n-BuTYL AND ISOBUTYL (2-METHYLPROPYL) GROUPS A RAIR spectrum of the n-butyl surface species on AI(100) has been reported at 335 K in the vCH31vCHzregion (203). A vCH3/vCH2 RAIR spectrum has been reported for the isobutyl group formed by the decomposition of triisobutylaluminum on Al(100) at 335 K (203), and a VEEL spectrum has been obtained from decomposition of the trialkylaluminum on Al(111) at 100 K (204). These alkyl surface species are stable to 450-500 K and then decompose to give the expected alkenes by P-H elimination.

VIBRATIONAL SPECTRA OF HYDROCARBONS

227

I. TERTIARY BUTYLGROUPS A VEEL spectrum of the species formed by the decomposition of tertbutyl (2-methyl-2-propyl) chloride and bromide on Cu(100) at 185 and 120 K, respectively (193), shows intensity anomalies similar to those discussed for 2-propyl groups when interpreted in terms of surface tert-butyl groups. Further RAIRS and VEELS work is required.

J. TRIMETHYLENEDIMETALLO GROUPS A VEEL spectrum characterizing the products of the thermal decomposition of 1,3-diiodopropane on Al(100) at 300 K has been reported (205). It probably derives from the presence of the five-membered (CH2)3M2ring skeleton at the surface. The spectrum consists of prominent bands which can reasonably be interpreted as follows: 2870 cm-' (s), vCH2s; 1420 cm-' (ms), CH2 scissors; 1160 cm-1 (ms), CH2 wag or twist; 101.5 (w), 890 cm-' (ms), vCC; 720 cm-' (s), CH2 rock; 590 cm-' (s), vCAI; and 390 cm-* (vs), vA1-I. In agreement with the dimetallocyclopentane formulation, other characteristic absorptions of CH3 groups, such as vCH3 as or SCH3s, were not observed, at least at VEEL resolution. Propene was desorbed at 510 K. ALLYLGROUP K. THE SURFACE A VEEL spectrum has been recorded for the products of the decomposition of allyl bromide (l-bromopropene) on Al(100) at 310 K (205). The spectrum indicated that the surface species retained its C =C double bond and was not a-bonded to the surface (vC=C, 1655 cm-'), but more detailed conclusions about its conformation could not be drawn. The decomposition of allyl chloride on Ag(ll0) (206) at 180 K has yielded a VEEL spectrum reasonably assigned to a a-ally1 species; at 300 K, adsorbed, 1,5-hexadiene has been identified. Similar results have been reported for Ag(llO)/O (207) together with other oxygenated products. L. SURFACE VINYLGROUPS The VEEL spectrum of a surface vinyl group derived from low-temperature decomposition of vinyl iodide on Pt(ll1) has been reported with features at 2920 cm-I (s), vCH; 1600 cm-' (w), vCC; 1380 cm-' (m), S= CH2 scissors; 1255 cm-' (mw), SC=CH; 955 cm-' (s), y=CH2; and 690 cm-' (mw), y=CH2 twist (208).The strength of the out-of-plane y=CH2 mode shows that the plane of the vinyl group is more nearly parallel than perpendicular to the surface, probably cr-bonded to one metal atom and a-bonded to another. This inference is confirmed by the relatively low wavenumbers of the vCH modes. However, in such a situation, the vCC wavenumber

228

NORMAN SHEPPARD AND CARLOS DE LA CRUZ

might have been expected nearer to 1500 than to 1600 cm-', with both the (T- and n--bondingto metal atoms taken into account. Possibly the a-bonding is weaker because of steric constraints. The vinyl group has been suggested to participate in decomposition reactions of ethene on metal surfaces, and so this is an important reference spectrum. It would be profitable to study spectra of such species on other metals. Zaera and Bernstein (209) recently attributed RAIRS bands at 1210, 1190, and (presumbly) ca. 950 cm-l of species formed from vinyl iodide adsorbed at low coverage on Pt(ll1) at 100 K to surface vinyl groups, in approximate agreement with the earlier VEEL data. They also suggested decomposition pathways of the vinyl group at higher temperatures, on the basis of the RAIRS data. M. HCCH SURFACE GROUPS Grassian and Pimentel (210) prepared such surface groups by thermal decomposition of cis- and trans-dichloroethenes at temperatures >200 K or by their photolysis at 110 K on Pt(ll1). An ethyne type B spectrum was obtained, as had also been obtained from the direct low-temperature adsorption of ethyne on this surface (Section II.B.l).

N.

SURFACE PHENYL GROUPS

VEEL spectra of surface phenyl groups have been obtained by ultraviolet photolysis of phenyl chloride on Ag(ll1) at 300 K (211) and by thermal decomposition of phenyl iodide on Cu(ll1) (212). The spectra are similar, and the strengths of the absorptions from the out-of-plane yCH modes, at 740 and 725 cm-I respectively, show that the phenyl group is also more near to parallel than to perpendicular to the metal surfaces, again probably because of a combination of (T- and a-bonding.

0. SURFACE CYCLOPROPYL GROUPS Martel et al. (213) reported what is considered to be a VEEL spectrum of the cyclopropyl group from electron bombardment of cyclopropane on Cu(ll0). P. SURFACE CYCLOHEXYL GROUPS

Xu and Koel (214) obtained what is believed to be the VEEL spectrum of a-bonded cyclohexyl by the electron bombardment of cyclohexane on Pt(ll1).

VIBRATIONAL SPECTRA OF HYDROCARBONS

V.

Cycloalkanes

A.

CYCLOHEXANE

229

1. Single-Crystal Work

We commence our discussion with a consideration of the vibrational spectrum of the much-studied cyclohexane molecule when adsorbed on various metal surfaces. The majority of published papers have been concerned with adsorption on metal single-crystal faces. VEEL spectra measured at low temperatures have been obtained for cyclohexane-derived species on Cu(ll1) (215-217), Ni(111) (7, 75, 218), Ni(l1O) (216), Ni[5(111) x (?lo)] ( 7 ) ,Pd(11l) (219),Pd(ll0) (216, 220), Pt(ll1) (75, 214, 221-223), Pt(llO)(l X 1) and hexagonal (224), and Ru(0001) (225,226).Higher resolution RAIR spectra, limited to the range 4000-800 cm-' and sometimes to the vCH region, have been described for cyclohexane on Cu(111) (191, 215, 217), Cu(100) (227), Ni(ll1) (228), Pt(ll1) (138, 229), Pt(100)(1 x 1) (230), and Mo(1lO) (230a). We first review the results characterizing the close-packed surfaces, i.e., fcc (111) or hcp (0001). All the VEEL spectra have a strong band near 2900 cm-'( vCH2 as) and a very broad and strong band, variable in position between 2750 and 2550 cm-', which is assigned to "soft-mode" vibrations of C-H bonds involved in agostic (hydrogen-bond-like) interactions (231) with surface metal atoms. In addition, medium-strength features occur near 1450, 1030, and 520 cm-l. All these, but particularly the soft-mode and ca. 520-cm-' bands, are reduced in relative intensity in off-specular spectra (222, 225),which implies that they are evidence of MSSR-active fully symmetrical modes. Additional features occur rather consistently in the VEEL spectra near 1260, 870, and 370 cm-', arising from impact-excited modes. The RAIR spectra representing adsorption on the close-packed C u ( l l l ) , Ni(l l l ) , and Pt(l1l) surfaces are very similar in profile and differ mainly in the variability in position of the soft-mode absorption. For Ni(l11) and P t ( l l l ) , the band positions at low coverage are nearly identical at 2920 (m, sh), 2903 (s), ca. 2890 (m, sh), 2845 (m), ca. 2700 (s, v.bd), 1445 (w), and 1030 cm-' (w). For C u ( l l l ) , the vCH values are somewhat higher, at 2926 and 2770 cm-I for the strong features. In each case, at higher coverages (but still less than a monolayer), the 2920-cm-' shoulder grows in parallel with additional absorptions at 2855,1455, and [for C u ( lll) ] 860 cm-', but without any further growth of the soft-mode absorptions. These results have been taken to signify a change in phase involving either tilted molecules in contact with the surface (123, 215) or the growth of a partial second layer before completion of the monolayer (228). The absorptions of the initially formed species can be assigned as 2903

230

N O R M A N SHEPPARD AND CARLOS DE LA CRUZ

( vCHz us), 2890 (nonagostic vCH), 2845 ( vCHz s), ca. 2700 (agostic vCH), 1445 (SCH: scissors), and 1030 cm-' (CH2 rocking). Together with the dipole-excited additional VEEL feature at 520 cm-' (SCCC), this list is essentially identical with that of the infrared-active modes of the isolated cyclohexane molecule (232),except for the additional agostic uCH feature. All these features agree with MSSR expectations (the selection rule is strictly applicable in RAIRS and greatly simplifies the spectrum to be expected) for an adsorbed cyclohexane complex of C3vsymmetry with the undissociated molecule in the usual chair form and with the median plane of the C6 skeleton parallel to the surface. Raval (218) has drawn the same conclusion from a RAIRS study of C(,HD,, on Ni(ll1). This result is in agreement with LEED studies of species on Pt(ll1) (233) and Ni(ll1) (234) and with an electron stimulated desorption ion angular distribution (ESDIAD) (235) investigation of Ru(0001). An angle-resolved ultraviolet photoelectron spectroscopy (ARUPES) investigation of Ni( 111) (234) is in good agreement but suggests that there may be slight tilting from parallelism caused by steric interactions between close-packed molecules in the ( d 7 X d 7 ) R19.1" surface array. There is general agreement that, as initially considered by Demuth et ul. ( 7 9 , the agostic phenomenon involves the three axial C-H bonds on one side of the ring interacting with the metal surface. The spectroscopic phenomena (a lowering of vCH and the very substantial broadening and intensification of the band) are closely similar to those associated with hydrogenbonding interactions of the O-H...O type in vibrational spectra such as those of alcohols. In the latter cases, thc interaction is between an electrondeficient H atom on t h e OH group with the electron-rich oxygen atom of another alcohol molecule. In the present case, the interaction would seem to be between a comparatively electron-rich H atom with the electrondeficient surface. It probably occurs in the form of a synergetic interaction between electron donation from the C-H bond into d orbitals of the metal atoms and back-donation from ri orbitals into an antibonding CH orbital. Each of these processes could contribute to the weakening of the C-H bonds indicated by their substantially lowered wavenumbers. On Mo(ll0) (230a) coadsorption with sulfur leads to an enhancement of the tilted species from cyclohexane with a reduction of the soft modes. This has been interpreted as implying that the latter originate from metal d-orbital donations to C H antibonding orbitals. However, i t seeins possible that steric effects from the presence of the coadsorbate sulfur may alternatively be a factor. If the threefold axis of the adsorbed cyclohexane passes through a threefold hollow in the (1 11) surface, each of the axial C-H bonds could interact with a surface metal atom. The separation between pairs of axial C-H

VIBRATIONAL SPECTRA OF HYDROCARBONS

231

bonds closely matches the nearest-neighbor metal-metal bond distances. Alternatively, the threefold axis could pass through a metal atom, with the C-H bonds interacting with a set of either fcc or hcp hollows. In each of the earlier LEED investigations, the authors assumed the latter case, but apparently without consideration of the alternative. The spectroscopists have often preferred the former structure, and to date theoretical ub initio calculations (236-238) do not seem to provide a clear-cut decision. Further LEED work done with intensity-voltage measurements to determine the distance between the median plane of the C6 skeleton and the surface plane of the metal would seem to be important for distinguishing between the cases of on-top or threefold hollow sites for the agostic interactions of the C-H bonds. The former should require a greater metal surface to C6 plane distance. Direct distinctions between the alternative hollowsite models and the twofold bridging sites may require the use of the more recently developed photoelectron diffraction technique. The low-temperature VEEL spectra of species formed from cyclohexane on Ni(ll0) and Pd(1lO) planes show the same features, including the soft modes, as on the corresponding (111) planes, implying that the Ch skeleton is again approximately parallel to the surface. In principle, C3Lsymmetry is not possible for the surface complex in this case, GCbeing the highest possible symmetry among surface sites. The interactions with the surface are clearly not strong enough to distort the symmetry of the adsorbed cyclohexane appreciably except via the localized agostic interactions. The VEEL spectrum of species formed from cyclohexane on Pt(100) (1 X 1) shows that soft-mode absorption is either absent or much weaker (ca. 2600 cm-I?), and there are more marked differences in the fingerprint region at wavenumbers less than 1500 cm-'. However. the intensity distribution in the main bands still implies approximate parallelism with the surface. The more detailed RAIR spectrum in the vCH region still retains some similarity to the spectrum of the species on Pt(ll1) but with more prominent low-coverage features at 2935 and 2856 cm-'. Once again, there is a similar phase change on increasing coverage within the monolayer. The soft-mode absorption of cyclohexane on the stepped Ni[S(ll 1) X (710)] surface is clearly present at ca. 2620 cm-', but with much reduced intensity compared with that on Ni(ll1) (7), implying, as expected, less convenient CH contacts with the metal surface. The band positions in the fingerprint region are similar, but notable changes in relative intensities, e.g., of the prominent ca. 520-cm-' absorption, suggest a nonparallel orientation with respect to the surface. Indeed the on-specular spectrum of the species on Ni[S(111) X ( i l O ) ] bears a remarkable resemblance to the offspecular spectrum of the species on Pt(11l) (222) whereby the modes with dipole changes parallel to the surface, i.e., parallel to the median C, plane,

232

NORMAN SHEPPARD AND CARLOS DE LA CRUZ

would be dominant in the spectrum. The triangular metal atom arrangement on this close-packed surface is clearly particularly favorable for agostic interactions. Cooper et al. (228) also studied the adsorption of cyclohexane on the Ni(ll1) surface with preadsorbed oxygen. The absorptions of the soft modes of the species Ni(111)(2 X 2)/O sharpen, move further down to 2613 cm-', and show several resolved components. This spectrum implies an increased strength of agostic interaction as would be expected if the coadsorbed oxygen reduced the electron density of the surface metal atoms so as to promote more aCH -+ d-orbital bonding at the expense of d + antibonding aCH electron transfer. This experimental finding favors the view of Kang and Anderson (236) that the former rather than the latter is more responsible for the weakening of the C-H bond during agostic interactions, a view not shared by other theory groups (237, 238). At higher oxygen coverage, ( d 3 X d 3 ) R30"/0, no agostic interactions remain, presumably for reasons of site blockage. The same is true of oxygen preadsorption on Ru(0001), as shown by Hoffmann and Upton (226), and is cited as evidence in favor of C H interactions with hollow rather than on-top sites. The latter view is also favored by the authors' theoretical calculations for Ni(l1l) and Ni(100) surfaces, although the conclusion that interaction would be stronger with the fourfold sites on the (100) surface does not appear to be supported by the experimental results described earlier for Pt(100) compared with Pt(l1l). The relative strengths of the agostic interactions of the C-H bonds with different metal surfaces can be semiquantitatively evaluated by the lowering of the agostic vCH wavenumbers relative to the value of 2885 cm-', which is the mean value between the as and s bond-stretching modes of alkane CH2 groups. For the close-packed surfaces, these lowerings are approximately as follows: Cu, 115 cm-'; Ni, 155 cm-'; Pd, 250 cm-'; Pt, 285 cm-'; and Ru, 305 cm-'. The chemical significance of these data [as originally pointed out by Demuth et al. (75)]is shown by the fact that on warming, the cyclohexane molecules desorb from Cu(ll1) and Ni(11l) without decomposition, but on P d ( l l l ) , P t ( l l l ) , and Ru(0001) they decompose on the surface, leading ultimately to surface-bound benzene. The (100) surfaces of Cu and Pd behave in this respect in the same way as the (111) surfaces, but the highindex [5(111) X (?lo)] plane of Ni also leads to dehydrogenation to benzene whereas this is not the case for Ni(ll1). On most surfaces, decomposition to benzene has occurred by ca. 300 K, and this is a significant factor in relation to the later discussion of room-temperature spectra of species formed from cyclohexane adsorption on oxide-supported metal catalysts. On Ru(0001), decomposition to benzene has occurred by 230 K, and on Ni[5(111) X (?lo)] by 225 K. The latter result implies that high-index rough surfaces are much more reactive in this respect.

VIBRATIONAL SPECTRA OF HYDROCARBONS

233

The VEEL spectra of the species formed from cyclohexane on Pt(ll1) show that at least two intermediate species occur along the decomposition pathway to benzene. These spectra are discussed in Sections V1.A and VI.C, in the context of spectra of species formed from adsorbed cyclohexene (239) and cyclo-1,3-hexadiene (240) on the same surface. On Pt(100) hex, in contrast to Pt(lll), most of the cyclohexane molecules desorb before conversion to benzene, but the latter was formed after adsorption at 300 K. An intermediate in the conversion of cyclohexane into benzene on Pt(100) (1 X l),stable between ca. 200 and 300 K, was recognized spectroscopically, but not structurally identified, by RAIRS (230) and by VEELS (224). It seems that there is a smooth transition from the spectrum of adsorbed cyclohexane on Pd(100) to that of benzene at temperatures exceeding 250 K without the detection of intermediate spectra (220). The vibration frequency of cyclohexane against the Cu(ll0) surface has been measured at 175 K by helium atom scattering (139). 2. Finely Divided Surfaces Infrared spectroscopic investigations of cyclohexane adsorbed on metal oxide-supported metals have been carried out in two contexts, namely, the initial adsorption of cyclohexane itself (86,241-247) or the initial adsorption of benzene followed by the addition of hydrogen to give cyclohexane (87, 248-251 ). All these spectra were observed at ambient temperature (ca. 300 K) or higher, under which conditions the single-crystal work indicates that cyclohexane on either Ni or Pt surfaces is dehydrogenated under vacuum to give benzene n-bonded to the metal surface. This was the early conclusion of Ward (86) (Fig. 8A) and of Palazov et al. (241-244) for cyclohexane adsorbed on Ni/Si02 and has been supported by much subsequent work with oxide-supported Ni or Pt on Si02 or A1203(Fig. 8B). The absorption bands of benzene, derived from the dehydrogenation of cyclohexane, occur at ca. 3040 (vCH aromatic) and 1390 cm-' (skeletal ring deformation) (247). Only Erkelens and Eggink-Du Burck (249) failed to observe the vCH absorption. Instead, on Ni/Si02 they observed a very broad and weak absorption between 3060 and 2760 cm-' after evacuation, which led them to postulate a mixture of adsorbed species including sp3, sp2, and n-bonded components. Sheppard et al. (see later, Fig. 10) showed early that there is a cyclohexane-benzene equilibrium on Pt/SiO, depending on the presence or absence of gas-phase hydrogen. Haaland (247;Fig. 8C) has most recently studied the cyclohexane on Pt/ A1203system in great detail, and our detailed discussion will principally relate to his work. The spectrum after evacuation showed a greater degree of conversation to benzene than in Fig. 8B, and this grew further on prolonged

234

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3100

2900

1500

1300

110

cm-'

FIG.8. Infrared spectra of cyclohexane adsorbed at room temperature on several metals followed by evacuation: (A) Ni/Si02; (B) Pt/Si02; (C) Pt/A1203.[(A) and (B) from Ref. 86; (C) reprinted from Ref. 247, Sur$ Sci. 111, D. M. Haaland, p. 555. Copyright 1981 with kind permission of Elsevier Science-NL, Sara Burgerhartstraat 25, 1055 KV Amsterdam, The Netherlands.]

evacuation. Initial adsorption before evacuation gave bands from undissociated cyclohexane at 2927 (vCH2as), 2852 (vCH2s), and 1450 cm-l (6CH2 scissors); the counterparts of the 2927- and 1450-cm-' absorptions of C6HI2 occurred at 2208 and 1085 cm-' in the spectrum of ChD12.After evacuation there were detailed differences between the spectrum obtained from chemisorbed benzene when the latter was adsorbed on a virgin Pt/A120, sample, i.e., one that had not previously been exposed to a hydrocarbon, and when it was derived from the adsorption and then evacuation of cyclohexane. The differences are that the vCH band from the 7i-bonded benzene complex on Pt/A1,0, occurs at 3050 (2281 from C6D6)cm-' on the virgin sample but at ca. 3031 (2260) cm-' if cyclohexane adsorption had previously taken place. Also, the latter spectrum consistently shows an additional weaker absorption at 2947 (2200) cm-', which varies in intensity relative to the 3031 (2260) cm-' band, depending on the experimental conditions. Prentice (87) obtained similar spectra in the vCH region from benzene adsorbed on mature Pt/Si02, i.e., a catalyst that had previously held adsorbed hydrocarbon species but that had been regenerated by reduction in hydrogen. Additional absorptions at 1392 (1268), 1272, and 1147 cm-', observed after evacuation (Fig. SC), are attributed to adsorbed n-bonded benzene on both

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types of sample. The A1203absorption cutoff at ca. 1100 cm-' precludes the observation of absorption bands at lower wavenumbers. In view of its position, Haaland attributed the 2947-cm-' band to the presence of an additional a-bonded species, which he demonstrated exchanges deuterium for hydrogen from the OH groups of the alumina support more rapidly than does the n-complex. H e attributed the shift in wavenumber of the vCH band from the benzene n-complex to the copresence of the a-bonded species. He also showed that thermal desorption more readily removes the latter, with a parallel shift of the 3031-cm-' band back to 3050 cm-'. In relation to the structural origin of the a-bonded species, we recall that a sharp band at 2941 cm-' was observed from the thermal decomposition of cyclohexane on a Pt(100)(1 X 1) surface (230). Both the vCH bands characteristic of the adsorbed species on Pt/A1203have considerable widths, probably due to a range of slightly different adsorption sites on the metal particles. Indeed partially resolved structure in the vCH absorptions of the n-complexes was attributed by Haaland to adsorption on different facets of the metal particles. As shown by Sheppard et al. (248;see also further discussion in Section VIII.A.2), the readdition of hydrogen to Pt/Si02 transforms the adsorbed benzene back into gas-phase plus physically adsorbed cyclohexane and possibly, in view of the persistence of the CH2 absorptions at 2922, 2850, and 1450 cm-' after multiple +H2/-H2 cycles, into some species with the cyclohexane ring attached to the surface by one or two C-M a-bonds (see Section V1.A on cyclohexene adsorption). The desorption during evacuation of large physically adsorbed hydrocarbons from the pores of metal oxide-supported metal catalysts can be a slow process. The absence of absorptions indicative of CH3 groups after hydrogenation shows that the n- and a-bonded species retain cyclic C6 skeletons. Reiicha also reported absorptions at 2940 and 2870 cm-' after room-temperature adsorption of cyclohexane on Ni/Si02 (252). Yates et al. (253-255) reported the generation of spectra with absorptions at 2937 and 2864 cm-' from a-bonded cyclohexyl groups by ultraviolet photochemical reaction of cyclohexane on Rh(C0)2/A1203catalysts. It is not clear whether this cyclohexyl species, which retains some thermal stability up to 600 K, is adsorbed on the Rh or A1203 surfaces. B. CYCLOPENTANE

1. Single-Crystal Work VEEL spectra of cyclopentane on Ru(0001) were observed by Hoffmann et al. (226, 256), and those of cyclopentane on Pt(l1l) were observed by Avery (257). Chesters and Gardner (138) obtained a RAIR spectrum in

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the vCH region of cyclopentane on Pt(11l). Avery (258) also reported, on the basis of VEEL evidence but as yet without published detail, that cyclopentane adsorbs intact on Ir(l1l) at low temperatures and desorbs without reaction on warming. The low-temperature VEEL spectra of the species on Ru(0001) at 170 K and on Pt(ll1) at 90 K are similar in contour and show broad and strong soft-mode absorptions at ca. 2610 and 2690 cm-', respectively. The strengths of these features imply that the median plane of the flexible C5 skeleton is approximately parallel to the metal surface so that multiple C-H...M contacts are once again possible. The RAIR investigation at ca. 140 K of cyclopentane on Pt(ll1) gives high resolution and shows the soft-mode absorption to be very broad. The positions of the nonagostic vCH absorptions and of the VEELS features at wavenumbers less than 1500 cm-' are as expected for the intact, nondissociatively adsorbed species (257, 259). In each case, as the temperature was raised to 200 K, a markedly changed VEEL spectrum was observed, which was attributed to the formation of cyclopentene by dehydrogenation (see Section V1.B). Avery's study of the adsorption of cyclopentane was continued to 260 K, whereby a much simpler spectrum was obtained, convincingly attributed to the formation of the v5-C5H5($-cyclopentadienyl) structure adsorbed flat on the surface. The 200 K spectrum of the species on Ru(0001) may even contain some features characteristic of this species (strong bands at 758 and 3057 cm-I). Preadsorption of oxygen on Ru(0001) led to a spectrum without softmode vCH absorptions.

2. Finely Divided Catalysts Ward (86) investigated the room-temperature adsorption of cyclopentane on Ni/Si02 and Pt/SiO, in the vCH region (Figs. 9A and 9B). After prolonged evacuations, only weak spectra were retained for the species on NiiSiO, in the vCH region, with absorptions near 2955 and 2870 cm-' typical of Cs-ring CH2groups (259 and ref 110, spectrum E2-2). The spectra of the species on Pt/SiO, were more complex (Fig. 9B), with probably at least two species present, characterized by sharp absorptions at 2956 and 2850 cm-', and broader, weaker bands at 2975,2920, and 2800 cm-', respectively. The gas phase after hydrogenation was principally cyclopentane. The single-crystal work done with Pt( 111) implies that on this metal much of the originally adsorbed cyclopentane should have been converted to vsCsHsat temperatures as low as room temperature. Although no absorptions were recorded at wavenumbers >3000 cm-' to support this inference, it should be recalled that the features in the vCH region in the VEEL spectra were mainly impact induced and that the MSSR could cause the dipolarexcited feature characteristic of this species that is active in the infrared

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3100

2900

cm-'

2900

237

2700

FIG.9. Infrared spectra of cyclopentane adsorbed on (A) Ni/Si02 and (B) Pt/Si02 and of cyclopropane adsorbed on (C) NilSi02, H-depleted, and (D) Pt/SiOz, H-depleted. [From Ref. 86.1

region to be very weak in the spectrum of the parallel-adsorbed surface species. Large intensity increases were observed on hydrogenation of the species on each metal caused by the generation of gas-phase and physically adsorbed cyclopentane with possibly some cT-bonded cyclopentyl species. The large intensification is now seen to be probably caused by the conversion of the very weak CH absorption of the planar C5H5species formed after evacuation at ambient temperatures to the much more strongly absorbing CH2 groups of cyclopentane. Atypically in comparison with the hydrogenation of the species formed in the adsorption of acyclic alkenes (Part I), it was Ni rather than Pt that gave the higher proportion of the gas-phase product. A single hydrogenation removed most of the adsorbed species from Ni/Si02as cyclopentane. Cyclopentane was again the principal gas-phase product formed from Pt/Si02, but several dehydrogenation/hydrogenation cycles were needed to remove most of the surface species. C. OTHERCYCLOALKANES 1. Single-Crystal Work Hoffmann and Upton (226) measured the VEEL spectra of the series of C,H2, cycloalkanes (n = 3-6, 8) on Ru(0001) in the vCH region. They found no sign of soft-mode phenomena for cyclopropane or cyclobutane and only weak such features for cyclooctane. The latter result was confirmed by Hostetler et al. (260) using RAIRS. The interbond C-C-C angles of 60" and ca. 90" for n = 3 and 4 would not provide convenient geometric requirements for multiple C-H...M interactions. The C6 ring undoubtedly

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provides the best situation for these, and the flexibility of the Cs ring provides another favorable situation. In the case of cyclooctane, it is clear from the spectrum that the majority of the CH2 groups cannot interact with the surface. These authors and also Avery (257), pointed out that there is a smooth increase of monolayer desorption temperatures with increasing n and that there do not seem to be anomalously high values for n = 5 and 6 to correspond to the presence of strong soft-mode features resulting from C-H-..M interactions for these two cycloalkanes. In an earlier review ( 4 ) , one of us suggested that these interactions, seemingly significant in terms of surface reactivity, may be partially balanced by greater steric interactions through closer contact with the metal surfaces. This would be particularly likely if the agostic interactions in cyclohexane proved to be with the threefold hollow sites. Complete VEEL spectra for cyclopropane adsorbed on Ru(0001) (261, 262) and Cu(ll0) (213) have been described and interpreted. These two spectra are similar in profile, including absorption bands at positions characteristic of in-plane modes such as SCH2 scissors, ca. 1470; CH2 wag, ca. 1030; C3breathing mode, ca. 1190; and C3 deformation, ca. 860 cm-’ (263). All these, except for the C3 breathing mode, are active in the infrared spectrum of cyclopropane itself in the gas phase and give dipole changes parallel to the C3 skeletal plane. Taking into account the MSSR of these, we infer that only the C3 breathing mode would be active if the molecule retained a C3 axis and was adsorbed with the C3 skeleton parallel to the metal surface. In fact, in the Ru(0001) case, the strong features from the coupled CH2 wagging and C3 deformation modes have been shown to be particularly dipole-active from off-specular measurements (261), and this strongly suggests that the C3 plane is at a high angle with respect to the surface. In each case, absorptions in the 3080- to 2980-cm-’ range [impact excited on Ru(0001)I are also consistent with the presence of an intact cyclopropane skeleton. Finally, there has been considerable discussion of the vibrational origin of a band at 570 cm-’ obtained for species on the Ru(0001) surface (261). If experimentally verified, there seems little doubt that this has to be a frustrated translation mode of the molecule perpendicular to the surface, as it is also highly dipolar in character. Its wavenumber is higher than might have been expected for a nondissociative bonding of cyclopropane to the surface. However, it should be noted that Zeise’s salt, involving r-bonding of ethene to Pt, does have a C 2 - - - Mvibrational mode near 500 cm-’ (264). The vibrational spectra are hence largely consistent with a structure in which one of the protruding “banana” bonds of the cyclopropane molecule interacts with a metal atom by synergic bonding analogous to that for n--bonded ethene, so that a C-C bond is essentially parallel, and the C3plane perpendicular, to the surface. However, the analysis of ARUPS

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239

results for species on Ru(0001) indicates that the symmetry of the adsorption complex is C, (c,) [or C,] rather than C,,, which might have been expected for the C3-perpendicular structure. In fact, however, the symmetry of such a complex, with the C3plane perpendicular to the surface, is strictly C, because the six hollow sites around a metal atom are in two different sets of three, one of which has another metal atom in the second layer. VEEL spectra have been published of methylcyclohexane on Pt( 111) (265) as a function of temperature. The monolayer at 165 K shows a very strong soft-mode centered at 2475 cm-', as expected by analogy with cyclohexane itself. At 295 K this had disappeared, and the new surface species was suggested to be a-allylic in character; it could alternatively involve c-bonding to the surface. As the temperature was ramped to 450 K, the spectrum transformed into a new spectrum with a strong band at 825 cm-'. The latter was interpreted as the yCH mode of a benzyl group lying approximately parallel to the surface. 2. Finely Divided Surfaces The adsorption of cyclopropanes at room temperature has been characterized by infrared spectroscopy for a number of silica-supported catalysts, viz., Ni (86),Pt (86),Pd (266),and Rh (91).The spectra are identical with those obtained from the adsorption of propene on the same metals. They give absorptions from CH3groups showing that the C3ring has been opened, and the nature of the spectra has already been discussed (140, and Part I, Section VI.C.1.b). Typical spectra of species formed from cyclopropane on Ni/Si02 and Pt/Si02, obtained by Ward at room temperature, are shown in Figs. 9C and 9D. Identical spectra have been reported for species formed from cis- and trans-1,2-dimethylcyclohexanesand o-xylene adsorbed on Ni/Si02at room temperature (242). A weak band at 3010 cm-' was assigned to aromatic vCH and one at 2915 cm-' to residual methyl groups on the aromatic ring. Hydrogenation led to regenerated dimethylcyclohexanes. The dehydrogenated aromatic species formed after reevacuation was considered to be cbonded to the surface, retained through dissociation of one of the CH3 groups. VI. Cycloalkenes A. CYCLOHEXENE

1. Single-Crystal Surfaces VEEL spectra have been obtained at low temperature for species formed from cyclohexene adsorbed on Pt(ll1) (239),Pd(ll0) (220),Ag(ll0) (267),

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NORMAN SHEPPARD AND CARLOS DE LA CRUZ

and Ag(llO)/O (268). A RAIR spectrum of the species on Cu(100) in the vCH region was published by Teplyakov and Bent (269) in connection with a study of the hydrogenation mechanism. By analogy with the species found from the adsorption of ethene on these surfaces, it would be anticipated that nondissociatively adsorbed cyclohexene would be di-a-bonded to Pt and .rr-bonded to Pd, Cu, and Ag. In agreement with these expectations, the VEEL spectra of the species on Pd(ll0) and Ag(ll0) are closely similar to each other and quite different from that of the species on Pt(ll1). The vCH RAIR spectrum of the Cu(100) species has also been interpreted in terms of a s-complex. The VEEL spectra of the species on the (110) faces exhibit all the features expected for a n-bonded species approximately parallel to the metal surface, i.e., showing vCH modes at wavenumbers above and below 3000 cm-’ ( = C H and CH2, respectively), an absorption at ca. 1630 cm-l ( v C = C ) [Ag (llo)], and very strong bands between 680 and 660 cm-’ (yCH=CH) (see 270). The quite different VEEL spectrum of the species on Pt(ll1) after adsorption at 95 K has prominent bands at 2890, 1450, 1080, 820, and 520 cm-’ (239),a pattern which is reminiscent of the spectrum of cyclohexane adsorbed on the same surface. It is therefore reasonably interpreted as evidence of the expected cis-di-a adsorbed structure with the C6 skeleton approximately parallel to the metal surface. At 200 K a new spectrum was observed, with a notably different profile in the fingerprint region (in particular, the strong bands at 820 and 550 cm-’ had disappeared). This spectrum is still, however, considered by Henn et al. (239) to be evidence of a C6HI0surface species because Bi-postdosing thermal desorption spectroscopy (BPTDS) still only gave gas-phase cyclohexene at this temperature. The 95 and 200 K spectra were attributed to LY- and /3-adsorbed cyclohexene. The intensity changes between these two, involving the disappearance of the “out-of-plane” 550-cm-l band, would seem to be consistent with a C6ring more closely perpendicular rather than horizontal with respect to the surface at 200 K. Possibly this is the trans-di-a species for which transformation from the cis form involves an energy of activation. Perhaps reduced steric interaction with the surface leads to a more energetically stable trans form. At this point, it is noteworthy that the spectrum of cyclohexane adsorbed at low temperature on Pt(ll1) also changes in the 200-230 K region to give a closely similar spectrum [the match is rather better with the cyclohexane spectrum from Land et aL at 200 K (222)than with that obtained by Bussell et al. at 230 K (221)].We conclude therefore that it is very probable that the first dissociation step of cyclohexane on Pt(ll1) gives trans-di-acyclohexene, the temperature of conversion and the possibility of adjacent

VIBRATIONAL SPECTRA OF HYDROCARBONS

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equatorial C-M bonds being more conducive to the formation of the trans rather than the cis form. This suggestion, that the 200 K spectrum is representative of trans-di-a adsorbed cyclohexene, is further strongly reinforced by the observation that the 200 K species formed both from cyclohexane and from cyclohexene undergo the same spectral transformations to give a further intermediate at temperatures greater than 260 K, leading finally to adsorbed benzene above at temperatures exceeding 300-330 K (222,239). It is natural to speculate whether the intermediate formed at ca. 260 K is indicative of adsorbed 1,3-cyclohexadiene. The VEEL spectrum of the latter on Pt(ll1) was recorded at 95 K by Hugenschmidt et al. (240).From the absence of absorptions between ca. 1500 and 1650 cm-', it was deduced that once again the molecule could be multiply a-bonded to the surface. The spectrum showed prominent features at 839, 1323, and 2887 cm-', which might correspond to certain features in the spectrum of the intermediate formed at ca. 260 K (855, 1355, and 2930 cm-'). However, the overall spectrum of the intermediate, which is observed consistently over a temperature range from 260 to ca. 330 K, is more complex. From the dipolar nature of most of its bands, Land et al. (222) deduced that this surface species was notably unsymmetrical, and Henn et al. (239) cited BPTDS evidence that it corresponded to a surface formula of C6H9.The lack of symmetry led Land et al. to suggest a cyclic surface species of formula (CH&(CH), involving a 1,2,3-allylicgroup. To judge from the spectra of adsorbed cyclohexene and of 1,3-cyclohexadiene, the multiple bonding to the surface is more likely to be of a u rather than a s nature on Pt. Laserinduced thermal desorption (LITD) has also provided evidence for the decomposition of cyclohexane on Pt(lll), first to cyclohexene (271)formed by an adsorbed C6H9intermediate, leading to benzene at temperatures exceeding 270 K (272). On heating, the spectrum of the species formed from cyclohexene on Pd(11O) gradually transformed into that of benzene at 250 K, without the appearance of a spectrum indicating an intermediate species (220). 2. Finely Divided Surfaces At an early stage, Palazov and Shopov et al. (241, 242) established that cyclohexene, like cyclohexane, was retained as s-adsorbed benzene after evacuation of a Ni/Si02 sample at room temperature. Patterson and Weaver (83)obtained SER spectra of a number of cycloalkenes, (namely, cyclopentene, cyclohexene, 1,4-~yclohexadiene,and 1,3cyclohexadiene) on the surface of a roughened gold electrode at potentials between -0.4 and +0.4 V versus the standard calomel electrode. In general, the vC= C modes were reduced in wavenumber, by 65 cm-' (1,3-cyclohexa-

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NORMAN SHEPPARD AND CARLOS DE LA CRUZ

diene) and by 120 cm-' (cyclopentene), with somewhat greater shifts at the positive potentials. The intensities of the v= CH modes were substantially reduced relative to those of the alkane groups compared with the situation with the uncomplexed cycloalkenes. A full spectrum illustrated for adsorbed 1,4-cyclohexadiene showed a strong enhancement of the intensity of the yHC =C H mode. These two observations were attributed to the effect of the MSSR as applied to H C = C H groups r-bonded parallel to the metal surface. In the case of 1,3-cyclohexadiene, there was evidence for one complexed double bond (uC= C, 1506 cm-', strong) and an uncomplexed one (a weak triplet centered at 1593 cm-I). B. CYCLOPENTENE

Avery (257,273,274) reported a detailed VEELS study of cyclopentene adsorbed on Pt(ll1) and noted preliminary results for Ir(l1l) (258). The spectrum observed with Pt(ll1) at 90 or 200 K is virtually identical with that obtained by heating adsorbed cyclopentane on the same surface to 200 K. The spectrum of the adsorbed species lacks the absorptions of the free cyclopentene molecule that are characteristic of the H C =CH group, viz., 3085, u=CH; 1630, K = C , and 700 cm-' y H C = C H (275). Furthermore, there was no absorption indicating a CH3 group (1380 cm-'; SCH3s), which would have indicated ring opening. It can therefore be reliably concluded, as expected, that the cyclopentene molecule is di-abonded to the Pt surface, probably in the cis form, as indicated by the presence of a uCH soft-mode absorption at ca. 2700 cm-l. The cis-di-a configuration of the flexible C j ring would be likely to allow methylene CH bonds close to the surface for agostic interactions. A strong band at 880 cm-' is similar to a strong absorption in the infrared spectrum of gas-phase cyclopentane (259). Spectra of organometallic r-complexes of cyclopentene on Pt and Pd (275) have been determined and, in addition to a strong absorption in the 900- to 880-cm-l region, they have another strong feature in the 860- to 830-cm-' region. The latter is assigned to the y H C = C H mode and is not expected to be present for the di-a structure. On warming the species on Pt( 111) to 370 K under UHV, Avery obtained a much simplified spectrum, with absorptions at 3060 (w), 3010 (w), 1245 (w), and 840 cm-' (vs) which could clearly be assigned to the presence of the C5H5,q5-cyclopentadienyl, surface species such as is well known in the spectra of sandwich compounds of the type (C5H5)2Fe(163). The bands from uC-C and in-plane 6 C H modes which appear in the spectra of the sandwich compounds are missing from the surface spectrum. This is as expected by the MSSR for q5-C5H5adsorbed parallel to the surface. The same spectrum was obtained at 260 K characterizing the species formed in the dehydrogenation of cyclopentane.

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A different decomposition pathway of cyclopentene on Ir(l11) was reported by Avery (258), but so far without details of the VEEL spectra. The initial spectrum of the nondissociatively adsorbed species was reported to differ from that of the species on Pt(lll), and when the temperature was 160 K, a species considered to be a dehydrogenated cyclic C5H6was formed; this was envisaged as a d i - d r species involving replacement of the CH bonds of the HC=CH group by CM. Analogous bonding to the surface was found for adsorbed cis- or trans-but-2-ene on Pt(ll1) at 300 K (102). At 400 K the C5H6 species was thought to be converted into cyclic C5H3rather than into the v5-C5H5found on Pt(111). However, the latter was identified as an unstable intermediate at temperature formed near 350 K during the transformation of initially adsorbed cyclopentadiene into the C5H3species. Ward (86) recorded an infrared spectrum of cyclopentene on Ni/Si02 in the vCH region. The absorptions at 2960 and 2875 cm-' are consistent with the persistence of CH, groups in a five-membered ring. Addition of hydrogen led to an intensified spectrum and to cyclopentane in the gas phase.

C. OTHERCYCLOALKENES The VEEL spectrum of 1,3-cyclohexadieneon Pt(l11) at 95 K has already been discussed in relation to the thermal evolution of cyclohexene adsorbed on that surface (Section VI.A.l and ref 240). Palazov and Shopov et al. (241, 242) reported that cyclohexadiene adsorbed on Ni/Si02 at room temperature decomposes into benzene on evacuation. SER spectra of 1,3and 1,4-cyclohexadiene have been obtained indicating n--complexes on a gold electrode (83),and the general results have been discussed earlier in Section VI.A.2. VEEL spectroscopic evidence has been cited (without details) for the disproportionation of cyclopentadiene adsorbed on Pt(11l) at 95 K to give a mixture of v5-C5H5and cyclopentene, followed by the further conversion of the cyclopentene fraction into v5-C5H5at 315 K. On Ir(lll), following adsorption at low temperatures, cyclopentadiene has been reported to transform at 350 K into CSH5 as an unstable intermediate before the final formation of what is considered to be a cyclic C5H3surface species at 450 K. VEELS studies have been made of the 1,3- and 1,5-cyclooctadienes and cyclooctatetraene adsorbed on Pt(11l) (260,276).The spectra recorded at 170 K are consistent with q4 interactions of the hydrocarbon with the surface (with spectra distinguishable from those of the 1,3- and 1,5-dienes themselves) involving losses of C = C groups. For cyclooctatetraene the

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surface structure is presumed to involve alternate C =C bonds with the molecular skeleton still in the tub form. There may be some question whether the 680-cm-l absorption observed at this temperature represents multilayer cyclooctatetraene (260),in comparison with the 221 K spectrum (276).At 373 K the spectrum changed to a simple form with prominent absorptions at 480,888, and 2990 cm-'. This spectrum was assigned to an $-adsorbed flat cyclooctatetraene molecule. The same spectrum was obtained from the 1,5and 1,3-0ctadienes at 371 and 453 K, respectively. Thermolysis of this species gave benzene in the gas phase, while a(CCH) was retained on the surface. A spectrum has also been published for cyclooctatetraene on Pd(l11) (277). Resonance Raman spectra have been reported for CGOon Ag and Ir surfaces (278).Well-defined spectra were observed for the species on Ag, which are similar to those observed for solid C60;the spectra of the species on Ir are broad and ill-defined.

VII.

General Comments about the Spectra of the Cycloalkanes and Cycloalkenes

The most remarkable feature of the spectra of cycloalkanes adsorbed on single-crystal metal surfaces is the observation of the strong bands of vCH soft modes of cyclohexane and of cyclopentane when the geometrical considerations are particularly favorable for the agostic interactions of C-H bonds (axial C H groups in the case of cyclohexane) with the metal surfaces. It is yet to be determined-by intensity-voltage studies of LEED or by PED?-whether the agostic interactions are with metal atoms or with hollows between the metal atoms of the close-packed surfaces. The strengths of the agostic interactions can be semiquantitatively compared for the species on the latter surfaces in terms of the wavenumber lowering relative to the values for C H bonds that are free of such interactions. The sequence is Cu, 115 cm-'; Ni, 155 cm-l; Pd, 250 cm-'; Pt, 285 cm-'; and Ru, 305 cm-'. In qualitative agreement with this sequence, cyclohexane adsorbed on Cu(ll1) and Ni(ll1) surfaces desorbs without decomposition on warming of the sample, but it dissociates to other species on the other metal surfaces. The decomposition pathway occurring on the surface has been followed in some detail for cyclohexane on P t ( l l l ) , and it is as follows: cyclohexane + cyclohexene -+ unidentified intermediate -+ benzene. For cyclopentane on Pt(ll1) the corresponding sequence is: cyclopentane + cyclopentene + cyclopentadienyl, qs-C5HS. Each dehydrogenation sequence ends with an aromatic species adsorbed n--bonded parallel to the metal surface (258).On Pt(ll1) cyclohexene and cyclopentene are adsorbed in the di-c form, but on P d ( l l l ) , A g ( l l l ) , and probably Cu(lOO), they are

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adsorbed as a-complexes, in agreement with the pattern of behavior of ethene bonded to the same surfaces. With respect to the atomic arrangements of the surfaces, the adsorption of cyclohexane occurs very similarly on (111) and (110) planes, in the former case as a nondissociative complex of symmetry C3"; as is often the case, the results on the (100) face [of Pt(100)J are qualitatively different. On Ni the (111) face is less reactive for cyclohexane dehydrogenation than the stepped and kinked [5(111) X (?lo)] plane. Cyclooctatetraene and the 1,5- and 1,3-cyclooctadienes on Pt(ll1) also gave a common product at higher temperatures, which, from the simplicity of the spectrum, is postulated to be flat-lying C8Hs. This possibly involves a-electron donation to the surface in order to overcome antiaromaticity. The most pertinent question remaining for the low-temperature adsorption of the cycloalkanes is whether the C-H..-M agostic interactions are with surface metal atoms or with hollows.

VIII. Aromatic Hydrocarbons

A. BENZENE 1. Spectra on Metal Single Crystals Low-temperature (up to ca. 300 K) VEEL spectra have been obtained for benzene adsorbed on single-crystal surfaces on the following faces: Ni(ll1) (9, 279-284); Ni(100) (280);Ni(ll0) (281, 285); Pd(ll1) (24, 25, 286-290); Pd(100) (24, 25, 286, 287); Pd(ll0) (291-293); Pt(ll1) (284, 294-297, 297a, 297b); Pt(ll0) (298, 299); Pt(100)(1 X 1) and hexagonal (224); Rh(ll1) (295, 300-302); Ru(0001) (303-306); O~(0001) (307); Re(0001) (308);and Ag(ll1) (309,310).Most of these experimental results have earlier been summarized schematically (4, Fig. 10) together with a discussion of the symmetry properties of the species adsorbed on different sites and of the consequent dipolar and impact selection rules. More recently, two papers have described higher resolution RAIRS results for benzene on Pt(ll1) (311, 312) and on Cu(ll0) (312). The latter spectra, obtained by Haq and King, were particularly informative in relation to VEELS results characterizing the same system obtained earlier by Lehwald et al. (284).Furthermore, Raman spectra have been obtained for C6D6 on Ag(ll1) and Ag(ll0) (313). There is general agreement, based on measurements of vibrational spectra taken at room temperature or below, that benzene adsorbs nondissocia6 ring oriented parallel or tively on single-crystal metal surfaces with the c near-parallel to the surface. Furthermore, there is a strong general resem-

246

NORMAN SHEPPARD AND CARLOS DE LA CRUZ

blance between the VEEL spectra, whatever the metal or the crystal face used [fcc ( l l l ) , (loo), (110), etc.], except for variations in position of the strongest bands in the 940- to 690-cm-' region for C6H6(700 to 500 cm-' for C6D6).This family resemblance is shown clearly in Figs. 10A (C6H6) and 10B (ChD6) of our earlier review ( 4 ) and implies that most of the spectral features derive from the adsorbate itself, with few modifications resulting from different site symmetries for the adsorption complexes. A benzene molecule adsorbed parallel to the surface will have its symmetry reduced from DGhfor the isolated molecule to at least C6, as a result of the one-sided perturbation by a flat metal surface. In fact, with an fcc (111) or hcp (0001) surface, C3" is the highest possible symmetry because adjacent threefold sites are of different types. The diffraction methods, tensor-LEED or PED, and also (for isolated sites) the more direct STM, have frequently indicated adsorption on C3, sites, usually of the ad type with symmetry planes dissecting C-C bonds. This is the case for Ni(ll1) (PED, high coverage, 314); Pt(l11) (STM, low coverage, 315); Ru(0001) (LEED, 316); Rh(lll), C6H6/C0 and C6H6/ 2CO (LEED, 295, 300, 301, 317-319); and Pd(lll), C6H6/2C0(LEED, 319). The C3" adsorption sites are usually associated with Kekult-type distortions of the Ch ring with alternating short and long C-C distances, i.e., cyclohexatriene in type, ranging from small [Pd(lll)], to moderate [Rh(lll), C6H6/2C0,1.37 and 1.50 5 0.15 A; (3 X 3) LEED pattern], to seemingly very large [Rh(lll), C6H6/C0, 1.30 and 1.80 ? 0.15 A; c ( 2 d 3 X 4) LEED pattern] differences. VEEL spectra have been measured for the last two cases on Rh(ll1) (295, 300, 301), and it has to be said that there were relatively slight differences between them or in comparison with spectra obtained on Rh(l1l) without CO [c(2d3 X 3) LEED pattern]. The most notable change was a lowering of the strong v4 absorption by 25 for the c(22/3 X 4) structure, which, on the face of it, implies weaker adsorption on the C3"site. Similarly, the coadsorption of benzene and CO on Ru(0001) to give an ordered structure leads to minor VEEL spectral changes, although in this case v4 is raised by 20 cm-' (306). In other cases, twofold bridge sites have been inferred with a symmetry of C, or less, i.e., Pt(ll1) (diffuse LEED, high coverage, 320); Pt(ll1) (STM, isolated sites, 315);Pt(lll), C6H6/2C0(LEED, 318);Pd(ll1) (LEED, 318);and Ni(ll1) (PED, low coverage, 314). In the following discussion of vibrational assignments, we shall for consistency use the Herzberg numbering (321), although a number of authors prefer to use the Wilson alternatives (322). For reference, the latter are given in parentheses: vl ( v2); v2 (q);v3 ( vj); v4 ( v l l ) ; vs ( v13); v6 ( ~ 1 2 ) ; v7 ( v 5 ) ; v8 (v4);

v9 (v14);

v10 (v15); v11 (v10); v12 ( v 2 0 ) ; v13 (v19); v14 (v18);

( v 7 ) ; v16 (v8); v17 (v9); v18

(vh); v19

(v17); v20 (v16).

vlS

VIBRATIONAL SPECTRA OF HYDROCARBONS

247

Continuing with spectroscopic considerations, under symmetry c6, the MSSR allows only the fundamentals v l, v2, and v4 to be observed onspecular in VEELS by the dipolar mechanism. These are respectively the vCH, vCC (breathing), and yCH out-of-plane modes that are symmetrical with respect to the sixfold axis; the corresponding band positions for unperturbed C6H6 (c6D6)benzene are 3062 (2293), 992 (943), and 673 (497) cm-l, respectively (323); v4 gives a very strong band in the gas-phase infrared spectrum and has its vibrational dipole moment perpendicular to the C6 plane. It is hence fully MSSR-allowed for parallel orientation with respect to the surface; v1 and v2 are forbidden in the gas-phase infrared spectrum and are hence expected to give weak features in the spectra of the adsorbed species; v4 is clearly responsible for the strongest bands in the spectra of adsorbed C6H6(C6D6)in the 940-690 (700-500) cm-’ regions. The metaldependent variability of these bands is discussed shortly. Consistently observed weak features near 3030 (2250) cm-l could readily be assigned to vl ,but other vCH ( vCD) modes could contribute to these VEEL features from impact excitation. With the possible exception of the weak feature at 1000 cm-l in the spectrum of ChH6on Ag(lll), there are no features in the on-specular spectra between 1000 and 900 cm-’ which merit obvious assignment to v2. Similar values for v2 have been reported for species on Cu(lll), Cu(lOO), and Cu(ll0) (324). Much of the “family resemblance” of the C6H6 (ChD6)VEEI, spectra arises from consistently observed weaker features which can very reasonably be correlated in wavenumber terms with corresponding fundamentals identified from gas-phase infrared and Raman spectra of C6H6and C6D6. The generally accepted assignments and wavenumbers for these (the mean value quoted can vary by 230 cm-l because of the limited resolution of VEELS) are as follows for C6H6 (C6D6);the corresponding band positions of the gas-phase fundamentals are given in square brackets: 1430 (1350), ~ 1 [1486 3 (1335)l; 1330 (1220), ~9 [1310 (1286)l; 1130 (830), ~ 1 [1150 0 (824)], 7 (867)l; ca. 550 (-), vz0I410 (352)l. Additional metal-dependent or ~ 1 [1178 features in the ranges 400-265 (370-270) cm-1 are doubtless associated with vCM modes. It should be noted that C-C bond-length variations around the c6 ring, as well as different patterns of normal mode mixings on sites of different symmetries, can cause variations in the band positions between the adsorbed species and the undistorted gas-phase molecules, and even between the relative C6H6and C6D6values for the adsorbed species. A case in point is the 3 adsorption on a C3,site would lead to mixing with skeletal mode ~ 1 whereby the higher wavenumber skeletal mode v16. The ensuing resonance would be 3 for C6H6more than for C6D6. expected to lower the ~ 1 value Modes ~ 1 3 ,v9, vl0 or v17, and vz0 have symmetry species El,, B2,, B2,,

248

NORMAN SHEPPARD AND CARLOS DE LA CRUZ

or E,, and E2,, respectively, for unperturbed benzene molecules. All of these features have high impact character, as shown by off-specular measurements, that is much greater than for the strong bands in the 930-690 modes should, to a first approximation, (700-500) cm-' region. El, and BzZL also be impact-forbidden on-specular for C,, symmetry but would be impact-allowed for C,, ( 4 , Table I). B2, modes are dipolar-allowed for C3, (cd). The occurrence of these features on-specular led Lehwald et at. (284) to postulate C,, (cd) sites for the adsorption of benzene on Pt(1ll) and Ni(11l). On the other hand, from recent VEELS studies made with variable incident electron energy (297, 297a), it was concluded that on Pt(ll1) the weaker spectral features from adsorbed benzene at 3000, 1410, and 1130 cm-' are brought up in intensity by resonance of the electron-beam energy with a b,,(n*) electronic transition at 2.7 eV. In the absence of this interaction, only the u4 mode has strong intensity, as would be consistent with the presence of C,, or C,, adsorption complexes. We now turn to the regions of high intensity associated with u4 modes, for which assignments are controversial. For simplicity, we first consider the results on the close-packed fcc (111) and hcp (0001) faces of the different metals. In the order of increasing wavenumbers, we have the approximate metal-dependent wavenumber sequence as follows, although the order is probably not definitive for band positions separated by less than 25 cm-I; where two band positions are recorded at higher coverages, that at higher wavenumber is consistently a strong shoulder on the lower wavenumber band: C6H6(C6D6): gas phase, 673 (497); Ag, 675 (-); Pd, 730/820 (515/ 620); Re, 740/845 (-565/-); Ni, 750/850 (540/645); Ru, 7.55/860 (550/-); Os, 760/850 (-/--); Rh, 800 (560); and Pt, 840/920 (600/715) cm-'. It is seen that there is a wavenumber sequence from a very slight perturbation relative to the gas-phase molecule, as expected for Ag, to values of 170 (105) cm-' higher for the other extreme case of Pt. Less symmetrical crystal planes have been investigated for Ni, Pd, and Pt, with the following results: Ni(100), 750/845 (540/645); Pd(100), 720/870 (520/675); Ni(llO), 700/845 (510/-); Pd(llO), 750/745/890 (505/560/680); Pt(100)(1 x l ) , 810/-, and hexagonal, 810/-900; Pt(llO)(l X 2), 910/655 cm-'. In the case of Pd(llO), there appears to be sound evidence for attributing the 705- and 745-cm-I bands to two different surface species, with the latter correlating well with the growth of a c(4 X 2) LEED pattern (291). The Pd(l1O) spectra also showed additional weak features for C6Hh (C6D6)at 1580 (1550) cm-' related to v16of benzene [1596 (1%2)], thereby providing additional confirmation that the spectra are indicative of nondissociatively adsorbed benzene molecules. The principal uncertainty in the literature concerns the assignment of the strong shoulder some 100-120 (105-115) cm-' higher in wavenumber

VIBRATIONAL SPECTRA OF HYDROCARBONS

249

relative to the main v4 band. These have been variously attributed to the following: (i) v4 of a second adsorbed C6H6(C6D6)species; (ii) a second out-of-plane yCH/yCD mode v l l ( E l g )of a single species adsorbed on a site of symmetry C, or less which is coupled with, and shares intensity with, v4; (iii) a partial contribution from the otherwise “missing” v2 CC breathing mode of a single species. We summarize in the following the principal VEEL evidence for or against these alternatives. a. Assignment ( i ) . The evidence relating to the presence or absence of a second adsorbed C6H6(C6D6)species remains controversial in a number of cases. Lehwald et al. (284) claimed that a second species was present at higher coverages on Ni(ll1) and Pt(ll1) because of coverage-dependent and/or temperature variability of the ratio of intensities of the low-wavenumber and high-wavenumber absorptions in the 930-690 (700-500) cm-l region. We note (vide infra) that this has recently been confirmed by a RAIRS study with Pt(ll1). Surprisingly, however, Jobic et aZ. (282) did not find coverage variability in the Ni(ll1) case. This led them to suggest a combination of assignments (ii) and (iii) on the assumption of the presence of a single adsorbed species. Waddill and Kesmodel (286) also drew the conclusion that a single adsorbed species is present on Pd(ll1). This led them to propose a general acceptance of assignment (ii). They supported the single-speciesassumption on the basis of TPD studies of sequential C6H6 and C6D6adsorption. As these were desorbed by TPD with equal probability, the authors concluded that the high-coverage and low-coverage species were the same, although they did acknowledge that rapid CSH6/C6D6exchange between different types of surface sites could lead to the same result. Subsequently, independent NEXAFS plus UPES/XPES work (325) has, however, led to the conclusion that on Pd(ll1) a second species is formed at high coverages that is tilted ca. 30” from surface parallelism. Furthermore, firm spectroscopic evidence has been found for the presence of two species on Pd(ll0) (291) as noted earlier. b. Assignment (ii). The two absorptions in question shift together to lower wavenumbers in the spectra of adsorbed C6H6(C6Dh)and are hence most probably indicative of similar yCH ( y CD) out-of-plane deformation modes. They are both highly dipolar active. If it is correct that in some cases they originate in a single adsorbed species, as suggested for Pd(ll1)

250

NORMAN SHEPPARD AND CARLOS DE LA CRUZ

by Waddill and Kesmodel (286), it is agreed that the high-wavenumber shoulder is most likely to be indicative of mode v l l , of symmetry El,, for unperturbed benzene. This can become dipolar active for a surface complex of symmetry C, or less and would then pick up intensity by resonance with v4. This assignment was supported by Aarts and Sassen (290), who also studied the spectrum of C6H5D. Fujisawa et al. (291) also supported the vll assignments of the high-wavenumber shoulders at 890 (690) cm-’ in their VEEL spectra of species on Pd(llO), possibly overlapping with v2. c. Assignment (iii). Jobic et al. (282),having to their satisfaction established that there is a single adsorbed species on N i ( l l l ) , suggested that part of the intensity of the high-wavenumber shoulder could be from the “missing” but formally dipolar-allowed, skeletal breathing mode, v2. They then measured the spectrum arising from the adsorption of the isotopically substituted 13C6H6molecule and interpreted a change in contour of the high-wavenumber shoulder in terms of the presence of a lower wavenumber band of the heavier isotopomer. Unfortunately, however, the limited resolution of the VEEL spectra meant that this conclusion depended on a resolution of overlapping absorptions from the postulated v2 and v l l .Jakob and Menzel (303) subsequently supported the v2 assignment on the basis of an analysis of spectra characterizing adsorption on Ru(0001), where once again the spectral evidence seemed to be consistent with the presence of a single adsorbed species. Huntley et al. (285) likewise interpreted the spectrum on Ni(ll0) in this manner. In the C6Dbspectra on Ni(ll1) and on other surfaces, Jobic et a/. (282) also pointed to a nonoverlapped band at ca. 830 cm-’ for assignment to v 2 ,which, at the correct position for alternative assignment to vl0 or ~ 1 (see 7 the earlier discussion), nevertheless seemed surprisingly strong in relation to the vlO/vl7counterparts in the C6H6 spectra. However, this intensity argument is not decisive because vl0 and vI7 are both 6CH (SCD) modes. In the C6D6 case, the amplitudes of the D-atom motions will be reduced relative to those of the H atoms, but the motion of the carbon atoms will have to be enhanced so that there is no net linear or angular momentum associated with the normal mode. If the carbon atoms are the strong impactscattering centers, this could possibly lead to enhanced intensity for the mode for C6D6compared with that for C6Hh. Other questions can be raised with respect to the v2 assignment. This assignment of the many C6D6 spectra implies that the mode is of very limited variability, whereas v4, as we have seen, varies gradually but substantially from metal to metal. If vz is to move downward from the value of 992 (943) cm-’ for benzene itself, it might have been expected that this too would be a gradual process. Fujisawa et a/. (291) also raised the question

VIBRATIONAL SPECTRA OF HYDROCARBONS

25 1

whether v2 values of 850 (830) cm-' are not unrealistically far removed from the gas-phase benzene value. The v2 breathing mode of organometallic models gives particularly strong features in the Raman spectra. Anson and Powell (326) obtained highquality Raman spectra of n-complexes (p1-C6H6)Cr(C0)3and also of (p3C6H6)O~3(C0)9. In the latter case, crystallography shows that the benzene site with a is complexed to three osmium atoms in an effectively C3, (ud) Kekul6 distortion of the C-C bonds of lengths 1.51 and 1.41 A with errors of ?0.04 A. The p1chromium complex gave v2 at 980 cm-', little removed from the value for the free benzene molecule, whereas the p3Os, complex gave v2 at 916 cm-', a lowering of 76 cm-' from the benzene value. This goes halfway toward the low-wavenumber shift of ca. 140 cm-' that is implied by the assignment of the 850 (830) cm-' features to v2.Unenhanced Raman spectra of C6H6 and C6D6on Ag(ll1) and Ag(ll0) show v2 unambiguously at 990 (945) cm-' (313). Hence the question is whether, and if so why, there is a large and consistent lowering of v2 on passing from a group IB (IUPAC group 11) metal such as Ag and Cu to the group VIII (IUPAC groups 8-10) metals. We note also that Raman spectra of benzene adsorbed on finely divided Ni/Si02 and Pt/Si02 (see later discussion) show very strong bands, clearly indicative of v2, at ca. 990 cm-'. There is, of course, the possibility that the predominant form of adsorbed benzene on the latter finely divided catalysts may be different from those on single crystals, e.g., from benzene adsorbed on top of single metal atoms, or that more than one type of adsorbed species could be present. At present, a definitive assignment of the 850 (830) cm-* features to v2 in the spectra on many group VIII metal single-crystal surfaces remains unresolved. d. RAZR Spectra. RAIRS would seem to have potential to resolve the important ambiguities that remain in the interpretation of the many VEEL spectra in the 940- to 690-cm-' region. The bands in question should be strong enough for measurement by RAIRS, and the high resolution of the technique should enable any overlapping bands from v4, vI1, or v2 modes to be separately identified. Furthermore, the evaluation of any 12C+ I3C isotopic shifts as well as the strict MSSR infrared selection rules should eliminate ambiguities caused by the presence of impact features in the VEEL spectra. To date, just two such systems, with adsorbates on Cu(ll0) at 85 K and on Pt(ll1) at 90, 220, and 300 K, have been investigated by RAIRS (312). A single absorption characteristic of v4 occurs for the species on Cu(ll0) at 685 cm-' with a slight shift to 683 cm-' as the coverage increases within a monolayer, compared with the gas-phase value of 673 cm-'. It is clear that the benzene is adsorbed parallel to the surface with only slight perturbations through v-bonding. At higher coverage weak additional

252

NORMAN SHEPPARD AND CARLOS DE LA CRUZ

absorptions at 3084 ( vlz), 3066 ( vl), 3042 ( q 5 )1479 , ( v13), and 1037 cm-' ( vI4) occur, as expected for a nonoriented and physically adsorbed second layer. The RAIRS results characterizing the adsorbate on Pt(ll1) are more complex and particularly interesting. They agree very well with the earlier VEELS data of Lehwald et al. (284) except for the inability of the VEEL spectra to resolve close-lying bands near 820 and 830 cm-'. At low coverage, at 90 or 300 K, bands were observed at ca. 900 and 830 cm-', and both grew with coverage. After a dose of 1 L, the former band reached its maximum, whereas the 830-cm-' band continued to grow, together with a shoulder at ca. 820 cm-'. A t high coverage, still within a monolayer, at 220 K (multilayer formation is readily identified by its v4 feature at ca. 680 cm-'), the 900-cm-' band lost intensity as the 820-cm-' band continued to grow until it slightly exceeded in intensity its ca. 830-cm-' counterpart. Haq and King (312) convincingly interpreted these spectra as evidence of the v4 modes of three different adsorbed species without any signs of additional features in the spectra indicating v2 or vll modes. STM studies of Pt(ll1) have also given isolated molecular features of different shapes, emphasizing that adsorbed species on different types of sites can coexist, with some of them having threefold symmetry. Haq and King (312)suggested that the 900-cm-' band may be associated with adsorption on twofold bridged sites and that the closely similar 830- and 820-cm-' bands are associated with adsorption on the fcc and hcp threefold hollow sites. Lowcoverage twofold bridge sites and high-coverage threefold sites (hcp) have been deduced from the PED study of benzene on Ni(ll1) (314).An alternative possibility for the 830- and 820-cm-l adsorptions is that they relate to flat-lying and slightly tilted molecules (the latter arising from high-coverage steric crowding), as has been suggested for 745- and 705-cm-l absorptions characterizing the adsorbate on Pd(ll0) (291).One looks forward to further RAIRS studies, particularly with Ni(ll1) and P d ( l l l ) , where there is still uncertainty about whether one or more surface species are present. e. Unusual Case. An exception to the "family resemblance" of VEEL spectra from benzene at low temperatures was observed for the species on Re(0001) at 120 K (308),whereby two vCHIVCD modes were observed, at 3050 (2295) and 2910 (2175) cm-l. As dissociation seems to be unlikely at such low temperatures, Tardy et al. (308) suggested that in this case adsorption occurs on an exceptionally unsymmetrical site. A t ambient temperatures a simpler spectrum indicated decomposition, probably to give an a(CCH) species such as is obtained by the decomposition of ethyne at 470 K on the same surface. A t 800 K the spectrum indicated the formation of a hydrogen-free graphitic overlayer.

VIBRATIONAL SPECTRA OF HYDROCARBONS

253

In most cases the spectra from nondissociatively adsorbed species are observed up to ambient temperatures, but the spectrum of the species on Os(OOO1) (307) changed in detail between 273 and 325 K and again at 382 K. In conjunction with TPD and ARUPES experiments, these changes were attributed to losses of hydrogen atoms to form surface phenyl groups (325 K) and then 1,2-disubstituted C6H4species (382 K). VEEL spectra of benzene adsorbed on (111) PtloNigO(281) and Pt78Ni22 (296) alloys were found to be similar to those of benzene on pure Ni(ll1) and P t ( l l l ) , respectively, except that in the Ni-rich case the main v4 absorption was shifted by 30-40 cm-’ to higher values. Preadsorption of cesium on Pd(ll0) (292,293) led to surface reconstruction of the missing-row type, i.e., Pd(llO)(l X 2), with the Cs atoms in the troughs. A t 90 K two strong bands of C6H6 (C6D6),at 685 (505) and 770 (595) cm-’, varied in intensity with coverage. The former, which occurs at low coverage, was attributed to adsorption on the Cs-free part of the surface, and the latter to adsorption near Cs atoms. However, even the low-wavenumber band was observed to be 20 cm-’ lower than that characterizing the adsorbate on the clean Pd(ll0) surface, showing that electron donation to metal from the adsorbed Cs atoms leads to weaker .rr-adsorption on the surface, as would be expected in view of the high electron density associated with the .rr-orbitals of benzene. Hallmark and Campion (313) did very well to observe unenhanced Raman spectra of C6D6 on Ag(ll1) and Ag(ll0) surfaces. They observed bands of the species on the Ag(ll1) surface indicating the vz (ring breathing) and v4 (yCD) modes, as is consistent with molecules lying parallel to the surface on C6,,or C3,,sites. Additional weak features indicating vll and vI9 (yCD) modes were observed for the species on Ag(ll0). It was concluded that the appearance of the four modes was consistent with a C2, site symmetry. All the observed band positions were within 15 cm-’ of the values for the corresponding modes for gaseous C6D6,indicating only minor perturbations from .rr-bondingto the surface. Jacob and Menzel investigated the coadsorption of benzene with hydrogeddeuterium (305) and with C O (306) on Ru(0001). In each case the benzene spectrum was little affected. In the C6D6/H2coadsorption, HID exchange occurred at ca. 350 K, a temperature that is notably lower than was required for benzene dissociation, which gave a TPD H2/HD/D2peak at ca. 450 K. In all cases the v C 0 frequency of the coadsorbate was very notably lowered compared with that when CO alone was adsorbed. Measurements have been made by helium atom scattering of the vibration of benzene against the Cu(ll0) surface near 150 K (139) and of the “frustrated-translation’’ mode of benzene parallel to the surface on Rh(ll1) at 160 K (327).

254

NORMAN SHEPPARD AND CARLOS DE LA CRUZ

2. Finely Divided Metals Early infrared studies of the adsorption of benzene on metal oxidesupported metal catalysts concentrated on the 3000-cm-' vCH region. Ward (86),Palazov and Shopov (243),and Takenaka and Sheppard (328),investigating both Ni/Si02 and Pt/SiO,, observed a weak and broad absorption at ca. 3040 cm-' after benzene adsorption followed by evacuation (Figs. 10,11A, and 11B). A possible weaker companion band sometimes occurred near 2960 cm-'. Erkelens and Eggink-Du Burck (249) also observed a weak aromatic vCH absorption after evacuation of Pt/SiO, and reported subsequent slow self-hydrogenation to give CH2groups. They observed no bands of adsorbed species on Pd/Si02 or Fe/Si02, but, working with Nil S i 0 2 or Cu/Si02, they observed atypical spectra consisting of extremely broad and weak absorptions, centered near 2950 cm-' but stretching from ca. 3060 to 2760 cm-'. There is general agreement that the addition of gas-phase hydrogen to a catalyst with adsorbed benzene leads to the replacement of the aromatic vCH absorption by a CH2-rich spectrum, with vCH2 absorptions at 2923 and 2852 cm-' and 6CH2 at 1460 cm-', with some cyclohexane being desorbed into the gas phase. It was shown early by Sheppard et al. (248) that progressive pumping led to the gradual replacement of the CH,-rich spectrum once again by the weak aromatic vCH absorption (Fig. 10 and inset of Fig. 11B). This could be repeated through several cycles, with the absorption bands gradually weakening. It was proposed that the CH2-rich spectrum is indicative of cyclohexane-ring species attached to the metal surface by a few C-M bonds which were slowly converted into cyclohexane itself on hydrogenation. The very large increase in observed intensity resulting from hydrogenation was at the time interpreted as evidence of the presence of dissociatively adsorbed benzene species, probably including surface C6 carbides, but it was noted that the strengths of aromatic vCH bands of n--species can also be weak and variable. We now interpret this intensity increase in terms of n-adsorbed ChH6species, whereby even the intrinsically weak vCH absorptions are further weakened by the MSSR because the C-H bonds are near-parallel to the surface. Subsequently, the Pt/SiO, and Pt/Al,O, systems have been investigated in more detail, with the latter metal oxide support providing the advantage of transmission down to ca. 1100 cm-' (87,246,329-331). Representative spectra are shown in Figs. 11D and 1 l E . We first discuss the most detailed spectra obtained by Haaland (329). Haaland studied the adsorption of both C6H6(Fig. 11E) and CbDhon a freshly reduced Pt/Al,O, catalyst and observed absorptions at ca. 3040 (2260), 1398 (1268), 1274 (1230?), and 1147 (-) cm-', where the band

255

VIBRATIONAL SPECTRA OF HYDROCARBONS Benzene on silica supported platinum 0.1c

Initial adsorption (30 Mm pumping)

0

d 0

m

C

-

1

0.2c

I

a

300 torr H, (16hours)

0

t 0

m

-P

5:

2

0

3 2 ;

,

:,:"_"

1500 ,

1400

10 min

0

0 0.20

0

3200

3000

2800

1500

1400

FIG.10. Infrared spectra of benzene adsorbed on Pt/Si02 at room temperature after 30 min of evacuation, after being in contact with 300 Torr H2 for 16 h (the dashed spectrum is from gas-phase cyclohexane), after pumping for 10, 30, and then 50 min, and after reentry of 300 Torr H2 for 16 h. [From Ref. 248.1

positions for C6D6 are given in parentheses. The strong v4 region falls, of course, under the A1203blackout. The preceding band positions comparewith those of gas-phase benzene [3068 (2287); 1486 (1335); 1310

256

NORMAN SHEPPARD AND CARLOS DE LA CRUZ

,:I Pt m

-

G. rn

r 3100

2900

a

N r

A

I

1400

1600

-

a c

c

A 1200

cm-'

Fro. 11. Infrared spectra of benzene adsorbed on (A) Ni/Si02; (B) PtiSiO, (inset after addition and then evacuation of Hz); (C) PtlSiO,; (D) PtlSiOz: (E) Pt/A1203.[(A) from Ref. 328; (B) from Ref. 248; (C) reprinted from Ref. 330, J. Mol. Sfrucr. 174, T. Szilagyi, p. 395. Copyright 1988 with kind permission of Elsevier Science-NL, Sara Burgerhartstraat 25, 1055 KV Amsterdam, The Netherlands; (D) from Ref. 331; (E) from Ref. 331: reprinted from Ref. 329, Surf. Sci. 102, D. M. Haaland, p. 405. Copyright 1981 with kind permisson of Elsevier Science-NL, Sara Burgerhartstraat 25, 1055 KV Amsterdam, The Netherlands.]

(1286); and 1150 (824) cm-'1 and with the VEEL Pt(ll1) and Pt(100) singlecrystal values, cited earlier, of 3030 (2250), 1430 (1350), 1330 (1220), and 1130 (830) cm-'. It is noted that the second and fourth band positions, which were assigned earlier to ~ 1 (&) 3 and to vl0 (&) or v17 modes, occur at lower values in the infrared spectra than in the averaged VEEL

VIBRATIONAL SPECTRA OF HYDROCARBONS

257

spectra determined in single-crystal work; for Pt(ll1) the latter values for u13 are 1420 (1350) cm-'. These bands are broad and poorly resolved in the VEEL spectra, and the infrared values are more precise, but the degree of difference raises once again the question of whether the same adsorption sites are dominant on the finely divided and single-crystal surfaces. This reinforces the importance of obtaining high-resolution RAIR spectra of the single-crystal surfaces. We see later, however, that there is better agreement between the VEEL and Raman spectral data for these bands. This, in conjunction with the infrared values, could imply split double degeneracies for these modes, with the infrared spectrum recording one component and the Raman and VEEL spectra the other. For the infrared spectra there is, of course, no impact mechanism available for exciting additional features from not completely symmetrical modes, and none of the in-plane ca. 1395 ( v I 3 ) ,1275 (vg), or 1147 cm-' ( vl0 or v I 7 )modes would be allowed if the MSSR applied strictly to parallel adsorption on (111) or (100) facets. They would, however, all become allowed on a C, site, such as would arise from adsorption on twofold bridges. The infrared spectra of alternative monosubstituted or ortho-disubstituted benzenes (the most likely dissociatively adsorbed species) would give rise to two additional strong bands between 1400 and 1620 cm-', and so the observed spectrum is again seen to be consistent with nondissociative adsorption. Haaland was also able to resolve the ca. 3040-cm-' vCH absorption into separate components at 3074, 3048, 3032, and 3011 cm-l. He attributed these to the presence of multiple adsorption sites, although such additional component bands were not observed for the other absorptions. It should be recalled, however, that the MSSR can break down for adsorbates on small-particle Pt catalysts (no mean particle size was given), as in the case of r-bonded ethene (Part I, Section VI.B.c), leading to all the vCH modes becoming active. Haaland also measured the infrared spectra of benzene adsorbed on Pt/Al,03 that had been regenerated after previous benzene/cyclohexane adsorptions (247); the surface was thought to retain structured carbonaceous deposits. In this case, the broad uCH feature was centered at ca. 3030 cm-' (with components at 3042, 3031, 3024, and 3014 cm-') rather than 3040 cm-' for the species on the freshly prepared catalyst, and a weaker companion band occurred at 2947 cm-l. The benzene absorption bands at wavenumbers <1500 cm-' were little changed in position but become more prominent in room-temperature spectra in which the 2947cm-' feature was weakened. Spectra measured over the range 300-650 K showed that the 2947-cm-' feature disappeared at 435 K, whereas the vCH aromatic bands retained considerable intensity at temperatures up to 560 K.

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NORMAN SHEPPARD AND CARLOS DE LA CRUZ

A t that stage the band center had changed back to ca. 3040 cm-', and considerable generation of surface C O implied that a reaction had occurred with oxygen in the support, possibly the OH groups. A vCH value of 2947 cm-I implies the presence of C H groups involving a-bonding to the metal surface. The fact that hydrogenation gave cyclohexane and no methyl groups shows that only intact C6 rings were present. These could have been in the form of separate (CH)6 cyclic species, with all carbons a-bonded to the surface, or alternatively (a possibility that seemingly was not explicitly considered by Haaland) of partial r-bonding and partial a-bonding of the same cyclic C6 species to the Pt surface. It should be noted that the 2947cm-' absorption represents a relatively minor proportion of e-bonded C H groups. This is because C-H bond for C-H bond, the absorption bands of sp3 CH groups are much stronger than those of sp'-hybridized ones (332, 333) and because the latter are further weakened by the MSSR for flatlying rings. Similar temperature-dependent spectra were obtained after evacuation following the adsorption of C6D12on Pt/A1203.However, it was also found that exchange with OH groups of the support led to a proportion of vCH absorptions at both ca. 3025 and 2947 cm-' (and corresponding vOD absorptions) in addition to the strong vCD absorption at 2260 and 2200 cm-'. Such exchange occurred even at room temperature and quantitatively favored the 2947-cm-' species. Szilagyi (330) also investigated the adsorption of benzene on Pt/SiO,. The spectrum of the sample after evacuation (Fig. 1 l C ) was similar to the others. Before evacuation, however, it exhibited some absorption of cyclohexane formed either by self-hydrogenation (possibly leading also to hydrogen-free carbonaceous surface species at this stage) or by the addition of residual hydrogen from the reduced PtlSiO, catalyst. The adsorption of C6D6also led to D -+ H exchange involving the OH groups of the support. This time the sequence of adsorption followed by evacuation and then by readdition of hydrogen did not lead to any formation of cyclohexane but instead to a more complex spectrum indicating retained alkyl species (possibly formed from cyclohexane rings with multiple C-M bonds). A second evacuation was not observed to lead to the re-formation of adsorbed benzene but instead to the complete disappearance of vCH absorptions which were regenerated (from carbonaceous residues) on the readdition of gasphase hydrogen, the sequence being repeatable through many cycles. Readdition of benzene led to a band near 3040 cm-' that was stronger than before, and the surface species could be transformed into cyclohexane upon addition of hydrogen. The latter adsorbed benzene molecules were considered to be adsorbed on nonflat sites so that they were less readily dehydrogenated to give carbonaceous residues.

VIBRATIONAL SPECTRA OF HYDROCARBONS

259

Reficha also measured infrared spectra of C ~ and H ~of C ~ adsorbed D ~ on Ni/A1203(252).Apart from absorption bands characteristic of physically adsorbed C6H6and C6D6,a low-wavenumber vCD band at 2050 cm-' was observed in the latter case and also assigned to a cT-bonded C6-ring species such as Haaland had considered. Primet et al. (250, 251) also showed that aromatic vCH absorptions characteristic of benzene-derived species on Pt/SiO, caused a marked lowering (from 2065 to 2030 cm-') of the strong bands characteristic of coadsorbed CO. The authors interpreted the vC0 shift as arising from donation of electrons from the benzene n--orbitals to the metal surface. Palazov (245) drew similar conclusions concerning benzene and C O coadsorption on Ni/SiOz. As mentioned earlier in connection with single-crystal work, several highquality Raman spectra have been obtained by Krasser et al. (334-338) for benzene adsorbed on Ni/Si02 and Pt/Si02. Examples of spectra are illustrated in Figs. 12A and 12B, and these were shown to involve some surface enhancement. Because of the weakness of Raman scattering from the highly polar metal oxide supports, these spectra cover a much wider wavenumber range than the infrared spectra. As expected, the intensity distribution in bands of modes common to both types of spectra is very different in the Raman spectra from that in the infrared spectra. The Raman spectra have characteristics expected for nondissociative adsorption but with particularly prominent features for the completely symmetrical v1 and v2 modes at 3060 (2290) and 990 (945) cm-l, respectively.

r nm c

m (D .N c

c

260

NORMAN SHEPPARD AND CARLOS DE LA CRUZ

We collect the infrared and Raman assignments for benzene on the finely divided catalysts in Table VII and compare the spectral data with these for liquid benzene (323,339).Common features within both types of spectra show good wavenumber agreement except for the features attributed to the ~ 1 (I&) 3 skeletal mode. In the infrared spectra, the values are 1398 (1268) cm-', and in the Raman spectra they are 1432 (1364) cm-'. Possibly the E l , degeneracy is split on the surface, with one component showing up well in the infrared and the other in the Raman spectrum. Table VII includes Raman spectra recorded with Pt/SiO, at two different exciting wavelengths to show some degree of variability of the spectra in this respect. As Krasser et al. showed (337), these spectra of species on Ni/Si02 and on Pt/SiO, can be very satisfactorily interpreted in terms of the presence of undissociated adsorbed C6H6 or C6D6on a C,, site. A moderate-strength feature at 864 cm-' would not have been active under the MSSR selection rules as applied to Raman spectra (340,341) for Chvsymmetry of the surface complex but would seem to be allowed for C3".The spectral characteristic of the species on Ni/Si02 and Pt/SiO, are very similar except for weaker bands in the 950- to 650-cm-' region where the variable position of v4 is expected on the basis of the V E E L investigations of single-crystal planes. At this qualitative level it might seem that the controversy over the position of the C6breathing mode, v 2 ,has been resolved in favor of 990 cm-' rather than the alternative of ca. 850 cm-' preferred under assignment (iii) (Section III.A.l). O n closer inspection, however, it is seen that the features in the Raman spectra most readily assigned to v4 (684 cm-' for species on Ni/ S i 0 2 and 738 cm-' for the species on Pt/Si02) are at notably different values from those derived from VEEL or RAIR spectra for adsorption on (111) planes. The latter gave doublets of 750/850 and 840/920 cm-', respectively. By comparison with the gas-phase value of 673 cm-' for v4, the Raman values imply in each case the presence of a species more weakly perturbed than that on the (111) planes. However, additional weak Raman features could fairly well represent the doublets (776/864 and 8641948 cm-I, respectively), although several of the latter Raman bands could alternatively be assigned to other benzene fundamentals. Such doublets also could contribute to the unexpected strength of the 864-cm-' feature. It is therefore possible that the spectra of the species on Ni/SiOz and Pt/SiO, are indications of two species, with the 990-cm-' feature being associated with the more abundant, but more weakly perturbed one, possibly adsorbed on single-atom sites on the rough surfaces of the metal particles. In that case the prominent 864-cm-' feature could still be evidence of v, of molecules adsorbed on three-metal-atom sites, and indeed similarly placed features also occur in the spectra of ChDh.The controversy over the location of the q on threefold sites therefore continues and would

TABLE VII Assignments of the Infrared and Raman Spectra of Benzene Adsorbed on Oxide-Supported Platinum and Nickel (in Comparison with Liquid State Data)

Infrared Mode" (323)

Wavenumbers (cm-') (323) 3068 3063 3062 3047 1596 1486 1326 1310 1178 1150 1038 1010 995 992 975 849 703 673 606 410

Rarnan

IR liq (339)

PtlA1203 (329)

3048 (s) 3063 (s) 3080 (s), 3039 ( s ) 3074 (m, sh), 3032 (s)

1479 (s)

-

1398 (m)

-

1310 (vw)

-

1274 (m)

Raman liq (339)

Pt/SiOz 3505 Ab

3063 (vs) (3046) 1605 (m), 1584 (m) -

-

-

975 (vw) -

0

-

-

672 (vs) 406 (w)

1147 (mw)

I

Ni/SiOz 5145 A (337)

3080 (w) 3060 (s) 3045 (s) 1600 (m), 1590 (m) 1432 (w)

3080 (w) 3060 (s) 3045 (s) 1605 (m), 1585 (m) 1440 (mw) 1326 (vw), 1320 (vw) 1262 (w) 1184 ( s ) 1161 (w) 1056 (vw) 995 (vs) 869 (vw) 864 (m)

-

1176 (m) 1176 (w) 1036 (s) 1010 (w, sh) -

Pt/SiOz 4880 A (337)

992 (vs) 850 (w) 606 (ms) -

3060 (s) 3040 (s) 1600 (m) 1432 (rn) 1355 (rn)? -

1176 (m) 1144 (vw)

-

990 (vs) 950 (vw)

664 (m)? 380 (rns)

-

1193 (rns) 1145 (vw) 1028 (vw) or 1028 (vw) 948 (w) 990 (vs) or 948 (w) 864 ( 4 900 (w), 735 (m) 604 (m) 360 (m)

a

Herzberg numbering of the modes. Personal communication from Prof. W. Krasser

776 (m), 684 (m) 603 (m) 407 (rn) 304 (w)

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NORMAN SHEPPARD AND CARLOS DE LA CRUZ

seem to be most directly resolvable by (difficult) experiments involving unenhanced Raman spectra of flat (111) single-crystal planes of metals such as Ni, Pd, or Pt-preferably Pt, for which the perturbation of v4 is strongest. Moskovits and Di Lella measured the SER spectra of benzene and of C6D6 on cold-deposited Ag (342) and later made an interesting study of parallel VEEL and SER spectra of C6H6adsorbed on such surfaces (343). For Ag deposited at 60 or 250 K, the SER spectrum is dominated, as might be expected, by the skeletal breathing mode, vz, at 994 cm-'. At 60 K other . Raman lines were observed, at 860 (vI1) and 380 cm-' ( v ~ ~A) feature indicating the out-of-plane mode at 694 cm-' v4 grew in relative intensity with increasing coverage and at 250 K this band was in evidence at low surface coverages. The VEEL spectrum of the surface at 250 K shows the profile usually found at low coverages for flat single-crystal planes, viz., the dominance of the out-of-plane mode v4 at 694 cm-'; the spectrum differs mainly from that of the species on Ag(ll1) (309,310),as might be expected, in a somewhat greater prominence of the weaker features that are usually attributed to off-specular, impact-excited, modes. However, as a result of increasing coverage, large intensity increases were seen in VEELS bands at 3065 and 1570 cm-' and in the 1200- to 1000-cm-' region. The latter type of spectrum has been reported for all coverages on the Ag film deposited at 60 K. At 250 K annealing might well have led to the generation of flat crystalline faces, and so the low-coverage VEEL spectra of the adsorbate on the 250 K surface were tentatively assigned to these. The very different high-coverage spectra can be attributed to adsorption on rough areas, with "off-specular" features being dominant. The 694-cm-' mode becomes Raman- and presumably SER-allowed for flat facets on the surface, according to the MSSR. Rougher surfaces could lead to additional SERS features. Similiar SER spectra of benzene on cold-deposited silver have also been reported by Otto et al. (99, 136) (Fig. 6D). A few very broad Raman bands at ca. 2960, 2750, ca. 1000 (shoulder), and ca. 800 cm-' have been reported for benzene adsorbed on Rh/A1203 (344). An v6-C6H6 bonding mode has been suggested. Evacuation and subsequent coadsorption with C O lead to a changed spectrum of the type identified earlier as indicative of-(CH =CH),- in polymers. SER spectra determined by two experimental procedures have also been reported for benzene on roughened gold electrodes (345, 346). Electrochemical roughening gave two bands representing the skeletal breathing mode v2, at ca. 965 and 975 cm-'; the latter band was displaced when bromide ion was coadsorbed. A single broad band was observed at 970 cm-' characterizing the adsorbate on the gold electrode surface prepared by electroplating. These spectra are again as expected for flat adsorption of the benzene molecules on the electrode surface. SER spectra have also

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263

been obtained for benzene adsorbed on vapor-deposited sodium, but only at multilayer coverage (347).Nevertheless, orientation effects with respect to the metal surface were observed. Incoherent inelastic neutron scattering (INS) has been used to obtain vibrational spectra of benzene adsorbed on Raney nickel (348), Raney platinum (283),and platinum black (349).Such spectra show no symmetryrelated selection rules, but the band intensities are mainly related to the mean square amplitudes of motion of the hydrogen atoms involved in the vibrations. In all cases it is agreed that benzene at less than monolayer coverage adsorbs intact, presumably in a parallel orientation. The values assigned to the out-of-plane v4 mode characterizing the species on Pt and determined by INS are more consistent with the earlier VEELS results from Pt black rather than from Raney Pt. The vibrational assignment adopted for the adsorbate on Raney Ni corresponds to assignment (iii) discussed earlier and remains uncertain. The INS experiments also provide results in the low-wavenumber region and have led to assignments of the vCM modes and those involving bending and torsional motions of the benzene molecule relative to the surface. Some disagreement remains between the experimental results reported for Raney Pt and Pt black.

B. SUBSTITUTED BENZENES 1. Spectra of Adsorbates on Metal Single Crystals Toluene adsorbed on the following metal single crystals has been investigated: Ni(lll), at 110 K (350); Pd(lll), at 180 K (289);Pt(lll), at 300 (289, 294) and 350 K (351);and Ru(0001) over a wide temperature range and also with coadsorbed CO (352, 353). With the exception of the work with Ni(lll), carried out by RAIRS, the spectra were obtained by VEELS. The results of low-temperature investigations of Ni(lll), Pd(lll), and Ru(0001) are consistent with nondissociative adsorption with the phenyl ring close to parallel to the surface. This is once again shown by the dominant strength of the in-phase yCH mode (analogous to u4 of benzene). Higher coverage RAIR spectra of the adsorbate on Ni(ll1) show the increasing growth in intensity of additional features which imply a tilting of the phenyl group to a limited degree. The wavenumbers of the yCH mode for toluene nondissociatively adsorbed on various metal surfaces follow a sequence similar to that observed for benzene. The latter are given in parentheses in the following list: Ni(lll), 732 (750); Pd(lll), 740 (730); Ru(OOOl), 760 (755); and Pt(lll), 840 (830). However, the values of 726 (672) cm-' for the two compounds in the liquid state differ by 50 cm-' indicating that the strength of v-bonding of toluene to the surface is systematically reduced because of the presence of the CH3 group.

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NORMAN SHEPPARD AND CARLOS DE LA CRUZ

In the cases of Pd(lll), Ru(0001), and Pt(lll), the assignments of bands to CH3 or C6H5groups have been well authenticated by the use of partially deuterium-substituted molecules. Another second moderate-strength band near 880 cm-' has been assigned to the skeletal breathing mode of the Cg ring of the species on Ru(0001) at 250 K in accordance with the same research group's assignment of the analogous band at 860 cm-' characterizing adsorbed benzene (352). However, it was acknowledged that such an assignment denotes a major change in coupling between vCC of the exocyclic C-C bond and the ring-breathing frequency of toluene. For the free molecule, these give rise to a coupled pair of absorptions, at 1209 and 785 cm-', compared with the uncoupled v2 value for benzene at 992 cm-' (354, 355). Combinations of TPD and VEELS results imply that by 350 K on Pt(ll1) and by 360 K on clean Ru(0001), or at 320 K in the presence of coadsorbed CO, a loss of hydrogen atoms has occurred from the CH3 group (351,352). Avery showed that at first a single hydrogen atom is lost from the adsorbate on Pt(ll1). From the position of the residual vCH absorption of CH2C6D5 (2940 cm-l), the hybridization at the CH2 carbon atom was shown to be sp3rather than sp2. The spectrum is well interpreted in terms of u-bonding of the surface to the CH2 group plus n-bonding of the phenyl group. When toluene was adsorbed on Ru(0001), the yCH mode at 760 cm-' at 250 K did not change in position appreciably after hydrogen loss, implying that the degree of n-bonding was relatively little changed. The limited resolution of VEELS does not enable a clear-cut spectroscopic distinction between CH3 or CH2 groups in the vCH region, i.e., between nondissociative and dissociative adsorption. The higher resolution of RAIRS would help in this respect, but RAIRS has not yet been applied to the problem. Coadsorption of toluene and CO on Ru(0001) leads to improved LEED patterns and different patterns of v C 0 absorptions, but not to appreciable changes in the VEEL spectra of toluene. Wilk et al. (356) made a VEELS, LEED, and TPD study of p- and oxylene adsorbed on Pt(ll1). At 245 K the VEEL spectra are as expected for n-bonded interactions of nondissociated molecules with the benzene ring parallel to the surface. TPD results measured with partially deuteriumsubstituted molecules imply that in each case two methyl hydrogens per molecule had dissociated when the temperature reached 370 K, and the spectra would be consistent with each CH3group being converted to CH2M. The appearance of a few in-plane C6-ring skeletal modes suggests some degree of tilting with respect to the surface. The spectra measured at 550 K suggest the overall decomposition of p-xylene to give species characterized by a(CCH) type spectra (see Section 1V.D). For o-xylene a much changed spectrum was observed at 470 K, which was interpreted to be evidence

VIBRATIONAL SPECTRA OF HYDROCARBONS

265

that the aromatic ring was perpendicular to the surface, multiply u-bonded via the carbon atoms of the original CH3 groups. The spectrum observed at 640 K is similar to that observed with p-xylene at 550 K. A weak RAIR spectrum has been obtained for o-xylene on Ni(100), but none is available for the m- or p-xylene (357). 2. Finely Divided Surfaces At an early stage Shopov and Palazov (242) showed that on Ni/Si02 at ambient temperatures o-xylene and the cis- and trans-1,2-dimethylcyclohexanes gave the same spectrum in the vCH region after the latter had been dehydrogenated by evacuation. The single absorption band at 2915 cm-' (Fig. 13B) was assigned to a CH3 group of an o-xylene skeleton a-bonded to the surface through dissociation of the second CH3group to give a CHMz group. The presence of a-bonds between the adsorbed species and the surface was deduced from the fact that the surface species could be cyclically hydrogenated and dehydrogenated with the species being retained by the surface throughout. The hydrogenated species were consistent with the presence of a-adsorbed methyl-substituted cyclohexanes. Davydov and Sheppard (358)carried out similar experiments characterizing the adsorption of toluene and m- and p-xylenes on Pt/SiO, (Figs. 13A, 13C, and 13D). Spectra, obtained after evacuation, indicated very weak absorptions in both the aromatic and aliphatic vCH regions, the latter being

3100

cm-'

2900

FIG.13. Infrared spectra of methyl-substituted benzenes on Ni/SiOz: (A) toluene; (B) o-xylene; (C) m-xylene; (D) p-xylene. [(A), (C), and (D) from Ref. 358; (B) reprinted from Ref. 242, Kinet. Katal. (Kinet. Catal. Transl.) 8, D . M. Shopov and A. N. Palazov, p. 862 (p. 732 Transl.).Copyright 1967with kind permission of Elsevier Science-NL,Sara Burgerhartstraat 25, 1055 KV Amsterdam, The Netherlands.]

266

NORMAN SHEPPARD AND CARLOS DE LA CRUZ

somewhat stronger for rn-xylene. Much stronger spectra were obtained after hydrogenation, with absorptions characteristic of methyl groups and CH2 groups of cyclohexane rings; after subsequent dehydrogenation, a residual spectrum was obtained in the aliphatic region, often together with very weak aromatic vCH absorptions. This spectrum was cyclically restored to that of the hydrogenated form on repeated cycles of addition and removal of H2, with little loss of intensity per cycle. In the presence of gas-phase H2, vPtH absorptions at 2130 cm-' were always present. For a flat-lying substituted benzene that retains an undissociated methyl group, the MSSR would predict that the vCH3 s mode and one of the vCH3 as modes would be very weak. Only one component of the vCH3 as mode, which usually gives broad and weak bands in the spectrum of the parent hydrocarbon, would be allowed. Such an absorption, at 2970 cm-', is only readily apparent in the spectrum of adsorbed rn-xylene (Fig. 13C). Dissociative adsorption to give a-bonded CH2M groups, as envisaged earlier in the case of toluene on Pt(ll1) at 350 K, might give a vCH2 s absorption at a lower wavenumber, ca. 2930 cm-l. It seems likely that this is the cause of the absorptions near this position in Figs. 13A, 13C, and 13D, although the overall weakness of these aliphatic vCH absorptions would be equally consistent with further dissociation to give CHM2groups. The fact that hydrogenation leads to little release of methyl-substituted cyclohexanes into the gas phase further supports the view that at room temperature the methyl-substituted benzenes are a-bonded to the surface via dissociation of the methyl groups. Gao et al. (346) obtained SER spectra of toluene, isopropylbenzene, and tert-butylbenzene on roughened gold electrodes. The spectra are reasonably interpreted in terms of flat-lying molecules, 7r-bonded to the surface via the aromatic ring. However, one imagines that the n-bonding must be weaker in the presence of the bulky tert-butyl substituent, and for this molecule all the bands in the spectrum of the liquid occur also in the SER spectrum. (The normal Raman spectrum shown in Fig. 2A of the Gao et al. paper (346) is that of isopropylbenzene rather than that of the indicated toluene.) C. GENERAL COMMENTS ON THE SPECTRA OF ADSORBED AROMATIC MOLECULES Avery (258) presented a general discussion of the role of aromaticity in determining the structures of adsorbed cyclic C5 and C6 hydrocarbons. In both cases the extra stability associated with an aromatic ring usually causes C6H6and C5H5to be the preferred structures formed upon dehydrogenation of the adsorbed cyclic alkanes. For example, although Pt a-bonding to the

VIBRATIONAL SPECTRA OF HYDROCARBONS

267

surface seems to be generally preferred over a-bonding for the linear alkenes or cyclic monoalkenes, the latter is still retained for the C6H6 and C5H5rings. However, in the adsorption of benzene itself, the species on the Pt(ll1) surface does show the strongest spectroscopic perturbation relative to that of the gas-phase benzene. Os(O001) may be an extreme case with a-bonding. Avery concluded that this is also the case for C5 rings on Ir(lll), whereby a u-bonded C5H3 species is more stable than a abonded C5H5species. The remaining (important) controversy concerning spectral interpretation of aromatic C6H6 surface species adsorbed parallel to the surface has to do with the identification of the wavenumber of the c 6 breathing mode, v2. With regard to the alkyl-substituted benzenes, it appears to be a general finding that C-H bond breaking occurs more readily at the alkyl substituent than at the aromatic ring, to give u-bonding to the surface through the former.

IX. Acyclic Alkenes: An Update since Part I

A number of recent and very useful reviews have discussed alkenes (among other hydrocarbon species) adsorbed on metal surfaces. Cremer and Somorjai (359) reviewed vibrational spectroscopic and other work on the structure and hydrogenation of ethene on Pt(lll), incorporating recent results from sum-frequency generation (SFG), a recent experimental technique that is discussed later. The principal conclusion (an important one) is that the weakly bonded a-complex of ethene on Pt(ll1) is most likely to be the active surface species involved in hydrogenation. Yagasaki and Masel (360) considered the effect of different crystal faces on catalytic reactions with particular reference to their own extensive work on VEEL spectra of ethene on different single-crystal faces of platinum. Vayssilov (361) discussed the possible structures and catalytic transformations of hydrocarbons on single-crystal metal surfaces in light of evidence from vibrational and other experimental techniques. Zaera (362) provided a valuable and comprehensive review of the coordination, structures, and reactivities of many types of hydrocarbon ligands on metal coordination or cluster compounds and also summarized principal structural conclusions for adsorbed species on single-crystal metal surfaces. Bradshaw (363) also briefly reviewed the relationships between the structures of ligands on metal cluster compounds and those of adsorbed species on metal surfaces. Further pertinent nonspectroscopic determinations of the structure of ethene on single-crystal metal surfaces have recently been published. Using photoelectron diffraction Bradshaw, Woodruff, and co-workers (363) have

268

NORMAN SHEPPARD AND CARLOS D E LA CRUZ

shown that on Ni(ll1) the C-C bond of the di-a-bonded species is approximately parallel to a metal-metal bond on the surface and of length 0.160 2 0.018 nm. This implies the presence, as expected from spectroscopy, of essentially a C-C single bond (364, 365). For ethene on Ni(llO), theory in conjunction with angle-resolved inverse photoemission (36%) suggests a structure, intermediate between di-a- and n-bonded, that is half-bridged on top of the metal-atom ridge. Vibrational spectroscopy (Part I, Section IV.B.3) suggests a di-cr type structure. The group of Bradshaw, Woodruff, and co-workers also studied ethene on Cu(ll0) by both PED and STM (366,367).They showed that the ethene molecules were parallel to and, at low coverage, bridged to a pair of metal atoms of the continuous rows of metal atoms in the (110) direction. The lowcoverage finding is particularly interesting and unexpected as the vibrational spectroscopic studies (218,368,369) imply the presence of n-bonded species on this surface. Such species are normally assumed to be bonded to a single metal atom in the on-top position. At temperatures near 100 K the typical n-complex ethene at low coverage is characterized by absorptions at 1534, 1275 (coupled vC=C and SCH2), and 904 cm-l (yCH), with the first two bands stronger than usual relative to the out-of-plane yCH absorption. At higher coverages, the first two absorptions moved to 1522 and 1257 cm-' (surprisingly implying slightly stronger n-bonding), but the yCH absorption, even more surprisingly, disappeared. This observation implies that the plane of the ethene molecules becomes very substantially tilted, if not perpendicular, relative to the metal surface. It has been suggested that an end-on adsorption of ethene with respect to the surface would best account for the complete lack of absorption characteristic of the yCH mode and the marked strengths of the bands from the coupled vC= C and SCH2 modes (368, 369), but this suggestion is not supported by the results of the higher coverage PED study (366),which are consistent with an on-top n-complex that is less than 24" out of alignment with the metal-atom ridge-a mystery remains! Coadsorption of ethene with oxygen leads once again to a more normal spectrum of a n-adsorbed species (369). A spectroscopic means of determining the direction of tilt of the ethene plane with respect to the surface is in principle available through the inspection of which of the vCH modes, in addition to the fully symmetrical one, becomes active. Unfortunately, several attempts to obtain absorptions in the 300O-cm-' region have failed. Another system that has received recent intensive study by both diffraction and spectroscopic methods is that of ethene on Pt(ll1). In recent papers, Somorjai et al. have investigated the adsorption of ethene by diffuse LEED to give the unexpected result that the di-aspecies [which on Pt(ll1) gives an extreme type I spectrum] is adsorbed across threefold sites with

VIBRATIONAL SPECTRA OF HYDROCARBONS

269

one C-M bond directed to a single Pt atom and the other one to a twofold bridged site (370).The molecule retains C, symmetry with the C-C bond tilted by some 22" with respect to the surface plane. This is an unexpected but, if correct, very significant finding. It is important to seek confirmation of this result for the species on Pt(ll1) and to reinvestigate the structures on the (111) surfaces of other metals. Such a bonding situation could perhaps help explain why ethene on Ni(ll1) shows normal and "soft" vCH absorptions, but the recent PED study of ethene on Ni(ll1) did not describe such an unsymmetrical bonding situation. Can this technique distinguish clearly between a C-C direction parallel to Ni-Ni bonds, or at 30" with respect to this, as required for a structure similar to that described on Pt(lll)? It is, however, to be noted that the adsorption of methyl groups on Ni(ll1) (and Cu(ll1)) appears more likely to occur at hollow sites than adsorption on Pt( 11l), where direct on-top a-adsorption is preferred (Section IV,B and Table 111). It is the di-aspecies on Pt(ll1) which is converted at higher temperature into ethylidyne. This conversion had also been investigated by infraredvisible sum-frequency generation (SFG) by Cremer et al. (371),a welcome first application of this new spectroscopic technique to hydrocarbon adsorption chemistry. They observed an absorption characteristic of an intermediate with a vCH3 band at 2957 cm-l and suggested that this arises from an ethylidene M2CHCH3(or possibly ethyl) species with its C-CH3 axis at an angle to the surface. It is very clear experimentally, as discussed in Part I, that ethylidyne on Pt(ll1) is formed from di-a-ethene and not directly from its T-bonded isomer. The SFG experimental technique involves the interaction of two lasers at the interface, one with a fixed visible frequency, %, and the other of variable frequency, in this context in the infrared region, YR , where vibrational absorptions are expected. It is the beam of frequency (VV + q ~ ) that is detected in the visible region, where detectors are particularly sensitive. The technique has the advantage, shared by second harmonic generation (SHG), of being sensitive only to the surface layer and not to uniform bulk phases. The method can therefore be advantageously applied in normal catalytic reactions, where there are considerable pressures of gaseous reagents in contact with the surface. Under these conditions, at 295 K with 35 Torr of C2H4and 110 Torr of H2, each of the ethylidyne, di-a and, Tcomplexes have been observed to be present on Pt(ll1) (371, 372). After the di-a species had completed its conversion to ethylidyne, just the first and third species remained, but the rate of hydrogenation was little changed. As ethylidyne is well known to be hydrogenated to give ethane much more slowly, it was concluded that the wcomplex is the reactive surface species in the ethene to ethane transformation, with the di-a species being notably

270

NORMAN SHEPPARD AND CARLOS DE LA CRUZ

less reactive. Under certain conditions, absorptions attributable to the surface ethyl intermediate were also observed. Investigating the same C2H4/Pt(lll)system, McDougall and Yates (373) determined the temperature dependence of the di-a to ethylidyne conversion and then ethylidyne decomposition, giving the approximate energies of activation of 17.5 and 29 kcal/mol, respectively. The same authors investigated the slow “hydrogenation” of the ethylidyne species (presumably at ambient temperature) and showed that this accelerated at a pressure of 0.1 mbar, in the pressure range where earlier spectroscopic work had shown the appearance of reversibly formed PtH groups absorbing near 2130 cm-l. The initial addition of a low pressure of H2 led to a shift of the strong ethylidyne absorption from about 1338 to 1336 cm-’ and at 0.1 mbar there was a further shift back to ca. 1342 cm-’. The authors appear to have interpreted this as a removal of the ethylidyne group (to give another CH3containing species?). Kubota et al. (373a) showed that at 112 K on Pt(ll1) a second layer of a v-species, with its molecules tilted toward the surface by rotation about the C-C axis, forms on top of the di-a first layer from ethene. No such layer forms on top of an ethylidyne-covered first layer. A recent study of ethene adsorbed on Fe(100) by Hung and Bernasek (48) showed adsorption as a di-u species (type I spectrum) at 100 K, but with a soft-mode component at ca. 2720 cm-’ in addition to the stronger 2985-cm-’ absorption in the vCH2 region. When the temperature had been raised to 253 K, the spectrum largely changed to that probably indicative of a mixture of CH and a(CCH) species (see also Section 1V.D). At 523 K, only adsorbed carbon remained on the surface. Ethene adsorbed on a c(2 X 2) Mn layer on Pd(100) gives a spectrum profile of type 1‘,as on Pd(100) itself (373b).Ethene on Mo(ll0) at 80 K (373c) gives a spectrum resembling that of a type B, di-a/v, ethyne species; i.e., dehydrogenation appears to have occurred. Preadsorption of oxygen gives a type I1 spectrum from a v-ethene species; preadsorbed carbon gives a spectrum resembling that of a di-a species, but exhibiting a lower than normal vCC mode at 1030 cm-’. Celio et al. (374) showed that adsorption of ethene on a finely divided Pt/Rh alloy supported on A1203led to a characteristic (vC=C/GCH2) absorption at 1292 cm-’, which is contrasted with adsorptions at 1204 or 1230 cm-’ characterizing the species on Pt/A1203 and Rh/A1203, respectively. Other absorptions indicative of the v-C2H4 surface species were little changed. Akemann and Otto (137) obtained SER spectra of the species on cold-deposited In, Cu, Ag, and Au. The two coupled vC=C/BCH2 modes were lowered relative to the value characterizing free ethene, in the order Cu = Au < Ag < In. More profound changes occurred in the SER

VIBRATIONAL SPECTRA OF HYDROCARBONS

271

spectrum of ethene on potassium, but the structural changes were not characterized. The adsorption of perdeuterioethene on a deuterium-covered Pt/SiO, catalyst was characterized by infrared spectroscopy (375). To produce the deuterium-covered catalysts, the sample was exposed to five doses of D2 at 673 K, followed by a short evacuation of the sample. The first spectrum resulting from the adsorption of perdeuterioethene on this sample showed bands associated with gas-phase C2D6and CzD4. This result indicates the occurrence of a degree of self-deuteration on the catalyst. The spectrum obtained after the removal of the gas shows bands of n--C2D4,di-a*-C2D4, CCD3, and a trace of CCD2H,ethylidyne species. The last species probably originated from a small amount of CD2= CDH within the initial C2D4 gas. The infrared spectrum measured after the addition of D2 indicates that all the adsorbed species were deuterated to CzD, and a trace of C2D5H. Cremer ef al. (202a) studied, by SFG in the vCH region, the evolution of propene on Pt(ll1) with increasing temperature from the di-cr species to propylidyne and then possibly to vinylmethylidyne. Ayre and Madix investigated the adsorption of 2-methylpropene (isobutylene) on Ag(ll0) at 115 K (376).The spectrum implied nondissociative adsorption (vC=C, 1605 cm-') but seemingly, to judge from the weakness of the yCH2 band at 890 cm-l, with the planar C4 skeleton substantially tilted with respect to the metal surface. Adsorption on Ag(ll0) with preadsorbed oxygen (0.25 monolayer) led at 115 K to a similar pattern of absorptions, which suffered pronounced changes when the temperature was raised to 224 K, when O H bands also became strong. The spectrum was suggested to be evidence of a 2-methyl-n-ally1species, CH3C(CH2),. A still different spectrum observed at 266 K was thought possibly to be evidence of C(CHz)3. At higher coverage on the clean surface, the spectrum is consistent with the C4plane being parallel to the surface, as has been more recently found . latter paper also gave a spectrum of butby RAIRS at 180 K ( 3 7 6 ~ )The 1-ene on clean Ag(ll0) at 190 K from a n--species with the three-atom C-C =C skeleton parallel to the surface but with the C-CH3 axis inclined at llo", i.e., in the form of the gauche isomer. Bertolini et al. (377) determined the VEEL spectra and NEXAFS of but-1-ene and 13-butadiene on Pd(ll1) and Pt(ll1). They concluded that at 300 K but-1-ene was n--bonded on Pd and di-cr-bonded on Pt, as is now becoming an expected pattern. 1,3-Butadiene was concluded to be di-n-bonded on Pd, but di-o-bonded to Pt via one C-C double bond. Hostetler et al. (378, 378a) concluded from RAIRS and VEELS that norbornadiene at 130-220 K was di-o-bonded to the Pt( 111) surface via one double bond, also forming an agostic soft-bond interaction of the bridgehead CH2 group

272

NORMAN SHEPPARD AND CARLOS D E LA CRUZ

with the surface. It was concluded that at 261 K the latter was converted into a C-M a-bond, and decomposition at 534 K led to adsorbed benzene.

X. The Reactivity of Surface Species: An Example. Kinetic Aspects of the lnterconversion and Hydrogenation of Ethene and Other CpH, Species on Platinum Surfaces A. INTRODUCTION Our article has concentrated on the relationships between Vibrational spectra and the structures of hydrocarbon species adsorbed on metals. Some aspects of reactivities have also been covered, such as the thermal evolution of species on single-crystal surfaces under the UHV conditions necessary for VEELS, the most widely used technique. Wider aspects of reactivity include the important subject of catalytic activity. In catalytic studies, vibrational spectroscopy can also play an important role, but in smaller proportion than in the study of chemisorption. For this reason, it would not be appropriate for us to cover a large fraction of such work in this article. Furthermore, an excellent outline of this broader subject has recently been presented by Zaera (362). Instead, we present a summary account of the kinetic aspects of perhaps the most studied system, namely, the interreactions of ethene and related C2 species, and their hydrogenations, on platinum surfaces. We consider such reactions occurring on both single-crystal faces and metal oxide-supported finely divided catalysts. We use the following general reaction scheme:

*

CH2=CH2(g)

C,H4(a)

intermediate 2H(a)

C2H3(a) + H(a)

4a

*

C2H4(a) C2H3(a) + H(a)

H2k)

(1) (2) (3)

4b

e intermediate +C2H4(a)

(4)

where (8) denotes species in the gas phase and (a) those in the adsorbed phase. Reaction (2) represents a dehydrogenation reaction whereas reactions (4) through (6) represent hydrogenation reactions (see later). In the absence of added gas-phase or preadsorbed hydrogen, reactions (4), (5), and (6) represent self-hydrogenation reactions with H(a) derived from the dehydrogenation of other adsorbed ethene species such as are represented in reaction (2).

VIBRATIONAL SPECTRA OF HYDROCARBONS

273

Under normal experimental conditions (moderate to good coverage and temperatures up to 300 K), the single stable dehydrogenated species observed spectroscopically on Pt(ll1) is ethylidyne, CCH3. Equations (2) and (4) could be generalized to include dehydrogenated species with two or fewer hydrogen atoms per C2unit, such as occur on Pt(100) (C=CH2) and on Ni(ll1) (HC=CH). It should be noted that (depending on the metal face, the coverage, and/ or the temperature) a given symbol can denote different surface species. For example, C2H4(a)could denote a di-a, di-cr*, n-, or n-* species. The last species, a weakly bound species, is sometimes observed in the presence of coadsorbed hydrogen. Which surface species is relevant has to be considered in each particular case.

B. SINGLE-CRYSTAL INVESTIGATIONS OF Pt(ll1) 1. Surface Reactions in the Absence of Added Gas-Phase Hydrogen a. General Comments. Figure 14 shows typical TPD diagrams (379) characterizing desorption from a sample of ethene adsorbed on Pt(ll1) as a function of temperature. Notice that there are three atomic mass unit (amu) signals measured between 200 and 400 K. These are related to C2H4 (amu 28, T,,, = 285 K), G H 6 (amu 30, T,,, = 295 K), and H2 (amu 2, TmaX = 310,510, and 650 K). From the Auger electron spectroscopic C/Pt peak-to-peak ratios as a function of temperature (see inset of Fig. 14), it was found that 54% of C2H4was reversibly adsorbed, 44% remained on the surface to dehydrogenate, and 2% was hydrogenated to C2H6 upon warming. Other data from Berlowitz et al. (380) agree well with this general picture. The higher temperature hydrogen desorption peaks represent dehydrogenation of the initially ethylidyne surface species. b. Kinetics. Tables VIII (381-389) and IX (381) include the basic kinetic data characterizing the reactivity of ethene and other C2H4-derived species on Pt(ll1). These data allow us to discuss our general scheme step by step. Reaction 1

CH2=CH2(g) F==+ C2H4(a)

The rate of ethene desorption can be expressed as Rate = 8C2H4(a)Al exp( -EiIRT), where the subscript or superscript 1 refers to reaction (1) and BC2H4(a), Al , and E: are the adsorbed ethene coverage, the preexponential factor, and activation energy, respectively. 8 is related to the number of metal atoms; e.g., 8 = 0.25 at saturation corresponds to one surface species per four Pt atoms.

274

NORMAN SHEPPARD AND CARLOS DE LA CRUZ

C,H4/Pt(l 11)

0.4

285

0.0 0

-

295

I

1

300

1

I

500

200 400 600 800 Temperature /K

C,H,(amu

28)

t

I

I

700

1

900

Temperature /K FIG.14. TPD diagrams of 1.1 L of ethene adsorbed on Pt(ll1) at 100 K. Inset: plot of C(272 eV)/Pt(237 eV) AES ratios against temperature of ethene adsorption at 100 (--) and 300 K (---). [Reprinted with permission from Ref. 379. Copyright 1988 American Chemical Society.]

Mitchell et al. (390) using nuclear reaction analysis (NRA), found 8 = 0.25 for the saturation adsorption of C2H4on Pt(ll1) at 100 K (also see 391). This result has been confirmed by a combined study done with NRA and XPES (392) and by STM (393-395). Furthermore, the value of 8 = 0.25 for C2H4saturation coverage at low temperature is in agreement with a Monte Carlo simulation of C2H4 adsorption on Pt(ll1) by Windham et al. (396),who showed that an ensemble of four Pt surface atoms is required to absorb one C&H4molecule. The TPD profiles of Fig. 14 show that there are two types of adsorbed C2H4on Pt(ll1). These can be referred to as reversibly and irreversibly adsorbed. The reversibly adsorbed C2H4is inferred to be coverage-depen-

TABLE VIII Kinetic Data for Ethene Desorption, Ethene Decomposition to give Ethylidyne, and Ethylidyne Formation from Ethene on Pt(l1I ) Surfaces Technique Sample

eC&(a)'

E, (kcal mol-I)

A (s-')

T (K)

A. Ethene Desorption 1065-1092 285-291 (311)"

Pt( 111)

h

9-12 (Wd

Pt(ll1) Pt(ll1) Pt( 111) Pt(1ll) Pt( 111)

1

14.9 15-17.3 17 18.3 (18.4 (18.6 18.4 18.3

B. Ethene Decomposition 4 x 10'0 -5 x 10" 5 x 10'2 9.9 x 1011 ii.1 x 1013 (5.5 x 10l1 10l3 (1013)e

to Ethylidyne 236-214 246-294 292-320 231-265 243-276

14.4 (16.7)d 14-16 17.9-18 17

C. Ethylidyne Formation 1.4 X 10'" (3.9 x 10")" (1-6.1) X 10'" 10'2-1013 (10'3)'

Pt(l11) Pt( 111)

1 h h h

Pt(ll1)

1

Pt(ll1) Pt(l1l) Pt(ll1)

1

h h

Perturbationh

Detection'

Reference

TPD

MS

379, 381, 382

LITD TPD TPD LITD LITD

FT-MS SIMS MS FT-MS IT-MS

385 383, 384 383, 384 386 386

TPD

MS RAIRS

382 389

from Ethene 230-280

T-NEXAFS

388

230-265 230-265 223-343

RAIRS RAIRS RAIRS

386, 387 386, 387 389

297 163-343

The term I or h indicates low or high coverage of adsorbed ethene, as inferred from ethene exposures. " TPD, temperature-programmed desorption; LITD, laser-induced thermal desorption; ' FTMS, Fourier-transform mass spectrometry; SIMS, secondary-ion mass spectrometry; MS, mass spectrometry; T-NEXAFS, transient near-edge X-ray absorption fine structure spectroscopy; RAIRS, reflection-absorption infrared spectroscopy. Data for perdeuterio species. 'Estimated value.

TABLE IX Kinetic Data for Ethane Derorption from Ethene Self-Hydrogenation on Pt(ll1) and for Ethane Desorption from a “Hydrogen-Covered” Pt(ll1) Surfuce Technique Sample

KzJ& (a)’

Pt(ll1)

h

Pt(ll1)

h

E, (kcal mol-’)

T (K)

Detection‘

Reference

A. Ethane Desorption from Ethene Self-Hydrogenation -18 302 TPD

MS

381

B. Ethane Desorption from Ethene Hydrogenation 252 TPD

MS

381

- 6

No/e: The superscripts a-c have the same meanings as in Table VIII.

Perturbationb

VIBRATIONAL SPECTRA OF HYDROCARBONS

277

dent, as no thermal desorption was observed at low ethene coverage (387). At high coverage, however, ethene starts to desorb at approximately 230 K (385, 386). At this temperature the irreversibly adsorbed C2H4 also starts to be converted into a dehydrogenated C2 species, probably because the free surface sites allow spaces for decomposition to give adsorbed hydrogen. RAIRS and VEELS have shown that the irreversibly adsorbed CzH4 takes the di-a configuration on Pt(lll), whereas the dehydrogenated species is ethylidyne, CCH3. AES measurements (385) indicate that the ethene desorption occurs at a much higher rate than the dehydrogenation process. The activation energy and preexponential factor for the first process have been estimated from TPD data on the basis of different models (379, 382). The values are 9-12 kcal mo1-l and 106.5-109.2s-l, respectively (Table VIIIA). The activation energy seems to be reasonable when compared with the activation energy value for the CzH4(a) decomposition process. However, the lower limit of the preexponential factor is very small and difficult to explain. It is possible that a particular BCzH4(a)value much lower than 0.25 is responsible for this value. 2a

Reaction 2

C,H4(a)

2b

intermediate

CzH3(a) + H(a)

As mentioned previously, the adsorbed ethene decomposition gives ethylidyne. Then the rate of intermediate formation is given by Rate = BC2H4(a)Aza exp(-E?/RT) and the rate of ethylidyne formation is given by Rate = Bintermediate AZbexp( -Eib/RT). Kinetic data have been obtained with laser-induced thermal desorptiod Fourier-transform mass spectrometry (LITD/FT-MS) (385, 386), secondary-ion mass spectrometry (SIMS) (383, 384), temperature-programmed desorption (TPD) (381-384), RAIRS (386, 387, 389), and transient nearedge X-ray absorption fine structure (T-NEXAFS) spectroscopy (388).The results are summarized in Tables VIIIB and VIIIC. It is important to point out that the more rapid mass-spectrometric techniques monitor species evolved such as CzH$, CzH;, and CzH$ (LITD/FT-MS), CH; (SIMS), and Hz (TPD) from the overall CzH4(a)dehydrogenation process on the surface, whereas the RAIRS and T-NEXAFS techniques monitor the disappearance of di-a-CzH4or the appearance of ethylidyne from the C2H4(a)dehydrogenation process. However, both sets of data show a general trend: the activation energies and preexponential factors for the C2H4(a) decomposition ( E p = 14.9-18.6 kcal mol-' and AZa = 4 X lo1'-9.9 X 1013 s-') and ethylidyne formation (E2b = 13-18.1 kcal mol-I and Azb = 1010-1013s-') increase with increasing adsorbed ethene coverage. This effect has been

278

NORMAN SHEPPARD AND CARLOS D E LA CRUZ

studied in detail with the LITD/FT-MS and RAIRS techniques by Erley et al. (386). These authors found on the one hand that below a critical adsorbed ethene coverage [estimated to be equal to 24 ? 2% of the full ethene Pt(ll1) saturation], the decomposition of C2H4(a) was fast and followed the expected first-order kinetics determined from the LITD/FTMS data (Fig. 15c). Above this critical value, a slower and pronouncedly non-first-order decomposition of C2H4(a) was observed on the basis of LITD/FT-MS data (Fig. 15b). On the other hand, the RAIRS results show strictly first-order kinetics for the ethylidyne buildup on Pt(ll1) for all adsorbed ethene coverages. These results also show that the latter process is always a slow one when compared with the C2H4(a)decomposition (Figs. 15a-c). There is good agreement with the RAIRS results of Mohsin et al. (387) and with the T-NEXAFS results of Gland et al. (388). From the foregoing results, it is clear that one or more intermediates occur on the Pt(1ll) surface with the short lifetimes that can be monitored by LITD and that these are dependent on both the ethene and ethylidyne coverages. It is possible that the ensemble effects observed for the C2H4/

~~

~

0.0

- 0.0

T= 265 K, 27% sat

- -0.5

-0.5

z -1.0

IR

D ._

- -1 .o = Ki C

D

(II

k

--1.5 Xi

-1.5

g

(II

3

-

r

v

E -2.0 z -I

LITD

--2.0 f

I -2.5

-3.0

a 1

0

~

1

100

'

200

1

b '

0I

'

100 I '

'

200

'

0

200

400

600

Time (s)

FIG. 15. Adsorbed ethene decomposition data (laser-induced thermal desorption; mass 27 against time) and ethylidyne formation data (IR signal against time) for selected coverages and reaction temperatures. [Reprinted from Ref. 386, Surt Sci. 301, W. Erley, Y . Li, D. P. Lang, and J. C. Hemminger, p. 177. Copyright 1994 with kind permission of Elsevier ScienceNL, Sara Burgerhartstraat 25, 1055 KV Amsterdam, The Netherlands.]

VIBRATIONAL SPECTRA OF HYDROCARBONS

279

Pt(ll1) system by STM (393, 394) are responsible for the stability of these intermediates. Recent SFG spectra (359, 371, 372) indicate that one such intermediate could be ethylidene, CHCH3, although some authors have preferred a CHCH2 species bonded to the surface through each carbon atom (386). It is not clear how the latter can be described as di-a; tri-a would seem to be more appropriate. This and the ethylidene intermediates could occur in sequence (Section X.D). Reaction 3

2Wa)

=H2(g)

The rate of molecular hydrogen desorption from the Pt(ll1) surface is given by Rate

=

S&,, A3 exp(-Ei/RT).

The adsorbed hydrogen coverage (&(a)) is at present unknown. The activation energy and preexponential factor can be taken from the work of ). are 9 kcal mol-’ and 0.075 s-l, respectively. Christmann et al. ( 3 9 6 ~ These Reaction 4

C2H4a) + H(a)

=C2H,(a),

followed by reactions ( 5 ) and (6). In this step, the well-established dehydrogenated C2H,(a) species is ethylidyne, CCH3. The hydrogenation of ethylidyne to give ethane is clearly an unfavored reaction on a Pt(l1l) surface unless enough adsorbed hydrogen is available from the presence of H2 in the gas phase. Reaction 5

C2H4(a) + H(a)

CZ&(a)

Reaction 6

C2H5(a) + H(a)

C2H6(g)

In the absence of reactant hydrogen, these two steps describe self-hydrogenation. To the best of our knowledge, only one kinetic parameter has been reported in the literature for the overall rate expression of ethene self-hydrogenation (381). This is the activation energy, which has been estimated from TPD data to be 18 kcal mol-l. 2. Surface Reactions in the Presence of Preadsorbed Hydrogen or Deuterium a. C2H4on “Hydrogen-Covered”Pt(ll1). TPD, taken after treating the Pt(ll1) surface with 25 L of H2 followed by 0.4 L of C2H4,gave a broad CzH6 signal centered at 252 K. The activation energy of this process was found to be ca. 6 kcal mol-’ (381).These values are 50 K and 12 kcal mol-’ lower than the corresponding values for C2H6desorption from C2H4 self-hydrogenation on Pt(ll1) (Table IX). The dependence of &H6 production on the surface coverage resulting from dosing H2 and C2H4 has been

280

NORMAN SHEPPARD AND CARLOS DE LA CRUZ

also examined for Pt(ll1) (381). Predosing the surface with increasing exposures of H2, followed by a CzH4 exposure of 0.4 L, gave an increase in G H 6 production and a decreasing temperature for the C2H6 formation. Exposing the surface to 30 L of H2 followed by increasing exposures of C2H4also gave an increase in C2H6production and a decreasing temperature for the C2H6desorption signal. It is clear, as expected, that steps (2) and hence (4) of our general scheme, which involve hydrogen-deficient surface species, are inhibited by the preadsorption of hydrogen on the surface. b. C2D4on “Deuterium-Covered’’ Pt(ll1). TPD taken after predosing the Pt(ll1) surface with 5 L of D2 followed by 0.5 L of C2D4gave a broad C2D6signal with a maximum of 297 K and a shoulder at 265 K (see Fig. 2 of ref 397). The former value is 74 K lower than the corresponding value for C2D6formation from C2D4 self-deuteration on Pt(ll1) (397). A comparison of both results with those of the C2H6desorption following the adsorption of C2H4on the “hydrogen-covered’’ or “hydrogen-free’’ Pt( 111) surface clearly shows the presence of H/D kinetic isotope effects. c. C2D4o n “Hydrogen-Covered’’ Pt(ll1). If C2D4is adsorbed on a “hydrogen-covered” Pt( 111) surface, hydrogeddeuterium-exchange reactions compete with hydrogenation (397). The TPD signal for C2D4H2. observed at 292 K, is (as expected) intermediate between the values for C2H6 (257 K) and C2D6(297 K). More importantly, the TPD maxima for C2D3H and C2D4H2occur at closely similar positions (297 and 292 K, respectively) with similar yields. These results strongly indicate that exchange and hydrogenation proceed through the formation of a common intermediate. This intermediate is almost certainly ethyl, as suggested many years ago by Kemball on the basis of mass-spectrometric analyses of isotopic mixtures (398). This process is today termed P-hydride elimination (362) (see Section X.B.5). 3. C2H4 Hydrogenation or Deuteration in the Presence of Gas-Phase H2 or Dz

a. C2H4 + H2. The dynamic hydrogenation of C2H4to give C2H6 has been investigated in the presence of an initially clean Pt(ll1) surface at temperatures between 300 and 373 K (399).The rate expression was found to be Rate

=

8

X

10’ e x p ( - 1 0 , 8 0 0 1 R T ) P ~ ~ ~ , P ~ ~ 1 ,

where the units of A and E are s-l and cal. mol.-l, respectively. The reaction did not exhibit self-poisoning under the reaction conditions.

VIBRATIONAL SPECTRA OF HYDROCARBONS

281

+

b. C2H4 DZ. The deuteration of C2H4to give C2H6.,D, products has been investigated in the presence of Pt(ll1) at temperatures between 300 and 373 K (399). The rate expression was found to be very similar to that for hydrogenation. However, the reaction with D2was 1.3 times slower than that with Hz. The C2H6.nDndistribution was also characterized by mass spectrometry. Most of the CzH6.,D, contained either one or two deuterium atoms per molecule, although minor products with up to six deuterium atoms were also present. The proportion of the C2H4D2species increased with temperature. c. C2H4 Hydrogenation on an Ethylidyne-Covered Pt(ll1) Surface. Zaera and Somorjai (399)reported an experiment in which a Pt(ll1) surface was saturated with ethene, at ambient temperature, in a two-level ultrahighvacuumlhigh-vacuum apparatus. This was then evacuated, and a C2H4 + H2 mixture was sequentially introduced into the high-pressure section. Circulation was started, and the Pt( 111) surface was heated to temperatures between 300 and 370 K. The ethane formation was followed by gas chromatography and/or mass spectrometry. Reaction rates were determined from the slope of the ethane accumulation curves as a function of time. LEED, HZTPD, and Auger data were obtained characterizing the Pt(ll1) surface after each reaction time in the ultrahigh-vacuum section. These data indicated the presence of ethylidyne on the surface. Furthermore, the hydrogenation activity of ethene was reported to be equal to that observed when starting with a clean Pt(ll1) surface. The lack of any important effect of ethylidyne on the ethene hydrogenation has presented something of a dilemma because (CCH,)Pt(lll) models imply that there is little room left for the adsorption and reaction of ethene (400). Surface hydrogen (or deuterium) atoms which could be present despite the high coverage of ethylidyne would somehow have to be transferred to ethene weakly adsorbed on top of the ethylidyne adlayer (399), or, alternatively, the ethylidyne species would need to move apart from each other under reaction conditions to allow ethene to reach the surface for reaction (400). Several pieces of experimental and theoretical evidence have been gathered to show that ethylidyne is essentially a spectator species during the hydrogenation of ethene on Pt(ll1) (359). Therefore, the first possibility is considered unlikely because of the large static ethylidyne coverage under reaction conditions. It was therefore important to explore the second possibility through ASED-MO calculations (401). These showed that (1) relatively little repulsive energy inhibited the movement of ethylidyne threefold surface site groups closer together than occurs in the (2 X 2) state superstructure and that (2) the activation energy needed to move the species

282

NORMAN SHEPPARD AND CARLOS DE LA CRUZ

through the intervening twofold surface sites was very similar to the activation energy for ethene hydrogenation. This feasible fluxionality on the surface might also account for the poorly ordered (2 X 2) LEED pattern (399) and STM images (394). More recently, SFG (which avoids interfering infrared signals from gasphase reactants or products) has provided direct evidence for the presence of r- and di-cr-ethene species, as well as ethlidyne, when a 110 Torr H2 and 35 Torr C2H4gas stream flowed over a Pt(ll1) surface, giving a hydrogenation turnover rate of ca. 275 s-* per adsorbed complex at 295 K (359, 371, 372). After saturation of the surface with ethylidyne, the same gas stream showed the continued presence of n-adsorbed species, but no div-CzH4, although a high turnover rate was maintained. It was concluded that the r-species is the catalytically active one, even on Pt(lll), and that it can indeed interact with the surface in the presence of a full coverage of ethylidyne. The ethylidyne fluxionality model was hence preferred. 4. Hydrogenation of an Ethylidyne-Covered Pt(ll1) Surface

McDougall and Yates (373) measured the rate of hydrogenation of an ethylidyne adlayer on a Pt(ll1) surface at H2 pressures exceeding 1 mbar (75 X lo-’ Torr). They followed the decay of the concentration of ethylidyne CCH3 per cm2 of (based on a saturation ethylidyne coverage of 6 X surface) by using the normalized integrated intensity of the RAIRS band at ca. 1339 cm-’. Analysis of these data gave initial rates ranging from 0.9 X lo1’ CCH3 cm-’s-’ at 20 mbar to 3 X lo1’ CCH3 cm-’ ss1 at 1333 mbar of H2. At pressures less than 1 mbar, exposure to H2 brought about only subtle changes in the 1339-cm-’ band of ethylidyne. Similar conclusions have also been reached on the basis of 14Cradiotracer measurements of the removal of ethylidyne by hydrogenation on a Pt(ll1) surface (402).

5. Reactivity of Surface Ethyl Groups on Pt(ll1) The possibility of the production of surface ethyl groups by decomposition of alkyl halides on metal surfaces (Section 1V.E) also opened up the possibility of spectroscopically studying their ranges of stability and reactivity. In a very illuminating paper, Lloyd et al. (192)(using results from both surface CZH5 and CzD5produced by photodecomposition) showed that on Pt(ll1) these groups decompose in the sequence

-

K C2H5 (160 K) -c230 -+ n(C2H4)

1230 K

di-a(C2H4)

CCH3

fragments.

VIBRATIONAL SPECTRA OF HYDROCARBONS

283

The precise transition temperatures might depend on the presence of the necessarily coadsorbed C1 atoms, but the decomposition sequence is probably independent of this. Ethene itself, after adsorption on Pt(ll1) at very low temperature, has shown a transition from T- to di-c+-bondingat ca. 50 K (403). As shown very successfully by Zaera (397), who used H/D isotopically substituted molecules, the decomposition of the ethyl group occurs by 0-hydride elimination followed by readsorption, as was envisaged years ago by Kemball (398) to account for patterns of H/D exchange in ethene and ethane produced from the C2H4D2reaction. This is now widely accepted as constituting the first decomposition step of surface alkyl groups on a variety of metal surfaces (199, 397). The kinetic results show that as the temperature is raised to room temperature, ethyl on a Pt(ll1) surface would only be a transient species, the more so in the presence of hydrogen. Backman and Masel (404) may have obtained a VEEL spectrum of ethyl on Pt(ll1) by the back-reaction of the slow hydrogenation of ethylidyne at 298 K.

C . ETHENEAND HYDROGENATION REACTIONS ON OXIDE-SUPPORTED Pt CATALYSTS 1. General Comments

Jackson et al. (405) reported C2H4pulsed-flow chemisorption measurements of ethene on Pt/A120, at 193,273, and 294 K. At 193 K, C2H6 began to desorb after the adsorption of ethene from four pulses injected over the catalyst. After seven pulses had been passed over the catalyst, an elution of C2H4, along with C2H6,began. By the ninth pulse, only C2H4 was eluted from the reactor. At 273 K, the first pulse of C2H4 resulted in 30% selfhydrogenation to C2H6,the remainder being retained by the catalyst. The second pulse produced 26% C2H6 while 20% was retained. No further self-hydrogenation was observed with subsequent pulses, although further adsorption occurred until effective saturation of the surface was achieved. At 294 K, adsorption of the first pulse of C2H4produced 63% C2H6, the remainder being retained by the catalyst; 26% of the second pulse was hydrogenated to G H , and a further 22.2% was retained. Further pulses of C2H4 produced no further hydrogenation, although some adsorption occurred until surface saturation was eventually achieved. These observations clearly indicate that the chemisorption of ethene is characterized by a primary process in which permanently retained hydrocarbon species are chemisorbed on the metal surface and by a secondary process which is responsible for the adsorption and then self-hydrogenation of the catalyti-

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NORMAN SHEPPARD AND CARLOS DE LA CRUZ

cally active species on the surface. This conclusion is in line with the findings from the GH,/single-crystal Pt surface systems (Section X.B.l).

2. IdentiJication of Chemisorbed Species Formed f r o m Adsorption of Ethene or Perdeuterioethene on Oxide-Supported Pt Catalysts Infrared and 2H NMR spectroscopic studies have shown that the adsorption of ethene or perdeuterioethene on Pt/Si02 and Pt/A1203 catalysts at low temperature (typically between 180 and 200 K) produces r-C2H4 (406-408), di-a-C2H4(407-409), and (seemingly on larger metal particles) di-a*-C2H4(407) chemisorbed species. On warming to ambient temperature, some of the r-C2H4 species desorb (407, 408) while the di-n-C2H4 species is converted to ethylidyne. The di-a*-C2H4species, when it occurs, seems to be stable from low to ambient temperatures. Comparisons with the vibrational spectroscopic studies of the adsorption and dehydrogenation of ethene on single-crystal Pt surfaces (Section X.B.l) show that the di-a-C2H4to ethylidyne conversion occurs on (111) facets of the Pt crystallites of the catalysts. It is considered that the di-a*-C2H4 species occur on metal sites on which this conversion is not allowed, perhaps on (loo), (110), or (210) facets. It is not clear whether the labile n-C2H4 species is formed on amorphous areas of the clean Pt particles or whether it occurs on sites which are affected by proximity of the metal oxide support (408);we favor the former possibility.

3. Kinetic Data for Ethylidyne Formation The activation energy and preexponential factor characterizing the perdeuterioethylidyne (CCD3) formation from C2D4 on a Pt/A1203catalyst have been measured at temperatures between 245 and 285 K by 'H NMR spectroscopy (409, 410). These values were found to be equal to 15 kcal mol-' and 9 X lo7 s-', respectively. The first value is in line with the activation energy of CCD3 formation on a Pt(ll1) surface at temperatures between 230 and 280 K (Table VIIIC). However, the value of the preexponential factor is much smaller than the one found for the Pt(ll1) case, i.e., 3.9 X 10" s-l. Perhaps the site heterogeneity of the small Pt particles of the catalyst is responsible for this discrepancy. 4. Evaluation of the Hydrogenation Reactivities of Chemisorbed Species on Oxide-Supported Pt Catalysts Figure 16 shows a series of infrared spectra which correspond to the hydrogenation of the T - C ~ Hdi-a-CzH4, ~, and ethylidyne species on a Pt/ A1203catalyst at 140 K, taken from a paper by Mohsin et al. (408).Notice

VIBRATIONAL SPECTRA OF HYDROCARBONS I

I

I 1100

285

1339 1200

1300

Wavenumber (cm-')

I 1400 2750

2950

3150

Wavenumber (cm-')

FIG. 16. Infrared spectra that show the effects on a Pt/AI2O3 catalyst, initially with 1 ~ C2H4, di-a-C2H4,and ethylidyne species, of a sequence of doses of H2: (a) sample under vacuum at 140 K; (b) after adding 5.3 X 10l6 molecules of H2; (c) after adding 2.4 X 10'' molecules of H2; (d) after warming to 200 K (e) after adding 6.3 X lo'* molecules of H2; (f) after evacuating the sample cell and warming to 200 K. [Reprinted with permission from Ref. 408. Copyright 1988 American Chemical Society.]

that the bands of the n-C2H4species at 1203,2955, and 3018 cm-' and the band of the di-a-C2H4species at 2912 cm-' were both rapidly removed at temperatures between 140 and 200 K. Furthermore, Mohsin et al. (411) showed that at ambient temperature, when only r-C2H4and ethylidyne species were present on the catalyst, hydrogenation led to rapid elimination of the r-C2H4species but to much slower removal of ethylidyne. It would be valuable to repeat kinetic studies at low temperatures so that only the di-a-C2H4and .rr-C2H4species are present. It should be noted that the r species on metal oxide-supported catalysts are probably on rough, non( l l l ) , surfaces of the metal particles. Using a similar Pt/SiO, catalyst, De La Cruz and Sheppard (197) showed that a very small dose of H2 (3 X Torr) at ambient temperature led to a weak but well-defined spectrum indicative of a possible ethyl surface species. Recently, McGee et al. (195) showed that the ethyl species on Pt/ A1203,derived from the thermal decomposition of ethyl chloride, decomposes to give the n-C2H4species and ethylidyne, although in this case the di-a-C2H4intermediate was not identified. Hensley and Kesmodel (412)

286

NORMAN SHEPPARD AND CARLOS DE LA CRUZ

obtained VEEL spectra of the di-a-C2H4 and ethylidyne species on Pt/ A1203.

5. Isotopic Scrambling between Ethylidyne Species on Oxide-Supported Pt Catalysts In addition to the hydrogenation reactivities of the n-C2H4 and di-uC2H4species reported in the previous section, there is infrared (413) and 2H NMR (410) evidence of isotopic scrambling between ethylidyne species on oxide-supported Pt catalysts. The infrared spectrum of the species formed from the adsorption of CH2= CD2 on a Pt/Si02catalyst at ambient temperature shows bands that can be associated with the following ethylidyne species: CCHD2,CCH2D, and CCH3 (413). The formation of CCHD2 and CCH2D species requires the transfer of one H (or D) atom to the other carbon atom of the parent ethene molecule, together with the initial adsorption of another H/D atom onto the Pt crystallites. The formation of CCH3 species implies that H/D interchange occurs between separate ethylidynes; i.e., the process is not totally intramolecular, probably facilitated by PtH/PtD mobility and exchange. The same phenomenon has been observed to occur on a Pt(ll1) surface (413) and may be unavoidable on Pt at temperatures near 300 K, at which ethylidyne is formed from the di-a-C2H4species. 2H NMR spectroscopic studies have shown that, starting initially from CCH3 and adsorbed deuterium atoms on a Pt/Al,O, catalyst, it is possible to produce CCH3,CCH2D,CCHD2,and CCD3(410).The activation energy and preexponential factor for the overall CCH3(a) + 3D(a)

CCD3(a) + 3H(a)

exchange were found to be equal to 14.3 kcal mo1-l and 4 respectively, at temperatures between 255 and 295 K.

X

lo7 s-l,

6. The Kinetic Hydrogenation or Deuteration of Ethene on OxideSupported Pt Catalysts a. Kinetic Studies of Hydrogenation. ethene is given by Rate

=

The rate of hydrogenation of

A exp( - E,/RT)PEzH4Ph2,

where A , E,, and a and b are the preexponential factor, activation energy, and reaction orders in the partial pressures of ethene (PCZH,)and hydrogen (P,,), respectively. The reaction involves contributions from reactions (5) and (6), with one or the other of these being the rate determining. A recent investigation of a Pt/Si02 catalyst by Cortright et al. (414)

VIBRATIONAL SPECTRA OF HYDROCARBONS

287

revealed that the reaction orders in C2H4 and in H2 were temperature dependent. As the temperature was raised from 223 to 336 K, the authors observed that the reaction varied from approximately half-order in H2 at low temperature to first-order at high temperature. The change in GH, order was also temperature dependent, with a more severe pressure dependence at high temperature (336 K) than at low temperature (223 K). The order in CzH4 varied approximately from -0.5 to zero with increasing ethene pressure. The zero-order behavior implies that the surface was saturated with ethene under these conditions. The variation in H2 order could be explained by a mechanism in which there are competitive and noncompetitive adsorption sites for hydrogenation. Hydrogen competes with ethene for sites at low C2H4 partial pressures (and high temperatures), but only the noncompetitive adsorption site for hydrogen is available at high G H 4 partial pressures (and low temperature). Table X (414-419) is a summary of these data with similar results from investigations of other (C2H4 + H2)/metal oxide-supported Pt catalysts. The table also includes analogous data for C2H4 + H2 reacting in the presence of Pt wires, evaporated Pt films, or Pt(ll1) single crystal. The overall agreement is most satisfactory for the different forms of platinum catalysts. Table X shows that the activation energy for the ethane formation from C2H4 + H2varies only between 8.6 and 10.8 kcal mo1-l for Pt/Si02 and PtlAl2O3catalysts, Pt wires and films, and single-crystal Pt(ll1). Some values between 11 and 16 kcal mol-' have also been reported (416, 417). These have been associated with contaminated surface (415). b. Deuterium Tracer Studies. Isotopic tracer studies have been conducted in which C2H4 was reacted with D2 in the presence of oxide-supported Pt catalysts, and the products were analyzed for the isotopic forms of gaseous ethane, ethene, and dihydrogen (417, 419, 420). These studies indicate the following: (i) the products included deuterated ethane with zero to six deuterium atoms; (ii) the maxima in the ethane isotopic distributions were concentrated in ethanes with two or one deuterium atoms; and (iii) broadened distributions were observed with increasing temperature. The mechanism that can account for the formation of the observed ethanes involves the reversible formation of an ethyl radical as the half-hydrogenated state (398).

c. Kinetidlnfrared Spectroscopic Studies. Sheppard et al. (421) were the first to perform spectroscopic studies of the adsorbed species on metal

TABLE X Typical Data for the Formation of Ethene from (CZH4+ HZ) on Several Oxide-Supported Pt Catalysts, Pt Wires, an Evaporated Pt Film, and a Pt(ll1) Surface Reaction order Sample Pt/SiO, Pt/SiOz Pt/SiOn PtlAIZ03 Pt/A1203 Pt wire Pt wire Pt wire Pt film Pt(ll1)

pH,

PC,H,

25 23

150 152

100 39 26 25 42 20

100 117 116 150 41 100

C2H4

HZ

E , (kcal mol-I)

-0.43

1.10

-0.5 -0.3 -0.5

1.2 1.0 1.2

8.6 8.9-9.1 9.7 9.9

-0.33 0 -0.6

1.03 1 1.31

A (s-')

10 10.7 10.8

T (K)

Reference

223-336 213-248 177-206 293 273

414 415 416 419 41 7 418 418 414 415 399

357-655 336 8 x 108

300-333

VIBRATIONAL SPECTRA OF HYDROCARBONS

289

surfaces under conditions of dynamic equilibrium, In this set of experiments the hydrogenation of CzH4 to CzH6 took place in a cell through which passed regulated room-temperature flows of the reactants over a cooled catalyst. The cell was of double-beam design to minimize the bands of the gas-phase components in the spectrum. However, the production of ethane in the vicinity of the catalyst and the lower temperature in the catalyst side of the cell led, nevertheless, to the appearance of its gas-phase bands at 2985 and 2896 cm-l. Low temperatures between ca. 190 and 215 K were used to give conveniently measurable reaction rates, but the actual temperatures measured in the gas phase near the catalyst were only of qualitative significance. The infrared beam passed through a sequence of four catalyst disks to enhance spectroscopic sensitivity, but this led to a sloping background in the 3000-cm-' region. Figure 17 shows the spectra obtained during the experiment. Spectrum A is essentially that of physically adsorbed ethane, and spectrum B is the spectrum of gas-phase ethane with its sharp prominent Q branches. These spectra correspond to C2H4/H2molar ratios of 0.23 and 0.48, respectively. When the CzH4/H2molar ratio was increased to 0.67 and further to 0.93, new bands were observed, at ca. 2920,2885, and 2790 cm-l. The 2920-cm-' band can be assigned to a di-a*-CzH4species, and the 2885- and 2790-cm-' bands are indicative of ethylidyne. On cutting off the supply of ethene (right-hand side of Fig. 17), there was clearly at first a rapid drop in temperature, leading to the condensation of ethane. This result shows that at equilibrium the temperature of the catalyst itself depends in part on the amount of heat supplied by the exothermic hydrogenation reaction. In another context (422),it has been shown that for this reason there can be a substantial increase in the temperature of the metal particles relative to that of the oxide support when a catalytic reaction is in progress. This is a reason for regarding the temperatures measured in the C2H4/H2gas stream close to the catalyst as only of qualitative significance. As the liquid ethane is flushed off by the continuing stream of hydrogen, it is clear that the absorptions of the di-cT*-C2H4species disappear rapidly, and those of ethylidyne disappear much more slowly, as described earlier in Section X.B.3.c. It had been originally hoped to detect the presence of a nreacting species, or even of an ethyl intermediate, in the spectra obtained under flowing conditions. Unfortunately, the unexpected uncanceled absorptions representative of the ethane products obscured the spectral regions expected for each of these. This problem has recently been overcome in principle by the use of SFG for a reaction occurring on Pt(lll), as this technique does not give signals from bulk phases (Section X.B.3.c). However, this technique is not applicable to finely divided catalysts, but only to flat surfaces. Measurements were also made of the conversions of ethene to ethane

/

n

z

U

60 I

n

W

1

z W J

W

c I

>

sssz-

m 0

N

d

0

m r n

m 0

VIBRATIONAL SPECTRA OF HYDROCARBONS

291

by gas chromatography, giving the values of 50,20,17, and 13%,respectively, corresponding to spectra A-D. These relative values do not conform to the normal kinetics of ethene hydrogen-atom addition (Table X), and it was concluded that the measured reaction rates most probably were determined by the rates of diffusion of reactants and products to or from this surface within the fine pores of the pressed-disk catalyst, rather than by the actual catalytic surface reactions. Soma (423) later performed excellent quantitative infrared measurements of the hydrogenation of ethene on a Pt/AlZO3catalyst which clearly showed a dominant contribution from the n--species. These measurements were carried out as follows: The amount of hydrogen adsorbed on Pt was estimated from the integrated band intensity of vPt-H at 2120 cm-'. The amount of n-CzH4 was estimated from the integrated band intensity of the coupled vC=C/GCH2 vibrational band at 1205 cm-'. The full coverages from hydrogen, & = 1,and n--CzH4,n-&, were determined from the respective saturated adsorptions at the reaction temperature under gas pressures of approximately 100 Torr. A typical result of the hydrogenation in the presence of Pt/AlZO3at 203 K is shown in Fig. 18. The changes of the absorption bands from those characteristic of T - C ~ H species ~ and dissociatively adsorbed hydrogen with time were so small compared with the reaction rate that the observed reaction was considered to be in the steady state. The rate of the reaction and the adsorbed amounts of hydrogen and n-C2H4species were measured during the course of the reaction with different ratios of H2 and CzH4 partial pressures (PH/PE)in the range between 0.5 and 2. For a given PHIPE,the amounts of adsorbed hydrogen and ethene depended somewhat on the first gas adsorbed. When hydrogen was first admitted and ethene was then added to initiate the reaction, the amount of adsorbed hydrogen was more than when ethene was preadsorbed. Other adsorbed species such as ethyl, ethylidyne, or di-n-*-C2H4species were not detected during the reaction at or after the removal of the gas phase by evacuation. The sum of & and & amounted to 80-90% of the surface coverage and sometimes exceeded 100%when hydrogen was first admitted. The dependences of the reaction rate upon CzH4 and H2 partial pressures are shown in Fig. 19. The partial pressure dependence of the reaction rate for both CzH4 and H2 showed maxima. This result indicates that a Langmuir-Hinshelwood type surface reaction, whereby the two reactants compete for surface sites, determines the reaction rate. However, Soma's assumption that only the n--species was active requires reinvestigation as ethene adsorption on the same type of catalyst under similar conditions (406) showed a substantial absorption at 1420 cm-l. At that time Soma assumed that this was also indicative of the n--species, but we now know that it could have been indicative of the di-u species.

292

NORMAN SHEPPARD AND CARLOS DE LA CRUZ

i a

Time (min)

I

I

0

evac

50

100

150

Time (min)

FIG. 18. Hydrogenation of ethene on Pt/A1203 at 203 K: (a) pressure change and (b) amounts of Pt-H and Pt(r-CZH4)adsorbed species during the reaction. & = 1 corresponds to 1.48 X molecules of surface Pt-H and & = 1 to 9.08 X 10'* molecules of Pt(n-C2H4). [From Ref. 823.1

7. Hydrogenation of Ethylidyne on Oxide-Supported Pt Catalysts In a hydrogen atmosphere at 195 K, (what we now know to be) the ethylidyne species was shown by Soma (406) to be more stable than the n--C2H4species because the band at 1337 cm-' (SCH3s) scarcely changed its intensity. However, the intensity of this band decreased in the presence of hydrogen when the cell temperature was raised to 243 K, and ethane was produced in the gas phase. Ethylidyne is therefore not strictly a spectator species, but its contribution (when present) to the rate of ethene hydrogenation is very low.

293

VIBRATIONAL SPECTRA OF HYDROCARBONS

,,I,

Ib I

la P, = 100Torr

200 P, Torr

0

100

200

P, Torr

FIG. 19. Pressure dependence of the hydrogenation rate of C2H4on PtlA120, at 203 K: (a) dependence on HZpressure when the ethene pressure ( P E )was 100 Torr; (B) dependence on CZH4pressure when the hydrogen pressure (PH)was 100 Torr. The symbols (0)and (+) indicate that H2was admitted first and then CzH4was added or vice versa, respectively. [From Ref. 423.1

D. REACTION MECHANISMS: PRESENT PERSPECTIVES Soma’s excellent infrared and kinetic study of ethene hydrogenation catalyzed by Pt/A1203 (423) showed clearly the dominant role played by the n-adsorbed ethene species and by the reversibly adsorbed hydrogen that occurs at higher pressures in the form of on-top PtH. It also pointed to a Langmuir-Hinshelwood mechanism as the n-adsorbed ethene was shown to compete with adsorbed H atoms for surface sites. Cremer et al. (359,371,372) have also recently shown by the use of SFG that the n--species, reactive for hydrogenation even on ethylidyne-covered Pt(ll1) surfaces, is much more reactive in catalysis than the di-cr one. A turnover number of 275 s-l per adsorbed n-ethene species has been measured for a 35 Torr CzH4and 100 Torr H2 mixture at 295 K in the presence of Pt(ll1). This measurement was made under conditions of approximately zero order reaction in ethene and 0.8 order reaction in hydrogen. It should also be recalled that a particularly weakly perturbed n-type species, designated n*, has been identified on Pt(ll0) and Pt(210) surfaces when preadsorbed hydrogen was present (424).On the basis of TPD work, Bowker et al. (425)recently provided evidence that such a species is primarily involved in ethene hydrogenation catalyzed by Rh(ll1). It remains to be seen whether this is more generally the case. A distinction between n and n* required the detection of the uC=C/S=CHZ absorption in the

294

NORMAN SHEPPARD AND CARLOS DE LA CRUZ

1600- to 1 5 0 0 - ~ m -region, ~ but this information was not available in the SFG Pt(1ll) study because of the, at present, limited wavenumber range of infrared lasers. It is seen that Horiuti and Polanyi’s 1934 mechanism for hydrogenation (426) has stood up very well to detailed spectroscopic scrutiny, with the proviso that the reactive ethene surface species is to be identified as of a T rather than of a di-a type. This mechanism is now as follows: C2H4(g) + H2(g) * (nC2H4)(a) + 2H(a) C2HS(a) + H(a)

+

C2H6(g).

The reversibility of the second-stage P-hydride elimination, involving an ethyl intermediate (now spectroscopically identified) also confirms Kemball’s early mechanism for H/D isotopic substitution (398). The sequence is as follows:

-

C2Hdg) + W g ) + ( G H d ( a ) + 2Wa) -+ CZH,D(~)+ D(a> C2H,D(g) + H(a) + D(a), with repetitions of the process for the attaining of higher degrees of substitution by deuterium. Di-a-ethene has also been separately detected by spectroscopy but found to play a different mechanistic role related to its stronger bonding to the metal surface. This, not the n-species, has been shown to be the precursor in the formation of ethylidyne by surface dehydrogenation on Pt(ll1) and probably on other facets with triangular metal sites. Ethylidene is now the favored intermediate in this reaction (359,371,372,427),which can formally be represented as C2H4(a)+ CHCH2(a) + H(a)+ CHCH3(a)-+ CCH3(a) + H(a). The recently proposed structure for the di-aspecies on Pt(l11) [obtained by diffuse LEED (370)],in which one of the C-Pt cr-bonds bridges a pair of metal atoms to give a structure of symmetry C,, would seem to be in a geometrically favorable position to lose a hydrogen atom. This would give a tri-a M2CH-CH2M species as opposed to the more normally proposed air vinyl. An alternative description of the first step in the preceding sequence could therefore be proposed as (M2)CHZ-CH2M

+ M2CH-CH2M.

Janssens and Zaera (198)suggested that at high coverage the first two steps may be reduced to a single concerted 1 + 2 internal shift of a H atom as a vacant site is required for the hydrogen atom released in forming the (CHCHz) intermediate. The combined LITD-MS and infrared study referred to in Section X.B.1.b implies that either the second or third step in

VIBRATIONAL SPECTRA OF HYDROCARBONS

295

the sequence is the slow (rate-determining) one on Pt(ll1). It should be recalled that platinum is a metal which strongly favors a-bonding of hydrocarbon species. On Pd, for example, which more strongly favors .rr-bonding, the first intermediate in the above sequence may well be a a/rvinyl species. Vibrational spectroscopy has also shown that ethylidyne is the dominant “carbonaceous” species up to temperatures >300 K and that this is hydrogenated only very slowly. This finding does not support the suggestion of Thomson and Webb (428) that the carbonaceous species is active in this reaction as a carrier of hydrogen. It is seen that vibrational spectroscopy, when the results are considered in conjunction with the results of other physical techniques, has already made useful contributions to the study of the kinetics of ethene-related reactions on platinum surfaces.

XI.

Looking Ahead: Some Suggested Priorities for Future Research

A. GENERAL COMMENTS ON VIBRATIONAL SPECTROSCOPIC EXPERIMENTAL TECHNIQUES 1. VE E L S The advantages of VEELS are its wide wavenumber coverage (4000-200 cm-’ ) and high sensitivity biased in the direction of low-wavenumber bands. Recent advances in experimental resolution, in resolution enhancement, and in Bayesian noise-reduction techniques will enable researchers to obtain better quality spectra in the future. Ready information about noncompletely-symmetrical modes through the impact mechanism is also available by VEELS. More use needs to be made of the impact selection rules at the specular direction as a function of the azimuthal direction of the plane of incidence with respect to symmetry directions on single-crystal surfaces, e.g., for ( l l l ) , (loo), and (110) faces, incident-plane measurements along the directions of in-contact rows of metal atoms, and at 30” (or 90”), 45”, and 90” angles to these directions, respectively. There is the probability that the use of different directions of the plane of incidence relative to such surface crystallographic directions is responsible for some of the spectral differences reported by different laboratories for the same systems, and perhaps, more specifically, for differences between some type I and type I‘ VEEL spectra of ethene and type A and type A‘ spectra of ethyne. VEEL spectra also most readily give low-wavenumber features of the “frustrated translation” or “frustrated rotation” types which are intimately related to the bonding of adsorbates to surfaces.

296

NORMAN SHEPPARD AND CARLOS DE LA CRUZ

2. RAIRS More studies in the low-wavenumber region (<800 cm-') would be valuable in making use of the higher resolution and strict MSSR compatibility of this technique relative to VEELS. The increasing use of synchrotron radiation for this purpose would appear to be well suited to single-crystal work. The capability of RAIRS to characterize species on single-crystal surfaces under higher (catalytically realistic) pressures with single or multiple adsorbates (e.g., for the hydrogenation of alkenes) has yet to be fully exploited. Polarization-modulation techniques should enable the compensation of gas-phase contributions to the spectrum except in cases of nearblackout gas-phase absorptions. Step-scan FTIR spectrometers will increasingly contribute to time-dependent studies and the identification of adsorption intermediates in catalytic processes with subsecond rates, particularly when combined with low-temperature work. 3. Sum-Frequency Generation (SFG)

Although as yet seemingly restricted to above ca. 1500 cm-' by the limited availability of tuneable infrared detectors, this technique also virtually eliminates gas-phase contributions to spectra. The pulsed lasers used also open up the possibilities of fast (nanosecond or less) kinetic studies of catalytic reactions. 4.

Raman Spectroscopy

Raman spectra of adsorbed species, when obtainable, are of great importance because of the very different intensity distributions among the observable modes (e.g., the skeletal breathing frequency of benzene) compared with those observed by infrared spectroscopy and because Raman spectra of species on oxide-supported metals have a much wider metal oxidetransparent wavenumber range than infrared spectra. Such unenhanced spectra remain extremely weak for species on single-crystal surfaces, but renewed efforts should be made with finely divided catalysts, possibly involving pulsed-laser operation to minimize adsorbate decomposition. Renewed efforts should be made to obtain SER and normal Raman spectra characterizing adsorption on surfaces of the transition metals such as Ni, Pd, or Pt, by use of controlled particle sizes or UV excitation, respectively.

5. Transmission and Diffuse-Reflectance Infrared Techniques In the infrared region, and possibly also by Raman spectroscopy, more studies should be made of catalysts and catalytic reactions carried out in situ. By use of conventional transmission infrared techniques, it is also

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nowadays much more feasible to study directly metal or bimetal, e.g., Rh/Pt, catalysts of low metal loading such as are used in practice in industry. More low-temperature infrared studies of finely divided catalysts are needed to provide cross-connections between work on these and single-crystal surfaces. The initial nondissociatively adsorbed species are more likely to be present, and first-formed species during reaction (such as ethyl from ethene and hydrogen) can be investigated before decomposition. The necessity to avoid condensation on the high-area metal oxide support may limit such work to low-molecular-weight hydrocarbons. Increasing developments in infrared microscopy might lead to spectroscopic studies of the adsorbed species on different faces of well-formed three-dimensional single crystals of dimensions of a fraction of a millimeter.

B. SUGGESTIONS FOR STUDIES OF PARTICULAR HYDROCARBON/METAL ADSORBATE/ADSORBENT INTERACTIONS We suggest some priority areas as follows: 1. Ethene [Page References to Part I (l)] On Single- Crystal Surfaces

a. To reinvestigate a low temperature species adsorbed on Ni(100) (p. 50 and Table VI) and on Ir(ll1) (p. 61) to obtain spectra of nondissociated species. b. To develop further the work of Masel’s group on species adsorbed on Pt(100), Pt(llO), and Pt(210) planes at higher resolution by using RAIRS or VEELS to differentiate better between coexisting di-c+and n-species (pp. 31 and 64). c. To carry out similar multiplane studies as in (b) for additional metals. d. To reinvestigate type I’ spectra of adsorbates on Ru(0001), Pd(100), and Fe(ll1) to check whether experimental variations in the VEEL spectra are related to different azimuthal directions of the plane of incidence relative to the surface crystallographic directions (pp. 16 and 63, and ref 17, Fig. 2). e. To determine whether the HCCH species obtained at room temperature on Ni(ll1) following the low-temperature adsorption of ethene transforms into ethylidyne (as found on Ni/Si02)on addition of hydrogen (p. 50). f. To study “soft” vCHIVCD modes of a di-cr species [Ni(lll)] or a Tspecies [Pd(lll)] (pp. 65 and 66; Table VI) by using the deuteriumsubstituted ethenes CH2CD2and cis- and trans-CHDCHD to investigate the direction of tilt that causes the agostic interaction.

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g. To study certain crystal planes, e.g., Pt(ll1) or Ir(lll), for adsorption at low temperatures on near-perfect or deliberately roughened (unannealed) surfaces to determine to what extent surface defects lead to spectra of different adsorbed species coexisting with those on flat areas (p. 66).

On Finely Divided Catalysts h. To reinvestigate Ni, Pd, and Rh at low temperatures to obtain better spectra in the vCH region, particularly of di-aspecies, and to determine how general is the di-a(rather than n) to ethylidyne transformation that is now well known to occur on (111) facets of Pt (pp. 44 and 52). i. To investigate low-temperature hydrogenation on several metals to better establish the spectrum of the surface ethyl species (pp. 38 and 39). It could be advantageous to repeat an elegant pulsed ethenehydrogenation experiment by Hattori and Burwell (429) using infrared spectroscopy. j. To determine the relative rates of hydrogenation of di-a and n-species at low temperature on several metals (pp. 39 and 40). k. To study systematically, e.g., for Pt and Pd, the effect of temperature of reduction on the capability of a catalyst to adsorb n-species (pp. 37 and 38). 1. To reinvestigate, e.g., for Ni and Pd, whether there are differences in the spectra, and hence adsorbed species, formed from “hydrogencovered” and “hydrogen-depleted” catalysts at low and room temperature (p. 53). 2. Higher Alkenes [Page References to Part I ( I ) ] a. To investigate the VEEL and/or RAIR spectra of propene and the linear butenes as a function of temperature and coverage on additional metals such as Pd(ll1) and Ni(ll1) and particularly to investigate the interconversion of the linear butenes as a function of the metal surface and temperature (p. 82). b. To use RAIRS to investigate the hydrogenation of the alkenes on different crystal faces of particular metals, e.g., Ni and Pt, to determine which planes give complete hydrogenation to the alkane and which retain carbonaceous residues with C/H compositions that are variable with the pressure of hydrogen. 3. Ethyne (Acetylene) [References to Part I1 (This Chapter)]

a. To make azimuthal VEEL measurements at low temperatures for different crystal planes, e.g., Ni(lll), to establish whether the impact selection rules can account for the type A and A‘ spectra measured in different laboratories (ref 4, Figs. 7 and 8) (Section 1I.B.l.a).

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b. To make measurements at low temperatures with finely divided catalysts, e.g., Ni/Al2O3and Pt/Al2O3,to obtain spectra following initial nondissociative adsorption and before polymerization (Sections II.B.3.a and II.B.4). c. To use RAIRS to investigate hydrogenation and polymerization on different crystal planes. 4. Higher Alkynes

a. To complete isotopic substitution experiments (H ---f D; 12C+ 13C) characterizing propyne and the linear butynes on (111) single-crystal planes, e.g., Ni, Pd, or Pt, for vibrational assignment purposes (Section II.C.l). b. To obtain spectra of nondissociatively adsorbed species on finely divided catalysts at low temperatures (Section II.C.2).

5. Acyclic Alkanes a. To use RAIRS to investigate the adsorption of methane, ethane, propane, and isobutane (2-methylpropane) on nonflat single-crystal surfaces, e.g., fcc (110), over a range of temperatures to establish when dissociative chemisorption sets in, such as has been observed for ethane and propane at room temperature on finely divided Pt (Section III.D.2). b. To make more systematic low-temperature studies of alkanes on finely divided catalysts.

6. Species Derived from the Dissociative Adsorption of Halogen- or Nitrogen-Substituted A lkanes or Alkenes a. To characterize more of the Cz and C3 species by RAIRS in conjunction with diffraction method.

7. Cycloalkanes, Cycloalkenes, and Aromatic Hydrocarbons a. To characterize, by the higher resolution of RAIRS, the intermediates involved in the interconversions of cyclohexane and benzene already detected by VEELS (Sections V.A.l and VI.A.l). b. To investigate RAIR spectra, down to 650 cm-', to obtain the v4 yCH band of benzene adsorbed on Ni(ll1) and Pd(ll1) so as to resolve the uncertainties in interpreting the VEEL spectra (Section VIII.A.l). c. To investigate the hydrogenation of benzene by RAIRS, e.g., on Pt(lll), to find which surfaces give complete hydrogenation to cyclohexane and which retain hydrogenated species. d. To make LEED I-V and perhaps PED measurements to determine whether agostically bonded cyclohexane is centered over metal atoms or hollows on (111) surfaces (Section V.A.l).

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XII. Conclusions In this review we have attempted to give a comprehensive account of the literature up to about the end of 1995, starting with the first pioneering attempts to obtain infrared spectra from hydrocarbons adsorbed on oxide-supported metal catalysts by Pliskin and Eischens in 1956. This is, of course, in reality an unattainable objective as the literature has greatly proliferated over the past 15 years and there are undoubtedly relevant articles in conference volumes, and possibly a few in the journal literature, that we have failed to locate. A motivation for going back as far as 1956 was that many of the earlier spectra were incorrectly interpreted. This situation became much improved after the advent of VEELS in the early 1970s, which enabled the study of vibrational spectra of adsorbates on flat single-crystal surfaces of known metal-atom arrangements as a function of temperature. For most chemists, a long-term motivation for work in this research field has always been the goal to understand the subject of heterogeneous catalysis, which is of fundamental importance and finds major use in industry for the efficient and selective attainment of desired products. In this respect, it is the understanding of the spectra obtained for adsorbates on finely divided metal catalysts that is of particular importance. For this reason, the figures in the two parts of this review are confined to such cases, including selected examples of SER spectra of species on rough surfaces of the coinage metals. However, the interpretation of these spectra depends on the understanding of the well-defined patterns of spectra obtained for species on single-crystal surfaces, of which the VEEL spectra were illustrated and discussed in the earlier review by one of us ( 4 ) . The present review updates the coverage of the single-crystal literature. It can reasonably be claimed that during the past few decades the several vibrational spectroscopies have played a major role in revolutionizing our understanding of the chemisorption of hydrocarbons on metal surfaces. Many insights have thereby been gained into related catalytic reactions. This is the main theme of our review. Increasingly, attention is now being directed to similar kinetic and mechanistic studies of surface-directed reactions. We have confined ourselves to providing a foretaste of such work in Section X by discussing some results for the specific case of ethene-based reactions, including hydrogenation, on platinum catalysts. The interpretation of the spectra of surface-adsorbed species, on singlecrystal surfaces in particular, is helped by complementary evidence derived from diffraction methods (LEED, PED) and from other nonvibrational spectroscopies (UPES, XPES, NEXAFS, SIMS, etc.). In particular, temperature-programmed desorption (TPD) is often measured in parallel with

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VEEL or RAIR spectroscopy to provide insight into temperature-dependent surface reactions. The TPD technique also provides information about the overall C/H composition of the hydrocarbon layers at different temperatures. Whereas we have made reference to, and taken into account, experimental results obtained from the use of such nonvibrational techniques in many cases, particularly when considering spectral interpretaions, it has not been feasible for us to systematically cover such papers that do not also include vibrational spectroscopic work. ACKNOWLEDGMENTS The authors are very grateful to the following persons who have given us permission to reproduce spectra or figures from their publications or theses, redrawn to a uniform format for comparison purposes: Prof. G. Blyholder, Dr. M. Clark, Dr. V. Y. Davydov, Prof. H. Dunken, Prof. R. P. Eischens, Prof. J. G. Ekerdt, (the late) Dr. J. Erkelens, Mr. T. Grimm 111, Prof. D. M. Haaland, (the late) Prof. R. M. Hexter, Dr. L. H. Little, Prof. A. Otto, Dr. H. A. Pearce, Dr. J. D. Prentice, Dr. G. Shahid, Prof. D. M. Shopov, Prof. T. Sziligyi, Prof. T. Takenaka, Dr. J. W. Ward, Prof. M. Weaver, and Prof. J. T. Yates, Jr. We also thank the following for permission to republish diagrams from their publications: Prof. J. C. Hemminger, Prof. B. E. Keol, Prof. W. Krasser, Dr. Y. Soma, and Prof. M. Trenary. We are also grateful to Prof. M. A. Chesters, Prof. J. G. Ekerdt, Dr. V. H. Grassian, Prof. W. Krasser, Dr. E. M. McCash, Dr. M. McCoustra, Dr. G. S. McDougall, Dr. P. Pudney, Dr. R. Raval, and Prof. M. Trenary for fruitful discussions and/or for giving us access to spectral data before publication. We are also indebted to the following copyright holders of publications for permission to reprint the spectra illustrated in this review from references cited in the figure captions: Academic Press, the American Chemical Society, the Deutscher Verlag fur Grundstoffeindustrie GmbH, Elsevier Science, the Institute of Petroleum in London, the International Union of Pure and Applied Chemistry, the Royal Society in London, and Springer-Verlag. One of us (N.S.) thanks the U.K. Science and Engineering Council for a series of research grants that supported the work of his laboratories in this area at the Universities of Cambridge and of East Anglia, Nonvich. We also thank the Royal Society, London, and the Consejo Nacional de Ciencia y Technologia of Venezuela for supporting two transatlantic exchange visits which assisted greatly in the writing of this review. C. De La C. also thanks La Universidad del Zulia for a study leave for this purpose. REFERENCES 1. Sheppard, N., and De La Cruz, C., Adv. Caral. 41, 1 (1996). 2. Pliskin, W. A,, and Eischens, R. P., J. Chem. Phys. 24,482 (1956). 3. Eischens, R. P., and Pliskin, W. A,, Adv. Catul. 10, 1 (1958). 4. Sheppard, N., Annu. Rev. Phys. Chem. 39,589 (1988). 5. Demuth, J. E., and Ibach, H., Surf Sci. 78,L238 (1978). 6. Demuth, J. E., and Ibach, H., Surf Sci. 85,365 (1979). 7. Lehwald, S., and Ibach, H., Surf: Sci. 89, 425 (1979). 8. Ibach, H., and Lehwald, S., J. Vac. Sci. Technol. 18, 625 (1981). 9. Bertolini, J. C., Massardier, J., and Dalmai-Imelik, G., J. Chem. SOC.,Faruduy Trans. 1 74, 1720 (1978). 10. Bertolini, J. C., Massardier, J., and Dalmai-Imelik, G., C. R. Hebd. Seances Acad. Sci. 285,515 (1977).

302 11. 12. 13. 14.

15. 16. 17. 18. 19. 20. 21. 22. 23.

24. 25. 26. 27. 28. 29. 30. 31. 32. 33. 34. 35. 36. 37. 38. 39. 40. 41. 42. 43. 44. 45. 46. 47. 48. 49. 50. 51. 52. 53.

54. 55.

NORMAN SHEPPARD AND CARLOS DE LA CRUZ Bertolini, J. C., and Rousseau, J., Surj Sci. 83, 531 (1979). Hammer, L., Ph.D. Thesis, University of Erlangen-Nurnberg, Germany (1985). Hammer, L., Hertlein, T., and Muller, K., Surf: Sci. 178, 693 (1986). Dotsch, B., Hammer, L., and Muller, K., Springer Ser. Surj Sci. 18, p. 212 (1989). Rudolf, P., Astaldi, G., Modesti, S . , and Rosei, R., Surj Sci. 220, L714 (1989). Dinardo, N. J., Demuth, J. E., and Avouris, Ph., J. Vuc. Sci. Technol., A [2] 1, 1244 (1983). Dinardo, N. J., Demuth, J. E., and Avouris, Ph., Phys. Rev. B 27, 5832 (1983). Zaera, F., and Hall, R. B., J. Phys. Chern. 91, 4318 (1987). Stroscio, J. A,, Bare, S. R., and Ho, W., Surj Sci. 148, 499 (1984). Bandy, B. J., Chesters, M. A,, Pemble, M. E., McDougall, G. S . , and Sheppard, N., SurJ Sci. 139, 87 (1984). Gates, J. A,, and Kesmodel, L. L., J. Chem. Phys. 76,4281 (1982). Gates, J. A,, and Kesmodel, L. L., Surj Sci. 124, 68 (1983). Kesmodel, L. L., J. Chern. Phys. 79,4646 (1983). Kesmodel, L. L., Waddill, G. D., and Gates, J. A,, Surj Set. 138, 464 (1984). Marchon, B., Surf: Sci. 162, 382 (1985). Timbell, P. J., Gellman, A. J., Lambert, R. M., and Willis, R. F., Surg. Sci. 206,339 (1988). Nascente, P. A. P., Van Hove, M. A,, and Somorjai, G. A., J. Vac. Sci. TechoL, A [2] 10,2342 (1992). Tardy, B., and Bertolini, J. C., J. Chim. Phys. 82, 407 (1985). Heitzinger, J. M., Gebhard, S . C., and Koel, B. E., J. Phys. Chem. 97, 5327 (1993). Yoshinobu, J., Sekitani, T., Onchi, T., and Nishijima, N. I.,J. Phys. Chern. 94,4269 (1990). Takaoka, T., Sekitani, T., Aruga, T., and Nishijima, M., Surf: Sci. 306, 179 (1994). Ibach, H., Hopster, H., and Sexton, B., Appl. Surf: Sci. 1, 1 (1978). Ibach, H., Hopster, H., and Sexton, B., Appl. Phys. 14, 21 (1977). Ibach, H., and Lehwald, J., J. Vac. Sci. Technol. 15,407 (1978). Avery, N. R., Lungmuir 4, 445 (1988). Hatzikos, G. H., and Masel, R. I., Surf: Sci. 185, 479 (1987). Mehandru, S. P., and Anderson, A. B., Appl. Surf: Sci. 19, 116 (1984). Dubois, L. H., Castner, D. G., and Somorjai, G. A,, J. Chern. Phys. 72, 5234 (1980). Mate, C. M., Kao, C. T., Bent, B. E., and Somorjai, G. A., Surf: Sci. 197, 183 (1988). Parmeter, J. E., Hills, M. M., and Weinberg, W. H.,J. A m . Chern. Soc. 108,3563 (1986). Jakob, P., Cassuto, A,, and Menzel, D., Surf: Sci. 187, 407 (1987). Sasaki, T., Kawada, F., Aruga, T., and Iwasawa, Y., Surt Sci. 278, 291 (1992). Bertolini, J. C., Tardy, B., and Ducros, R., C. R. Hebd. Seances Acad. Sci. 298,107 (1984). Marinova, T. S . , and Kostov, K. L., Surf: Sci. 181, 573 (1987). Kostov, K. L., and Marinova, T. S . , Surf: Sci. 184, 359 (1987). Steip, U., Tsai, M.-C., Kuppers, J., and Ertl, G., Surf: Sci. 147, 65 (1984). Erley, W., Baro, A. M., and Ibach, H., Surf: Sci. 120, 273 (1982). Hung, W. H., and Bernasek, S . L., Surf: Sci. 339, 272 (1995). Backx, C., Fuerbacher, B. F., Fitton, B., and Willis, R. F., Surf: Sci. 63, 193 (1977). Hamilton, J. C., Swanson, N., Waelawski, B. J., and Cellotta, R. J., J. Chern. Phys. 74, 4156 (1981). Backx, C., Willis, R. F., Fuerbacher, B., and Fitton, B., Surf: Sci. 68, 516 (1977). Backx, C., and Willis, R. F., Chem. Phys. Lett. 53, 471 (1978). Chesters, M. A,, and McCash, E. M., J. Electron Spectrosc. Relat. Phenom. 44, 99 (1987). Marinova, T. S., and Stefanov, P. K., SurJ Sci. 191, 66 (1987). Marinova, T. S . , and Stefanov, P. K., Surf: Sci. 219, 490 (1989).

VIBRATIONAL SPECTRA OF HYDROCARBONS

303

Avery, N. R., J. Am. Chem. Soc. 107, 6711 (1985). Stuve, E. M., Madix, R. J., and Sexton, B. A., Surf: Sci. 123,491 (1982). Madix, R. J., Appl. Surf: Sci. 14, 41 (1982-1983). Bao, S., Hofmann, Ph., Schindler, K.-M., Fritzsche, V., Bradshaw, A. M., Woodruff, D. P., Casado, C., and Asensio, M. C., Surj Sci. 307-309,722 (1994). 60. Bao, S., Schindler, K.-M., Hofmann, Ph., Fritzsche, V., Bradshaw, A. M., and Woodruff, D. P., Surf: Sci. 291, 295 (1993). 61. Anson, C. E., Keiller, B. T., Oxton, I. A., Powell, D. B., and Sheppard, N., J. Chem. SOC.,Chem. Commun., 470 (1983). 62. Gentle, T. M., and Muetterties, E. L., J. Phys. Chem. 87, 2469 (1983). 63. Arvanitis, D., Baberschke, K., Wenzel, L., and Dobler, U., Phys. Rev. Lett. 57, 3175 (1986). 64. Weinelt, M., Huber, W., Zebisch, P., Steinruck, H. P., Ulbricht, P., Birkenheuer, U., Boettger, J. C., and Rosch, N., J. Chem. Phys. 102, 9709 (1995). 65. Ormerod, R. M., Lambert, R. M., Hoffmann, H., Zaera, F., Wang, L. P., Bennett, D. W., and Tysoe, W. T., J. Phys. Chem. 98, 2134 (1994). 66. Land, T. H., Michely, T., Behm, R. J., Hemminger, J. C., and Comsa, G., J. Chem. Phys. 97,6774 (1992). 67. Evans, J., and McNulty, G. S., J. Chem. Soc., Dalton Trans., 79 (1984). 68. Little, L. H., Sheppard, N., and Yates, D. J. C., Proc. R. Soc. London Ser., A 259, 242 (1960). 69. Lapinski, M. P., and Ekerdt, J. G., J. Phys. Chem. 94, 4599 (1990). 70. Sheppard, N., and Ward, J. W., J. Catal. 15, 50 (1969). 71. Erkelens, J., and Wosten, W. J., J. Catal. 54, 143 (1978). 72. Krasser, W., Fadini, A,, and Renouprez, A., J. Mol. Strucr. 60, 427 (1980). 73. Nash, G. P., and De Sieno, R. P., J. Phys. Chem. 69, 2139 (1965). 74. Ito, M., Mori, Y., Kato, T., and Suetaka, W., Appl. Surf: Sci. 2, 543 (1979). 75. Demuth, J. E., Ibach, H., and Lehwald, S., Phys. Rev. Lett. 40, 1044 (1978). 76. Bobrov, A. V., Kimelfeld, Ja. M., and Mostovaya, L. M., J. Mol. Struct. 60, 431 (1980). 77. Clark, M., Ph.D. Thesis, University of Cambridge (1960). 78. Dunken, H., Schmidt, K., and Hobert, H., Z. Chem. 4,312 (1964). 79. Beebe, T. P., Jr., Albert, M. R., and Yates, J. T., Jr., J. Catal. 96, 1 (1985). 80. Sokolova, N. P., Kazakova, L. A., and Borisenko, L. A., Zh. Fir. Khim. 44,2657 (Eng. ed., p. 1515) (1970). 81. Bayman, A,, Hansma, P. K., Kaska, W. C., and Dubois, L. H., Appl. Surf: Sci. 14, 194 (1982). 82. Parker, W. L., Seidle, A. R., and Hexter, R. M., Langmuir 4, 999 (1988). 83. Patterson, M. L., and Weaver, M. J., J. Phys. Chem. 89, 5046 (1985). 84. Shirakawa, H., Ito, T., and Ikeda, S., Polym. J. 4, 460 (1973). 85. Rives-Arnau, V., and Sheppard, N., J. Chem. Soc., Faraday Trans I 76, 394 (1980). 86. Ward, J. W., Ph.D. Thesis, University of Cambridge (1962). 87. Prentice, J.D., Ph.D. Thesis, University of East Anglia (1977). 88. Randhava, S. S., and Rehmet, A,, Trans. Faraday Soc. 66,235 (1970). 89. Szilagyi, T., Vib. Spectrosc. 2, 29 (1991). 90. Kavtaradze, N. N., and Sokolova, N. P., Zh. Fiz. Khim. 44, 171 (1970). 91. Pearce, H. A,, Ph.D. Thesis, University of East Anglia (1974). 92. Sheppard, N., Chenery, D. H., Lesiunas, A., Prentice, J. D., Primet, M., and Pearce, H. A., in “Molecular Spectroscopy of Condensed Phases” (M. Grosmann, S. G. Elkornoss, and J. Ringeissen, eds.), p. 345. Elsevier, Amsterdam, 1976. 93. Shahid, G., and Sheppard, N., J. Chem. Soc., Faraday Trans. 90,507 (1994). 56. 57. 58. 59.

304

NORMAN SHEPPARD AND CARLOS D E LA CRUZ

Parker, W. L., Siedle, A. R., and Hexter, R. M., J. Am. Chem. SOC. 107,264 (1985). Fielchenfeld, H., Luckier, M., Efron, L., and Willner, B., Surf: Sci. 268, 127 (1992). Blyholder, G., and Wyatt, W. V . ,J. Phys. Chem. 78, 618 (1974). Pockrand, I., Pettenkofer, C., and Otto, A,, J. Electron Spectrosc. Relat. Phenom. 29, 409 (1983). 98. Pettenkofer, C., Mrozek, I., Bornemann, T., and Otto, A., Surf: Sci. 188, 519 (1987). 99. Mrozek, J., and Otto, A., J . Electron Spectrosc. Relat. Phenom. 54/55, 895 (1990). 100. Manzel, K., Schulze, W., and Moskovits, M., Chem. Phys. Lett. 85, 183 (1982). 101. Fielchenfeld, H., and Weaver, M. J., J. Phys. Chem. 93, 4276 (1989). 102. Avery, N. R., and Sheppard, N., Proc. R. SOC.London, Ser. A 405, 27 (1986). 103. Pudney, P., Ph.D. Thesis, University of East Anglia (1989). 104. Fricke, A,, Graupner, H., Hammer, L., and Miiller, K., Surf: Sci. 272, 182 (1992). 105. Chesters, M. A., and McCash, E. M., J. Electron Spectrosc. Relat. Phenom. 44, 99 (1987). 106. Chesters, M. A., J. Mol. Sfruct. 173, 405 (1988). I06a. Roberts, A. J., Haq, S., and Raval, R., J. Chem. SOC.,Faraday Trans. 92, 4823 (1996). 107. Bateman, J. E., and Chesters, M. A,, personal communication, 1996. 108. Bent, B. E., Mate, C. M., Crowell, J. E., Koel, B. E., and Somorjai, G . A,, J. Phys. Chem. 91, 1493 (1987). 109. Raval, R., and Cooper, E., personal communication, 1996. 110. Schrader, B., “Ramanhfrared Atlas of Organic Compounds,” 2nd ed., Spectrum C1-14. VCH Verlagsges., Weinhein, 1989. 111. Sheppard, N., Pure Appl. Chem. 4,71 (1962). 112. Morrow, B. A., and Sheppard, N., Proc. R. SOC.London, Ser. A 311, 415 (1969). 113. Jackson, S. D., and Casey, N. J., J. Chem. Soc., Faraday Trans. 91,3269 (1995). 114. Abrantes, L. M., Fleischmann, M., Hill, I. R., Peter, L. M., Mengoli, M., and Zotti, G., J. Electroanal. Chem. 164, 177 (1984). 114a. Shorthouse, L. J., Haq, S., and Raval, R., Surf: Sci. 368, 296 (1996). 115. Ceyer, S. T., Annu. Rev. Phys. Chem. 39,479 (1988). 116. Beckerle, J. D., Yang, Q. Y., Johnson, A. D., and Ceyer, S . T., J. Chem. Phys. 86, 7236 (1987). 117. Lee, M. B., Yang, Q. Y., Tang, S. L., and Ceyer, S. T., J. Chem. Phys. 85, 1693 (1986). 118. Lee, M. B., Yang, Q. Y., and Ceyer, S . T., J. Chem. Phys. 87, 2724 (1987). 119. Yang, Q. Y., Maynard, K. J., Johnson, A. D., and Ceyer, S. T., J. Chem. Phys. 102, 7734 (1995). 120. Ceyer, S. T., Beckerle, J. D., Lee, M. B., Tang, S . L., Yang, Q. Y., and Hines. M. A., J. Vac. Sci. Technol., A [2] 5, 501 (1987). 121. Yang, Q. Y., Johnson, A. D., Maynard, K. J., and Ceyer, S. T., J. Am. Chem. SOC. 111, 8748 (1989). 122. Oakes, D. J., McCoustra, M. R. S., and Chesters, M. A,, Faraday Discuss. 96,325 (1993). 122a. Yoshinobu, J., Ogasawara, H., and Kawai, M., Phys. Rev. Lett. 75, 2176 (1995). 123. Camplin, J. R., Cook, J. C., and McCash, E. M., J. Chem. SOC., Faraday Trans. 91, 3563 (1995). 124. Wu, M.-C., and Goodman, D. W., J. Am. Chem. SOC.116, 1364 (1994). 125. Kuijpers, E. G. M., Breedjik, A. K., Van der Wal, V. J. J., and Gem, J. W., J . Catal. 72, 210 (1981). 126. Marzouk, H. A., Bradley, E. B., and Arunkumar, K. A., Spectrosc. Lett. 18, 189 (1985). 127. Horn, K., and Pritchard, J., Surf: Sci. 52, 437 (1975). 128. Chesters, M. A,, in “Analytical Applications of Spectroscopy” (C. S. Creaser and A. M. C. Davies, eds.), p. 201. Royal SOC.Chem., London, 1988. 94. 95. 96. 97.

VIBRATIONAL SPECTRA O F HYDROCARBONS

305

129. Horn, K., in “Proceedings of Vibrations in Adsorbed Layers Conference” (H. Ibach, ed.) p. 140. Kernforschungsanlage, Julich GmbH, Germany, 1978. 130. Chesters, M. A,, Gardner, P., and McCash, E. M., Surf: Sci. 209, 89 (1989). 131. McDougall, G. S., Ph.D. Thesis, University of East Anglia (1984). 132. De La Cruz, C., and Sheppard, N., J. Cutul. 127,445 (1991). 133. Trautmann, S., Baerns, M., and Auroux, A., J. Cutul. 136, 613 (1992). 134. Alimardonov, E. A., Gass, A. M., Kapusta, 0. I., and Klimin, S. A,, Sov. Phys.-JETP Lett. 41, 424 (1985). 135. Brings, R., Mrozek, I., and Otto, A,, J. Rumun Spectrosc. 22, 119 (1991). 136. Otto, A., Mrozek, I., Grabhorn, H., and Akemann, W., J. Phys. Condens. Mutter 4, 1143 (1992). 137. Akemann, W., and Otto, A., Lungmuir 11, 1196 (1995). 138. Chesters, M. A., and Gardner, P., Spectrochim. Actu, Part A 46,1011 (1990). 138u. Pawela-Crew, J., and Madix, R. J., Surf: Sci. 337, L800 (1995). 139. Witte, G., and Wo11, Ch., J. Chem. Phys. 103, 5860 (1995). 140. Shahid, G., and Sheppard, N., Spectrochim. Actu, Part A 46,999 (1990). 141. Brogan, M. S., Cairns, J. A., Dines, T. J., and Rochester, C. H., J. Chem. Soc., Faruduy Trans. 91, 733 (1995). 142. Zhou, X . L., Zhu, X.-Y., and White, J. M., Acc. Chem. Res. 23, 327 (1990). 143. Zhou, X. L., Zhu, X.-Y., and White, J. M., Sur$ Sci. Rep. 13, 73 (1991). 144. Zaera, F., Acc. Chem. Res. 25, 260 (1992). 145. Henderson, M. A., Mitchell, G. E., and White, J. M., Surf: Sci. 184, L325 (1987). 146. Henderson, M. A., Mitchell, G. E., and White, J. M., Surf Sci. 248, 279 (1991). 147. Fan, J., and Trenary, M., Lungmuir 10,3649 (1994). 148. Zaera, F., and Hoffmann, H., J. Phys. Chem. 95, 6297 (1991). 149. Zaera, F., Lungmuir 7, 1998 (1991). 150. Radhakrishnan, G., Stenzel, W., Hemmen, R., Conrad, H., and Bradshaw, A. M., J. Chem. Phys. 95,3930 (1991). 151. Zhou, Y., Feng, W. M., Henderson, M. A,, Roop, B., and White, J. M., J. Am. Chem. Soc. 110,4447 (1988). 152. Roop, B., Lloyd, K. G . , Costello, S. A,, Campion, A,, and White, J. M., J. Chem. Phys. 91,513 (1989). 153. Pansoy-Hjelvik, M. E., Xu, R., Gao, Q., Weller, K., Feher, F., and Hemminger, J. C., Surf: Sci. 312, 97 (1994). 154. Fairbrother, D. H., Peng, X. D., Viswanathan, R., Stair, P. C., Trenary, M., and Fan, J., Surt Sci. 285, LA55 (1993). 155. Fairbrother, D. H., Peng, X. D., Trenary, M., and Stair, P. C., J. Chem. SOC.,Furuday Trans. 91, 3619 (1995). 156. Oakes, D. J., Wender, J. C., Chesters, M. A,, and McCoustra, M. R. S., J. Electron Spectrosc. Relat. Phenom. 64165,477 (1993). 157. Zhou, Y., Henderson, M. A., Feng, W. M., and White, J. M., Surf Sci. 224, 386 (1989). 158. Lin, J.-L., and Bent, B. E.. J. Vuc. Sci. Technol., A [2] 10, 2202 (1992). 159. Lin, J.-L., and Bent, B. E., Chem. Phys. Lett. 194, 208 (1992). 160. Chiang, C.-M., and Bent, B. E., Surf Sci. 279, 79 (1992). 161. Lamont, C. L. A., Conrad, H., and Bradshaw, A. M., Surf: Sci. 280,79 (1993). 162. Lin, J.-L., Chiang, C.-M., Jenks, C. J., Yang, M. X., Wentzlaff, T. H., and Bent, B. E., J. Cutul. 147, 250 (1994). 163. Nakamoto, K., “Infrared and Raman Spectra of Inorganic and Coordination Compounds,’’ 4th ed. Wiley-Interscience, New York, 1986. 164. Schule, J., Siegbahn, P., and Wahlgren, U., J. Chem. Phys. 89, 6982 (1988).

306

NORMAN SHEPPARD AND CARLOS D E LA CRUZ

165. Siegbahn, P. E., and Panas, I., Surf: Sci. 240, 37 (1990). 166. Yang, H., and Whitten, J. L., J. Am. Chem. Soc. 113, 6442 (1991); Surf. Sci. 255, 193

(1991). Zheng, C., Apeloig, Y., and Hoffman, R., J. Am. Chem. SOC.110, 749 (1988). Krohmer, P., and Goubeau, J., 2. Anorg. Allg. Chem. 369, 238 (1969). Kress, J., and Novak, A,, J. Organomet. Chem. 121, 7 (1976). Gay, A. P., Can. J. Chem. 41, 1511 (1963). Peng, X. D., Vismanathan, R., Smudde, G. H., Jr., and Stair, P. C., Rev. Sci. Insfrum. 63, 3930 (1992). 172. Rasko, J., Bontovics, J., and Solymosi, F., J. Caral. 143, 138 (1993). 173. McBreen, P. H., Erley, W., and Ibach, H., Surf: Sci. 148, 292 (1984). 174. Henderson, M. A,, Radloff, P. L., White, J. M., and Mims, C. A., J. Phys. Chem. 92, 4111 (1988). 275. Henderson, M. A,, Radloff, P. L., Greenlief, C. M., White, J. M., and Mims. C. A., J. Vac. Sci. Technol., A [2] 6, 769 (1988). 176. Mitchell, G. E., Radloff, P. L., Greenlief, C. M., Henderson, M. A., and White, J. M., Surf: Sci. 183, 403 (1987). 177. George, P. M., Avery, N. R., Weinberg, W. H., and Tebbe, F. N., J . Am. Chem. Soc. 105, 1393 (1983). 178. Hills, M. M., Parmenter, J. E., Mullins, C. B., and Weinberg, W.H., J. Am. Chem. SOC. 108, 3554 (1986). 179. Monim, S. S., Venus, D., Roy, D., and McBreen, P. H., J. Am. Chem. Soc. 111,4106(1989). 180. Zhou, X.-L., Liu, 2.-M., Kiss, J., Sloan, D. W., and White, J. M., J. Am. Chem. Soc. 117, 3565 (1995). 181. Solymosi, F., and Klivtnyi, G., J. Phys. Chem. 99,8950 (1995); Surf: Sci. 342,168 (1995). 182. Weldon, M. K., and Friend, C. M., Surf: Sci. 321, L202 (1994). 183. Erley, W., McBreen, P. H., and Ibach, H., J. Catal. 84, 229 (1983). 184. Oxton, I. A., Powell, D. B., Sheppard, N., Burgess, K., Johnson, B. F. G., and Lewis, J., J. Chem. Soc., Chem. Commun. 719 (1982). 185. Holgrem, J. S., and Shapley, J. R., Oganometailics 4, 793 (1985). 186. Chang, S. C., Kafafi, Z . H., Hauge, R. H., Whitmire, K. H., Billups, W. E., and Margrave, J. L., lnorg. Chem. 25, 4530 (1986). 187. Chang, S. C., Kafafi, Z. H., Hange, R. H., Billups, W. E., and Margrave, J. L., J . Am. Chem. Soc. 107, 1447 (1985). 188. Solymosi, F., and Rasko, J., J. Caral. 155, 74 (1995). 189. Howard, M. W., Kettle, S. F. A., Oxton, I. A., Powell, D. B., Sheppard, N., and Skinner, P., J. Chem. Soc., Faraday Trans. 2 77, 397 (1981). 190. Eady, C. R., Johnson, B. F. G., and Lewis, J., J. Chem. SOC.,Dalton Trans. 477 (1977). 191. Demuth, J. E., and Ibach, H., Surf: Sci. 77, L238 (1978). 192. Lloyd, K. B., Roop, B., Campion, A,, and White, J. M., Surf: Sci. 214, 227 (1989). 193. Lin, J.-L., Teplyakov, A. V., and Bent, B. E., J. Phys. Chem. 100, 10721 (1996). 194. Hoffmann, H., Griffiths, P. R., and Zaera, F., Surf. Sci. 262, 141 (1992); Zaera, F., Hoffmann, H., and Griffiths, P. R., Vacuum 41,735 (1990). 195. McGee, K. C., Driessen, M. D., and Grassian, V. H., J. Catal. 157, 730 (1995). 196. Driessen, M. D., and Grassian, V. H., unpublished work (personal communication). 197. De La Cruz, C., and Sheppard, N., J. Mol. Struct. 247, 2.5 (1991). 198. Janssens, T. V. W., and Zaera, F.. J. Phys. Chem. 100, 14118 (1996). 19Y. Bent, B. E., Nuzzo, R. G., Zegarski, B. R., and Dubois, L. H., J. Am. Chem. SOC. 113, 1137 (1991). 200. Lin, J.-L., and Bent, B. E., J. Phys. Chem. 96, 8529 (1992).

167. 168. 169. 170. 171.

VIBRATIONAL SPECTRA O F HYDROCARBONS

307

201. Forbes, J. G., and Gellman, A. J., J. Am. Chem. SOC. 115, 6277 (1993). 202. Jenks, C. J., Bent, B. E., Bernstein, N., and Zaera, F., J. Am. Chem. Soc. 115,308 (1993). 202a. Cremer, P. S., Su, X., Shen, Y. R., and Somorjai, G. A., J. Phys. Chem. 100, 16302 (1996). 203. Zegarski, B. R., and Dubois, L. H., Surf: Sci. 262, L129 (1992). 204. Bent, B. E., Nuzzo, R. G., and Dubois, L. H., J. Am. Chem. Soc. 111, 1634 (1989). 205. Bent, B. E., Nuzzo, R. G., Zegarski, B. R., and Dubois, L. H., J. Am. Chem. SOC. 113, 1143 (1991). 206. Carter, R. N., Anton, A. B., and Apai, G., J. Am. Chem. Soc. 114,4410 (1992). 207. Carter, R. N., Anton, A. B., and Apai, G., Surf: Sci. 290, 319 (1993). 208. Liu, 2.-M., Zhou, X.-L., Buchanan, D. A,, Kiss, J., and White, J. M., J. Am. Chem. SOC. 114, 2031 (1997). 209. Zaera, F., and Bernstein, N., J. Am. Chem. Soc. 116,4881 (1994). 210. Grassian, V. H., and Pimentel, G. C., J. Chem. Phys. 88, 4478, 4484 (1988). 211. Song, Y., Gardner, P., Conrad, H., Bradshaw, A. M., and White, J. W., Surf: Sci. 248, L279 (1991). 212. Xi, M., and Bent, B. E., Surf: Sci. 278,19 (1992). 213. Martel, R., Rochefort, A., and McBreen, P. H., J. Am. Chem. Soc. 116,5965 (1994). 214. Xu, C., and Koel, B. E., Surf: Sci. 292, L803 (1993). 215. Chesters, M. A,, Parker, S. F., and Raval, R., J. Electron Spectrosc. Relat. Phenom. 39, 155 (1986). 216. Raval, R., and Chesters, M. A., Surt Sci. 219, L505 (1989). 217. Raval, R., Parker, S. F., and Chesters, M. A,, Surf: Sci. 289, 227 (1993). 218. Raval, R., Surf: Sci. 331-333, 1 (1995). 219. Wadill, G. D., and Kesmodel, L. L., Chem. Phys. Lett. 128, 208 (1986). 220. Raval, R., Pemble, M. E., and Chesters, M. A., Surf: Sci. 210, 187 (1989). 221. Bussell, M. E., Henn, F. C., and Campbell, C. T., J. Phys. Chem. 96, 5978 (1992). 222. Land, D. P., Erley, W., and Ibach, H., Surf: Sci. 289, 237 (1993). 223. Xu, C., Tsai, Y.-L., and Koel, B. E., J. Phys. Chem. 98, 585 (1994). 224. Lamont, C. L. A,, Borbach, M., Stenzel, W., Conrad, H., and Bradshaw, A. M., Chem. Phys. Lett. 230, 265 (1994). 225. Hoffmann, F. M., Felter, T. E., Thiel, P. A., and Weinberg, W. H., Surf: Sci. 130,173 (1983). 226. Hoffmann, F. M., and Upton, T. H., J. Phys. Chem. 88, 6209 (1984). 227. Chesters, M. A,, Parker, S. F., and Raval, R., Surf: Sci. 165, 179 (1986). 228. Cooper, E., Coats, A. M., and Raval, R., J. Chem. SOC.,Faraday Trans. 91,3703 (1995). 229. Chesters, M. A,, in “Analytical Applications of Spectroscopy” (C. S. Creaser and A. M. C. Davies, eds.), p. 201. Royal Society of Chemistry, London, 1988. 230. Martin, R., Gardner, P., Tushaus, M., Bonev, Ch., Bradshaw, A. M., and Jones, T. S., J. Electron Spectrosc. Relat. Phenom. 54/55, 773 (1990). 230a.Weldon, M. K., Uvdal, P., Friend, C. M., and Wiegand, B. C., Surf: SCL 355, 71 (1996). 231. Brookhart, M., and Green, M. L. H., J. Organomet. Chem. 250, 395 (1987). 232. Snyder, R. G., and Schachtschneider, J. H., Spectrochim. Acta Part A 21, 169 (1965). 233. Firment, L. E., and Somorjai, G. A,, J. Chem. Phys. 66,2901 (1977). 234. Huber, W., Zebisch, P., Bornemann, T., and Steinriich, H.-P., Surf: Sci. 239,353 (1990). 235. Madey, T. E., and Yates, J. T., Jr., Surf: Sci. 76, 397 (1978). 236. Kang, D. B., and Anderson, A. B., J. Am. Chem. Soc. 107,7858 (1985). 237. Saillard, J.-Y., and Hoffmann, R., J. Am. Chem. SOC. 106, 2006 (1984). 238. Shustorovich, E., Baetzold, R. C., and Muetterties, E. L., J. Phys. Chem. 87,1100 (1983). 239. Henn, F. C., Diaz, A. L., Bussell, M. E., Hugenschmidt, M. B., Domagala, M. E., and Campbell, C. T., J. Phys. Chem. 96, 5965 (1992).

308

NORMAN SHEPPARD AND CARLOS D E LA CRUZ

240. Hugenschmidt, M. B., Diaz, A. L., and Campbell, C. T., J. Phys. Chem. 96, 5974 (1992). 241. Palazov, A. N., Andreev, A. A., and Shopov, D. M., C. R. Acud. Bulg. Sci. 18,1145 (1965). 242. Shopov, D. M., and Palazov, A. N., Kinet. Katul. 8, 862 (1967); Kiner. Cufal (English Transl.), p. 732 (1967). 243. Palazov, A. N., and Shopov, D. M., Comm. Dept. Chem., Bulg. Acad. Sci. 1,175 (1968). 244. Shopov, D. M., and Palazov, A. N., C. R. Acud. Sci. Bulg. 22, 181 (1969). 245. Palazov, A., J. Carul. 30, 13 (1973). 246. Baumarten, E., and Weinstrauch, F., Spectrochim. Actu, Part A 35, 1323 (1979). 247. Haaland, D. M., Surf: Sci.111,555 (1981). 248. Sheppard, N., Avery, N. R., Clark, M., Morrow, B. A,, Smart, R. St. C., Takenaka, T., and Ward, J. W., in “Molecular Spectroscopy” (P. Hepple, ed.), p. 97. Inst. Petrol., London, 1968. 249. Erkelens, J., and Eggink-Du Burck, S. H., J. Catal. 15, 62 (1967). 250. Primet, M., Basset, J. M., Mathieu, M.-V., and Prettre, M., J. Catal. 29, 213 (1973). 251. Basset, J. M., Dalmai-Imelik, G., Primet, M., and Mutin, R., J. C a f d . 37, 22 (1975). 252. Reiicha, R., Collect. Czech. Chem. Commun. 37, 1607 (1972). 253. Ballinger, T. H., and Yates, J. T., Jr., J. Am. Chem. Soc. 114, 10074 (1992). 254. Ballinger, T. H., and Yates, J. T., Jr., J. Phys. Chem. 96, 9979 (1992). 255. Wong, J. C. S., and Yates, J. T., Jr., J. Am. Chem. Soc. 116, 1610 (1994). 256. Hoffmann, F. M., O’Brien, E. V., Hrbek, J., and De Paola, R. A,, J. Elecron Spectrosc. Relut. Phenom. 29, 301 (1983). 257. Avery, N. R., Surf Sci. 163,357 (1985). 258. Avery, N. R., in “Surface Science: Principles and Applications” (R. F. Howe, R. N. Lamb, and K. Wandelt,eds.), Springer Proc. Phys., Vol. 73, p. 256. Springer-Verlag, Berlin and Heidelberg, 1993. 259. Miller, F. A,, and Inskeep, R. G., J. Chem. Phys. 18,1519 (1950). 260. Hostetler, M. J., Nuzzo, R. G., Girolami, G. S., and Dubois, L. H., J. Phys. Chem. 98, 2952 (1994). 262. Hoffmann, F. M., Felter, T. E., and Weinberg, W. H., J . Chem. Phys. 76, 3799 (1982). 262. Felter, T. E., Hoffmann, F. M., Thiel, P. A,, and Weinberg, W. H., Surf Sci. 130, 163 (1983). 263. Duncan, J. L., and McKean, D. C., J. Mol. Spectrosc. 27, 117 (1968). 264. Anson, C. E., Sheppard, N., Powell, D. B., Bender, B. R., and Norton, J. R.. J . Chem. Soc., Furuduy Trans. 90, 1449 (1994); Table 2. 265. Xu, C., Koel, B. E., Newton, M. A., Frei, N. A., and Campbell, C. T., J. Phys. Chem. 99, 16670 (1995). 266. Avery, N. R., J. Caral. 19, 15 (1970). 267. Roberts, J. D., and Madix, R. J., data listed in ref. 268. 268. Roberts, J. D., and Madix, R. J., Surf Sci. 226, L71 (1990). 269. Teplyakov, A. V., and Bent, B. E., J. Chem. SOC.,Faruday Trans. 91,3645 (1995). 270. Netto, N., Di Lauro, C., Castellucci, E., and Califano, S., Specrrochim. Acra. Part A 23, 1763 (1967). 271. Land, D. P., Pettiette-Hall, C. L., McIver, R. T., Jr., and Hemminger, J. C., J. Am. Chem. Soc. 111, 5970 (1989). 272. Pettiette-Hall, C. L., Land, D. P., McIver, R. T., Jr., and Hemminger, J. C., J. Am. Chem. Soc. 113, 2755 (1991). 273. Avery, N. R., SurJ Sci. 137, L109 (1984). 274. Avery, N. R., Surf- Sci. 146,363 (1984).

VIBRATIONAL SPECTRA OF HYDROCARBONS

309

275. Wertz, D. W., Bocian, D. F., and Hazouri, M. J., Spectrochim. Acta, PartA 29,1439 (1973). 276. Hostetler, M. J., Dubois, L. H., Nuzzo, R. G., and Girolami, G. S., J. Am. Chem. Soc. 115, 2044 (1993). 277. Lee, A. F., Baddeley, C. J., Hardacre, C., and Lambert, R. M., J. A m . Chem. Soc. 117, 7719 (1995). 278. Rosenberg, A,, Phys. Rev. B 51, 1961 (1995). 279. Bertolini, J. C., Dalmai-Imelik, G., and Rousseau, J., Surf Sci. 67, 478 (1977). 280. Bertolini, J. C., and Rousseau, J., Surf Sci. 89, 467 (1979). 281. Bertolini, J. C., Massardier, J., and Tardy, B., J. Chim. Phys. 78, 939 (1981). 282. Jobic, H., Tardy, B., and Bertolini, J. C., J. Electron Spectrosc. Relat. Phenom. 38, 55 (1986). 283. Jobic, H., and Renouprez, A., Surf Sci. 111, 53 (1981). 284. Lehwald, S., Ibach, H., and Demuth, J. E., Surf Sci. 78, 577 (1978). 285. Huntley, D. R., Jordan, S. L., and Grimm, F. A,, J. Phys. Chem. 96, 1409 (1992). 286. Waddill, G. D., and Kesmodel, L. L., Phys. Rev. B 31,4940 (1985). 287. Waddill, G. D., and Kesmodel, L. L., Phys. Rev. B 32, 2107 (1985). 288. Waddill, G. D., and Kesmodel, L. L., Chem. Phys. Lett. 128, 208 (1986). 289. Grassian, V. H., and Muetterties, E. L., J. Phys. Chem. 91, 389 (1987). 290. Aarts, J. F. M., and Sassen, N. R. M., Surf Sci. 214, 257 (1989). 291. Fujisawa, M., Sekitani, T., Morikawa, Y., and Nishijima, M., J. Phys. Chem. 95, 7415 (1991). 292. Nishijima, M., Fujisawa, M., Takoaka, T., and Sekitani, T., Surf Sci. 283, 121 (1993). 293. Fujisawa, M., Takaoka, T., Seikitani, T., and Nishijima, M., J. Phys. Chem. 96, 4602 (1992). 294. Abon, M., Bertolini, J. C., Billy, J., Massardier, J., and Tardy, B., Surf Sci. 162,395 (1985). 295. Mate, C. M., and Somorjai, G. A,, Surf Sci. 160, 542 (1985). 296. Massardier, J., Tardy, B., Abon, M., and Bertolini, J. C., Surf Sci. 126, 154 (1983). 297. CEmic, F., Dippel, O., and Hasselbrink, E., Surf: Sci. 342, 101 (1995). 297a. Dippel, O., C b i c , F., and Hasselbrink, E., Surf Sci. 357-358, 190 (1996). 2976. Schenk, A,, and Kuppers, J., Surf Sci. 352-354,133 (1996). 298. Surman, M., Bare, S. R., Hofmann, P., and King, D. A., Surf Sci. 126,349 (1983). 299. Bare, S. R., Hofmann, P., Nyberg, G. L., Surman, M., and King, D. A,, unpublished work cited in Duncan and McKean (263). 300. Koel, B. E., and Somorjai, G. A,, J. Electron Spectrosc. Relat. Phenom. 29, 287 (1983). 301. Koel, B. E., Crowell, J. E., Mate, C. M., and Somorjai, G. A,, J. Phys. Chem. 88, 1988 (1984). 302. Koel, B. E., Crowell, J. E., Bent, B. E., Mate, C. M., and Somorjai, G. A., J. Phys. Chem. 90, 2949 (1986). 303. Jakob, P., and Menzel, D., Surf: Sci. 201, 503 (1988). 304. Jakob, P., and Menzel, D., Surf Sci. 220, 70 (1989). 305. Jacob, P., and Menzel, D., Langmuir 7, 134 (1991). 306. Jacob, P., and Menzel, D., Surf Sci. 235, 15 (1990). 307. Graen, H. H., Neumann, M., Wambach, J., and Freund, H. J., Chem. Phys. Lett. 165, 137 (1990). 308. Tardy, B., Bertolini, J. C., and Ducros, R., Bull. Soc. Chim.Fr. 3, 313 (1985). 309. Avouris, Ph., and Demuth, J. E., J. Chem. Phys. 75, 4783 (1981). 310. Demuth, J. E., Avouris, Ph., and Schmeisser, D., J. Electron Sectrosc. Relat. Phenom. 29, 163 (1983). 311. Schweitzer, E., and Bradshaw, A. M.,J. Electron Spectrosc. Relat. Phenom. 38,141 (1986). 312. Haq, S., and King, D. A,, J. Phys. Chem. 100,16957 (1996).

310

NORMAN SHEPPARD AND CARLOS DE LA CRUZ

313. Hallmark, V. M., and Campion, A., Chem. Phys. Lett. 110,561 (1984); J. Chem. Phys. 84,2933 (1986). 314. Schaff, O., Fernandez, V., Hofmann, Ph., Schindler, K. M., Theobald, A., Fritzsche, V., Bradshaw, A. M., Davis, R., and Woodruff, D. P., Surf: Sci. 348,89 (1996). 315. Weiss, P. S., and Eigler, D. M., Phys. Rev. Lett. 71,3139 (1993); Sautet, P., and Bocquet, M.-L., Surf: Sci. 304, L445 (1994). 316. Stellwag, C., Held, G., and Menzel, D., Surf Sci. 325, L379 (1995). 317. Van Hove, M. A., Lin, R. F., and Somorjai, G. A., J. Am. Chem. Soc. 108,2532 (1986). 318. Lin, R. F., Blackman, G. S., Van Hove, M. A,, and Somorjai, G. A., Acta Crystallogr. Sect. B: Struct. SCLB43, 368 (1987). 319. Barbieri, A,, Van Hove, M. A,, and Somorjai, G. A,, Surf: Sci. 306,261 (1994). 320. Wander, A., Held, G., Hwang, R. Q., Blackman, G. S., Xu, M. L., de Andres, P., Van Hove, M. A,, and Somorjai, G. A,, Surf: Sci. 249,21 (1991). 321. Herzberg, G., “Infrared and Raman Spectra of Polyatomic Molecules,” pp. 364-365. Van Nostrand, New York, 1945. 322. Wilson, E. B., Jr., Decius, J. C., and Cross, P. C., “Molecular Vibrations,” p. 249. McGrawHill, New York, 1955; Wilson, E. B., Jr., Phys. Rev. 45,706 (1934). 323. Shimanouchi, T., “Tables of Molecular Vibration Frequencies,” Consolidated Vol. 1, pp. 152, 153. National Bureau of Standards, Washington, DC, 1972. 324. Avery, N. R., personal communication, 1987. 325. Hoffmann, H., Zaera, F., Ormerod, R. M., Lambert, R. M., Wang, L. P., and Tysoe, W. T., Surf Sci. 232, 259 (1988). 326. Anson, C. E., and Powell, D. B., personal communication; Anson, C. E., Ph.D. Thesis, University of East Anglia (1988). 327. Witte, G., Range, H., Toennies, J. P., and Woll, Ch., Phys. Rev. Lett. 71, 1063 (1993). 328. Takenaka, T., and Sheppard, N., unpublished work, 1966. 329. Haaland, D. M., Surf Sci. 102, 405 (1981). 330. Szilhgyi, T., J. Mol. Struct. 174, 395 (1988). 331. Grimm, W. M., 111, M.Sc. Thesis, University of East Anglia (1981). 332. Francis, S. A., J. Chem. Phys. 18, 861 (1950). 333. Golicke, R. C., Mills, I. M., Person, W. P., and Crawford, B., Jr., J. Chem. Phys. 25, 1266 (1956). 334. Krasser, W., Ervens, H., Fadini, A,, and Renouprez, A. J., J. Raman Spectrosc. 9, 80 (1980). 335. Krasser, W., J. Mol. Struct. 80, 187 (1982). 336. Krasser, W., and Renouprez, A. J., Solid State Commun. 41, 231 (1982). 337. Krasser, W., Bechthold, P. S., and Kettler, U., Z. Anal. Chem. 314, 319 (1983). 338. Krasser, W., Kojnok, J., and Guerts, J., J. Raman Spectrosc. 22, 777 (1991). 339. Ref. 110, Spectrum F1-01. 340. Hexter, R. M., and Albrecht, M. G., Spectrochim. Acta, Part A 35, 233 (1979). 341. Nichols, H., and Hexter, R. M., J. Chem. Phys. 73, 965 (1980). 342. Moskovits, M., and Di Lella, D. P., J. Chem. Phys. 73, 6068 (1980). 343. Wolkow, R. A,, and Moskovits, M., J. Chem. Phys. 84,5196 (1986). 344. Parker, W. L., Hexter, R. M., and Seidle, A. R., J. Am. Chem. SOC. 107, 4584 (1985). 345. Gao, P., and Weaver, M. J., J. Phys. Chem. 89, 5040 (1985). 346. Gao, P., Davies, J. P., and Weaver, M. J., J. Phys. Chem. 94, 6858 (1990). 347. Lund, P. A., Tevault, D. E., and Smardzewski, R. R., J. Phys. Chem. 88, 1731 (1984); Chem. Phys. Lett. 89,508 (1982). 348. Jobic, H., Tomkinson, J., Candy, J. P., Fouilloux, P., and Renouprez, A. J., Surf Sci. 95, 496 (1980).

VIBRATIONAL SPECTRA OF HYDROCARBONS 349. 350. 351. 352. 353. 354. 355.

311

Graham, D., and Howard, J., J. Chem. SOC. Faraday Trans. I 80,3365 (1984). Coats, A. M., Cooper, E., and Raval, R., Surf: Sci. 307, 89 (1994). Avery, N. R., J. Chem. Soc., Chem. Commun., p. 153 (1988). Rauscher, H., Jakobs, P., Menzel, D., and Lloyd, D. R., Surf: Sci. 256,27 (1991). Rauscher, H., and Menzel, D., Surf: Sci. 342, 155 (1995). Wilmshurst, J. K., and Bernstein, H. J., Can. J. Chem. 35, 911 (1957). Varsinyi, G . ,“Assignments for Vibrational Spectra of 700 Benzene Derivatives,” Hilger, London (1974). 356. Wilk, D. E., Stanners, C. D., Shen, Y. R., and Somorjai, G . A,, Surf: Sci. 280,298 (1993). 357. Myers, A. K . , and Benziger, J. B., Langmuir 3, 414 (1987). 358. Davydov, V. Y . , and Sheppard, N. (unpublished work, 1967). 359. Cremer, P. S., and Somorjai, G. A., J. Chem Soc. Faraday Trans. 91, 3671 (1995). 360. Yagasaki, E., and Masel, R. I., “Catalysis” (Royal Society of Chemistry Specialist Periodical Report) 11, 1655 (1994). 361. Vayssilov, G . N . , Adv. Colloid & Interface Science 47, 25 (1993). 362. Zaera, F., Chem. Rev. 95, 2651 (1995). 363. Bradshaw, A. M., Surf: Sci. 331-333,978 (1995). 364. Bao, S., Hofmann, Ph., Schindler, K.-M., Fritzsche, V., Bradshaw, A. M., Woodruff, D. P., Casado, C., and Asensio, M. C., J. Phys. Condens. Matt. 6, L93 (1994). 365. Bao, S . , Hofrnann, Ph., Schindler, K.-M., Fritzsche, V., Bradshaw, A. M., Woodruff, D. P., Casado, C., and Asensio, M. C., Surf: Sci. 323, 19 (1995). 365a. Gutdeutsch, U., Birkenheuer, U., Bertel, E., Cramez, J., Boettger, J. C., and Rosch, N., Surf: Sci. 345, 331 (1996). 366. Schaff, O., Stampfl, A. P. J., Hofmann, Ph., Bao, S., Schindler, K-M., Bradshaw, A.M., Davis, R., Woodruff, D. P., and Fritzsche, V., Surj Sci. 343, 201 (1995). 367. Buisset, J., Rust, H.-P., Schweizer, F. K., Cramer, L., and Bradshaw, A. M., Phys. Rev. B. 54, 10373 (1996). 368. Jenks, C . J., Bent, B. E., Bernstein, N., and Zaera, F., Surf: Sci. 277, L89 (1992). 369. Kubota, J . , Kondo, J. N., Dornen, K., and Hirose, C., J. Phys. Chem. 98,7653 (1994). 370. Doll, R., Gerken, C. A,, Van Hove, M. A., and Somorjai, G. A., Surf: Sci. 374, 151 (1997). 371. Cremer, P., Stanners, C., Niemantsverdriet, J. W., Shen, Y. R., and Somorjai, G. A., Surf: Sci. 328, 111 (1995). 372. Cremer, P. S . , Su, X., Shen, Y. R., and Somorjai, G. A,, J. Am. Chem. Soc. 118,2942 (1996); Catal. Lett. 40, 143 (1996). 373. McDougall, G., and Yates, H., in “Catalysis and Surface Characterization,” Spec. Publ. 114, p. 109. Royal Society of Chemistry, London, 1992. 373a. Kubota, J., Ichihara, S., Kondo, J. N., Domen, K., and Hirose, C., Surf: Sci. 357-358, 634 (1996). 3736. Jia, J . F., Wu, K., Lu, S.-H., Zhao, R. G., Wei, X.-M., Wu, S.-C., and Wang, D.-Z., Surf: Sci. 338,69 (1995). 373c. Fruhberger, B., and Chen, R. G., Surf: Sci. 342,38 (1995). 374. Celio, H . , Trenary, M., and Robota, H. J., J. Phys. Chem. 99, 6024 (1995). 375. De La Cruz, C., and Sheppard, N., Catal. Lett. 37, 47 (1996). 376. Ayre, C. R., and Madix, R. J., Surf: Sci. 262, 51 (1992). 376a. Pawela-Crew, J., Madix, R. J., and Vasquez, N., Surf: Sci. 340, 119 (1996). 377. Bertolini, J. C., Cassuto, A,, Jugnet, Y., Massardier, J., Tardy, B., and Tourillon, G., Surf: Sci. 349, 88 (1986). 378. Hostetler, M. J., Nuzzo, R. G., Girolami, G. S., and Dubois, L. H., Organometullics 14, 3377 (1995).

312

NORMAN SHEPPARD AND CARLOS D E LA CRUZ

378u. Hostetler, M. J., Nuzzo, R. G., and Girolami, G. S . , J. Am. Chem. Soc. 117,1814 (1996). 379. Windham, R. G., Bartram, M. E., and Koel, B. E., J. Phys. Chem. 92,2862 (1988). 380. Berlowitz, P., Megiris, C., Butt, J. B., and Kung, H. H., Lungmuir 1,206 (1985). 381. Godbey, D., Zaera, F., Yeates, R., and Somorjai, G. A., Surf Sci. 167, 150 (1986). 382. Salmeron, M., and Somorjai, G. A., J. Phys. Chem. 86,341 (1982). 383. Creighton, J. R., and White, J. M., Surf Sci. 129, 327 (1983). 384. Ogle, K. M., Creighton, J. R., Akhter, S., and White, J. M., Surf Sci. 169, 246 (1986). 385. Pettiette-Hall, C. L., Land, D. P., McIver, R. T., Jr., and Hemminger, J. C., J. Phys. Chem. 94, 1948 (1990). 386. Erley, W., Li, Y., Lang, D. P., and Hemminger, J. C., Surf Sci. 301, 177 (1994). 387. Mohsin, S . B., Trenary, M., and Robota, H., Chem. Phys. Lett. 154,511 (1989). 388. Gland, J. L., Zaera, F., Fisher, D. A., Carr, R. G., and Kollin, E. B., Chem. Phys. Lett. 151, 227 (1988). 389. Deleted in proof. 390. Mitchell, I. V., Lennard, W. N., Griffiths, K., Massouri, G. R., and Huppertz. J. W., Sure Sci. 256, L598 (1991). 391. There is another possibility in the literature. Freyer et ul. [Freyer, N., Pirug, G., and Bonzel, H. P., Surf. Sci. 125, 327 (1983); 126, 487 (1983)l used quantitative XPES to measure the saturation coverage of C2H4adsorbed on P t ( l l l ) by comparing the areas under the C(1s) peak of adsorbed C2H4and of adsorbed CO. Their results indicated a CzH4saturation coverage of 0.5 at 100K. Possible explanations for this result are discussed in Griffiths et ul. (392). 392. Griffiths, K., Lennard, W. N., Mitchell, I. V., Norton, P. R., Pirug, G., and Bonzel, H. P., Surf Sci. 284, L389 (1993). 393. Land, T. A., Michely, T., Behm, P. J., Hemminger, J. C., and Comsa, G. A,, Appl. Phys. A53, 414 (1991). 394. Land, T. A., Michely, T., Behm, R. J., Hemminger, J. C.. and Cosma, G . A., J. Chem. Phys. 97, 6774 (1992). 395. Land, T. A,, Michley, T., Behm, R. J., Hemminger, J. C., and Cosma, G. A,, Surf Sci. 264, 261 (1992). 396. Windham, R. G., Koel, B. E., and Paffett, M. T., Lungmuir 4, 1113 (1988). 396u. Christmann, K., Ertl, G., and Pignet, T., Surf Sci. 54, 365 (1976). 397. Zaera, F., J. Phys. Chem. 94, 5090, 8350 (1990); J. Am. Chem. Soc. 111, 8744 (1989); Surf Sci. 219, 453 (1989). 398. Kemball, C., Adv. Cutul. 11, 223 (1959); Card Rev. 5, 33 (1971). 399. Zaera, F., and Somorjai, G. A,, J. Am. Chem. Soc. 106,2288 (1984). 400. Somorjai, G. A., Van Hove, M. A., and Bent, B. L., J. Phys. Chem. 92, 973 (1988). 401. Anderson, A. B., and Choe, S. J., J. Phys. Chem. 93, 6145 (1989). 402. Davis, S. M., Zaera, F., Gordon, B. E., and Somorjai, G. A,, J. Cutul. 92, 240 (1985). 403. Hugenschmidt, M. B., Dolle, P., Jupille, J., and Cassuto, A,, . I . Vuc. Sci. Technol., A [2] 7, 3312 (1989). 404. Backman, A. L., and Masel, R. I., J. Vuc. Sci. Technol., A [2] 9, 1789 (1991). 405. Jackson, S. D., Glanville, B. M., Willis, J., McLellan, G. D., Webb, G., Moyes, R. B., Simpson, S . , Wells, P. B., and Whyman, R., J. Cutul. 139, 221 (1993). 406. Soma, Y., J. Cutul. 59, 239 (1979). 407. De La Cruz, C., and Sheppard, N., J. Chem. Soc., Chem. Commun., p. 1854 (1987). 408. Mohsin, S . B., Trenary, M., and Robota, H. J., J. Phys. Chem. 92, 5229 (1988). 409. Wang, P., Slichter, C. P., and Sinfelt, J. H., J. Phys. Chem. 89, 3606 (1986). 410. Zax, D. B., Kling, C. A., Slichter, C. P., and Sinfelt, J. H., J. Phys. Chem. 93, 5009 (1989).

VIBRATIONAL SPECTRA OF HYDROCARBONS

313

411. Mohsin, S. B., Trenary, M., and Robota, H. J., J. Phys. Chem. 95,6657 (1991). 412. Hensley, D. A,, and Kesmodel, L. L., J. Phys. Chem. 95, 1368 (1991). 413. Chesters, M. A,, De La Cruz, C., Gardner, P., McCash, E. M., Prentice, J. D., and Sheppard, N., J. Electron Spectrosc. Relat. Phenom. 54/55, 739 (1990). 414. Cortright, R. D., Goddard, S. A,, Rekoske, J. E., and Dumesic, J. A., J. Catal. 127, 342 (1991). 415. Schlatter, J. C., and Boudart, M., J. Catal. 24, 482 (1972). 416. Dorling, T. A., Eastlake, M. J., and Moss, R. L., J. Catal. 14, 23 (1969). 417. Bond, G. C., Phillipson, J. J., Wells, P. B., and Winterbottom, J. M., Trans. Faraday Soc. 60, 1847 (1964). 418. Kazanskii, V. B., and Strunin, Kinet. Catal. 1, 517 (1960). 419. Bond, G. C., Trans. Faraday SOC.52, 1235 (1956). 420. Rekoske, J. E., Cortright, R. D., Goddard, S. A,, Sharma, S. B., and Dumesic, J. A,, J. Phys. Chem. 96, 1880 (1992). 421. Sheppard, N., Awry, N. R., Morrow, B. A,, and Young, R. P., in “Chemisorption and Catalysis” (P. Hepple, ed.), pp. 135-147. Inst. Petrol., London, 1971. 422. Kember, D. R., and Sheppard, N., J. Chem. Sac., Faraday Trans. 2 77, 1321 (1981). 423. Soma, Y., J. Catal. 75, 267 (1982). 424. Yagasaki, E., and Masel, R. I., Surf Sci. 226, 51 (1990); J. Vac. Sci. Technol., A [2] 8, 2616 (1990). 425. Bowker, M., Gland, J. L., Joyner, R. W., Li, Y., Slinko, M. M., and Whyman, R., Catal. Lett. 25,293 (1994). 426. Horiuti, J., and Polanyi, M., Trans. Faraday Soc. 30, 1164 (1934). 427. Zaera, F., Janssens, T. V. W., and Ofner, H., Surf Sci. 368, 371 (1996). 428. Thomson, S. J., and Webb, G., .I Chem. . Soc., Chem. Commun. 526 (1976). 429. Hattori, T., and Bunvell, R. L., Jr., J. Phys. Chem. 83, 241 (1979).

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ADVANCES IN CATALYSIS, VOLUME 42

Application of Combined X-Ray Diffraction and Absorption Techniques for in Situ Catalyst Characterization BJERNE S. CLAUSEN,' HENRIK TOPSaE,' AND RONALD FRAHM2 Haldor Topsoe Research Loboratories DK-2800 Lyngby, Denmark Institut fiir Angewandte Physik Heinrich-Heine- Universitat 0-40225 DiisseldorJ Germany

I.

Introduction

Information characterizing the structural arrangements of the atoms in a catalyst is essential for the understanding of many catalytic phenomena. Structural information is often obtained from diffraction techniques, and X-ray diffraction (XRD) has for years been the most widely used structural technique in catalytic research. It is therefore not surprising that the first volume of Advances in Catalysis, published in 1948, contained a review of the application of XRD to the study of catalysts (1). In recent years, it has been realized that techniques based on X-ray absorption provide important additional possibilities for catalyst characterization. Techniques such as X-ray absorption fine structure (XAFS) spectroscopy have had a significant impact on catalyst research. For example, the application of these techniques has for the first time allowed structural descriptions of many catalysts which, because of the presence of microcrystalline structures (nanophase particles) or amorphous phases, cannot be elucidated by XRD. Some of the potential benefits of X-ray absorption techniques in studies of catalysts were pointed out in a pioneering article by Van Nordstrand ( 2 ) ,but widespread application of X-ray absorption spectroscopy first began only after Sayers, Stern, and Lytle (3) had provided a theoretical underAbbreviations: CCD, charge-coupled device; DEXAFS, dispersive extended X-ray absorption fine structure; EXAFS, extended X-ray absorption fine structure; QEXAFS, quick extended X-ray absorption fine structure; TPR, temperature-programmed reaction; XRD, X-ray diffraction; XAFS, X-ray absorption fine structure.

315 Copyright 0 1998 by Academic Press All rights of reproduction in any form reserved. 0360-0564/98 $25.00

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standing of the extended X-ray absorption fine structure (EXAFS). Reviews in this series ( 4 , 5 )have highlighted the possible applications and subsequent progress in the quantitative understanding of XAFS data. X-ray absorption experiments usually require access to high-intensity synchrotron radiation. A number of synchrotron facilities are now available worldwide, and their usefulness is continuously improving, opening the possibility for several new types of studies. Experimental and theoretical progress has been reviewed (&lo), and only some recent developments related to catalysis are discussed here. One of the inherent advantages of both X-ray absorption and diffraction techniques is the possibility to perform experiments while the catalyst is in its working environment. This is important because such in situ studies assure that the structural information is directly related to the actual working catalyst. Thus, direct correlations can be established between the structural and chemical features and the performance of the catalyst under different processing conditions. These important links are generally referred to as structure-activity relationships. Consequently, one of the focal points of this review is a discussion of the recent advances allowing in situ X-ray experiments (Section 111). Apart from the many advantages of X-ray diffraction and X-ray absorption spectroscopy, each method, as discussed in Sections I1 and IV, is also characterized by some important limitations in the investigation of catalysts. These limitations to a large extent are complementary, and they can therefore be overcome by using a combination of both techniques. The various approaches for performing experiments with the combined techniques are described in Section V, and some examples are given in Section VI. X-ray diffraction and absorption techniques have rarely been used to investigate time-dependent changes. However, it will be shown that the techniques may also be used to characterize the dynamics of catalytic processes (Sections IV-VII), and in recent years the time resolution has improved by several orders of magnitude. This improvement has opened the door to new types of quantitative studies of the dynamical behavior of catalysts, and such information is often necessary for a detailed description and understanding of catalytic phenomena. Some possible future applications and advances in the techniques are summarized in Section VII.

II.

X-Ray Diffraction and Absorption Spectroscopy Applied to Catalyst Characterization

This section includes a general discussion of the applicability of X-ray diffraction and X-ray absorption techniques for detailed structural charac-

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317

terization of various catalysts. To illustrate the advantages and limitations of the methods, the essential features of each are first reviewed. X-ray diffraction patterns from typical catalyst powders give information about the interplanar lattice spacings through the Bragg equation 2d sin 8 = nh,

(1) where d is the interplanar spacing, 8 is the angle between the lattice plane and both the incident and the diffracted X-ray beam, A is the wavelength of the X-rays, and n is the order of the Bragg reflection. Combined with the fact that the intensities of diffraction lines depend on the arrangement of atoms in the unit cell of the crystal lattice (the crystal structure), this information in principle provides an almost unique description of the nature of the crystalline phases present. However, in practice, the interpretation of the diffraction patterns may not always be trivial. Catalysts often contain many different phases, and this complexity may result, for example, in overlapping peaks and complications in the data interpretation. The presence of broadened diffraction lines may complicate the analysis, but on the other hand also provides information about the particle sizes (D)of small crystallites, for example, through the Scherrer formula Kh D=-------

B cos 8 ’

where K is the Scherrer constant, which is dependent on the crystallite shape and in most cases is close to 0.9, and B is the line width in radians at half-maximum after correction for instrumental broadening. From a more complete line shape Fourier analysis, information about the particle size distribution may also be provided ( I I , I 2 ) . In typical catalysts, phenomena such as inhomogeneities and strainlstress may contribute to line broadening, and it is not trivial to separate the different contributions (11-13). When synchrotron radiation is used, the experimental line width is drastically reduced by the extremely small divergence of the beam, and the diffractograms of catalysts can be more easily interpreted (see Section IV) . The most serious limitation of XRD as a catalyst characterization method is often related to the fact that many of the phases present in a catalyst may not give rise to any well-defined diffraction line at all. Absence of a diffraction pattern is a consequence of the requirement that a structure must contain a periodicity extending more than about 2-3 nm to yield a diffraction pattern measurable in a sense of the Bragg equation [Eq. (l)]. Thus, particles or domains with sizes smaller than 2-3 nm will appear to be X-ray amorphous in XRD experiments; i.e., they do not exhibit sharp diffraction lines.

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These limitations no longer pertain in X-ray absorption spectroscopy, whereby inner atomic shells are ionized and photoelectrons are ejected. An absorption spectrum consists of the absorption edge itself, the nearedge region (XANES), extending from the absorption threshold to about 50 eV beyond the edge, and the EXAFS oscillations, which may be observable (in special cases) up to about 1500 eV beyond the edge. The fine structure is a consequence of the interference between the outgoing photoelectron wave and the part that is backscattered from neighboring atoms. It directly probes the local atomic environment of the absorbing atom. According to the plane-wave, single-scattering approximation, the EXAFS ) to j coordination shells can be expressed by the oscillations ~ ( k due equation

(-?)

exp( -2k2aj2) exp - sin[2kR, + 6,(k)], i

(3)

where k is the photoelectron wave number, N j is the coordination number in the jth shell, Ri is the interatomic distance, F,(k) is the backscattering amplitude, aj is the mean relative vibrational amplitude, h is the electron mean free path, and S, is a phase shift associated with the interaction of the photoelectron with the potential of the absorbing and the backscattering atom. As a consequence of the limited mean free path of the photoelectrons characterized by A, EXAFS is useful in providing information about the local short-range order around the absorbing atom. Typically, the structural information is obtained by Fourier-transforming the X ( k ) into R-space, giving a radial distribution function from which coordination numbers and interatomic distances are determined. In principle, EXAFS information may be obtained for most or all of the elements in a catalyst. Thus, for multicomponent samples, the characterization of local surroundings for all (or almost all) the elements may be obtained. However, we stress that the radial distribution function cannot be transformed into a unique three-dimensional structure. Therefore, the EXAFS technique is not ideal for providing such information and the data representing materials consisting of several different phases may often be too difficult to analyze meaningfully. Recently, significant progress has been made in the theoretical understanding of the XANES part of the absorption spectrum. This part is dominated by multiple scattering events and may thus provide information characterizing the local geometry of the atoms neighboring the absorbing atoms. By use of modern analysis packages such as the FEFF and XCURVE codes (14,15),which take into account multiple scattering, it is now possible to extract some information about the three-dimensional structure of mate-

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rials. However, the presence of a number of phases together makes unambiguous interpretation difficult. On the basis of the foregoing points, we illustrate the advantages and limitations of XRD and EXAFS by discussing how they provide information to describe the different types of catalyst structures shown schematically in Fig. 1. The different cases are distinguished by the presence or absence of long-range and/or short-range order, because this is the key issue determining the applicability of the different X-ray techniques. A typical pattern representing a catalyst that contains only structures with local order (i.e., one or a few well-defined coordination shells) is illustrated in Fig. la. This situation could represent immobilized complexes, atoms in ionexchange positions, solid solutions, very small clusters, promoter atoms at surfaces, etc. It is clear that the complete lack of long-range order makes XRD a rather useless characterization tool for such catalysts. On the other hand, XAFS is an ideal technique in this case, and this technique may often be the only tool that can provide the needed insight. In contrast, Fig. l b illustrates a pattern of a catalyst with a broad distribution in the dimensions of structural order. This could, for example, represent

0.1

1

10

100

Dimension (size), nm

FIG. 1. Schematic representation of different size distributions of structures present in typical catalysts. The X-ray amorphous region is indicated by the shaded area.

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small nanoclusters with a distribution in their sizes. There are many examples of such catalysts; the nanoclusters or particles could be located inside a zeolite or on the surface of a support, being present as oxides, sulfides, or metals. As long as the dimensions of order are less than about 2-3 nm (the XRD cutoff), the applicability of the techniques is the same as already stated (i.e., XRD provides little or no insight). The unique advantages of XAFS in studies of catalysts represented by the patterns shown in Figs. l a and l b have been exploited many times (6, 9, 16). Figure l c illustrates an example which could correspond to the presence of larger particles which are typically encountered in many medium to low surface area catalysts. In the example, all the structures have dimensions larger than the XRD cutoff and thus XRD provides detailed insight into all the structures present. In investigations of such catalysts, XRD is the preferred technique, and XAFS does not provide much-if anyadditional insight. In many real catalysts, none of the simplified situations described in Figs. la, lb, or l c may apply alone. Rather, one often encounters complex samples incorporating structures with both short-range and long-range order; i.e., one may have the simultaneous presence of structures above and below the XRD cutoff (Fig. Id). Typical examples are Co-Mo/A1203 hydrotreating catalysts (I 7). These catalysts generally have the major fraction of the Co atoms atomically dispersed at the edges of small MoSz clusters (the Co-Mo-S structures). Furthermore, a small fraction of the Co atoms may be located in the alumina lattice and, especially at high Co loadings, 1- to 5-nm Cogs8particles may also form. The MoSz clusters are typically present as 1- to 2-nm singlelayer structures, and the A1203 crystallites are about 4 nm in dimension. Thus, for hydrotreating catalysts, XRD provides a diffuse picture (consisting of very broad diffraction lines of MoSz and the alumina support (17, 18). However, a wealth of additional structural information can be provided by XAFS (16, 29). Zeolite-supported metal catalysts with atoms and clusters existing both inside and outside the zeolite channels are often examples of the situation depicted in Fig. Id. Again, XAFS has been an important tool for characterizing such catalysts (4-10). It is clear from the foregoing discussion that the X-ray absorption and diffraction techniques to a large extent complement each other and that a much more complete structural description can be obtained by combining the two methods. Recently, a number of approaches have been developed which allow such combined measurements to be performed simultaneously on the same sample (Section V).

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In Situ Approaches

111.

One of the main advantages of both X-ray absorption and diffraction techniques is that they offer the possibility of carrying out relevant in situ experiments. Most materials are relatively transparent to high-energy Xrays, and that allows the construction of in siru cells when appropriate windows are used. Figure 2a illustrates the general principle of a design used in several investigations for performing in situ XRD (20-27). The catalyst in the form of a flat sample was mounted on a heating stage, and the windows were often cooled to protect them and/or the seals. If one uses the so-called 8-Osetup, both the X-ray source and detector are moved, and the advantage is that the sample is stationary and may be kept flat (26, 27). The very simple and inexpensive in siru cell (26) shown in Fig. 2b is easily adapted to most commercial in-house X-ray powder diffractometers. To allow pretreatments at higher temperatures than most window materials can stand, the sample can be moved to a quartz heating zone. The collection a

Gas out

b Detector

-

Gas out

Temperature controller

FIG.2. Schematic representation of typical in situ cells used in XRD experiments: (a) general design principles; (b) commercial-type X-ray in situ cell [adapted from Jung and Thomson (26)].

322

BJERNE S. CLAUSEN, HENRIK TOPS0E, AND RONALD FRAHM

time for a low-noise diffractogram is typically of the order of minutes to hours. The cells used for in situ EXAFS experiments (7, 28-30) have in many cases been designed on the basis of principles similar to those shown in Fig. 2. An example is shown in Fig. 3 (29),which clearly illustrates the necessity to place the windows away from the central heated zone of the catalyst sample in order to allow water cooling of the windows and seals. All the XRD and EXAFS cells discussed in the preceding paragraph have some severe drawbacks. The range of pressures and temperatures is typically quite limited. Furthermore, the local surroundings of the catalyst are very different from those encountered in normal catalytic reactors, whereby gas flows through a catalyst bed. Thus, the cell will unavoidably exhibit some unusual concentration, flow, and temperature gradients and

FIG. 3. Cross section of the side view of a typical EXAFS in situ cell: (A) main body made of stainless steel; (B) beryllium windows; (C) convection baffles; (D) gas inlet; (E) liquid inlet; (F) gas and liquid outlet; (G) cooling jackets; (J) thermocouple [reproduced from Dalla Betta ef al. (29), with the permission of the American Institute of Physics].

In Situ CATALYST CHARACTERIZATION

323

these are difficult to estimate. Thus, it is difficult to ensure that the pretreatments of catalyst samples have been carried out in a well-defined manner, and the resulting state of the catalyst may be quite different from that of catalysts pretreated in normal catalytic test reactors. Furthermore, it is difficult to ensure that the conditions during catalysis are identical to those encountered in the actual reactors. Consequently, there may not be a clear connection between the structural information obtained from such cells and the structures of catalysts working in typical flow reactors. One approach used to overcome these difficulties has been to adopt the same geometry in the in situ cell as that used in a fixed-bed, tubular flow reactor (31). It is not possible to use normal tubular flow test reactors directly as usual XRD or EXAFS cells because their dimensions are so large and the reactor walls are so thick that the transmitted X-rays are attenuated too much. A convenient solution to both of these problems is to miniaturize the tubular flow reactor and to use capillaries as combined cells and catalytic reactors (32). Such cells also allow experiments at high pressures; for example, high-quality quartz glass capillaries of a diameter of 0.5 mm and a wall thickness of 10 pm can stand a pressure of about 60-80 bar. This approach is illustrated in Fig. 4. The catalyst sample in powder form is arranged as a normal catalyst fixed bed, and during the experiment, gases may pass through the bed such that unwanted gradients can be minimized. Depending on the construction material of the capillary, temperatures exceeding 1000°C and pressures exceeding 100 bar can be achieved. The heating/cooling may be done indirectly, and temperature uniformity can be better than 1°C (32). The low mass and heat capacity also allow quite rapid transients to be followed. This feature is of the utmost importance since chemical reactions in the kinetically controlled regime can be investigated, whereas slower variations in the temperature result in a quasi-steady state of the catalyst. The catalytic activities obtained in such in situ cells are similar to those measured in separate ideal tubular flow capillary tube catalyst ,“-“-pm>-o!yp

stainless steel tube

7

FIG.4. Schematic drawing of combined capillary in situ cell and catalytic fixed bed reactor [adapted from Clausen et al. (32)].

324

BJERNE S. CLAUSEN, HENRIK TOPSBE, AND RONALD FRAHM TABLE I Comparison of Methanol Synthesis Data Obtained in rhe in Situ Cell with Those of a Laboratory Pilot Reactor Concentration (~01%)" In situ cell Component

0.1 MPa

Laboratory pilot reactor

3 MPa

3 MPa

ma. 4.5 4.3 0.3

86.93 4.85 4.23 0.34

~

Hz

co coz CH,OH

n.a. 4.6 4.7 0.01

Note: n.a., not analyzed. a Cu/ZnO/A120, catalyst, SV = 30,000, T = 493 K, inlet gas composition: 4.6% CO, 4.7% COz,3.1% Ar, balance H2.

reactors (Table I) and this important feature makes possible the comparisons of meaningful structural and catalytic information. The preceding setup allows both X-ray diffraction (32) and absorption experiments (33,34).The capillary geometry was used until about 30 years ago for ex situ XRD studies in connection with the placement of Lindemann tubes in powder Debye-Scherrer cameras. At that time, films were used to detect the diffracted X-rays. Today, this cumbersome technique has been almost completely replaced as modern detectors are used.

IV. Recent Advances in X-Ray Diffraction and X-Ray Absorption Techniques A. XRD As discussed, XRD has for many years been the standard, everyday characterization method for solid catalysts, and in almost every laboratory in this field there is access to an X-ray diffractometer. This instrument allows a wide variety of different characterizations, but there are also limitations of such equipment. For example, the limited resolution of an in-house diffractometer may often be insufficient for a detailed analysis. This point is illustrated in Fig. 5a, which shows the diffractogram of an industrial type steam-reforming catalyst consisting of nickel crystallites on a spinel support (35).The Ni(ll1) and the spineI(400) lines overlap so that a detailed analysis is impossible. This problem can be overcome if the XRD

325

In Siru CATALYST CHARACTERIZATION

a

J, -80

-70

-60

-50

-40

-30

2 0 (degree)

.-b S u)

a,

c C

-27

-25

-23

-21

-19

2 0 (degree)

FIG.5. (a) In-house Cu K a XRD diffractogram of a nickel-based steam-reforming catalyst supported on a spinel in the reduced state. (b) High-resolution XRD patterns obtained by using synchrotron radiation. Only the most important nickel lines of the steam-reforming catalyst after reduction, steaming at 790"C, steaming at 89OoC,and industrial use are shown. Sp denotes peaks from the MgAlZO4spinel in the substrate, whereas A1 denotes peaks from the alumina phase. Adapted from Niemann e t d (35).Copyright 1990 Societa Italiana di Fisica.

experiments are performed at a synchrotron radiation facility, since the experimental line width in this case is narrower by an order of magnitude than that measured with a typical in-house diffractometer. Figure 5b shows the corresponding diffractogram obtained by use of synchrotron radiation. The nickel and the spinel lines are now well separated. The significantly reduced experimental line width enables one to characterize in detail the changes taking place in the active nickel crystallites. In general, the reduced experimental line width opens new opportunities to perform line profile analysis.

326

BJERNE S. CLAUSEN, HENRIK TOPSBE, AND RONALD FRAHM

The relatively low intensity of a normal in-house facility based on X-ray tubes makes detailed analysis of dilute catalysts difficult, and a high-quality diffractogram may take several hours to record. The intensity can be improved by an order of magnitude by use of a rotating-anode X-ray generator, but the intensity available from a synchrotron facility is many orders of magnitude higher. In an in-house facility, a diffractogram is typically obtained by scanning the diffraction angle, and in most cases both the sample and the detector are rotated. As a result, the different parts of the diffractogram are not recorded at the same time. Consequently, meaningful timeresolved on-line studies are very difficult unless the changes occur over a time scale of hours or days. In view of this difficulty in performing in situ experiments, X-ray diagrams have typically been recorded after quenching of the sample to room temperature or by stepwise increasing of the temperature to the desired value and then keeping the temperature constant while the diffractogram was accumulated (20, 35-39). It was shown (40) that the aforementioned problem may be overcome by recording the whole diffraction diagram simultaneously by using the socalled energy-dispersive X-ray diffraction technique. This technique also has the advantage of reducing the data acquisition time, eliminating the problem arising from the presence of K a doublets and reducing the fluorescence radiation. Because the sample is fixed, the construction of in situ equipment is facilitated. The energy resolution of the energy-dispersive detector, which is much poorer than the resolution of an angular-dispersive detector, is the main limitation of this technique and in some cases may be a serious drawback.

31

33

35

37

39

41

43

2 0 (degree) FIG.6. Three-dimensional representation of in situ XRD patterns recorded during reduction of a Cu/ZnO methanol synthesis catalyst. The collection time per diagram was 60 s [adapted from Clausen el al. (32)].

In Siiu CATALYST CHARACTERIZATION

327

The availability of high-intensity synchrotron radiation can significantly reduce the time necessary for recording a diffractogram. The many modern detectors, such as image plates, CCD cameras, and positionsensitive detectors, can now reduce the data collection time for typical catalysts to a few seconds and sometimes even significantly less. This development makes possible time-resolved XRD investigations of catalysts. An example is illustrated in Fig. 6, which shows a three-dimensional plot of XRD diagrams recorded during reduction of a copper-based methanol synthesis catalyst (32). A position-sensitive detector was used. This type of detector is continuously improving and has recently been developed to cover large Bragg angles, enabling essentially all diffraction lines to be recorded without moving the detector. In their resolution and sensitivity, modern position-sensitive detectors for powder diffraction match traditional point detectors well, as shown by their successful use in Rietveld refinement. B. EXAFS X-ray absorption spectra are most commonly recorded in transmission or fluorescence mode. Double crystal or channel-cut crystal monochromators are used to select X-rays with a specific energy (defined by the Bragg angle of the Si crystals that are usually used) from the white beam of synchrotron radiation. The whole absorption spectrum is obtained by rotating the monochromator step by step through the angular range. The absorption or fluorescence yield from the sample is measured at each step, resulting in an absorption/fluorescence spectrum as a function of energy. The time needed to record a high-quality EXAFS spectrum in this mode typically ranges from about 15 min to 1 h, depending on the properties of the sample (e.g., composition and homogeneity), the spectrometer (e.g., mechanical stability and detector efficiency), and the synchrotron radiation source (e.g., electron/positron current, stability, and divergence). Obviously, it is a prerequisite that the samples have structures and chemical states that do not change on the time scale of a typical EXAFS measurement since otherwise the first and last parts of the spectrum would contain information from different states of the sample, making interpretation of the spectra almost impossible. When such structural or chemical unstable samples are used and when the dynamic properties are of interest, it is necessary to record the EXAFS spectra on a shorter time scale, say a few minutes, seconds, or even milliseconds, depending on the time scale of the dynamic process. Several years ago, it was demonstrated that time-resolved EXAFS measurements could be performed by use of the energy-dispersive method (41),

328

BJERNE S. CLAUSEN, HENRIK TOPSDE, AND RONALD FRAHM

also abbreviated DEXAFS. This method is appealing in principle, but for catalyst characterizations it has some disadvantages which so far have limited the widespread use of the technique. A bent crystal monochromator is used to resolve spatially photons of different energies and to focus these so that they-ideally-hit the sample at a single point. The sample is thus illuminated by a “white” photon beam, and a transmission EXAFS spectrum can be recorded by use of a position-sensitive detector behind the sample. The setup is attractive from the point of view that no mechanical movements of, for example, heavy turntables are involved. Reasonably high-quality spectra of highly concentrated materials (typically metal films) have been recorded on the time scale of milliseconds (42-45). However, a point focus is difficult to obtain in practice because of nonideal curvature of the monochromator crystal. This means that photons with different energies do not hit the sample at the same point, and as a result extreme sample uniformity is required to avoid additional background oscillations from interfering with the EXAFS. With powder materials such as catalysts, it is difficult to obtain spectra of sufficient quality to extract meaningful quantitative information. Nevertheless, the results of such experiments may provide important qualitative information about catalyst dynamics. Typically, the best results have been obtained with ion-exchanged zeolites (45) since these are very homogeneous samples. Figure 7 shows recent DEXAFS results representing the oscillating oxidation reaction of a Pt-exchanged zeolite (45).Although DEXAFS is ideal from a conceptional point of view for time-resolved experiments, it is difficult with the presently available capabilities to obtain high-quality absorption spectra. In the case of dilute

11.52

11.55

11.58

11.61

11.64

Photon energy (keV)

FIG.7. Variations in the near-edge fine structure at the Pt L, edge of a Pt/zeolite catalyst during CO oxidation exhibiting chemical oscillations [adapted from Hagelstein et al. (45)].

I n Situ CATALYST CHARACTERIZATION

329

specimens, fluorescence detection is necessary; however, there is no satisfactory way to collect such data by using the DEXAFS setup. A special technique, quick-EXAFS or QEXAFS, has been developed to reduce the time required to collect an EXAFS spectrum. Frahm (46, 47) found that, by use of a mechanically stable double-crystal monochromator, high-quality spectra can be recorded by moving the monochromator in a continuous manner instead of moving it in discrete steps. The reasoning behind this procedure is that in most cases it is not counting statistics that determine the time needed to collect a spectrum. A large fraction of the time between each step is spent on accelerating and decelerating the monochromator turntable and on waiting for the mechanical vibrations in the whole monochromator system to damp out. It is clear that, by a continuous movement of the monochromator, these problems are largely eliminated and the time to record a spectrum is significantly reduced. Another limitation may be the dead time introduced by the data acquisition system. Typically, the time needed to access the computer to transfer and subsequently store the information of each data point takes only fractions of a second for most computers. For a spectrum with about 1000 data points, however, this may add up to several seconds, irrespective of the way the monochromator is moved. These computer dead time problems can largely be avoided either by optimizing the computer input routines (48) or by utilizing a multichannel analyzer which is read out by the computer only after each complete scan (49). As shown in more detail in Section VII, scan times of the order of a second for a full spectrum (i.e., one covering about 1000 eV) can now be obtained with the QEXAFS technique. The limit in time resolution of the QEXAFS technique is presently determined by the capabilities of the stepping motors and the mechanical stability of the monochromators, which put strict limits on how fast one can repeatedly accelerate and decelerate the monochromator turntable between the QEXAFS scans (46-49). A promising solution to overcome this limitation has been given by Frahm (50). Instead of moving the whole turntable to position the Bragg angle of the monochromator crystal, the turntable is instead kept fixed and the positioning is realized by use of tilt tables actuated by piezoelectric crystals to rotate the monochromator crystal. In the case of a double-crystal monochromator, it is necessary to synchronously move the two piezoelectric crystals turning both monochromator crystals. Since piezoelectric crystals can operate at extremely high frequencies, it is possible with this piezo-EXAFS technique to reach a time resolution of the order of milliseconds for a spectrum (50, 51). Although there are still problems that need to be solved regarding the well-known hysteresis effects in the piezos, their dynamic range, etc., the technique has

330

BJERNE S. CLAUSEN, HENRIK TOPSOE, AND RONALD FRAHM

. c-

-c

0

v

10.0

nn -4.0

-2.0

0.0

2.0

4.0

Piezo voltage [V]

FIG. 8. Piezo-QEXAFS spectra at the Cu K-edge recorded during reduction of a Cu/ ZnO/Alz03 methanol synthesis catalyst. The recording time was 50 mskcan, and only every 40th scan is shown (51).

already proven its potential in ultrafast measurements of time-dependent phenomena in catalysts (Fig. 8). The fact that much shorter collection times are possible by using DEXAFS, QEXAFS, or piezo-QEXAFS means that many advanced dynamic phenomena can now be studied and one does not have to be concerned with problems in connection with sample changes during the measurements.

V.

Combined EXAFS/XRD Methods

To date, mainly two groups have been responsible for the development of the combined EXAFWXRD methods, namely the group of Clausen, TopsQe,Niemann, and co-workers, who developed the combined EXAFS/ XRD and QEXAFWXRD techniques, and the group of Thomas, Greaves, and co-workers, who developed the DEXAFWXRD method. In the following, some of the developments are discussed briefly. For chemically stable systems for which time resolution is not needed, Clausen et al. (52) obtained excellent spectral quality by combining normal EXAFS with XRD simply by recording the two data sets sequentially for the catalyst sample in the same setup. In this approach (Fig. 9) a diffractometer was mounted between the first and second ionization chambers in a standard EXAFS spectrometer setup, and the EXAFS was then recorded by step-scanning the monochromator through the energy region of interest (33).The XRD pattern can be acquired at the most convenient wavelength,

In Siru CATALYST CHARACTERIZATION

331

ization chamber 2

FIG.9. Schematic drawing of the apparatus for combined QEXAFWXRD in situ experiments [adapted from Clausen et al. (33)].

typically at the start of the EXAFS scan, i.e., at an energy less than that of the absorption edge to minimize absorption and to avoid a high background resulting from the fluorescence radiation. This setup provides high-quality EXAFS and XRD data characterizing the same sample. By employing special catalyst cells, in situ experiments are also possible (Section 111). This approach has the advantage of being relatively easy to implement at existing EXAFS spectrometers. This research group also evaluated the DEXAFS technique (44) but concluded that because of the drawbacks of this technique in investigations of catalysts (Section IV), it was more advantageous to focus on the QEXAFS technique. The group of Thomas (53) showed, however, that qualitative data can be obtained by using the DEXAFS technique in combination with XRD (Section VI). In this approach (Fig. lo), the XRD patterns were measured alongside the DEXAFS by using a curved position-sensitive detector placed vertically above the sample stage at the focus of the dispersive monochromator. Because a large band-pass is required for the absorption spectra whereas a narrow band-pass is needed for high-quality XRD patterns, a movable slit is necessary for acquisition of the two types of data. This requirement obviously limits the time resolution of the combined setup. Nonetheless, this type of information may in many cases be sufficient to give an indication of the changes in catalyst structures. However, the Thomas group (54) later also adopted the combined QEXAFWXRD tech-

332

BJERNE S. CLAUSEN, HENRIK TOPSBE, AND RONALD FRAHM Bent Si(ll1) energy-dispersing crystal

Synchrotron

source

9

Sample in environmental cell

Cryostat

FIG. 10. Apparatus for combined DEXAFSlXRD in situ experiments. [Adapted with permission from Nature, Ref. 53. Copyright 1991 Macmillan Magazines Limited.]

nique first developed by Clausen et al. (33) and described in Section VI. The use of QEXAFS provided much improved absorption spectra relative to those obtained with DEXAFS without sacrificing much in time resolution. It was even possible to obtain highly accurate quantitative information characterizing the phase changes during reduction of a methanol synthesis catalyst (Section VI). To take advantage of the time resolution of the QEXAFWXRD technique, it is important to have in situ cells which do not have undesired gradients or inhomogeneities and which allow rapid changes in the conditions. The in situ capillary appears to offer advantages in this regard because changes taking place in narrow temperature or time ranges can be investigated (Section VI).

VI.

Examples of in Situ Combined EXAFS/XRD Investigations

Although combined EXAFS/XRD has been found to be very useful in studies of dynamic phenomena in catalysts, only a limited number of experiments have been reported so far. The difficulty in designing and constructing in situ cells that can be used both in XRD and in X-ray absorption spectroscopy is probably the most important limitation hindering more widespread use. Here, we briefly review some of the work performed with the combined EXAFS/XRD techniques. It has recently been shown (34) that the commonly used EXAFS analysis procedures [Eq. ( 3 ) ] may introduce large errors in the determination of the coordination numbers of small particles. The fact that the motion of the atoms in small particles is highly anharmonic, especially at high tempera-

333

In Situ CATALYST CHARACTERIZATION

tures, gives rise to asymmetric pair distribution functions with broad tails (55). This part of the pair distribution function contributes to the lowwavenumber part of the EXAFS spectrum, which normally cannot be included in the EXAFS data analysis. The absence of low momentum transfer information and the nontransferability of amplitude functions between the reference bulk material and the small particles have the effect that the coordination numbers obtained by data fitting are systematically too small. An improved procedure was suggested on the basis of molecular dynamics simulations, and the combined EXAFS/XRD method was used to verify experimentally the new procedure (34, 56). The measurements were performed by using silica-supported copper catalysts, and the fact that the same catalyst samples were measured under identical conditions ensured that the two sets of data indeed resulted from identical structures. The EXAFS spectra showed the presence of small metallic copper particles and the mean particle size estimated from the apparent coordination numbers was between 10 and 12 A for the samples investigated. A line profile analysis of the Cu(ll1) diffraction line revealed an average XRD crystallite size of 30-35 A based on the Scherrer formula [Eq. (2)]. Thus, the standard EXAFS analysis exhibits systematically smaller particle sizes compared to the XRD technique. The discrepancy in particle sizes can be quantitatively accounted for by applying the new procedure (Fig. 11) for correcting the EXAFS coordination numbers obtained from the standard analysis. The

12

t

.....)....

a

5 c

11

280K

-

0 .c tu .-c

z

0 0

lo

c

c

$

2a

9 9

10

11

12

True coordination number FIG.11. Relationship between the apparent first-shell coordination number determined from standard EXAFS analysis and the true coordination number derived from the structures obtained from molecular dynamics simulations for copper particles [adapted from Clausen et al. ( 3 4 ) )

334

BJERNE S. CLAUSEN, HENRIK TOPSBE, AND RONALD FRAHM

XRD diagram corresponding to the corrected EXAFS particle sizes (solid line in Fig. 12) is in good agreement with the experimentally obtained XRD results. As an example of the potentials of the combined QEXAFS/XRD method for in situ, time-resolved studies, the information that can be obtained regarding the solid-state transformations taking place during activation (reduction) of a Cu/ZnO/A1203methanol synthesis catalyst is presented (33, 57). A three-dimensional plot of the raw QEXAFS data recorded in situ during reduction is shown in Fig. 13a. The reduction of the copper oxide phase takes place within an extremely narrow temperature range. The spectra typical of CuO are observed to change within a few degrees Kelvin into spectra that are typical of metallic copper. In Fig. 13b, the simultaneously recorded XRD diffractograms are shown for the 2 0 angular region where the Cu(ll1) peak appears. The integrated intensity of this peak as a function of the reduction temperature is shown in Fig. 14. The changes in the QEXAFS occurring during reduction are illustrated by plotting the intensity of a specific feature (9040 eV) in the absorption spectrum as a function of the reduction temperature (Fig. 14). A comparison of the curves for XRD and QEXAFS reveals that QEXAFS detects copper metal at a slightly lower temperature than the XRD. This comparison indicates that very small copper clusters are formed initially. They can be

_____

Standard EXAFS analysis

- New procedure

35

37

39

41

43

20 (degree)

FIG. 12. Cu(ll1) diffraction line from a reduced Cu/Si02 catalyst acquired by using the combined XRDiEXAFS setup. The dashed line shows the calculated broadening of the Cu(ll1) peak corresponding to particles with the mean size determined by use of the standard EXAFS analysis. The full line shows the results when the new procedure ( 3 4 ) was used to estimate the copper particle size [adapted from Clausen et al. ( 3 4 ) ] .

I n Situ CATALYST CHARACTERIZATION

335

9000

Energy (eV) C1.I 1

I

37

38

39

40

1 1:

41

2 0 (degree)

FIG.13. (a) Raw QEXAFS data near the Cu K-edge obtained in ~ i t uduring reduction of a Cu/Zn0/A1203catalyst. The recording time for each spectrum was 120 s. (b) I n sifu XRD patterns of the Cu(ll1) line recorded simultaneously with the QEXAFS data in (a). The recording time for each X R D patlern was 90 s [adapted from Clausen et al. (3.?)].

observed by EXAFS spectroscopy, but they are still too small to bc observed with XRD (i.e., the clusters appear to be X-ray amorphous). The fact that the intensities of the curves for QEXAFS and for XRD change smoothly from onc level to another without any additional features shows that the reduction of CuO to metallic copper does not involve the presence of an intermediate phase, such as Cu20, in any appreciable amount. The combined EXAFSiXRD investigations of Clausen et al. ( 5 8 )characterizing the gas phase-induced wettinginonwetting phenomena o f metallic copper particles on ZnO illustrate that dynamic, morphological changes may take place during catalysis. Figure 15a shows the Fourier-transformed

336

BJERNE S. CLAUSEN, HENRIK TOPSBE, AND RONALD FRAHM

0

1

QEXAFS (9040eV) XRD (Cu(ll1))

I

4 0

,*= !-{-

,

,

434

436

0.

428

430 432 Temperature (K)

FIG. 14. Changes in the X-ray absorption feature at 9040 eV (open squares, scaled to coincide with the XRD data at low and high temperatures) and in the integrated intensity of the Cu(ll1) line (solid circles) occurring during reduction of a Cu/Zn0/AI2O1 catalyst [adapted from Clausen et al. (33)].

EXAFS spectra above the Cu K-edge of a Cu/ZnO catalyst. The spectra, which show the typical features of relatively small metallic copper particles, were obtained in situ during synthesis conditions at ambient pressure and in the gas mixtures indicated in the figures. The changes in the Cu-Cu coordination number upon changing the oxidation potential of the synthesis gas (Fig. 15b) were interpreted in terms of a change in the particle dispersion associated with a wetting/nonwetting phenomenon of the small copper particles on the ZnO support. An increase in the oxidation potential (wet conditions) gave rise to a nonwetting condition; i.e., the copper particles become spherical, giving relatively high coordination numbers. The subsequent lowering of the water partial pressure (dry conditions) results in wetting of the support; the copper particles assumed a disk-like shape, resulting in relatively low coordination numbers, i.e., relatively many low-coordinated surface atoms. Because the copper particles are so small, the simultaneously recorded XRD patterns did not provide unambiguous support for this interpretation. On the other hand, a simple explanation for the observed results could be provided on the basis of a calculation of the relative surface and interface free energies of a particle. The description of the change in particle morphology is based on the Wulff-constructed particles (58). This gives the surface area and distribution of surface planes for a particle as a function of the contact surface free energy between particle and support. The change in contact surface free energy is related to the changes in the number of

337

In Situ CATALYST CHARACTERIZATION

E

P fn C

0.1

0.0

Distance (A) 4.9% 4.9%

co, cop,

4.9%

co,

cop,

10%

5% cop, 90% H p

I

1

90%Hp

2

I

I

I

I

I

3

4

5

6

7

Experiment number FIG.15. (a) Fourier transforms of the in situ EXAFS spectra above the Cu K-edge for a Cu/ZnO catalyst under the influence of the gaseous environments stated in the figure. (b) Variation in the apparent Cu-Cu coordination number with changes in the gaseous environment [adapted from Clausen et al. (SS)].

oxygen vacancies at the interface between the copper particle and the ZnO support (59). In terms of catalytic kinetics, the implications of the dynamic changes in catalyst morphology during methanol synthesis are dramatic. Figure 16a shows the agreement between the predictions of a static microkinetic model and the measured rates of methanol synthesis catalyzed by Cu/ZnO/A1203

338

BJERNE S. CLAUSEN, HENRIK TOPSBE, AND RONALD FRAHM

1

2

3

4

5

6

7

a

7

a

Calculated rate of MeOH synthesis (x 10-6 rnol/s/g catalyst)

a 7 -

b

6 5 4 3 2 1 -

1

2

3

4

5

6

Calculated rate of MeOH synthesis (x 10-6 rnol/s/g catalyst) FIG.16. (a) Comparison of the calculated rate with the measured rate of methanol synthesis catalyzed by Cu/ZnO/Al,O,. The calculated rate was obtained from a static microkinetic model. (b) The corresponding comparison estimated by use of a dynamic microkinetic model [adapted from Ovesen et al. (59)].

(59). In general, the microkinetic model is able to explain the magnitude of the rate quite well, but the data are grouped into two families. The two families of points in Fig. 16a are characterized by having different reduction potentials, and it is likely that the difficulty in obtaining agreement by using the static model is a consequence of the fact that the assumption of a constant number of catalytically active sites is not valid. Thus, a dynamic microkinetic model of the methanol synthesis was developed (59) that accounts for the changes in particle shape (and surface area) with changes

In Situ CATALYST CHARACTERIZATION

339

in the gaseous environment and also takes into account that the rate of methanol synthesis is different on each of the three low-index facets (59). The result of using this dynamic microkinetic model is shown in Fig. 16b. A much better representation of the data is obtained. Thomas et al. (53, 54) investigated the decomposition and subsequent reduction of aurichalcite and the ZnO-catalyzed solid-state conversion of ion-exchanged zeolite into cordierite. The decomposition of aurichalcite into mixed oxides of copper and zinc and the subsequent reduction of the

FIG. 17. In sifu combined DEXAFWXRD measurements recorded during the thermal reduction of a CuO :ZnO mixed oxide: (a) DEXAFS spectra; (b) diffraction data. [Adapted with permission from Nature, Ref. 53. Copyright 1991 Macmillan Magazines Limited.]

340

BJERNE S. CLAUSEN, HENRIK TOPSBE, AND RONALD FRAHM

metal oxide species to produce ZnO-supported metallic copper particles were followed by using combined DEXAFWXRD (Fig. 17). The quality of the absorption spectra was limited by sample inhomogeneities, making a quantitative analysis difficult. Changes in the catalyst structure were clearly observed but curvature in the background not originating from the EXAFS was also evident. VII. Outlook At the present time much effort is being devoted to tailor-making of new nanomaterials with specific catalytic properties. In this quest for constantly decreasing the dimensions of the catalytically active components, one will unavoidably encounter materials that will be partly or completely X-ray amorphous. The present review has shown that the combined EXAFS/ XRD techniques are uniquely well suited for providing the necessary structural understanding. Thus, in view of the trend in catalyst technologies and advances in technique developments, the application of the combined techniques will no doubt play an increasing role in future catalyst characterization efforts. We now briefly discuss some likely applications and technique developments which involve the X-ray techniques discussed presently.

FIG.18. Atomic structure of a small binary metal particle of a Ni-Au alloy as predicted from Monte Carlo simulations (60).

341

I n Situ CATALYST CHARACTERIZATION

a

0 0

2

4

6

8

Interatomic distance (A) 4.5

4.0 Y

25, d

.4!

3.5

m r v 1 3 $ ’ 2 3.0 EE? .=

-g o

2.5

cn c

2.0

.as o! c

-

1.5

0

100

200

300

400

500

Temperature (“C) FIG. 19. Combined QEXAFS and temperature-programmed sulfiding results of a Mo/ AI2O3catalyst during sulfiding in a H2S/Ar gas mixture: (a) Fourier transforms of the in situ EXAFS spectra above the Mo K-edge; (b) variation in the H2S concentration in the gas outlet from the in situ EXAFS cell as simultaneously recorded by a mass spectrometer (61).

Small particles of binary alloys have been investigated in detail in static EXAFS experiments, but if information about the dynamic behavior of the alloy composition and the segregation phenomena is desired, timeresolved combined EXAFSlXRD studies are necessary. Figure 18 shows the atomic structure of a small binary particle of a Ni-Au alloy as predicted from Monte Carlo simulations (60). Ni and Au do not form a miscible alloy in the bulk but can form a stable alloy at the surface. The structural and chemical changes that occur when such particles are exposed to different

342

BJERNE S. CLAUSEN, HENRIK TOPSBE, AND RONALD FRAHM

catalytic process conditions can be elucidated in a rather unique way by such time-resolved studies (60). Apart from combining EXAFS with XRD, it is also of interest to combine one or both of the X-ray techniques with other characterization techniques, such as thermal analysis techniques [temperature-programmed reaction (TPR), scanning calorimetry, etc.]. Figure 19 shows preliminary results from combined QEXAFS and temperature-programmed sulfiding. The combined technique was used to investigate the sulfiding mechanism of a Mo/ A1203hydrotreating catalyst (61). Figure 19a shows QEXAFS spectra at the Mo K-edge recorded during the sulfiding. The transformation from a structure in which molybdenum is surrounded by oxygen atoms to a MoS2like structure is clearly seen. In Fig. 19b the corresponding H2S mass spectrometer profile is shown and both the consumption and production of H2S can be correlated with specific features in the absorption spectra. Improved time resolution in the millisecond regime and lower noise as a result of the increased X-ray intensity of the new synchrotron radiation facilities undoubtedly will result in many more exciting developments in elucidating dynamic phenomena in catalysis. ACKNOWLEDGMENTS T. B. Zunic is gratefully acknowledged for helpful discussions during preparation of the manuscript. The authors are also grateful to HASYLAB, Deutsches Elektronen-Synchrotron, Hamburg, Germany, for making available the beam time needed to develop the time-resolved techniques.

REFERENCES 1. 2. 3. 4.

5. 6.

7. 8.

9.

10. 11.

Jellinek, M. H., and Frankuchen, I., A d v . Cutul. 1,257 (1948). Van Nordstrand, R. A,, A d v . Cutul. U,149 (1960). Sayers, D. E., Stern, E. A,, and Lytle, F. W., Phys. Rev. Lett. 27, 1204 (1971). Bart, J. C. J., A d v . Cutul. 34, 203 (1986). Bart, J. C. J., and Vlaic, G., A d v . Cutul. 35, 1 (1987). Pettifer, R. F., in “Characterization of Catalysts” (J. M. Thomas and R. M. Lambert, eds.), p. 264. Plenum, New York, 1980. Lytle, F. W., Via, H., and Sinfelt, J. H., in “Synchrotron Radiation Research” (H. Winick and S. Doniach, eds.), Chapter 12. Plenum, New York, 1980. Gallezot, P., in “Catalysis Science and Technology” (3. R. Anderson and M. Boudart, eds.), Vol. 5. Springer-Verlag, Berlin, 1984. Prins, R., and Koningsberger, D. C., in “X-ray Absorption: Principles, Applications, Techniques of EXAFS, SEXAFS and XANES” (D. C. Koningsberger and R. Prins, eds.), p. 321. Wiley, New York, 1988. Niemantsverdriet, J. W., “Spectroscopy in Catalysis: An Introduction,” Chapter 6. VCH, Weinheim, 1993. Warren, B. E., and Averbach, B. L., J. Appl. Pbys. 21,595 (1950); 23,497 (1952).

In Siru CATALYST CHARACTERIZATION

343

12. Hosemann, R., Preisinger, A., and Vogel, W., Ber. Bunsen ges. Phys. Chem. 70,797 (1966). 13. Klug, H. P., and Alexander, L. E., “X-Ray Diffraction Procedures.” Wiley, New York, 1974. 14. Rehr, J. J., Mustre de Leon, J., Zabinsky, S. I., and Albers, R. C., J. Am. Chern. Soc. 113, 5135 (1991). 15. Gurman, J., Binsted, N., and Ross, I., J. Phys. C 17, 143 (1986). 16. Clausen, B. S., and Topsee, H., in “X-Ray Absorption Fine Structure for Catalysts and Surfaces” (Y. Iwasawa, ed.), Chapter 7.5. World Scientific, Singapore, 1996. 17. Topsee, H., Clausen, B. S., and Massoth, F., in “Catalysis Science and Technology” (J. R. Anderson and M. Boudart, eds.), Vol. 11. Springer-Verlag, Berlin, 1996. 18. Topsee, H., in “Proceedings of the NATO Advanced Study Institute on Surface Properties and Catalysis by Non-Metals: Oxides, Sulfides and other Transitions Metal Comopunds” (J. P. Bonnelle, B. Delmon, and E. G. Derouane, eds.), p. 329. Reidel, Dordrecht, The Netherlands, 1983. 19. Clausen, B. S., Topsee, H., Candia, R., Villadsen, J., Lengeler, B., Als-Nielsen, J., and Christensen, F., J. Phys. Chem. 85, 3868 (1982). 20. Nandi, R. K., Pitchai, R., Wong, S. S., Cohen, J. B., Bunvell, R. L., Jr., and Butt, J. B., J. Card. 70, 298 (1981). 21. Shiryaev, P. A,, Kushnerev, M. Ya., Shashkin, D. P., and Krylov, 0. V., Kinet. Karal. 23(5), 1280 (1982). 22. Topsere, H., and Villadsen, J., unpublished. 23. Nix, R. M., Rayment, T., Lambert, R. M., Jennings, J. R., and Owen, G., J. Catal. 106, 216 (1987). 24. Mamott, G. T., Barnes, P., Tarling, S. E., Jones, S. L., and Norman, C. J., Powder Diffr. 3,234 (1988). 25. Ritter, J. J., Powder Diffr. 3, 30 (1988). 26. Jung, H., and Thomson, W. J., J. Card. 128, 218 (1991). 27. Herzog, B., Ilkenhans, Th., and Schlogl, R., Fresenius’ J. Anal. Chern. 345,247 (1994). 28. Clausen, B. S., Lengeler, B., Candia, R., Als-Nielsen, J., and Topsee, H., Bull. Soc. Chirn. Belg. 90, 1249 (1981). 29. Dalla Betta, R. A., Boudart, M., Foger, K., Loffler, D. G., and Sanchez-Arrieta, J., Rev. Sci. Instrum. 55, 1910 (1984). 30. Kampers, F. W. H., Maas, T. M. J., van Grondelle, J., Brinkgreve, P., and Koningsberger, D. C., Rev. Sci. Instrum. 60(8), 2635 (1989). 31. Clausen, B. S., and Topsee, H., Catal. Today 9,189 (1991). 32. Clausen, B. S., Steffensen, G., Fabius, B., Villadsen, J., Feidenhans’l, R., and Topsee, H., J. Caral. 132, 524 (1991). 33. Clausen, B. S., Gribzk, L., Steffensen, G., Hansen, P. L., and Topsae, H., Card Lett. 20,23 (1993). 34. Clausen, B. S., Gribzk, L., Topsae, H., Hansen, L. B., Stoltze, P., Nerskov, J. K., and Nielsen, 0. H., J. Caral. 141, 368 (1993). 35. Niemann, W., Bak Hansen, J. H., Clausen, B. S., Fastrup, B., Fabius, B., Villadsen, J., Wroblewski, Th., and Topsee, H., in “2nd European Conference on Progress in X-ray Synchrotron Radiation Research” (A. Balerna, E. Bernieri, and S. Mobilio, eds.), Vol. 25, p. 575. SIF, Bologna, 1990. 36. Rashkin, J. A., and Pierron, E. D., J. Caral. 6, 322 (1966). 37. Krylov, 0. V., and Margolis, L. Ya., Int. Rev. Phys. Chem. 3, 305 (1983). 38. Morozova, 0. S., Kryukova, G. N., Ziborov, A. V., and Plyasova, L. M., J. Catal. 144, 50 (1993). 39. Herzog, B., Herein, D., and Schlogl, R., Appl. Catal. A: Gen. 141, 71 (1996).

344

BJERNE S. CLAUSEN, HENRIK TOPSBE, AND RONALD FRAHM

40. Genvard, L., M ~ r u pS., , and Topsere, H., J. Appl. Phys. 47(3), 822 (1976). 41. Matsushita, T., and Phizackerley, R. P., Jpn. J. Appl. Phys. 20, 2223 (1981). 42. Lee, J. M., Paesler, M. A,, Sayers, D. E., and Fontaine, A,, Physica B (Amsterdam) 158B, 52 (1989). 43. Sayers, D. E., Bazin, D., Dexpert, H., Jucha, A., Dartyge, E., Fontaine, A., and Lagarde, P., in “EXAFS and Near Edge Structure 111” (K. 0. Hodgson, B. Hedman, and J. E. Penner-Hahn, eds.), p. 209. Springer-Verlag, Berlin, 1984. 44. Hagelstein, M., Cunis, S., Frahm, R., Niernann, W., and Rabe, P., in “2nd European Conference on Progress in X-Ray Synchrotron Radiation Research” (A. Balerna, E. Bernieri, and S. Mobilio, eds.), Vol. 25, p. 407. SIF, Bologna, 1990. 45. Hagelstein, M., Ressler, T., Hatje, U., and Forster, H., to be published. 46, Frahm, R., Nucl. Instrum. Methods Phys. Res., Sect. A 270, 578 (1988). 47. Frahrn, R., Rev. Sci. Instrum. 60,2515 (1989). 48. Frahm, R., Doctoral Dissertation, University of Rostock (1992). 49. Als-Nielsen, J., Grubel, G., and Clausen, B. S., Nucl. Instrum. Methods Phys. Res., Sect. B 97, 522 (1995). 50. Frahm, R., Synchrotron Radial. News 8(2), 38 (1995). 51. Frahm, R., Clausen, B. S., Molenbroek, A. M., and Steffensen, G., in “Proceedings of the 9th International Conference on X-Ray Absorption Fine Structure, Grenoble, France, 1996,” Abstr. No. 630 (1996). 52. Clausen, B. S., GrBbzk, L., Steffensen, G., and Topsee, H., in “HASYLAB Annual Report,” p. 495 (1991). 53. Couves, J. W., Thomas, J. M., Waller, D., Jones, T. H., Dent, A. J., Derbyshire, G. E., and Greaves, G. N., Nature (London) 354,465 (1991). 54. Thomas, J. M., and Greaves, G. N., Catal. Lett. 20,337 (1993); Sankar, G., Wright, P. A,, Natarajan, S., Thomas, J. M., Greaves, G. N., Dent, A. J., Dobson, B. R., Ramsdale, C . A., and Jones, R. H., J. Phys. Chem. 97,9550 (1993). 55. Hansen, L. B., Stoltze, P., Ngrskov, J. K., Clausen, B. S., and Niernann, W., Phys. Rev. Lett. 64,3155 (1990). 56. Clausen, B. S., Topsee, H., Hansen, L. B., Stoltze, P., and Norskov, 3. K., Catal. Today 21, 49 (1994). 57. Clausen, B. S., and Topsrae, H., Synchrotron Radiut. News 7, 32 (1994). 58. Clausen, B. S., Schiotz, J., Gr%brek,L., Ovesen, C. V., Jacobsen, K. W., NWskov, J. K., and Topsoe, H., Top. Curd 1, 367 (1994). 59. Ovesen, C. V., Clausen, B. S., Schietz, J., Stoltze, P., Topsee, H., and N~rskov,J. K., J. Catal. 168, 133 (1997). 60. Molenbroek, A. M., Steffensen, G . , Hyldtoft, J., Clausen, B. S., and NGrskov, J. K., to be published. 61. Clausen, B. S., Pleth Nielsen, L., Molenbroek, A. M., and Topsoe, H., to be published.

ADVANCES IN CATALYSIS, VOLUME 42

Present State of the Art and Future Challenges in the Hydrodesulfurization of Polyaromatic Sulfur Compounds D. DUAYNE WHITEHURST, TAKAAKI ISODA, AND ISAO MOCHIDA Institute of Advanced Material Study Kyushu University Fukuoka 816, Jnpan

1.

Introduction

The field of hydrodesulfurization (HDS) has been the subject of numerous past and recent reviews (1-10). One wonders why another is needed at the present time. However, the urgency of understanding one aspect of this area, the deep desulfurization of middle-distillate oils, calls for a review that brings together the composition and chemistry of the organic species that are difficult to desulfurize (mostly substituted dibenzothiophenes), the mechanisms involved, and the detailed chemistry of the surfaces of available catalysts. By summarizing this information in a systematic way, it is hoped that new approaches can be suggested that will help industry to meet the demands of upcoming legislation that will limit the salability of fuels with more than 0.05% S. Because of their high energy densities and convenient physical form, petroleum products are presently being consumed in vast quantities, and this consumption continues to grow at alarming rates. Inevitably, this high consumption is having a major impact on the global environment. Most notably, transportation fuels, the major petroleum products, are receiving the highest scrutiny, because the pollution from exhaust gas is difficult to control within vehicles as space is limited and because of the difficulty of continually monitoring emissions on individual vehicles. Pollutants of major concern in exhaust gases include SO,, CO, NO,, particulates, trace elements, olefins, aromatic hydrocarbons, and their interaction products. One key feature in environmental control is the level of SO, in the exhaust as it contributes to acid rain, poisons catalysts in catalytic converters, is a major fraction of particulates, and is an integral component in the cycle of atmospheric gas chemistry that leads to ozone production and smog. 345 Copyright 0 1998 by Academic Press. All rights of reproduction in any form reserved. 0360-0564/98 $25.00

346 D. DUAYNE WHITEHURST, TAKAAKI ISODA, AND ISAO MOCHIDA Although it is possible to improve catalytic converters in gasolinepowered vehicles and to develop exhaust control devices for diesel-fueled vehicles and even aircraft, development in these areas is slow. Extensive research efforts in these areas are desperately needed and should be strongly encouraged. However, in reality, implementation of these devices will be far in the future and the problems of today are increasing. For these reasons, the burden of environmental protection is falling on the manufacturers of the fuels to produce products that will have minimal impact on the environment, and this translates to very low levels of sulfur in fuels. So-called “deep refining” of petroleum products is thus required, although the extent of this refining may vary, depending on the type of fuel and its content of contaminants. Gasoline-fueled vehicles have been greatly improved in the past few decades with the development and installation of three-way catalytic systems. Lean burning in gasoline engines, which would improve thermal efficiency, will require considerably more improvements in three-way catalysts. Diesel engines, on the other hand, have no exhaust gas cleanup equipment and the potential for future installation still appears uncertain. Accordingly, in the past decade, research has been focused on the deep refining of middle distillates or gas oils, especially deep desulfurization. The U.S. Clean Air Act formulated very severe regulations on petroleum products such as gasoline and gas oil to protect the urban atmosphere, because further improvements in vehicles were becoming limited. Influenced by this Clean Air Act, the Japanese government became concerned about the degradation of the environment in the early 1970s. Prior to the rest of the world, Japan imposed severe regulations on fuel composition. Such regulations stimulated the development of several new environmental protection technologies in Japan. The Japanese Central Council for Control of Environmental Pollution proposed a schedule of phased reduction of NO, and particulates in emissions. Thus, the sulfur content of gas oil was mandated to be reduced to 0.20 wt% in 1993 and further to 0.05 wt% by 1997. It may well be that by early in the 21st century, the sulfur level in fuels will have to be reduced to 0.01 wt%. To put into perspective the magnitude of chemical transformation that will be required to meet these new standards, Fig, 1shows the gas chromatographic analyses of a typical Iranian gas oil containing 1.2% sulfur (A and C) and a hydrotreated gas oil containing 0.25% sulfur (B and D). Figures 1A and 1B present the total hydrocarbon distribution as detected by flame ionization detectors (FID). Figures 1C and 1D show the analyses of the sulfur species in these gas oils as detected by sulfur-specificflame photometric detectors. From these figures it can be clearly seen that the number of sulfur species is relatively few and that, on hydrotreating, some species are

1.2%S - Gas Oil

0.25% S - Gas Oil

(A) FID

(B) FID

b U Refractory S-Compounds

U Refractory S-Compounds

(D) FPD

__

FIG.1

Compositions of gas oils hydrotreated to different levels.

348 D. DUAYNE WHITEHURST, TAKAAKI ISODA, AND ISAO MOCHIDA easily removed but several species appear to be refractory. To reduce the sulfur content of the present 0.25% sulfur gas oils to the new 0.05% sulfur standards, 75% of these refractory species will have to be desulfurized, preferably without altering any of the other components of the gas oil, which constitute more than 98% of the total gas oil. This will be the challenge for newly evolving catalysts and processes. The Japanese refining industry responded to the challenge of meeting the 1993 standards by installation of new capacity. In an attempt to meet the more severe 1997 standards, industrial, governmental, and academic laboratories have conducted a great deal of research to identify means to achieve these new regulations with minimal increases in the cost of processing. The U.S. and Japanese refining industries have been able to satisfy the present regulations for 0.05% S in diesel fuel. This was accomplished partly by developing new more highly active catalysts and partly by virtue of a fortunate economic situation in which low-sulfur crude feedstocks were reasonable low-cost alternatives. In addition, at the time of the initial changes in standards, there was also a depression in the price of construction and of high-pressure reactors. Thus, the Japanese refining industry was able to install new capacity of high-temperature, high-pressure reactors to meet the imposed standards with moderate investment. This situation may change in the future, as Europe is now beginning to face the same requirements for much lower sulfur content in diesel fuels and will also have to consider installation of new high-pressure hydrotreaters. It is expected that the availability of low-sulfur crudes may become limited in the future due to the demands of cleaner fuels in all of the developed countries. The need for the installation of new capacity throughout the world will also lead to escalation of the costs of construction of new highpressure reactors. If similar standards are imposed on even slightly heavier fuels such as home heating fuels or if the standards become even more strict for diesel fuels, industry may not be able to meet these standards with existing equipment. Thus, there is a continuing need for new HDS technology, particularly in the areas of developing more active catalysts and novel process options that can meet the new standards at lower pressures with present equipment. Gas oil refining currently operates at moderate temperatures (340360°C) and hydrogen pressures of 3.0-5.0 MPa, usually with CoMo/A1203 catalysts. These conditions are unable to achieve the 0.05% S specification of gas oil. Increasing the temperature can achieve the S goal, but the color of the produced oil becomes degraded at the elevated temperature due to unwanted side reactions. Consequently, it is necessary to understand the HDS reactivity of refractory sulfur species in the gas oil and to clarify possible inhibition mechanisms that may limit HDS reactions. The design and development of more effective HDS catalysts that will minimize the increase of the costs to the refinery are severely needed.

POLYAROMATIC SULFUR COMPOUNDS

349

The present review summarizes contemporary views of the problems, achievements, and prospects involved in the deep desulfurization of gas oils, including identification and reactivity of sulfur species in the feed, the reaction pathways and mechanisms, activity and selectivity of the conventional catalysts, and concerns of fluorescence color production. Process schemes and guidelines for the development of the next-generation catalysts for improved deep desulfurization technology based on these discussions are also proposed. The structure and nature of the active sites of current catalysts will not be extensively covered in this review, because several excellent reviews have been published on these subjects within the past two years (1-3). Some major problems associated with deep desulfurization of gas oil are as follows:

1. In a practical HDS process for gas oil, both aromatic species existing in the feed and various types of sulfur compounds compete for the active sites on the catalyst surface. Moreover, H2S and some other hydrocarbons produced in the early stages of the desulfurization appear to inhibit the HDS of the less reactive sulfur species. The reactivities of refractory sulfur compounds and the effects of inhibitors in gas oils need to be fully understood for the development of an improved economical desulfurization process. 2. The reaction mechanisms for conversion of the refractory sulfur species such as 4,6-dimethyldibenzothiophene (4,6-DMDBT) in gas oil and the inhibition of their HDS by coexistent species need clarification. Novel mechanistic routes that enhance the reactivity or reduce the inhibition are necessary for designing more efficient deep HDS catalysts and processes. 3. The reactivities toward HDS of various sulfur species in the various stages of their respective reaction pathways need to be described from the viewpoint of their molecular and electronic structures, or quantum chemical descriptions, when the HDS reactivities of the important sulfur species become available. 4. To meet the demand for low-sulfur gas oil, guidelines for the design of new catalysts for such extensive desulfurization while maintaining product quality and minimizing the increase in costs to refineries are urgently needed .

II. Description of Systematic Approach for Describing the HDS Phenomenon

Past reviews have described in detail the chemistry and mechanisms involved in the hydrodesulfurization of the most reactive organic species found in petroleum fuels (1-5). A vast amount of knowledge has been

350 D. DUAYNE WHITEHURST, TAKAAKI ISODA, AND ISAO MOCHIDA amassed in this area because hydrotreating is perhaps the largest volume process in today’s commercial refineries. From the outset, the present authors apologize for any omissions of references to contributions to this important area by past researchers. The most comprehensive review to date is clearly the contribution by Topsoe, Clausen, and Massoth ( I ) , and the reader is referred to that review for additional information. The purpose of the present review is to complement this bank of knowledge with emphasis specifically on the most refractory sulfur-containing components of petroleum streams. This particular topic, though limited, has much need for clarification and focus so that those entering the field can grasp the magnitude of the problems involved. For readers to understand the specific difficulties involved, the descriptions of organic mechanisms, reaction selectivities, and the catalytic surface phenomenon will be expressed in a systematic manner. Attempts are made where possible to relate well-established organometallic catalytic mechanisms to observations with present-day heterogeneous commercial catalysts. Others have expressed a similar approach, and the reader is referred to an excellent review by Sfinchez-Delgado ( 4 ) that deals specifically with this topic. We quote from that article the following: “. . . a great deal of information has been derived from relatively simple experiments on discrete metal complexes; in this way, knowledge at the molecular level has frequently been obtained, which is rarely possible in heterogeneous and surface chemistry studies. This is of paramount importance if one remembers that any form of catalysis deals essentially with making, breaking and rearranging chemical bonds, and is therefore always a molecular process.” It is only within the past 5 years that understanding of the exact nature of the most effective catalysts for HDS has begun to be developed, and the reader is referred to recent reviews by Topsfie, Clausen, and Massoth ( I ) and Startsev (2) for detailed discussions. This topic will be discussed briefly in Section IV. For the time being, it is sufficient to say that the most effective catalysts have been described as consisting of a specific stoichiometric combination of nickel or cobalt with molybdenum or tungsten that exists as a sulfide containing one Ni or Co atom in combination with two molybdenum or tungsten atoms, chemically anchored to the surface of a molybdenum sulfide crystallite which in turn is bound to the surface of a solid support (generally alumina or silica-alumina) (2). Such combinations are referred to as “promoted” Mo or W catalysts (1-10). The active catalytic species has recently been termed Ni-Mo-S or SBMS ( 2 , I I ) . These unique species exhibit orders of magnitude higher activities than similar catalysts composed of the individual metal sulfides. Though the reason for their high activity remains a mystery and has prompted considerable debate, they have been used for purifying feedstocks for over 70 years. The classic work

351

POLYAROMATIC SULFUR COMPOUNDS

of Chianelli ( 3 ) did provide theoretical explanations for specific electronic effects of certain combinations of different metal sulfides that offer the correct balance of bond strengths for metal-S and metal-SH2. This optimal balance of bond strengths allows the overall HDS reaction to proceed in a facile manner and thus provides more active catalytic sites. In the present review, we will illustrate this species with the following abbreviated structures to aid in discussions of the detailed chemistry of the catalytic surface.

Ni/Mo/S CRYSTALLITE

~

-:st:

€IDS ACTIVE SITE H

:s

/H

/Ni\s

s 0- P

The figure on the left represents a single M022S44 crystallite with the active Ni component bound to the surface and the crystallite anchored to the alumina surface through Mo-0-A1 and Mo-S-A1 bonds. In this representation, H2S is shown dissociatively adsorbed on the Ni. The figure on the right is a shorthand notation. Organic mechanisms and pathways will be related to the chemistry occurring on such sites in terms of known homogeneous organometallic catalytic mechanisms. In describing catalytic activities and selectivities and the inhibition phenomenon, we will use a common format, where possible, which is based on a common reaction pathway scheme as outlined in Scheme 1.In contrast to the simple one- and two-ring sulfur species from which direct sulfur extrusion is rather facile, in the HDS of multiring aromatic sulfur compounds such as dibenzothiophene derivatives, the observed products are often produced via more than one reaction pathway. We will not discuss the pathways that are specific for thiophene and benzothiophene as this is well represented in the literature ( I , 5, 8, 9) and, in any event, they are not pertinent to the reaction pathways involved in deep HDS processes whereby all of the highly reactive sulfur compounds have already been completely converted. For highly substituted dibenzothiophenes, ring hydrogenation prior to sulfur extrusion is the major route to hydrocarbon production as, relative to the parent molecules, aliphatic substituents on aromatic ring carbons

352

D. DUAYNE WHITEHURST, TAKAAKI ISODA, AND ISAO MOCHIDA

adjacent to the sulfur atom impose severe steric hindrance toward bonding to the catalyst surface and to the production of appropriate intermediate species. This creates difficulties in determining individual rate constants as the same product may be produced either by hydrogenating one of the sulfur compound’s aromatic rings (kHsl) followed by desulfurization (kDl) or by removal of sulfur without ring hydrogenation (k,,,,) followed by hydrogenation of one of the phenyl rings of the desulfurized product (kHP,),as illustrated in Scheme 1. For proper kinetic evaluation of the conversion selectivity in such a case, it is necessary to find means for distinguishing the relative contributions of the various pathways to the formation of products. A systematic method for distinguishing parallel and sequential formation of products will be described, and results of HDS of such multiaromatic sulfur compounds will be discussed in these terms. For ease of understanding the following discussion, we define the various rate constants as follows. Rate constant kD0

kDl kD2 kHSi kHsi

kwl kHp,

Reaction concerned Desulfurization without ring hydrogenation Desulfurization of one ring of the hydrogenated sulfur compound Desulfurization of the fully saturated sulfur compound Hydrogenation of one ring of the sulfur compound Hydrogenation of the second ring of the sulfur compound Hydrogenation of one phenyl ring of desulfurized biphenyl Hydrogenation of the second phenyl ring of biphenyl

Rate constants such as kD represent rate constants involving sulfur removal from molecules, and rate constants such as kHP and kHS represent rate constants involving hydrogenation of an aromatic ring. The subscripts 0, 1, and 2 indicate that, for the desulfurization process, the product is a

Where R=JI, CH, or Other

SCHEME1. Reaction pathways in the HDS of dibenzothiophenes.

POLYAROMATIC SULFUR COMPOUNDS

353

molecule in which none, one, or two of the aromatic rings have been hydrogenated, respectively. The subscripts 1 and 2 for the hydrogenation processes indicate the sequence of hydrogenation of the two aromatic rings. For example, kHS,indicates the rate constant associated with the hydrogenation of the first aromatic ring. This scheme applies to all polyaromatic thiophene-based HDS conversion processes; only the rate constants are different. With the more reactive dibenzothiophenes, the rate of extraction of sulfur from the first hydrogenated intermediate (kD,) is so high that products derived from the fully saturated dibenzothiophene are not observed. Similarly, the rate of hydrogenation of single aromatic rings (kHp2)is generally much lower than those of the other reactions, and dicyclohexanes are generally produced in only trace amounts. As the parent dibenzothiophene becomes less reactive, due to alkyl substituents, the importance of kHs2,kD,, and kHp, increases, and they cannot be neglected. If used consistently, this kinetic treatment can be used to assess the effects of the degree of ring condensation and the chemical nature, the number, and the position of substitution of substituents on the aromatic rings on the activities and selectivities of catalysts. The kinetics become more complicated on acidic catalysts as alkyl isomerization, transalkylation, and dealkylation may also take place prior to hydrogenation or desulfurization. With these catalysts, these additional reaction pathways must also be included in the overall reaction network.

111.

Composition of Sulfur Species in Middle-Distillate Oils

A. GENERAL COMPOSITIONAL FEATURES OF GASOILS As discussed in the previous section, gas oils are complicated mixtures of several different hydrocarbon classes contaminated with small amounts of sulfur-containing species that lower product quality and limit marketability of the whole gas oil (see Fig. 1). The concentration of these sulfur species must be lowered significantly to meet present standards and, in the future, they must be nearly completely removed. To understand the difficulties in such conversions, it is helpful to consider the detailed composition of gas oils, from both the viewpoint of the desirable components of gas oils which are not to be converted and that of the sulfur-containing species which must be treated to extinction. With the analytical procedures available today, it is now possible to determine in fine detail the composition of all of the hydrocarbon components as well as the offending sulfur species. In this section we discuss the general features of gas oils and address the present state of knowledge of

354 D. DUAYNE WHITEHURST, TAKAAKI ISODA, AND ISAO MOCHIDA the content of different classes of sulfur species and their relative reactivity in conventional hydrodesulfurization processes. The term “gas oil” refers to a hydrocarbon liquid having a particular boiling range. There are two major types of gas oils distinguished by their boiling ranges. Atmospheric distillates have initial boiling points of 228232°C and final boiling points of 339-377°C as defined by ASTM. Higherboiling gas oils must be distilled under vacuum and have boiling ranges of about 340-530°C (vacuum gas oils). This review will concentrate on the TABLE I Properties of Gas Oils

Gas oil ~~~

~~

Japanese diesel

European LCO

Vacuum gas oil

Total sulfur (wt%) Density (g/ml at 15°C) Pour point (“C) Boiling range (ASTM) (“C) IBP 5% 10% 20% 30% 40% 50% 60% 70% 80% 90% 95% FBP Composition (FAI, ~01%) Saturates Olefins Aromatics Polars Identified sulfur types (ppm S) Alkylbenzothiophenes Alkyldibenzothiophenes Alkylphenanthrothiophenes Alkylphenylbenzothiophenes Alk ylbenzonaphthothiophenes Alkylpolyaromatic thiophenes

0.706 0.84 -10

1.007 0.859

2.45 0.926

232 258 267 276 282 287 292 297 303 311 321 330 339

228 255 265 275 282 289 296 306 316 330 350 366 377

205 341 378 410 434 450 473 494 510 530 >530 >530 >530

78 5 17

58 4.6 37.4

12

Number of compounds identified % of total sulfur identified

160 74

2500 2700

76 12 890 1630 570 80 1210

61

123 19

POLYAROMATIC SULFUR COMPOUNDS

355

composition of atmospheric gas oils as these materials are the ones for which regulations will be made extremely strict (<0.05% S or 500 ppm) in the very near future. Commercial products falling within this boiling range include transportation diesel fuels, jet fuels, and home heating fuels. The consumption of such fuels is about the same or slightly larger than that of automotive gasoline, and so the ability to produce these products with acceptable product specifications is extremely important. Some general features of some representative commercial fuels available today, i.e., a Japanese diesel fuel, a European light cycle oil (LCO), and a typical Middle Eastern vacuum gas oil (VGO), are provided in Table I. As can be seen from the table, the diesel fuel is primarily composed of saturated hydrocarbons, with aromatic hydrocarbons representing the second largest class of components. LCO contains nearly equal amounts of saturates and aromatics, and vacuum gas oils are predominantly aromatic. The sulfur-containing species represent about 5, 5 , and 19% of all the molecules in diesel fuel, LCO, and VGO, respectively. The current trend is to increasingly add more LCO derived from thermal and catalytic cracking to these fuels, which will increase the aromatic contents of these fuels. The atmospheric distillates do not contain significant amounts of other, more polar, molecules (identified as polars); however, the VGO contains about 12% of such materials. These polar compounds contain oxygen and nitrogen functional groups. Original OIL Separation by LC- 1 (Silica gel) Paraffin (OIL I )

Aromatic Single Ring (OIL 2)

Aromatic Double and Triple Rings (OIL 3

+

PASC Separation by LC-2 (PdCI2/Silica gel Ligand Exchange Chromatography)

Aromatic Double and Triple Rings (OIL 3-1)

Thiophene Single Ring (OIL 3-2)

Thiophene Double and Triple Rings (OIL 3-3)

FIG.2. Separation and analysis of polyaromatic sulfur-containing compounds (PASC). Oils 3-1, 3-2, and 3-3 were analyzed by gas chromatography-mass spectroscopy (GC-MS) and gas chromatography-atomic emission detection (GC-AED). Reproduced from Ref. 12, with permission.

GC-AED chromatogram of whole gas oil

I

/

2.0 E4

AED-S

0, 10

20

30

Time (min)

Liquid chromatogram /Fraction

of whole gas oil

*

Fraction 2

AED-C

10,000 8000

O

7000

5000

3 0 E42 0 E4

AED-S 10

20

Paraffins (61%)*

1000

30 Time (min)

10

20

20

Monoaromatics

DI- and Triaromatics

(21 %)*

(18%)*

FIG.3 . GC-AED profiles of gas oil fractions (22). Estimated percentages from LC analysis. Reproduced from Ref. 12, with permission.

POLYAROMATIC SULFUR COMPOUNDS

Compounds

357

Structures

Alkylnaphthalenes (n = 0-5)

a / . C "

Alkyltetralins (n = 1-3)

Alkylbiphenyls (n = 0-4)

Alkylfluorenes (n = 1,2)

QQcn

Alkylphenanthrenes (n = 0-3)

Dimethyl benzothiophenes

Dimethyldibenzothiophenes (n = 0-3)

4-Methyldibenzothiophene 4.6-Dirnethyldibenzothiophene

FIG.4. Typical polyaromatics in gas oil. Reproduced from Ref. 12, with permission.

B. DETAILED COMPOSITION OF MOLECULAR SPECIES IN GASOILS

With today's advanced analytical procedures, it is possible to describe the composition of these fuels in considerable detail. By combining several sequenced liquid chromatographic separations with gas chromatographymass spectroscopy and by using specific gas chromatographic detectors for sulfur compounds, it has been possible to identify the majority of individual sulfur species in some fuels (12-19). A typical separation scheme is shown

TABLE I1 Polyaromatic Hydrocarbons Identijied in Gas Oil (12) w

'A 00

Compound Naphthalene Methyltetrahydronaphthalene 2-Methylnaphthalene 1-Methylnaphthalene Dimethyltetrahydronaphthalene Dimethyltetrahydronaphthalene Dimethyltetrahydronaphthalene Dimethyltetrahydronaphthalene Dimethyltetrahydronaphthalene Biphenyl 2-Ethylnaphthalene C2-Benzothiophene 2,5- or 2,7-Dimethylnaphthalene 1,7-Dimethylnaphthalene

Measured retention index 200.0 214.4 220.2 223.2 224.6 229.0 229.5 231.5 232.2 236.4 239.0 237.7 240.5 242.8

Literature retention index 200.0 220.2 223.0

236.4 238.6 240.3 242.8

Compound C?-Naphthalenes C2-Biphenyls 9-Methylfluorene C4-Naphthalenes

2-Methylfluorene 1-Methylfluorme C2-Biphenyl C,-Naphthalene C2-Biphenyls

Measured retention index 270.0, 270.4 271.4, 272.0 273.0 273.5 275.2, 276.2 276.7, 279.5 280.4, 280.9 283.7. 286.1 287.5 288.0 288.9 289.5 290.2, 291.4 292.1, 292.9

Literature retention index

273.8

288.4 289.2

1,3- or 1,6-Dimethylnaphthalene Methylbiphenyl 1,4- or 2,3-Dimethylnaphthalene 1,5-Dimethylnaphthalene Trimethyltetrahydronaphthalene 1,2-Dimethylnaphthalene Cz-Biphenyl Trimethyltetrahydronaphthalene Methylbiphenyl C3-Naphthalenes

VI w

W

2,3,6-Trimethylnaphthalene C3-Naphthalene G-Biphenyl Fluorene

243.4 245.2 246.3 246.8 247.7 248.5 250.5 251.2 254.0 255.7. 256.3 256.9, 257.8 258.8, 259.8 260.9, 261.9 262.6, 264.0 265.0 266.3 269.0 269.7

243.3 246.0 246.9 248.5

265.1

269.7

C3-Naphthalene Dibenzothiophene C3-Naphthalene Phenanthrene C4-Biphenyl C2-Fluorene C4-Biphenyl C2-Fluorene C4-Biphenyl Methyldibenzothiophene C4-Biphenyl 3-Methylphenanthrene 2-Methylphenanthrene 9- or 4-Methylphenanthrene 1-Methylphenanthrene MW 208 C2-Dibenzothiophene G-Phenanthrenes C3-Phenanthrenes

294.8 295.8 297.6 300.0 304.9 305.4, 306.8 307.4 308.4 310.7 312.1 316.4 318.0 319.7 322.7 323.6 326.6, 328.7 329.4 335.1 335.9, 338.1 340.5, 343.1 354.1, 356.8

296.0 300.0

319.2 319.9 322.8 323.5

360 D. DUAYNE WHITEHURST, TAKAAKI ISODA, AND ISAO MOCHIDA in Fig. 2 (12,18).With this procedure, it is possible to follow the kinetics of the disappearance of individual sulfur species during hydrodesulfurization reactions (14, 15). Some pioneering work by the Chemical Inspection and Testing Institute (12, 13) is illustrated in Fig. 3, where a light gas oil was separated by conventional liquid chromatography contacting silica gel to yield three fractions: saturates, single-ring aromatic compounds, and multiring aromatic compounds. The third fraction was further separated into a sulfurfree fraction and sulfur-compound-containing fractions having differing degrees of aromatic ring condensation, using a procedure developed by Nishioka et al. (16). This latter method utilizes ligand exchange interactions of PdCI2 supported on silica gel to effect the sulfur species separation. It can be seen in Fig. 3 that the saturate fraction consists primarily of linear paraffins having between 12 and 25 carbons. The monoaromatic fraction is much more complicated as is the diaromatic fraction. Almost no sulfur species are found in the saturates or monoaromatic fractions. Thus, the sulfur species that must be removed from these fuels are found in multiring aromatic structures. The sulfur-free aromatic fraction was shown to be composed primarily of five classes of alkyl-substituted aromatic ring structures. These are illustrated in Fig. 4, and the individual components are enumerated in Table 11. More than 70 individual sulfur species were further identified by gas chromatography-atomic emission detection (GC-AED) and by gas chromatography-mass spectroscopy (GC-MS) (12).Figure 5 illustrates the composition of the sulfur components of the gas oil. The major components are individually identified on the figure. It can be seen that alkylbenzothiophenes and dibenzothiophenes are the major components. The size of the alkyl substituents ranges up to 16 carbons on benzothiophene and up to 7 carbons on dibenzothiophene. OF DIFFERENT SULFURSPECIES C. REACTIVITY

As this gas oil was desulfurized with a conventional CoMo/A1203catalyst, the most refractory sulfur compounds were found to be 4-methyldibenzothiophene and 4,6-dimethyldibenzothiophene,consistent with previous reports (5-8) on the relative reactivity of various sulfur species. Using separations and analytical procedures similar to those already described, the diesel fuel shown in Table I was characterized and individual rate constants were obtained for more than 60 sulfur compounds (14).As shown in Table I, the concentrations of alkylbenzothiophenes and alkyldibenzothiophenes were about equal in the diesel fuel. Table I11 summarizes the individual sulfur species that were identified, their contribution to the

9

10

11

12

13

14

15

16

17

18

19

20

21

22

23

Retention time (min)

FIG.5. GC-AED chromatography of the gas oil from an Arabian crude (12).Reproduced with permission of Ref. 12.

362 D. DUAYNE WHITEHURST, TAKAAKI ISODA, AND ISAO MOCHIDA TABLE 111 Content of Sulfur Compounds in Diesel Fuel and Their Pseudo-First-Order Reaction Constants

Number

Retention time (min)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31

15.78 17.36 17.67 17.82 17.97 18.18 19.27 19.41 19.68 19.97 20.03 20.73 20.90 21.03 21.14 21.26 21.39 21.57 21.83 21.98 22.12 22.32 22.71 23.00 23.13 23.38 23.98 24.30 24.67 25.01 25.20

Sulfur compound" MBT CZ-BT-1 G-BT-2 CZ-BT-3 C~-BT~ Cz-BT-4 C,-BT-1 C3-BT-2 CS-BT-3 C3-BT-4 C3-BT-5 C4-BT-1 C4-BT-2 C4-BT-3 C4-BT-4 C4-BT-5 C4-BT-6 C4-BT-7 C4-BT-8 C4-BT-9 C,-BT-10 C4-BT-11 Cs-BT-I CS-BT-2 CS-BT-3 CS-BT-4 C6-BT-1 DBT C6-BT-2 C6-BT-3 C6-BT-4

Rate constant (min-')

Sulfur content in feed (ppm)

CoMo

NiMo

Classification

11 47 35 74 28 122 46 75 106 171 141 42 34 61 82 65 96 88 44 76 52 67 106 92 76 92 72 169 73 89 36

>0.20 >0.20 >0.20 20.20 >0.20 0.26 >0.20 >0.20 >0.20 0.091 0.25 >0.20 >0.20 0.22 0.27 10.20 0.23 0.079 0.21 >0.20 >0.20 >0.20 0.25 0.21 0.15 0.22 0.15 0.058 >0.20 0.058 >0.20

>0.20 >0.20 >0.20 10.20 >0.20 0.39 >0.20 >0.20 >0.20 0.093 >0.25 >0.20 >0.20 0.22 0.26 >0.20 0.22 0.061 0.18 >0.20 >0.20 >0.20 0.25 0.21 0.14 0.25 0.25 0.057 10.20 0.19 >0.20

1 1 1 1 1 1 1 1 1 2 1 1 1 1 1 1 1 2 1 1 1 1 1 1 1 1 1

2 1

2 or 1 1

total sulfur content, and their individual rate constants for reaction catalyzed by a NiMo/A1203at 360°C and 2.9 MPa hydrogen pressure. In this table, we use a shorthand description for the various compounds, which is as follows: (number of substituent carbons)-(aromatic ring core structure)(isomer number). Thus, C2-BT-l represents the first GC isomer of benzothiophene containing alkyl substituents with two carbons, etc.

363

POLYAROMATIC SULFUR COMPOUNDS TABLE I11 (continued)

Number

Retention time (min)

32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60 61

25.36 25.71 25.96 26.24 26.60 26.74 27.05 27.28 28.04 28.27 28.71 28.81 29.08 29.25 29.54 29.96 30.18 30.83 31.09 31.34 31.69 32.06 32.54 32.91 33.11 33.60 33.86 34.17 34.49 35.21

Sulfur compound" C6-BT-5 CG-BT-6 C7-BT-1 4-MDBT MDBT-1 MDBTd MDBT-2 C7-BT-2 CZ-DBT-1 4,6-DMDBT Cx-DBT-2 CZ-DBT-3 G-DBT-4 Cz-DBT-5 G-DBT-6 C3-DBT-1 CyDBT-2 C3-DBT-3 C3-DBT-4 C3-DBT-5 C3-DBT-6 C3-DBT-7 C4-DBT-1 C4-DBT-2 C4-DBT-3 C4-DBT-4 C4-DBT-5

C,-DBP Cy DBT-1 CS-DBT-2

Sulfur content in feed (ppm) 53 51 78 209 173 57 105 80 89 146 37 194 103 133 128 77 81 120 65 99 96 131 100 52 87 51 57 37 61 66

Rate constant (min-') CoMo

NiMo

Classification

20.20 >0.20 0.054 0.018 0.063

>0.20 >0.20 0.034 0.020 0.065

0.071 >0.20 0.014 0.006 0.030 0.020 0.064 0.018 0.034 0.019 0.007 0.011 0.021 0.027 0.011 0.012 0.022

0.068 >0.20 0.017 0.008 0.031 0.022 0.062 0.021 0.032 0.020 0.010 0.013 0.020 0.023 0.012 0.014 0.022

2 1 3 4 3 3 2 3 3 3

0.024

0.024

3

0.009

0.009

4

4

4 3 3

4 3 or 4 3

* Compound nomenclature is abbreviated as follows: (number of C in substituent)-(parent aromatic ring)-(isomer number).

It was observed that the overall kinetics for the desulfurization could be described by lumping the rate constants for the individual sulfur species into four reactivity groups as shown in Fig. 6. These groups are listed in decreasing order of reactivity in Fig. 7. The relative contributions of the four groups were 39,20,26, and 15% for groups 1,2,3, and 4, respectively. Thus, if the new sulfur standards are to be met with this feed, groups 1,2,

364 D. DUAYNE WHITEHURST, TAKAAKI ISODA, AND ISAO MOCHIDA Representative compounds

/

h

#

C2-BT-4 C3-BT-5 C5-BT-1

A

DBT

1

2

A MDBT-1 0

0-

Cz-DBT-4

0 '

2. c

0 4-MDBT

1

+

Ci-DBT-5

0 4.6-DMDBT

0

10

20

30

40

50

60

70

Time (min)

FIG. 6. Pseudo-first-order plots of some sulfur compounds. Reaction conditions: 360°C, 2.9 MPa, NiMo catalyst. Reprinted with permission from Ref. 14, Ma et al. (1994). Copyright 1994 American Chemical Society.

and 3 must be removed completely and about half of the least reactive materials (group 4) must also be converted. Detailed analysis of the VGO indicates that lowering the sulfur content of this feed to the 0.05% level should be much more difficult than for the other two feeds shown in Table I. The dibenzothiophenes in this material are present at about twice the concentration of alkylbenzothiophenes, and it has a much higher content of alkyl-substituted polyaromatic rings. The reactivity of thiophenes fused to more condensed aromatic ring systems might be suspected to be less than those of dibenzothiophenes; however, it has been observed that the more condensed ring systems are in fact more easily desulfurized than dibenzothiophenes (14, 17),particularly when Ni-Mo catalysts are used. This observation may be rationalized in several ways. Molecules containing more than three condensed aromatic rings are easily hydrogenated. In the hydrogenated multiple aromatic ring systems, steric hindrance will be less severe due to ring puckering of the hydroaromatic, as discussed later. There is also the possibility that many of the highly condensed ring systems have the thiophene ring as the terminal ring rather than in the center of the molecule, in analogy with many biologically derived thiophene materials. Such structures will have HDS rate constants similar to that of benzothiophene, which is higher than that of dibenzothiophene. Alternatively, a higher degree of ring condensation does not neces-

POLYAROMATIC SULFUR COMPOUNDS

365

Group I - Benzothiophenes with no substituents in the 2- or 7-position Relalive k = 1

A1kyb.Q 6

!--Alkyl I - 2 2 7

S

Group 2 - Dibenzothiophenes with no substituents in the 4-or 6-position Relative k = 0.23

Group 3 - Dibenzothiophenes with 1 substituent in the 4- position Relatlve k = 0 080

\

Alkyl Group 4- Dihenzothiophenes with 2 substituents in the 4- and 6-positions Relative k = 0.028

FIG. 7. Kinetic reactivity groups in desulfurization (360"C, 2.9 MPa). Reprinted with permission from Ref. 14, M a er 01. (1994). Copyright 1994 American Chemical Society.

sarily lead to higher steric hindrance as the ring systems may be bent as in 1.2-benzanthracene rather than straight as in 2,3-benzanthracene.

D. THESIGNIFICANCE OF OTHER FEEDCOMPONENTS TO THE DESULFURIZATION PROCESS There has been a great deal of work reported on the relative rates of desulfurization of model compounds (5). However, much of this information was obtained with individual sulfur species in inert solvents. As already noted, real gas oils contain a wide variety of materials that may influence

366

D. DUAYNE WHITEHURST, TAKAAKI ISODA, AND ISAO MOCHIDA

the rates or selectivities in desulfurization reactions of polyaromatic sulfur compounds (PASC). As shown in Table I, the PASCs in gas oils represent a relatively small fraction of the total mixture. If these PASCs must be removed to near extinction, it is important to consider how the majority of the feed components will be affected by the process. High temperatures may induce cracking of coexisting fragile molecules such as linear paraffins, or alkyl groups on aromatic rings may be cleaved. Such reactions can lead to severe yield losses of the desired products. High temperatures also induce dehydrogenation and condensation reactions that produce polyaromatic hydrocarbons and gums and/or impart undesirable color to the products. High hydrogen pressures will lead to hydrogenation of the non-sulfur-containing aromatics and thus consume more hydrogen than desired. In addition, the non-sulfur-containing compounds may affect the desulfurization rate by competitive sorption on the catalytic sites. It has been reported that diaromatic ring compounds do in fact inhibit desulfurization of a variety of PASCs (18-20) and especially the group 4 type PASCs (21). In the case of VGO, the polar compounds such as nitrogen- and oxygencontaining compounds are known to have strong inhibiting effects on both hydrogenation and hydrodesulfurization reactions (18-20). In the following sections, we will integrate the compositional information presented in this section into discussions of reaction rates, inhibition, changes in selectivity, and considerations of process alternatives for achieving ultralow-sulfur levels. IV. Conventional HDS Processes and Catalysts

A. CONVENTIONAL PROCESS SCHEMES In seeking new and improved ways for achieving the ultralow levels of sulfur in the fuels of the future, it is important to understand the nature of the sulfur compounds that are to be converted (especially PASCs), as described in Section 111. It is equally important to understand how these transformations occur through interactions with catalytic surface species, the pathways involved during these transformations, and the associated kinetic and thermodynamic limitations. These considerations dictate the process conditions and reactor process configurations that must be used to promote such transformations. In this section, we describe the reactor configurations and process conditions being used today; what is known about the catalyst compositions, structure, and chemistry; and what is known about the chemistry and reaction pathways for conversion of PASCs in conventional HDS processes.

POLYAROMATIC SULFUR COMPOUNDS

367

Most of today’s distillate HDS processes consist of fixed-bed, down-flow reactors configured in a manner similar to that shown in Fig. 8 ( I ) . It should be noted that hydrogen is used in excess and is recirculated after scrubbing out the H2S byproduct. Care must be used in the scrubbing operation as it is necessary to maintain a low but optimum level of H2S in the recycle stream to maintain catalyst stability and activity. The consequence of this H2S requirement when hydrotreating PASCs to extinction is discussed in more detail in later sections, but at this point it should be mentioned that H2S is a strong inhibitor of HDS for PASCs. The installed capacities for hydrotreating distillates are predominantly moderate-pressure reactors (up to 3 MPa). Typical conditions used in today’s commercial processes are summarized in Table IV (1). In the U.S., the Clean Air Act mandated that low-emission fuels will have to be developed for future use. Industry responded quickly, and by 1994 typical diesel fuels in the U.S. contained 0.05% S, with average cetane numbers of 42 and 31-37% aromatics. California imposed stricter standards, requiring 0.05% S and a minimum of 48 cetane with an emission that did not exceed that of a 10% aromatic fuel. This is the present standard for California Air Resources Board (CARB) certification. Through the development of improved processing and additives that lower emissions, Chevron was the

separator

FIG.8. Scheme of a typical heavy gas oil desulfurizer unit ( I ) . Reproduced with permission from “Petroleum Processing Handbook” (J. J. McKetta, ed.), Marcel Dekker, New York, 1992.

368 D. DUAYNE WHITEHURST, TAKAAKI ISODA, AND ISAO MOCHIDA TABLE IV Typical Process Conditions for Various Hydrotreating Processes" Feed

Process

Temperature ("C)

HZ pressure W a )

LHSV

H2 consumption (Nm'/rn3)

Naphtha Kerosine Atrn gas oil Vac gas oil Atm residue Vac gas oil Residue

HDT HDT HDT HDT HDS HDC HDC

320 330 340 360 370-410 380-410 400-440

1-2 2-3 2.5-4 5-9 8-1 3 9-14 10-15

3-8 2-5 1.5-4 1-2 0.2-0.5 1-2 0.2-0.5

2-10 5-10 20-40 50-80 100-175 150-300 150-300

a

From Ref. I , with permission of Springer-Verlag.

first to commercialize a fuel that met these standards, although the diesel fuel contained about 20% aromatics (22). This new fuel was about 6q/gal more expensive than the previous diesel fuel. It is unlikely that such strict standards will be widely mandated, but pressures toward increasingly cleaner fuels are mounting. To meet the same low-sulfur standard for atmospheric gas oils will be much more difficult, as discussed shortly. Presently, there is an ample supply of low-sulfur crudes. The price spread between low- and high-sulfur crudes is also low. With these low-sulfur crudes, the present specifications for diesel fuels can be met by the developed countries that have imposed the new low-sulfur restrictions without straining the supply. However, as the competitive demand for these lowsulfur crudes increases in the future, their value will escalate and higher sulfur crude supplies will have to be considered even though processing will be much more difficult. In Europe and Japan the 0.05% S standard is nearing implementation, but meeting these standards will be difficult in Europe without the construction of higher-pressure andlor second-stage reactors. As discussed earlier, recent installation of new HDS reactors, at depressed prices, allowed the Japanese refining industry to achieve the 1993 targets of 0.20% S in gas oils, and it is now prepared to meet the 0.05% S target when specifications are implemented, as long as low-sulfur crudes are available. As the supply of these crudes declines, the importance of new, more active catalysts and/ or new process configurations will increase. Activity can be increased by increasing reaction temperature, but temperature limits are being approached with today's catalysts, as severe side reactions, such as cracking of paraffins and alkylaromatics, occur at temperatures only about 40°C higher than those used presently. In addition, product quality suffers due to gum and polyaromatic hydrocarbon formation andlor color development.

POLYAROMATIC SULFUR COMPOUNDS

369

Higher temperatures also lead to more rapid aging, due to catalyst sintering and coking. Even if new noncracking catalysts can be developed, thermodynamic equilibrium constraints on partially hydrogenated intermediates will be the next obstacle at more elevated temperatures. Such equilibrium limitations are significant in denitrogenation of polyaromatic rings where nitrogen removal requires prior ring saturation. Nitrogen compounds are severe catalyst poisons (because of adsorption), and raising the temperature to overcome adsorption limitations leads to new limitations imposed by thermodynamics, according to which ring saturation is not favored at high temperatures (20).This point is discussed in more detail in the section on limitations in conventional processes. To put these problems into perspective, as already discussed, the PASCs remaining at the 0.20% S level are lower in reactivity, by a factor of 10-50, than the sulfur compounds that are now removed in lowering the sulfur level from 1.2 to 0.20%. Even with catalysts 2 times as active, the reactor volume may have to be doubled to convert the required 75% of the leastreactive PASCs to achieve the 0.05% S target. Based on the composition of a typical gas oil and using the first-order rate constants for the different classes of sulfur compounds (14), a theoretical HDS severity plot is presented in Fig. 9. It can be seen that reactor volumes will have to be increased by about a factor of 4 to meet the new specifications unless much more active catalysts can be developed than are presently available. Thus, industry has a major challenge to develop new catalysts and processes for future clean-fuel production. In the following discussion, we review the state of knowledge of present-day catalyst compositions and surface chemistry and of mechanistic reaction pathways to aid in identifying where improvements can be made in catalyst composition and process configurations. PATHWAYS AND MECHANISMS B. HDS REACTION As discussed in Section 111,when the sulfur content is lowered from 0.20 to 0.05%, the chemistry of HDS of gas oils is essentially the chemistry of alkyl-substituted dibenzothiophenes. Though gas oils initially contain mostly alkyl-substituted benzothiophenes, these are completely removed by the time 0.20% S is achieved. Thus, this review will deal predominantly with the reaction pathways involved in the HDS of alkyl-substituted dibenzothiophenes. There are many excellent reviews on reaction pathways of the more reactive sulfur species such as thiophenes and benzothiophenes (2, 5, 8, 23, 24), and the reader is referred to those reviews for information on the reaction pathways and mechanisms of HDS for the more reactive

370 D. DUAYNE WHITEHURST, TAKAAKI ISODA, AND ISAO MOCHIDA

10,500 10,000 9500 9000

8500 8000 7500 L

7000

3

5

6500

5

6000

m

5500 Q

5000 4500 .~ 4000~~

Present specification

3500~~ 3000~~ 2500

Future specification

2000

1 4

1 500 .1000

-

500 ~0' 0

500

1000

1500

2000

2500

30

FIG.9. Simulated desulfurization to meet 0.05% specification

sulfur compounds. The most recent and comprehensive is the review by Girgis and Gates (5). Though it has long been known that the more highly condensed thiophene structures, such as dibenzothiophene and especially their alkyl-substituted derivatives, have low HDS reactivity (21, 25-29), it is only recently that study in this area has intensified. Researchers throughout the world are now actively seeking understanding of the fundamental causes of low reactivity and attempting to find means to circumvent the problems (29-33). For perspective, Table V presents the reported relative reactivities of selected thiophene derivatives and related compounds (5, 26, 34). In this table, the dibenzothiophene rate constant has been renormalized to equal

371

POLYAROMATIC SULFUR COMPOUNDS TABLE V HDS Reactivity of Thiophene Derivatives (5, 26, 34) (Relative Rate Constants for CoMo/A1203,300"C, 7-10 MPa) Batch reactor

Flow reactor

2250

very fast

1330

400

100

130

260

530

100

372 D. DUAYNE WHITEHURST. TAKAAKI ISODA, AND ISAO MOCHIDA 100 and all of the other rate constants were adjusted accordingly. It is curious to note that as the thiophene ring systems become more condensed beyond dibenzothiophene, the reactivity increases. This phenomenon has been observed in the HDS of commercial gas oils, where it was found that the least reactive sulfur compounds in gas oil were the middle distillate range sulfur compounds (predominantly alkyldibenzothiophenes) ( I 7).This observation is most likely explained by the fact that highly condensed aromatic ring systems are more readily hydrogenated and sulfur extraction from the resulting hydrothiophene derivatives is more facile, especially if hydrogenation occurs on the ring adjacent to the sulfur (see Table V). It is generally agreed that there is a common sequence of mechanistic pathways for all PASCs as shown earlier in Scheme 1.Though the partially hydrogenated thiophene intermediate is illustrated in that scheme as a tetrahydro derivative, it has been shown that this intermediate is, in reality, an equilibrium mixture of the tetra- and hexahydrodibenzothiophenes. The rate of equilibration is at least 10 times greater than the other associated rates. Thus, the pair can be treated as a single kinetic species (26). Unfortunately, even though this sequential pathway is generally recognized by researchers in the field, there is no common way to report the relative rates. This leads to confusion to those who read and try to compare results from different references. It is suggested that the following designation of rate constants be adopted t o avoid future confusion. The subscripts of the rate constants are meant to represent the chemical transformations involved and are thus more intuitive for comparing results with dibenzothiophenes having different degrees of substitution. The rate constants of concern are summarized as follows. Rate constant

kDo ko, kD2 k~s, kHsl k~p,

k~~~

Reaction concerned Desulfurization without ring hydrogenation Desulfurization of one ring of the hydrogenated sulfur compound Desulfurization of the fully saturated sulfur compound Hydrogenation of one ring of the sulfur compound Hydrogenation of the second ring of the sulfur compound Hydrogenation of one phenyl ring of desulfurized biphenyl Hydrogenation of the second phenyl ring of biphenyl

Abbreviation in tables and figures

-

6S6+ 6-6 @Cy, CyJCy,

6S6

-+ -+

+Cy,

Cy,-Cy, 4SCY6

4SCY6 4 CY6SCY6

4-4 $-cy6

--f

d-CY, CYCCy6

If one takes, as a reference, the rate constant for direct desulfurization of unsubstituted dibenzothiophene (k,,,,) and renormalizes this value to be

POLYAROMATIC SULFUR COMPOUNDS

373

100, then all other rate constants can be visualized on a relative basis for comparison between different authors, different catalysts, and conditions. Comparisons of reactivity for other alkyl-substituted dibenzothiophenes can also be made if the same reaction conditions and catalysts are used for the substituted and unsubstituted compounds. Throughout this report, we will discuss the rate constants of various sulfur-containing compounds using renormalized values, for ease of comparison. Table VI presents the renormalized relative rate constants reported by a variety of authors for the conversion of dibenzothiophenes (26, 28, 29, 31,32). At the bottom of this table, selectivities are presented. It is common to describe a catalyst’s selectivity as the relative contributions from a direct extraction of the sulfur atom (kD,)and the contribution from hydrogenation prior to desulfurization (kHsl,the rate-determining step). As can be seen from this table, there does not appear to be a general agreement between different authors on the relative contributions of the different pathways, even when using similar catalysts and similar conditions. Differences in the reported relative contributions of these two routes vary by more than two orders of magnitude for the same type of catalyst. Comparisons between NiMoS,/A1203 and CoMoS,/Al2O3 and between results obtained under very different reaction conditions are of course expected to be different, and reasonable trends are observed, as discussed later. The roots of the unexpected differences appear to lie in the fact that the reported rate constants were obtained by using curve-fitting techniques with insufficient data, not placing reasonable chemically based boundaries on the initially assumed rate constants, and not including all of the mechanistic pathways. Most researchers have found pseudo-first-order behavior for the various steps, and so it is possible to match theoretical curves with data to obtain the best rate constant values. Unfortunately, in most instances, too few data points were obtained to generate a unique theoretical fit. It is absolutely imperative that data be obtained for at least four conversion levels that are well spaced in the conversion matrix and extend to over 95%conversion. The partially hydrogenated dibenzothiophene intermediates are most often never detected as their desulfurization rates are extremely high (kD, and kDJ.The cyclohexylbenzenes and bicyclohexyls can arise from two different routes, and the concentrations of their precursors (biphenyl and cyclohexylbiphenyl, respectively) pass through maximum values that can easily be calculated from the relative values of the formation and conversion rate constants. However, unique values for these relative rates can only be predicted if data are available well prior to and well beyond the times of maximum concentrations for these intermediates, because minor experimental errors can confuse curve-fitting optimization. An alternative approach would be to independently establish rate con-

TABLE VI Comparison of Relative Rate Constants for Dibenzothiophenes by Different Authors"

Vrinat 32 NiMo/AlzOl 290 5

100

Hou a 11a 26 CoMo/Al2O3 300 10.2

lsoda 31 CoMo/A1203 320 2.5

Geneste 29 CoMo/Al,Oq 300 5

Vrinat NiMo/Alp03 32

Isoda NiMo/A1203 31

NiMo/AI2O? 320 2.5

Aubert 28 NiMo/AI2O3 340 7

290 5

320 2.5

100

100

100

100

100

100

100

Isoda 31

3750

5.2

106

22

0.15

32

0

370

12

0

0

33

11

0

46

0

0

670 5.9 25.0

3.1

0.3

I .4

33 19.2

"For comparison, in all data sets k,,,, IS set equal to 100.

0.9

4.5 3.7 1.8

2.2 18.5

POLYAROMATIC SULFUR COMPOUNDS

375

stants for biphenyl and cyclohexylbenzene hydrogenation rates (kHpland kHPJ so as to set boundaries on the assumed rate constants that are being fit. Ideally, competitive reactions should be employed that allow direct measurement of the uncertain reactions. This has been done by several authors, but the independently measured rate constants were often not used in setting boundaries in the computer fitting of the complete matrix. Thus, vastly different relative rate constant values in the fitted matrix were obtained, as can be seen in Table VI. It is true that completely independent experiments may be confused by subtle effects of inhibition by H2S, which is a byproduct in desulfurization reactions, but as discussed later, the magnitude of the inhibition at the levels produced in the experiments is not large enough to result in such drastic changes in the hydrogenation rates of biphenyl. A few well-chosen experiments in which competitive test reactions are conducted simultaneously can provide definitive information that can be used in setting reasonable limits on the ratios of the important rate constants. This is illustrated later. For perspective, Figs. 10a and 10b summarize some relative rate constants and thermodynamic equilibria for the hydrogenation of representative aromatic hydrocarbons (5, 35). Both figures present normalized rate constant data in which the rate constant for biphenyl hydrogenation is set to 100. The data in Fig. 10a were obtained under typical HDS conditions with CoMoS,/A1203. Figure 10b refers to temperatures that are higher and are more typical of coal liquefaction conditions, but they are useful for illustration of how relative rates and thermodynamic equilibria are affected by different conditions. Both sets of data suggest that biphenyls should be less reactive than naphthalenes but more reactive than simple monoaromatic ring systems and that more condensed aromatic systems are even more easily hydrogenated. There have been reports issued in which the hydrogenation rate constants for dibenzothiophene (kHsl) have been assumed to be several orders of magnitude smaller than the hydrogenation rate constant for biphenyl (kHP,),and good fits of the data were obtained. However, equally good fits of the same data can also be obtained if one assumes that the rate constants for hydrogenation of dibenzothiophene and biphenyl are about the same. Thus, before proceeding, one must make some rational assumptions to set reasonable boundaries on the rate constants that are being fit. As mentioned earlier, the desulfurization rates of tetra- and hexahydrodibenzothiophenes are high relative to rates of other reactions in the overall HDS reaction scheme. Because of this, these intermediates are often not observed experimentally. Thus, the origins of cyclohexylbenzene and bicyclohexane are confused. When attempting to deconvolute all of the rate

376 D. DUAYNE WHITEHURST, TAKAAKI ISODA, AND ISAO MOCHIDA a 4

96

78

22

3

33

97

9

58

FIG. 10. (a) Relative hydrogenation reactivities of selected aromatic hydrocarbons (CoMoSJA1203, 7.5 MPa, 325°C). Normalized rate constants are shown above the arrows, and equilibrium concentrations at the specified conditions are shown below the compounds (5). (b, facing page) Relative hydrogenation reactivities of selected aromatic hydrocarbons (CoMoS,/A1203,69 MPa, 427°C).Rate constants are shown above the arrows, and equilibrium concentrations at 69 and 3.1 MPa are shown below the compounds (35).

constants with limited data, it is very helpful to know what relative rates seem reasonable for hydrogenation of dibenzothiophenes and their corresponding biphenyls. Also, at high temperatures, the concentrations of the intermediate partially hydrogenated sulfur compounds may become equilibrium limited. In this case, the aforementioned simple model becomes more complicated. As discussed later, elucidation of these thermodynamic equilibria for alkyldibenzothiophenes is needed, and future work in this area is recommended. It is believed that conventional HDS catalysts possess two very different types of catalytic sites which contribute to the different pathways described earlier. One induces the direct extraction of sulfur (kD,,),and the other catalyzes aromatic ring hydrogenation (IcHs,) ( 1 ) .This topic is discussed in

377

POLYAROMATIC SULFUR COMPOUNDS

b 3.1 69 MPa

I

69 MPa

22

78

3.1 MPa

65

35

69 MPa

22

78

3.1 MPa

60

40

69 MPa

10

34

29

27

FIG.10. (continued)

more detail later. Thus, all reacting species (intermediates as well as starting materials) must compete for adsorption on a particular type of site and most likely are transformed through the same fundamental mechanisms. Many authors have observed typical Langmuir-Hinshelwood type kinetic behavior, which supports this assumption (5). The behavior of each compound, whether an intermediate in a complex matrix or reacting as an individual species, should be the same. It should be kept in mind that the relative adsorption constants (K,)for different reactants may not be the same for the direct desulfurization and hydrogenation sites (see Section VI for more details). Thus, knowledge of the fundamental reactivities and adsorption constants of dibenzothiophenes and their corresponding

378

D. DUAYNE WHITEHURST, TAKAAKI ISODA, AND ISAO MOCHIDA

diphenyls is very important in the optimization of curve fitting of these complex reaction schemes. Reported relative rates (kHsIIkHpl)are quite varied. As shown in Fig. 11,this ratio has been estimated to be from 1.5 (35) to as little as 0.13 (23) for some multiaromatic thiophene derivatives. Phenylnaphthalene was also reported to have a very different preference for the position of naphthalene ring hydrogenation in independent studies as compared to being an intermediate in the reaction matrix (36, 37). Molecular orbital calculations (HOMO-HOMO-3) predict that the bond (CoMo/A1203,300% 7MPa) 1361

I

i

410

loo

(CoMo/A1203,32fC, 7.5 MPa) 1371 12

~

=

% 29

0 28

~

(NiMo/AI26, 250°C, 4MPa) 1231

I

14000

&

-0 380

1

100

0-

570

440

& 0

0

FIG.11. Reaction pathways in the HDS of multiaromatic thiophene derivatives. All rate constants are normalized to kDO= 100for ease of comparison. (Top) Reprinted with permission from Ref. 36. Copyright 1981 American Chemical Society; (middle) from Ref. 37; (bottom) from Ref. 23.

POLYAROMATIC SULFUR COMPOUNDS

379

orders of benzothiophenes are slightly higher (0.014) than those of the corresponding biphenyls. It has been shown that enhancing the bond order in a series of benzothiophenes by 0.05 results in a 10-fold increase in hydrogenation reactivity (38). Thus, one might expect that kHSlwould be about three times greater than kHP,in the dibenzothiophene matrix. An alternative method for setting guidelines on rate constant assumptions is from the halfwave potential in electrochemical reduction of organic compounds in nonaqueous solution. This technique has proven quite useful in predicting the relative catalytic hydrogenation reactivities for a variety of aromatic hydrocarbons (39,40).Table VII shows the halfwave potentials for a variety of aromatic compounds. These data indicate that the benzothiophene ring system should be more easily reduced than the corresponding naphthalene ring systems but that biphenyl ring systems and dibenzothiophene ring systems should have comparable reactivities toward catalytic reduction. One might expect that dibenzothiophene may have greater reactivity than biphenyl in catalytic hydrogenation (kHsIvs kHpl)because of a stronger adsorption on the catalyst, but a much lower rate constant seems unreasonable. Competitive reaction studies have borne out this assumption (41). Perhaps the most definitive method for establishing guidelines on relative reactivities is through carefully selected competitive reactions with model compounds. For example, if one desires to know the relative rates of hydrogenation of dibenzothiophene and biphenyl, it is useful to add an analog of biphenyl to the reaction mixture in a competitive experiment. In this way, one can simultaneously determine (kDo+ kHsl)for dibenzothiophene and k,, for the biphenyl analog, under conditions having the same solvent effects and H2S byproduct inhibition. In separate experiments, the relative values of kHp,for biphenyl and the analog may then be determined and this ratio may be applied in the competitive HDS experiment to set guidelines for assigning initial values for kHs, and kHp,prior to curve fitting. For example, data obtained by this technique are shown in Table VIII and Fig. 12a for HDS of dibenzothiophene catalyzed by a CoMo/carbon (41). In these experiments, 3,3'-dimethylbiphenyl was used as the biphenyl analog, as the chemistries are identical and only small differences in adsorption behavior are expected. As shown in Table VIII, the ratio of rate constants ((kD, + kHsI)/kHpl)was estimated to be 62. With this ratio held constant, the other rate constants in the HDS pathway matrix were estimated using curve-fitting techniques. The graphs in Fig. 12 (41, 41a) represent product selectivity plots in which the relative concentrations of products of HDS are plotted against the percentage conversion of the starting material. The lines represent the calculated product distributions for sequences of first-order reactions having

380 D. DUAYNE WHITEHURST, TAKAAKI ISODA, AND ISAO MOCHIDA TABLE VII Halfwave Potentials of Model Compounds (39, 40)

Halfwave potential El12

Q I

(V)

Diffusion current I d (MA)

-2.75

-2.82 (Q

=

1.68)

-2.5-2.8 (6-4)

no wave up to -2.85 H -2.5

a mO" /

268

-2.59

158

-2.49, -2.57

157

-2.61

292

-2.5

269

-2.55

216

/

381

POLYAROMATIC SULFUR COMPOUNDS TABLE VIII Competitive Rate Constants in HDS Reaction (41) Exp 1

Exp 2

Calc

the values for the relative rate constants for the various reactions shown in the pathways above the graphs. The points are the observed experimental values. As can be seen in Fig. 12a, this technique provides very satisfying results. It should be noted that the estimated kHS,IkHp,for this catalyst is 7, consistent with the theoretical arguments presented earlier.

a

b (NiMo/A1203, 320°C ,2.5MPa)

(CoMo/Carbon, 300°C. 2.9MPa)

Q-p*Q-p

Q-$l-"-QQ 1 3300 k",,

k,,

kD,

Product selectivity 90

1

'00

1

kD, 3300

Product selectivity 90 1

I

6

'3

.-",

701

,A

g 60

5

50

5 40 2

73

g 30 a 20

s 10 0

10 20 30 40 50 60 70 80 90 100

YOConversion of A

0

0

10 20 30 40 50 60 70 80 90 100

% Conversion of A

FIG.12. Changes in HDS reaction pathways of dibenzothiophene with catalyst type. All rate constants are normalized to kDo = 100 for ease of comparison (41, 41a).

382 D. DUAYNE WHITEHURST, TAKAAKI ISODA, AND ISAO MOCHIDA The graphs shown in Fig. 12 are selectivity plots developed by assuming first-order behavior for all reactions. It should be mentioned that satisfactory estimation of the various rate constants requires that comparable fits must also be obtained for kinetic plots in which the product composition is plotted against reaction time using ail of the rate constants obtained from the selectivity plot. The data discussed in this section satisfy this criterion, as illustrated in Fig. 13. The excellent agreement between the calculated curves and the data clearly demonstrates that all reactions exhibit pseudo-first-order kinetics under a given set of conditions. Figure 12b illustrates an example of estimating relative rate constants for HDS of dibenzothiophene catalyzed by a NiMoS,/A1203 using wellspaced conversion experiments. The data point at very high conversion is essential in obtaining a unique solution to the overall matrix of rate constants as the data obtained below 90% conversion could be simulated with

a

I

k,

100

.-c

g

70

50

6 40

3300

Change in composition with time

1

""I\

c 70 0

a

E

.

00

.. .

Qc '

I

0

8

30 20 10

0 0

1

k,,

0 .c .In

60

(NiMo/Al,O,, 320°C , 2.5hWa)

3300

Change in composition with time

80

.-$

b

(CoMo/Carbon, 300°C. 2.9MPa)

20

40

60

80

Reaction time (min)

101 3

-

-+

80

100

Reaction time (min)

FIG.13. First-order time dependence in the HDS of dibenzothiophene. All rate constants are normalized to ko, = 100 for ease of comparison (41, 41a).

POLYAROMATIC SULFUR COMPOUNDS

383

many quite different sets of rate constants. Note again, kHsl/kHp,is greater than 1. The results presented in Figs. 12a and 12b provide one other piece of information. As discussed in more detail later, NiMoS,/A1203 has a relatively high activity for aromatic ring hydrogenation. This indicates that the catalyst of choice for high-conversion HDS may change as PASCs become more refractory due to steric hindrance imposed by alkyl substituents. This is discussed in more detail in the next section. If one uses the aforementioned guidelines for the expected relative rates of hydrogenation of dibenzothiophene to give biphenyl (kHsl/kHpI),then it might be appropriate to reconsider the data reported in the literature. An attempt has been made to recalculate reported literature data with these new assumptions, and the results are presented in Table IX. These newly derived relative rate constants were obtained by extracting data from the reported graphs and may not be completely accurate. However, the calculated predictions were in quite good agreement with the reported data (comparable to the fits shown in Fig. 12). The newly derived relative rate constants should be compared with those given in Table VI. The data reported by Geneste (28) were not recalculated, as the authors reported zero-order kinetics in their studies and the fitting technique described here is not appropriate. The comparisons between Tables VI and IX are admittedly crude, but the new data (Table IX) are now consistent, with kHs, being larger than kHp, and NiMoS,/AI2O3 being a more active hydrogenation catalyst than C o M o S x / A l ~ 0Another ~. point can be made from the results shown in Table IX. As dibenzothiophene becomes more substituted, particularly in positions adjacent to the sulfur atom, the importance of hydrogenation activity becomes greater. This is discussed further in the next section.

C. EFFECTS OF ALKYLSUBSTITUENTS ON HDS BEHAVIOR It has been known for many years that the ease of removal of sulfur from a thiophene compound is affected by the presence of alkyl substituents near the sulfur atom ( I , 2, 5, 8, 9). Over 25 years ago, Givens and Venuto clearly showed that the position and number of substituents present on benzothiophene had a strong influence on both the overall reactivity and the degree of desulfurization (25).Although that effort was directed toward hydrocracking applications, it still has relevance to today’s HDS processes. Table X summarizes their work. Though there is a general trend of reduced reactivity with degree of substitution, it appeared that substitution on the thiophene double bond in the 3 position had the greatest detrimental effect. As it is very easy to hydrogenate the thiophene double bond and the

TABLE IX Recalculated Relative Rate Constants for Different Catalysts and Conditions" Starting compound

Author Original reference Catalyst Temperature ("C) Hz pressure (MPa) Reactant -+ product

4S# 5 4-4 6Scy6

A ccy6

4s 4

4s4

4s 4

&4

@tJ

&4

H3C4s4

H3C&4

Vrinat 32 NiMo/Alz03

Isoda 31 NiMo/Alz03

Aubert 28 NiMo/Alz03

Houalla 26 CoMo/AhO3

Isodab 31 CoMo/Al,O3

Geneste' 29 CoMoiAlZO,

Vrinat 32 NiMo/Alz03

Isoda 31 NiMo/Alz03

290

320 2.5

340

5

7

3M) 10.2

320 2.5

300 5

290 5

320 2.5

100

100

100

100

100

100

3900

3900

2500

405

2592

3000

23

23

31

5.4

370

210

15

15

29

4

250

85

0.3

0.5

ko

cybscy6 A Cy6-Cy6

29

~~~~

~

~

For comparison. in all data sets kl,,, is set equal to 100. Insufficient data to attempt recalculation. ' Zero-order kinetics, not appropriate for recalculation

POLYAROMATIC SULFUR COMPOUNDS

385

TABLE X Reactivities of Methyl-Substituted Benzothiophenes (25) (CoMo/Alz03, 400"C, I atm)

Yield desulfurized product

Relative

(%I

Desulfurization selectivity

91

100

91

1.oo

74

89

66

0.448

60

95

57

0.350

43

75

32

0.160

54

80

43

0.233

47

51

24

0.114

39

38

15

0.067

34

16

0.072

Conversion

Q-&

p

p 0

kdes

S

dihydrothiophene is orders of magnitude more reactive than the fully aromatic precursor, it may well be that the effects observed reflected a reduction in the relative contribution of hydrogenative HDS vs direct extraction of the sulfur atom. As one proceeds to more condensed thiophene derivatives (such as dibenzothiophene), the effects of alkyl substitution become much more signifi-

386 D. DUAYNE WHITEHURST, TAKAAKI ISODA, AND ISAO MOCHIDA TABLE XI Effects of Methyl Substitution on Relative Reactivity in HDS" Author Reference Catalyst Temperature ("C) H2 pressure (MPa)

Houalla 1, 26 CoMo/A1203 300 10.2

Vrinat 32 CoMo/A1203 290 5

Vrinat 32 NiMo/AlzOl 290 5

Isoda 21 NiMo/Alz03 320 2.5

Starting compound DBT 4-MDBT 4,6-DMDBT 2.8-DMDBT 3,7-DMDBT

100 9 6.7 91 48

100

100 29

100 38 11

25

'For comparison, in all data sets the DMB conversion rate is set equal 9

1

6

4

to 100.

cant. Table XI summarizes the reports of several authors that clearly show that alkyl substitution adjacent to the sulfur atom is by far the most damaging toward reactivity. The early work of Houalla clearly showed that ring substitution in remote positions such as the 2, 7, and 8 positions did little to reduce the reactivity (26). It has been proposed that steric inhibition of adsorption on catalyst surfaces is the major cause of such reduced reactivity (17, 26). Recent computer modeling clearly shows that for molecules such as 4,6dimethyldibenzothiophene (4,6-DMDBT), the methyl groups interfere with catalyst-molecule interactions as the sulfur atom adsorbs primarily through a one-point attachment and the dibenzothiophene ring system is nearly perpendicular to the catalyst surface. Hydrogenation of one ring of the 4,6-DMDBT causes the rigid planar aromatic structure to pucker and allow much better interaction between the sulfur atom and the catalyst surface (17). The result of this behavior is that dibenzothiophenes with methyl substituents in the 4 and/or 6 positions force the reaction matrix to favor the hydrogenative pathway (21).Thus, NiMoS, /A1203catalysts may be more desirable than their cobalt counterparts as they have higher hydrogenation activities. This is reflected in the relative reactivities of the methyl-substituted dibenzothiophenes in the presence of CoMoS, /A1203and NiMoS, /A1203 shown in Table XI. Whereas substitution in the 4 position lowers the reactivity of

387

POLYAROMATIC SULFUR COMPOUNDS

benzothiophene with CoMoS, IAI2O3by a factor of 11,with NiMoS,/A1203 catalyst the reactivity was only lowered by a factor of about 2.6. Using the computer-fitting procedures of the reaction pathways described in the previous section, the relative rate constants for HDS were determined for dibenzothiophene (DBT), 4-methyldibenzothiophene (4-MDBT), and 4,6-DMDBT under typical gas oil HDS conditions with NiMoS,/AI2O3 catalysts. The various rate constants are summarized in Table XI1 and the reaction pathways and associated rate constants for 4-MDBT and 4,6DMDBT are provided in Figs. 14 and 15, respectively. Some general trends may be observed in this table. First, it can be seen that the direct sulfur extraction rate constant ( k D , ) is much more sensitive to the presence of methyl substituents adjacent to the sulfur atom than is ) either the hydrogenation rate of the sulfur-containing compound ( k H s lor the hydrogenation rate of the corresponding biphenyl derivative (kHP,).In fact, the rate of hydrogenation of 4-MDBT is a little higher than that of

TABLE XI1 HDS Rate Constants for Alkyldibenzothiophenes over NiMoS,/A120j (Rate Constant X 105 s-' . (g of catalyst)-', 2.5 MPa H2, 325"C)y Starting compound

-

DBT

4-MDBT

4,6-DMDBT

@d'kD"- #+4

86.1

14.4

6.6

2840

432

293

Reactant

d'scy6

product

-% &Cy6 kD

488

cy6scy6 A cy6-cyS

+s+ k 4SCY6 d'scy, 4-4 &cy6

18.9

30

3cy6scy6 k,,

d'-cY6

21.5 17.6

16.4

12.3

kH'., cy6-cy6

5.8 3.4

Catalyst selectivities

bo J~Hs, b0+ kHs,/ICHP, k H S , ikHPl a

4.6 6.4 1.2

Data recalculated from Isoda et a(. (2Z).

0.5 3.6 2.4

0.3 4.8 3.7

388

D. DUAYNE WHITEHURST, TAKAAKI ISODA, AND ISAO MOCHIDA

FIG. 14. Relative rate constants in the HDS of 4-methyldibenzothiophenecatalyzed by NiMoS,/A1203 at 32YCand2.5 MPa of Hz .The rate constants above the arrows are normalized to kDo = 100 for ease of comparison. The absolute rate constant for 4-MDBT conversion was s-* . (g of catalyst)-'. 44.4 x

Change in compositionwlth time

Product selectivity

0 % Conversion of A

FIG 15 Kinetics and reaction pathways in the HDS of 4,6-dlmethyldibenzothiophene (320°C, 2 5 MPa, NiMo/A1203)(41a)

POLYAROMATIC SULFUR COMPOUNDS

389

the unsubstituted DBT. The effect of methyl substituents on the rates of reduction of biphenyls is not large, but dimethylbiphenyl does have the lowest rate. This is consistent with reports on the effects of methyl substituents on the hydrogenation rates of single aromatic rings, which showed very little difference in rates for substituted and unsubstituted benzene (42). The result in the changes in these rate constants is that the preference for reaction pathway changes from direct extraction being strongly preferred in DBT to hydrogenative HDS being highly preferred for 4,6-DMDBT. In the case of 4,6-DMDBT, it was possible to determine the rate constants for direct extraction of sulfur from the fully saturated sulfur-containing ring system (kD,) and for the secondary hydrogenation of the tetrahydroAs might be expected, the rate dibenzothiophene intermediate (kHS2). constants for direct sulfur extraction follow a clear trend in which k,, < k,, < k,, . The reverse trend is observed in the aromatic ring hydrogenation rates, kHs,> k,, and kHpl> kHp,, which is consistent with the literature (see Fig. 10) (5, 35). As all of the direct extraction rates were low relative to the hydrogenative HDS rates, even at low conversions, the fully saturated compound dimethylcyclohexylcyclohexanewas observed in the product mixture. As discussed in later sections, this fact indicates that the major cause of rate reduction by adjoining alkyl substituents may not be due to lowering adsorption constants but could well be due to steric limitations in the oxidative addition of the C-S bond to the catalytic site, as discussed later. A more subtle explanation could be adsorption-disguised kinetics whereby the intermediate is not released from the catalyst surface but remains adsorbed so that further conversion proceeds. This was not observed for unsubstituted dibenzothiophenes, however. Having established reliable values for all of the important rate constants as a function of alkyl substitution on dibenzothiophenes, it is now possible to examine critically how these rate constants (and associated changes in product selectivity) are affected by other components of commercial gas oils and by the H2S that is produced during the HDS process. It is also possible to evaluate how these various rate constants are affected by changes in catalyst composition and by process conditions. Knowledge of the details of these effects can lead to novel catalyst modifications and process configurations that may be able to reach the new stricter standards of 0.05% S. These topics are discussed in later sections. However, for perspective, we will first summarize what is known about present-day catalyst compositions and catalytic mechanisms that bring about the transformations observed in HDS processes.

390 D. DUAYNE WHITEHURST, TAKAAKI ISODA, AND ISAO MOCHIDA AND STRUCTURE OF PRESENT-DAY HDS CATALYSTS D. COMPOSITION

1. Structure and Class$catioras of the Co(Ni)-Mo-S Species

Although Co(Ni)MoS, /A1203and related catalysts have been used commercially for more than 70 years, the exact chemical nature and structure of the active species are still not known. Many models and theories have been put forth and argued in the literature for almost the same length of time. There have been many excellent reviews that summarize in detail the evolution of theories of the structures of these catalysts (1-3, 9, 10) and there is little need to repeat in detail the contents of those reviews in this article. However, for ease of understanding the following discussion, the various models that have been proposed are illustrated in Fig. 16. We briefly discuss these models as there is a need to understand as much as possible about the nature of the best catalysts today to see if there is some theoretical limitation to increasing the activity of Co(Ni)MoS, /support catalysts to the levels that will be required for meeting the new 0.05% S or stricter specifications through catalyst replacement in existing units. Such new catalysts will have to have activities at least an order of magnitude greater than those of today’s best catalysts. Keep in mind that there has been more than 50 years of intensive study and continual improvement to reach the activities available today, and so the probability of increasing catalyst activity more than an order of magnitude seems quite low. To understand present thinking about the catalyst structure, the following discussion will concentrate on Co(Ni)MoS,/A120, as this combination has received the most attention and is the closest to being understood. Similar catalysts composed of Co(Ni)WS, /A1203are also known and commercially used and silica-based supports have some special applications. However, for ease of describing what is known about catalyst structures and chemistry, this review will focus on the Co(Ni)MoS,/A1203 system. Although we have represented the composition as Co(Ni)MoS, , in operating catalysts the S/Mo stoichiometry is very near 2. The importance of the small deviation from this stoichiometry will become apparent as we discuss the potential for increasing the catalytic activity beyond the present values. A brief description of the different theories will now be given to show what has been proposed and rejected as well as what is known and what is not known. One curious observation is that high activities for direct sulfur extraction from thiophene derivatives are only exhibited by metal sulfides that form stacked lamellar crystallites, similar t o graphite structures (1-3). MoS2 is classic in this regard and has found applications as a high-temperature lubricant with properties very similar to those of graphite. The other widely used metal sulfide in HDS is WS2, which also forms lamellar crystal struc-

Diameter

Sites I rim

\ n lavers stacking height

MoS2 Rim-Edge Model (46)

Contact Synergy Model (53)

I

MoS2 Monolayer Model (50)

a

ntercdation Models (43,44,51, 52)

d-

-I-

NilMoiS CRYSTALLITE

,,

HDS

+ o=5 0 5 H0.w

0 = Ni,Co

Co-Mo-S Model

e

( I 1,5340)

Sulfide Bimetallic Species Models (2)

I

Surface Complex Model

FIG.16. Proposed models for Co(Ni)MoS,/Alz03 type catalysts. Reproduced with permission of cited references. e, Reprinted with permission from Ref. 60. Copyright 1986 American Chemical Society.

392 D. DUAYNE WHITEHURST, TAKAAKI ISODA, AND ISAO MOCHIDA tures. These lamellae may be envisioned as a sandwich of the metal between two sulfur layers. The sulfur layers constitute the basal plane of the stacked layered structure (see Figs. 16a, 16d, 16f, and 16g). This model will be important to keep in mind in the following discussions. The exact crystal geometry of this stacked MoS2 sandwich is also important to the catalytic activity and selectivity. Although the importance of the edges of basal planes in MoSz catalysts has been known for many years (43-47), a recent study by Daage and Chianelli clearly demonstrated that there are two types of sites on unpromoted MoS2 crystals (47).These sites were termed “rim” and “edge” sites and refer to sites associated with the edges of the terminating layers of the crystal stack (rim sites) and the sites associated with the edges of the interior layers of the stacks (edge sites), as illustrated in Fig. 16a. The ratio of the number of rim and edge type sites was shown to be related to the ratio of hydrogenation (IcHs,) and direct sulfur extraction (k,,,,) rate constants. In their experiments, the hydrogenation sites were found to correlate with the rim sites and the direct sulfur extraction sites correlated with the edge sites. Co(Ni)MoS, /A1203 type catalysts are known as “promoted” MoS, / Al2O3catalysts. This means that small amounts of an added sulfide of a second metal, Co or Ni, induces major improvements in the activity of the catalyst. Analogously promoted catalysts are also known for supported WS, catalysts. The first patent relating to this promotion was by I. G. Farben in 1928 (48). Originally, it was thought that the promoter somehow increased the activity of the MoS, as neither CoS, nor NiS, exhibited significant activity alone. As early as 1943, A. C. Byrns demonstrated that the promoter had to be in intimate contact with the MoS, as mixtures of separately supported metal sulfides showed no increase in activity (49). From that time on, theories to explain this phenomenon have flourished. Without such theories to test and either confirm or reject, fundamental understanding cannot advance. It was initially believed that the promoter and Mo sulfides were individual crystallites in intimate contact (touching) and that the promoter aided hydrogen activation. Then it was proposed that the crystallites may not need be in direct contact as hydrogen spillover to the support could accomplish the same objective (see Fig. 16b). However, slowly it became apparent that the promoter was not effective as a separate sulfide crystallite but was actually only effective if it was present in some form on the surface of the MoS, crystallites (1-3). An early proposal suggested that the promoter is bonded to the support, which would lead to higher stability of a deposited MoSz monolayer (50), as illustrated in Fig. 16c. However, the chemistry was subsequently found to be more subtle.

POLYAROMATIC SULFUR COMPOUNDS

393

The nature of the surface promoter species has been debated for many years. As MoSz crystals exist as layered structures, two models evolved. One proposed that the promoter was intercalated deep within the bulk of the MoS2 layers (51) and another proposed that bulk intercalation was not thermodynamically stable and a “surface-intercalated” structure was more likely (52) (see Fig. 16d). Both proposals related the promotion to crystal surface reconstruction and solid-state chemistry and are valid only for multilayered structures. In fact, the most active species are now known to be monolayers and short layered stacks of small MoS2 crystals bound to the support surface, and the promoter decorates the edges of these small crystallites (I, 2, 53), as illustrated in Fig. 17. Instead of the active species being “promoted” MoS2, it is now known that the most active catalytic sites are in fact the added Co or Ni in the form of a new species, in which the Co or Ni is bound to the surface edges of very small MoS2 crystallites through sulfide bridges. It has also been established that the basal plane plays no catalytic role in HDS ( I , 2). Though many metal sulfides have been investigated as promoters, only Co and Ni have been found to be particularly effective. The chemical origin of this specific promotion by Co or Ni is discussed later. These unique species have been given many names. Originally, Topsoe referred to them as Co : MoS2 (53),but later changed the terminology to Co-Mo-S (11) (Fig. 16e). The exact stoichiometry of this Co-Mo-S was not originally defined, but it was established that the observed maximum in stoichiometry of added Co or Ni to MoS2 (Co/(Co + Mo) = 0.3) was related to the point at which all of the available edge positions were covered by the Co or Ni. Another name that has been used is “sulfide bimetallic species” (SBMS), coined by Startsev (2); it refers to the active species as a surface-bound bimetallic species having the stoichiometry of Co/Mo = 1/2 in which Co is bonded to two surface Mo atoms through sulfide bridges, as illustrated in Fig. 16f. There appear to be two or three types of sites having these structures. These have been termed “edge” and “corner” sites and relate to the position of the promoter atom on the periphery of the MoS2crystallite. Unfortunately, the term “edge site” has been used by different authors with slightly different meanings (46, 53-60). As discussed earlier, Chianelli studied unpromoted MoSz and his “edge site” definition refers to Mo sites on the exterior surface MoS2 crystallites on which the site is found on an internal layer of a stack (see Fig. 16a). Topsoe defined “edge site” as any promoter bound to surface Mo in the plane of a MoS2 slab (see Figs. 16e and 16f). A complication in this terminology is that some of the most effective commercial HDS catalysts used today contain large portions of the promoter bound to the surface of monolayers of MoS2. In this instance, there

394 D. DUAYNE WHITEHURST, TAKAAKI ISODA, AND ISAO MOCHIDA

a

b

Promotion of HDS activity is linked to the Co atoms present as Co-Mo-S

L

0

0.1 0.2 0.3 0.4 C Co in Co-Mo-S (mmol g-’)

FIG.17. The many phases of cobalt in commercial catalysts (57, 65, 66). Reproduced with permission of cited references.

would be no “edge sites” as defined by Chianelli (only rim sites) (46),but in Tops~e’snomenclature the edge site terminology is still appropriate (53-60). However, the simple conclusion that Chianelli’s edge sites are fundamentally different from Topsoe’s edge sites may not be valid, since Topsoe showed conclusively that the original active Mo sites in the MoS2

POLYAROMATIC SULFUR COMPOUNDS

395

crystallites are, in fact, replaced (or covered) by the promoter, producing the new Co(Ni)-Mo-S species, which have higher activity than the original Mo sites they replaced (11, 53). This matter is more important than one of nomenclature, in that Topsae postulated that the edge sites (of his definition) are responsible for hydrogenation activity (kHs,). Chianelli, on the other hand, concluded that the rim sites are responsible for hydrogenation and his edge sites are active for direct sulfur extraction. Topsae proposed that “corner sites” are responsible for direct sulfur extraction (kD,)(53-60), but the exact nature of corner sites is not known. What is known is that the active sites for sulfur removal constitute only about 10% of all of the Co(Ni)-Mo-S sites as identified by Mossbauer emission spectroscopy (MES) (57). Thus, there is something special about some of the Co-Mo-S sites. Further study in this area is greatly needed to clarify this issue, and it is recommended that, in the future, authors use terminology in a uniform manner. Some suggestions for standardization are made in later discussions. To further confuse the issue, there also appear to be two different types of Co(Ni)-Mo-S sites having different activities for both direct sulfur removal and hydrogenation. These are commonly classified in the literature as Type 1 and Type 11, as originally defined by Topsae (11, 55-57). Type I sites are found on monolayer MoS2 slabs and Type I1 sites occur on multilayered slabs (I, 55-57). Because of either geometric or electronic factors induced by the support, Type I sites are less active than Type I1 sites. Preference for their production appears to be related to the reagents and conditions used in the catalyst preparation. Type I sites were found to transform into Type I1 sites at high sulfiding temperatures (>875 K) (55, 56). More detail about the nature of Type I and Type I1 sites is provided in the sections on catalyst preparation methods and the nature of the active sites. 2. Detailed Characterization of Co(Ni)-Mo-S Sites With improvements in the preparation of more active HDS catalysts, MoS2 crystallites became smaller, and traditional physical techniques for characterization such as X-ray diffraction (XRD), scanning electron microscopy (SEM), and transmission electron microscopy (TEM) became limited. In fact, today’s best catalysts do not exhibit XRD patterns, and the active catalyst particles can no longer be observed directly by TEM. Thus, new techniques were required to provide structural information about Co(Ni)Mo-S catalysts. As modern surface science characterization procedures evolved, they were immediately applied to the study of CoMoS,-based

396 D. DUAYNE

WHITEHURST, TAKAAKI ISODA, AND ISAO MOCHIDA

catalysts with the hope that they could aid in the elucidation of the structure of these unusual surface-bound species (52-66). The first breakthrough came

from Mossbauer emission spectroscopy (MES) of 57Co-labeledmaterials. With this technique, Topsoe and co-workers observed that all highly active HDS catalysts exhibited a unique doublet MES (54-60). Structures exhibiting this MES doublet have now been observed by many different researchers and have received many names in the literature, as discussed earlier ( I , 2). Examples of this doublet can be seen in Figs. 17a and 18a (56-61). In commercial catalysts, not all of the promoter added to the formulation results in the formation of the unique Co(Ni)-Mo-S species. Figure 17a illustrates the various materials that have been identified in commercial Co-Mo catalysts. The figure also shows how these various forms of cobalt a

1.98 -

Intensity (-lo6counts)

SEQ Co, Mo

-)/m ‘I

1.881.95 -

b 1.91 7 2.73 -

7

NTAA

CoMoS

2.72;

1.01-

0.97 1.85 1.81-

-6

-4 -2 0 2 4 Doppler velocity (mm s-’)

Catalyst structure (MES) (61)

6

YOCo in CoMoS Catalyst activity (61)

FIG.18. Relationship of preparation method to Co(Ni)-Mo-S/Al*O, structure and activity. (a) Catalyst structure (MES) (61);(b) catalyst activity (61).SEQ = sequential impregnation: COIM = coimpregnation. Modified and reproduced with the permission of Ref. 61.

POLYAROMATIC SULFUR COMPOUNDS

397

may be distinguished by their MES (57, 65, 66). A t low levels of added cobalt, support interactions predominate and the cobalt associates with the alumina, forming spinels. As more cobalt is added, it preferentially interacts with the edges of the MoS2 crystallite surfaces, forming the desired active species. When all of the edges are covered, cobalt forms a separate phase of the stable sulfide Cogss. Only the Co-Mo-S structure is active in HDS, as shown in Fig. 17b, where the deconvoluted MES spectral area of the Co-Mo-S is plotted against the observed activity of the catalysts for the desulfurization of thiophene. Similar correlations were found for hydrogenation of butene (55, 56). Thus, the observed HDS activity of different catalysts as a function of the CoiMo ratio will vary depending on the distribution of the cobalt in these various forms. It can be seen that the degree of dispersion of the MoSz will be an important factor in determining this distribution. Optimal ratios have been reported to vary in the range 0.2-0.8, depending on the author and method of preparation ( I ) . In an ideal structure such as SBMS, the ratio should be 0.5 (2). The absolute amount of Mo and Co or Ni is also important. As discussed later, the best catalysts are produced by procedures that provide a complete monolayer coverage of the support surface by the Mo in oxidic form ( I ) . It is only within the past 4 years that the detailed structural elucidation of the Co(Ni)-Mo-S species has been achieved through the use of highenergy techniques such as X-ray photoelectron spectroscopy (XPS), extended X-ray adsorption fine structure (EXAFS) spectroscopy, and X-ray absorption near-edge structure (XANES) spectroscopy (61-64). These characterization procedures have been complemented by newly developed high-resolution transmission electron microscopy (HRTEM), which can now provide microscopic resolutions in the range of 10 A or smaller. Through the use of these techniques, it has been shown that the MoSz portion of the Co(Ni)-Mo-S complex has a structure essentially identical to that of conventional MoS, (slabs of Mo sandwiched between sulfur layers that constitute the basal planes of the layered structure). EXAFS has provided detailed information about the local environment of the active Co and Ni sites and the Mo atoms to which they are attached in terms of the types of atoms within two atomic shells away from the atom being characterized. Cobalt and nickel were shown to be definitely bonded to the surfaces of small MoS2 crystallites. Representative structures for the environments of Mo and Co are illustrated in the following diagram. In such structures, Mo has a coordination number (CN) of 6, with six nearneighbor sulfur atoms, three nearby Mo atoms, and one nearby Co or Ni atom. Co-S configurations were either CN = 5 (square pyramidal) or CN = 6 (octahedral), with either one nearby Mo atom (low HDS activity) or two nearby Mo atoms (high HDS activity) (62).

398 D. DUAYNE WHITEHURST, TAKAAKI ISODA, AND ISAO MOCHIDA

3

CO-Co C N = O

Mo-Co C N = 1

Co-Mo CN = 2

Mo-S

Co-S

Mo-MO CN

CN=6

CN = 6

Nickel, on the other hand, on alumina and on silica supports was found to have only five nearby sulfurs (square pyramidal) with Ni-Mo coordination numbers from 1 to 1.5. Ni-Mo-S supported on carbon was observed to have Ni-S coordination numbers of 6 in a trigonal-prismatic configuration. In addition, Ni (at low Ni concentrations) was found to have one nearby Ni, which could indicate that, in some catalysts, Ni is present as pairs on the MoS2 surface. The overall structure of the Ni-Mo-S was believed to be similar to that of millerite (i.e., Ni is located in the center of the MoS edge in a square-pyramidal configuration, with one sulfur extending perpendicular to the surface) (62-64). One caution is offered with respect to the preceding conclusions. Although all of these data were obtained after the catalysts were sulfided at relevant temperatures, prior to analysis, all of the samples were flushed with He or H2 at about 400°C. Any H2S not strongly coordinated could have been removed in this way, and the CN observed on analysis could be lower than that in the operating catalyst. Some indication of this potential problem has been reported, whereby the CN was observed to vary with the composition of the flushing gas (63). It is recommended that in future studies of this nature, researchers should try to retain H2S coordinated to the operating catalyst. 3. Sensitivity to Preparation Method

It seems quite surprising that a structure as sophisticated as Co-Mo-S or SBMS could arise by merely coimpregnating salts of Mo and Co or Ni on alumina and then calcining and sulfiding, as is done in many commercial catalyst preparations. The most thermodynamically stable species under HDS conditions are the individual sulfides MoS2, CoyS8, and Ni3S2 (I). The interaction of the promoter with the MoS2 edges must be very strong as it forms readily and survives for long periods of use in HDS processes.

POLYAROMATIC SULFUR COMPOUNDS

399

It has been proposed that in the oxide state, prior to final sulfiding, cobalt having an octahedral configuration is very selectively converted to COMo-S (57).Thus, the reagents used, the sequence of metal incorporation, and the calcination and sulfiding temperatures employed are all critical in producing catalysts of high HDS activity. Much of the early literature on promoted catalysts followed the experience gained from preparation of MoS2-onlycatalysts, whereby high activity was associated with high Mo dispersion and small crystallite size. The most active materials per Mo atom resulted from preparations producing complete monolayer coverage of the support with the Mo in the oxide form (67-75). It was also shown that incomplete sulfiding produced catalysts with lower activity. Molybdenum bonded to oxygen is difficult to reduce from + 6 to +4 and any Mo in the final catalyst containing Mo-0 bonds is believed to be inactive for HDS (67-75), for example, surface-bound species such as A1-0-Mo. The reagents used, the sulfiding gas composition, and the sulfiding temperature are all important to the activity of the finished catalyst. Quite surprisingly, it was found that hydrogen is not the major contributor to reduction of Mo from + 6 to +4. Hydrogen sulfide alone is sufficient, and small amounts of water in the sulfiding gas (3%) actually aid the exchange of sulfur for oxygen and reduction. Under wet conditions, sulfiding can be accomplished at moderate temperatures (400-500 K). Under dry conditions, sulfiding requires higher temperatures (>700 K) (71). Prereduction of the molybdenum actually makes the sulfidization more difficult (75).Thus, there is much art in the preparation of just the supported MoS2 crystallites, but with the proper techniques, the support surface can be covered with very highly dispersed MoSz having crystallite sizes as little as 10 A. For perspective, this amounts to associations of as few as seven Mo atoms. The relevance of this to HDS of PASCs is discussed in the next section. Producing MoS2in high dispersion is only one requirement in the production of promoted catalysts. The methods used to achieve selective bonding of the promoter to the MoS2 crystallite surface have also developed into a fine art. It has been shown that sequential incorporation of Mo and then Co(Ni) provides more active catalysts (56),and the activity may be further improved by using organometallic reagents in which the Mo is specifically bonded to the surface of the support and sulfided and the Co or Ni is then bonded to the Mo with organometallic reagents (2). Thus, there does seem to be promise for producing catalysts of improved activity by developing improved catalyst preparation procedures. Recent studies have shown that, by careful selection of the starting reagents and solvents, it is now possible to synthesize Co(Ni)-Mo-S species on any support without the requirement of sequential impregnations (61).

400 D. DUAYNE WHITEHURST, TAKAAKI ISODA, AND ISAO MOCHIDA The ratio of promoter ion to Mo ion is still important in such syntheses as excess promoter (e.g., Co) can result in the stoichiometric formation of C o M o 0 4 . Unfortunately, the Co in this compound is converted primarily into inactive Cogs8on sulfidation. A t the appropriate reagent ratio, using ammoniacal solutions of salts, one can achieve near 100% Type I Co(Ni)Mo-S. However, with phosphoric acid-containing solutions, high yields of Type I1 are obtained, which have higher HDS activities. This observation may be closely related to the fact that Mo exists as MOO:- at p H values >9, whereas at p H less than 5 , Mo exists as polymeric anions, such as Mo70$; or Mo80;; (76, 77). Phosphate is also known to modify the surface of alumina by bonding to A1-OH groups; this lowers the potential for MoS2 bonding to the surface (61, 76) and induces the formation of multilayered Co(Ni)-Mo-S (Type 11) structures. Titania-supported Co(Ni)-Mo-S catalysts are reported to be more active than alumina-supported catalysts, supposedly because support interactions are not as strong ( I , 78). The most selective synthesis reported for high-activity catalysts (100% Type 11) utilizes a multidentate ligand additive (nitrilotriacetic acid (NTAA)) in the formulation (61). Figures 18a and 18b show the relationships between preparation method and catalyst structure and activity, ‘respectively. Type I1 catalysts were observed to be about two times as active as Type I catalysts, and alumina- and silica-supported catalysts exhibited the same specific activities for HDS (see Fig. 18b). For perspective, present commercial catalysts are prepared by procedures very similar to those shown in Fig. 18b for ammoniacal solutions or phosphate-modified preparations. Using the NTAA procedure, catalysts supported on carbon have been found to be about two times more active than similarly prepared catalysts on silica or alumina supports (79). Thus, it appears that catalyst syntheses are reaching near-optimal activities, which are about four times those of present-day commercial catalysts. This activity may, however, not be high enough to achieve future 0.05% sulfur specifications for fuels through catalyst replacement alone. Novel approaches may lead to more active catalysts. Structures like the SBMS model closely resemble the structures of heterogeneous catalysts made by chemically bonding soluble “homogeneous” organometallic complex catalysts to surfaces. Such chemically bonded (attached) metal complexes are known to exhibit the same catalytic activity as their soluble counterparts but can be used in fixed-bed, continuous-flow reactors in either the liquid or gas phase (80-84). In addition, such chemically bonded surface complexes have the potential for a much higher activity per unit reactor volume than their homogeneous counterparts, as they have no solubility limitations. Organometallic complex catalysts are generally limited to less than 10-3-M solutions, whereas their surface-bound heterogeneous analogs

POLYAROMATIC SULFUR COMPOUNDS

401

can function at the equivalent of 1-5-M concentration in fixed-bed reactors (81, 82). By comparison, commercial HDS catalysts in fixed-bed processes contain 1-2 mol of Mo and 0.5-1 mol of Co or Ni per liter of reactor volume. The principles involved in the synthesis of surface-bound complex catalysts could very well offer guidance in producing improved HDS catalysts. The pioneering work of Startsev and co-workers in this area is quite commendable, and more work of this nature is highly encouraged. Some excellent reviews of the principles and applications of catalysts of this nature are available for reference (83, 84). With homogeneous organometallic catalysts, the nature of the attaching ligand greatly influences the activity and selectivity of the catalyst, both sterically and electronically. With their anchored counterparts, the attaching ligand plays an important role in extending the catalytic site away from the support surface, which helps to avoid steric limitations imposed on intermediates by the solid surface (81, 84). Thus, the syntheses of SBMS type catalysts by Startsev were based on these principles (2). The active Co or Ni site is electronically influenced by the underlying MoS2 (ligand), and the MoS2 crystallite extends the active site away from the support surface so as to allow sterically unimpeded progress through the catalytic sequence. The importance of this feature is expanded in the section on geometric considerations. Using the analogy of supported organometallic catalysts, a model of a surface-bound HDS catalyst is illustrated in Fig. 16g. The chemistry of organometallic metal complexes is well understood, and many researchers have sought analogy between HDS catalysts (SBMS) and soluble organometallic complex chemistry (4).Anchored organometallic catalysts offer a bridge between these two disciplines and may point the way to improved HDS catalyst preparations. 4. Electronic Basis for Promotion As discussed in the preceding section, only Co and Ni have been found to be effective in enhancing the activity of commercial HDS catalysts. There have been several theories put forward to explain why these two metals are so unique (1-3, 67, 85-88). All of these assume that the active catalytic site is the Co(Ni) bonded to the MoSz crystallite surface through sulfide linkages. The enhanced activity is related to the strength of the peripheral Co(Ni)-S bond, which is influenced by electron transfer from the Co(Ni) to the 4d shell of the underlying Mo. This phenomenon is different for different promoting metals, and periodic relationships between HDS activity and the various metals have been observed (67, 85-88). Although there are differences in the assumed important parameter responsible for this effect, all of the proposed theories exhibit good correlations between the assumed important parameter and the HDS activity, as shown in Fig. 19.

402

D. DUAYNE WHITEHURST, TAKAAKI ISODA, AND ISAO MOCHIDA

4

-

Measured Calculated

-

800

-

600

-

400

-

200

3-

2-

1-

-2.0

. A

2 -

"

'

OL

6

'

I

.

NiMoS COMOS

-3.0-

0) x

a,

'

-1

-

-2

-

-4.0 -

-0

'0

3 L

-5.0

c r?

1sf, -

0' ;

Ta W R ; ;I A ; Nb Mo Tc Ru Rh Pd Ag V Cr Mn Fe Co Ni Cu

- 20 - 40

-3

-

-7.0 2

3

4

5

6

7

8

9

10

Number of d - electrons

Minimum M-S Bond Energy Tops+ etal. ( I , 85, 86)

Optimal LIH, of M-S Bond Startsev etal. ( 2 )

FIG. 19. Periodic trends in HDS promotion. Reproduced with permission of cited references. (Ref. I , with permission from Springer-Verlag.)

The various theories put forth are as follows:

1. Optimal Heat of Formation of M-S Bond [Chianelli et al. (3, 87)]. High activity is the result of a balance between the strength of the H&M bond and the strength of the thiophene S-M bond. If the H2S-M bond is

POLYAROMATIC SULFUR COMPOUNDS

403

too strong, the H2S will not be released. If the thiophene S-M bond is too weak, the sulfur compound will not be adsorbed. 2. Minimum M-S Bond Energy [Topsoe et al. (1, 85, 86)]. Low M-S bond energy allows for the maximum number of unoccupied coordination sites for reaction. 3. Method of Interesting Bonds [Startsev et al. (2, 88)l.A proper balance of M-S bond strengths is required. M-S bonds that are either too strong (Ti-S) or too weak (Cu-S) will correspond to low activity, 4. Density Function Electronic Structure [Smit and Johnson (67)].High activity is the result of electron transfer from Co(Ni) to Mo and corresponds to the removal of (T antibonding metal d-sulfur 3p electrons from Co(Ni). Optimal activity is the result of the right oxidation state of the Co(Ni) and the M-S bond length. The significance of these various theories to the problems addressed in this review relates to the probability of discovering new catalysts with significantly higher activity than today’s known catalysts. Many combinations have been tried, but as yet no significantly higher activities have been found, and many combinations were actually observed to be antagonistic to HDS activity (1-3, 67, 85-88). If the optimal catalytic properties for the promoter-base metal sulfide combination are already achieved with today’s commercial catalysts with the specific Co-Mo-S or SBMS structures, then the only way in which catalytic activity can be increased is to generate as many of these structures as possible. If there is still room for improvement as suggested by the model of Topsae, then intrinsically more active catalysts are possible. Work by Topsae does indicate that only about 10% of the Co(Ni)-Mo-S species, identified by MES, are actually active for sulfur removal (56, 57). Thus, there may be the possibility of increasing activity by increasing the proportion of the truly active form of Co(Ni)-Mo-S in the catalysts. Alternatively, it is known that for monometallic sulfides, Rh and Ir are almost two orders of magnitude more active than Mo, Co, or Ni sulfides. However, today’s promoted catalysts are approaching the activities of these highly active sulfides on a per gram of catalyst basis. Attempts to promote these more active monometallic metal sulfides have so far not provided exceptional catalysts. Unfortunately, these highly active species are also highly expensive, and exceptionally high activities with good stability would have to be achieved for their commercialization. The area of identifying new promoter-base sulfide combinations with very high HDS activities is a very important area for future research in view of the need for considerably more active catalysts in the future.

404 D. DUAYNE WHITEHURST, TAKAAKI ISODA, AND ISAO MOCHIDA 5. Number of Active Sites As discussed in the previous section, an important question is, how many sites are actually present in today’s best catalysts? This question has been addressed by many researchers throughout the world for many years and so far there is no definitive answer. The main approach to counting the number of sites has been to try to relate some chemisorption behavior with observed HDS activity. Thiophene HDS is generally the reaction of choice ( I ) ; unfortunately, with this molecule, the HYD sites cannot be estimated reliably, and as discussed earlier, for dialkyldibenzothiophenes the hydrogenation route is preferred. Reagents used for chemisorption measurements , H2S,NO, H 2 ,and thiophene ( I , II,68-70). Interprehave included 0 2CO, tation of such studies has been the subject of considerable debate, but reports by Burch and Collins did show some very interesting correlations between C O sorbed at room temperature and HDS activity. On the basis of their observed CO sorption capacities of about 30 pmol/g and making the assumption that the average MoS2 crystal size was about 25 A (33 Mo atoms), they estimated that there was only one highly active (promoted) site per crystallite and that the activity of this unique species was much higher than that of an equivalent site in unpromoted MoS2 (68). Topsoe and Topsoe investigated NO adsorption on HDS catalysts (11) and observed higher absolute levels of adsorption than did Burch and Collins. However, they were able to distinguish the amount of NO adsorbed on Co or Ni from the amount of NO adsorbed on Mo by infrared spectroscopy. As Ni or Co was added to the catalyst, the total NO sorption capacity increased. Interestingly, however, they observed that as Co or Ni was added to the catalyst formulation, some of the sites on Mo that had adsorbed N O originally were lost and replaced by sites on Co or Ni that adsorbed NO. These “replaced” sites were believed to be those created when Co(Ni)Mo-S (or SBMS) type sites were formed. The activity increase for thiophene HDS was found to correlate well with the number of these newly formed Co(Ni)-Mo-S sites. On the basis of the amount of NO adsorbed and estimations, by EXAFS, of the number of edge sites in the MoS2 crystals, Topsoe’s group concluded that only about 10% of the edge sites were active (56, 57). These calculations led the authors to conclude that there were two different forms of Co-Mo-S sites, which they called “corner sites,” which had high desulfurization activity, and “edge sites,” which had low desulfurization activity (56). For comparison of these results with results of other authors, one may assume that the NO-Mo and NO-Co or NO-Ni infrared adsorption coefficients are the same. With this assumption, one can transform Topsoe’s data into a form resembling those of Burch and Collins. Figure 20 shows such

POLYAROMATIC SULFUR COMPOUNDS

405

1.4

1.2

2 t

1

B

t

Lo 0.8

.-

I

.-

2 v1

@

0.6 0.4

0.2

0 0

10

20

30

40

50

60

Micromoles of NO on PromoterlCrystallite

FIG.20. Increase in HDS activity with added promoter. Data recalculated from Ref. 56.

a transformation. It should be noted that both Co and Ni sites have very similar activities per site and that a maximum NO adsorption occurs at about 60 ,umol/g. It is believed that 2 mol of NO is adsorbed on each site or NO is adsorbed as a dimer (11). Thus, again one must conclude that there were about 30 pmol/g of sites and, assuming the average MoS2crystal size was about 20 A, there again must have been only one active site on each crystallite. This new site is calculated to be about 30 times more active than the original Mo site that it replaced. The agreement between this result and the earlier proposal by Burch and Collins (68) is quite remarkable, although probably fortuitous. Other authors using EXAFS measurements of the concentration of active Co atoms in a pure Co-Mo-S/A1203 Type I1 catalyst have estimated that the promoted active site is about 22 times more active than a Mo/A1203site (62). Further studies in this area are highly encouraged as knowledge of the number of active sites can offer guidance as to the ultimate potential activity of new catalytic materials. Unfortunately, thiophene has been the prime indicator of HDS activity for the majority of the reported correlations between catalytic activity and catalyst structure. This is not an appropriate model compound for determining how many active hydrogenation sites are present in the catalyst

406 D. DUAYNE WHITEHURST, TAKAAKI ISODA, AND ISAO MOCHIDA because the direct sulfur extraction route is orders of magnitude greater than that of hydrogenation. As discussed in the previous section, with 4methyldibenzothiophene or 4,6-dimethyldibenzothiophene,the relative values for the direct sulfur extraction (kD,,) and the hydrogenative route for desulfurization (IcHsI) are close in value, and so their individual contributions are easier to determine than with thiophene. It is suggested that, with these more hindered molecules, it may be possible to relate the various rate constants to the sorption capacities of appropriate reagents and determine the number of active sites for both hydrogenation and direct desulfurization. Such studies could provide guidance in the synthesis of new, more active catalysts for the deep desulfurization of gas oils. 6. Geometric Considerations In previous sections, it was shown that with the best preparation procedures available today, it is now possible to make HDS catalysts having MoS2 crystallites of as small as 10 A,or assemblies of seven Mo atoms. Type I Co-Mo-S crystallites are single-layer crystals having a thickness of Type I1 Co-Mo-S crystallites have been prepared as stacks of about 6 small crystallites with a height to diameter ratio of about 1.5-3 (or 3-5 stacked layers 20-30 A high) (62).This dimension should be compared to the molecular dimensions of the PASCs they are to convert. For example, dibenzothiophene (DBT) has dimensions of approximately 3 X 8 X 12.2 or about the same size as a Co-Mo-S Type I crystallite! With alkyl substituents, the thickness is slightly larger and the width is closer to 10 A. Thus, one must stop to consider how much steric restriction is imposed on the reacting molecules by the surface of the support, particularly with those Type I crystals that lie flat on the support surface. Chianelli was one of the first to point out the importance of crystal geometry when considering desulfurization of larger molecules such as DBT (46, 47). As an example, Fig. 21 illustrates several Type I and Type I1 crystallites and 4,6-dimethyldibenzothiophene (4,6-DMDBT) molecules in their approximate sizes. The various definitions of types of sites by different authors are also included. It can be seen in this figure that alignment of 4,6-DMDBT along the edge of a Type I crystal is quite limited geometrically. If the interaction is not in the plane of the MoSz, the limitation will not be so severe, but access to the active site is still limited to approach from only one side, and the reacting molecule cannot approach the catalyst perpendicularly, as the molecule is much wider than the Type I layer thickness. Type I crystals, standing perpendicular to the alumina surface, allow a higher probability site access, but this crystal alignment would most likely occur in crystals bonded to the alumina surface by Al-O-Mo or A1-S-Mo bonds.

A.

A,

POLYAROMATIC SULFUR COMPOUNDS

407

FIG.21. Geometric considerations in the HDS of dialkyldibenzothiophenes.

As discussed earlier, if bonding occurs through A1-0-Mo linkages, the Mo sites in that vicinity will most likely be inactive as they are more difficult to reduce to a low-valent state. The sites on the Type I1 crystal, illustrated in Fig. 21, are much more accessible. Even about 75% of all of the edge sites can be approached by a 4,6-DMDBT molecule in a perpendicular alignment with the alumina surface. Thus, for either geometric or electronic reasons, Type I sites should be expected to have lower activity than Type I1 sites. The higher activity of Co-Mo-S supported on carbons may also be due to increased access to the active sites, as it is postulated that the MoS2 slabs

408 D. DUAYNE WHITEHURST, TAKAAKI ISODA, AND ISAO MOCHIDA in such catalysts are aligned parallel with the basal plane of the graphite layers (2). The EXAFS analysis of carbon-supported Co-Mo-S materials indicates that their structure is more consistent with a single-layer slab than a stack, even though the activity is much higher (61, 62). Thus, either the bonding to carbon surfaces does less damage electronically or the geometry of such catalysts is more favorable for reactant access. A definitive answer to this question would be very useful in assessing the potential for improved activity with new Co-Mo-S catalysts.

7. The Two-Site Dilemma Throughout the previous discussions, HDS catalysts were described as containing two different types of catalytic sites, one that facilitates direct sulfur extraction and another that facilitates hydrogenation. This could easily be rationalized in catalysts of a few years ago wherein the distribution of the promoter in the catalyst surface was uncertain, the crystals of MoS2 were large, and the composition of the support was variable. However, as catalysts have been improved, the crystallite sizes have been reduced to as small as seven Mo atoms in a cluster, and the stoichiometry of promoter to Mo is optimized at 1/2. The surface of the support is now carefully controlled, and the stacking of MoSz can be dictated with reasonable accuracy. With such improved catalysts, it now becomes difficult to surmise how two different types of sites can exist, each with a different composition and function. The difficulty may be seen by considering what is presently believed to be the structure of a Type I Co(Ni)-Mo-S crystal. In the most active form, it consists of a cluster of about seven Mo atoms (61,62).Building a molecular model of such a structure is quite informative. Assuming this cluster has a structure and composition analogous to those of conventional MoS2, then the Mo atoms are six-coordinate, having trigonal-bipyramidal geometry, and the Mo is sandwiched between two inert layers of sulfur atoms, which constitute the basal plane or the 0100 plane. An illustration of such a structure is presented in Fig. 22a. The illustration actually contains more than seven Mo atoms as it is easier to visualize. It can be seen that the edges have two different arrangements of Mo-sulfur bonding. The TO10 face is composed of sulfur atoms that bridge adjacent Mo atoms, and the sulfur atoms form a square around the Mo atoms. There are no uncoordinated orbitals on the Mo on this face and therefore no place to conveniently bond to Co or Ni or to provide an -SH group, if necessary. Bridging >SH groups such as those that were described in the section on organometallic clusters are possible on this face.

POLYAROMATIC SULFUR COMPOUNDS

409

-0 1 0 1

FIG.22. (a) The three faces of MoS2. (b) A possible representation of a Co-Mo-S site.

The 1010 face, on the other hand, has alternating bridged sulfurs and two vacant orbitals on Mo that extend away from the edge. These are available for interaction with other species. If these orbitals are not occupied by some ligand, they would represent a coordinately unsaturated site (cus) in organometallic complex terminology, or a vacancy in the conventional heterogeneous catalysis terminology (see Fig. 23). In this configuration, a cluster of seven Mo atoms forms a hexagonal flat crystal with three identical bonding positions at the three 1070 faces. Each Mo has four orbitals extending away from the 1010 face, forming essentially a square of four sulfide groups on which the Co(Ni) can bond. This face is the one where most researchers believe SBMS or Co-Mo-S species are formed ( I , 2). At a

410 D. DUAYNE WHITEHURST, TAKAAKI ISODA, AND ISAO MOCHIDA Sequential deS

S1 I H Mo-S-

7

Simultaneous deS

7 y

H

H\

MO-S-

Hydrogenation

sI

Mo

H\SI Mo-S-MO-S-MO

i]

H\ s I I M 0- S- Mo -S-MO

slH

I \

n

s

uI

II Mo-S-Mo-S-MO

I

S S S I1 I1 II MO-S-MO -S-Mo

H\

+ H H\ /

S/H S I I I Mo - S - M O - ~ - M ~ S

H

H\

S

I Mo-S-Mo

-

H2S

I

s/ -S-Mo

I

H\

+H*

0

S I I Mo-S-MO-S-MO

O=Vacancy

FIG.23. Vacancy models of the HDS mechanism.

S II

H

s‘ I

POLYAROMATIC SULFUR COMPOUNDS

41 1

2/1 stoichiometry of S/Mo, for a seven-Mo-atom cluster, there will be 12 sulfurs in the basal planes and 12 extending empty orbitals. All of the Mo atoms on the 1010 edge would be identical. To achieve the appropriate SlMo stoichiometry, two additional bridging sulfurs connecting four of these orbitals would be needed, and the remaining eight orbitals would be vacant (cus). This structure is somewhat strained, and the bridging -S- on the 1070 faces would most likely be easily broken by the dissociative addition of either H2 or H2S. If the Co or Ni were to be bonded to these extended orbitals through sulfide bridges, a SBMS site could be generated. It may be that the two different sites noted in the literature ( I , 2) arise from just how the Co(Ni) links to the MoS2 crystallites through these peripheral -S-bridges. This question will now be addressed. To have enough coordination positions on the Co(Ni) available for catalysis, the number of connecting linkages to Mo should not be excessive. The complex must also be very stable or the catalyst would not withstand the rigors of continued use in commercial processes. It is known that the Co(Ni) is associated with two Mo atoms (62);thus a single linking bond attachment will not be considered. Likewise, four linkages (possible only for squarepyramidal geometry, CN = 5 ) leave only one vacant orbital on the Co(Ni) and would not be expected to allow either dihydrogen dissociation or oxidative addition to progress mechanistically. Thus, only two or three linking -S- structures seem reasonable. Model structures using octahedral Co(Ni) indicate that all bonding between the four extending orbitals on each 1010 face must be through adjacent positions (cis) on the Co(Ni). Trans bonding is possible for some configurations with larger clusters of Mo or stacked crystals. As the concern is with the Type I Co(Ni)-Mo-S site, discussions here will concentrate on cis attachment. Models shows that for two attaching linkages between Mo-S- and Co or Ni, there are three possible geometric configurations. There would be four orbitals available for bonding to reactants. Thus, modes of bonding up to q4 for thiophene compounds are possible (see Fig. 24). Oxidative additions and hydrogen dissociation could proceed readily. However, in all of these two-linking-bond configurations, there is always a potential -SH group in very close proximity to one of the Co(Ni) bonding orbitals. It is believed that this would readily lead to three linking bond attachments as the most stable configuration for the SBMS species. Unfortunately, if this is true, no matter how the Co or Ni attaches to these three linkages, the structure of the resultant SBMS complex will have only one geometry! An illustration of this type of site geometry is provided in Fig. 22b. In this figure, thiophene is shown bonded to one of the three nonlinked orbitals of Co, and a hydride and H2S molecule occupy the other

412 D. DUAYNE WHITEHURST, TAKAAKI ISODA, AND ISAO MOCHIDA

Q M

q5

C-J pM

r16

s

M

w l ' c , 71) c=c

p-1'

s. 1'

FIG.24. Bonding modes for thiophene compounds.

two orbitals. The nonbridging orbital on Mo is shown as being occupied by an -SH group, which could be the source of a proton in proton-assisted mechanisms, as will be discussed later. There are three identical SBMS sites of this type possible on a crystallite containing a cluster of seven Mo atoms (one at each 1070 face). If all three faces contain one Co, the overall crystallite would have a Co(Ni)/Mo ratio of approximately 1/2, which is known to be optimal. However, all three will be equivalent in terms of the bonding relationship between Co and Mo. Each will be in very close proximity to a potential -SH group on one of the attaching Mo atoms, which could possibly participate cooperatively in either hydrogenation or C-S bond cleavage reaction mechanisms. The three remaining bonding orbitals on Co(Ni) can bond to thiophene reactants in modes up q3,and mechanisms such as those shown in Figs. 25 and 26 can all proceed without requiring that any of the three linking bonds be broken. Thus comes the dilemma. With such a structure, there cannot be two different types of sites with different geometries in terms of the Co-Mo-S or SBMS composition. At slightly larger crystallite sizes, it is possible to build structures that have high and low probabilities of proximity to other potential surface -SH groups and could constitute, in principle, two different kinds of sites. If the crystals take on a triangular morphology, rather than hexagonal, it is possible to construct structures that have only one to three of these sites per crystal by making the sides of the triangle the TO10 face. This would limit the potential number of active sites to only three per crystal, independent of the number of Mo atoms. Thus, the observation of Topsoe, that for crystals of about 20 A, only 10% of the Co-Mo-S sites on some catalysts

POLYAROMATIC SULFUR COMPOUNDS

413

Reversible Coordination and Ligand Replacement

Hydrogen Dissociation (activation)

- u

Hydrogen Transfer to an Unsaturated Bond

-

f

Hydrogen Addition

(endo)

Sulfur-CarbonBond Cleavage

(Melallaliuabenrene)

FIG.25. Important HDS reaction steps demonstrated with organometallic complexes.

are active (56, 57), could have a geometric origin. Similarly, the conclusions of Burch and Collins, that their catalysts had only one active site per crystallite of 33 Mo atoms (68),could also be explained in this way. Stacking the small seven-Mo crystallites, as in Type I1 Co-Mo-S, allows another degree of freedom in which a catalytic site may bridge two MoS2 layers, and bridging may be cis or trans for an octahedral metal atom. Similarly,the probability of proximity to -SH groups becomes higher in stacked crystals. If one accepts the logic of the nature of the Co-Mo-S or SBMS site as described, a fundamental question arises. Can Type I Co-Mo-S sites, having only one geometric configuration, perform two different functions? If so, how is this possible? One possible explanation could be that the

414 D. DUAYNE WHITEHURST, TAKAAKI ISODA, AND ISAO MOCHIDA same site can bond to the reactant sulfur compound in one of two ways. Examining the structure shown in Fig. 22b more carefully, one can see that the three remaining bonding orbitals on Co(Ni) may not be geometrically equivalent. One is in close proximity to the remaining -SH group on Mo. We will refer to this orbital as Y. The other two orbitals are equivalent and remote from the adjacent -SH group. We will call these R. Thus, if the substrate coordinates to the Y orbital, close to the -SH group via the #S mode (as shown in Fig. 22b), cooperative mechanisms are possible. If it coordinates in this way to the other R orbitals, cooperative mechanisms are not possible. However, it is known that 7i-bonding is quite probable with thiophene compounds, and both types of orbitals may be involved. Examining the two orbital combinations shows that only two types of pairs are possible, X2 and nRR. Again, bonding that involves the Y orbital will have a neighboring -SH group for cooperative mechanisms whereas the RnR pair will not. Thus, one distinction that can be drawn between sites is proximity to the adjoining -SH groups. Another way in which the three-linkage model could perform two different tasks could be controlled by the thiophene molecule itself. The reactant could bond via the sulfur lone pair and exhibit one mechanistic route, or it could bond via the T system and exhibit the other mechanistic route. Organometallic analogs of HDS catalysts that have the same geometric structures but differ in the central metal or types of ligands have been reported to behave in two different ways (89).Thus, it is possible for one geometric configuration of a catalytic site to exhibit two different behaviors (due to the lability of different C-S bonds, steric effects, the presence of different ligands on the Co(Ni), etc.), and the preferred mode of conversion could be primarily dependent on the nature of the reacting molecules rather than the catalyst structure. In summary, there are several conceivable ways in which a seven-Moatom cluster having octahedral Co or Mo linked to the periphery by three -S- bonds could exhibit two different behaviors. The following lists the suggestions discussed in this section and some additional suggestions: The number of -S- linkages between the Co(Ni) and the MoS2 edge may vary. The number of orbitals on Co(Ni) involved in reactant coordination may vary. The orbitals on Co(Ni) involved in coordinating to the reactant may have different orientations relative to adjoining -SH groups. The reactant itself could dictate the mechanism by which it converts, depending on how it coordinates to the active site, e.g., if it prefers to bond

POLYAROMATIC SULFUR COMPOUNDS

415

to the Co(Ni) by more than one bonding mode, for example through S (77l-S) or through the system ($-C=C or q6). Even for equivalent sites, the oxidation state of the Co(Ni) may be variable and different mechanisms may be preferred. The number of available bonding orbitals on Co(Ni) could vary with conditions: (i) the CN could change from 5 to 6; (ii) at CN = 6 the complex could be either octahedral or trigonal bipyramidal; (iii) the number of vacant orbitals may vary if equilibrium adsorption is rapid and competitive between the sulfur compound and other ligands, such as H2S (see Fig. 22b), where one orbital is occupied by a coordinated H2S molecule. The nature of ligands on Co(Ni) other than the reactant could be different, such as H2S, -SH, -H, aromatic hydrocarbon, or a second reactant molecule. In the preceding discussion,we assumed that the Mo atoms had a trigonalbipyramidal geometry. If this changes for very small crystals, it may be possible that an octahedral geometry could be preferred. In such a case, the preceding arguments become even more confusing as all of the vacant orbitals on Mo become equivalent on all faces. The formation of two different SBMS sites could only occur on stacked crystals, as described earlier. At present, there is no answer to the dilemma that arises in the case of single crystallites of MoS2 containing very few Mo atoms. However, as the fields of organometallic complex chemistry and further improvements in the synthesis of supported Co-Mo-S or SBMS catalysts come closer and closer together, the answer may emerge. 8. Estimates of Potential for Increased Activity

To estimate how much improvement may be possible, we can take as a first estimate the conclusions of Burch and Collins (68) and/or the modified data shown in Fig. 20-only one active site per MoS2 crystallite. The worst scenario would be that today’s catalysts now have the maximum possible activity per site. In such a case, the only improvement possible is to increase the number of MoS2 crystallites on the catalyst surface or to increase the number of the special “corner sites” noted by Topsgie. In the Burch and Collins model, 33 Mo atoms per crystallite were assumed. However, improved preparation procedures have reduced crystal sizes in experimental catalysts to less than 10 or only about seven Mo atoms per crystallite. The maximum improvement would be to reduce the number of Mo atoms in an active crystallite to two (assuming the SBMS structures proposed by Startsev [2]). Thus, an improvement by a factor of only about 3 might be expected, relative to the most active reported catalysts.

A

416 D. DUAYNE WHITEHURST, TAKAAKI ISODA, AND ISAO MOCHIDA However, it is believed that in active A1203-supported catalysts (especially Type I) a portion of the MoSz is actually chemically bonded to the A1203surface through Mo-0-A1 bonds (55-57, 71-75). Were it not for such bonding, the crystallites would grow larger during continued use under HDS conditions and the catalysts would deactivate. Such bonded Mo-0-A1 species are difficult to reduce and most likely cannot function as HDS catalysts (55-57, 71-75). Thus, if one must sacrifice the activity of at least one bonding layer, then the potential for improvement is limited. Identifying supports that can bind Mo yet retain HDS activity for the bonding atoms is an important goal for new catalyst improvement. Carbon-supported Co(Ni)MoS, catalysts have been reported to have a higher activity per Co(Ni) than comparable A1203-supportedcatalysts (2, 4Ib, 57, 61, 62, 79), and this phenomenon may be responsible. Alternatively, higher activity for carbon-supported catalysts could be due to geometric considerations, as discussed earlier. If the Co(Ni)-Mo-S crystallites lie flat on the surface of silica and alumina supports but line up parallel with the graphitic layers on carbon supports, then carbons may provide less steric restrictions for adsorption or for transformations of reaction intermediates. Determining the actual number of active sites on carbon-supported catalysts would help to clarify this issue.

E. THE MECHANISM OF CATALYTIC HYDRODESULFURIZATION 1. General Considerations In the previous sections, we have discussed the pathways by which PASCs are desulfurized and what is presently known about the structure of the active species in supported Co(Ni)-Mo-S catalysts. In this section, we discuss the chemical reactions and intermediates involved in the catalytic sequence that results in desulfurization. This sequence is often called the catalytic mechanism. Several theories have been put forward to explain how the Co(Ni)-Mo-S catalytic species actually removes sulfur from very stable aromatic sulfur compounds. It is still not known whether one or more metal atoms are involved in the sequence or if peripheral -SH groups participate in the chemical transformations. Proposed schemes have ranged from (i) pairs of activated -SH groups attacking the sulfur compound with no participation of the metal at all to (ii) the requirement for multiple metal sites, each with its own function, to (iii) catalysis in which all transformations occur on a single metal atom in exact analogy to reactions of homogeneous organometallic complexes. Likewise, there is still no general agreement as to how dihydrogen is activated by Co(Ni)-Mo-S catalysts. It is still unclear whether dihydrogen dissociation is homolytic (giving 2 H ) or heterolytic

POLYAROMATIC SULFUR COMPOUNDS

417

(giving H+ and H-). These are quite basic questions and need to be clarified if new or improved catalysts are to be discovered.

2. Conventional Heterogeneous HDS Catalysts The fundamental transformations in HDS are the same for all mechanisms; only the details of how these occur will vary with the mechanism proposed. The primary phenomena involved are summarized as follows:

1. Adsorption (coordination) of the sulfur compound to the active site 2. Hydrogenation of unsaturated C = C bonds 3. Cleavage of two carbon-sulfur bonds (sequential or simultaneous) 4. Addition of hydrogen to the broken bonds of both sulfur and carbon 5. Release of the hydrocarbon product from the catalytic site 6. Release of the H2S from the site The sequence of these phenomena may not necessarily occur in the order shown. As discussed earlier, it is believed that there are two different types of sites which have different functions, direct sulfur extraction and hydrogenation. Certainly, a site capable of extracting sulfur from the parent aromatic molecule can also extract sulfur from its hydrogenated derivative, although the rate may be higher. In the following discussion, we refer to these phenomena as steps in various mechanistic pathways. Prior to about 1981, when the unique species Co(Ni)-Mo-S or SBMS was first identified, mechanisms proposed for HDS envisioned that reactions occurred through interactions between the organosulfur molecule and Mo orbital “vacancies” on the external surface of MoS2 crystallites or between the organosulfur molecule and peripheral Mo-SH groups, as illustrated in Fig. 23. The olefinic or acetylenic byproducts were assumed to be hydrogenated much faster than the desulfurization reactions occurred; thus they may or may not be observed in reaction products. Vacancies were later called coordinately unsaturated sites (cus). This is more in line with terminology used in organometallic chemistry. In view of the present understanding of the nature of the active sites, SBMS or Co(Ni)-Mo-S, the following discussion describes mechanisms in terms of catalysis by organometallic complexes. The references available on this topic are too numerous to mention, and the mechanisms are very well understood. A particularly useful reference is the book by Candlin, Taylor, and Thompson (90),although there are many others that can be consulted. 3. Known Organometallic Analogs of HDS Catalysts

Several excellent reviews and articles have been written recently that relate known organometallic chemistry mechanisms to the transformations

418 D. DUAYNE WHITEHURST, TAKAAKI ISODA, AND ISAO MOCHIDA that occur in commercial HDS processes (4, 7, 89, 91, 92). There is little need to present all of the information found in those reviews in this paper as they are well documented, and the reader is referred to those reviews for more detailed information. However, a few selected points will be discussed that are pertinent to the HDS of dialkyldibenzothiophenes. There are various ways in which thiophene compounds can bond to metals, as shown in Fig. 24. The 7'-S type bonding refers to attachment of the thiophene compound to the metal through the lone electron pair on the sulfur atom. a-Bonding of carbon to the metal is called ql-C type bonding. All other bonding modes shown in the figure refer to bonding between the metal and some a orbitals of unsaturated C=C bonds. The significance of the different bonding modes is described later. For the present, it should be mentioned that the 7I-S and q3,q1-Sbonding modes are the ones believed to be important in C-S bond cleavage reactions. The a-bonding modes activate the unsaturated bond toward nucleophilic attack (by such species as hydride). For condensed thiophene ring systems, such as benzothiophene, the preferred mode of bonding is q6,through the benzene ring. This activates the benzene ring for nucleophilic attack but does not directly lead to C-S bond cleavage (4, 89, 91, 92). The terminology used in describing the mechanistic steps involved in the transformation of species bonded to the central metal atom is illustrated with relevance to HDS reactions in Figs. 25 and 26 using one possible configuration for SBMS in the following abbreviated description for the catalytic metal site:

Co-Mo-S or

SBMS

Abbreviation M = Co or Ni L = ligands( 1-4 )

The mechanistic steps shown in Figs. 25 and 26 have been demonstrated with soluble organometallic complexes (4, 89-102). Some of the reactions have been shown to be truly parts of a catalytic cycle, and others were observed to produce stable products which were isolated and characterized ( 4 ) . It should be kept in mind that organometallic chemistry studies generally involve reactions at low temperatures (<150°C), and reactions that do not proceed at these low temperatures could well occur at the higher

Y /

420 D. DUAYNE WHITEHURST, TAKAAKI ISODA, AND ISAO MOCHIDA

temperatures employed in commercial HDS processes. Also, the illustrations of “hydride” attack on soluble organometallic complexes involve a soluble hydride source, such as LiA1H4. Analogous reactions on heterogeneous catalysts would have to involve hydride sources at adjoining metal sites. Similarly, “proton” attack for soluble materials involves an added external soluble acid. Heterogeneous analogs would have to involve adjacent -SH groups. Many of these studies utilized noble metals such as Ir, Os, Rh, Ru, or Re, whereas others used more conventional metals such as Mn, Fe, Mo, or Co. The particular metal on which the observations were made is not important at this point. What is important is that all of the important steps required for direct sulfur removal and hydrogenation of thiophene and more condensed derivatives have been shown to occur with soluble metal complexes. Thus, organometallic complex chemistry can be of great value in elucidating the mechanisms involved in conventional HDS processes and perhaps can point the way to improved catalyst formulations. The significance of reversible ligand coordination and replacement is discussed in detail in later sections dealing with inhibition and rate-limiting steps. Soluble metal complexes have been known for almost 30 years to be active catalysts for simple olefin hydrogenation reactions and need little discussion. Hydrogen activation is facile and reactions proceed smoothly and with high selectivities. The addition of dihydrogen to the unsaturated bond is generally stereospecific endo, and the two added hydrogens are cis to each other in the product. Perhaps the most important discovery with relevance to dialkyldibenzothiophene conversion is the discovery that the C-S bond cleavage, a ringopening reaction, can proceed directly on organometallic complexes without the requirement of prior hydrogenation of the thiophene ring (4, 89, 91-95). The product of this ring opening is a rather stable pseudoaromatic class of compounds called metallathiabenzenes (as illustrated in Fig. 25). These monomeric organometallic intermediates have been isolated and their chemistries have been elucidated, as illustrated in Fig. 26 (reaction paths 111-V) (4, 89, 91-95). This mode of thiophene conversion is the only known pathway that is relevant to dibenzothiophenes and their alkyl derivatives. Many of the intermediate metallathiabenzenes react with hydrogen gas at low temperatures to yield the fully saturated desulfurized hydrocarbon is high yield (4, 89, 98). Another important feature exhibited by these compounds is that many of them will add dihydrogen to the metal center and produce a stable hydride-containing product. This stable product can be destabilized by the addition of an external neutral ligand that forces the hydride to transfer to the carbon bonded to the metal, thus breaking the

POLYAROMATIC SULFUR COMPOUNDS

42 1

metal-carbon bond ( 4 ) .Such processes are referred to as facilitated reductive elimination. Generally, facilitated reductive elimination reactions are induced with phosphines or carbon monoxide. However, it is not unreasonable to suspect that H2S could serve the same function. If this is true, then in addition to the conventional role of H2S as a means to prevent the reduction of Mo to metal, there may be a fundamental chemical basis for observations that H2S is actually beneficial at low levels in HDS processes ( I 03). Pathway I11 of Fig. 26 has been demonstrated for thiophene and benzothiophene with Ir complexes ( 4 ) and for all thiophenes, including dibenzothiophene, with Rh complexes (94, 95). These oxidative additions appear to be influenced by substituents present on the carbon atoms adjacent to the sulfur atom. Insertion between sulfur and the unsubstituted carbon is highly preferred. For 2-methylthiophene the exclusive product is the 1-5 bond insertion product, whereas for 3-methylthiophene, no preference for insertion was observed (1-2 and 1-5 bond insertion products were equal). In competitive studies, thiophene was found to be about twice as reactive as 2,5-dimethylthiophene. This behavior is similar to that observed for relative reaction rates of substituted thiophenes observed with conventional HDS catalysts. Thus steric limitations can occur, even with monomeric, homogeneous catalysts. Several of these metal complexes that readily form metallathiabenzenes have been shown to be true homogeneous catalysts for the conversion of benzothiophene to ethylbenzene (4). Pathway IV has been demonstrated only for thiophenes and substituted thiophenes. Pathway V has not been demonstrated but could be involved in reactions conducted in the presence of acids. This point is expanded later because of the potential involvement of acidic -SH groups on HDS catalysts. The other pathways for hydrogenation and C-S bond cleavage need some discussion. Pathways I and IT are well documented for thiophenes and alkyl-substituted thiophenes. However, there are no examples reported for benzothiophene or dibenzothiophene. Coordination of the thiophene ring via S and C=C a-bonding activates the ring for nucleophilic addition. It has been clearly demonstrated that conventional soluble hydride sources such as LiAlH4 can add hydride to these complexes. The hydride addition can proceed in one of two ways. Either the addition can partially saturate the thiophene ring in the 2-position, as shown in pathway I, or it may induce C-S bond cleavage, as shown in pathway 11. Substituents present in the 2and 5-positions greatly affect the rates of reaction (TH > 2-MTH = 3-MTH B 2,5-DMTH) (89). This clearly demonstrates that steric restrictions on transition states can occur even for monomeric, organometallic complexes.

422 D. DUAYNE WHITEHURST, TAKAAKI ISODA, AND ISAO MOCHIDA As shown in Fig. 26, pathway I, the partially saturated product can be further converted to 2,3-dihydrothiophene by proton attack on the intermediate (89, 97). The strength of the acid is quite important as weak acids do not induce this transformation. Acids having strengths similar to or higher than that of chloroacetic acid are required. Of relevance to this is the observation that >SH bridges between Mo atoms in metal clusters have acid strengths of this magnitude (89). Such bridging groups are quite probable on the edges of MoS2 crystallites, and so the potential for cooperative interactions between adjacent metal and -SH centers may be quite reasonable. At the elevated temperatures of commercial HDS processes, even nonbridged -SH groups could serve this function, although they are weaker acids. If the partially hydrogenated 2,3-dihydrothiophene intermediate undergoes isomerization to the 2,5-dihydro isomer, the product decomposes rapidly at 120°C to give butadiene in high yield. This resembles the “simultaneous” pathway discussed for HDS mechanisms catalyzed by MoS2 (see Fig. 23) and has been observed in Re/A1203-catalyzed conversions of 2,3dihydrothiophene (97). As in the case of metallathiabenzenes, intermediates such as those shown in pathway 11 (Fig. 26) can also react with dihydrogen to produce the desulfurized fully saturated hydrocarbon product; for example, the ringopened product from 2,5-dimethylthiophene is cleanly converted to n-hexane on reaction with dihydrogen (27 atm, 110OC) (89). The modes of coordination of the aromatic sulfur compounds discussed all involve bonding to either the sulfur lone pair or the 7~ system of the thiophene ring system by 2, 4, or 5 IT electrons. It has been shown that, for benzothiophene, the preferred mode of bonding is through the IT system of the benzene ring (89). This generally does not activate the ring, but in one report, IT-bonding through the benzene ring of benzothiophene labilizes that ring toward nucleophilic hydride attack and a partially hydrogenated benzene ring results in a mechanism resembling that of pathway I (98). This observation is quite important, as a major route to desulfurization of dialkyldibenzothiophenes is via hydrogenation of the benzene ring prior to desulfurization. Further work in this area is encouraged. All of the preceding discussion has concentrated on the chemistry of monomeric metal complexes. However, in recent years, there have been several very interesting developments using clusters of metals linked through sulfide bridges (99-102). Jones and co-workers have found clusters of Rh and of Co induce C-S bond cleavage (loo), and Adams and coworkers have observed that clusters of 0 s can cleave both C-S and C-C bonds (101). Clusters containing two molybdenum atoms linked by four sulfur bridges have been observed to exhibit interesting properties. The

POLYAROMATIC SULFUR COMPOUNDS

423

sulfide bridges appear to be able to dissociate dihydrogen reversibly to produce bridged >SH groups. These >SH groups can add to activated olefins and alkynes, and if one group is a bridging CH3S< group, true homogeneous catalysts result (102). Unfortunately, these catalysts are only active enough to hydrogenate activated olefins and alkynes. However, the results do give some credence to earlier speculation that the -SH groups on MoS2 crystals were potential hydrogen activation sites. Perhaps the most exciting cluster materials made to date are those illustrated in Fig. 27. These are bimetallic clusters containing two Mo and two Co atoms connected by three sulfur bridges (99). The two metals are believed to serve different functions. Cobalt is very effective in extracting sulfur from organic compounds but does not activate dihydrogen well. Molybdenum is a good activator of dihydrogen, but it has a low activity for sulfur extraction. Thus, working in a cooperative fashion, the cluster compound is more effective than either metal alone. Like SBMS, these clusters are much more active for desulfurization than clusters containing only Mo (99). These clusters have been found to stoichiometrically extract sulfur from a wide variety of thiols and thiophene derivatives. The ease of sulfur extraction, as shown in Fig. 26, is thiophenol > thiophene > 2-MTH = 3-MTH > 2,5-DMTH. This is again in reverse order to that expected for metal-sulfur coordination bond strength but in the expected order of decreasing reactivity with increasing steric limitation. To date, the reactions of these clusters with sulfur compounds have been stoichiometric. The starting cluster could not be regenerated by dihydrogen, but the intermediates bound to Co were hydrogenated to the corresponding hydrocarbons. The reaction From Thiophene

RsX' K

C, to C, olefins and paraffins

s=c=o

c=o

R _R' _ _% -C o w H H 100 CH3 H 30 CH3 CH3 2o FIG.27. Desulfurization with organometallic clusters. Reprinted with permission from Ref. 99. Copyright 1994 American Chemical Society.

424 D. DUAYNE WHITEHURST, TAKAAKI ISODA, AND ISAO MOCHIDA

mechanism proposed was almost identical to that proposed by Startsev for HDS with heterogeneous SBMS catalysts (2). It was believed to proceed via coordination of the thiophene through the sulfur atom, oxidative addition via pathway I11 producing a metallathiabenzene, a series of dissociative additions of hydrogen to the Mo, transfer of hydrogen to Co, reductive eliminations of the hydrogenated intermediate (as in pathway 111), and finally release of the hydrocarbon. Unfortunately, with these clusters, the extracted sulfur remained as a stable compound (as shown in Fig. 27). Regeneration was thus shown to be the difficult step. However, some success was found by reaction of the sulfur-containing cluster with CO to produce COS. The authors concluded that the rate-limiting step in commercial desulfurization processes may be removal of sulfur from the surface of the catalyst. However, with conventional heterogeneous MoS2-based catalysts, if the extracted sulfur is in the form of M=S, it should be easily converted to H2S, as this is the intermediate in low-temperature sulfiding, and once the sulfur is present as adsorbed H2S, it is easily desorbed (2). In the previous section, which discussed the origin of promotion in commercial HDS catalysts, it was theorized that the strength of the metal-sulfur bond had to be at an optimal level. Metal-sulfur bonds that were too strong prevented completion of the catalytic cycle and those that were too weak were unfavorable for the sulfur compound coordination ( I , 2, 67, 85-88). The cluster complexes synthesized to date must not have achieved this optimal level. More research with a wider variety of clusters is encouraged. It is known that SBMS or Co(Ni)-Mo-S formulations are optimal at 2 M o / l Co, and so work in the future to synthesize and test clusters having a stoichiometry closer to that of the known optimum for commercial HDS catalysts is warranted. Also, by changing the nature of the ligands on the metals, their electronic properties may be altered to be more favorable for sulfur removal. Although all of the preceding discussion appears quite positive, there are several other points that need to be addressed. The catalytic hydrogenation of single-ring aromatic hydrocarbons has not been demonstrated with soluble metal complexes and, as was discussed earlier, single-ring hydrogenation is a significant side reaction in the HDS of dialkyldibenzothiophenes (e.g., conversion of biphenyls to cyclohexylbenzenes). It would also be useful to have data on the relative rates of benzothiophene or dibenzothiophene desulfurization in competition with hydrogenation of diaromatic hydrocarbons such as naphthalenes. A potential major accomplishment for new catalyst synthesis would be the discovery of materials that will desulfurize dialkyldibenzothiophenes and not hydrogenate diaromatic hydrocarbons. This would be of great benefit for limiting hydrogen consumption in HDS processes.

POLYAROMATIC SULFUR COMPOUNDS

425

A significant observation in the preceding discussions was that, in many instances, the reactions could be made to progress only by the addition of an external nucleophile (such as hydride) or by the addition of an external proton to some stable intermediate. These observations are quite important as they may be providing an indication of which steps of the catalytic sequence in commercial HDS catalysts are rate limiting. Many early authors suggesting HDS mechanisms have proposed the involvement of adjacent -SH groups in cooperation with the reactions occurring on the SBMS site (see Fig. 23). If such groups are required, it may be necessary to synthesize catalytic sites in which metal and protonic sites are in close proximity. In addition, from the work of Angelici, the acid strength of the protons seems to be important (89).Finding ways to increase the acid strength of bridging \ A H groups or replacing them with stronger acid functions may be of benefit in designing improved HDS catalysts. V. Computational Aids to Mechanistic Understanding A. GENERAL COMMENTS

In the preceding section, we discussed reports by various authors that attempted to develop correlations between the basicity of the sulfur atom in thiophene compounds and the reactivities of those compounds in desulfurization with homogeneous organometallic complexes. Several examples were noted whereby the trend in the rates of reaction was just the opposite of that which would be expected if one assumed that a higher calculated S-metal bond strength would lead to higher reactivity (4, 89, 99). For example, even with very reactive bimetallic cluster complexes of Co and Mo, the reactivity decreased with increased substitution on the carbon atoms adjacent to the thiophene sulfur atom. It was proposed that steric hindrance was the cause of lower reactivity even with these soluble catalysts (99).This is reminiscent of the results presented in Tables X and XI, showing that the HDS rates of benzothiophenes and dibenzothiophenes decreased as alkyl groups were substituted on carbon atoms adjacent to the sulfur atom (21,25,26,30,32). In this section, we present information from quantum chemical calculations that contributes to the understanding of the relative roles of electronic and steric effects in the reactivity of dibenzothiophenes having different degrees of substitution. OF METAL-SULFUR COORDINATION BOND B. CALCULATION STRENGTH (ADSORPTION)

The reactivity of an organic sulfur compound in an HDS process depends on several factors. The compound must first be adsorbed onto the catalyst

426 D. DUAYNE WHITEHURST, TAKAAKI ISODA, AND ISAO MOCHIDA surface. This step in the process will be affected by the strength of the bond formed upon adsorption and by competition of other species that can bond (coordinate) to the same active site. It has been observed by many researchers that HDS processes exhibit Langmuir-Hinshelwood kinetics, which is consistent with this competitive mechanism for adsorption ( I , 2, 5). The strength of the coordinate bond depends on a balance between the ability of the molecule to donate electrons to the metal site, either from the lone pair on sulfur (basicity) or from some 71 system on the molecule, and steric factors that limit this interaction. The previous section cited several instances in which steric factors were the controlling characteristic in reactions of organometallic complexes (89, 102). It is now possible to calculate theoretically the heats of adsorption of sulfur compounds and their entropies of interaction from first principles, and several authors have reported such calculations for a variety of sulfur compounds, including thiophenes and substituted thiophenes (2, 89, 102). Table XI11 presents some calculations for the interaction of SBMS type catalytic sites with relevant sulfur compounds (2). The significance of these calculations is that they predict that, as might be expected, tetrahydrothiophene is much more strongly adsorbed than thiophene. However, surprisingly, thiophene was predicted to be more strongly adsorbed than H2S. This means that H2S should be easily displaced from the catalytic site by thiophene and that the tetrahydrothiophene would not be released until it is converted to the corresponding hydrocarbon. We discuss in more detail in the following section the effects of inhibitors in the HDS of dialkyldibenzothiophenes. However, in the present context, H2S would not be considered a significant inhibitor for adsorption of simple thiophenes. Steric factors could lower the overall strength of the sulfur-metal coordinate bond, and highly substituted alkylthiophenes may be more susceptible to competitive adsorption than unsubstituted thiophenes. This does not mean that H2S may not be an inhibitor in the overall rate of conversion as the rateTABLE XI11 Enthalpies of S-Containing Molecule Adsorption on SBMS of Different Compositions in Terms of the Method of Interacting Bonds (2) AH (kcalimol) of molecule adsorbed SBMS composition

H2S

Thiophene

THT

Co/MoS2 Ni/MoS, co/ws2 NilWS,

11.5 12.2 11.9 12.6

21.2 22.1 21.1 22.6

35.1 36.1 35.8 36.8

427

POLYAROMATIC SULFUR COMPOUNDS

determining step may involve other equilibria or could be related to slower processes that occur between adsorbed species. The theoretical calculations described have recently been supported by an extraordinary kinetic analysis conducted by Vanrysellberghe and Froment of the HDS of dibenzothiophene (104). That work provides the enthalpies and entropies of adsorption and the equilibrium adsorption constants of H 2 , H2S, dibenzothiophene, biphenyl, and cyclohexylbenzene under typical HDS conditions for CoMo/A1203catalysts. This work supports the assumption that there are two different types of catalytic sites, one for direct desulfurization (termed a) and one for hydrogenation (termed T). Table XIV summarizes the values obtained experimentally for adsorption constants of the various reactants and products, using the LangmuirHinshelwood approach. As described in more detail in Section VI, this kinetic model assumes that the reactants compete for adsorption on the active site. This competitive adsorption influences the overall reaction rate in a negative way (inhibition). It is clear from the data presented in Table XIV that H2S, though adsorbed competitively with dibenzothiophene, is not a major inhibitor for dibenzothiophene adsorption. Dibenzothiophene was shown to be preferentially adsorbed relative to biphenyl on both the cr and T sites. It is surprising that no adsorption of H2S was noted on the hydrogenation site (T) since it is known to be a strong inhibitor for many aromatic hydrogenations. A TABLE XIV Adsorption Equilibrium Constants and Rate Coeficients in the HDS of Dibenzothiophene Catalyzed by CoMo/AIz03 (104)a Adsorption equilibrium constants (m3/kmol) Compound Hz HzS Dibenzothiophene Biphenyl Cyclohexylbenzene

Direct desulfurization site ( a ) 0.707 62.79 75.69 9.54

Hydrogenation site (7) 0.0142 2.52 1.41 0.334

~~~~~~~

Reaction

4w 4.w &4

-+

4-4

+

4SCY6

4'-cy6

Rate coefficient (kmol . (kg of catalyst)-' . h-') 0.158 0.308 1.69

Relative rate w a d ' k) 100 19.9 6.5

Consult the reference for the form of the rate equation. Reprinted with permission from Ref. 104. Copyright 1996 American Chemical Society.

428 D. DUAYNE WHITEHURST, TAKAAKI ISODA, AND ISAO MOCHIDA major finding was that H2 was very weakly adsorbed by either site and it was concluded that competitive equilibrium adsorption between hydrogen and benzothiophene, or other components in the reaction mixture, controlled the overall rate of product formation for both direct desulfurization and hydrogenation (104). The significance of these findings to the effects of inhibitors in HDS and the observed Langmuir-Hinshelwood kinetics is discussed in more detail in the next section. As Vanrysselberghe and Froment’s data suggest that rate-controlling HDS processes are associated with either hydrogen dissociation or adsorption, it would be useful to have available theoretical calculations on the energetics of adsorption of H2 and the bond dissociation energies of H2 on different Co-Mo-S or SBMS compositions and/or configurations. Such information has been derived for a wide variety of organometallic complexes by Angelici (105). Those studies showed that the energetics of hydrogen dissociation were quite different for homolytic and heterolytic dissociation of hydrogen. For heterolytic dissociation, metals having higher basicity underwent oxidative addition (via proton addition) more readily. The nature of the metal and its attached ligands were very important to its basicity. Homolytic dissociation appeared less sensitive to the nature of the ligands. Extension of similar computations to supported clusters of SBMS type catalysts could be very helpful in determining the importance of the catalyst’s geometry and the importance of the proximity of -SH groups to the active site. Some calculations relating to this have been carried out for the compound CoMoS4 (106). Those studies concluded that dihydrogen could be dissociated homolytically by either the Mo or the associated sulfurs with about equal probability. More studies in this area are needed. An alternate possibility that does not seem to have been addressed for soluble organometallic complexes is the heterolytic dissociation of dihydrogen by a bridging M-S-M bond. This type of dissociation would produce a M-H (hydride) and an -SH (proton) in close proximity. The previous discussion of organometallic complex analogs for HDS catalysts showed that one of the major mechanistic routes for desulfurization was by sequential attack on the coordinated sulfur compound by a hydride and then a proton. This heterolytic dihydrogen dissociation mechanism has been proposed by Kasztelan and Guillaume for the hydrogenation of toluene catalyzed by MoS2/A1203under typical HDS conditions in connection with studies of the inhibition of hydrogenation by H2S (107). These authors showed that hydrogenation was not inhibited below 50 Pa of H2S partial pressure, moderate inhibition was observed between 50 and 60,000 Pa (0 to - 1/2 order), and no further inhibition occurred at higher H2S partial pressures. The rationale proposed for this behavior was that the ratecontrolling step involved hydride addition from the Mo at low H2S partial

429

POLYAROMATIC SULFUR COMPOUNDS

pressures and proton addition from an adjoining -SH group at high H2S partial pressures. As discussed previously, there are many mechanistic steps required for the removal of sulfur from a thiophene molecule, and it is quite possible that the rate-controlling one may change, depending on the reaction conditions, the presence of molecules competing for adsorption, and the nature of the molecule being converted. Computational modeling of such processes on SBMS type structures could be quite instructive. OF ELECTRON DENSITY ON SULFUR AND C. CALCULATION C=C BONDORDERS

OF

Correlation of reactivity with the electronic properties of dialkyldibenzothiophenes is somewhat more complicated because these molecules may undergo sulfur extraction via two different pathways. The sulfur extraction can proceed either directly without aromatic ring hydrogenation (kDJ or after hydrogenation of one or both of the aromatic benzene rings (kD, or k,,), as discussed in the previous section (5,15,17, 21,30, 31). In the direct extraction route, the electron density on the sulfur atom is believed to be related to the reactivity of the molecule toward sulfur removal, provided that steric limitations are not present. This is also the case with homogeneous organometallic complex oxidative additions (4,89,92).In the hydrogenative route, reactivity is related to the highest bond order in the molecule. This was shown by comparison of the measured relative rates of desulfurization of different sulfur compounds and quantum mechanical calculations using the MOPAC-PM3 procedure (38).These initial calculations provide only a first approximation as they assume that all orbitals have equal importance. A more thorough treatment using the frontier orbital theory could provide more information; however, even with these less sophisticated calculations, some very interesting correlations were noted (38). Figure 28 presents the calculated electron densities and bond orders of a variety of thiophenes, benzothiophenes, and dibenzothiophenes. It can be seen in the figure that the electron density on sulfur actually increases as the thiophene ring system becomes more condensed, e.g., with the parent molecules thiophene < benzothiophene < dibenzothiophene. This has also been noted by Rauchfuss (92). Thus, if electron density were the only controlling factor, one might expect dibenzothiophenes to be more reactive than thiophenes. However, as was discussed previously, the reverse is true. Thus other factors must also be important. For a given ring system, the electron density on sulfur was found to correlate quite well with measured reactivities. Similar trends and sensitivities to electron densities were observed for all thiophene ring systems (38). This is illustrated in Fig. 29, where the rate constants are normalized to

430 D. DUAYNE WHITEHURST, TAKAAKI ISODA, AND ISAO MOCHIDA Group number (Electron density range): I (5.690-5.780)

I1 (5.900-6.000)

Ill (>6.000)

Single-cyclic sulfur compounds and dialkyl sulfide:

Two-cyclic sulfur compounds:

FIG.28. Electron density (the data next to the numbers) on the sulfur atom and the bond order (the data next to the bond) of representative heterocyclic sulfur compounds and same of their hydrogenated derivatives. Reprinted with permission from Ref. 38. Ma et al. (1995). Copyright 1995 American Chemical Society.

bring the parent ring systems to similar values. It can be seen in this figure that there is a striking deviation from the correlation for the alkylsubstituted dibenzothiophenes. Therefore, it may be safely assumed that steric factors outweigh electronic factors for this series. At present, the precise reason for low reactivity of alkylbenzothiophenes and alkyldibenzothiophenes is not definitely known. It is clear that steric factors are indeed important, and it has been proposed that steric hindrance lowers the adsorption constant for these species (5, 17, 25, 26). Molecular

POLYAROMATIC SULFUR COMPOUNDS

431

Group number (Electron density range):

I (5.690-5.780)

I1 (5.900-6.000)

111(>6.000)

Three-cyclicsulfur compounds:

FIG.28. (continued)

orbital calculations do indicate that bulky methyl groups would be in close proximity to a planar surface containing the catalytically active site (17, 30, 31). It is known that thiophene bonds to metal surfaces at an angle of 130" rather than perpendicular to the planar metal surface (92).However, if C-S bond cleavage proceeds through the formation of the metallathiabenzene intermediate, described in the previous section, then the transition state could well be perpendicular. An alternate mechanism for sulfur removal involves the initial hydrogenation of one of the benzene rings prior to sulfur removal. This reduces steric restrictions by forcing the methyl group out

432

D. DUAYNE WHITEHURST, TAKAAKI ISODA, AND ISAO MOCHIDA

Group number (Electron density range): l(5.690-5.780) I1 (5.900-6.000)

111 (%.OOO)

Four-cyclic sulfur compounds:

a

1.

0

1.444

0

,

Iy]

s

1.1

x.9

1.416

(28) XlS

-The bond with bond order more than 1.60 o The bond with bond order between 1.50 and 1.60 FIG.28. (continued)

of the plane of the sulfur ring system (puckering). It also increases the electron density on the sulfur, which strengthens the coordination of the molecule to the catalytic site. The carbon-sulfur bond becomes more labile and desulfurization proceeds more rapidly. This route is particularly favored with NiMo/A1203catalysts, which are known to be much more active for hydrogenation (21). In view of the mechanisms observed with organometallic complexes, steric restrictions imposed by bulky substituents during oxidative addition or hydride attack should also be considered as rate determining. One observation that tends to support steric restrictions to oxidative addition is with NiMo/AI2O3 catalysts, in which a significant product observed in the HDS of 4,6-dimethyldibenzothiophene is 3,3’-dimethylbicyclohexyl (DMCC) (21), as shown in Fig. 15. The relative rate of hydrogenation of the dimethylcyclohexylbenzene (DMCB) is low, and so the DMCC must have arisen from continued hydrogenation of the partially hydrogenated intermediate dimethyltetra(or hexa)hydrodibenzothiophene. The fully sat-

433

POLYAROMATIC SULFUR COMPOUNDS

5.6

5.7

5.8

5.9

6.0

6.1

6.2

Electron density of the sulfur atom FIG.29. Correlation between relative hydrogenolysis reactivity and the electron densities on sulfur. (0)Reaction conditions: 300"C, 71 atm, sulfided Co0-Mo03/A1203 (5). (1) DBT; (2) benzo[b]naphtho[2,3-d]thiophene; (3) 7,8,9,10-tetrahydrobenzo[b]naphtho[2,3-d]thiophene; (4) 5b,6,11,1la-tetrahydrobenzo[b]naphtho[2,3-d]thiophene.(0)Reaction conditions: 450°C, 1 atm, sulfided Co0-Mo03/A1203(6). ( 5 ) Thiophene; (6) tetrahydrothiophene; (7) benzothiophene; (8) 2,3-dihydrobenzothiophene.(+) Reaction conditions: 360°C, 2.9 MPa, sulfided Ni0-MoO3/AI2O3(3). (9) 1-Methyl-DBT; (10) 2- or 3-methyl-DBT; (11) 4-methylDBT; (12) 4,6-dimethyl-DBT. (A) Reaction conditions: 300"C, 102 atm, sulfided COO-Moo3/ A1203(8).(13) 2J-Dirnethyl-DBT (14) 3,7-dimethyl-DBT; (15) 4-methyl-DBT; (16) 4,6-dimethyl-DBT. Reprinted with permission from Ref. 38, Ma e l al. (1995). Copyright 1995 American Chemical Society.

urated product (cyclohexylcyclohexane) was not observed in the HDS of dibenzothiophene with the same NiMo/A1203 catalyst, and so direct desulfurization of the intermediate hexahydrodibenzothiophene must have been very fast relative to the further hydrogenation of that intermediate (21).Further research in this area to distinguish between steric limitations to oxidative addition vs steric limitations to sorption is needed to clarify this issue. As mentioned earlier, the MOPAC-PM3 calculations also helped to determine the importance of bond order in the hydrogenative route to desulfurization. Figure 28 shows the calculated bond orders of all bonds in a wide variety of thiophenes, benzothiophenes, and dibenzothiophenes (38). These values were correlated with the rates of desulfurization of sterically hindered alkyl-substituted benzothiophenes and alkyl-substituted 1,l'-diox-

434 D. DUAYNE WHITEHURST, TAKAAKI ISODA, AND ISAO MOCHIDA

obenzothiophenes reported by Geneste (29). Clear trends were noted, as shown in Fig. 30. These compounds were believed to undergo desulfurization primarily by hydrogenation of the thiophene ring prior to desulfurization, and so the overall rates of desulfurization would be expected to relate to the ease of the hydrogenation of the first unsaturated bond in the molecule. A higher bond order is expected to correspond to a position that is more labile to hydrogenative attack. The correlation shown in Fig. 30 supports this assumption and indicates that, for benzothiophenes, enhancement of the bond order by 0.05 units will increase the rate of hydrogenation by a factor of about 10. Thus, theoretical calculations can be a very useful tool in mechanistic understanding. As it is now known that the most active Co-Mo-S or SBMS catalysts consist of very small clusters of Mo atoms (as few as seven), it is within the realm of practical computational procedures to completely model the catalyst/S-molecule interactions. Assumptions made about steric crowding around the catalytic site may be quite different for such small clusters as the catalytic site is not an extended planar surface, as discussed in the previous section. Future work in this area should provide new insight into the true limitations in HDS of dialkyldibenzothiophenes.

1.7

1.8

1.9

2.0

Bond order FIG. 30. Relative hydrogenation reactivity of methyl-substituted benzothiophenes and methyl-substituted 1,l ‘-dioxodibenzothiophenes versus their Cz-C3 bond order. (0)Benzothiophenes: (5) BT; (7) 2-methyl-BT; (8) 3-methyl-BT; (9) 2,3-dimethyl-BT. (0)1,l-Dioxobenzothiophenes (BT02): (11) BTOz ; (29) 2-methyl-BT02; (30) 3-methyl-BTOZ; (31) 2,8-dimethyl-BT02. Reprinted with permission from Ref. 38, Ma et a!. (1995). Copyright 1995 American Chemical Society.

POLYAROMATIC SULFUR COMPOUNDS

VI.

435

Limitations in Conventional HDS Processes

A. GENERAL COMMENTS Previous discussions focused on the potential for improvement in the activity and selectivity of HDS catalysts. It was concluded that there may be limits to how much the activity can be improved for conventional HDS catalyst compositions. Thus, it may be necessary to look for alternatives that can help achieve the new, stricter standards of 0.05% sulfur in fuel. In seeking alternatives, it is useful to assess what might limit the options that are being considered. Such limitations include safety regulations for the present process equipment, thermodynamic constraints on mechanistic intermediates, limitations on process conditions and alternative mechanistic routes imposed by the major components of the fuel being processed, and the influence of non-thiophene components on the rates of desulfurization and/or hydrogenation (inhibitors). Knowledge of these limitations is important when seeking alternatives to present process conditions and configurations.

B. PROCESS EQUIPMENT LIMITATIONS When new demands warrant changes in a refineries product slate, the most desirable process change for the refiner is none at all. The next most desirable change to make is to replace the catalyst in an existing reactor and to continue with the same process conditions. This minimizes the cost to the refinery and causes the least amount of downtime. As discussed earlier, it may be difficult to improve HDS catalysts to the point that such a simple change will allow the new standards to be met in many of the existing facilities. Lowering the flow rate through the reactor could allow more time for conversion. However, this is never considered, as constant daily product production is the key to refinery profits. If higher hydrogen pressures could be used, the rates of desulfurization could be substantially increased. However, this is a limited option. As discussed in the beginning of this report, some refineries were able to purchase new high-pressure reactors during a time of low equipment and construction costs. However, new construction will not benefit from this luxury. Many of the presently installed reactors were designed for moderate pressures, less than 5 MPa. It would therefore be desirable to devise new processes around these pressures. Increasing temperature is another means to increase reaction rate. This is the lowest-cost process alternative to achieve higher rates as long as no

436 D. DUAYNE WHITEHURST, TAKAAKI ISODA, AND ISAO MOCHIDA

losses in product yield are incurred. However, there are practical limits to this route, which are imposed by pressure-temperature safety standards. For most reactors, temperature increases of about 50°C are within reason. As discussed in a later section, H2S is an inhibitor for the catalytic site responsible for direct sulfur extraction. Thus, if the H2S partial pressure could be lowered in the reactor, the desulfurization rate could be increased. The simplest means to achieve this goal is through increased hydrogen recycle rates or increasing the hydrogenifeed ratio. Such changes are expensive and can in some instances lower the overall thoughput of the feed. If none of these options are available, new equipment may be necessary. In the last section of this review, a number of novel processes schemes are discussed that have the potential for meeting the new standards while overcoming the limitations presented in this section. C . THERMODYNAMIC LIMITATIONS ON MECHANISTIC INTERMEDIATES

There are no real thermodynamic limits in the removal of sulfur from any organic sulfur compound by reaction with hydrogen (1, 2, 5). There are, however, limits on the overall rates of conversion that may be achieved by increasing the temperature of the reaction. A classic limitation in rates is the result of the inverse relationship between adsorption on a catalytic surface and temperature. This may be a problem with dialkyldibenzothiophenes, which have steric limitations for adsorption. A more subtle limitation is encountered in the hydrogenative route to desulfurization. As shown in Fig. 31, at moderate pressures (3 MPa) and temperatures exceeding about 340"C, the hydrogenation of dibenzothiophene starts to become thermodynamically limited. Daage and Chianelli appear to be the only authors in this field that have considered this constraint (47). They showed that the conversion of dibenzothiophene at 350°C and 3-MPa hydrogen pressure was limited to about 20% by thermodynamic equilibrium. This is not a serious impediment to the overall rate of conversion via the hydrogenative route in their case. However, with dialkyldibenzothiophenes, if higher temperatures are necessary to increase conversion rates, this equilibrium limitation may become more severe. In equilibriumcontrolled reactions, such as A [B] + C, if the equilibrium concentration of the intermediate [B] is less than 2% and the conversions of A -+ B and B + C occur on two different catalytic sites, then the overall rate of formation of C will be controlled by the diffusion rate of B from site 1 to site 2. In such a situation, staged reactions or physical mixtures of catalysts with different functions will not help. The only way to overcome this limitation is to design catalysts in which the two catalytic sites are extremely

-

437

POLYAROMATIC SULFUR COMPOUNDS

0

100

300

200

400

500

T, "C FIG.31. Thermodynamic equilibrium in the hydrogenation of DBT to H,-DBT (2.9-MPa hydrogen pressure).

close to each other. This principle is well known in petroleum refining in such processes as catalytic reforming and has been demonstrated for combinations of conventional Bronsted acids and noble metal hydrogenation catalysts for reactions in which the equilibrium concentration of the M (108). intermediates [B] was limited to as little as For most studies done at conventional temperatures (<340°C), equilibrium limitation is not a major problem. However, at higher temperatures, especially with alkyl-substituted dibenzothiophenes, the rate constants, estimated by curve fitting, may become affected unless this phenomenon is taken into account. This complicates determining the rate constants using the model described in Section IV, where only sequential irreversible firstorder reactions were considered. If equilibrium is established much more rapidly than subsequent reactions, the fitted rate constants for hydrogenation (kHS,,kHs2,kHP,,and k H F 2 ) will still be able to predict the product distributions vs time or conversion, but will not be strictly correct. For the most accurate description of the overall reaction sequence, thermodynamic equilibria must be included in the rate constant determinations. This will become a significant problem if high temperatures must be used to lower the sulfur content of fuels to less than 0.05%. More studies to establish the thermodynamic equilibria of alkyldibenzothiophenes are needed for guidance in designing improved processes that require higher temperatures than are presently employed.

438 D. DUAYNE WHITEHURST, TAKAAKI ISODA, AND ISAO MOCHIDA

Another consequence of increased reaction temperatures is the shift in the equilibrium concentrations of polyaromatic hydrocarbons and their hydrogenated derivatives. At high temperatures, the fully aromatic hydrocarbons are thermodynamically favored, especially at the pressures of today’s HDS reactors. As the concentration of polyaromatics increases, the fluorescence of the finished fuel increases. This may cause problems with other specifications, as discussed shortly. D. FEEDSTOCK COMPOSITION LIMITATIONS

1. Susceptibility of Parafins and Alkylbenzenes toward Cracking Perhaps the major obstacles to increasing the process severity in HDS processes to meet the new standards are the non-sulfur-containing components of gas oils. It should be remembered that diesel fuels containing only 0.2% sulfur are 99% non-sulfur-containing compounds. Table I shows that these non-sulfur compounds in diesel fuels and gas oils are composed of 60-80% paraffins and 20-40% aromatics with small amounts of olefins. Figure 1 shows the distribution of the various components by boiling point. The sharp spikes in this analysis are the linear paraffins in the fuel. These are the most valuable components of diesel fuels as they provide the highest cetane number of the various components of the fuel. Paraffins and alkylaromatics are easily cracked thermally and are very reactive in the presence of acidic catalysts. Polyarornatic hydrocarbons are prone to condensation reactions, especially with acidic solids. Thus, when the sulfur content is lowered from 0.2 to 0.05%, one must convert 75% of very refractory materials and not harm any of the other 99% of the feed being processed. This is a severe challenge for new processes. Any new approach must have extremely high selectivity for sulfur removal without yield loss. 2. DifJiculty in Hydrogenation of Single Aromatic Rings

The new specifications not only limit the concentration of sulfur to 0.05% but also specify that the fuel should have the combustion properties of a 10%or lower aromatics-containing fuel and have a minimum cetane number of 40. Fuels that have more than 10% aromatics are now able to meet these specifications through additives (22). However, as smoke emission standards become more restrictive, the aromatics content of diesel fuels may have to be lowered to a true value of 10% or less. A very thorough review of the consequences of this potential problem has recently been written by Stanislaus and Cooper, which covers in detail aromatic hydrocarbon hydroprocessing kinetics, thermodynamics, catalyst compositions, and mechanisms (209).There is little need to repeat the details of that report

POLYAROMATIC SULFUR COMPOUNDS

439

in the present work, and the reader is referred to that review for a more thorough treatment. The main point to be made here is that thermodynamics again limit the production of low-aromatics fuels, as the standard is based on volume percent aromatics. HDS processes presently do a very good job of saturating multiaromatic rings to give single aromatic rings. However, single aromatic ring saturation is kinetically very slow, much slower than PASC desulfurization. In an attempt to remove PASCs from diesel fuels by raising the reaction temperature, the partially hydrogenated multiringed aromatics will dehydrogenate back to polyaromatics. Hydrocracking is proposed as one of the ways to remove multiringed materials, with the consequent losses in yield. 3. Potential for Polymerization and Color Formation Even if cracking can be minimized, there are other problems that arise with high-severity processes, especially with acidic catalysts. The olefinic feed components and byproducts of cracking can polymerize to form materials that produce gums via oxidation during storage. Another problem that is particularly important in Japan is color development in severely processed diesel fuels (110-119). Products having a yellow tinge or exhibiting fluorescence are rejected by distributors. The color specification is a semiquantitative method of comparisons with a standard set of colored fuels, called the Seibolt Color. Fluorescence was found to correlate with Seibolt Color in HDS products, and it has been suggested that the fluorescence intensity may be a more quantitative method for this measurement (119). In addition, it requires much less training for reliable measurements. Because of the criticality of this specification in Japan, a substantial amount of work has been done in recent years to determine the cause of this color development and to devise process schemes that overcome the problem (110-119). The various process schemes are discussed Section VII. The origin of color formation has been investigated by several research groups (118, 119). Isolation of the color bodies was achieved by highpressure liquid chromatography, and it was found that major contributors to the fluorescence color were high-boiling (>340°C) polyaromatic ring systems. The materials that exhibited the most intense fluorescence were anthracene, fluoranthene, and their alkyl derivatives. Fluoranthenes were observed to have fluorescence intensities about three times higher than anthracenes. Benzophenanthrenes, naphthacenes, 2,3-benzofluorenes, and pyrenes were the next most fluorescent, but they had much lower intensities. The other fuel components had no fluorescence (119). The most fluorescent materials constitute only about 1.2% of the diesel fuel, and it is not known whether they are produced during high-tempera-

440 D. DUAYNE WHITEHURST, TAKAAKI ISODA, AND ISAO MOCHIDA

ture HDS processing or whether they are inherent materials in the original feedstocks. The intensity of fluorescence responds to processing temperature, and low-temperature secondary treatment of high-severity HDS processes has been found to be a very effective method for removing fluorescence color after high-temperature HDS processing to achieve <0.05% sulfur (116, 119). The thermodynamic equilibria of the most intensely fluorescent product, fluoranthene (FLU), are not well established for the conditions relevant to HDS processes, but it is known that 1,2,3,1Ob-tetrahydrofluoranthene (4HFL) is the first major kinetic hydrogenation product during hydrotreating (120). If one assumes that the thermodynamic equilibrium of FLU H 4HFL is similar to that of naphthalene ++tetralin, then one can obtain a reasonably good correlation between the observed fluorescence and the % FLU in the FLU/4HFL pair as a function of the hydroprocessing temperature. This includes first-stage hydrotreating, severe HDS, and all subsequent lower hydrofinishing steps. It thus appears that a major contributor to fluorescence color may be existing condensed polyaromatic ring systems in the initial feedstocks and that they may not necessarily be produced in the process. If high temperatures are necessary for achieving the 0.05% sulfur specification, it may also be necessary to subsequently hydrotreat the low-sulfur product at a lower temperature to remove fluorescence color by shifting the equilibrium concentrations of polyaromatic rings to their hydrogenated derivatives, as discussed in the next section.

E. REDUCTION IN HDS REACTION RATESBY FEEDCOMPONENTS AND BYPRODUCTS 1. General Considerations It is generally accepted that both hydrogenation and sulfur extraction rates in HDS processes are lowered by the presence of certain other components in the feed or products, and numerous studies have addressed this problem (1-3, 5-9, 33, 41u, 104, 107, 109, 121-137). This reduction in rate is generally referred to as inhibition. This phenomenon is of extreme concern in Japan, where the new standards will be imposed in the very near term, and considerable effort has been devoted to reducing the problems associated with other feed components that further limit the rates of desulfurization of the most refractory components of diesel fuels and gas oils. Reports have issued mainly from three groups: Kabe and co-workers (13, 18,19,128), Nagai and co-workers (128-130), and Mochida and co-workers (14, 15, 17, 21, 30, 31, 38, 41, 41a, 114-117, 119, 133-137). Those reports are emphasized in this section. To provide some guidance to the reader as to the difficulty in assessing this limitation, the following text discusses

POLYAROMATIC SULFUR COMPOUNDS

441

some of the aspects involved in obtaining quantitative results with respect to the inhibition of desulfurization of alkyldibenzothiophenes. 2. Mathematical Modeling of Inhibition

The mathematical description of inhibition generally follows the kinetic form originally described by Langmuir (121).With time, more sophisticated revised treatments were developed, the most popular of which are the Langmuir-Hinshelwood or Langmuir-Hinshelwood-Hougen-Watson kinetic expressions (1-3, 5-9, 23, 104, 107, 109). The application of these kinetic forms in HDS reactions has been the subject of numerous in-depth studies and reviews as cited in the aforementioned references, and the reader is referred to those references for detailed information. In particular, hydrogenation of aromatic compounds in diesel fuels is discussed in great detail by Stanislaus and Cooper (109).A complete compilation of inhibition studies in HDS processes up to 1990 is available in a review by Girgis and Gates (5). More recent work by Froment (104) and by Vrinat (32, 33) is also quite instructive. This section discusses the inhibition phenomenon with specific reference to its influence on the conversion rates of dialkyldibenzothiophenes. The kinetic description of inhibition effects of even the parent molecule, thiophene, is quite complicated, and the complications become even greater as the thiophene core is fused to other aromatic rings and/or substituted with alkyl groups. In commercial processes the fact that there are many different sulfur species that are simultaneously being converted makes describing inhibition with a single equation an almost impossible task. A particularly relevant comment to this effect was made by Stanislaus and Cooper and is quoted here (109): In general, the use of Langmuir-Hinshelwood-Hougen-Watson (LHHW)-type of rate equation for representing the hydrogenation kinetics of industrial feedstocks is complicated, and there are too many coefficients that are difficult to determine. Therefore, simple power law models have been used by most researchers to fit kinetic data and to obtain kinetic parameters.

This statement could well be expanded to include studies describing the kinetics of model compounds. In reviewing the literature, one finds that there are almost as many kinetic representations as there are researchers and/or model compounds. Even the same authors have found it necessary to use different equations to describe the different responses to inhibitors for closely related sulfur species such as thiophene, benzothiophene, and dibenzothiophene (104, 122, 123). The inhibiting effect of H2S for the hydrogenation of a simple molecule, such as toluene, has been found to require extremely complex equations to adequately describe mathemati-

442 D. DUAYNE WHITEHURST, TAKAAKI ISODA, AND ISAO MOCHIDA cally the fact that the functionality of the effect varies with the absolute level of H2S in the reactor (107). This raises some serious problems for people just entering this field, and perhaps it might be worthwhile to consider what chemical phenomena are responsible for the observed inhibiting effects that have and are being reported. The basic premise of the original kinetic description of inhibition was that, for a reaction to proceed on a surface, one or more of the reactants (A) must be adsorbed on that surface in reversible equilibrium with the external solution, having an equilibrium adsorption constant of KA , and the adsorbed species must undergo some transformation involving one or more adsorbed intermediates ( a ) in the rate-limiting step, which leads to product formation. The product must desorb for the reaction cycle to be complete. If other species in the reaction mixture (I) can compete for the same adsorption site, the concentration of the adsorbed reactant (AAD)on the surface will be lower than when only pure reactant A is present. Thus, the rate of conversion will depend on the fraction of the adsorption sites covered by the reactant (0,) rather than the actual concentration of the reactant in solution, and the observed rate coefficient (kobs)will be different from the true rate coefficient (ktrue).In its simplest form the kinetic expression for this phenomenon in a first-order reaction can be described as follows: rateobs= kOdA1= k t r u e e A

0,

+ KA[A] + K&i] + . . . + K1,[1,]).

= KA[A]/(~

(1) (2)

If the reaction is nth order in the adsorbed reactant A, then this term is raised to the power n. The term in the denominator can be envisioned as relating to the degree of inhibition, as shown in Eq. (3), and will vary between 0 (complete inhibition) and 1 (no inhibition). If K,[A] becomes very large in comparison with the sum of the other terms in the denominator, Eq. (2) will reduce to a value of 1, and the rate will become zero order in the reactant A. Degree of inhibition = 1/(1

+ KA[A] + KI,[I~]+ . . . + KI,[~,])

(3)

In HDS reactions, the rates of conversion are generally observed to be first order in hydrogen as well as in the sulfur compound, so Eq. (2) expands to include hydrogen with its own fractional coverage function (eH2). rateobs= kObSIAl[H,l= k t r u e 0 A 0 H 2

(4)

For this reason many authors have separated the functionality of hydrogen from that of the sulfur compound being investigated. Another problem specific to hydrogen is that the concentration of hydrogen in the liquid at

POLYAROMATIC SULFUR COMPOUNDS

443

the catalyst surface is difficult to assess, and so most authors describe hydrogen concentration in terms of the applied hydrogen pressure rather than concentration at the catalyst surface. For vapor phase processes this is perfectly adequate, but for liquid phase reactants, the instantaneous hydrogen concentration in liquid-filled catalyst pores is governed by hydrogen solubility. Hydrogen exhibits an abnormal temperature-solubility relationship as the solubility increases with temperature, which is beneficial for hydrogenation processes. The solubility limit does, however, vary with the nature of the liquid. For perspective, hydrogen solubilities have been measured for coal liquefaction solvents (124)at about 300°C and 5 MPa. The observed solubility was about 0.1 wt% or about 0.5 mol/L. Coal liquefaction solvents are much more aromatic than diesel fuels but as a first approximation the solubilities should be similar. At high hydrogen concentrations, the hydrogen adsorption term (KHZ[H2])could be similar in magnitude to the adsorption term of the trace amounts of the sulfur compound reactant (K,[A]). Thus, the kinetic expressions for inhibition become quite complicated. It has also been reported that, in HDS reactions of dibenzothiophenes, the inhibition kinetics of the direct sulfur extraction route (kDJ require a different mathematical treatment from those of the hydrogenation route (kHsI)(104). A possible explanation for this observation could be that the rate-limiting steps in the two processes may involve the interaction of different numbers of adsorbed intermediates [nin Eq. (2)]. Vanrysselberghe and Froment proposed that the rate-limiting step in direct sulfur extraction from dibenzothiophene involves the simultaneous interaction of an adsorbed dibenzothiophene intermediate with two adsorbed hydrogen atoms (104). The question of why different molecules require different inhibition kinetic equations remains to be answered. It could be that the rate-determining mechanistic steps are different for the different molecules. Considering the mechanisms for soluble metal complex analogs of HDS catalysts, discussed in Section IV, it can readily be seen that different molecules could well exhibit different rate-limiting mechanistic steps. For example, even for the simplest case of direct extraction of sulfur from dibenzothiophene with Co-Mo-S/A1203, several intermediates are involved. Using the mechanism proposed by SAnchez-Delgadofor a single metal site mechanism (see Fig. 26, pathway 111), the steps would be the following: 1. 2. 3. 4.

Adsorption of DBT Oxidative addition of DBT onto Co Dihydrogen dissociation on either Co or Mo (with transfer to Co) Hydrogen insertion into the Co-C bond (reductive elimination)

444 D. DUAYNE WHITEHURST, TAKAAKI ISODA, AND ISAO MOCHIDA

5. Hydrogen insertion of a second hydrogen atom into the Co-S bond 5a. Or repetition of steps 3 and 4 to cleave the Co-S (T bond (at this point the intermediate would be 2-mercaptobiphenyl) 6. Repetition of steps 2-5 to produce biphenyl and Co-SH2 7. Desorption of biphenyl 8. Desorption of H2S It should be kept in mind that, for this simple case, three coordination positions must be available on the Co for the mechanisms to proceed. If one or more of these coordination sites are occupied by a nonparticipating molecule such as an aromatic feed component or H2S,the rate of progress through this sequence will be reduced. Similarly, the DBT and H2 must be adsorbed to enter the sequence, and if competing molecules occupy coordination sites on Co, the rate of progress through the sequence will also be reduced. The latter point is particularly significant when one tries to explain why DBT is self-inhibiting. If the rate-determining step requires adsorbed (or dissociated) dihydrogen, then a second molecule occupying a Co coordination position would inhibit dihydrogen adsorption. Thus, lower rates would be the result. Several authors have suggested that competitive dihydrogen adsorption is critical to the overall rate of desulfurization (104).It is believed that many of the reported observations for inhibition of desulfurization are, in fact, the result of limitations on dihydrogen adsorption rather than inhibition of adsorption of the sulfur compound being investigated. This point will be expanded later. Following similar logic, if the mechanistic route follows the pathway in which an adsorbed molecule is first attacked by a hydride ion and then a proton (Fig. 25, pathways I, 11, and IV), the rate-limiting step could depend on the availability of either a hydride ion or a proton in the vicinity of the adsorbed intermediate. The lability of the molecule toward hydride attack would also be a function of the degree of electron donation from the T system to the metal to which it is adsorbed and to steric limitations imposed by alkyl substituents on the aromatic ring systems. In these pathways, two or three bonding orbitals must be available on the Co for the mechanisms to proceed. Occupation of these orbitals by some nonreacting molecules would result in inhibition. The foregoing simple arguments pertain to single-site mechanisms. If multiple sites are involved in the reaction mechanisms, then the problem of occupation of bonding orbitals on one or more sites by nonreacting molecules further complicates the picture. It has been suggested that, even with homogeneous metal clusters of Mo and Co, dihydrogen dissociation may occur on Mo, and the dissociated hydrogen may be transferred to

POLYAROMATIC SULFUR COMPOUNDS

445

the adjoining Co, where it then inserts between Co-S or Co-C bonds (102).

3. Attempted Determination of LHHW Parameters in HDS Reactions From the preceding discussion, it can be seen that the mathematical description of the chemical transformations involved in product formation can be extremely difficult. However, knowledge of the response of HDS reaction rates to different kinds of feed components and byproducts is extremely important for designing new processes that will allow refineries to meet the stringent standards of the future. The following text attempts to summarize the observations reported in the literature on the effects of inhibitors on the hydrodesulfurization rates of alkyl-substituted dibenzothiophenes. It is quite possible that many reports have been overlooked, and the present authors apologize for any oversights that may have occurred in this review. Many authors have attempted to determine the Langmuir-Hinshelwood parameters for various thiophene derivatives and relevant inhibitors. It is generally agreed that two different sets of values for adsorption constants are necessary to describe the kinetic behavior of HDS catalysts (1,5, 104). One set describes the site that catalyzes direct desulfurization and the other set describes the site that catalyzes hydrogenation. However, there does not seem to be a general agreement in the absolute or even the relative values for the adsorption constants of different compounds, as discussed in the following text. In addition, different modifications of the LangmuirHinshelwood equation have been reported by different groups for the same compounds reacting in the presence of the same catalyst under comparable conditions. These modifications differ mostly in the exponential term ( n ) that may appear in expressions comparable to Eqs. (2) and (4). An attempt has been made to summarize the available literature for comparison of adsorption constants and forms of the equations used. Table XV presents a number of parameters reported by different authors for several model compounds on CoMo/A1203in the temperature range 235350°C (5,33,104,122,123,125-127). The data presented include the adsorption equilibrium constants at the temperatures employed in the studies and the exponential term ( n ) of the denominator function of the 8 parameter that was used in the calculation. The numbers shown in parentheses, relating to the value of n, indicate that the hydrogen adsorption term (KH[H2])is expressed as the square root of this product in the denominator. Data are presented for both the direct sulfur extraction site (CT)and the hydrogenation site (T).

TABLE XV Reported Langmuir-Hinshelwood Parameters in the HDS of Thiophene Derivatives” Equilibrium Adsorption Constants and Exponential Terms for the Desuifuriration Site (u) Gates (126) (CoMolAl,O,, 300°C) n

K,,

Gales (126) (CoMo/A120,, 350°C)

K,

n

Compound H2 H2S Thiophene Benzothiophene Dihydrobenzothiophene Dibenzothiophene 4-Methyldibenzothiophene Butene Biphenyl Cyclohexylbenzene Ethylbenzene

Gates (127) (CoMo/A1203, 275°C) n

K,

1 2

1.8 870

2

11.0

Satterlield (125) (CoMo/A1203, 235°C)

Satterfield (125) (CoMo/A120,, 251°C)

Satterfield (125) (CoMo/AIzO,,

265°C)

Compound Hz H2S Thiophene Benzothiophenc Dihydrobenzothiophcnr Dibenzothiophene 4-Methyldibenzi)lhiophene Butene Biphenyl Cyclohexylbenzene Ethylbenzene ~

Gates (120) (CoMo/A120,. 350°C)

n

K,

n

K,

2 2

32.2 25.h

2 2

27.2 12

Gates ( 1 2 7 ) (CoMo/A120i, 275°C) n

K,

K,

n

K,

n

K,,

n

K,,

n

K,,

n

K,

2 2

0.31 0.43

2 2

0.13 0.23

2 2

0.056 0.25

3.5 3 3

0.536 91.2 13.7

3.5 3

0.358 211

3.5 3

0.702 62.8

3 3

19 19

3

75.7

Sattertield (125) (CoMo/A1203, 235°C)

Satterfield (125) (CoMo/Ai20,, 251°C)

(T)

Satterfield (125) Van Panis (122. 123) Van Panjs (122, 12.7) Froment (104) (CoMo/A120i, (CoMo/A1203. (CoMo/Al20,, (CoMo/A120,. 260°C) 300°C) 265°C) 260°C)

n

K,

n

K,

n

K,

1

0091

I

0.098

I I

0.019 0.012

1 1

0.013 0

0

n

K.

2.5

0.0602

2

8.87

n

3 I

23.0

2

~

n

0

3.5 3

0

3

2.52

3 3

low

0.0142

2.0s

8.87

4.6

3 ~

K,

K.

7.5 2

2

Froment (104) (CoMo/A120,, 300°C)

n

Equilibrium Adsorption Constants and Exponential Terms for the Hydrogenation Site Gates (126) (CoMo/A1203, 300°C)

Van Parijs (122, 123) Van Panjs (122, 123) (CoMo/A120,, (CoMo/AI2Oi, 260°C) 260°C)

~

“Consult cited references for units of K , and K,; relative values of these paranicters should he compared. hut not absolute values.

2Y4.0

l.6Y

POLYAROMATIC SULFUR COMPOUNDS

447

It may be seen that only a limited amount of data of this kind was found for comparison, and so it is difficult at present to draw many general conclusions. In addition, there does not appear to be agreement as to which compound in a reaction mixture has the strongest interaction with the catalyst. Most of the recent publications do agree that hydrogen has a very low equilibrium adsorption constant and its adsorption would be severely inhibited on both types of sites by H2S, aromatic compounds, and thiophene derivatives. From that point on, there is little agreement. The most recent, and perhaps the most thorough, study is that by Froment and Vanrysselberghe (104). Perhaps the largest discrepancies in reported results are the relative values for the adsorption constants of H2Sand thiophene molecules (THs, including thiophene, benzothiophene, and dibenzothiophene). The reported prefer> THs ence for adsorption on the direct desulfurization site ranges from H2S> (122,123,125)to about the same (104) to H2S << THs (125). Satterfield’s studies indicated that, as the temperature was increased, the preference for adsorption of THs becomes larger (125), but the differences between authors is far more than can be explained by the different temperatures of their experiments. The various parameters are summarized in Table XV. The report of Froment may provide the best guidelines at present (104).That report indicates the following relative preferences for adsorption on the direct desulfurization site (v):

KAD(CT site): Dibenzothiophene > H2S >> biphenyl > > H2. Reported adsorption preferences on the hydrogenation site for different compounds provide no better guidance than was found for the direct desulfurization site. Again, the most recent reports agree that dihydrogen adsorption is very low relative to that of the other species being considered. Surprisingly, Froment found no indication for inhibition of hydrogenation of aromatic rings by H2S (104). As discussed earlier, H2S has been shown to be a strong inhibitor for toluene hydrogenation when present in concentrations normally found in HDS processes. Satterfield reported that thiophene and H2S had similar adsorption equilibrium constants on the hydrogenation site and both declined equally with increasing temperature (125). As far as the other reactive species in HDS conversions are concerned, the report by Froment offers the best guidance for the relative adsorption constants on the hydrogenation site (7) and indicates the following preferences (104):

KAD(T site): Dibenzothiophene > biphenyl >> H2. The relative adsorption constants for H2S and dibenzothiophenes are discussed in more detail later.

448 D. DUAYNE WHITEHURST, TAKAAKI ISODA, AND ISAO MOCHIDA Obtaining adsorption constant data in complicated reaction systems, such as HDS processes, is difficult as can be seen from the preceding discussion. It is often more instructive to determine the relative adsorption behaviors for competing materials in binary mixtures. This has been done by many authors and this approach is discussed next. 4. Inhibition Studies with Binary Mixtures

To determine the relative inhibitive effects of different materials on a given reaction, it is often easiest to conduct sets of experiments in which the reactant of interest is converted under common conditions with and without the addition of an inhibiting material. Different additives can then be compared as to their relative effects on the conversion and selectivity of the reactant conversion. Some very useful information of this nature has been obtained by Nagai for the relative equilibrium adsorption constants of different sulfur compounds on MoS2catalysts during the denitrogenation of perhydroacridine (PHA) (129, 130). It was assumed that the sulfur compounds were adsorbed competitively with the PHA on the denitrogenation site. Table XVI presents the relative adsorption constants obtained by fitting the LHHW parameters in competitive experiments with binary mixtures of PHA and various sulfur compounds. For ease of comparison, the observed competitive adsorption constants have also been normalized to a common value for perhydroacridine in the last column of the table. These data indicate that the most strongly adsorbed sulfur compound on the site responsible for denitrogenation is dibenzothiophene. The trend within the thiophene derivative series is in agreement with the theoretical calculations for electron density on sulfur presented in Figs. 28 and 29. Studies like this could be quite useful in determining whether the reason alkyldibenzothiophenes have low reactivities is related to limitations in TABLE X V I Competitive Adsorption Constants in H D N of Perhydroacridine (Mo/Al2OS,288-360"C, 10 MPa) (129, 130)

Compound

n

Dibenzothiophene Benzothiophene Thiophene Dimethyl sulfide Ethanethiol H2

2 2 2 2 2 2

'K

=

KPHA (Perhydroacridine)

Ks (S compound)

1350 720 560 910 830

16.9 6.9 5.2 3.2 1.o

1000 for perhydroacridine.

KH,

Relative K" 12.5 9.6 9.3 3.5 0.0

<1

POLYAROMATIC SULFUR COMPOUNDS

449

adsorption or to limitations in some mechanistic step such as oxidative addition of the adsorbed species to the active site. Another approach is to conduct competitive experiments with binary mixtures in which the complete reaction pathway is developed according to a reaction scheme like that of Scheme 1 described in the beginning of this review or like those shown in Figs. 12-15. Much of the confusion found in past reports of the kinetics of dibenzothiophene and its alkylated derivatives has come from incomplete deconvolution of the reaction network. Selectivity is often reported as the ratio of the yields of biphenyls (direct sulfur extraction) to the yields of cyclohexylbenzenes (hydrogenative route). As discussed in Section IV, cyclohexylbenzenes are produced via two different routes and, unfortunately, even low-conversion studies do not circumvent this confusion. To illustrate how conclusions can often be confused if the wrong model is used, some examples of reported competitive inhibition experiments will be discussed. It has been reported that H2S is a strong inhibitor for the direct sulfur extraction site but does not inhibit the hydrogenation site (127, 128). It will be shown later that the major inhibiting effect of H2S is on the hydrogenation of biphenyl to cyclohexylbenzene. Thus, the product distribution changes but relative contributions of the direct sulfur extraction route (kDJto the hydrogenative route (kHsl)do not change as much as is inferred from the change in yields of biphenyl and cyclohexylbenzene. This can be seen by recalculating the data reported by Vrinat characterizing the effects of H2S inhibition on benzothiophene and 4-methyldibenzothiophene (32). The reported data were sufficient to develop rate constants for all of the relevant steps in a manner similar to that shown in Fig. 12. The relative rate constants for hydrogenation of benzothiophene (kHsl) and biphenyl (kHpl)were in line with the relative values observed in Fig. 12 and the data recently reported by Froment (104) (see Table XV), which established that benzothiophene is more readily hydrogenated than biphenyl. The deconvoluted rate constants for HDS of benzothiophene and 4-methyldibenzothiophene in the presence and absence of 6/1 H2S are presented in Table XVII. Examining Table XVII, one can see that it is true that, for dibenzothiophene, the direct desulfurization route (kDo)suffers the greatest inhibition by H,S; only 7.5% of the original activity remained in the inhibited case. However, the hydrogenative route (kHsl) was also severely inhibited by H2S;only 25% of the original activity remained. The largest inhibition was, in fact, in the hydrogenation of biphenyl ( k H p , ) . These effectswere even more dramatic in the case of 4-methyldibenzothiophene. For that compound, kDoand kHs, were both equally inhibited, so that the selectivity was not changed by H2S inhibition. The biphenyl hydrogenation (kHp,),on the other hand, suffered severely. These results clearly

TABLE XVII Effect of H2S on the Pseudo-First-Order Rate Constants for Dibenzothiophene and 4-Methyldibenzothiophene (NiMo/A1zO.l,5 MPa H2, 290°C) (33)" Starting compound

-

Additive Overall rate constant x Reactant product

4s4%

Starting compound

DBT

DBT

Fractional activity

None 15.9

611 H2Sh 1.2

0.075

4-MDBT

4-MDBT

Fractional activity

None 4.5

611 H2Sh 0.70

0.156

4-4

13.30

0.56

0.042

0.91

0.13

0.145

@CY, 4 +CY,

398.00

14.00

0.035

25.37

5.54

0.218

4 s 4 k @SCy,

2.66

0.66

0.248

3.62

0.53

0.146

+ f p k @-Cy,

2.66

0.09

0.035

1.81

0.03

0.014

0.25 27.9 2.0 2.5

0.25 42.0 20.3 25.4

k"

Catalyst selectivities kD,,lkHs, k ~I b, ( , km, / h i p , (kn,, + k H s , ) / k w

5.0 29.9 1.o 6.0

0.85 25.0 7.0 13.0

Rate constants were obtained by recalculation of original data. 6/1 indicates the mole ratio of H2S and the benzothiopbene compound.

POLYAROMATIC SULFUR COMPOUNDS

451

show that the hydrogenation site is subject to inhibition by H2S. Thus, unless the reaction matrix is completely deconvoluted, the conclusions one draws may be questioned. Similar effects are noted for the inhibition of 4,6-dimethyldibenzothiophene (4,6-DMDBT) by naphthalene (21, 133, 134). Table XVIII shows the recalculated rate constants for the HDS of 4,6-DMDBT in the presence and absence of 10% naphthalene in decalin solvent with a NiMo/A1203 catalyst at 320°C and 2.5-MPa hydrogen pressure. The table shows that the selectivity for 4,6-DMDBT reactions was not severely changed by naphthalene inhibition; all rates were lowered by about the same amount. The activity for hydrogenation of aromatic hydrocarbons was inhibited more than for the hydrogenation of sulfur-containing compounds. Interestingly, the desulfurization of the hydrogenated derivatives of 4,6-DMDBT was inhibited less than that of the fully aromatic parent.

TABLE XVllI Effect of 10% Naphthaleiie on the Pseudo-First-Order Rate Constants in the HDS of 4,6-Dirnethyldibenzothiophene(NiMo/A1203,2.5 MPa H2, 320°C) (21, 41a, 133) Starting compound 4,6-DMDBT

4,6-DMDBT

Fractional activity

None 29.6

10% NAPH 9.6

0.324 0.324

7.0

2.6

0.366

308

160

0.519

514

321

0.625

22.6

7.1

0.314

18.5

5.8

0.314

6.2

1.6

0.258

3.6

0.98

0.271

0.3 44.0 3.6 4.8

0.4 62.5 4.4 6.0

452 D. DUAYNE WHITEHURST, TAKAAKI ISODA, AND ISAO MOCHIDA This observation is consistent with stronger bonding between the more basic sulfur species and the desulfurization site as discussed earlier in and shown in Table XI11 (2) and Figs. 28 and 29 (17, 38). The relative adsorption constants for hydrogen, naphthalene, and dibenzothiophene have been estimated by determining how the rate constants, as described in Scheme 1, change as a function of the naphthalene concentration in HDS reactions of dibenzothiophene with NiMo/A1203catalyst at 320°C and 2.5-MPa hydrogen pressure (21, 233, 134). The original data were used to determine the rate constants as described in Section IV, and the LHHW parameters were then estimated. The values obtained in this way are presented in Table XIX. These values should be considered only as estimates as data were limited, but the trends are believed to be valid. Examining the table, one can see that the relative preference for sorptions are the following: Desulfurization site: Dibenzothiophene >> naphthalene > hydrogen Hydrogenation site: naphthalene > dibenzothiophene >> hydrogen. The changes in the various rate constants with increasing naphthalene concentration are in general agreement with the more accurate data shown in Table XVIII (compare rate constants with and without 10%naphthalene). However, it appears that the direct extraction of sulfur from dibenzothiophene is less susceptible to inhibition than that of 4,6-DMDBT but that

TABLE XIX Changes in Relative Pseudo-First-Order Constants for HDS of Dibenzothiophene with Increasing Naphthalene Concentrations (NiMo/Al2Oj,320"C, 2.5 MPa H 2 ) (21, 41a, 133)

(mol/L)

NAPH (mol/L)

Hza (mol/L)

0 .5 10

0.0146 0.0147 0.0148

0 0.351 0.707

0.26 0.26 0.26

30

0.0151

2.166

0.26

DBT %

NAPH

koOh

kD1"

kHsIh

hih

84.7 61.8 46.9 24.5

1694 1226 954

20.3 6.2 3.3

468

1.7

1.5.2 4.3 2.3 1.2

Estimated relative adsorption constants n

KDBT KNAPH KH,

2 25 0.4 0.2

2 2.5 0.4 0.2

1 4 8 0.02

1 4 8 0.02

a Based on estimated hydrogen solubility. Rate constants were obtained by recalculation of original data.

POLYAROMATIC SULFUR COMPOUNDS

453

dibenzothiophene hydrogenation is more susceptible to inhibition than that Of 4,6-DMDBT. These results show that, in equimolar concentrations, naphthalene would not be considered as a strong inhibitor toward direct sulfur extraction (b,,) for PASCs. However, as discussed earlier, the content of di- and trinuclear aromatics in diesel fuels and gas oils can be as high as 20-30%. whereas the level of sulfur compounds in today's diesel fuels is only 0.2% sulfur, or about 1 wt% PASCs. So the competition for the active site by aromatic hydrocarbons is very strong. Their effect on the direct desulfurization route will lower the rate to about one-third of the noninhibited rate in the case of dibenzothiophene and would lower that of 4,6-DMDBT even more. The hydrogenative desulfurization route (kHs,)presents a somewhat different picture. Dibenzothiophene hydrogenation is very strongly inhibited by aromatics. A t the 20% aromatics level, the rate of hydrogenation would be expected to be only about one-tenth that of the uninhibited rate. Although less susceptible to inhibition, the hydrogenation of 4,6-DMDBT would be only about one-fifth that of the uninhibited rate. This is particularly harmful in the case of NiMo/A1203-catalyzed desulfurization of alkylsubstituted dibenzothiophenes because, with this catalyst, hydrogenation prior to desulfurization is the most important mechanistic pathway, as described by Mochida and co-workers (14, 17, 21, 41a, 133). This observation suggests that Ni-Mo-S catalysts would be preferred over Co-Mo-S catalysts for the HDS of 4,6-DMDBT, as Co-Mo-S is not a very active hydrogenation catalyst. However, in the presence of aromatics, the hydrogenative route suffers the most inhibition, and the advantage of Ni-Mo-S catalysts disappears (133). Unfortunately, diaromatic hydrocarbons are not the only potential hydrocarbon inhibitors present in gas oils and diesel fuels. Triaromatic hydrocarbons are also present in significant amounts (see Fig. 2) (12). It is known that triaromatics, such as phenanthrene, are even stronger inhibitors than diaromatics for the HDS of thiophene compounds. Equilibrium adsorption constants for phenanthrene and naphthalene have been reported to be 65 and 11 atm-', respectively (231).In Iranian gas oil, triaromatics have been reported to be present at about one-tenth the concentration of diaromatics (109).Thus, the contribution to inhibition of HDS reactions by triaromatics (KT,,[Tri]) could be about the same as that from diaromatics, even though triaromatics are present in smaller amounts. In support of this suggestion, Kabe and co-workers have demonstrated that phenanthrene is a fairly strong inhibitor of the hydrogenation site in dibenzothiophene HDS (128).However, in gas oils other trace components such as nitrogen- and oxygen-containing materials can be even stronger

454 D. DUAYNE WHITEHURST, TAKAAKI ISODA, AND ISAO MOCHIDA

inhibitors in the HDS of dibenzothiophene. The inhibiting effects of polar nitrogen- and oxygen-containing compounds on HDS and hydrogenation reactions are well documented in extensive reviews by others (5, 20, 109, 130, 132) and will not be discussed in detail here, as diesel fuels and gas oils that have been hydrotreated to a level of 0.2% sulfur contain only very small amounts of such materials. The reader is referred to those prior reviews for more information. However, for perspective, the following order of the degree of inhibition for both the direct extraction and hydrogenative desulfurization routes for dibenzothiophene has been reported as follows (130): acridine > carbazole > phenothiazine > dicyclohexylamine > phenanthrene.

5. Summary of Observations Refating to Inhibition in HDS Reactions

In summary, the following conclusions can be drawn relating to the limitations of inhibition of HDS reactions by other feed components. The two most important classes of inhibitors in HDS reactions of diesel fuels and gas oils, when sulfur levels are reduced from 0.2% S to lower levels, are aromatic hydrocarbons, which are natural components present in the fuels, and H2S, which is produced as a product of the HDS reaction itself. Inhibition by aromatic hydrocarbons is most severe for dialkyldibenzothiophenes as these materials are preferentially desulfurized by hydrogenation of the aromatic ring prior to sulfur extraction and as aromatics are more strongly adsorbed on the hydrogenation site of Co(Ni)-Mo-S catalysts than are dibenzothiophenes. Inhibition by H2S severely inhibits both the hydrogenation site and the sulfur extraction site of Co(Ni)-Mo-S catalysts, but the inhibition is greater for the sulfur extraction site. The degree of inhibition is less for alkyl-substituted dibenzothiophenes than for the unsubstituted ring system, but the absolute rates of desulfurization of alkyldibenzothiophenes are so low that any inhibition is a major problem when attempting to meet the new 0.05% S specifications. It appears that the inhibition by both H2S and aromatic hydrocarbons involves competition between the inhibitor and the reactive sulfur compound for adsorption on the active site. However, inhibition could also be the result of occupancy of one or more of the bonding orbitals of the Co(Ni) by some nonreacting molecule, such as H2S or naphthalene. This would prevent the oxidative addition of the thiophene ring to the Co(Ni) in a mechanistic sequence such as that described in Figs.

POLYAROMATIC SULFUR COMPOUNDS

455

25 and 26b. For example, in Fig. 22b, the coordinated H2S would prevent thiophene from oxidatively adding to Co. The least strongly sorbed reactant in HDS processes is dihydrogen. It could well be that the major reason for rate reduction by inhibitors is prevention of dihydrogen adsorption by competing molecules. This would include the sulfur-containing reactant as well. This would explain the phenomenon of self-inhibition in HDS reactions. More research in the area of inhibition of the reactions of alkyldibenzothiophenes is strongly recommended to aid in finding the means to overcome inhibition limitations and provide suggestions for new catalysts and/or processes that will be able to meet the new stricter low-sulfur specifications. Some novel concepts for improved HDS processes that address these specific problems are discussed in the next section. VII.

Novel Approaches for Deep Desulfurization

A. REVIEWOF DIFFICULTIES In Sections 111,IV, and VI, several problems were identified that present obstacles to the achievement of the ultralow levels of sulfur demanded by new specifications for fuels in the not too distant future. These may be categorized into the following groups: Low reactivity of alkyl-substituted dibenzothiophenes (PASCs) that are substituted in positions adjacent to the sulfur atom Insufficient activity of present catalysts to overcome the low-reactivity PASCs using existing moderate-pressure ( 4MPa) hydrotreating reactors Inhibition of desulfurization reactions by aromatic hydrocarbons present in the fuels being treated Inhibition of desulfurization reactions by H2Sproduced as a byproduct in the HDS process Excessive hydrogen consumption due to nonselective hydrogenation of aromatic hydrocarbon components in the fuels being hydrotreated Temperature limitations in HDS processes imposed by thermodynamic limitations on concentrations of intermediates Temperature limitations in HDS processes imposed by product quality degradation such as color-fluorescence formation Temperature limitations in HDS processes imposed by side reactions of the major feed components that lead to lower product yields Reluctance of refiners to purchase new process equipment that would ease these pro'cess limitations

456 D. DUAYNE WHITEHURST, TAKAAKI ISODA, AND ISAO MOCHIDA In this section, a number of published novel approaches are described that address these specific problems. Many new processes have been claimed, and it is certain that not all of these have been identified by the present authors. Several other reviews have some information relating to this specific topic ( I , 109, 112,113), and it is suggested that those references be consulted as well. In addition, there have been two recent major symposia held by the American Chemical Society, Division of Petroleum Chemistry, that have dealt specifically with this topic (Washington, DC, August 1994, and Orlando, FL, August 1996). Papers presented at those symposia are also quite instructive. B. NOVELAPPROACHES THATHAVEBEENREPORTED The approaches can be divided into several categories, as follows: New catalysts with higher activities Catalysts that offer alternative reaction pathways for desulfurization Staged process operations Novel reactor designs Alternative approaches for lowering sulfur levels 1. New Catalysts with Higher Activities

Several recent reports have identified new procedures for preparing Co(Ni)-Mo-S/A1203 type catalysts that provide much higher dispersions of the base catalyst (MoS2) and procedures that generate the active component SBMS or Co(Ni)-Mo-S with high selectivity (1, 2, 56, 61, 76-79). These procedures were reviewed in Section IV. There are reports that the use of other supports, such as Ti02 or carbon, results in catalysts with higher activities than those of catalysts that can be made on conventional supports, such as alumina, tungsta, and silica ( I , 2,41,61,62, 79).Acidity in the support has been reported to enhance hydrogenation activity, and so attempts have been made to increase acidity through incorporation of silica-alumina, zeolites, and additives ( I , 109). Additives such as phosphoric acid, fluoride, and others can help to reduce the interactions between the Co(Ni)-Mo-S crystallites and the support, with moderate improvements in activities (1, 76, 77). Of these approaches, carbon-based catalysts seem to offer the greatest hope for novel catalysts with higher activities.To date, the varieties of carbons investigated have been extremely limited, considering the present state of the art in modifying the chemical and electrical properties of carbons. It is anticipated that future work in this area will provide some attractive new materials. Such catalysts may not be useful for fixed-bed operations with raw

POLYAROMATIC SULFUR COMPOUNDS

457

feeds, as regeneration will be problematic. However, when used as the finishing catalyst in a staged operation, they may be quite adequate, especially in treating 0.2% S diesel fuels that are already very clean. 2.

Catalysts That Offer Alternative Reaction Pathways for Desulfurization

As described in Section IV.B, dibenzothiophenes, when substituted in positions adjacent to the sulfur atom, have reduced activity for direct sulfur extraction. As a result, catalysts that promote aromatic ring hydrogenation offer another route to desulfurization, as the partially hydrogenated ring presents much less steric restrictions to adsorption via qJ-S type bonding (17,21) or to oxidative addition to form a metallathiabenzene intermediate, as discussed in Section IV.E.3. In addition, the metal-S coordination bond strength is increased by increasing the electron density on sulfur, and the C-S bonds in hydrothiophenes are much weaker. Unfortunately, the conventional catalysts that have the highest aromatic hydrogenation activities (Ni-Mo-S) are very susceptible to inhibition by the aromatic hydrocarbon components present in large amounts in today’s diesel fuels and gas oils. It would therefore be desirable to identify novel catalysts that have high hydrogenation activity and selectivities for aromatic sulfur compound hydrogenations in the presence of large molar excesses of aromatic hydrocarbons. The acidity of the support may help in this regard. Enhanced hydrogenation activity is discussed in the review by Stanislaus and Cooper (109), and several processes have already been announced that claim to provide the desired activity (109, 138-143). However, most of these processes use noble metals that have limited tolerance to sulfur, and they are not suitable for feeds containing 0.2% (or 2000 ppm) S. There is some promise for the use of more sulfur-tolerant catalysts such as RuSz (109, 144-148). In addition, RuSz is more selective for hydrogenation of sulfur-containing aromatics in preference to aromatic hydrocarbons. Some preliminary work in this area has been published by Mochida and co-workers (135, 136). That work showed that even physical blends of RuS2/A1203and CoMo/A1203were more effective for the desulfurization of 4,6-DMDBT, in the presence of 10% naphthalene in decane solvent, than would be expected from the statistical average of the combination of the two catalysts, as shown in Fig. 32a. With RuS2/A1203alone, tetrahydrodimethyldibenzothiophene was observed to be the only product and approached an equilibrium value of about 10%. When CoMo/Al2O3and RuS2/ Alz03catalyst particles were both added to the reaction mixture, the rate of conversion of 4,6-DMDBT increased, and the tetrahydro intermediate was observed in much smaller amounts. Thus, RuS2/A1203provided a

458 D. DUAYNE WHITEHURST. TAKAAKI ISODA, AND ISAO MOCHIDA

b

a

-.

0

Reaction time (h)

5 10 Concentration of naphthalene (wt%)

FIG.32. Improved reactivity for 4,6-dimethyldibenzothiopheneconversion in the presence of naphthalene inhibitor (320°C, 2.5 MPa). Reprinted with permission from Refs. 135 and 136, Isoda et al. (1996). Copyright 1996 American Chemical Society.

means to increase the rate of hydrogenation of 4,6-DMDBT to yield an intermediate that was easily desulfurized by the CoMo/AI2O3 catalyst. Reactions that employed mixtures of NiMo/A1203 and RuS2/A1203did not offer the same benefit, as NiMo/A1203is, itself, an active hydrogenation catalyst. However, in the presence of aromatic hydrocarbons, NiMo/A1203 is nonselective, and the aromatic hydrocarbons are hydrogenated in preference to 4,6-DMDBT. Thus, hydrogen consumption with NiMo/A1203is excessive. As discussed earlier, the hydrogenation of 4,6-DMDBT to the tetrahydro intermediate is thermodynamically limited at high temperatures. Higher hydrogenation activity could allow the use of lower temperatures with improved thermodynamic equilibria; however, inhibition by aromatic hydrocarbons present in the feed may force refiners to employ higher temperatures. In such a system, the overall rate of conversion of 4,6-DMDBT can possibly be increased by bringing the two different catalytic functions into closer proximity. Thus, there could be benefits to having both RuS2 and Co-Mo-S within the same catalyst particle. This was demonstrated by coimpregnating Ru, Co, and Mo into an A1203 support (136). Th'is new ternary sulfide catalyst was indeed found to be superior to the physical blends and was much more resistant to inhibition, as shown in Figs. 32a and 32b. Unfortunately, the initial promise of this approach was followed by disappointment, as the ternary sulfide catalyst rapidly lost activity with extended use. XPS analyses of used catalysts showed that Ru catalyzed the reduction of Mo, which led eventually to crystal growth of the MoS2, which resulted in loss of activity. Perhaps other supports having higher surface areas or

459

POLYAROMATIC SULFUR COMPOUNDS

improved metal-support interactions could overcome this difficulty. This approach of increasing the rate and selectivity of aromatic ring hydrogenation while reducing the sensitivity of catalysts to inhibition is, in principle, a worthwhile area of study, and more research in this area is needed. Another novel catalyst modification has been suggested in which the active Co-Mo-S catalyst is used in combination with an acidic catalyst such as a zeolite. This combination has the potential of opening another reaction pathway by isomerization of the alkyl groups on molecules such as 4,6-DMDBT to positions that do not sterically interfere with adsorption or oxidative addition. This is illustrated in Fig. 33. Gates and co-workers reported many years ago that the 2,8- and 3,7-dimethyldibenzothiophenes are much more easily desulfurized than 4,6-DMDBT (see Table XII) (26). Therefore, a combination of an isomerization catalyst and a desulfurization catalyst could be synergistic for removing dialkylbenzothiophenes. Several authors have pursued this approach and indeed observed that desulfurization of 4.6-DMDBT was increased when acidic zeolites were used in combination with conventional HDS catalysts (30, 31, 33, 137, 149-151). Figure 34 shows that there can be a great acceleration in the conversion of 4,6-DMDBT through the use of a hybrid catalyst consisting of CoMo impregnated into a composite containing 5% NiY zeolite and alumina (137). Low temperatures had to be employed, as at temperatures exceeding about 340"C, severe color fluorescence occurred in the product.

y \

CH,

m

b_ CH3

CH,

+

m

c

3

+W

c

4

CH,

-

C, - C, hydrocarbons

CH,

CH,

CH,

Cti3

FIG.33. Reaction pathways of 4,6-DMDBT conversion in the presence of zeolite containing CoMo/A1,03 catalyst: (a) HDS with isomerization route; (b) hydrocracking route; (c) direct desulfurization route.

460 D. DUAYNE WHITEHURST, TAKAAKI ISODA, AND ISAO MOCHIDA P

g

CoMolAlrOa + N I - H Y P

0

30 60 90 120 Reaction Time (min)

FIG.34. Improved HDS activity of 4,6-dimethyldibenzothiophenein the presence of a hyhrid CoMo/Alz03 and Ni-HY zeolite catalyst (270°C, 2.4 MPa, 0.1 wt% of reactant in 10 g of decane, CoMo/Al,O,:Ni-HY = 1.0:0.05 9). Reprinted with permission from Ref. 137, Isoda et al. (1996). Copyright 1996 American Chemical Society.

At 270°C desulfurization was enhanced and several isomeric dimethyldibenzothiophenes were observed in the product mixtures, showing that isomerization had indeed proceeded. Dibenzothiophenes containing more and less than two methyl groups were also observed in the products. This indicates that transalkylation was also a significant reaction. However, the major products of 4,6-DMDBT conversion appeared to be the result of cracking reactions. The decane solvent used in this study also was cracked to a significant extent, indicating that the acidity of this catalyst may have been too high, and significant yield losses in diesel fuel would be expected with this catalyst. Continuous bench-scale HDS operations with gas oil flowing over this catalyst demonstrated that no significant deactivation occurred in 700 h of operation (151). Studies have also been reported in which the type of zeolite was varied. ZSM-5 was compared with HY as an acid function in composites with Co and NiMo/A1203 (33, 149). As might be expected, the ZSM-5 additive did little to improve the HDS of 4,6-DMDBT, as the pore size is too small to allow facile entry to 4,6-DMDBT. Mechanistic studies were conducted at higher temperature (360”C), at which cracking was severe. More than 80% of the desulfurization of 4,6-DMDBT was found to occur via center cracking of the C-C bond joining the aromatic rings. Toluene was a major product. Unfortunately, if such a scheme were to be applied to gas oil or diesel fuel, the losses in desired product would be excessive. It has been claimed that, by adjusting the acidity, a stable commercial catalyst can be manufactured that contains a “small amount of zeolite”

461

PQLYAROMATIC SULFUR COMPOUNDS

.3

LHSV:4.0 l/h

*.

140

.*

1 developed catalyst

m

2

100

80

0

0.05

0.1

0.15

0.2

Sulfur in Product (mass YO) FIG.35. HDS performance of a developed catalyst. Reprinted from Cuful. Today 35; M. Yumoto, K. Usiui, K. Watanabe, K. Idei, and H. Yamazaki; 45. 0 1995 with kind permission of Elsevier Science-NL, Sara Burgerhartstraat 25, 1055 KV Amsterdam, The Netherlands.

and exhibits enhanced desulfurization activity for gas oil (150). Sulfur contents of less than 0.05% S could be produced, as shown in Fig. 35. The exact nature of the zeolite was not specified, nor was the extent of cracking of the gas oil being desulfurized, but the reported results look encouraging. It has also been noted that noble metals ion-exchanged into zeolites show improved desulfurization activity in the presence of H2S. Improved liberation of the H2Sfrom the metal was claimed to be enhanced by zeolite. Identifying new zeolite additives for HDS catalyst composites is a challenging area for future research. The majority of the components in gas oils and diesel fuels are quite labile to cracking, especially alkyl groups on aromatic rings. Transalkylation of alkyl groups between aromatic rings can be either desirable or disastrous, depending on whether the alkyl groups are transferred from or to the dibenzothiophene nucleus. Also if olefin intermediates in cracking side reactions alkylate the dibenzothiophene nucleus, the reactivity of the resultant products could be much lower than that of the parent sulfur compound. Discovering low-temperature processes that avoid cracking or other undesired side reactions while promoting isomerization of the alkyl groups on the refractory dialkyldibenzothiophenes is theoretically possible and should receive attention in the future. Thus, there is ample room for research into new zeolite compositions that can selectively do only those reactions that will lower the reactivity of the sulfur compounds in either simultaneous or sequential HDS processes. Research in this area is recommended.

462

D. DUAYNE WHITEHURST, TAKAAKI ISODA, AND ISAO MOCHIDA

3. Staged Process Operations

If new investment is considered, then there are many new approaches to reducing the sulfur level in gas oils and diesel fuels by means of sequentially staged operations. The first stage can be a mild-severity selective process to modify the feed so as to make desulfurization more facile, or it may be a severe process to lower the sulfur level but produce a product that must be further treated to meet some other specification. Examples of the first approach were mentioned in the previous discussion of new catalysts that offer alternative reaction pathways, such as isomerization of alkyl groups on dialkylbenzothiophenes into less sterically hindering positions or partial hydrogenation of the aromatic rings of alkyldibenzothiophenes so that subsequent HDS processing will be more facile. In the latter case, the hydrogenation process must be conducted at low temperature so that thermodynamic limitations are not problematic. Subsequent second stages should not reverse the benefits achieved in the first stage. There are relatively few examples of this approach for improved HDS processes. The second approach, high severity followed by a finishing step, is more common, and several examples of this have been reported (14,109,114-116, 138-140). The major concern in this approach is that high severity with present commercial catalysts produces color fluorescence in the low-sulfur product, and hydrofinishing is necessary to restore quality. An example of exploratory studies to identify the optimal conditions in a two-stage process using conventional catalysts is shown in Table XX (114). The table shows TABLE X X Two-Stage Desulfurization of Diesel Oil at Moderate Pressure (5 MPa) (114) Reaction condition“ (“C-kg/cm2-h)

Catalyst First

Second

First

Second

CoMo NiMo CoMo NiMo CoMo CoMo CoMo CoMo CoMo CoMo NiMo CoMo

CoMo NiMo NiMo CoMo NiMo NiMo NiMo NiMo NiMo CoMo NiMo NiMo

320-SO-1 320-50-1 320-SO-1 320-50-1 300-SO-1 340-SO-1 360-SO-1 320-SO-1 320-SO-1 340-50-1 340-SO-1 340-50-1

320-50-1 320-50-1 320-50-1 320-50-1 320-50-1 320-50-1 320-50-1 300-50-1 340-50-1 340-50-1 340-50-1 340-50-1

~~

~

Temperature

=

~~

Sulfur content (wt%)

0.080 0.074 0.049 <0.05

0.074 0.051 0.028 0.109 0.049 0.026 0.024 0.027 ~~

“C; pressure = kg/cm*; reaction time

=

hr

Color transparent transparent transparent light fluorescence transparent transparent transparent transparent transparent light green transparent transparent

POLYAROMATIC SULFUR COMPOUNDS

463

that the best combination in a two-stage moderate-pressure (5 MPa) process that can produce transparent colorless products with less than 0.05% S consists of higher temperature (340-360°C) CoMo/A1203catalyst in the first stage, followed by a NiMo/A1203-catalystin the second stage at milder temperatures (320°C). Each step requires about 1h of residence time. Thus, the reactor volume may be larger than desired. An alternative approach is to use three stages with shorter residence times in each stage (116).This approach is shown in Table XXI. The first two stages are high-temperature/ shorter times (360°C, 30 min, CoMo/A1203,then 360°C, 20 min, NiMo/ A1203)and a final lower-temperature finishing stage (260"C, 10 min, NiMo/ AI2O3). The combination of these three steps requires less than 1 h of residence time, and so the reactor sizes are reduced considerably relative to those of the two-stage approach. A further advantage of this three-stage approach is that the reaction pressure can be lowered from 5 to 3 MPa while still achieving a high degree of desulfurization with no loss in quality due to color-fluorescence formation. Another advantage of staged operations is that the reaction gas may be changed between stages. As discussed earlier, H2S is a severe inhibitor for HDS of alkyldibenzothiophenes. Figure 9 illustrates the progress of desulfurization in a typical HDS process. The first sulfur to be removed is quite easy, and HDS reactivity is not seriously lowered by H2S inhibition. However, if the H2S remains in the reactor, the refractory sulfur compounds, although less sensitive to inhibition, become even more difficult to convert. One approach to relieve this situation is to stage the overall process and remove the H2S produced in the first stage so that the second stage can continue with less inhibition. The refractory sulfur compounds can even inhibit their own TABLE XXI Properties and Composition of Feed and Product Oils f r o m Three-Stage HDS at 2.9 M P a (116) Stage Catalyst Condition (Y-min) S content (wt%) RFI H distribution (%) Ha" H, H, H,

Feed

First CoMo 360-30

Second NiMo 360-20

Third NiMo 260-10

0.706 34.9

0.092 25.7

0.035 30.2

0.016 5.0

5.8 7.0 60.7 26.5

5.7 6.6 60.0 27.7

5.4 6.5 60.7 27.4

4.7 5.6 61.0 28.0

a Hydrogen distributions are for hydrogens on aromatic rings (Ha), hydrogens one carbon removed from an aromatic ring (Hm),hydrogens two carbons removed from aromatic rings (H,). and remote aliphatic hydrogens (H?).

464

D. DUAYNE WHITEHURST, TAKAAKI ISODA, AND ISAO MOCHIDA a

100

4-MDBT

Effect of

0 Time (min)

20

40

60

Time (min)

FIG.36. Removal of byproduct H2S between stages increases HDS reactivity in sequential reactors (second-stage conditions: 360"C, 2.9 MPa, 30 min residence time). (0)Desulfurization performance in normally sequenced stages; (0)desulfurization performance when H2S is removed between stages. Zero time indicates product composition after the first stage at 360"C, 2.9 MPa, and 30 min residence time. Figure modified and reproduced with permission from Ref. 14. Copyright 1994 American Chemical Society.

conversion in the same way. Figure 36 shows the beneficial effect of removing H2S after partial conversion of 4-MDBT and 4,6-DMDBT.

4. Novel Reactor Designs Knowledge of the composition and reactivity of the different boiling ranges of gas oils and diesel fuels can provide guidance in the design of new reactor configurations that may be more effective for HDS processes. The composition of gas oil, described in Section IV and Table 111, showed that the different reactivity classes for HDS varied with boiling points. The most difficult sulfur compounds to desulfurize were found in the higher boiling range fractions (14,15).It has been shown that NiMo/A1203is more effective for desulfurization of the higher boiling fractions as it possesses a higher hydrogenation activity and provides an alternative pathway for HDS of the dialkyldibenzothiophenes that occur in the higher boiling range fractions. Figure 37 illustrates this observation (17). Thus, if the light and heavy boiling fractions are separated by distillation and processed separately, the overall process is more efficient. Figure 38 illustrates one potential schematic diagram of such a scheme (153). If it would be possible to combine distillation with HDS conversion in a single reactor, a very efficient

465

POLYAROMATIC SULFUR COMPOUNDS I

*--

c

I

-

CoMo

El NiMo

h

E

80

.-

g 8 3i

60

40 F-1

F-2

F-3

F-4

F-5

T-F

Fraction

FIG.37. HDS reactivities of different boiling fractions of gas oil. Reproduced with permission of Ref. 17.

-

~alllngrange i 220 370 'C SUlf", :1.02 wt % Fluorereence Intendty :83.

Heavier traction

Lighter fraction Boilmg range ; 220.3DO (45.6wt %I Sulfur ; 0.632 wt %

'c

-

Boiling range : 300 370 .C (54.4wt %) Sulfur i 1.325 wt %

p

1 33

FIG.38. Comparison between the conventional single-stage hydrodesulfurization process and co-curreut/countercurrent reaction system (153).

466 D. DUAYNE WHITEHURST, TAKAAKI ISODA, AND ISAO MOCHIDA process may result. For example, if the feed is introduced in the midpoint of the reactor and hydrogen is introduced at the bottom, then light fractions would be processed up-flow with co-current hydrogen and H2S produced in the lower portion of the reactor. The heavy fractions would be processed down-flow, countercurrent to the hydrogen, and the least-reactive sulfur compounds would experience the least amount of H2S inhibition. The design of such novel reactors could be of benefit to the specific problems associated with deep desulfurization of gas oils and diesel fuels. Recently, such a process (SynSat) has been announced as available for commercial application (154).

5. Alternative Approaches for Lowering Sulfur Levels If the specifications for diesel fuels become even more strict in the future and sulfur levels less than 0.05% S (500 ppm) must be achieved, it may be more reasonable to remove the sulfur by means other than hydrotreating. At the level of 500 ppm, selective stoichiometric reactions or selective sorption may be worth considering, even if the sorptions are irreversible. It should be recalled that benzene was for many years freed of thiophene by stoichiometric removal by Raney nickel, as pointed out by Startsev (2). There have been a few investigations of selective sorbents for trace amounts of aromatic sulfur compounds in gas oils. It has been claimed that certain active carbons exhibit the required selectivity and may be regenerated easily (152). More research in this area is encouraged.

VIII. Concluding Remarks The increasing demands for cleaner, more environmentally friendly fuels have prompted major interest in the fundamental understanding of one of the oldest refinery processes in use today, hydrodesulfurization. New standards are being set that are difficult to achieve with today’s catalysts and installed equipment. This has presented a challenge to researchers throughout the world to develop new technology in a field that is quite mature and has been continually improved over a period exceeding 72 years. This challenge is now being addressed in several ways. The traditional approach of improvement through incremental change is reaching a limit as to how much more can be done with conventional catalysts and process configurations. The urgency of the problem of meeting a 0.05% S specification in diesel fuels with no degradation in other fuel qualities, such as color or cetane number, has forced researchers to seek deeper understanding of the fundamental chemistry and structure of HDS catalysts, the mechanisms through which they operate, and the reaction pathways that the organic sulfur species undergo.

POLYAROMATIC SULFUR COMPOUNDS

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The solution seems to lie in the interdisciplinary cooperative research efforts of scientists and engineers having very different backgrounds. This particular problem has provided a focal point for diverse groups to share information and exchange ideas. In particular, organometallic chemists, familiar with the detailed structure and mechanisms of soluble metal complexes; surface scientists, familiar with the characterization of the exact nature of the surface species; and heterogeneous catalysis scientists and engineers, familiar with kinetics, mechanistic determinations on solids, and process optimization, are beginning to realize that the phenomenon observable with single metal or bimetallic soluble species is in fact identical with the phenomenon observable on solid catalysts composed of mixtures of metal sulfides bonded to the surface of the solid. Materials scientists, familiar with the synthesis and modification of carbons, metal oxides and sulfides, ceramics, and zeolites, can also be important partners in the development of novel catalysts having improved properties. As these disciplines merge closer and closer together, new understanding is evolving that is certain to lead to new and improved commercial catalysts. The opportunities for novel, innovative research in this area are numerous and the seriousness of the problem is such that efforts will be and should be devoted to providing the new processes necessary for preserving and/or improving the environment of the earth while allowing the standard of living of mankind to advance. REFEREN cEs

I . Topsoe, H., Clausen, B. S., and Massoth, F. E., in “Catalysis Science and Technology” (J. R. Anderson and M. Boudart, eds.), Vol. 11. Springer-Verlag, New York, 1996. 2. Startsev, A. N., Catul. Rev.-Sci. Eng. 37(3), 353 (1995). 3. Chianelli, R. R., Daage, M., and Ledoux, M. J., Adv. Cutal. 40, 177 (1994). 4. Sanchez-Delgado, R. A,, J. Mol. Catal. 86, 287 (1994). 5. Girgis, M. J., and Gates, B. C., Ind. Eng. Chern. Res. 30(9), 2021 (1991). 6. Harris, S., and Chianelli, R. R., in “Theoretical Aspects of Heterogeneous Catalysis” (J. B. Moffat, ed.), Catalyst Ser., p. 206. Van Nostrand-Reinhold, New York, 1990. 7. Pnns, R., de Beer, V. H. J., and Somorjai, G. A,, Catul. Rev.-&?. Eng. 31(12), 1 (1989). 8. Gates, B. C., Katzer, J. R., and Schuit, G. C. A,, “Chemistry of Catalytic Processes,” pp. 390-445. McGraw-Hill, New York, 1979. 9. Massoth, F. E., Adv. Catal. 27, 265 (1978). 10. Weisser, O., and Landa, S., “Sulfide Catalysts, Their Properties and Applications.” Pergamon, New York, 1973. 11. Topsoe, N.-Y., and Topsee, H., J. Cutul. 84,386 (1983). 12. Chemical Inspection and Testing Institute of Japan Report on “Advanced Refining Technologies for Gas Oil Distillate” (1989). 13. Kabe, T., Ishihara, A,, and Tajima, H., Ind. Eng. Chem. Res. 31, 1577 (1992). 14. Ma, X., Sakanishi, K., and Mochida, I., Ind. Eng. Chern. Res. 33, 218 (1994). 15. Ma, X., Sakanishi, K., Isoda, T., and Mochida, I., Fuel 76, 329 (1997). 16. Nishioka, M., Campbell, R. M., Lee, M. L., and Castle, R. N., Fuel 65, 270 (1986).

468 D. DUAYNE WHITEHURST, TAKAAKI ISODA, AND ISAO MOCHIDA 17. Ma, X., Sakanishi, K., Isoda, T., and Mochida, I., Prepr. Div. Pet. Chem., A m . Chem.

SOC. 39, 622 (1994). 18. Kabe, T., Ishihara, A,, Nomura, M., Itoh, T., and Qi, P., Chem. Lett., 2233 (1991). 19. Nagai, M., and Kabe, T., J. Catal. 81, 440 (1983). 20. Whitehurst, D. D., Mitchell, T. O., Farcasiu, M., and Dickert, J. J., Jr., “Exploratory Studies in Catalytic Coal Liquefaction,” Report on EPRI Project 779-18, Electr. Power Res. Inst., Palo Alto, CA, 1979. 21. Isoda, T., Ma, X., and Mochida, I., Prep. Div. Pet. Chem., Am. Chem. SOC.39,584 (1994). 22. Piel, W. J., Karsa, L. J., and Kesling, H. S., Jr., Prepr. Div. Pet. Chem., Am. Chem. Soc. 39, 520 (1994). 23. Vrinat, M. L., Appl. Caial. 6, 137 (1983). 24. Schuit, G. C. A,, and Gates, B. C., AZChE J. 19,417 (1973). 25. Givens, E. N., and Venuto, P. B., Prepr. Div. Pet. Chem., Am. Chem. Soc. 15(4), A183 (1970). 26. Houalla, M., Nag, N. K., Sapre, A. V., Broderick, D. H., and Gates, B. C., AZChE J. 24, 1015 (1978). 2 7. Houalla, M., Broderick, D. H., Sapre, A. V., Nag, N. K., de Beer, V. H. J., Gates, B. C., and Kwart, H., J. Caial. 61, 523 (1980). 28. Aubert, C., Durand, R., Geneste, P., and Mareau, C., J . Catal. 97, 169 (1986). 29. Geneste, P., Amblard, P., Bonnet, M., and Graffin, P., J. Catal. 61, 115 (1980). 30. Isoda, T., Nagoa, S., Korai, Y., and Mochida, I., Prepr. Div. Pet. Chem., A m . Chem. SOC.40, 559 (1996). 31. Isoda, T., Nagoa, S., Korai, Y., and Mochida, I., Prepr. Div. Pet. Chem., A m . Chem. SOC.40, 563 (1996). 32. Lamure-Meille, V., Schultz, E., Lemarie, M., and Vrinat, M., Appl. Catul. 131,143 (1995). 33. Landau, M. V., Berger, D., and Herskowitz, M., J. Catal. 158,236 (1996). 34. Nag, N. K., Sapre, A. V., Broderick, D. H., and Gates, B. C., J. Catal. 57, 509 (1979). 35. Whitehurst, D. D., Mitchell, T. O., and Farcasiu, M., “Coal Liquefaction: The Chemistry and Technology of Thermal Processes,’’ p. 165. Academic Press, New York, 1980. 36. Sapre, A. V., Broderick, D. H., Frankel, D., Gates, B. C., and Nag, N. K., AIChE J. 26, 690 (1980). 3 7. Sapre, A. V., and Gates, B. C., Ind. Eng. Chem. Process Des. Dev. 20, 68 (1981). 38. Ma, X., Sakanishi, K., Isoda, T., and Mochida, I., Energy Fuels 9, 33 (1995). 39. Whitehurst, D. D., Farcasiu, M., Mitchell, T. O., and Dickert, J. J., Jr., “The Nature and Origin of Asphaltenes in Processed Coals,” EPRI Ann. Rep. (March 1976-February 1977), AF-480, pp. 6-21 to 6-30. Elect. Power Res. Inst., Palo Alto, CA, 1977. 40. Whitehurst, D. D., Mitchell, T. O., and Farcasiu, M., “Coal Liquefaction: The Chemistry and Technology of Thermal Processes,” pp. 42-44. Academic Press, New York, 1980. 41. Farag, H., Whitehurst, D . D., Sakanishi, K., and Mochida, I., Prepr. Div. Pet. Chem., Am. Chem. Soc. (1997). 41u Isoda, T., Ma, X., and Mochida, I., Jpn. Pet. Inst. 37, 368 (1994). 42. Mochida, I., and Yoneda, Y., J. Catul. 11,183 (1968). 43. Farragher, A. L., and Cossee, P., Catal. Proc. Znt. Congr. 5th 1972 (1973). 44. Voorhoeve, R. J. H., and Stuiver, J. C. M., J. Cutal. 23, 228 (1971). 45. Ratnasamy, P., and Silvasanker, S., Catul. Rev.-Sci. Eng. 22, 401 (1980). 46. Chianelli, R. R., Cufal. Rev.-&% Eng. 26, 361 (1984). 4 7. Daage, M., and Chianelli, R. R., J. Caial. 149, 414 (1994). 48. Farben, I. G., Br. Pat. 315,439 (1928). 49. Byrns, A. C., Bradley, W. E., and Lee, M. E., Znd. Eng. Chem. 35, 1160 (1943). 50. Schuit, G. C. A,, and Gates, B. C., A l C h E J. 19, 417 (1973).

POLYAROMATIC SULFUR COMPOUNDS

469

51. Voorhoeve, R. J. H., J. Catal. 23,236 (1971). 52. Farragher, A. L., “Symposium on the Role of Solid State Chemistry in Catalysis,” New

Orleans Meeting (1977). 53. Clausen, B. S., Morup, S., Topsoe, H., and Candia, R., J. Phys. (Orsay, Fr.) 37,249 (1976). 54. Topsoe, H., Clausen, B., Candia, R., Wivel, C., and Morup, S., J. Catal. 68,433 (1981).

55. Candia, R., Sorenson, O., Villadsen, J., Topsoe, N., Clausen, B. S., and Topsoe, H., Bull. SOC. Chim. Belg. 93, 763 (1984). 56. Wivel, C., Clausen, B., Candia, R., Morup, S., and Topsoe, H., J. Cafal. 87,497 (1984). 57. Topsoe, H., and Clausen, B., Appl. Catal. 25, 273 (1986). 58. Topsoe, H., Clausen, B. S . , Candia, R., Weivel, C., and Morup, A., Bull. SOC. Chim. Belg. 90, 1189 (1981); J. Catal. 68, 433 (1981). 59. Topsoe, H., in “Surface Properties and Catalysis by Non-metals” (J. P. Bonnelle, B. Delmon, and E. Derouane, eds.), p. 329. Reidel, Dordrecht, The Netherlands, 1983. 60. Topsoe, H., Clausen, B. S., Topsoe, N., and Pedersen, E., Ind. Eng. Chem. Fundam. 25, 25 (1986). 61. van Veen, J. A. R., Gerkema, E., van der Kraan, A. M., Hendriks, P. A. J. M., and Beens, H., J. Catal. 133, 112 (1992). 62. Bouwens, S. M. A,, van Zon, F. B. M., van Dijk, M. P., van der Kraan, A. M., de Beer, V. H. J., van Veen, J. A. R., and Koningsberger, D. C., J. Cafal. 146, 375 (1994). 63. Louwers, S. P. A., and Prins, R., J. Catal. 133, 94 (1992). 64. Louwers, S. P. A,, and Prins, R., J. Catal. 139, 525 (1993). 65. Delmon, B., Bull. SOC. Chim. Belg. 88, 979 (1976). 66. Pirotte, D., Zabala, J. M., Grange, P., and Delmon, B., Bull. SOC. Chim. Belg. 90, 1239 (1981). 67. Smit, S. S., and Johnson, K . H., Prepr. Div. Pet. Chem., Am. Chem. SOC.39,532 (1994). 68. Burch, R., and Collins, A., Appl. Catal. 17, 273 (1985). 69. Polz, J., Zeilinger, H., Miiller, B., and Knozinger, H., J. Catal. 120, 22 (1989). 70. Kabe, T., Qian, W., Wang, G., and Ishihara, A., Prepr. Div. Pet. Chem., Am. Chern. SOC. 41,554 (1996). 71. Arnoldy, P., van den Heijkant, J. A. M., de Bok, G. D., and Mouljin, J. A., J. Catal. 92, 35 (1985). 72. Massoth, F. E., J. Catal. 30, 204 (1973). 73. de Beer, V. H. J., van Sint Fiet, T. M. H., van der Steen, G. H. A. M., Zawaga, A. C., and Schuit, G. C. A., J. Catal. 35, 297 (1974). 74. de Beer, V. H. J., van der Aalst, M. J. M., Machiels, C. J., and Schuit, G. C. A,, J. Catal. 43,73 (1976). 75. Angelici, R. J., Acc. Chem. Res. 21, 387 (1988). 76. Kim, S. I., and Woo, S . I., J. Cafal. 133, 124 (1992). 77. Wang, L., and Hall, K., J. Cacal. 77, 579 (1987). 78. Okamoto, Y., Maezawa, A,, and Imanaka, T., J. Catal. l 2 0 , 2 9 (1989). 79. van Veen, J. A. R., Gerkema, E., van der Kraan, A. M., and Knoester, A,, J. Chem. SOC., Chern. Commun., 1684 (1987). 80. Haag, W. O., and Whitehurst, D. D., Catal., Proc. In?. Congr., Sfh, 1972, p. 465 (1973). 81. Manassen, J., and Whitehurst, D. D., in “Catalysis: Progress in Research,” p. 177. Plenum, New York, 1973. 82. Whitehurst, D. D., CHEMTECH 10, 44 (1980). 83. Basset, J., Grubbs. R., Hegedus, L., Osborn, J., and Stille, J . K., in “International Workshop on Heterophase Attached Homogeneous Catalysts, Grenoble, Nov. 6-9” (1977). 84. Yermakov, Y . I., Kuznetsov, B. N., and Zakharov, V. A,, Stud. Surf: Sci. Catal. 8, (1981).

470 D. DUAYNE WHITEHURST, TAKAAKI ISODA, AND ISAO MOCHIDA 85. Topsoe, H., Clausen, B. S., Topsoe, N.-Y., Hyldoft, J., and Norskov, J. K., 39, 638 (1993). 86. Topsoe, H., Clausen, B. S., Topsoe, N.-Y., Hyldoft, J., Nbrskov, J., Ovesen, C. V., and Jacobsen, C. J. H., Bull. Soc. Chim. Belg. 104, 283 (1995). 87. Pecoraro, T. A,, and Chianelli, R. R., J. Catal. 67, 430 (1981). 88. Bulakov, N. N., and Startsev, A. N., Mendeleev Comm. 4, 173 (1993). 89. Chen, J., Daniels, L. M., and Angelici, R. J., J. Am. Chem. Soc. 112, 199 (1990). 90. Candlin, J. P., Taylor, K. A,, and Thompson, D. T., “Reactions of Transition Metal Complexes.” Elsevier, New York, 1968. 91. Angelici, R. J., Coord. Chem. Rev. 105, 61 (1990). 92. Rauchfuss, T. B., Prog. Inorg. Chem. 39,259 (1991). 93. Wiegang, B. C., and Friend, C. M., Chem. Rev. 92, 1 (1992). 94. Jones, W. D., and Feher, F. C., Acc. Chem. Res. 22,91 (1989). 95. Jones, W. D., and Dong, L., J. Am. Chem. Soc. 113,559 (1991). 96. Wiegang, B. C., and Friend, C. M., Chem. Rev. 92, 1 (1992). 97. Sauer, N. N., Markel, E. J., Schrader, G. L., and Angelici, R. J., J. Catal. 117,295 (1989). 98. Lopez, R., Peter, R., and Zdrazil, M., J. Catal. 73, 406 (1982). 99. Raiz, U., Curnow, O., and Curtis, M. D., J. Am. Chem. Soc. 116, 4357 (1994). 100. Jones, W. D., and Chin, R., J. A m . Chem. Soc. 114,9581 (1992), and references therein. 101. Adams, R. D., Pompeo, M. P., Wu, W., and Yamamoto, J. H., J. Am. Chem. Soc. 115, 8307 (1993). 102. Caswitt, C. J., and Rakowski Du Bois, M., J. Am. Chem. Soc. 108, 5482 (1986). 103. Leglise, J., van Gestel, J., and Duchet, J.-C., Prep. Div. Pet. Chem., Am. Chem. Soc. 39, 533 (1994). 104. Vanrysselberghe, V., and Froment, G. F., Ind. Eng. Chem. Res. 35,3311 (1996). 105. Wang, D., and Angelici, R. J., J. Am. Chem. Soc. 118, 935 (1996). 106. Simpson, H. D., Prep. Div. Pet. Chem., Am. Chem. SOC.39, 563 (1994). 107. Kasztelan, S., and Guillaume, D., Ind. Eng. Chem. Res. 33, 203 (1994). 108. Haag, W. O., and Whitehurst, D. D., US. Pat. 3,954,883 (1976). 109. Stanislaus, A,, and Cooper, B. H., Catal. Rev.-Sci. Eng. 36, 75 (1994). 110. Takatsuka, T., and Wada, Y., Shokubai 33, 306 (1991). 111. Takatsuka, T., Wada, Y., Suzuki, H., Komatsu, S., and Morimura, Y., J . Jpn. Pet. lnst. 35, 179 (1992). 112. Wada, Y., and Takatsuka, T., Petrotech (Tokyo) 13,418 (1990). 113. Tsutsumoto, N., Petrotech (Tokyo) 13,998 (1990). 114. Sakanishi, K., Ando, M., Abe, S., and Mochida, I., Sekiyu Gakkaishi 34, 553 (1991). 115. Sakanishi, K., Ando, M., and Mochida, I., Sekiyu Gakkaishi 35, 402 (1992). 116. Ma., X., Sakanishi, K., and Mochida, I., Fuel 73, 1667 (1994). 117. Ma., X., Sakanishi, K., Isoda, T., and Mochida, I., Ind. Eng. Chem. Res. 34,748 (1995). 118. Wakita, M., Yanazawa, K., and Matunaga, A,, Sekiyu Gakkaishi, 37, 561 (1994). 119. Ma., X., Sakanishi, K., Isoda, T., Nagao, S., and Mochida, I., Eng. Fuel 10, 91 (1996). 120. Lapinas, A. T., Klein, M. T., Gates, B. C., Macris, A,, and Lyons, J. E., Ind. Eng. Chem. Res. 26, 1026 (1987). 121. Langmuir, I., J. Am. Chem. Soc. 38, 2221 (1914). 122. Van Parjis, I. A,, and Froment, G. F., Ind. Eng. Chem. Prod. Res. Dev. 25,431 (1986). 123. Van Parjis, I. A,, and Froment, G. F., Ind. Eng. Chem. Prod. Res. Dev. 25,437 (1986). 124. Prather, J. W., Ahangar, A. M., Pitts, W. S., Henley, J. P., Tarrer, A. R., and Guin, J. A,, Znd. Eng. Chem. Process Des. Dev. 16, 267 (1977). 125. Satterfield, C. N., and Roberts, G. W., AIChE J. 14, 159 (1968). 126. Sapre, A. V., and Gates, B. C., Ind. Eng. Chem. Proc. Des. Dev. 21, 86 (1982).

POLYAROMATIC SULFUR COMPOUNDS 127. 128. 129. 130. 131. 132. 133. 134. 135. 136. 137. 138. 139.

140. 141. 142. 143. 144. 145. 146. 147. 148. 149. 150. 151.

152. 153. 154.

471

Broderick, D. H., and Gates, B. C., AIChE J. 27, 663 (1981). Nagai, M., and Kabe, T., J. Catal. 81, 440 (1983). Nagai, M., Sato, T., and Aiba, A., J. Card. 97, 52 (1986). Nagai, M., Masunaga, A,, and Hana-oka, N., J. Catal. 101, 284 (1986). La Vopa, V., and !;atterfield, C. N., J. Catal. 110, 375 (1988). Rollmann, L. D., .I. Catal. 46, 243 (1977). Isoda, T., Ma, X., and Mochida, I., J. Jpn. Pet. Inst. 37, 506 (1994). Isoda, T., Ma, X., Nagao, S., and Mochida, I., J. Jpn. Pet. lnst. 38, 25 (1995). Isoda, T., Nagao, S., Ma, X., Korai, Y., and Mochida, I., Energy Fuels 10, 482 (1996). Isoda, T., Nagao, S.,Ma, X., Korai, Y., and Mochida, I., Energy Fuels 10,487 (1996). Isoda, T., Nagao, S., Ma, X., Korai, Y., and Mochida, I., Energy Fuels 10,1078 (1996). Suchanek, A. J., Oil Gas J. May 7, p. 109 (1990). Naber, J. E., and Stork, W. H. J., in “Proceedings of the First Japan-EC Joint Workshop on the Frontiers of Catalytic Science and Technology for Alternative Energy and Global Environmental Protection, Tokyo” (1991). Peries, J. P., Billon, A,, Hennico, A., and Kressmann, S., NPRA 89th Ann. Meet., San Antonio, TX, March 17-19 (1991). Cooper, B. H., Stanislaus, A., and Hannerup, P. N., Prepr. Pup.-Am. Chem. SOC., Div. Fuel Chem 3’1,41 (1992). Sogaard-Anderson, P., Cooper, B. H., and Hannerup, P. N., NPRA 90th Annu. Meet., New Orleans, LA. March, AM-92-50 (1992). Cooper, B. H., Sogaard-Anderson, P., and Hannerup, P. N., AIChE Spring Natl. Meet., Houston, TX (1993). Lacroix, M. J., Boutarfa, M., Guisse, M., J. Guillard, C., Vrinat, M., and Breysse, M., J. Catal. 120, 473 (1989). Geantet, C., Gobi-dos, S., De Los Reyes, J. A., Cattenot, M., Vrinat, M., and Beyesse, M., Catal. Today .LO, 665 (1991). Breysse, M., Alfonso, J., Lacroix, M., Portefaix, J. L., and Vrinat, M., Bull. SOC.Chim. Belg. 108, (1991). Gobolos, S., Lacroix, M., Decamp, T., Vrinat, M., and Breysse, M., Bull. SOC. Chim. Belg. 108, 907 (1991). Koussathana, M., Vamvouka, D., Economou, H., and Verykios, X., Appl. Catal. 7, 283 (1991). Yitzhaki, D., Landau, M. V., Berger, D., and Herskowitz, M., Appl. Catal. A 122, 99 (1995). Yumoto, M., Usiui, K., Watanabe, K., Idei, K., and Yamazaki, H., Catal. Today 35, 45 (1995). Fujieda, S., Morinaga, K., and Kondo, A,, Petrotech (Tokyo) 18, 296 (1995). Savage, D. W., and Kaul, B. K., U.S. Pat. 5,454,933 (1995). Mochida, I., Sakanishi, K., Ma, X., Nagao, S., and Isoda, T., Catal. Today 29,185 (1997). Sucharek, A,, Prep. Div. Petr. Chem., Am. Chem. SOC.41,583 (1996).

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ADVANCES IN CATALYSIS, VOLUME 42

MuItiphase Homogeneous Catalysis EIIRGIT DRIESSEN-HOLSCHER Insrrtut fur Teciinische Chemre icnd Petrolchemie Rheinlsch-Westfrilrsche Technrsche Hochschule Aachen 52056 Aachen, Germany

1.

Introduction

Catalysis in liquid-liquid biphasic systems has developed recently into a subject of great practical interest because it provides an attractive solution to the problems of separation of catalysts from products and of catalyst recycle in homogeneous transition metal complex catalysis. Two-phase systems consist of two immiscible solvents, e.g., an aqueous phase or another polar phase containing the catalyst and an organic phase containing the products. The reaction is homogeneous, and the recovery of the catalyst is facilitated by simple phase separation. With biphasic catalysis, the usual advantages of homogeneous over heterogeneous catalysis can be achieved: mild reaction conditions high activity arid selectivity modulation of Ihe coordination sphere of the transition metal by variation of central atoms or ligands for specific applications and special steric or electronic properties Furthermore, one of the advantages with heterogeneous catalysis can also be achieved: economical separation of products from catalysts Abbreviations: acac, ocetylacetonate; amphos, class of diphenylalkylphosphanes,with quaternary amines; (S,S)-BDPP, 2,4-bis(diphenylphosphino)pentane; BINAS-Na, bis[[(disulfonatophenyl)phosphino]methyl]tetrasulfonatobinaphthenesodium salt; BPPM, 4-(diphenylphosphino)-2-[(diphenylphosphino)methy~Jpyrrolidine-l-carboxylicacid 1,l-dimethylethyl ester; COD, 1,5-cyclooctadiene;DMSO, dimethyl sulfoxide; dppb, 1,4-bis(dipheny1phosphino)butane; dppbts, 1,4-bis(diphenylphosphino)butane rn-fetrasulfonated; edta, ethylenediaminetetraacetate; en, ethylenediamine; HMF, 5-(hydroxyrnethyl)-2-furfural; PTA, 1,3,5-triaza-7-phosphaatlamantane; SDS, sodium dodecyl sulfate; tppds, friphenylphosphane m-disulfonated; tppms, niphenylphosphane m-monosulfonated; tppts, friphenylphosphane m-trisulfonated. 473 Copynght C 1998 bv Academc Press All rights of reproduction in any form reserved 0360 0564198 $25 00

474

BIRGIT DRIESSEN-HOLSCHER

The present article deals with several aspects of multiphase catalysis. Since 1991, one new review per year has been published in this field (1-8) and others appeared earlier (9-13). This review deals mainly with recent developments and summarizes briefly the literature published before 1994. The status and the potential of this kind of catalysis for new applications are discussed.

It. The Principle of Multiphase Catalysis The general principle of two-phase catalysis in polar solvents, for example, in water, is shown in the simplified diagram of Fig. 1.The metal complex catalyst, which can be solubilized by hydrophilic ligands, converts the reactants A + B into the product C. The product is more soluble in the second than in the first phase and can be separated from the catalyst medium by simple phase separation. Excellent mixing and contacting of the two phases are necessary for efficient catalytic reaction, and thus the reactor is normally well stirred. This is the simplest case of two-phase catalysis because the solubilities of the product C and of the catalyst are so different that a nearly perfect phase separation results. This kind of process would be ideal for industrial Reaction: A

I

S I t- catalyst

catalyst

+

B -C

S I1 recvcle

I

0I catalyst

1

C product

S I + catalyst recycle

two-phase reactor

phase separator

distillation

FIG.1. Schematic representation of process illustrating biphasic (liquid-liquid) catalysis. Either phase may be the continuous phase in the reactor, which is usually stirred to maintain good contact between the two phases.

475

MLJLTIPHASE HOMOGENEOUS CATALYSIS

application. The catalyst phase is directly recyclable, and the purification of the products is normally by simple distillation. If the product C is partly dissolved in the catalyst phase S I, an extraction is necessary (Fig. 2a). In this case the extraction is performed with a phase S I1 (e.g., petroleurn distillate) which is almost completely insoluble in the catalyst phase S I (e.g., water). Reaction:

a

A +B-C

r-

S I + catalyst ---I P.

t

B

S I

-A

t

catalyst

+

catalyst + A, B

d

I

C

recycle S I + catalyst

reactor

extraction

distillation

Reaction: A + B - C

b

A+ B

S + catalysi

-

c+s

C

+catalyst /

-

productC

S

+

d

catalyst

reactor

phase separator

FIG.2. (a) Simplified flow diagram for two-phase catalysis with product(s) C that is (are) soluble in the catalyst medium S I. (b) Simplified flow diagram for two-phase catalysis whereby the second phase is formed [product(s)] during the process.

476

BIRGIT DRIESSEN-HOLSCHER

It is also possible that the second phase may form during the reaction, because the product(s) C is (are) not soluble in the original catalyst phase S (Fig. 2b). A phase separation enables continuous return of the catalyst to the reactor.

111.

Aspects of Mass Transfer in Multiphase Catalysis

It is important to study chemical kinetics of catalysis in dispersed liquidliquid systems. Therefore it is necessary to consider the following (14-16): which phase is continuous and which is dispersed the interfacial area and mass transfer coefficients characterizing the transport between phases the reaction regime-whether the chemical kinetics are fast or slow relative to mass transfer Although the chemical engineering principles are well established, there has been little characterization of the transport-reaction processes in biphasic catalysis. Research is needed.

IV. Reactions in Water as the Catalyst Phase

A.

CLASSES OF WATER-SOLUBLE LIGANDS AND METALCOMPLEXES

1. Anionic Ligands By the introduction of polar anionic groups, including sulfonate, carboxylate, and phosphonate groups, ligands and metal complexes incorporating these ligands can be converted into water-soluble derivatives. The functional groups that form anions in water and render the compounds water soluble are presented in Scheme 1. a. Ligands with Sulfonate Groups. Sulfonated phosphanes are the most widely used ligands in water-soluble metal complexes and two-phase catalysis. The sodium salts of these compounds can be obtained by the sulfonation of phosphanes with oleum and subsequent neutralization

Lig-COO-M' Lig-P( O)-(O-M'),

SCHEME 1. Functional groups in water-soluble anionic ligands (M'

=

Cation).

MlJLTIPHASE HOMOGENEOUS CATALYSIS

477

with NaOH ( I 7). These classical direct sulfonation techniques have some problems because of the oxidation potential of oleum and because a mixture of mono-., di-, and trisulfonated compounds may result. An appreciable amount of the phosphane is converted to the phosphane oxide, which substantially complicates the purification of the desired phosphane. The direct sulfonation with sulfuric acid discovered by Hanson et al. ( I 8) is more effective. Another method by which the formation of phosphane oxide is strongly inhibited was published in 1995; the use of B(OH)3 in combination with sulfur trioxide hinders the oxidation and offers the possibility to control the number and the position of sulfonate groups ( I 9). The introduction of a sulfonate group by reaction with oleum is not limited to arylphosphanes. Tris(o-phenylalkyl)phosphanes, P[(CH2), (C6H5)]3(n = 1, Z!, 3, and 6), can be sulfonated in the para position and to a lesser extent in the meta position (18). The technique of sulfonating water-insoluble ligand precursors can be applied to bidentate, polydentate, and chiral phosplianes (20-24); the compounds 1-3 are presented in Scheme 2 as examples. Ammonium salts of sulfonated phenyl phosphites, which are surprisingly resistant to hydrolysis, have been prepared (25).The compounds are more water stable than triphenyl phosphite, but they decompose completely within 24 h. Other routes to sulfonated phosphanes are the nucleophilic substitution of phosphides [Eq. (l)], which was first published in 1995 (26), and the reaction of haloge-natedarenesulfonates with PH derivatives (27) [Eq. (2)]. S03Na I

PPh2 S03K

I

478

BIRGIT DRIESSEN-HOLSCHER

CH, -PAr,-,P h,

CH2-PArz,Ph,

Na0,S

1

bisbis

2

norbos

R = p-C6H4-S03Na SCHEME2. Examples for polydentate and chiral sulfonated phosphanes.

The latest developments to produce sulfonated, water-soluble phosphanes are characterized by improvements of direct functionalizations to build up tailor-made molecules. b. Ligunds with Curboxylute Groups. Phosphanes with one carboxylic group have been known since 1952 (28,29) but their water solubility is too

MULTIPHASE HOMOGENEOUS CATALYSIS

479

small for them to be useful for catalysis in water. For this purpose, ligands with at least two carboxylate groups are necessary. The phosphane analogue of ethylenediaminetetraacetic acid is obtained as an air-stable monohydrate of the tetrasodium salt (30). The phosphorylated derivative of maleic acid 4 was synthesized in 1993 (31).

0

Water-soluble carboxylates of phosphanorbornadienes can be obtained by [4 + 21 cycloaddition of 3,4-dimethylphosphole to maleic anhydride, subsequent rearrangement, and alkaline hydrolysis (32). Alkali metal ph osphides serve as useful intermediates in the synthesis of new water-soluble phosphanes in reactions with appropriate organic halides. A general route to carboxyphenylphosphanes has been reported (33). Reaction of LiPPh2 with itaconic acid yielded a monodentate ligand with two carboxylate groups, Ph2PCH2CH(COOLi)(CH2COOLi)(34). The binding of functionalized chiral ligands to water-soluble polymers has also been shown; a (dipheny1phosphino)pyrrolidine derivative reacts with poly(acry1ic acid) to form a macroligand that is useful in biphasic reduction (35). c. Ligands with Phosphonate Groups. Ligands containing phosphonate groups were first reported in 1992. Roundhill et al. (36, 37) described the preparation of [(diphenylphosphino)alkyl]phosphonates. The highly water-soluble triphenylphosphine monophosphonate represents the first example of a compound containing an aryl ring which is substituted by both phosphine and anionic phosphonate moieties. The reaction sequence is shown in Scheme 3 (38). 2. Cationic Ligaizds

Phosphanes with quaternized aminoalkyl and aminoaryl groups or with phosphonium groups can also provide water solubility. Before the nitrogen atom can be alkylated, the phosphorus center has to be protected by oxida-

480

BIRGIT DRIESSEN-HOLSCHER

SCHEME3. Synthesis of phosphonatophosphanes.

tion or by coordination to a metal. The last step of this route is a reduction or decomplexation which leads to the desired phosphane ligand (Scheme 4). The amphos ligand was first synthesized by Baird et al., who took advantage of the synthesis principle presented in Scheme 4 (39).Several groups used this method to synthesize other phosphanes (40). The reaction of PH3 with Me2N(CH&C1 in superbasic media (KOH/ DMSO) was investigated. The tertiary phosphane, upon protonation, affords the extremely water-soluble ligand {P[(CH2)2NHMe2]3}C13 (solubility: 1450 g/1000 ml of water at room temperature) (41). Cationic guanidinophosphanes (42) and tertiary arylphosphanes containing one or two benzylic amino groups have also been synthesized (43). (Phosphinoalky1)phosphonium salts were first prepared through monoquaternization of bisphosphanes in 1991 (44). The synthesis proceeds according to Eqs. (3) and (4).

Me3tP*PPh2

X-

(4)

wNMe2 2. CH3' PhzP -NMe3'1

PhzP

3. HSiC13

SCHEME4. Synthesis of (aminoalky1)phosphanes made water soluble by quaternization of the N atom.

MULTIPHASE HOMOGENEOUS CATALYSIS

481

Ligands with triimethylammonium groups can be prepared by the regioselective addition of diphenylphosphane to conjugated alkenes with good yields. For example, the alkene CH2= CHC(0)OCH2CH2NMe3+I-,an acrylic ester, has been used (45).The authors did not report the water solubility. The growing number of publications in recent years underlines the rapidly increasing interest in this class of compounds.

3 . Neutral Ligantls

The class of neutral ligands includes phosphanes with hydroxyalkyl and polyether substituents and sugar derivatives of phosphanes. Phosphanes with polyether substituents have noticeable water solubility if the degree of polycondensation n is at least 16 up to 32. They are of interest as recoverable, nonionic ligands. Several examples of this class of compound have been synthesized (structures 5-7) (46). However, the synthetic methodology is limited to the reactions between phosphorus compounds and polyether derivatives.

An alternative approach was reported in 1997. A series of novel watersoluble polyether-substituted triphenylphosphanes was prepared by means of the ethoxylatiori of mono-, bis-, and tris(p-hydroxypheny1)phosphanes (47). They manifest inverse temperature-dependent solubility in water that enables them to act as thermoregulated phase-transfer ligands. Crown-ether-substituted ligands were first investigated by Okano et al. (structure 8), who prepared arylphosphanes with these groups (48).

482

BIRGIT DRIESSEN-HOLSCHER

PPh2

Hydroxyalkyl-substituted phosphanes were first described by Chatt et al. in 1973. To date, only relatively few water-soluble alkylphosphorus ligands have been developed. The acid-catalyzed addition of PH3to formaldehyde was a large-scale industrial process to produce the flame-retardant salt of the phosphonium ion [P(CH,OH),]+. In 1958, it was reported in a patent (49) that tris(hydroxymethy1)phosphane was the product when the same reaction was carried out by using a PtC12 catalyst instead of a protic acid. Tris( hydroxymethy1)phosphane is being tested in the form of metal complexes as a catalyst for the addition of PH3 to formaldehyde (50). Tris(2hydroxyethy1)phosphane has been known since 1979, but the chemistry has not been developed until recently (52). Higher tris(hydroxya1kyl)phosphanes have not been investigated much; they are known only from the patent literature. They can be synthesized by the radical addition of PH3 to functionalized alkene, e.g., to ally1 alcohol, and they are soluble in water and in alcohols (52).The coordination chemistry of these basic ligands with several transition metals was investigated by Pringle et al. (53). Chelating alkylphosphanes that are water soluble by virtue of having hydroxy end groups were reported recently. The catalytic formylation of H2PC6H4PH2and of H2PCH2CH2PH2in the presence of formaldehyde in aqueous media yields hydroxymethyl bistphosphanes) in near-quantitative yields (54). Hydroxypropyl-substituted compounds were prepared by radical addition of allylic substrates to 1,2-diphosphinoethane in methanol (55). A new triphosphane with four hydroxymethyl end groups was synthesized by the formylation of PhP(CH2CH2PH& (56). The compound is water soluble and remarkably resistant to oxidation in aqueous media. Sugar derivatives of phosphanes can be made by reaction of ethylene glycol derivatives with sugar diacetonides to form monoallyl ethers, which can be brominated at the double bond. These compounds react with phenyland diphenylphosphanes to yield ligands with a sugar function in the side chain (57). Phosphinated glucopyranosides were prepared by Oehme and

MIJLTIPHASE HOMOGENEOUS CATALYSIS

483

Selke (58) for hydrogenation reactions with rhodium complexes containing these ligands. 1,3,5-Triaza-7-phosphaadamantane (9), also a nonionic ligand, was prepared from the reaction of formaldehyde, tris(hydroxymethyl)phosphane, and hexamethylenetetramine (59).

4. Metal Salts as 'Water-Soluble Catalysts Saltlike metal catalysts without hydrophilic phosphane ligands can be used for reactions in water. For example, aqueous RuC& and Pd(edta) are water soluble, and many other metal complexes that coordinate water as a ligand are soluble in water (60). These metal complexes have one important disadvantage when they are used as catalysts in water. It is the problem of leaching, which means that the catalyst can be extracted from the aqueous phase into the organic or product phase. The hydrophilicity of a metal salt or an ionic complex is not high, and so polar products may coordinate and transport them into the second phase. This topic is not considered in this review; a summary was reported by Kalck and Monteil (2). A recent book about this subject was written by Martell and Hancock (61).

B. CARBONYLATION REACTIONS IN WATER 1. Hydroformylation The 0x0 synthesis (hydroformylation), which was discovered by Otto Roelen in 1938 at Ruhrchemie AG, is one of the most important metalcomplex-catalyzed reactions of alkenes. Today, the worldwide capacity for 0x0 products is about 6 million tondyear. An important development in the past 15 years in hydroformylation technology was the introduction of biphasic homogeneous catalysis. Kuntz (62) expressed the basic idea of a new generation of water-soluble 0x0 catalysts with triphenylphosphane trisulfonate (tppts as the sodium salt) as a ligand for a rhodium-complex-catalyzed hydroformylation process. Ruhrchemie AG adapted the idea on the basis of research done at Rh6nePoulenc and developed it into an industrially viable process, which was

484

BIRGIT DRIESSEN-HOLSCHER

Ar-rn-CHO ArH

9

HZICO H20

\ C3H6

SCHEME5. Deactivation mechanism of the Rh(1) catalyst in the Ruhrchemie/RhAnePoulenc process.

-

recently licensed (62). The hydrophilic 0x0 catalyst [HRh(CO)(tppts),] is so far the ideal metal complex catalyst for biphasic hydroformylation of propene [Eq. ( 5 ) ] in industry (Section VI).

[cat.]

+H21Co

H20

‘ 0

4

0

4.

Although the standard Rh(I)/tppts catalyst has extremely good long-term stability, a slight loss of activity occurs after several years’ use. The deactivation mechanism of the Ruhrchemie/Rh6ne-Poulenc catalyst has been clarified in detail (63) and is presented in Scheme 5.

MULTIPHASE HOMOGENEOUS CATALYSIS

485

The development of ligands for biphasic 0x0 processes is ongoing; some selected examples of recent publications are cited here. A detailed review of hydroformylation was presented by Beller et al. (64). A major improvement in both catalytic activity and selectivity of propene hydroformylation in the organic-aqueous biphasic system was achieved by using the newly s#ynthesizedBINAS-Na ligand 10 in combination with rhodium(II1) acetate (65). S03Na

I

S03Na

10 BINAS-Na Ar = CsH4-3-SO3Na

The data obtained with this new ligand represent a 1Zfold increase in activity compared with what is obtained with tppts, with a further increase in selectivity to the n-aldehydes. The use of water-soluble catalysts in the biphasic hydroformylation of higher alkenes is limited by the solubility of these reactants in water. Consequently, the reactivities of higher alkene homologues are at least an order of magnitude less than that of propene. The use of surface-active agents (66), or ligands which have surface-active properties themselves (67),may enhance solubility but will have a negative influence on the efficiency of separation, which can result in rhodium loss. In some cases, water can also interfere with catalysis and can strongly influence the catalytic reaction (68). Another approach to solubilize the higher alkenes in the aqueous catalyst phase is to use cyclodextrins for a better mass transfer between the aqueous and organic phases (69). In the hydroformylation of l-decene, cyclodextrins shomed good results. Cyclodextrins increase the solubility of lipophilic reactants in the aqueous phase. The future will show whether

486

BIRGIT DRIESSEN-HOLSCHER

the use of a further component (cyclodextrin) in biphasic catalysis is still acceptable for industrial application. Reetz et al. synthesized phosphane ligands having the cyclodextrin group in the side chain of the ligand (11) (70). These ligands were tested in the hydroformylation of higher alkenes, e.g., 1-octene. The cyclodextrin group in the ligand enables an easier mass transport of the lipophilic octene into the aqueous catalyst phase. -S, , , - , , , Ph2P

p-cyciodextrine

Rhodium catalysts for the hydroformylation of n-tetradecene-1 were synthesized that are soluble in methanol as well as in water. After reaction took place in methanol solution, the product mixture was separated by the addition of water; the aqueous phase contained the catalyst (71). Several runs with the same catalyst showed that the conversion decreased, indicating losses of rhodium. Recently, a new rhodium recycling system was described that takes advantage of amphiphilic ligands such as PhzArP (Ar = 3-hydroxyphenyl, 4-carboxyphenyl). The corresponding rhodium complexes are active in the hydroformylation of 1-octene and can be separated from the products by acidic or basic extraction into water. After neutralization of the aqueous phase, the rhodium species could be extracted into a new batch of octene, with toluene as a solvent. The recovered catalyst retained only up to 87% of its activity (72). Attempts have been made to apply the oxoreaction for the synthesis of fine chemicals. Hydroformylation of methyl acrylate to give an a-branched aldehyde is useful in providing chiral building blocks after enzymatic hydrogenation [Eq. (6)].

aim product

target product

MIJLTIPHASE HOMOGENEOUS CATALYSIS

487

Two groups recently investigated this reaction by using rhodium catalysts and water-soluble phosphanes. When prepared from [R h(aca~)(CO)~] tppms was used a$.a ligand in water-toluene mixtures, the rhodium was transferred into the organic phase (73). The hydroformylation of methyl acrylate was faster and more selective under similar conditions in biphasic systems when the ligands were PPh3, tppts, dppb, and its tetrasulfonate dppbts. The alp ratio of aldehydes was always greater than 100, and in a few cases greater than 200 (74). Homologous o-alkenecarboxylic acid methyl esters were hydroformylated with the water-soluble rhodiumltppts catalyst in an aqueous-organic two-phase system (75). By addition of surfactants a miceliar system was created. Cationic surfactants are better than anionic or nonionic surfactants for this hydroformylation. Another attempt to apply the 0x0 reaction to the synthesis of fine chemicals was made by the hydroformylation of styrene. The chiral catalyst, a rhodium complex of a surfactant phosphane 12, gave no optical induction (76). This result agrees with the known poor enantioselection ability of other complexes with monodentate chiral phosphanes.

2. Carbonylation Reactions The transition metal-complex-catalyzed carbonylation of benzyl halides to yield phenylacetic acids is an extensively studied reaction [Eq. (7)].

Two-phase systems are applied, with the catalyst and the reactants in the organic solvent; an alkaline aqueous phase is used to dissolve the product as the sodium salt. An excess of base is necessary to neutralize the acid formed during the reaction and to deprotonate the phenyiacetic acid for dissolution in the aqueous phase. Although this process has significant disadvantages-high catalyst loading, the need for phase transfer agents,

488

BIRGIT DRIESSEN-HOLSCHER

metal loss, and a limitation to benzyl bromides as substrates-Montedison is said to practice it (77). The use of water-soluble catalysts in this reaction has hardly been investigated. Ruthenium/edta (78) and cobalt/tppts (79) catalysts have been described. The use of palladium/tppms catalyst was also reported (80). When edta and tppms are used as ligands, leaching of the metal by the product stream takes place. In the case of the cobalt/tppts catalyst, a high CO partial pressure and a catalyst concentration of >8 mol% are necessary. The reason for this effect is not clear. The carbonylation of bromobenzene with palladium/tppts complexes was reported by Monteil and Kalck (81). Some of the aforementioned disadvantages were alleviated in a new process for carbonylation of substituted benzyl chlorides (82). The reaction was carried out in a two-phase system in the presence of CO at atmospheric pressure; yields of phenylacetic acids of 80-94% were reported. The palladium catalyst contains tppts or BINAS-Na, 10, to allow water solubility. The maximum turnover frequency was found to be 135 h-l, and the lifetime of the catalyst increased as a result of continuous addition of reactants. The carbonylation of renewable 5(hydroxymethyl)-2-furfural (HMF) was reported to take place under aqueous-phase catalytic conditions with [Pd(tppt~)~] as a catalyst [Eq. (S)] (83). HOJ

+ c O [cat] HOOC H*/Hz0 50°C

Selective carbonylation of HMF was observed to yield 5-formylfuran-2acetic acid as the sole carbonylation product; the only byproduct was 5-methyl-2-furfural (MF). The activity and selectivity were both found to be strongly influenced by the tppts/Pd ratio; the maximum efficiency was observed for tppts/Pd = 6. Replacement of tppts by ligands containing fewer sulfonate groups, e.g., tppds or tppms, led to a dramatic drop in the catalytic activity. Furthermore, it was found that the selective carbonylation of benzyl alcohol to phenylacetic acid also took place in the presence of the catalyst (84). C. HYDROGENATIONS

Hydrogenation was reviewed recently by Chaloner et af. (85).The history of aqueous metal complex catalysis started with the hydrogenation reaction. Recent work has focused on the search for mechanistic clues to understand the changes in reaction rates and selectivities of aqueous hydrogenation reactions compared with the nonaqueous analogues.

MULTIPHASE HOMOGENEOUS CATALYSIS

489

1. Hydrogenations with Achiral Reactants

The ruthenium cluster [Ru2(7f-C6H6)H6]CI2is a catalyst for fumaric acid hydrogenation in aqueous solutions, with a turnover frequency of 35 h-' at 50°C (86). The water-soluble Ru(I1) complex [Ru(.r"-C6H6)(CH3CN),](BF4)2 Catalyzed the biphasic hydrogenation of alkenes and ketones with retention of the catalyst in the aqueous phase (87).However, the ruthenium complex moved to the organic phase when benzaldehyde was hydrogenated. In a benzene-D20 system, H-D exchange was observed between H2 and D20. Both monohydridic pathway and a dihydridic pathway are possible for hydrogen activation, and these two different catalytic cycles influence the yield and product distribution. Monohydrides play an important role in the following rhodium-complexcatalyzed hydrogenations in aqueous media. The catalyst precursor is [RhCl(PTA),], which gives the catalytically active species (HRh(PTA),] formed by dehydrochlorination of the primary product of H2 oxidative addition (88). The complex is an active catalyst for several reactants, e.g., olefinic and 0x0 acids, ally1 alcohol, and sulfostyrene. The hydrogen transfer reaction from aqueous formate to unsaturated aldehydes is also catalyzed by [RhCl(PTA),] (89). The selectivity for the reduction of the C-C bond is high, and the catalyst can be recycled. These (90), results are in contrast to those observed with [ R u C ~ ~ ( P T Acatalyst )~] showing high selectivities for the reduction of the C = O bond. The selective hydrogenation of a#-unsaturated aldehydes to give the corresponding unsaturated alcohols [Eq. (9)] was investigated with the ruthenium complex catalysts, initially present as [Ru(H)(Cl)(tppts),] or [Ru(H)2(tppts)41 ( 9 0

When unsaturated ketones are used as reactants, the C=C bond is preferentially reduced. Most of the complexes are transformed and thus deactivated after their first catalytic run. The phosphane (tppts) of the complexes underwent reactions with the organic products, giving phosphonium salts, which are responsible for the deactivation. The analysis of the aqueous phases shows that the recycling would be difficult in many cases. Complexes of ruthenium, [HRu(CO)Cl(tppms),] .2H20 and [HRu(CO) Cl(tppt~)~], were reported to be catalysts for the same hydrogenation reaction (92). The metal complexes were not pure; rather, they were used in the presence of the free sulfonated phosphanes and their respective oxides.

490

BIRGIT DRIESSEN-HOLSCHER

The hydrogenation of cinnamaldehyde was investigated with the first osmium/tppms complexes that are water soluble, [ O ~ H ~ ( t p p r n s )[OsHCl ~], ( C O ) ( t p p m ~ ) ~and ] , [ O ~ C l ( t p p m s ) ~ ( p - C(93). l ) ] ~ There is a clear advantage in using the aqueous biphasic systems over their homogeneous PPh3 analogues because product separation and catalyst recycling are efficient. The selectivity for formation of the unsaturated alcohol is considerably greater for the biphasic system than for the single-phase system. The complex [ n ~ e r - I r H ~ C l ( P M ewas ~ ) ~used ] as a catalyst for the hydrogenation of alkynes and alkenes in water, and water-soluble ethylenediamine (en) complexes of iridium, [Ir(COD)(en)]Cl, were found to be excellent catalysts for aqueous hydrogenations (94). It would be interesting to determine the loss of iridium during application of these complexes in biphasic catalysis. The reduction of aromatic nitro compounds to the corresponding amines was catalyzed by [ R U ~ ( C O )in~ combination ~] with aliphatic amine cocatalysts (95). A mixture of diglyme and water was used as a solvent, turnover frequencies were about 5000 h-l, and a C O partial pressure of 20-50 atm was necessary. The reaction is highly selective for aromatic amines. It was speculated that the reaction proceeds via an intramolecular hydrogen transfer in a hydrido-metal-nitrene intermediate without prior formation of H2 in the water gas shift reaction. 2. Asymmetric Hydrogenations The surface-active diphosphane 12 was applied in the hydrogenation of methyl a-acetamidocinnamate [Eq. (lo)] with [RhCl(COD)]2as the catalyst precursor in homogeneous methanolic solution and, alternatively, in ethyl acetate-water biphasic systems (96).

In the one-phase reaction, complete conversion and ee values of about 72% were reached. In the biphasic system, the rhodium complex of the surfactant ligand 12 showed considerably higher activity than in the onephase system, while retaining enantioselectivity (96). These results agree with results of earlier work that micelle-forming ligands enhance the solubility of lipophilic reactants in water. Enantioselective hydrogenations of dehydroamino acid derivatives are also catalyzed by rhodium complexes of phosphinated glucopyranosides (97). [Rh(Me-a-glup-OH)(COD)]BF4 and [Rh(Ph-P-glup-OH)( COD)]BF4

MULTIPHASE HOMOGENEOUS CATALYSIS

491

were used in aquea'us systems in the presence of sodium dodecyl sulfate (SDS) at room temperature and of H2at atmospheric pressure. The addition of SDS increased bath the rate and the enantioselectivity of the hydrogenations, e.g., a shortening of reaction half-time from 390 to 6 min and an increase of the ee value from 83 to 97%. These changes are attributed to micellar effects. It was found that high enantioselectivity was favored by micelle-forming agents which are known to contain less water in their micellar cores. Recently, carbohydrate amphiphiles have been tested in the asymmetric hydrogenation of (;?)-methyl a-acetamidocinnamate in water (98). With a rhodium(1)-BPPM complex, 50% of the reactant was converted in 5 min, and enantioselectivities up to 96% were observed. A comparison of amphiphiles with alkyl chains of different lengths showed that micelle-forming properties, hydrophilic-lipophilic balance, and the structure caused by hydrogen bonding in the head group may be responsible for these effects. Asymmetric aqueous hydrogenolysis of sodium cis-epoxysuccinate to give sodium hydroxysuccinate was catalyzed by rhodium(1) complexes of sulfonated (S,S)-BDPP (99).Deuterium-labeling experiments showed that both hydrogen and water participate as reactants in the aqueous hydrogenolysis, and the reaction proceeds via the direct C-0 bond cleavage of the epoxy group. Rhodaoxetane-BDPP species were identified as catalytic intermediates in high-pressure NMR experiments. The highest ee (40%) was achieved by tlne use of nonsulfonated BDPP; the enantioselectivity decreased slightly as the number of sulfonate groups on the ligand increased. D. TELOMERIZATIONS

Some homogeneously catalyzed telomerizations, i.e., dimerizations of dienes coupled with the addition of a nucleophile [Eq. (ll)], have been carried out in twophase systems. One example has found industrial application, the synthesis of 1,7-octadienol from butadiene and water (Section VI).

/ M

4-

Nu-H

[cat.] ___Ic

+

Nu

(11)

The telomerization of dienes in a two-phase system was first described in a patent (100). Water was used as the solvent for the catalyst, with sulfonated phosphane ligands providing the water solubility. Water, alcohols, phenols, acids, amines, and acetylacetic acid were used as nucleophiles.

492

BIRGIT DRIESSEN-HOLSCHER

Water-soluble quaternary ammonium phosphanes have been used as ligands for palladium in the telomerization with methanol under two-phase conditions (101). The telomerization of sucrose with butadiene was catalyzed in aqueous solution by palladium acetate and tppts (102). The sucrose conversion was about 96%, but octadienyl ethers of different degrees of alkylation were also formed. Trialkylamines are used as additives in the telomerization of butadiene and water in a two-phase system (103). The catalyst comprises a palladium salt and tppms or tppts. The amines may build cationic surfactants under catalytic conditions and be capable of micelle formation. The products include up to five telomerization products (alcohols, alkenes, and ethers), and thus the reaction is nonselective. Primary amines can be synthesized selectively by the catalytic two-phase telomerization of butadiene and ammonia [Eq. (12)], provided that the two amines are extracted from the aqueous catalyst phase with a polar solvent immediately after their synthesis (104).

2M +NH3

[cat.] /

HZOTToluene

+

NH2

I

In this way, the consecutive reactions to give higher amines in the aqueous phase are avoided; in contrast, these reactions are unavoidable in the homogeneous one-phase catalysis. The catalyst system consists of palladium acetate/tppts dissolved in water; the second phase is an organic solvent such as toluene. E. RING-OPENING METATHESIS POLYMERIZATION AND ISOMERIZATION The ring-opening metathesis polymerization (ROMP) of cyclic alkenes yields polymers that still contain all double bonds [Eq. (13)].

MULTIPHASE HOMOGENEOUS CATALYSIS

493

A well-accepted meclhanism starts with alkylidenemetal complexes as active catalytic species that insert the cyclic alkene to build a metallocyclobutane intermediate (105). Today, industrial and academic researchers are seeking structurally defined alkylidenemetal complexes that are capable of polymerizing monomers wit'h functional groups, e.g., oxanorbornene derivatives (106). These products may be used as speciality polymers and can be produced in aqueous media, with the advantages of good temperature control and good mixing resulting from polymer solubility in water. The group of Grubbs has investigated this reaction in depth. Following their early successes with hydrated RuC13 and with [Ru(Hz0)6]*+catalysts for ROMP of cyclic alkenes (107), they showed that a water-stable ruthenium carbene, derived from [ R L I C ~ ~ ( P Pand ~ ~diphenylcyclopropene, )~] was an extremely effective catalyst for the reaction with functionalized cyclic alkenes. The catalyst gave living polymers and allowed the synthesis of various block copolymers. Substitution of triphenylphosphane by tppts provided a water-soluble catalyst 13 with retention of its activity for the polymerization of water-soluble monomers (108).

Recently, this group published a route to new water-soluble, aliphatic phosphanes that build complexes of the same structure as 13 (109). Cy2P (CH2)2N(CH3);Cl-.CyzP(CHz)2S0;Nat, and others were prepared from air-stable, borane-protected precursors. The steric and electronic parameters of these new phosphanes were investigated. The preparation of a carbohydrate-functionalized polymer was made possible by the aqueous ROMP of carbohydrate-bearing 7-oxanorbornene with aqueous RuC13 as a catalyst (110). Furthermore, neoglycopolymers were generated via aqueous ring-opening metathesis polymerization with the aforementioned catalyst (111).These polymers were tested as inhibitors of the erythrocyte agglutinating activity of the carbohydrate-binding protein concanavalin A. Mechanistic studies of the rearrangement activity of the ring-opening metathesis polymerization catalyst [Ru(H20),Jzt were reported for unfunctionalized alkenes (112). The mechanism was found to be intermolecular, the alkene isomerization proceeding through an addition-elimination mechanism with a metal hydride catalytic species. This interpretation was

494

BIRGIT DRIESSEN-HOLSCHER

supported by the use of specifically deuterated substrates, by deuterium crossover experiments, and by carrying out the reaction in D20. F. OTHERCARBON-CARBON BOND-FORMING REACTIONS The past four years have been characterized by more and more diverse applications of aqueous organometallic chemistry for the synthesis of fine chemicals. Useful classical transformations have been realized in aqueous solutions. The recycle of the transition metal catalyst has sometimes been possible, but the main importance at this early stage lies in the reactions themselves. The palladium(I1)-catalyzed reaction of haloarenes with alkenes and alkynes (Heck-type reactions) in aqueous media has become known only in the preceding few years [Eq. (14)l.

X = I, Br, CI R = COOH, COOMe, CN,NO2, OH, Br, CI

Usually, the catalyst is prepared in situ from palladium(I1) salts, a tertiary phosphane, and a base (e.g., K2C03). Both inter- and intramolecular couplings have been investigated (113). Pioneering work in this area was done by Bumagin et al. (114), and several groups are working on this subject today. The most recent publications are cited here. The Heck arylation of ethene with iodoarenes was investigated with a preformed [PdC12(tppms)2]catalyst and gave the product styrenes in 60100% yield (115).In contrast, when the in situ system Pd(OAc)2 + 2tppms was used as a precursor, the conversion was less effective and yielded the hydration products. 0- and p-vinyltoluenes were prepared from bromotoluenes and ethene in dimethylformamide-water mixtures (116). In this case, potassium carbonate was used instead of triethylamine as the base. A double Heck reaction of a reactant for which P-hydride elimination is possible was described recently (117). The reaction was catalyzed by Pd(0) with phenanthroline ligand in aqueous media. The intramolecular version of Heck-type couplings was investigated for the first time in water, which led to a dramatic change in regioselectivity (118). With water-soluble Pdltppts catalysts, the generally observed exo

MUL.TIPHASE HOMOGENEOUS CATALYSIS

495

process was reversed in favor of the regioselective formation of endocyclized compounds. Another C-C coupling reaction is the copolymerization of ethene and carbon monoxide (IJ!9).The presence of water increased the copolymerization productivity up to 3.7 times, but the reason for the increase is not yet understood. Gr.

ORGANIC REACTIONS IN WATER

Classical organic reactions that have been carried out in water include, among others, the Diels-Alder reaction, the Claisen rearrangement, aldol condensations, Michael additions, and nucleophilic substitutions. In the Diels-Alder reaction, for example, water has been found to increase the reaction rate and to #enhancethe endoselectivity (120).Two reviews summarize the results for organic reactions in water (121). The potential of organic reactions compatible with or even promoted by water is not yet fully exploited. This is a good subject for future research.

V.

Multiphase Reactions with Solvents Other Than Water as a Catalyst Phase

A. IONICLIQUIDSAS CATALYST PHASES Room-temperature ionic liquids have been investigated as a new class of nonaqueous solvents for two-phase catalytic transformations. The class of orga.nochloroa1uminateionic liquids, typically a mixture of a quaternary ammonium salt such as 1,3-dialkylimidazoliumchloride with aluminum chloride (14), is the most widely explored system (122).

These liquids have been studied primarily for their applications as electrolytes in electrochetnical technologies such as electroplating, batteries, and alloy preparations, They have excellent chemical and thermal stabilities and are good solvents for highly charged complex ions of high or low oxidation states. The Lewis acidities can be varied with the composition of the liquid.

496

BIRGIT DRIESSEN-HOLSCHER

Chauvin and Olivier-Bourbigou (123) classified ionic liquids according to the complexing ability of their anions because they influence the solvation and complexing ability of ionic liquids. One problem is the instability of several ionic liquids in water, which reduces their potential for application in catalytic reactions. This subject is under investigation, and a series of novel air- and water-stable low-melting salts has recently been prepared (124). . , It was shown that room-temperature molten salts derived from the combination of 1,3-dialkylimidazoliumchloride and A1CI3can be used as solvents in two-phase catalytic dimerization of propene to give hexenes catalyzed by Ni(I1) compounds (125). The effects of phosphane ligands coordinated to nickel and operating variables were also investigated (126).The dimerization products separate as an organic layer above the molten salt. This reaction has been carried out with n-butenes as the reactant and cationic nickel complex catalysts dissolved in organochloroaluminate liquids (127). Several patents of BP Chemicals relate to the use of ionic liquids in catalysis. The polymerization of alkenes in ionic liquids was claimed (128), as was the alkylation of aromatic hydrocarbons with alkenes in the presence of an ionic liquid (129). Hydrogenation, isomerization, and hydroformylation of 1-pentene with cationic rhodium complexes were catalyzed in molten l-n-butyl-l-methylimidazolium salts (130).The ionic liquid can be recycled without significant loss of activity and the products isolated by simple phase separation. Recently, the air- and water-stable combinations of 1-n-butyl-3-methylimidazolium chloride with sodium tetrafluoroborate or sodium hexafluorophosphate have been prepared. The rhodium complexes [RhCl(PPh,),] and (Rh(COD),](BF,) are completely soluble in these ionic liquids and catalyze the hydrogenation of cyclohexene in a typical two-phase reaction with numbers of turnovers of up to 6000 (131).

B. DIOLSAS CATALYST PHASES The most prominent example of a two-phase process taking place in diols is the Shell oligomerization process, which is described in the following section. The reaction is carried out in 1,4-butanediol. Oligoethylene glycolsoluble cobalt catalysts have been synthesized for use in two-phase systems. Ritter et al. (132)prepared cobalt clusters with polyethylene glycol substituents (a-bonded fragments) for the hydroformylation of 1-hexene in liquid oligoethylene glycol. The ethylene glycol phase, including the cobalt cluster, can be recycled, but the nliso ratio of the product aldehydes decreased over the course of several runs.

MULTIPHASE HOMOGENEOUS CATALYSIS

497

C. FLUORINATED COMPOUNDS AS CATALYST PHASES The application of perfluorous polyethers in biphasic catalysis was first described by Vogt (.Z33),who also synthesized ligands based on hexafluoropropene oxide oligomers to create metal complexes that are soluble in the perfluorous polyethers. The solvophobic properties of the fluorous solvent were successfully incorporated in the metal complexes; catalytic oligomerization and polymerization reactions with nickel and cobalt complexes were demonstrated. In 1994, Horvath and Rabai (134) reported the so-called “fluorous biphase hydroformylation of alkenes.” Their fluorous biphase system (FBS) consists of a fluorous phase (mostly perfluorinated alkanes, ethers, and tertiary amines) containing a dissolved catalyst and another phase, which could be any common organic or inorganic solvent with limited solubility in the fluorocarbon. An FBS-compatible catalyst contains enough fluorous moieties that it is soluble preferentially in the fluorous phase and can be used for all hydrophobic alkenes in hydroformylation. The product aldehydes have a lower solubility than the alkenes in the fluorous phase. One metal complex used as a catalyst is [HRh(CO){P(CH2CHz(CF2)5CF3)}3]. The concept of fluorous biphasic catalysis is novel and appealing, but several questions remain regarding, for example, the activity and the lifetime of the catalyst, the costs, toxicities, and problems with the ozone depletion potential of the fluorine-containing compounds (135).This kind of process is not likely to find industrial application in the near future.

VI. Industrial Applications The concept of two-phase catalysis was first realized industrially in the Shell higher olefins,orocess (SHOP) (136),in which ethylene is oligomerized at 80-120” and 70-140 bar to give higher molecular weight, linear alkenes (C4-CI6). The process runs in a polar phase of lP-butanediol that also contains the organonickel catalyst formed from [Ni(COD),] and a (dipheny1phosphino)carbonic acid. The products separate as a second transparent liquid phase above the 1,4-butanediol and can therefore be removed easily. After this simple catalyst-product separation, catalyst traces are washed out of the alkene phase in a phase separator. Distillation of the products follows for purification of the alkenes, which are either marketed directly or fed into the next two SHOP catalytic steps of isomerization and metathesis. A SHOP flow diagram is shown in Fig. 3. This process now occupies a key position in alkene conversion technology.

498

BIRGIT DRIESSEN-HOLSCHER

reactor

phase separator

distillation

c10-14

FIG.3. Flow diagram of the Shell higher olefins process (SHOP) (136).

The history of aqueous industrial two-phase catalysis began in 1984 when the hydroformylation of propene in water was first carried out in the plants of Ruhrchemie AG. The development of the aqueous two-phase process was completely atypical in that the initial research work was done by RhBnePoulenc but the development work was done by the former Ruhrchemie (today part of Hoechst AG). A rather long time elapsed before further fundamental work was begun in academic laboratories. Two plants now produce 300,000 tons/year of butyraldehyde. Depending on the quality of the propene starting material, 99% conversion and a crude aldehyde product with an n/iso ratio of 95/5 can be attained (137). The water-soluble catalyst precursor is [HRh(CO)(tppts),]. The aldehydes form a second phase because they are not soluble in water. Side reactions, such as hydrogenation of the alkene and the formation of condensation products by aldol condensation, are insignificant. The loss of expensive rhodium in the product phase is very small. Some other advantages of the process concern the engineering part of the plant. No emissions are produced during the process, and the heat of the reaction can be used for distillation of the products. A flow diagram of the process is shown in Fig. 4. A further development of this successful technology was achieved to take advantage of the available feedstock base of butene isomers (raffinate 11) for the preparation of n-C5products (n-valeraldehyde, n-isoamyl alcohol, and n-valeric acid). In December 1995 production of n-valeraldehyde was started up in a new plant at HoechstlRuhrchemie (138). Generally, there are strong restrictions in the application of the two-phase catalytic processes to higher alkenes (Section IV.B.l), but the adaptation to butenes was possible with little modification of the process developed for propene.

MULTIPHASE HOMOGENEOUS CATALYSIS

499

n- Butanal vapors

3 Liquidkapour separator 't

FIG.4. Flow diagram of the RuhrchemieiRhBne-Poulenc process (137): 1, continuousflow, stirred tank reactor; 2, phase separator; 3, stripping column; 4, distillation column; 5, heat exchanger; 6, falling film evaporator; 7, liquid-vapor separator.

Another, similar propene hydroformylation process with rhodium and monosulfonated triphenylphosphane (tppms) was reported recently by Union Carbide (139). Capacity data are not available. On the industrial level, aqueous two-phase systems are used more often than nonaqueous two-phase systems. The Kuraray Co. operates a pilot plant for the hydrodimerization of 1,3-butadiene in a two-phase system with a Pd/tppms catalyst (140). The reaction is carried out in sulfolanewater, from which the products, the octadienols, separate. The final products can be octanol or nonanediol made by subsequent isomerization and hydroformylation. The capacity of the Kuraray process is about 5000 tondyear. RhBne-Poulenc uses carbon-carbon coupling for an efficient route to geranylacetone with a water-soluble Rhltppts catalyst (141). The addition of myrcene to acetylacetic acid methyl ester is regioselective (>99%) [Eq. (1511. Numerous diems can be used as reactants, e.g., isoprene, myrcene, and farnesene, and several compounds can be used as active methylene compounds. The reaction proceeds in an aqueous liquid-liquid system, with the conversion regulated by the time of contact between the phases, which is controlled by the stirring. The organic products are easily separated by simple decantation, and the aqueous phase containing the catalyst can be recycled. This reaction was industrialized to produce intermediates for vitamin E such as geranylacetone. The capacity is about 1000 tonslyear.

500

BIRGIT DRIESSEN-HOLSCHER

Myrcen

HZO, -CO,, -CH,OH

i

0

RhBne-Poulenc operates another biphasic process, the hydrogenation of a,@-unsaturatedaldehydes (64).The catalyst is readily made from hydrated RuC13 and tppts in water. The hydrogenation of various reactants (cinnamaldehyde, crotonaldehyde, or prenal) proceeds smoothly at low temperatures and under moderate partial pressures. It is possible to recycle the aqueous catalytic phase. The process is said to operate in a pilot plant, but the capacities are not known. In summary, six different industrial processes involving biphasic catalysis are known, and this new technology has proved be superior to traditional

MULTIPHASE HOMOGENEOUS CATALYSIS

501

one-phase processes. Increasing commercialization of new biphasic catalytic processes will surely follow.

VII.

Summary and Outlook

Multiphase catalysis with product-catalyst separation resulting from simple phase separation is now well established as one of the most novel and useful innovations in homogeneous catalysis. Multiphase catalysis is a technological breakthrough. The concept bridges the chasm between conventional homogeneous and heterogeneous catalysis. In contrast, the earlier idea of using solid-supported catalysts to bridge this chasm has been largely unsuccessful on the technological scale. Notwithstanding the successes of multiphase homogeneous catalysis, many questions remain and research is needed. Reaction mechanisms and kinetics must be investigated to explain the observed changes in activity and selectivity when water-soluble catalysts are applied. The synthesis of new water-soluble ligands, especially chiral derivatives, is an important challenge. The cost of valuable chiral ligands is often even higher than that of the transition metals themselves, and in multiphase catalysis both the ligand and the metal can be recycled. Investigations of aqueous biphasic reactions should include closer examinations of, for example, mass tramfer between phases and the role of salt effects, micelleforming agents, and cosolvents. The major problem associated with aqueous catalysis is the limited and often very low solubility of certain organic reactants in water. Much work is needed to find practical solutions for these hydrophobic reactants. Possibilities deserving further attention include the application of fluorous biphasic catalysis or nonaqueous ionic liquid catalysis. The potential of organic reactions compatible with or even promoted by water is not yet fully exploited. Strong growth is anticipated in both industrial and academic research on biphasic catalysis. REFERENCES 1. Barton, M., and Atwood, J. D., J. Coord. Chem. 24,43 (1991). 2. Kalck, P., and Monteil, F., Adv. Organomet. Chem. 34, 219 (1992). 3. Herrmann, W. A., and Kohlpaintner, C. W . ,Angew. Chem. 105, 1524 (1993). 4. Horvath, I. T., and Joo, F., eds. NATO ASZ Ser. 3, 5 (1995). 5. Roundhill, D. M., Adv. Organomet. Chem. 38, 155 (1995). 6. Cornils, B., and VJiebus, E., CHEMTECH 33 (1995). 7. Papadogianakis, G., and Sheldon, R. A,, New J. Chem. 20, 175 (1996).

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8. Joo, F., and Katho, A., J. Mol. Catal. 116, 3 (1997). 9. Joo, F., and Toth, Z., J. Mol. Catal. 8, 369 (1980). 10. Behr, A,, and Keim, W., Erdol. Erdgas. Kohle 103,126 (1987). 11. Sinou, D., Bull. Soc. Chim. Fr. 3,480 (1987). 12. Kuntz, E., CHEMTECH 17,570 (1987). 13. Southern, T. G., Polyhedron 8,407 (1989). 14. Atherton, J. H., Res. Chem. Kinet. 2, 193 (1994). 15. Alper, E., “Mass Transfer with Chemical Reaction in Multiphase Systems,” NATO AS1 Ser., Vol. I, Nijhoff Publishers, The Hague, 1983. 16. Hanna, J., and Noble, R. D., Chem. Rev. 85, 583 (1985). 17. Kuntz, E. G., FR-B 2 314 190 (RhBne-Poulenc Ifid.), (1975). 18. Hanson, B. E., Organometallics 12, 164 (1993). 19. Herrmann, W. A., Angew. Chem. 107,893 (1995). 20. Alario, F., Amrani, Y . , Colleuille, Y . ,Dang, T. P., Jenck, J., Morel, D., and Sinou, D., J. Chem. Soc., Chem. Commun. 202 (1986). 21. Amrani, Y . ,Lecomte, L., Sinou, D., Bakos, J., and Toth, I., Organometallics 8,542 (1989). 22. Lecomte, L., and Sinou, D., Phosphorus, Sulfur Silicon Relal. Elem. 53, 239 (1990). 23. Laghmari, M., and Sinou, D., J. Mol. Catal. 66, L15 (1991). 24. Bartik, T., Ding, H., Bartik, B., and Hanson, B. E., J. Mol. Catal. 98, 117 (1995). 25. Fell, B., Papdogianakis, G., Konkol, W., Weber, J., and Bahrmann, H., J. Prakt. Chem. 335, 75 (1993). 26. Schindlbauer, H., Monatsh. Chem. 96,2051 (1965). 27. Stelzer, G., Angew. Chem. 105,1097 (1993). 28. Mann, F. G., and Millar, I. T., J. Chem. Soc., 4453 (1952). 29. Issleib, K., and Thomas, G., Chem. Ber. 93, 803 (1960). 30. Padlahova, J., and Podlaha, J., Collect. Czech. Chem. Commun. 45, 2049 (1980). 31. Avey, A,, Schut, D. M., Weakley, T. J. R., and Tyler, D. R., Znorg. Chem. 32,233 (1993). 32. Mathey, F., and Mercier, F., J. Organornet. Chem. 462, 103 (1993). 33. Ravindar, V., Hemling, H., Schumann, H., and Blum, J., Synth. Commun. 22,841 (1992). 34. Fremy, G., Castanet, Y., Grzybek, R., Monflier, E., Mortreux, A,, Trzeciak, A. M., and Ziolkowski, J., J. Organornet. Chern. 505, 11 (1995). 35. Malmstrom, T., and Andersson, C., Chem. Commun., 1135 (1996). 36. Ganguly, S., Mague, J. T., and Roundhill, D. M., Znorg. Chem. 31, 3500 (1992). 37. Ganguly, S., and Roundhill, D. M., Organometallics 12, 4825 (1993). 38. Schull, T. L., Fettinger, J. C., and Knight, D. A., J. Chem. Soc., Chem. Commun., 1487 (1995). 39. Smith, R. T., and Baird, M. C., Transition Met. Chem. (London) 6, 197 (1981), Smith, R. T., Ungar, R. K., and Baird, M. C., ibid 7,288 (1982); Smith, R. T., and Baird, M. C., Znorg. Chim. Acta 62,135 (1982). 40. Toth, I., and Hanson, B. E., Organometallics 12, 1506 (1993); Toth, I., Hanson, B. E., and Davis, M. E., Catal. Lett. 5,183 (1990); Pfeiffer, G., Chan, S., Bendayan, A., Waegell, B., and Zahra, J. P., J. Mol. Catal. 59, 1 (1990); Nagel, U., and Kinzel, E., Chem. Ber. 119, 1731 (1986). 41. Bitterer, F., Kucken, S . , and Stelzer, O., Chem. Ber. 128, 275 (1995). 42. Dibowski, H., and Schmidtchen, F. P., Tetrahedron 51, 2325 (1995). 43. Buhling, A,, Kamer, P. C. J., and van Leeuwen, P. W. N. M., J. Mol. Catal. 98,69 (1995). 44. Renaud, E., Russell, R. B., Fortier, S., Brown, S . J., and Baird, M. C.. J. Organomet. Chem. 419,403 (1991). 45. Lavenot, L., Bortoletto, M. H., Roucoux, A,, Larpent, C., and Patin, H., J. Organornet. Chem. 509,9 (1996).

MULTIPHASE HOMOGENEOUS CATALYSIS

503

46. Amrani, Y., and Sinou, D., J. MoZ. Catal. 24,231 (1984); Wilson, M. E., Nuzzo, R. G., and Whitesides, G. M., J. Am. Chem. SOC.100,2269 (1978); Amrani, Y., and Sinou, D., J. Mol. Cafal.36,319 (1986); Nuzzo, R. G., Haynie, S. L., Wilson, M. E., and Whitesides, G. M., J. Org. Chem. 46,2861 (1981); Bergbreiter, D. E., Zhang, L., and Mariagnanam, V. M., J. Am. Chem. SOC.11.5, 9295 (1993). 47. Jin, Z., Zheng, X.. and Fell, B., J. Mol. Catal. 116, 55 (1997). 48. Okano, T., Iwahara, M., Suzuki, T., Konishi, H., and Kiji, J., Chem. Lett., 1467 (1986); Okano, T., Iwahara, M., Konishi, H., and Kiji, J., J. Organomet. Chem. 346,267 (1988). 49. Reuter, M., and Orthner, L., Ger. Pat. 1,035,135 (1958). 50. Harrison, K. N., Hoye, P. A. T., Orpen, A. G., Pringle, P. G., and Smith, M. B., J. Chem. SOC., Chem. Commun. 1096 (1989). 51. Tzschach, A., Radke, W., and Uhlig, W., Z . Chem. 19, 252 (1979). 52. Drucker, A., and Grayson, M., U.S. Pat. 3,489,811 (1967). 53. Horvath, I. T., and Joo, F., eds. NATO ASZSer. 3, 111 (1995). 54. Sreenivasa Reddy, V., Katti, K. V., and Barnes, C. L., J. Chem. SOC., Dulfon Trans., 1301 (1996). 55. Baxley, G. T., Weakley, T. J. R., Miller, W. K., Lyon, D. K., and Tyler, D. R., J. Mol. Cafa2.116,191 (1997); Baxley, J. T., Miller, W. K., Lyon, D. K., Miller, B. E.. Nieckarz, G. F., Weakley, T. J. R., and Tyler, D. R., Znorg. Chem. 35, 6688 (1996). 56. Smith, C. J., Reddy, V. S., and Katti, K. V., Chem. Commun., 2557 (1996). 57. Mitchell, T. N., and Heesche-Wagner, K., J. Organomet. Chem. 436, 43 (1992). 58. Kumar, A., Oehme, G., Roque, J. P., Schwarze, M., and Selke, R., Angew. Chem., Znf. Ed. Engl. 33,2197 (1994); Flach, H. N., Grassert, Y., and Oehme, G., Mucromol. Chem. Phys. 195,3289 (1994). 59. Darensbourg, D. J., Joo, F., Kannisto, M., Katho, A., Reibenspies, J. H., and Daigle, D. J.,Inorg. Chem. 33,200 (1994);Darensbourg,D. J.,J. Organornet. Chem. w 9 9 (1995). 60. Koelle, U., Coord. Chem. Rev. 135/136,623 (1994). 61. Martell, E. A., and Hancock, R. D., “Metal Complexes in Aqueous Solutions.” Plenum, New York, 1996. 62. European Chemical News, January 15, p. 29 (1995); Europa-Chemie 1 , l O (1995). 63. Kulpe, J. A., PhD Thesis, Technical University of Munich (1989). 64. Beller, M., Cornil:;, B., Frohning, C. D., and Kohlpaintner, C. W., 1. Mol. Cafal. A 104, 17 (1995). 65. Herrmann, W. A,, Kohlpaintner, C. W., Manetsberger, R. W., Bahrmann, H., and Kottmann, H., J. Mol. Card A 97, 65 (1995); Bahrmann, H., Bach, H., Frohning, C. D., Kleiner, H. J., Lappe, P., Peters, D., Regnat, D., and Herrmann, W. A,, ibid. 116,49 (1997). 66. Russel, M. J. H., Platinum Met. Rev. 32,179 (1988); Bahrmann, H., and Lappe, P., Eur. Pat. Appl. EP 602:, 463 (1994). 67. Ding, H., Hanson, B. E., Bartik, T., and Bartik, B., Organometullics W, 3761 (1994); Fell, B., and Papadogianakis, G., J. MoZ. Cafal. 66, 143 (1991). 68. Bartik, T., Bartik, B., and Hanson, B. E., J. MoZ. Cutul. 85,121 (1993); Sinou, D., Safi, M., Claver, C . , and Masdeu, A., ibid. 68, L9 (1991); Laghmari, M., and Sinou, D., ibid. 66, L15 (1991); Bakos, J., Karaivanov, R., Laghmari, M., and Sinou, D., Organometallics 13,2951 (1994). 69. Monflier, E., Fremy, G., Castanet, Y., and Mortreux, A., Angew. Chem. 107,2450 (1995); Anderson, J. R., Campi, E. M., and Jackson, W. R., Catal. Lett. 9,55 (1991). 70. Reetz, M. T., and Rudolph, J., Tetrahedron Asymmetry 4, 2405 (1993); Reetz, M. T., and Waldvogel, S . R., Angew. Chem. 109, 870 (1997). 71. Kanagasabapathy, S., Xia, Z., Papadogianakis, G., and Fell, B., J. Prakt. Chem. 337, 446 (1995).

504

BIRGIT DRIESSEN-HOLSCHER

72. Buhling, A,, Kamer, P. C. J., van Leeuwen, P. W. N. M., and Elgersma, J. W., J. Mol. Cufal. A 116,297 (1997). 73. Fremy, G., Castanet, Y . ,Grzybek, R., Monflier, E., Mortreux, A,, Trzeciak, A. M., and Ziolkowski, J. J., J. Organomet. Chem. 505, 11 (1995). 74. Fremy, G., Monflier, E., Carpentier, J. F., Castanet, Y . ,and Mortreux, A., Angew. Chem., Int. Ed. Engl. 34, 1474 (1995). 75. Fell, B., Schobben, C., and Papadogianakis, G., J. Mol. Catul. A 101, 179 (1995). 76. Bartik, T., Ding, H., Bartik, B., and Hanson, B. E., J. Mol. Cutul. A 98, 117 (1995). 77. Cassar, L., Chim. Ind. (Milan) 67, 256 (1985). 78. Khan, M. M. T., Halligudi, S. B., and Abdi, S. H. R., J. Mol. Cutul. 44, 179 (1988). 79. Mvnflier, E., Mortreux, A., and Petit, F., Appl. Cuful. A 102, 53 (1993); Monflier, E., and Mortreux, A,, J. Mol. Catul. 88, 195 (1994). 80. Okano, T., Hayaschi, T., and Kiji, J., Bull. Chem. Soc. Jpn. 67, 2339 (1994). 8I. Monteil, F., and Kalck, P., J. Organomet. Chem. 482, 45 (1994). 82. Kohlpaintner, C. W., and Beller, J., J. Mol. Cutul. A 116, 159 (1997). 83. Papadogianakis, G., Maat, L., and Sheldon, R. A,, J. Chem. Soc., Chem. Commun., 2659 (1994). 84. Papadogianakis, G., Maat, L., and Sheldon R. A,, J. Mol. Caful. A 116, 179 (1997). 85. Chaloner, P. A,, Esteruelas, M. A,, Joo, F., and Oro, L. A. “Homogeneous Hydrogenation, in Catalysis by Metal Complexes,” Vol. 15. Kluwer Academic Publishers, Dordrecht, The Netherlands, 1994. 86. Meister, G., Rheinwald, G., Stoeckli-Evans, H., and Suss-Fink. G., J. Chem. Soc., Dalton Trans., 3215 (1994). 87. Chan, W.-C., Lau, C.-P., Cheng, L., and Leung, Y.-S., J . Orgunornet Chem. 464, 103 (1995). 88. Joo, F., Nadasdi, L., Benyei, A. C., and Darensbourg, D. J., J. Organomet. Chem. 512, 45 (1996). 89. Darensbourg, D. J., Stafford, N. W., Jvo, F., and Reibenspies, J. H., J. Orgunornet. Chem. 488,99 (1995). 90. Darensbourg, D. J., Joo, F., Kannisto, M., Katho, A,, and Reibenspies, J. H., Inorg. Chem. 33, 200 (1994). 91. Hernandez, M., and Kalck, P., J. Mol. Catal. A 116,131 (1997). 92. Andriollo, A., Carrasquel, J., Marino, J., Lopez,F. A.,Paez, D. E., Rojas, I., and Valencia, N., J. Mol. Catal. A 116, 157 (1997). 93. Sanchez-Delgado, R. A,, Medina, M., Lopez-Linarez, F., and Fuentes, A., J. Mol. Cutul. A 116, 167 (1997). 94. Merola, J. S., Franks, M. A., Chiril, P., and Pafford, R., 209th Meet. Am. Chem. Soc., Anaheim, CA, INOR 416 (1995). 95. Nomura, K., J. Mol. Cutul. A 95, 203 (1995). 96. Ding, H., Hanson, B. E., and Bakos, J., J. Angew. Chem., Int. Ed. Engl. 34,1645 (1995). 97. Kumar, A., Oehme, G I Roque, J. P., Schwarze, M., and Selke, R., Angew. Chem., Int. Ed. Engl. 33,2197 (1994); Flach, H. N., Grassert, Y., and Oehme, G., Mucromol. Chem. Phys. 195, 3289 (1994). 98. Grassert, I., Vill, V., and Oehme, G., J. Mol. Catul. A 116, 231 (1997). 99. Bakos, J., Orosz, A., Cserepi, S., Toth, I., and Sinou, D., J. Mol. Cuful. A 116,85 (1997). 100. Kuntz, E., DE-B 27 335 16 (Rh6ne-Poulenc) (1978). 101. Pfeiffer, G., Chan, S., Bendayan, A., Waegell, B., and Zahra, J.-P., J. Mol. Catal. 59, 1 (1990). 102. Pennequin, J., Mortreux, A., Petit, F., Mentech, B., and Thiriet, B., FR-B 2693 188 (1994).

MULTIPHASE HOMOGENEOUS CATALYSIS

505

103. Monflier, E., Bourdauducq, P., Couturier, J.-L., Kervennal, J., and Mortreux, A., J. Mol. Catal. A 97, 29 (1995). 104. Prinz, T., Keim, W., and Driessen-Holscher, B., Angew. Chem., Int. Ed. Engl. 35, 1708 (1996). 105. Herisson, J. L., and Chauvin, Y., Makromol. Chem. 141, 161 (1970). 106. Herrmann, W. A,, Chem. Unserer Zeit 30, 152 (1996). 107. Lynn, D. M., Kanaoka, S., and Grubbs, R. H., J. Am. Chem. Soc. 118,784 (1996). 108. Grubbs, R. H., N A T O A S I Ser. 3(5), 15 (1995). 109. Mohr, B., Lynn, D. M., and Grubbs, R. H., Organometallics 15,4317 (1996). 110. Mortell, K. H., Gingras, M., and Kiessling, L. L., J. Am. Chem. Soc. 116, 12053 (1994). 111. Schuster, M. C., “ortell, K. H., Hegeman, A. D., and Kiessling, L. L. J. Mol. Catal. A 116, 209 (1997). 112. Karlen, T., and Lu.di, A., J. Am. Chem. SOC.116, 11375 (1994). 113. Grotjahn, D. B., and Zhang, X., N A T O A S I Ser. 3(5), 123 (1995). 114. Bumagin, N. A., More, P. G., and Beletskaya, I. P., J. Organomet. Chem. 371,397 (1989). 115. Kiji, J., Okano, T., and Hasegawa, T., J. Mol. Catal. A 97, 73 (1995). 116. DeVries, R. A., arid Mendoza, A., Organometallics l3,2405 (1994). 117. Grotjahn, D. B., and Zhang, X., J. Mol. Catal. A 116, 99 (1997). 118. Lemaire-Audoire, S., Savignac, M., Dupuis, C., and Genet, J.-P., Tetrahedron Lett. 37, 2003 (1996). 119. Vavasori, A,, and Toniolo, L., J. Mol. Catal. A 110, 13 (1996). 120. Rideout, D. C., and Breslow, R., J. Am. Chem. Soc. 102,7816 (1980). 121. Li, C.-J., Chem. Rev. 93, 2023 (1993); Lubineau, A,, Synthesis, 741 (1994). 122. Hussey, C. L., Pure Appl. Chem. 60, 1763 (1988); Mamantov, G., and Popov, A. I., “Chemistry of Noriaqueous Solutions, Current Progress,” p. 277. VCH, Weinheim, 1994. 123. Chauvin, Y., and Olivier-Bourbigou, H., CHEMTECH, 26 (1995). 124. Wilkes, J. S., and Zaworotko, M. J., J. Chem. Soc., Chem. Commun., p. 965 (1992). 125. Chauvin, Y., Gilbert, B., and Guibard, I.,J. Chem. Soc., Chem. Commun., p. 1715 (1990). 126. Chauvin, Y., Einloft, S., and Olivier, H., Ind. Eng. Chem. Res. 34, 1149 (1995). 127. Dupont, J., J. Catal. 165, 275 (1997). 128. Ambler, P. W., Hodgson, P. K. G., and Stewart, N. J., EP 0 558 187 A1 (BP Chemicals) (1993); Abdul-Sada, A. K., Ambler, P. W., Hodgson, P. K. G., Seddon, K. R., and Szewart, N., International Pat. WO 95121871 (1995). 129. Abdul-Sada, A. K., Atkins, M. P., Ellis, B., Hodgson, P. K. G., Morgan, M. L. M., and Seddon, K. R., 1nl:ernational Pat. WO 95/21806 (1995). 130. Chauvin, Y., Musrimann, L., and Olivier, H., Angew. Chem. 107, 2941 (1995). 131. Suarez, P. A. Z., Diullius, J. E. L., Einloft, S., D e Souza, R. F., and Dupont, J., Polyhedron 15, 1217 (1996). 132. Ritter, U., Winkhofer, N., Schmidt, H.-G., and Roesky, H. W., Angew. Chem., Int. Ed. Engl. 35, 524 (1996); Ritter, U., GIT Fachz. Lab., 969 (1996). 133. Vogt, M., Ph.D. Thesis, RWTH Aachen (1991). 134. Horvath, I. T., an.d Rabai, J., Science 266, 72 (1994). 135. Gladysz, J . A., Science 266,55 (1994); Cornils, B., Angew. Chem. 107,1709 (1995). 136. Keim, W., Shryne, T. M., Bauer, R. S., Chung, H., Glockner, P. W., and van Zwet, H., DE-P 2 054 009 (:Shell Int. Res.), (1969); Keim, W., Chem.-1ng:Tech. 56, 850 (1984). 137. Wiebus, E., and Cornils, B., Chem.-1ng.-Tech. 66, 916 (1994). 138. Hoechst AG, Patent DE-OS 4 333 323 (1993). 139. Haggin, J., Chem. Eng. News, April, p. 25 (1995). 140. Tokitoh, Y., and Yoshimura, N. (Kuraray Co.), U.S. Pat. 5,057,631 (1991); E P 0,436,226 (1990). 141. Mercier, C., and Chabardes, P., Pure Appl. Ckem. 66, 1509 (1994).

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Index A Absorption spectroscopy catalysis characterization with X-ray diffraction, 315-321, 340-342 combined EXAFSKRD methods, 330-340 in situ, 321-324, 332-340 limitation, 318 techniques, 327 Acetaldehyde, NMR study, 162 Acetone, mesityl oxide from, 162 Acetonitrile, 168 Acetylene magnetic anisotropy, 120-122 vibrational spectra, 183-202, 298-299 Acetylium ion derivation, 139 NMR, 127-128 Achiral reactants, hydrogenation with, 489-498 Acid strength, 119 Acyclic alkanes, vibrational spectra, 207214,299 Acyclic alkenes, vibrational spectra, 264267, 298 Acyclic alkynes, vibrational spectra, 183207, 298-299 Acylium ions chemical shift tensor, 135 Koch-Haaf reaction, 160 NMR, 128 as probe molecules, 139, 160 Adsorption sulfur compounds, 426-429,447 zeolites, simulations, 50-84 Aldehydes, unsaturated, hydrogenation of, 489, 500 Alkali metal phosphides, 479 Alkanes cracking hydrodesulfurization and, 438 in zeolites, 102- 106, 158

diffusion in silicalite, 34, 35 sorption on zeolites, 71-80 vibrational spectra acyclic, 207-214, 299 cycloalkanes, 229-239, 299 Alkenecarboxylic acid methyl esters, hydroformylation of, 487 Alkenes biphasic hydroformylation of, 485-486 biphasic hydrogenation of, 489 cyclic ring-opening methathesis polymeriza tion (ROMP), 492-493 vibrational spectra, 239-245, 299 haloarenes, reaction with, 494 hydroformylation of, 497 Shell higher olefins process (SHOP), 497-498 vibrational spectra acyclic alkenes, 264-267, 298 cycloalkenes, 239-245, 299 Alkylbenzenes, cracking, hydrodesulfurization and, 438 Alkyl carbenium ions, as probe molecules, 154-157 Alkylphosphanes, chelating, 482 Alkynes acyclic, vibrational spectra, 183-207, 298-299 haloarenes, reaction with, 494 Allene, vibrational spectra, 206-207 Ally1 alcohol, HZSM-5 studies, 143-144, 145 AIIyl group, as hydrocarbon surface species, vibrational spectra, 228 Aluminum chloride-l,3-dialkylimidazolium chloride catalyst system, 495-496 Amines, aromatic, hydrogenation of, 490 (Aminoalkyl)phosphanes, synthesis, 480 Ammonia, 172-174 Amphos ligand, 480 507

508

INDEX

Anionic ligands, rnultiphase catalysis, 476-479 Arenium ions, as probe molecules, 150-154 Argon diffusion in zeolites, 15-19 sorption on zeolites, 60-61 Aromatic compounds diffusion in zeolites, 40-50 hydrogenation of aromatic amines, 490 of single aromatic rings, 438-439 sorption on zeolites, 80 vibrational spectra, aromatic hydrocarbons, 245-267,299 Arylphosphanes, crown-ether-substituted ligands, 481 ASM-5, alkane sorption on, 71 Asymmetric hydrogenations, 490-491 Asymmetry factor, 123-124

B Benzaldehyde, trityl cation from, 148, 159 Benzene diffusion in zeolites, 40-51 sorption on zeolites, 80-83 trityl cation from, 147, 148 vibrational spectra, 245-263 Benzenium ion, chemical shift tensors, 124, 150 Benzothiophenes electron densities and bond orders, 429-431 hydrodesulfurization, 351-353 Benzoylium ion, chemical shift tensors, 124 Bicyclohexyls, 373 Biphasic systems, catalysis, see Multiphase homogeneous catalysis Bond activation by zeolites, 2, 84-87, 106-107 hydrocarbons, 98-106 methanol, 87-98 quantum mechanical simulations, 2 Bond strength, metal-sulfur coordination, 425 BPTDS (bi-postdosing thermal desorption spectroscopy), cyclohexene, 240 Bromobenzene, carbonylation of, 488 Bromonium ion, 154

Butadiene, telomerization with sucrose, 492 Butane cracking in zeolites, 104 diffusion in zeolites, 37-38, 39 sorption on zeolites, 71-72 vibrational spectra, 213 Butene isomers diffusion in DAF-1, 36 sorption on zeolites, 78 tert-Butyl cation, 116, 157 Butyl groups, as hydrocarbon surface species, vibrational spectra, 226, 227 1-n-Butyl-1-methylimidazolium salts, 496 But-2-yne. vibrational spectra, 202-204

C

Carbenium ions, 115, 143 chemical shift tensors, 124-125 fragments in zeolites, 92-93 history, 116 superacids, 117 Carbocations, 115 Carbon-carbon bond-forming reactions, Heck arylation, 494-495 Carbonium ions, 115 as probe molecules, 157-160 Carbonylation reactions, biphasic, 487-488 Carboxyphenylhosphanes, 479 Cartwheel motion, intercage migration, 45-46 Catalysis by solid acids, 115-174 by zeolites, simulations, 1-107 dehydrosulfurization, 351-353, 416-425, 456-461 multiphasic, 473-501 Catalysts in sztu characterization, X-ray diffusion and absorption spectroscopy, 315-342 metal-S and metal-SH,, 351 Cationic ligands, rnultiphase catalysis, 479-481 CAVERN device, 138, 141 Chabazite, methanol adsorption in, 91 Chalcogenonium ions, as probe molecules, 161-162 Chemical shift, 120-122 computational chemistry, 129-137

509

INDEX molecular structure and, 129-133 theoretical calculations, 133-137 theory, 122-129 Chemical shift anisotropies (CSA), 123-124 Chemical shift tensor, 124-125, 133-135 Cinnamaldehyde, hydrogenation of, 490 Cloverite, xenon diffusion in, 14 Cluster model adsorbate-zeolite :;ystems, 2, 84-107 HZSM-5, 131, 133 Cobalt, ethyne spectl-a on finely divided metal, 199 Cobalt catalysts, oligoethylene glycol-soluble, 496 Cobalt(nicke1)-molybdenum-sulfur catalysts, 417 geometric considei-ations, 406-408 with higher activities, 456 mechanism, 401-403 number of active sites, 404-406 potential for increased activity, 415-416 preparation, 398-401 site characterization, 395-398 structure and classification, 390-395 two-site dilemma, 408-415 Color formation, in hydrodesulfurization, 439-440 Computational chemistry, chemical shift, 129-137 Configurational-bias Monte Carlo method (CB-MC), adsorption in zeolites, 5253, 72, 74, 75 Copper, ethyne spectra on finely divided metal, 199 Cracking, in zeolites, 102-106, 158 Cyclic alkenes, ring-opening methathesis polymerization (ROMP), 492-493 Cycloalkanes, vibrational spectra, 229-239, 299 Cycloalkenes, vibrational spectra, 239-245, 299 Cyclodextrins, use in hydroformylation, 485-486 Cyclohexadiene, vibrational spectra, 243 Cyclohexane, vibrational spectra, 229-235 Cyclohexene, vibrational spectra, 239-242 Cyclohexylbenzenes, 373 Cyclohexyl group, as hydrocarbon surface species, vibrational spectra, 228

Cyclooctatetraene, vibrational spectra, 243-244 Cyclopentadiene, vibrational spectra, 243 Cyclopentane, vibrational spectra, 235-237 Cyclopentene, vibrational spectra, 242-243 Cyclopentenyl cations, as probe molecule, 140-143 Cyclopropane, vibrational spectra, 237-239 Cyclopropyl group, as hydrocarbon surface species, vibrational spectra, 228 D

DAF-I, butene diffusion in, 36 1-Decene, hydroformylation of, 485-486 Deep desulfurization, 345-349, 466-467 kinetics, 352, 363-365, 372, 427, 436-438, 441-443 limitations, 349, 435, 455-456 color formation, 439-440 feedstock composition, 438-440 process equipment, 435-436 reaction rate inhibition, 351,440-455 thermodynamics, 436-438 new approaches to, 456-466 process schemes, 366-369 alkyl substituent effect on, 385-389 catalyst structure and composition, 350, 390-416 catalytic mechanism, 351-353, 416-425 reaction mechanism, 369-383, 384 quantum chemical calculations electron density calculations, 429-434 metal-sulfur coordination bond strength, 425-429 sulfur species in middle-distillate oils gas oil composition, 353-360 other feed components in, 366 reactivity, 360-365 systematic approach, 349-353 Density functional theory (DFT) chemical shift, 131-133, 134 sorption on zeolites, 86, 90, 92 sum-over-state (SOS) method, 134 DEXGFS (dispersive EXAFS), 327, 328, 330-340 Dialkyldibenzothiophenes, electronic properties of, 429

510

INDEX

1,3-Dialkylimidazolium chloride-aluminum chloride catalyst system, 495-496 Dibenzothiophene alkyl substitution, 389, 457 hydrodesulfurization, 370, 373-374, 427 hydrogenation rate constant, 375 molecular structure, 406 Dibenzothiophenes, electron densities and bond orders, 429-431 Dienes, dimerization, 491-492 Diffusion, in zeolites, simulations, 2, 3-50 Dimethylcyclohexanes, vibrational spectra, 239 Dimethylcyclohexylbenzene (DMCB), 432 4,6-Dimethyldibenzothiophene(4,6DMDBT), desulfurization of, 386, 387,389,406,457-460 Dimethyl ether, formation in zeolites, 95-98 Dimethylphenyl carbenium ion, 146 Diols, as catalyst phase, 496

[(Diphenylphosphino)alkyI]phosphonates, 479 Diphosphane, hydrogenation catalyst, 490 Double Heck reaction, 494 E Energy minimization methods (EM) zeolite adsorption simulations, 53, 56, 71, 78 zeolite “forced” diffusion simulations, 4, 7-8 Environmental issues, petroleum products, 346, 348 Ethane cracking in zeolites, 102-103 diffusion in silicalite, 34-35 vibrational spectra, 210-212 Ethene activation in zeolites, 101-102 diffusion in silicalite, 35 Heck arylation of, 494 hydrogenation reactions o n metal oxidesupported platinum catalysts, 283293, 297-298 sorption o n zeolites, 78-79 Ethylbenzenium ion, chemical shift tensors, 150

Ethyl cation, 154, 155 Ethylene, Shell higher olefins process (SHOP), 497-498 Ethylenediaminetetraacetic acid, phosphane analog, 479 Ethyl groups, as hydrocarbon surface species, vibrational spectra, 221 -224, 282-283 Ethylidene groups, as hydrocarbon surface species, vibrational spectra, 224 Ethylidyne ethene hydrogenation and, 281, 292-293 hydrogenation, 292-293 vibrational spectra, 189 Ethyne diffusion in silicalite, 35 vibrational spectra, 183-202, 298-299 EU-1, structure, 30, 41 EXAFS (extended X-ray absorption fine structure), 316, 318, 319 Co-Mo-S, carbon supported, 408 Co(Ni)-Mo-S catalysts, 397 in situ, 322, 323 technique, 327-330 combined with XRD, 330-340 F Faujasites, sorption on alkanes, 71 benzene, 80-81 methane, 62-64 single atoms, 54-55 Ferrierite lattice, 19 First principles methods, adsorption in zeolites, 91 Fluoranthene (FLU), thermodynamic equilibria of, 440 Fluorinated compounds, as catalyst phases, 497 p-Fluoronitrobenzene, 168, 171 Fluorous biphase system (FBS), 497 Force field, diffusion in zeolites, 6 G Gas chromatography-atomic emission detection (GC-AED), polyaromatic sulfur-containing compounds, 355, 356, 360, 361

INDEX Gas chromatograph) -mass spectrometry (GC-MS) chemical shift, 130 polyaromatic sulfur-containing compounds, 355, 3!i6, 360 sorption on zeolites, 58, 60-61, 69, 82 Gas oil composition, 353-360 deep desulfurization, 345-349, 466-467 catalysis, 350-353, 390-425 limitations, 349,435, 455-456 new approaches to, 456-466 process schemes, 366-425 quantum chemical calculations, 425-434 reaction mechanism, 369-383, 384 systematic approach, 349-353 defined, 354 gas chromatographic analyses, 346, 347 properties, 354, 355 reactivity of sulfur species, 360-365 refining, 348 GIAO method (gauge-incIuding atomic orbitals), chemical shift calculation, 133, 134 Gold, ethyne spectra on finely divided metal. 201

H Haloarenes, reaction with alkenes and alkynes, 494 Hammett acidity, superacids, 117 Heck arylation, 494--495 Heptane, sorption on zeolites, 76 Hexafluoropropene oxide oligomers, 497 Hexane isomers diffusion in zeolitm, 37-38, 39 sorption on zeolites, 71, 72, 76 HREELS (high-resolution energy loss spectroscopy), see E E L S (vibrational electron energy loss spectroscopy) Hydrocarbons activation in zeolites, 98-106 diffusion in zeolites, 34-40 polyaromatic, in gas oil, 358-360 sorption on zeolites, 70-80 vibrational spectra, aromatic hydrocarbons, 245-267. 299

511

Hydrocarbon surface species, vibrational spectra, 214 ally1 group, 227 t-butyl groups, 227 n-butyl and isobutyl groups, 226 cyclohexyl group, 228 cyclopropyl group, 228 ethyl groups, 221-224,282-283 ethylidene groups, 224 HCCH group, 228 methylene groups, 219-220 methyl groups, 214-219 methylidyne groups, 220-221 phenyl group, 228 1-propyl and 2-propyl groups, 224-225 trimethylenedimetallo groups, 227 vinyl group, 227-228 Hydrodesulfurization, 345-349, 466-467 catalysts, 350, 390-416, 456-461 computational aids, 425-434 kinetics, 352, 363-365, 372, 427, 436-438, 441-443 limitations, 349, 435, 455-456 color formation, 439-440 feedstock composition, 438-440 process equipment, 435-436, 464-466 reaction rate inhibition, 351, 440-455 thermodynamics, 436-438 new approaches to, 456-466 process schemes, 366-369 alkyl substituent effect on, 363, 385-389 catalysts, 350, 390-416, 456-461 catalytic mechanism, 351-353, 416-425 reaction mechanism, 369-383, 384 staged process operations, 462-464 quantum chemical calculations electron density calculations, 429-434 metal-sulfur coordination bond strength, 425-429 sulfur species in middle-distillate oils gas oil composition, 353-360 other feed components in, 366 reactivity, 360-365 systematic approach, 349-353 Hydroformylation, 483-487 Hydrogenation with achiral reactants, 489-498 asymmetric, 490-491 multiphase catalysis, 488-492 single aromatic rings, 438-439

512

INDEX

5-(Hydroxymethyl)-2-furfural (HMF), carbonylation of, 488 HZSM-5 acetone reaction with ammonia on, 166-167 ally1 alcohol on, 143 cluster model, 131, 133 cyclopentenyl cations in, 142 MAS NMR spectra, 120, 121 phenylindanyl cation in, 146 propene reaction on, 155 trimethyloxonium cation on, 161

I IGLO method (individual gauge for localized orbitals), chemical shift calculation, 133-134 Indanyl cations, as probe molecule, 144- 147 Infrared methods, 296 adsorption of benzene, 254-260 Inhibition, hydrodesulfurization, 351, 440-441 binary mixture studies, 448-454 mathematical modeling, 441-444 Iodoarenes, Heck arylation of ethene with, 494 Ionic liquids, multiphase catalysis with, 495-496 Iridium complex catalyst, hydrogenation of, 490 IRRAS, see RAIRS (reflection-absorption infrared spectroscopy) Isobutyl groups, as hydrocarbon surface species, vibrational spectra, 226 Isopropyl cation, 154, 155 chemical shift tensors, 124 geometries, 135, 136

K Ketones, biphasic hydrogenation of, 489 Kinetics deep desulfurization, 352, 363-365, 372, 427,436-438,441-443 reactivity of ethene on platinum, 273279, 284

Kirkwood-Muller approach, 8 Koch-Haaf reaction, acylium ions, 160

L

Langmuir-Hinshelwood mechanism, 293, 446, 447 Larmor frequency, 127 LCO (light cycle oil), 355 LEED (low-energy electron diffraction), 185, 231 Lennard-Jones function, diffusion in zeolites, 8, 9 Liquid-liquid biphasic systems, catalysis, see Multiphase homogeneous catalysis

M Magic angle spinning (MAS), 125-126 Magnetogyric ratio, 122 Mathematic modeling, hydrodesulfurization inhibition, 441-444 Metal complexes, water-soluble catalysts, 483 Metallathiabenzenes formation of, 421 reaction with hydrogen gas, 420 Metals, vibrational spectra of adsorbed hydrocarbons allene, 206-207 butane isomers, 213-214 ethane, 210-212 ethene, 298 ethyne, 191-202 methane, 210 phenylacetylene, 206 propadiene, 206-207 propyne, 205-206 Methane dehydrogenation of, 100-101 diffusion in zeolites, 20-34 hydrogen exchange in zeolites, 98-100 sorption on zeolites, 71 faujasites, 62-64 mordenite, 65-66 silicalite, 66-70 zeolite A, 64-65 vibrational spectra, 207-210

513

INDEX Methanol activation and reaction in zeolites dehydration, 92--93 dimethyl ether formation, 95-98 proton transfer, 87-91 methanol-zeolite interaction, 106 sorption on zeolites, 78-79, 91 Methyl a-acetamidocinnamade, hydrogenation of, 490, 491 Methyl acrylate, hydroformylation of, 486-487 Methylcyclohexane, vibrational spectra, 239 Methylcyclopentyl cation, 157 Methylene groups, as hydrocarbon surface species, vibrational spectra, 219-220 Methyl groups, as hydrocarbon surface species, vibrational spectra, 214-219 Methylidyne groups, as hydrocarbon surface species, vibrational spectra, 220-221 Metropolis Monte Carlo method, sorption on zeolites, 62, 66 Middle-distillate oils composition, 353-,360 deep desulfurization, 345-349, 466-467 catalysis, 350-3'83, 390-425 limitations, 349, 435, 455-456 new approaches to, 456-466 process schemes, 366-425 quantum chemic:al calculations, 425-434 reaction mechanism, 369-383,384 reactivity of sulfur species, 360-365 systematic approach, 349-353 Molecular dynamics diffusion in zeolites, 2, 4-6 argon, 20 aromatic compounds, 40-50 hydrocarbons, 3'1-40 methane, 26, 32 noble gases, 24 propane, 36 xenon, 9-11,13 limitations, 51 sorption on zeolites benzene, 81 butane, 72 methane, 58, 64--65, 66, 69 Monte Carlo simulations, adsorption in zeolites, 2, 51-52

Mordenite methane sorption on, 65-66 structure, 30 MP2 method, 135 MSD, diffusion in zeolites, 15, 33 MSSR (metal-surface selection rule), 208 Multiphase homogeneous catalysis, 473474, 501 aqueous reactions anionic ligands, 476-479 carbon-carbon bond-forming reactions, 494-495 carbonylation reactions, 487-488 cationic ligands, 479-481 hydroformylation, 483-487, 498 hydrogenations, 488-491 metal salts as catalysis, 482-487 neutral ligands, 481-482 organic reactions, 495 0x0 synthesis, 483-487 ring-opening metathesis polymerization and isomerization, 492-494 telomerizations, 491-492 diols as catalyst phase, 496 fluorinated compounds as catalyst phase, 497 industrial applications, 497-501 ionic liquids as catalyst phase, 495-496 mass transfer, 474 principles, 474-476

N Neopentane, vibrational spectra, 212, 213 Neutral ligands, multiphase catalysis, 481-482 NEXAFS (near-edge X-ray absorption finestructure), ethyne, 187, 189 Nickel, ethyne spectra on finely divided metal, 192-194 Nitrogen-containing compounds, as probe molecules, 165 Nitromethane, as probe molecule, 167 NMR, solid acidity study, 115-174 chemical shift, 120-137 computational chemistry, 129-137 probe molecules, 139-174 sample preparation, 137-139

514

INDEX

0 1,7-Octadienol, synthesis of, 491 1-Octene, hydroformylation of, 486 Organochloroaluminate ionic liquids, as catalysts, 495-496 Organometallic catalysts, hydrodesulfurization, 413, 417-425, 432 Oxonium ions, 139-140 0 x 0 synthesis, 483

P Palladium, ethyne spectra on finely divided metal, 194-196 1-Pentene, ionic liquid catalyst for, 496 Perfluorous polyethers, in biphasic catalysis, 497 Petroleum products environmental issues, 346, 348 hydrodesulfurization, 349-349,466-467 catalysis, 350-353, 390-425 kinetics, 352, 363-364, 372, 427, 436438, 441-443 limitations, 349, 435, 455-456 new approaches to, 456-466 process schemes, 366-425 reaction mechanism, 369-383, 384 sulfur species in middle-distillate oils, 353-366 systematic approach, 349-353 Phenyl group, as hydrocarbon surface species, vibrational spectra, 228 Phenylindanyl cation, NMR spectra, 146 Phosphanes aliphatic, 493 hydroxyalkyl-substituted, 482 quaternary ammonium, 492 sugar derivatives, 482-483 sulfonated, 476-478 Phosphanorbornadienes, water-soluble carboxylates, 479 Phosphines, as probe molecules, 170, 172 (Phosphinoalky1)phosphonium salts, 480 Piezo-QEXAFS, 330-344 Platinum, spectra on finely divided metal ethene, 273-283,297-298 ethyne, 197 Polanyi principle, 87 Polyacetylene, vibrational spectra, 196

Polyaromatic sulfur compounds, hydrodesulfurization, 345-349, 466-467 catalysis, 350-353, 390-426 kinetics, 352, 363-364, 372, 427, 436438,44-443 limitations, 349, 435, 455-456 new approaches, 456-466 process schemes, 355-425 quantum chemical calculations, 425-434 reaction mechanism, 369-383, 384 sulfur species in middle-distillate oils, 364-366 systematic approach, 349-353 Polyethers, perfluorous, in biphasic catalysis, 497 Polymerization, in hydrodesulfurization, 439-440 Probe molecules, 119 NMR solid acidity studies, 139-140 acylium ions, 139, 160 aldehydes, 162-163 alkyl carbenium ions, 154-157 ally1 cation, 143-144 ammonia, 172-174 arenium ions, 150-154 carbonium ions, 157-160 chalcogenonium ions, 161-162 cyclopentenyl cations, 140-143 indanyl cations, 144-147 ketones, 162, 163-165 nitrogen-containing compounds, 165-170 phosphines, 170, 172 trityl cation, 147-150 Promoted MoSJaluminum oxide catalysts, 392 Propadiene, vibrational spectra, 206-207 Propane cracking in zeolites, 104 diffusion in silicalite, 34-35 diffusion in zeolites, 36 sorption on zeolites, 71 Propene CAVERN device, 138, 141 hydroformylation, 484 hydroformylation of, 498 1,3-label scrambling in, 156 Propionylium ion, chemical shift tensors, 124

515

INDEX Propyl groups, as hydrocarbon surface species, vibrational spectra, 224-226 Propyne, vibrational spectra, 202, 204-206 Proton transfer, in bond activation of methanol in zeolites, 87-91 Pulsed field gradient (PFG) NMR, zeolite diffusion coefficients, 5, 47 Pyridine as probe molecule, 165 sorption on zeolites, 80

‘Q QEXAFS (quick EXAFS), 329, 330-340 Quantum mechanical (QM) simulations, bond activation, 2,84-107

R

Radial distribution function (RDF), 55 RAIRS (reflection-absorption infrared spectroscopy), 181, 296 benzene, 249 but-2-yne, 202, 204 cyclohexane, 229, 230, 231, 233 cyclohexene, 240 cyclopentane, 235-236 hydrocarbon surface species butyl groups, 226 methyl groups, 218 propyl groups, 224-226 vinyl groups, 22E methane, 209 neopentane, 212 propadiene, 206-207 toluene, 263 xylene, 264-265 Raman spectroscopy, 296 Rhodaoxetane-BDPI’ species, 491 Rhodium, ethyne spectra on finely divided metal, 198-199 Rhodium complexes, as asymmetric hydrogenation catalyst, 490 Ring effect, 16-17, 10 Ring-opening methathesis polymerization (ROMP), 492-493 Rotational diffusivity, 29 Ruthenium complex catalyst, hydrogenation of, 489

Ruthenium sulfide, as desulfurization catalyst, 458 S

SBMS, 417 Selenonium ion, 161 SFG (sum-frequency generation), 296 Shallow-bed CAVERN device, 138, 140 Shell higher olefins process (SHOP), 497-498 Shielding constant, 122 Silicalite diffusion in argon, 18-19 aromatic compounds, 41 hydrocarbons, 34-40 methane, 20, 23, 28, 29, 30, 33 xenon, 11-14 sorption on alkanes, 71 benzene, 81-83 hydrocarbons, 70-80 methane, 66-70 xenon, 55-56 Silver, ethyne spectra on finely divided metal, 199-201 Simulations, 1 bond activation by zeolites, 2, 84-107 zeolite adsorption, 2, 50-84 zeolite diffusion, 2, 3-50 Single atoms adsorption on zeolites, 53-62 diffusion in zeolites, 8-19 Single crystal spectral studies but-2-yne, 202-203 ethane, 210 ethyne ambient and higher temperatures, 189-191 low-temperature, 183-189 higher acyclic alkanes, 212-213 higher acyclic alkynes, 202-205 methane, 207-210 neopentane, 212-213 propyne, 202, 204-205 Single crystal studies, ethene on platinum, 273-283,297-298 Skateboard intercage migration, 45-46 Sodalite, methanol adsorption in, 91

516

INDEX

Sodium cis-epoxysuccinate, hydrogenolysis of, 491 Solid acidity, NMR study, 115-120 chemical shift, 120-137 computational chemistry, 129-137 probe molecules, 139-174 sample preparation, 137-139 Structure, chemical shift, 129-133 Styrene hydroformylation of, 487 MAS NMR experiments, 144-147 Substituted benzenes, vibrational spectra, 263-266 Sucrose, telomerization with butadiene, 492 Sulfonated phenyl phosphites, ammonium salts of, 477 Sulfonated phosphanes, 476-478 Sulfonium ion, 161 Sulfur compounds, polyaromatic, see Polyaromatic sulfur compounds Sum-over-state (SOS) DFT method, 134 Superacids, 116-117

T Telomerizations, two-phase, 491-492 Tetradecene-1, hydroformylation of, 486, 487 Theta-1, structure, 40, 41 Thiophene, hydrodesulfurization, 351, 370, 371, 383, 405 Thiophenes bonding modes, 411,412 electron densities and bond orders, 429-431 Toluene disproportionation, 152 vibrational spectra, 263 Toluenium ion, chemical shift tensors, 150, 151 TF'D (temperature-programmed desorption) ethene, 273-282 ethyne, 185 xylenes, 264 Transition state theory (TST) diffusion in zeolites, 4, 6-7, 48 sorption on zeolites, 92 TRAPDOOR experiment, 166 1,3,5-Triaza-7-phosphaadamantane, 483

Trimethylchalcogenonium ions, 162 Trimethylenedimetallo groups, as hydrocarbon surface species, vibrational spectra, 228 Trimethyloxonium cation, 161 Trimethylphosphine, as probe molecules, 170, 173 Trimethylselenonium ion, 162 Triphenylphosphanes, polyether-substituted, 481 Triphenylphosphine monophosphonate, 479 Tris(2-hydroxyethyl)phosphane,482 Tris(o-phenylalkyl)phosphanes,sulfonation, 477 Trityl cation, as probe molecule, 147-150 TST, xenon diffusion in silicalite, 13-14 Two-phase catalysis, see Multiphase homogeneous catalysis

U Ultra-shallow-bed CAVERN device, 138, 141 UPS (ultraviolet photoelectron spectroscopy), ethyne, 189

V n-Valeraldehyde, production of, 498 VEELS (vibrational electron energy loss spectroscopy), 181, 295 benzene, 243-248, 249,252,253, 260, 262. 263 but-2-yne, 202, 203, 205 cycloalkanes, 237-238, 239 cycloalkenes, 243 cyclohexane, 229, 231, 233 cyclohexene, 239, 240, 241 cyclopentane, 235-236 cyclopentene, 242 ethyne, 185, 186, 191 hydrocarbon surface species ally1 groups, 227 butyl groups, 226, 227 cyclohexyl groups, 228 cyclopropyl group, 228 ethyl groups, 221, 222, 224 methylene groups, 219-220 methyl groups, 215-217 rnethylidyne groups, 221

517

INDEX phenyl groups, 228 propyl groups, 224,225 vinyl groups, 22?, 228 methane, 208-209 propadiene, 207 toluene, 263 tnmethylenedimetallo groups, 227 Velocity autocorrelation function (VAF), 38 VGO (vacuum gas od), 355, 366 Vibrational spectra, 1.82, 300-301 acyclic alkanes, 207, 214, 299 butane isomers, 213 ethane, 210-212 methane, 207-210 neopentane, 212..213 acyclic alkenes, 264-267, 298 acyclic alkynes, 183, 299 but-2-yne, 202-204 ethyne (acetylene), 193-202, 298-299 propadiene (allene), 206-207 propyne, 202, 204-206 aromatic hydrocarbons, 266-267, 299 benzene, 245-263 substituted benzenes, 263-266 toluene, 263 xylene, 264-265 cycloalkanes, 244-245, 299 cyclohexane, 229-235 cyclopentane, 235-237 cyclopropane, 237-239 dimethylcyclohexanes, 239 methylcyclohexane, 239 cycloalkenes, 244-245, 299 cyclohexadiene, :243 cyclohexene, 239-242 cyclooctatetraene, 243-244 cyclopentadiene, 243 cyclopentene, 242-243 future priority areas, 297-299 hydrocarbon surfac:e species, 214 ally1 group, 227 butyl groups, 226, 227 cyclohexyl group, 228 cyclopropyl group, 228 ethene reactivity on platinum surfaces, 272-295 ethyl groups, 221 -224, 282-283 ethylidene groups, 224 HCCH group, 228

methylene groups, 219-220 methyl groups, 214-219 methylidyne groups, 220-221 phenyl group, 228 1-propyl and 2-propyl groups, 224-226 trimethylenedimetallo groups, 227 vinyl group, 228-229 techniques infrared techniques, 298-299 RAIRS, 296 Raman spectroscopy, 296 sum-frequency generation (SFG), 296 VEELS, 298 Vinyl group, as hydrocarbon surface species, vibrational spectra, 227-228 Vinyltoluenes, preparation, 494 VPI-5, 19

w Water diffusion in zeolite, 19-20 sorption on zeolite, 62 Water-soluble ligands, multiphase catalysis, 476-482

X Xenon diffusion in zeolites, 9-14 sorption on zeolites, 54-65 X-ray diffractometer, 324 XRD (X-ray diffraction), 317 catalysis characterization with absorption spectroscopy, 315-321, 340-342 combined EXAFS/XRD methods, 330-340 in situ, 321-324, 332-340 limitation, 317 technique, 324-327 Xylene, vibrational spectra, 264-265 Xylene isomers, diffusion in zeolites, 41-42, 44-45

L

Zeolite A methane sorption in, 64-65 xenon sorption on, 56-61

518

INDEX

Zeolite adsorption, simulations, 50-51, 83-84 aromatics, 80-83 combined Monte Carlolenergy minimization, 53 configurational-bias Monte Carlo method, 52-53 hydrocarbons, 70-80 methane, 62-70 Monte Carlo method, 51-52 single atoms, 53-62 water, 62 Zeolite diffusion, simulations, 2, 3-4 benzene and aromatics, 40-50 energy minimization, 7-8 hydrocarbons, 34-40 methane, 20-34 molecular dynamics, 4-6 single atoms, 8-19 transition state theory, 6-7 water, 19-20 Zeolite rho, xenon sorption on, 61-62 Zeolites, 1 catalysis acylium ions, 139, 160 aldehydes, 162-163 alkyl carbenium ions, 154-157

ally1 cations, 143-144 ammonia, 172-174 arenium ions, 150-154 carbonium ions, 157-160 chalcogenonium ions, 161-162 cyclopentenyl cations, 140-143 indanyl cations, 144-147 ketones, 162, 163-165 nitrogen-containing compounds, 165-170 phosphines, 170, 172 trityl cation, 147-150 CAVERN device, 138, 141 cluster calculations, 2, 84-107 reaction in pores of, 1-2 simulations, 2-3 adsorption, 2, 50, 85 bond activation, 2, 84-107 diffusion, 3-50 transport, 2 solid acids, 118-119, 120 Zero-point energy (ZPE) corrections, 134 ZK4, methane diffusion in, 25 ZSM-5 hydrocarbon sorption on, 79, 80 methane sorption on, 68