Advances in Catalysis, Volume 31

ADVANCES IN CATALYSIS VOLUME 31

Advisory Board

M. CALVIN

M. BOUDART Stunford, California

Berkeley, California

V. B. KAZANSKY Moscow, U.S.S.R.

G. A. SOMORJAI Berkeley, Californiu

A . OZAKI Tr,kyo, Japan

l? H. EMMETT Portland, Oregon

G.-M. SCHWAB Munich, Germany

R. UGO Milun, ltuly

ADVANCES IN CATALYSIS VOLUME 31

Edited by

D. D. ELEY The University Nottinghum, England

PAULB. WEISZ

HERMAN PINES Northwestern University Evunston, Illinois

Mohil Research und Development Corporufion Princeton, New Jersey

1982

ACADEMIC PRESS A Subsidiary of Harcourt Brace jovanovich, Publishers

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82 83 84 85

9 8 76 5 4 3 2 1

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

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

vii ix

Introduction . . . . . . . . . . . . . ........... .......... Oxidation Reactions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .......... Hydrogenation and Dehydrogenation ......... . . . .......... Oligomerization Reactions ................................ .......... Carbonylation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .......... Hydroformy lation . . . . . . . . . . . . . . . . . . . ... ... . . . . . . . . . . Methanation ................................ .......... Conversion of Synthesis Gas to Hydrocarbons . . . . . . . . . . . . . . .......... Miscellaneous Reactions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . General Conclusions and Future Prospects .. . . . . . . . . . . . . . . . . . . . . . . . . . References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

2 10 18 24 39 46 50 55 58 61 68

Nonacid Catalysis with Zeolites I . E . MAXWELL

I. I1 . 111.

I v. V. VI . VII . VIII . IX . X.

Characterization and Reactivity of Mononuclear Oxygen Species on Oxide Surfaces M . CHEA N D A . J . TENCH

I. I1. I11. IV. V. VI . VII .

..................................................... of 0 by EPR and Optical Spectroscopy . . The Formation and Stability of 0- Ions ............................... Aggregate 0- Species ....................... . . . . . . . . . . . . . Reactivity of the 0- Ion ................................... The Surface Oxide Ion 0:; . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Characterization and Reactivity of M=O . .. Appendix . The EPR Parameters of 0- Ions ............................ References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . ~

78 79 90 95 101 107 124 126 128

Sulfur Poisoning of Metals C . H . BARTHOLOMEW. F? K . AGRAWAL. AND J . R . KATZER

I. I1 . 111.

Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . The Nature of Metal-Sulfur Bonds ................................... Sulfur Adsorption on Metals ......................................... V

136 137 143

vi

CONTENTS

IV.

V. VI . VII .

Effects of Sulfur on Adsorption of Other Molecules .................... Effects of Sulfur on Catalytic Activity and Selectivity Properties of Metals . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Regeneration and Hydrogenolysis . . . . . . . . . . .. . . . . . . . . . . . . . . . . . . . . . . . . Conclusions and Recommendations . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

172 187 229 233 235

Methanol Synthesis K . KLlER I. I1 .

Introduction ........................................................ Catalyst Selection .............. per Metal . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . The Activity of Pure Zinc Oxide ....................... The Binary Copper-Zinc Oxide System . . . . . . . . . . . . . . . . . Other Binary Catalysts . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Ternary and Quaternary Catalysts . . . . . . . . . . .. . . . . . . . . . . . . . . . . . . . . . . . . References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

243 251 254 257 257 287 291 296 310

AUTHOR INDEX ............................................................... SUBJECT INDEX ............................................................... CONTENTS OF PREVIOUS VOLUMES ...............................................

315 334 345

I" .

I v. V. VI . VII . VIII .

Contributors Numbers in parenrheses indicaie the puges on which the uuihors’ contributions begin.

I? K. AGRAWAL, School of Chemical Engineering, Georgia Institute of Technology, Atlanta, Georgia 30332 (135) C . H . BARTHOLOMEW, Department of Chemical Engineering, Brigham Young University, Provo, Utah 84602 (135) M. CHE,Laboratoire de Chimie des Solides, E R 133, C N R S , Universitk de Paris VZ, 75230 Paris-Cedex 05, France (77) J. R. KATzEn, Center f o r Catalytic Science and Technology, Department of Chemical Engineering, University of Delaware, Newark, Delaware 19711 (135) K . KLIEn, Department of Chemistry, Center f o r Surface and Coatings Research, Lehigh University, Bethlehem, Pennsylvania 18015 (243) I. E. MAXWELL, KoninklijkelShell-Laboratorium, Amsterdam, Shell Research B . V., Amsterdam, The Netherlands (1) A. J. TENCH,Chemistry Division, Atomic Energy Research Establishment, Harwell, Oxfordshire OX11 ORA, United Kingdom (77)

vii

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Preface A significant part of the effort in catalysis research is based on an understanding that a more complete knowledge of the nature and behavior of reactive intermediates or chemisorbed complexes will enable us to produce better catalysts for industrial processes. The supporting cash flow is certainly based on this assumption. This volume of Advances in Catalysis presents thoroughly written accounts of four different sectors of this effort. The opening article by I. E. Maxwell covers the very interesting broad range of developments that are resulting from the incorporation of metal ions into the familiar X- and Y-type zeolites, and the more recent revolutionary shape-selective narrow pore types. It is perhaps fanciful to speculate that if the walls of these pores had some controlled flexibility, we should have a pretty good approximation to an enzyme-active site. Next, M. Che and A. J. Tench show how electron paramagnetic resonance in particular has revolutionized our knowledge of 0-,on both simple oxides and the mixed oxides with a special function in catalytic oxidation. A later article will deal with more complex species, such as 0;. The article by C. H. Bartholomew, P. K. Agrawal, and J. R . Katzer summarizes our current knowledge of sulfur poisoning on metals, on the firm basis provided by recent low-energy electron diffraction studies of sulfur-metal surface structures. I believe that the late E. B. Maxted of Bristol, who pioneered work in this area, would have been delighted with present progress. Finally, K. Klier’s report on methanol synthesis once again reveals the large strides being made by the use of the latest physical techniques. In addition, he shows how absolute rate data-by allowing arguments based on entropies of adsorption and activation-can further mechanism studies in what is, after all (at least to this writer), a relatively complex industrial reaction.

D. D. ELEY

ix

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A D V A N C E S I N C A T A L Y S I S . V O L I I M k 31

Nonacid Catalysis w i t h Zeolites I . E . MAXWELL Koninklijke ~SlwllLtrhoratoriimni. Ani.or~rcltrrn ShcTll Rr.wrri~hB . V. Arn.vicrdmw, Thc Ni~tlir~rlcmdc

I . Introduction . . . . . . . . . . . . A . Scope . . . . . . . . . . . . B . Zeolite Composition and Structure . . . . C . Zeolite Ion Exchange and Acidity . . . . D . Zeolite Cation Siting . . . . . . . . E . Diffusion in Zeolites . . . . . . . . . . . . . . . . . I1 . Oxidation Reactions A . Carbon Monoxide Oxidation . . . . . . B . Alkane Oxidation . . . . . . . . . C . Alkene Oxidation . . . . . . . . . D . Oxidation of Nitrogen-Containing Compounds E . Oxidation of Sulfur-Containing Compounds . F . Miscellaneous Oxidation Reactions . . . . I l i . Hydrogenation and Dehydrogenation . . . . . A . Hydrogenation . . . . . . . . . . B . Dehydrogenation . . . . . . . . . IV . Oligomerization Reactions . . . . . . . . A . Ethylene . . . . . . . . . . . B . Acetylene . . . . . . . . . . . C. Propylene . . . . . . . . . . . D. Cyclopropenes . . . . . . . . . . E . Butadiene . . . . . . . . . . . F. n-Butene . . . . . . . . . . . G . Isobutylene . . . . . . . . . . . V . Carbonylation . . . . . . . . . . . A . Methanol . . . . . . . . . . . B . Ethanol and Higher Alcohols . . . . . . C . Ethylene . . . . . . . . . . . VI . Hydroformylation . . . . . . . . . . A . Cobalt Zeolites . . . . . . . . . . B . Rhodium Zeolites . . . . . . . . . VII . Methanation . . . . . . . . . . . . A . Palladium Zeolites . . . . . . . . . B . NickelZeolites . . . . . . . . . . C . Ruthenium Zeolites . . . . . . . . VIII . Conversion of SynthesisGas to Hydrocarbons . .

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2 2 3 5 6 9 10 10 13 14 16 17 18 18 19 21 24 24 29 30 31 32 36 37 39 39 45 46 46 47 49 SO 51 52 53 55

I Copyright s 1982 by Acadcrnic Press. Inc All rizhiz ofrepr, ~ i u c t i o ni n a n y form rrservcd ISBN n- I 2-no7x31-7

2

I. E. MAXWELL

Miscellaneous Reactions . . . . . . A. Water-GasShift . . . . . . . B. Kolbel-Engelhardt Reaction . . . . C. Water Splitting. . . . . . . . X. General Conclusions and Future Prospects . A. Activity and Selectivity . . . . . B. Stability. . . . . . . . . . C. Active Species and Reaction Mechanisms D. Future Prospects . . . . . . . References . . . . . . . . . .

IX.

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58 58 59 59 61 61 62 64 66 68

Introduction

A.

SCOPE

In recent years there have been considerable academic and industrial research efforts carried out in the field of zeolite catalysis. The main part of this work has, however, been related to reactions where the zeolite is used as a solid acid, e.g., isomerization, cracking, hydrocracking, etc. This is hardly surprising since this area represents the major application to date of zeolites as industrial process catalysts. Reviews on zeolite catalysis have therefore also in general tended to concentrate more on the work related to acid catalysis ( I -7a). In principle, zeolites offer considerable scope for surface modification. The ion-exchange properties and the periodicity (crystallinity) of zeolites, for example, enable functional metal ions or complexes to be stabilized in a variety of oxidation states and coordination geometries on specific surfaces in a well-defined pore structure. This represents a valuable starting point in the design of catalytic surfaces. Although these characteristics of zeolites have been recognized for some time, only a relatively small effort has been carried out in this area. This is probably due to the remarkable success of zeolites in solid acid catalysis. More recently, however, there has been renewed interest in the applications of zeolites in nonacid catalysis. In particular, zeolites have been successfully applied to reactions such as oxidation, reduction, olefin oligomerization, carbonylation, hydroformylation, synthesis gas conversion, and water splitting. The purpose of this review is to discuss the above applications in some depth with the intention of comparing where possible the advantages or

3

NONACID CATALYSIS WITH ZEOLITES

disadvantages of zeolites with alternative, perhaps more conventional catalysts for these reactions. A greater emphasis is given to the most recent and more novel applications of zeolites to heterogeneous catalysis. Attention is also focused on probable reaction mechanisms and active sites. A final section is included where the unique features of the zeolites for nonacid catalysis are highlighted and future research prospects in this area are evaluated. The technical literature, scientific publications, and patent literature have all been searched-mainly covering the period from 1972 up to and including the first half of 1980.

B.

ZEOLITE COMPOSITION AND STRUCTURE

A very brief discussion of the structure and chemistry of zeolites relevant to this review is now given. For a more comprehensive discussion of this subject the reader is referred to an excellent book by Breck (8). Zeolites are crystalline aluminosilicates (synthetic and natural), with a chemical composition which corresponds to the general formula

where M"' is the cation which balances the negative charge associated with the framework aluminum ions. The framework ions, Si4+ and A13+, are each tetrahedrally coordinated to four oxygen anions. The periodic threedimensional network, which is so characteristic of zeolites, is formed by linking the (SiO,) and (A10,)- tetrahedra through shared oxygen ions. These tetrahedra tend to form rings, containing from four to twelve tetrahedral units. Such rings normally form the entrances to channels or cages in

TABLE 1 Composificinund Pore Purumefrrs of Sonic Zeoliics

Type

Unit-cell composition

Zeolite A Zeolite X Zeolite Y Mordenite

Na,,(A102),z(Si0,),, Na,,(AIO,),,(SiO,),,,, Na,,(AIO,),,(SiO,),,, Na,(AlO2),(Si0,),,,

~

I A = 1 O - I nm. From Ref. 2.

Void volume (mIjmI)

Pore diameter

(Ay

Thermal decomposition temp.' ( 'C)

Si/AI ratio

0.47 0.50 0.48 0.28

4.2 7.4 7.4 6.7 x 7.0

700 772 793 I000

1 .0 1.23 2.43 5.0

4

I. E. MAXWELL

A

FIG.1. Schematic diagram of zeolite A. (Reproduced from Ref. 2 with permission from the author and Elsevier Scientific Publishing Company, Amsterdam )

zeolites and thus define the pore diameter for a particular structure. The composition pore volumes and pore diameters for a number of common zeolites relevant to this review are given in Table 1. Zeolites A, X , and Y all consist of tetrahedra linked to form cubooctahedra or so-called sodalite cage units. When these units are linked through four-membered rings zeolite A is formed (see Fig. l), whereas linking via the six-membered rings results in zeolites X and Y (see Fig. 2). The latter two zeolites only differ in the Si/Al ratio. Mordenite has a channel-like pore structure in which the basic building blocks consist of five-membered rings. A view of the mordenite structure perpendicular to the main channels is shown in Fig. 3. In general, zeolites have good thermal stability, but this is further improved by increasing the Si/Al ratio as shown in Table 1. The hydrothermal stability of zeolites also increases with decreasing aluminum content. F

FIG.2. Schematic diagram of zeolites X and Y . (Reproduced from Ref. 2 with permission from the author and Elsevier Scientific Publishing Company, Amsterdam.)

NONACID CATALYSIS WITH ZEOLITES

5

Frc. 3. Section through the mordenite structure (Z) perpendicular to the main channels. (Reproduced from Ref. 2 with permission from the author and Elsevier Scientific Publishing Company. Amsterdam.)

A N D ACIDITY C. ZEOLITEON EXCHANGE

The framework charge-compensating cations in a zeolite, which for synthetic zeolites are normally sodium ions, can be exchanged for other cations of different type and/or valency. However, care must be taken during ion exchange to avoid strongly acidic solutions which can lead to proton exchange with the zeolite metal cations or even structure collapse. For example, zeolites A, X, and Y decompose in 0.1 N HCI. The more silica-rich zeolites such as mordenite are, however, stable under such conditions. Acidity can be introduced into a zeolite in a number of different ways: (1) ion exchange with NH:,

followed by thermal decomposition, i.e.,

NHiZ-

+

H*Z-

+ NH,(q)

(2) hydrolysis of ion-exchanged polyvalent cations followed by partial dehydration, i.e., M(H,O):+Z-

+

M(OH)'"- " + Z -

+ H + Z - + (x - I)H,O

(3) direct proton exchange, i.e., Na+Z

+ H+

-+

H'Z-

+ Na'

(4') reduction of exchanged metal ions to a lower valency state, i.e., M"+Z- + :H, + M("-')+Z- + H + Z -

The so-formed proton or Bronsted acid sites can be further dehydroxylated

6

1. E. MAXWELL

to form Lewis acid sites, i.e., H

\Si/O\Al / \ / \

0

0 0

H

Si 0 0 ’ ‘0

0 ’

‘0

o/

‘0

As will be seen in the ensuing discussion, for most of the catalytic reactions discussed in this review, surface acidity leads to undesirable side reactions. In these cases care must be taken during catalyst preparation to avoid the introduction of acidity via any one of the above-described routes.

D.

ZEOLITE C A T I O N SITING

In most of the reactions discussed the active entity of the zeolite catalysts is introduced via ion exchange. Thus a knowledge of the possible siting of cations is a prerequisite for an understanding of the location and nature of the active sites in zeolites. In this respect the periodicity of the internal surface of the zeolites provides an almost unique opportunity to study the surface composition in considerable detail using powerful analytical methods such as X-ray diffraction. In zeolite A the six-membered rings (eight in total) facing onto the main cavity are the preferred cation sites (see Fig. 1). If there are more than eight cations per unit cell, the remainder generally lie in the plane of the eightmembered ring entrance to the large cavity, but are displaced off-center so that there is an asymmetrical ionic interaction with the oxygen anions. In recent years Seff and co-workers (Y) have extensively studied cation siting in zeolite A using single-crystal X-ray diffraction techniques. In favorable cases these workers have also been able to obtain detailed information on the interactions between cations and absorbate molecules. Two examples are shown in Fig. 4, where the adsorption complexes formed when acetylene and NO are adsorbed in Co(I1)A have been resolved. In the former case it is proposed that a weak complex is formed via an induced dipole interaction with the polarizable n-orbitals of the acetylene molecule. For the NO complex there is good evidence for electron transfer resulting in a complex between CO(II1) and NO-. In both cases the organic molecules

NONACID CATALYSIS WITH ZEOLITES

7

FIG.4. (a) Complex formation between Coz+ and acetylene in zeolite A ; (b) complex formation between NO and C o 2 + in zeolite A. (Reprinted with permission from Ref. Y. Copyright 1976 American Chemical Society.)

interact with the cations sited at the six-membered rings. The interaction results in a displacement of the cation toward the large cavity to give a more tetrahedral-like coordination with the organic fragment. Although zeolite A is less interesting than larger pore diameter zeolites from the point of view of catalysis (pore diameter is 4.2 A), these model studies d o provide some insight into the types of cation adsorbate interactions, which may also occur in larger pore zeolites. Cation siting in zeolites X and Y is more complex due to the larger number of potential sites. In general, most of the cations lie along the crystallographic threefold axes of the cubic unit cell. Various positions along these axes have been arbitrarily given Roman numbers as shown in Fig. 5. Only the site I1 cations on these axes, which are located at the six-membered rings (cf. zeolite A), are able to interact with adsorbate molecules present in the supercage. Access to the smaller sodalite cages is severely limited by the small diameter (2.2 A) of the six-ring entrance. An additional site (so-called 111) has been identified in the supercage, near the four-membered rings which

I . E. MAXWELI

FIG. 5 . Faujasite framework and cation siting. (Roman numerals indicate cation sites; arabic numbers show oxygen-atom numbering scheme.)

link adjacent sodalite cages (see Fig. 5). There is good evidence that ions located at these sites are the active entities for hexane dehydrocyclization over TeNaX (10) and butadiene cyclodimerization over CuNaY (11). The location of site 111 at the entrance to the supercages, together with the unsaturated coordination geometry of the cations, provides an ideal sile for interaction with adsorbate molecules. The siting of cations in mordenite is generally less well understood than that in the zeolites described above. Smith and co-workers (12) have, however, in recent years carried out a number of single-crystal X-ray analyses on various cation-exchanged forms of mordenite. These workers correctly emphasize, however, that the cation population densities are subject to unknown errors due to pseudosymmetry. The alkali metal ions are distributed over four major sites, namely: site I, at site 11, in site IV, at site VI, at

the end of the side pocket off the main channel; the side pocket at the center of an eight-ring; the junction of the side pocket with the main channel; the wall of the main channel.

Only Na' ions were found in site I due to space restrictions.

9

NONACID CATALYSIS WITH ZEOLITES

Alternative methods to ion exchange may be employed to introduce active metal components into zeolites such as pore volume impregnation and vapor-phase adsorption of volatile compounds. In these cases the siting of such species in the zeolite pore structure is generally less well defined.

E.

DIFFUSION I N ZEOLITM

In view of the importance of diffusion in zeolites it would be an omission not to include a brief discussion on this subject. This is perhaps the least well understood, but it is often directly related to the unique catalytic properties of these materials. Weisz (13) has lucidly described diffusion phenomena in zeolites using a plot of diffusivity against pore size, as shown in Fig. 6. Zeolites with pore diameters in the range of 4-9 8, are shown to provide a region of diffusivity beyond the regular and Knudsen regions, which Weisz has termed the configurational regime. This is the region, where molecules D

(c m2/ sec) I

lo-2

REGULAR

lob

CONFIGURATIONAL

do I6l2 10"

J I

10

I

I

loo

lo00

v (ANGSTROMS)

I I

I 10

(pm)

PORE SIZE

FIG.6 . Diffusivity D and size of aperture (pore). (Reprinted with pcrmission from Ref. 13. Copyright 1973 American Chemical Society.)

10

1. E. MAXWELL

must diffuse through spaces of near-molecular dimensions and is thus of considerable importance in shape-selective catalysis. As shown in Fig. 6, the configurational region spans an enormous range of diffusivities (approximately ten orders of magnitude). In this region subtle differences in configurational structure can have a large effect on diffusivity. For example cis- and trans-butene have diffusion coefficients which differ by two orders of magnitude in CaA (14). Even more subtle effects are possible such as a periodic variation in diffusivities with increasing carbon number for n-paraffins in erionite (15). Thus it is evident that zeolites offer considerable potential for steering reaction selectivity on the basis of differences in molecular shape. These possibilities extend far beyond the more familiar molecular sieve effects where bulky molecules are simply excluded entry into the zeolite cavities due to pore diameter restrictions. The foregoing discussion refers solely to intraparticle diffusivity (micropore diffusion) as distinct from interparticle effects (macropore diffusion). Since a practical zeolite catalyst will consist of composite particles, each containing a large number of individual zeolite crystals, it is important to make a clear distinction between these two types of diffusion. In some cases macropore diffusion may be important in determining the overall reaction kinetics but will obviously not introduce or affect shape selectivity in any way. Although diffusivity is often important in zeolite catalysis, other factors may also be crucial in determining shape selectivity. Recent work by Post (15a), for example, has shown that the shape selectivity behavior observed for the relative cracking rates of hexane isomers over H-ZSM 5 zeolite (see Section VlII) could not be understood on the basis of their measured diffusivities. Spatial restrictions imposed on transition-state species formed within the zeolite pores provide a possible explanation for the observed results.

II. A.

Oxidation Reactions

CARBON MONOXIDE OXIDATION

The oxidation of CO is widely used as a test reaction for oxidation catalysts because of its simplicity. Thus, there is quite an extensive literature on CO oxidation using various zeolite catalysts. The parent (sodium forms) of zeolites show very little oxidation activity as might be expected and therefore the majority of the studies have concentrated on transition metal ionexchanged forms.

NONACID CATALYSIS WITH ZEOLITES

11

It is of particular interest to compare the relative activities of transition metal ion-exchanged zeolites with, for example, the corresponding oxides in order to gain some insight into the influence of the zeolite lattice. Boreskov rt al. (16) compared the specific activities of CuY and CuO for CO oxidation calculated on the basis of surface copper concentrations. Although the specific activity of CuY increases with increasing copper level, even at 16 wt.% Cu the activity is 2.5 orders of magnitude less than that of CuO (see Fig. 7). A similar behavior has also been demonstrated (17, 18) for Fe, Ni, Co, and Cr ions exchanged into zeolite Y and their corresponding oxides (i.e., a-Fe,O,, NiO, Co,O,, and Cr,O,, respectively). In addition, the activation energies for the transition metal ion-exchanged zeolites are considerably higher than those for the corresponding oxides ( I 7, 18) (e.g., CuY, 19.5 kcal/mol'; CuO 13 kcal/mol; NiY, 26 kcal/mol; NiO, 9.5 kcal/mol ; COY, 16 kcal/mol ; Co,O,, 1 1 kcal/mol). Iron-containing zeolites are somewhat expectional in that the specific activity of FeY is independent of the iron content (18, 1 Y ) (in the range 0.4-5 wt.%) and although less active than a-Fe,O,, the activation energies for the zeolite and oxide catalysts are the same ( 1 8 kcal/mol). It has also been shown (20)that for transition metal ion-exchanged zeolites X and Y, the activities for CO oxidation increase exponentially with in5 -

I

f"

cuo

4 -

2

6

10

14 %

tu-

FIG.7. Comparison of specific activity of CuY and CuO for CO oxidation at 450°C. (Reproduced from Ref. 16 with permission from the authors and Plenum Publishing Corporation, New York.)

' 1 kcal = 4.2 k J .

12

I . E. M A X W t L L

creasing metal ion standard oxidation potential (i.e., Cu+ > Fez+ > Cr3+ > CuZf z Ni2+). More recently it has been shown (21-24) that the equilibrated pH of the transition metal ion-exchange solution is also critical in determining the specific activity of zeolite catalysts. The results obtained for the CuY system (21)are shown in Fig. 8. The influence of both transition metal ion loading and ion-exchange solution pH has been attributed to the formation of catalytically active metal oxygen bridge species within the zeolite cavities (i.e., M"+-02 --M"+ where the anion corresponds to extralattice (i.e., nonzeolite framework) oxide ions. The formation of such species would be expected to be favored by high metal loadings and hydrolysis conditions during ion exchange, as is observed. In fact there is some direct evidence from ESR (21),Mossbauer spectroscopy ( 2 5 ) , IR spectroscopy (26), and magnetic measurements (23) to support the existence of these oxygen-bridged species. Boreskov (18) has proposed a model for transition metal compounds in which the rate of oxidation is assumed to be determined by the rate of electron transfer between oxygen and the transition metal ion. This process is further assumed to be facilitated with increasing degree of covalency of the metaloxygen bond. Thus the more covalent transition metal oxides are more active than the rather ionic metal ion-exchanged zeolites. The oxygenbridged species as described above is considered to be more covalent in character, and hence more active for oxidation catalysis than the transition tog (SPECIFIC ACTIVITY) cm3/(sec otorn Cu)

4

6

8

PH

FIG.8. Specific activity of CuY zeolite as a function of ion-exchange solution pH. (Rcproduced from Ref. 21 with permission from the authors and Plenum Publishing Corporation, New York.)

NONACID C'ATAI.YSIS WITH ZEOLITES

13

metal ion-framework oxygen interaction. However, this model has not been sufficiently developed to provide a detailed mechanistic scheme which would explain the observed reaction kinetics (16, 20,23),which are generally first order in carbon monoxide and fractional or zero order in oxygen concentration. Beyer ct a/. (27) carried out a detailed kinetic study on CO oxidation over CuNaY zeolite after various pretreatments. The results obtained were explained in terms of the relative concentrations of the species C u 2 + , Cu', and Cuo present in each catalyst. Paetow and Riekert (28, 2Y) in a careful study have compared the relative activities of Cu2+-exchanged zeolite T and mordenite with various coppercontaining compounds. On the basis of turnover numbers per CO chemisorption site the Cu2+-exchanged zeolites are 2-4 orders of magnitude less active than CuO, CuMn,O,, and CuCr,O,. This was considered to be consistent with the involvement of lattice oxygen as a n intermediate which is easier to remove in oxides than zeolites. An interesting application of zeolite-catalyzed CO oxidation has developed pertaining to fluid cracking catalysts (FCC). This was based on the discovery by Chen and Weisz (30) that minute amounts (0-100 ppm) of platinum-group metals incorporated into zeolites produced catalysts which were highly active for CO conversion under FCC regeneration conditions. It was further demonstrated that these oxidation promoters could be incorporated directly into the zeolite FC catalysts without any adverse effects on cracking selectivity (31). This promoter technology has now been fully developed and is currently successfully applied in commercial practice. Numerous patents (32) have appeared in recent years relating to this particular application. It would be of interest, particularly in view of the low metal loadings, to determine to what extent the zeolite matrix contributes to the exceptional oxidation activity of these catalysts. To date, the appropriate studies do not appear to have been carried out.

B.

ALKANE OXIDATION

Metal ion-exchanged forms of zeolite X are active catalysts for the complete oxidation of methane to carbon dioxide and water (33,34). In general the platinum metal ion-exchanged forms (e.g., Pt, Pd, Ir) are considerably more active than the first-row transition metal ion forms (e.g., Cu, Mn, Cr, Fe, Ni). The kinetics were best described as first order in methane and zero order in oxygen. There is no general agreement on the mechanism and neither (33, 34) has the possible existence of bridged-oxygen species been considered.

14

1. E. MAXWELL

The oxidative dehydrogenatioii of cyclohexane to benzene has been studied more extensively. Transition metal ion-exchanged forms of zeolite Y have been shown (34-39) to be particularly active catalysts for this reaction. Although the platinum metal ions exhibii the highest activity, CuY was found to be the most selective for benzene formation (38,39). Moshida et ul. (40) carried out a comparative study of cyclohexane oxidative dehydrogenation using Cu'Y (prepared via Cu+ exchange in liquid ammonia) and Cu2+Y. Both catalysts exhibited quite good benzene selectivities ( > 90%), but interestingly showed different reaction kinetics. The reaction orders in oxygen were unity and one-half for Cu2'Y and Cu' Y , respectively. This was interpreted in terms of the active species being molecular and dissociatively adsorbed oxygen, respectively. Such a marked difference in mechanism for benzene formation does not seem very probable. The different kinetics may simply arise due to a change in the rate-determining step for oxygen activation. The possible role of bridged-oxygen species has also been neglected, which is likely to be particularly important for Cu'Y under oxidation conditions. Minachev et UI.(41, 42) have recently examined alkali metal ion forms of various zeolites (A, X, Y, L, chabazite, erionite, and mordenite) for cyclohexane oxidative dehydrogenation. Not surprisingly these alkali metal ion forms are considerably less active than those containing transition metal ions (reaction temperatures of approximately 300" and 450"C, respectively). Further, cyclohexene rather than benzene is the predominant product (selectivity to cyclohexane 67-84"/,), particularly with small-pore zeolites. In fact, NaA was the most active zeolite tested ( 4 2 ) ,which strongly suggests that the reaction is simply occurring on the outer surface of the zeolite crystallites.

OXIDATION C. ALKENE Transition metal ion-exchanged zeolites are active catalysts for alkene oxidation but generally result in deep oxidation to carbon dioxide and water (43-45). In common with CO and alkane oxidation, the platinum metal ions are more active than the first-row transition metal ions. Mochida et al. (43) have been able to correlate the catalytic activity of ion-exchanged Y zeolites for propylene oxidation with a so-called Y parameter as shown in Fig. 9. This parameter was considered to express the tendency of the metal ion toward the formation of a dative n-bond with propylene. Further, it was shown that with increasing Y factor there was a decrease in reaction order, which was considered evidence of increased propylene adsorption. In a more recent study of CuX zeolites, Gentry et ul. (45) found some evidence

NONACID CATALYSlS WITH ZEOLITES

15

c .-

FIG.9. Correlation between the catalytic activity and the parameter Y in propylene oxidation. Y = 10(In/In+l ) ( r / n l ' z ) ,where I, is the nth ionization potential, Y is the ionic radius of the metal ion, and n is its formal charge. (Reproduced from Ref. 43 with permission from the authors.)

for the participation of Bronsted acidity as well as Cu2+ ions in the total oxidation of propylene. A rather interesting application of zeolite-based alkene oxidation catalysis has been demonstrated by Japanese workers (46, 47). In particular, a P d 2 + , C u 2 + Yzeolite was shown to be an active and stable heterogeneous oxidation catalyst which is analogous to the well-known homogeneous Wacker catalyst system containing PdC1, and CuC1, (48). Under Wacker conditions (i.e., alkene/O,/H,O) the zeolite Y catalyst was shown to convert ethylene to acetaldehyde and propylene to acetone with selectivities in excess of 90% with CO, as the major by-product. The reaction mechanism was also shown to parallel that of the homogeneous system, where Pd2+ is reduced to Pdo and is catalytically reoxidized by Cu2., i.e.,

+ C,H, + H,O 2(H+Z-)] + 2[CuZf,22-1 2[Cu+Z-, H + Z - ] + to,

[Pdz+,22-1 [Pd',

+

-+

-+

+ [Pd', 2(H+Z-)] [Pdzf, 22-1 + 2[CuiZ-, H'Z-1 2[Cu2+,22-1 + H,O

CH,COCH,

The most active catalysts were obtained by preexchange with C u 2 + followed by postexchange with P d 2 + . This would seem to indicate the importance of cation siting in promoting electron transfer which was assumed to occur via adsorbed water molecules which bridge PdO and Cu2+ through a hydrogen bond. In principle, the zeolite catalyst system would offer advantages over the existing homogeneous catalyst, particularly with respect to corrosion due to the absence of HCl and chlorine-containing by-products. However, acetaldehyde and acetic-acid production via ethylene has recently become less economically attractive compared to methanol carboxylation chemistry.

16

I. E. M A X W E L L

Further, acetone via Wacker chemistry must compete with 2-propanol dehydrogenation and coproduction in the Hock phenol process (48). I t remains, however, as a very interesting example of selective oxidation using mixed transition metal ions in a zeolite matrix. The extent to which the zeolite matrix directly or indirectly facilitates the electron transfer step between Pd' and Cu2+ has not been examined. This would seem worthy of study within the general concept of electron transfer reactions in heterogeneous catalysis. In the absence of Pd2+ but in the presence of steam, Mochida et ul. (4Y, 50) showed that propylene could be oxidized over C u 2 + Y to yield a mixture of products such as 2-propano1, acetone, and acrolein. D. OXIDATION OF NITROGEN -CONTAINING COMPOUNDS The oxidation of N O and NO, is of industrial importance for cleaning combustion-flue gases. Transition metal ion-exchanged zeolites have been shown (51) to be highly active catalysts for this reaction. The relative activities are shown in Fig. 10, from which it can be seen that equilibrium conversions of NO to NO, can be achieved with C u 2 + X at reaction temperatures as low as 350°C. Kinetic studies showed that the reaction rates were fractional order in both NO and 0,. The following reaction mechanism was therefore proposed, for example, with C u 2 + X ,

This mechanism would seem to be quite plausible since there is good evidence for the formation of both nitrosyl (51-55) and oxygen-bridge complexes (16-26) with transition metal ions in zeolites. Interestingly, for a C r 3 + Y catalyst (51), the activity was enhanced by H,O and there was no deactivation by SO,, indicating that such a catalyst would likely be quite robust under practical feed conditions. Williamson et al. (56) showed that C u 2 + Ywas an active catalyst for the oxidation of NH, to N, and H,O. A mechanism was proposed involving the intermediate formation of an amine complex [Cu(NH,),]'+. The N H , reduction of Cu2+ and Cu' in this complex was proposed as the slow step with reoxidation via 0, being very rapid. This mechanism was consistent with the kinetic expressions which were shown to be first order in N H , and zero order in 0,.

NONACID CATALYSIS WITH Z;EOLITt!S 100-

.

*

.\.

801 .c

17

Loaded on NaX

‘-,Equilibrium

I

I

1

600

700

I

Reaction temperature (K)

FIG. 10. Activities of transition metal ions loaded on zeolite X (space velocity 23,400 cm3(g catalyst)- hr- NO 300 ppm, 0 , 6 . 4 vol”/,). (Reproduced from Ref. 51 with permission from the authors and the American Chemical Society.)

’

’,

E. OXIDATION OF SULFUR-CONTAINING COMPOUNDS Mars and co-workers (57-60) have systematically investigated the behavior of various porous materials including zeolites toward the catalytic oxidation of H,S with 0, to elemental sulfur. This process is attractive for H,S removal from gas streams with low H,S contents. These workers showed that the product sulfur, which is adsorbed in the pores, is the catalytically active entity, i.e., the reaction is autocatalytic. Sulfur radicals proved to be the active sites for oxygen chemisorption. The pore diameter of the support was shown to be important in determining the catalytic activity. The catalytic sulfur was found to be most active in supports with pore diameters in the range 5-10 A. Thus NaX was found to be a highly suitable support which was competitive with industrial active charcoals with regard to both activity and selectivity towards sulfur formation. More recent mechanistic studies (61) on H,S oxidation over NaX and NaY have confirmed the previous findings of Mars and co-workers (57-60). However, significant SO, formation was also observed, particularly at reaction temperatures in excess of 70 C. Complete oxidation of H,S to SO, can be achieved by incorporating a metal function into the zeolite. For example, vanadium on mordenite has been claimed (62) to be a very effective catalyst for this reaction. Pearce and Lunsford (63) demonstrated that M n 2 + Y was an active catalyst for SO, oxidation at ambient temperature. The product SO:- was,

18

I. E. MAXWELL

however, strongly adsorbed in the zeolite pores, poisoning the catalytic activity. Interestingly, there was quite good evidence that hydrated Mn" ions sited in the supercage were the catalytic entities.

F.

MISCELLANEOUS OXIDATION REACTIONS

Transition metal ion-exchanged zeolites X and Y are active catalysts for hydrogen oxidation (64, 65). There is in fact a close parallel in terms of relative activity and kinetics between hydrogen oxidation and carbon monoxide oxidation over these catalysts. The platinum metal ion-exchanged forms are the most active, followed by the first-row transition metal ions. The reaction rates are first order in hydrogen and close to zero order in oxygen (65). Further, there is evidence, particularly for Cu2+Y (65), that Cu2+-0-Cu2+-bridge species are the active sites. In this case the zeolite catalyst is also intrinsically less active than the corresponding oxide, CuO. Various ion-exchanged forms of zeolite Y have been investigated as catalysts for the oxidation of molecules such as y-butyrolactone (66),tetralin (67),aromatic amines (68),benzene (69),methanol (70),benzyl alcohol (71), and pyrocatechol (72). In general, transition metal ions or complexes have been incorporated into the zeolite. Space restrictions do not permit a detailed account of these studies. However, in general the selectivities to the more useful partial oxidation products are rather poor. The patent literature claims that olefins can be partially oxidized to epoxides (73) or hydroxy epoxides (74) and alcohols may be oxidized to ketones or aldehydes (75) using various metal ion-exchanged zeolites. In the examples given, the selectivities or conversion levels to the desired products are not particularly attractive. Metal ion-exchanged zeolites do, however, appear to be quite useful catalysts for effluent treatment. For example, Cu2+Xand C u 2 + Yare claimed to be good catalysts for the total oxidation (incineration) of chlorinated organic compounds (76).

111.

Hydrogenation and Dehydrogenation

Since hydrogenation and dehydrogenation have been extensively covered in a number of previous review articles (1-7), the discussion here is largely confined to recent developments.

NONACID CATALYSIS WITH ZEOLITES

19

A. HYDROGENATION In the past it has been generally accepted that alkaline earth and certain trivalent metal ion-exchanged zeolites are inactive for hydrogenation reactions. This was hardly surprising since these zeolites do not normally possess conventional sites where hydrogen molecules can be activated. Soviet workers have claimed, however, that such zeolites, (A, Y, mordenite, omega, chabazite) which evidently do not contain transition metal ion or other impurities are indeed active catalysts for the hydrogenation of alkenes, aromatics (77-89), and oxygenates (90-93). The activities were shown to be dependent on both the cation and zeolites structures. The cations were proposed as the active sites, whereby hydrogen dissociation is induced by the dipole formed between the cation and oxygen of the zeolite framework, e.g., Na'Z-

+ H, ti NaH H . 2

There is, however, no direct evidence for the above equilibrium, and further, these results have not yet been confirmed by other workers. This would seem to be very worthwhile doing in order to clarify this situation. Conventional zeolite-based hydrogenation catalysts are prepared by ion exchange with a transition metal ion followed by reduction. As previously discussed, the reduction step leads to the simultaneous formation of acid sites, i.e., M"+Z + ( n / 2 ) H 2+ Mo

+ nZH

Recently, a nickel zeolite hydrogenation catalyst has been prepared by a novel route (94) involving the adsorption and decomposition of nickel carbonyl onto NaX, which would not be expected to result in the formation of acid sites. In general, the platinum metal-containing zeolites are more active than those containing other transition metals. For example, in zeolite Y the following activity series has been found, PtNaY > PdNaY >> NiNaY

The methods of ion exchange and subsequent reduction are also important in determining the final metal dispersion and hence the catalyst activity (77,95-100). Zeolite-based hydrogenation catalysts containing platinum and palladium have increased resistance toward sulfur poisoning (101- 1 0 4 , and a higher activity (95, 105) than many other supports. In recent years there has been some effort devoted to attempt to explain this phenomenon. Although there is general agreement that the catalytic surface of the zeolites most probably

20

1. E. MAXWELL

exhibits weaker bonding toward the electronegative sulfur atoms, a detailed understanding does not yet exist. At least three proposals have been made, namely : The presence of an atomic dispersion of the metal (101); (2) the formation of a charge-transfer complex between the metal and strongly acidic O H groups (102,103), e.g., M" . . . HOZ", resulting in an electron-deficient metal atom; (3) the presence of small (< 10 A), electron-deficient metal clusters inside the zeolite cages (104). (1)

Figueras et ul. (105) found some direct evidence for electron-deficient palladium clusters on various cation-exchanged forms of zeolite Y from C O adsorption experiments. In particular, a correlation was observed between the turnover number for benzene hydrogenation and the CO stretching frequency. The shift toward higher frequency with increasing support acidity was considered as evidence for increased electron acceptor properties of the support. Further studies will, however, be required to provide a more detailed understanding of this phenomenon. Other recent studies have included zeolites containing copper (106), nickel (107-110), rhodium (111, 112), rhenium (113), and ruthenium (114, 115). For nickel-containing zeolites (e.g., NiY) three types of nickel were identified (1 101, e.g., nickel particles, in the range 20-500 8, o n the outer surfaces of the zeolite crystalli tes ; ( 2 ) small metallic clusters retained inside the zeolite cavities; (3) unreduced Ni2+ cations inside the zeolite cavities. (1)

The catalytic behavior is, not surprisingly, dependent on the relative concentration of these different types of nickel. There was evidence that type (1) nickel particles were similar in behavior to this metal on other supports. However, type ( 2 ) nickel clusters exhibited low activity toward benzene hydrogenation. Japanese workers (112) recently carried out a comprehensive study of the catalytic activity of RhY as a function of degree of reduction. Three different oxidation states of rhodium (Rh", Rh', Rho) were distinguished by using the X-ray photoelectron spectroscopy (XPS) technique. It was shown that Rh' was active for both ethylene hydrogenation and dimerization, whereas Rho was active for hydrogenation of both ethylene and acetylene. The demonstration of hydrogenation activity for Rh' Y is particularly interesting in view of the analogy with homogeneous catalyses where complexes containing formally monovalent rhodium [e.g., RhCI(PPh,), and RhH(Co)(PPh,), (116)] are highly active catalysts for the hydrogenation of olefins.

NONACID CATALYSIS WITH ZEOLITES

21

Coughlan et al. (114,115) recently prepared a series of ruthenium catalysts on zeolites A, X, Y , L, mordenite, and compared their catalytic activities for benzene hydrogenation. Perhaps the most remarkable result was the correlation between the turnover number and the metal surface area (determined by hydrogen adsorption), as shown in Fig. 1 1. This was considered to provide evidence for an influence of metal crystallite size or activity. This is rather surprising since other workers have shown (117) that for Pt, Pd, and Ni on various supports such as silica, alumina, silica-alumina, and silicamagnesia, benzene hydrogenation activity is independent of the metal particle size and the support used. However, the interpretation given for the correlation in Fig. 1 1 is based on the assumption that all the ruthenium metal accessible to hydrogen is also accessible to benzene molecules, which was not conclusively demonstrated. In this respect it is interesting to note that ruthenium in zeolite L exhibited the highest turnover number. There is an extensive patent literature on the use of zeolite based hydrogenation catalysts. Recent examples include aromatics hydrogenation in the presence of sulfur ( ] I d ) , pour-point reduction (119), and general hydroprocessing (120).

B. DEHYDROGENATION Most studies have used the dehydrogenation of cyclohexane to benzene as a test reaction. Kubo et al. (121, 122) found a good correlation between

i . z

/

i;; Slrn'g'

(TURNOVER NUMBER)

FIG.1 I . Dependence of the catalytic activity (turnover number) of ruthenium A, X, Y, L, and mordenite zeolites for the hydrogenation of benzene on the surface area S of the ruthenium mordenite-4; 0, A-2; A, X - I ; A, X-2; m, Y-l ; W, Y-4; 0, L-2; @, L-4. (Rezeolite: 0, produced from Ref. 115 with permission from the authors.)

22

I. E. MAXWELL

the amount of hydrogen adsorbed and the catalytic activity of PtY zeolites for cyclohexane dehydrogenation. In addition, the catalyst precalcination temperature prior to reduction had a marked influence on activity. As shown in Fig. 12 an optimum activity was obtained for a calcination temperature of 300"C, which corresponded to platinum particles in the range of 20-50 A in diameter. These workers concluded that very small platinum particles (<20 diam.) in zeolite Y had reduced hydrogen adsorption capacity and hence reduced dehydrogenation activity. Dehydrogenation activity has been demonstrated for Rh, Co, and Ni forms of zeolites X and Y (123-125). Both cyclohexane and tetralin dehydrogenation to benzene and naphthalene, respectively, have been used as test reactions. For N i x zeolites, unreduced Ni2+ ions were considered (124) to be the active centers. The incorporation of Ca2+ ions into the zeolite

40

' -

30 ..

-ae

I

N

r

I

0"

20-

.-0 L Y)

al C

0

10 -

100

200

300

400

500

Calcination temperature ("C)

FIG. 12. Dehydrogenation of cyclohexane over 0.24 wt.% Pt--Nay at 300°C. (Reproduced from Ref. 122 with permission from the authors and Bulletin of the Chemical Society of Japan.)

NONACID CATALYSIS WITH ZEOLITES

23

resulted in an increase in activity, which was attributed to an increase in the concentration of accessible Ni2+ sites. Corma et al. (126) found that PtNaY was an active but rather unstable catalyst for methylcyclohexane dehydrogenation to toluene. These workers studied both the dehydrogenation and the catalyst decay kinetics. I t was concluded that the reaction occurs via a series of consecutive partial dehydrogenation steps, the first of which was rate determining. Further, catalyst deactivation was caused by coke deposition from partially unsaturated precursor molecules. This section would be most incomplete without discussing the extensive studies of Mobil workers on the dehydrocyclization activity of TeNaX (127-131). Miale and Weisz (127) first showed that the incorporation of electronegative elements such as sulfur, selenium, and tellurium in NaX resulted in enhanced catalytic activity for n-hexane conversion and that depending on the element present, the catalytic characteristics shift from carbon-carbon cracking to dehydrocyclization. In fact, TeNaX was shown to be a very selective catalyst for parafin dehydrocyclization. Comprehensive studies (128-130) involving the influence of hydrogen, single-crystal X-ray diffraction, and diffuse reflectance spectroscopy produced a remarkably detailed picture of the most likely aromatization site in the zeolite. As shown in Fig. 13 this was shown to be a telluride ion (Te2-) located in the supercage and coordinated to two Na' ions at sites I1 and 111. In a final paper (131) in this series the novel TeNaX catalyst was compared with Cr,O,/Al,O, catalyst for n-hexane dehydrocyclization activity. Both

FIG.13. Projection ( I , T, 0) of TeNaX structure. Solid lines, silicon-aluminum framework; circles, oxygen framework. (Reproduced from Ref. 129 with permission from the authors.)

24

I. E. MAXWELL

catalysts showed the same dependence of reaction rate on hydrogen pressure. Further, the kinetic data were consistent with the stepwise dehydrogenation mechanism which had been previously established for Cr20,/A1,0, catalysts (132). IV.

01igomerization Reactions A . ETHYLENE

1. Dimerizution A number of zeolite-based catalysts are active for the dimerization of ethylene. The major products are n-butenes ( 1-butene, truns-2-butene, cis-2-butene), i.e.,

The isomer ratios are often close to the equilibrium composition at the particular reaction temperature, indicating isomerization as well as dimerization catalytic activity. The two most extensively studied zeolite catalysts contain nickel and rhodium, incorporated via ion exchange, and will be discussed separately. Ethylene dimerization catalysis has, however, been more thoroughly investigated for the broader range of homogeneous catalysts. For example, active metal complexes containing titanium, nickel, iron, cobalt, rhodium, ruthenium, and palladium, are all known (133). Where possible, comparisons will be made with the relevant homogeneous catalyst systems. a. Nickel Zeolites. Heterogeneous catalysts prepared by the impregnation of Ni2+ on amorphous aluminosilicate (followed by calcination to yield dispersed NiO) have been known for some time (134-1 37). Typically, n-butenes are obtained with selectivities of about 80% (based on ethylene reacted) at temperatures of 275-300 'C and atmospheric pressure. In a series of publications, Russian workers (135) reported on the comparative activities and selectivities of a variety of Ni2'-exchanged zeolites, together with NiO on amorphous aluminosilicate. In addition, some attempts were made to define the nature of the active site. Catalysts were prepared (138,139) by both ion exchange and impregnation of zeolites with N i 2 + . Ion exchange resulted in the most active but least selective (approximately 25% selectivity to n-butenes under conditions com-

NONACID CATALYSIS WITH ZEOLITES

25

parable to those described above) catalyst. The degree of acidity of the zeolite substrate was varied, i.e., C a y , CaX, HY, HM (mordenite), but no obvious correlation between acidity and activity or selectivity was found. Treatments with hydrogen (at 480°C) resulted in complete and partial losses in activity (139) for impregnated and ion-exchanged catalysts, respectively. This could be correlated with the degree of reduction which was incomplete under these conditions for the ion-exhanged catalysts. It was therefore concluded that the active site involved Ni2+ ions bonded to the oxygen anions of the zeolite framework and that nickel metal was inactive in the temperature range 150-250°C. Yashimo et al. (140) found that NiY (70% exchanged, equivalent to 9 wt. % Ni) was very active and selective for ethylene dimerization at only 20'C and 200 Torr ethylene pressure. The higher nickel loading would not seem to be sufficient to account for this very marked difference in activity as found between the Russian and Japanese workers. It may be relevant that the latter workers carried out their exchange using NiCl, [the former workers used Ni(NO,),]. Halide ions were shown to be essential to the ethylene dimerization activity of Rh(1) complexes (141a). Further work will, however, be required to clarify this aspect. The mechanism of ethylene dimerization for heterogeneous nickelcontaining catalyst systems does not appear to be at present well understood. b. Rhodium Zeolites. Rhodium complexes are well known for their homogeneous catalytic activity for a variety of reactions with olefins (133). In general, they exhibit high activity and specificity for these reactions which include olefin oligomerization. For example, the simple hydrated chloride of rhodium (RCI,.3H20) was shown by Alderson e f af. (1416) in 1965 to catalyze the dimerization of simple olefins including ethylene with high selectivity. However, it was some ten years later before studies on a zeolitebased heterogeneous rhodium catalyst were published. Under very mild conditions (0-20"C, 200 Torr ethylene pressure), ethylene was shown to be selectively dimerized to n-butenes over RhY (140). As shown in Fig. 14, 1-butene was formed initially but further isomerized to an equilibrium composition of n-butenes with increasing reaction time. In a comparative experiment using HY as a typical solid-acid catalyst, no ethylene conversion was measurable up to 200'C, and at higher temperatures unselective polymerization and cracking reactions occurred. This provided good evidence that the selective dimerization over RhY did not proceed via a carbenium ion mechanism. The same authors (140) then carried out a number of experiments where the catalyst pretreatment conditions were varied and reagents were added to the reaction mixture (i.e., ammonia, pyridine, and CO) in order to elucidate the type of active site. Rather surprisingly, it was concluded that

26

1. E. MAXWELL

0

20

40 60 reaction

80 100 120 time (min)

140

i0

FIG.14. Composition change of produced n-butenes with reaction time. Reaction conditions: catalyst, RhY (0.30 g) activated by evacuation at 300°C for 1 hr; temperature, 0 C ; ethylene initial pressure, 200 Torr. (Reproduced from Ref. 140 with permission from the I-butene; 0 ,trans-2-butene; @, cis-2-butene. authors.) 0,

the active sites in RhY were highly dispersed zero-valent rhodium. This conclusion contrasts with studies on the homogeneous rhodium system (141) which indicated that a monovalent rhodium complex was the active entity. More recently, other Japanese workers ( 112) also investigated the active sites in RhY for ethylene dimerization. These workers found that there was a distinct optimum activation temperature with respect to catalytic activity as shown in Fig. 15. XPS studies on the same catalyst (50%exchanged RhY) showed that there were substantial shifts in the Rh 3d line intensities and binding energies as a function of activation temperature. By comparison with known compounds the XPS spectra were interpreted in terms of a composition of trivalent, monovalent, and zero-valent oxidation states of rhodium. Deconvolution of the spectra gave the relative concentrations of each oxidation state for each activation temperature, as shown in Fig. 16. Comparison of Figs. 15 and 16 shows that there is a good correlation between the ethylene dimerization activity and the Rh( I ) concentration at various activation temperatures, with a maximum at 250 C. Contrary to suggestions by the previous workers (140), it was therefore proposed that Rh(1) was the active oxidation state for the ethylene dimerization reaction.

27

NONACID CATALYSIS WITH ZEOLITES

-

0.5

c

0.4

0.2

-

. .-C

E

- 0.3 &

P

2

0 ql

2

0.2

0.1

-

z

P

r

f 0.1 K

I 0

100

200

300

400

0 500

Activation Temperature ("C)

FIG. 15. Ethylene dimerization activity for RhNaY catalyst as a function of activation temperature. (Reproduced from Ref. 112 with permission from the authors.)

100

80

-ae

.-5 60 +?i

U 0

.-

40

3 20

0 100

200

300

400

500

600

Activation Temperature ("C)

FIG. 16. Oxidation state of Rh in the RhY (50%) catalyst as a function of activation temRh(l) (308.3 eV); and 0 , R h metal (307.5 eV). perature: 0. Rh(I1l) (Rh 3d 5/2, 310.2 eV); 0, (Reproduced from Ref. 112 with permission from the authors.)

28

I . E. MAXWELL

Clearly, with good evidence for Rh(1) as the active species in RhY, there is now a consistency between the homogeneous (141) and heterogeneous catalyst systems. A mechanism for the homogeneous dimerization reactions has been proposed by Cramer (141),which involves the formation of anionic Rh(1) complexes. The presence of such species inside the cavities of the anionic zeolite framework does not seem very probable. Further studies will be required to elucidate the mechanism and details of the active species for the RhY catalyst. c. Palladium Zeolites. Palladium complexes, such as tetrachlorobis(ethylene) dipalladium (PdCI,C,H,), , are known homogeneous catalysts for the dimerization of ethylene to n-butenes (133). It is therefore not so surprising that Pd zeolites have recently been shown to exhibit activity for this reaction (142-146). Both impregnated and ion-exchanged catalysts have been prepared, the latter by exchange using aqueous solutions of [Pd(NH,),]CI,. N a mordenite, N a y , CaNaY, and CaNaX were used as substrates. Catalytic activity for ethylene oligomerization was observed in the temperature range 50-200°C [atmospheric pressure, gas hourly space velocity (GHSV) 900 hr-’1. However, in general the selectivity to n-butenes was rather poor (maximum 60%), even at relatively low ethylene conversions. Hydrogen treatments were found to deactivate the catalyst due to the reduction of Pd” to PdO. This result together with some XPS data was used to arrive at the conclusion that cationic palladium was the active species. Under milder conditions (20”C, 200 Torr ethylene) other workers (140) found PdY (also prepared via ion exchange with [Pd(NH,),]CI,) to be inactive for ethylene dimerization. In general, the active homogeneous palladium complexes contain a halide ion (133), typically chloride, in the coordination sphere. The influence of halide ions on the Pd zeolite catalysts would therefore seem worthy of study. d. Other Metal Zeolites. In addition to the previously discussed catalysts (i.e., NiY, RhY, PdY), Yashima et al. (140)also examined C r y , RuY, COY, Fey, MnY, CdY, and ZnY for catalytic activity under mild conditions (20”C, 200 Torr ethylene). All these catalysts, with the exception of PdY, were prepared by ion exchange using the corresponding transition metal chloride salts. Chromium Y zeolite exhibited high activity, but poor selectivity to butenes. A batch experiment using benzene as a solvent showed that highly crystalline polyethylene was the major product, which will be discussed in more detail later. Ruthenium Y zeolite showed selective dimerization activity to n-butenes, but was less active than RhY and NiY. The remaining metal ion-exchanged zeolites were all inactive under the chosen reaction conditions.

NONACID CATALYSIS WITH ZEOLITES

29

2. Polymerization During screening studies for oligomerization activity of various transition metal ion-exchanged Y zeolites, Yashima et al. (140) found that CrNaY (82% ion exchanged with CrCl,) was highly active for ethylene polymerization. A more detailed study (147) showed that the polyethylene produced had a high melting point, high molecular weight, high density, and a linear chain structure without branching. In fact, the properties of the polymer product were considered to be similar to those obtained using chromium oxides in the Phillips process. On the basis of spectroscopic studies and due to the influence of various calcination procedures on catalytic activity, it was proposed that C r 2 + ions in the zeolite cages were the active sites for the polymerization reaction. However, Russian workers (148) carefully prepared Crz -exchanged zeolite Y under oxygen-free conditions and found that this material was inactive for ethylene polymerization. A similar lack of consensus as to the active species exists for the Phillips catalysts where both Cr3+ (149, 150) and CrZf (151-153) have been proposed to be the active oxidation states. Further studies will be necessary to resolve this question. More recently, Czech workers (154) have shown that the polymerization activity of chromium zeolite Y catalysts can be considerably enhanced when ion exchange is carried out using ultrastable HY. These workers propose that this is due to an improvement in the stability of the active valence state of chromium in the ultrastable form of the zeolite. Cobalt-containing zeolites have been studied for polymerization of ethylene (155-157). The catalysts which were prepared by precipitating cobalt carbonate together with zeolites A, X , Y, and mordenite were not very selective, yielding large amounts of ethane as well as C, and C, hydrocarbons. +

B. ACETYLENE Kruerke (158) appears to have been the first to screen various metal ionexchanged zeolites for acetylene oligomerization activity. Although detailed data have not been published, these studies showed that N i 2 + -and Co2+exchanged zeolite Y were very active for this reaction, producing benzene at near room temperature. The Mn2+-exchanged form of zeolite Y was only slightly active, whereas the Na', C a 2 + , Z n 2 + , and Cu2+ forms were all inactive for this reaction. French workers (159,160) more recently screened a wide range of first-row transition metal ion-exchanged Y zeolites and concluded that the active cations for acetylene trimerization were those with an even number of partially filled d orbitals, i.e., d8(Ni2+,Co'), d6(Fe2+),

30

1. E. MAXWELL

d4(Cr2+)were active, whereas d'O(Zn2+, Cu+), d9((Cu2+),d7(Co2+),and d S ( M n 2 + Fe3+) , were found to be inactive. This activity behavior seems to be consistent with a model suggested by Union Carbide workers ( 1 5 9 , where the reaction path involves two acetylene molecules simultaneously coordinated to the metal cation. These studies (159,160) were unfortunately only based on infrared spectroscopic measurements rather than microreactor experiments and thus provide relatively little information with regard to catalytic activity, selectivity, and stability. The NiY system was examined in some detail and a correlation was found between catalytic activity and the number of Ni2+ ions in the supercages. This accounted for the fact that NiY was only active when the exchange level exceeded 12 Ni2+ cations per unit cell. At lower exchange levels these cations are preferentially located in the hexagonal prisms and sodalite cages, which are inaccessible to acetylene molecules. The NiY zeolite was also shown to be active for the cyclotrimerization of propyne with 1,2,4-trimethyIbenzene being the main product. The activities of the above-mentioned transition metal ions for acetylene trimerization are not so surprising since simple salts and complexes of these metals have been known for some time to catalyze this reaction (161, 162). However, the tetramer, cyclooctatetraene, is the principal product in homogeneous catalysis, particularly when simple salts such as nickel formate and acetate are used as catalysts (161). The predominance of the trimer product, benzene, for the zeolite Y catalysts might be indicative of a stereoselective effect on product distribution, possibly due to the spatial restrictions imposed on the reaction transition-state complex inside the zeolite cages.

C. PROPYLENE 1.

Nickel Zeolites

Recently, a large variety of ion-exchanged zeolites of type X and Y were examined (163) for propylene oligomerization activity. These included L a y , L a x , CeX, MgX, NiY, COY, AIY, MgX, MnY, NIX, COX, and CaX zeolites which were tested in a fixed bed reactor at 190'C. With the exception of Nix, all the zeolites tested showed rather unselective hydrodimerization activity leading to a wide variety of paraffinic products. The appearance of saturated C, , C,, C, , and C, products was indicative of cracking reactions was well as hydrogen transfer. Nix, however, showed a high selectivity resulring in 95.5';.:, dimers under the reaction conditions. The product composition of propylene dimers was studied as a function ofcontact time in order to distinguish between primary and secondary products. 3-Methylpentenes were shown to be primary prod-

NONACID CATALYSIS WITH ZEOLITES

31

ucts and not formed via secondary isomerization of either hexenes or 2-methylpentenes. This indicated that a mechanism was operative similar to that found (164) for NiO on silica-alumina. Cyclobutane derivatives were invoked as intermediates in order to explain the primary product distribution. In contrast to NIX, NiY showed very poor selectivity for dimerization; the major products resulted from hydrodimerization and cracking reactions. It was concluded that these unselective reactions occurred over strong acid sites which were present on all the studied catalysts with the exception of N i x .

2 . Alkali Metal Zeolites In a patent to Asahi Electro-Chemical Co. (165), alkali metal- (i.e., Li, Na, K, Rb, or Cs) containing zeolites are claimed as active and selective catalysts for dimerizing C,-C, a-olefins. An example is given in which a catalyst is made by adding potassium metal (30 wt.%) to calcined KA. In a batch experiment at 150°C (heptane solvent) hexenes are obtained at the rate of 0.2 g (g catalyst)-' hr-'. By-product formation is claimed to be inhibited under these conditions.

D. CYCLOPROPENES Zeolites have been shown to exhibit a rather novel type of catalysis for the cycloaddition reaction of cyclopropenes to give tricyclohexanes (166), i.e.,

For a variety of methyl-substituted cyclopropene molecules, and a variety of zeolites (e.g., NaCaA, NaA, CaA, NaX, and HY), the reaction selectivity to dimer, diene, or polymer was found to be determined by at least three factors, namely : (1) type of active site, e.g., K A > NaA; ( 2 ) size of reactant molecule (i.e., number of methyl substituents); (3) shape and size of zeolite pores. Dimerization was found to occur only at a specific ratio of the size of the cylopropene (or methyl-substituted) molecule to the diameter of the pores. The best results were obtained with KA and NaA; HY resulted in polymerization in all cases. An ionic dimerization mechanism was proposed whereby the spatial constraints within the zeolite (with the correct size ratio) impede the approach of a third olefin molecule which would lead to polymerization.

32

I . E. MAXWELL

E. BUTADIENE Butadiene dimerization catalysts have been quite extensively studied, although the major effort has been concentrated on homogeneous catalyst systems using complexes containing nickel (167). iron (168), cobalt (169, and palladium (170). The products are normally the cyclodimers, 4-vinylcyclohexene and cycloocta- 1 ,5-diene, together with some trimer, cyclododeca-l,5,9-triene and possibly a small amount of divinylcyclobutane depending on the conditions and the particular catalyst, i.e.,

Until recently when a number of these homogeneous catalysts were heterogenized (I71), the only known heterogeneous catalysts were based on metal ion-exchanged zeolites. 1. Copper Zeolites The first work in this area appeared in the form of two patents assigned to Union Carbide in 1969/1970 ( I72, 173). These patents described methods of preparation of monovalent copper-containing zeolites which were claimed to be active and selective catalysts for the cyclodimerization of butadiene to 4-vinylcyclohexene (VCH), i.e.,

The same catalysts were also claimed to catalyze the coupling reaction between butadiene and acetylene to produce 1,4-cyclohexadiene, i.e.,

Both reactions could be carried out under mild conditions (100-1 10°C and atmospheric pressure). The very high selectivity to VCH in the former reaction is of particular interest since the only other butadiene dimerization catalysts (167) known at that time were homogeneous and in general gave, in addition to VCH, several other dimerization products (e.g., 1,4-cyclo-

NONACID CATALYSIS WITH ZEOLITES

33

octadiene and divinylcyclobutane). VCH is of potential commercial interest since it can be readily converted to styrene via oxidative dehydrogenation. Two methods of preparing Cu+ zeolites have been described ( I 72-1 74), namely, (1)

direct exchange of the zeolite with Cu’ ions (via Cul in liquid ammonia), and (2) initial exchange with Cu2+ ions followed by mild reduction to the Cu+ form. Catalysts prepared via the latter method were found to be more rapidly deactivated during reaction. This is almost certainly due to the formation of Bronsted or Lewis acid sites during the reduction step (175, 176), i.e.,

(from reducing Ht agent)

J 0

Si 0’ ‘0

A1 0 ’

Hf 0

\o

cui 0

0

Si

o/ \o

’0

\o

The acid sites so formed catalyze the polymerization of butadiene which leads to catalyst deactivation. Maxwell et al. (177, 178) studied the deactivation of reduced C u 2 + Y catalysts for butadiene cyclodimerization in some detail. This work showed that the catalyst stability could be markedly improved by using NH, as a reducing agent and choosing the activation conditions such that excess NH, remains selectively chemisorbed on the zeolite acidic sites. Further, the Cu2’Y-derived catalyst was thermally stable to 850°C and was therefore able to withstand a regeneration procedure which involved a polymer burn-off at 550°C. By contrast, the catalysts prepared by direct exchange with monovalent copper, i.e., C u + Y , formed CuO irreversibly when heated above 330°C. An X-ray structure analysis was carried out (179) on a single crystal of natural faujasite which had been exchanged with CuZc ions, dehydrated, and then exposed to butadiene. The major effect of adsorbing butadiene was to induce a migration of copper cations to site 111, located at the pore entrances to the supercages (see Fig. 17). The unsaturated coordination of

34

I . E. MAXWELL

Si, A t

FIG. 17. Perspective view showing the siting of Cu(II1) cations at the pore entrance to the supercage (Cu2+-exchanged faujasite, dehydrated at I SWC, butadiene adsorbed). The occupancy factors are such that there is approximately one Cu(11l) cation per two pore entrances.

these cations to the zeolite framework and their ideal location for interaction with adsorbate molecules led to the suggestion that these were most probably the sites where the butadiene cyclodimerization reaction occurs. However, the zeolite is not a unique substrate for this reaction, as is indicated in a recent patent (180),where it is shown that a Cu+-exchanged montmorillonite clay and synthetic amorphous aluminosilicate will also catalyze butadiene cyclodimerization with high selectivities to VCH ( 2 95:'<). Preexchange of these aluminosilicates with Cs+ ions was claimed to increase catalyst stability. This is most probably explained by a reduction in surface acidity resulting from the alkali metal ion exchange. The remarkable feature of all the above-described catalysts is their ability to give very high selectivities to VCH, in contrast to the homogeneous

NONACID CATALYSIS WITH ZEOLITES

35

catalyst systems (167). These homogeneous catalyst studies showed that the presence of at least two strong ligands in the coordination sphere of the metal ion was required for the formation of the o-allyl, n-ally1 complex, a precursor to VCH, i.e.,

l+

The possibility therefore existed that for these Cu+ aluminosilicate catalysts, oxygen anions fulfill this strong ligand role. To check this proposal a Cu+ complex was prepared (179) which contained extremely weak ligands, namely, cuprous trifluoromethanesulfonate. Surprisingly, this complex under homogeneous reaction conditions was also found to catalyze butadiene cyclodimerization selectively to give VCH (90%). It was therefore concluded (179) that in this case the selectivity to VCH probably resided in the electronic structure of the cuprous ions rather than any particular ligand effect of the zeolite framework. 2 . Other Metal Zeolites Soviet workers (181) studied the cyclodimerization of butadiene on various cationic forms of zeolite X. However, it is noteworthy that the temperature range where catalytic activity was observed (300-6OO'C)was considerably higher than that for the previously discussed Cu' zeolite catalysts (90- 1 lO'C). Interestingly, under these conditions the parent form of the zeolite, i.e., NaX, is also on active catalyst and forms VCH in high yields (99%). Ion exchange of NaX with Cu2+,Ni2+, Cr3+, and Co2+ (0.5-1%) apparently had no marked effect on activity. However, cation exchange with rhodium resulted in a 2-2.5-fold increase in the rate of butadiene cyclodimerization. The catalytic activity of NaX could possibly be due to a concentration effect of the reactant molecules in the zeolite pores. This would be expected to be particularly pronounced if the reaction was second order in butadiene. Such is the case for the thermal reaction of butadiene at 650'C (182).Furthermore, VCH is the major reaction product only at short residence times, probably because this molecule has the lowest standard heat of formation of the possible cyclodimerization products (183). The activity and selectivity of NaX for butadiene dimerization could possibly be due to rate enhancement of the thermal dimerization reaction under the prevalent conditions (i.e., relatively high temperature, short residence times).

36

1. E. MAXWELL

A patent (184) describes zeolite-based catalysts for the direct production of ethylbenzene from butadiene, i.e.,

The reaction was carried out at very high temperatures (500°C) using various cation-exchanged forms of mordenite (i.e., Ni", C u 2 + ,Z n 2 + ,C o 2 + ,C r 3 + , Pb2+, H', Bi3+, M n 2 + ) .Comparative experiments showed that the C o 2 + exchanged zeolite gave the highest selectivity to ethylbenzene (54.3%) and the lowest level of solid residue on the catalyst under the chosen reaction conditions. The reaction almost certainly involves the cyclodimerization of butadiene to VCH followed by dehydrogenation to form ethylbenzene. It is not clear from the patent whether hydrogen is a product of the reaction or whether hydrogen transfer occurs. The formation of some ethylcyclohexane suggests that the latter reaction is occurring to some extent. The rather low selectivities and relatively high extent of solids formation render these catalysts unattractive for a single-step route from butadiene to ethylbenzene.

F. II-BUTENE n-Butene dimerization to octenes is of potential commercial interest, since the octenes are useful intermediates in the preparation of higher alcohols (e.g., nonanols) via hydroformylation. The higher alcohols are themselves intermediates for the synthesis of biodegradable detergents such as alkylsulfates. The traditional heterogeneous catalysts for the oligomerization of lower olefins are supported COO and NiO. More recently, however, nickel-exchanged zeolites have attracted attention as catalysts for these reactions. In a patent (185), ion-exchanged zeolites containing group VIII metal ions were claimed as catalysts for the dimerization of lower olefins. In fact, only nickel-exchanged zeolite X was given in the examples. An interesting additional feature of the patent was the pretreatment of NIX (5% Ni by weight) with an organic or inorganic base in order to improve the catalyst performance. The base pretreatment increases both activity and selectivity. The latter effect is almost certainly due to neutralization of residual zeolite acidity which would catalyze undesirable side reactions such as oligomerization, hydrooligomerization, polymerization, and cracking. More detailed information on 1-butene dimerization over N i x is given in a publication (186) by the same authors of the above-mentioned patent

NONACID CATALYSIS WITH ZEOLITES

31

MOLAR FRACTION AT REACTOR EXIT

‘Or

0

-I

10 20 30 r ( g hr/mol) RESIDENCE TIME

FIG.18. 1-Butene dimerization over a NiX/Li,O catalyst at 180°C. (Reproduced from Ref. 187 with permission from the authors and La Chimica c L’lndustria.)

(185). For these studies the NIX precursor catalyst has been impregnated with molten lithium acetate. The catalyst displayed remarkably good selectivity to octenes. A typical set of results is shown in Fig. 18, where the reactant and product concentrations are given as a function of contact time, 7. The results indicate that isomerization of 1-butene to a n equilibrium mixture with 2-butenes (cis and trans) occurs rapidly, followed by the slower dimerization of butenes to octenes. The reaction was carried out under mixed phase conditions with the 1-butene reactant in the gas phase and the octene products in the liquid phase (trickle-bed reactor). Under the conditions the reaction rate was found to be mass-transfer controlled. In a subsequent paper (187) the same authors studied the kinetics of propylene/ 1-butene codimerization over the same NiX/Li,O catalyst. As might be expected, the selectivity to heptenes is not particularly high due to competitive reactions such as self-dimerization of propylene and butene.

G . ISOBUTYLENE

Isobutylene is present as 20-30% of the C , fraction from the naphtha cracking process. A number of different upgrading reactions with isobutylene have been carried out industrially (with and without prior separation from the C , fraction). One of these includes the acid-catalyzed oligomerization

38

1. E. MAXWELL

to dimers and trimers (48)which are useful intermediates for the manufacture of plasticizers. Russian workers, in particular, have been systematically studying the use of zeolites as catalysts for isobutylene oligomerization for some years. The activities and selectivities of a variety of ion-exchanged zeolites have been examined by these workers (188-1 91). A number of general trends can be discerned, i.e., (1) X-type zeolites are more active than their zeolite Y analogs. (2) Polyvalent cation-exchanged forms are more active than their monovalent analogs. (3) The activity order for the alkali metal ion forms is inversely proportional to the ionic radius, i.e., Li' > Na' > K + > Rb'. (4) The hydrogen forms of zeolite X and Y are generally more active, but less selective, than their cation-exchanged forms. Side reactions include hydrogen transfer (resulting in the formation of coke and paraffinic products), double-bond migration and disproportionation. ( 5 ) In general, the catalysts are not stable due to the build-up of coke. However, regeneration can be accomplished by a coke burn-off at 480°C in a stream of air. (6) CO, has been found to have a promoting effect (191-193) on the overall catalyst activity, but decreases oligomerization selectivity. Although the reactions have been generally described in terms of a carbenium ion mechanism, this does not altogether explain the catalytic behavior of the alkali metal ion-exchanged zeolites or the selectivity behavior. An ionic mechanism of the type previously described for cyclopropene dimerization would seem to be more appropriate for the alkali metal ionexchanged zeolites, where the activity does seem to correlate qualitatively with the electrostatic field ( e / v )exerted by the cation. The CO, effect has been explained (IYf) in terms of the formation of alkali carbonate or bicarbonate species, where the cation is replaced by a proton derived from a water molecule. This proposal does indeed account for both the observed increase in activity and decrease in selectivity which is similar to proton-exchange forms. Rhein and Clarke (194) screened a wide variety of alkali metal ioncontaining zeolites (A, X, Y, and L) for isobutylene polymerization activity under mild conditions (room temperature, in an autoclave). In general the activities were rather low, requiring some days to achieve reasonable polymer yields. The polymers obtained were characterized by a bimodal molecular weight distribution. The low and high molecular weight peaks were at approximately 250 and 3000-5000, respectively. Polymerization only oc-

NONACID CATALYSIS WITH ZEOLITES

39

curred in the presence o f zeolites with pore-entrance diameters in excess of 5 A (i.e., zeolites 3A and 4A were inactive), which is close to the kinetic diameter of isobutylene. This does seem to indicate that the reaction occurs at least in part within the pores of the zeolite. Polymer yields for different zeolites were on the order of 5A > X > Y > L. Ion exchange with various transition metal ions gave the following order for polymer yield: Pd > Pt x Ni > Cu.

V.

Carbonylation

Carbonylation of methanol has in recent years become a commercially important route for the production of acetic acid and methyl acetate. Industrial catalysts are at present homogeneous, based on cobalt and more recently rhodium compounds. The cobalt catalysts are less active (195) and require more severe operating conditions (i.e., 250°C, 650-750 atm) than the rhodium-based catalysts (196) ( 1 70-25O0C, 7- 14 atm). It is likely that synthesis gas and methanol (also made from synthesis gas) could become important basic building blocks of the petrochemical industry in the near future, particularly as coal and natural gas gain use as feedstock materials. Such a development would be expected to increase the relative importance of carbonylation reactions in industrial chemistry.

A. METHANOL Since the discovery of highly active homogeneous rhodium catalysts for methanol carbonylation by Monsanto workers (f 96), there has been considerable industrial and academic research effort in this area. The catalyst system consists not only of a rhodium complex, but also includes a halogen promoter, preferably containing iodine, typically methyl iodide, which is regenerated at the end of the catalytic cycle. The reaction is highly selective, giving high yields (99%) of carbonylated products (i.e., acetic acid or methyl acetate, depending on reaction conditions). The trace by-products are usually dimethyl ether and acetaldehyde. In 1970 Monsanto ( I Y 6 ) brought the first acetic acid plant on stream which was based on this new homogeneous catalyst system. The important advantages of the rhodium catalyst are the higher methanol conversion level and lower operating pressure compared to the previous cobalt catalyst system developed by BASF (195).

40

I. E. MAXWELL

The reaction mechanism proposed by Roth et ul. (196) is as follows: CH,OH

+ HI # CH,I + H,O

+ CH-J CH,Rh(III)IL, CH,Rh(III)IL, + CO#CH,Rh(III)ICOL, Rh(I)L,

CH,Rh(III)ICOL,c-LCH,CORh(III)IL, CH,CORh(III)IL,

+ H,O#Rh(I)L,

+ CH,COOH + HI

(1)

(2) (3)

(4) (5)

where L, are ligands which may be of different types. Step (1) involves the formation of methyl iodide, which then reacts with the rhodium complex Rh(I)L, by oxidative addition in a rate-determining step (2) to form a methylrhodium(II1) complex. Carbon monoxide is incorporated into the coordination sphere in step (3) and via an insertion reaction a rhodium acyl complex is formed in step (4). The final step involves hydrolysis of the acyl complex to form acetic acid and regeneration of the original rhodium complex Rh(I)L, and HI. Typical rhodium compounds which are active precursors for this reaction include RhCl, , Rh,O, , RhCl(CO)(PPh,), , and Rh(CO),CI, . Although the homogeneous catalyst systems have been successfully applied in commercial practice, some intrinsic problems associated with catalyst separation remain. This has led to considerable interest in the development of a suitable heterogeneous analog. Rhodium compounds have been heterogenenized on substrates such as carbon (197), alumina (198, 19Y), and synthetic polymers (200). More recently, zeolites have also attracted quite some attention as a support material for carbonylation catalysis, as is discussed later. Scurrell (201) recently briefly reviewed the literature on heterogenized homogeneous rhodium catalysts for methanol carbonylation up to 1976. 1. Rhodium Zeolites a. Zeolite X . Russian workers appear to have been the first to prepare a rhodium-zeolite catalyst (202)(i.e., RhNaX) and to show that this material was highly active, stable, and selective for methanol carbonylation. This catalyst was prepared by impregnating NaX with an aqueous RhC1, solution neutralized with 10% NH, to pH 4.5. At atmospheric pressure and 250% methyl acetate was obtained with a 87-90% selectivity (dimethyl ether was the sole by-product) at a methanol conversion level of about 79%. The specific activity of this catalyst (- 50 g methyl acetate (g Rh)-' hr- '), which contained 0.2 wt. % Rh, well exceeded that of 3 wt. Rh supported on activated carbon (197) (8-18 g methyl acetate (g Rh)-' hr-') (197). In a subsequent publication, Nefedov et ul. (203) examined the influence

NONACID CATALYSIS WITH ZEOLITES

41

of various catalyst and process parameters for this reaction. It is interesting that NaX was found to be inactive for dimethyl ether formation in the absence of methyl iodide but quite active in the presence of the latter (98% conversion of CH,OH, at 250"C, GHSV 0.64 hr- '). Although this was not stated in the publication, it would appear that NaX is a catalyst for the dehydrohalogenation reaction, i.e., CH,OH

+ CHJ

~t

CH,0CH3

+ HI

The specific activity for both methyl acetate and dimethyl ether as a function of rhodium level on the catalyst was measured and is shown in Fig. 19. Clearly, there is a marked decrease in specific activity for both products with increasing rhodium level in the range 0.5-1 wt.% Rh. The optimum in terms of catalyst efficiency was considered to occur in the range 0.25-0.5 wt.% Rh. The reaction rate was found to be zero order in both CO and CH,OH partial pressures, as has also been found for the homogeneous catalyst system (196). Russian workers have also investigated the influence of different support materials on methanol carbonylation activity and selectivity. A summary of the results obtained is given in Table 11. It would appear from these data that the selectivity to methyl acetate decreases with increasing support acidity. These results are consistent with those of other workers (204) who also found rather poor carbonylation selectivities (< 50%) for rhodium complexes supported on y-alumina, due to ether formation. Thus NaX is shown to be the best support resulting in a catalyst with both high activity and selectivity. In a comparative study (205), RhNaX catalysts were prepared by both ion exchange and impregnation with RhCI,. The former catalyst was superior with regard to both activity and selectivity, provided that the chloride ions were thoroughly removed following ion exchange. A

0.2

0.6

1 .o

Rh. %

FIG.19. Specific activity of RhNaX for methanol carbonylation as a function of rhodium concentration. (Reproduced from Ref. 203 with permission from the authors and Plenum Publishing Corporation, New York.)

42

I . E. MAXWELL

TABLE 11 E f f i c , t of Support Type on Activity rrnd Selec.riui/v of Imprtgtiuted Rhodium Cutulysts ,fbr Methunol Crirhotiylririon“

Support

CH,OH conversion (mol

Selectivity” to CH,COOCH, (mol”/,)

NaX NaY NaA CaX

78.0 80.5 31.6 60

74 59 59

‘x)

1

Support NaM A1203

Amorphous alumino-silicate

CH,OH conversion (mol%)

Selectivity” to CH,COOCH, (mol ‘i;)

34.6 57.5 43.6

I1 17 0

~

“0.5 wt. ”/, Rh, 250”C, GHSV 0.67 hr-’, atmospheric pressure, CO/CH,OH = 1.5-1.6, CH,I/CH,OH = 0.25. I, Selectivity, S , to CH,COOCH, was calculated using S = [ T ~ ~ ~ +~ rrlher)Ir ~ J ( where ~ ~ rEltrr= rate of formation of ester; = rate of formation of ether (sole by-product).

similar influence of chloride ions was observed for carbon-supported rhodium catalysts (198, 206,207). Treatment of the ion-exchanged RhNaX catalyst with hydrogen (205) results in almost complete reduction of Rh3+ to Rho with a consequent much reduced carbonylation activity, whereas ether formation remains virtually unaffected. These results demonstrate that the dehydration activity is a function of the support only and that cationic rather than zero-valent rhodium is the active entity for the carbonylation reaction. More recently, Christensen et al. (208) prepared RhNaX by immersion of NaX into an aqueous solution of RhCI, , probably resulting in some ion exchange as well as impregnation of rhodium. The results obtained for methanol carbonylation with respect to both specific activity and selectivity were very similar to those obtained by Nefedov et a/. (203). b. Zeolite Y . Japanese workers have very recently prepared RhNaY catalysts (209) and compared their methanol carbonylation activity with rhodium supported on A1,0,, SO,-Al,O,, SiO,, and a cation-exchange resin (Amberlite 200 C ) . RhNaY was considerably more active than all the other catalysts, with rhodium or SiO, and Si0,-AI,O, being almost completely inactive for carbonylation. Three RhNaY catalysts were also prepared (209) by exchange with the chloride, nitrate, and sulfate salts of rhodium and shown to have very similar activities. The optimum rhodium level, with respect to specific activity, for RhNaY catalysts was found (209)to be about 0.6 wt. ‘%, Rh (4% replacement of Na’ assuming Rh3’), which corresponds to only approximately one rhodium site per unit cell of zeolite Y [i.e., an approximate stoichiometry of RhNa,o(A1O,),,(SiO,)l,,l. This implies that the average distance between

~

,

~

~

NONACID CATALYSIS WITH ZEOLITES

43

rhodium sites is about 25 A. Catalysts containing higher rhodium levels (2-7 wt. % Rh) not only led to decreased specific activity (i.e., carbonylation rate per gram of Rh), but also exhibited decreased stability. c. Kinetics and Mechanism. For both RhNaX (impregnated) (203)and RhNaY (210), the rate of carbonylation was found to be zero order in the CH,OH and CO partial pressures. At low CH,I/CH,OH ratios ( < 0.2), the rate of carbonylation is first order in the CH,I partial pressure for both X (203, 211) and Y (210) zeolite catalysts, thus leading to the rate expression

r

= KPCH31P:OP:H,0H

This is in fact analogous to that found for the homogeneous rhodium complexes (196,212). At higher CH,I/CH,OH ratios (> 0.2) there is evidence that, at least for the zeolite catalysts, the reaction order in CH,I decreases dramatically (203,211). Christensen et al. (211)also examined the effect of CH,I partial pressure on CH,OCH, formation over RhNaX. A rate expression of the following form was found : rether

=

Kp:2j

for CH,I/CH,OH, 0.05-0.23. This would seem to support earlier studies (203)on RhNaX which indicated that dimethyl ether formation was facilitated by methyl iodide. Activation energies for methanol carbonylation for RhNaX, RhNaY, polymer-supported RhCI(CO)(PPh,), ,and homogeneous RhCl, are compared in Table 111. There is clearly quite good agreement between these values for both heterogeneous and homogeneous catalyst systems. Thus the kinetic data indicate that similar reaction mechanisms are occurring. The above rate expression is also consistent with the previously discussed mechanism where the oxidative addition of CH,I to a Rh(1) complex is the TABLE 111 Compurison of Acriuution Energies f o r Methunol Curbonylution with Vurious Homogeneous and Heterogeneous Catalysts

Catalyst RhCI, (homogeneous) RhNaX RhNaY Polymer supported RhCI(CO)(PPh, ),

Activation energy (kJ mol-')

Reference

61.5

212

60 56.5 55

211 210 201

44

I. E. MAXWELL

rate determining step. The subsequent insertion of CO to form an acyl Rh(II1) complex implies that CH,I rather than CH,OH is carbonylated. In order to gain direct evidence of this or otherwise, Takahashi et al. (210) studied carbonylation over RhNaY using deuterated methanol (i.e., CD,OD/CH,I). Despite some slow alkyl exchange between CD,OH and CH,I, it could be demonstrated by analysis of the CH,CO/(CH,CO CD,DO) ratio as a function of methanol conversion that indeed there is not direct carbonylation of CH,OH and that this occurs via CH,I, which is regenerated at the end of the catalytic cycle. There is also direct spectroscopic evidence for the formation of a rhodium acyl complex as an intermediate during methanol carbonylation over RhNaX. Scurrell (213) showed that infrared bands characteristic of the complex CH,CORh(III)ICOL, were formed at 2085 and 1710 cm- ' after exposure of RhNaX to CO and CH,I at 100°C. With regard to specific activity, it is instructive to compare these rates for a variety of heterogeneous catalysts under similar conditions to those shown in Table IV. Clearly, RhNaY, prepared by ion exchange has the highest specific activity. This is indicative of a high degree of dispersion and good accessibility of these ions to reactant molecules. However, when this specific rate is compared with the best estimate (complicated by the strong pressure dependence in the range 0-2 atm) available from homogeneous catalysis

+

TABLE IV Comparison qf Specific A ctioities,for Mrlliutiol Curbonylcrtiun '

wirh Various Hereroyrnrous Rhodium Cutulysrs"

Catalyst

Activation energyh (kJ/mol)

RhCI,/A120, Rh-cation exchange resin RhNaX (impregnation (immersion) RhNaY (ion exchange) RhCI, (homogeneous)

Rh loading (wt. %)

58 58

0.43 0.37

58 60 56.5

0.5 0.6 0.43

61.7

-

Specific activity (g ester (g Rh)-' h - ' ) 4.3

8.0

21 17 65

-

300

CH,I/CH,OH

Reference

0.1 0.1

20Y

0. I 0.12 0.09

203 211 209

-0.1

209

212

" 200 C and 1 atm.

* The values for the activation energy have been used to correct rate constants to a common temperature of 200'C.

NONACID CATALYSIS WITH ZEOLITES

45

(212) using RhC1, , RhNaY heterogeneous catalyst is evidently inferior. Thus it would appear that for the zeolite, either the activity per rhodium site is intrinsically lower or only a fraction of the rhodium ions is accessible to the reactant molecules in the supercages. Further work will be required to elucidate this point. The relatively large average distance between rhodium sites (- 25 A) for maximum catalyst specific activity as found for RhNaY (210) strongly suggests that isolated rhodium species are the active entities. The first-order dependence on rhodium concentration suggests that this is also the case in the homogeneous system (196). However, by contrast, a second-order rate dependence of the Rh(1) complex concentration was found for polymersupported RhCI(CO)(PPh,), (201). A transition state involving two Rh(1) sites was invoked to explain these kinetic observations. d. Metal Promoters. Russian workers have recently screened a variety of transition metal oxides and chlorides for possible promoter effects on methanol carbonylation. The addition of < 1% CuCl,, FeCl,, NiCl,, CrC1, , or CoC1, to RhNaX is claimed (214) to increase the yield to methyl acetate and decrease ether formation. Iron oxide (Fe,O,) addition to RhNaX (215) was found to have an optimal promoter effect with a Rh/Fe ratio of 2 . Indeed, this catalyst gave a good selectivity to methyl acetate (94%) at quite a high activity level (126 g ester (g Rh)- hr- at 210°C). A similar promoting action was found for CuO (216) with an optimum at a Rh/Cu ratio of unity. The mechanisms of these promoter effects do not appear to be understood at present.

2. Other Metal Zeolites Nefedov et al. (217) screened a variety of first-, second-, and third-row transition metal ions (i.e., V, Cr, Fe, Co, Ni, Cu, Mo, Rh, Pd, Ag, Ce, Hf, W, Pt, Au), impregnated as metal salts (0.5 wt. % metal) onto NaX. With the exception of rhodium none of these metal ions showed significant methanol carbonylation activity.

B. ETHANOLAND HIGHER ALCOHOLS Carbonylation activity of ethanol using ethyl iodide as a promoter has been investigated using RhNaX as a catalyst (208,211). Some carbonylation activity does occur, with formation of ethyl propionate; however, ether and also olefin formation results in rather poor selectivities to the ester. The selectivity decreases with increasing ethanol conversion. Olefin formation was ascribed to a dehydrohalogenation reaction, i.e., CH,CH,I

-+

CH,=CH

+ HI

46

1. E. MAXWELL

The observed lack of an overall consumption of ethyl iodide was attributed to reaction of HI with ethanol (211), i.e., HI

+ CH,CH,OH

--c

CH,CH,I

+ H,O

A comparison of the absolute rates of methanol and ethanol carbonylation (211) indicated that the poor selectivity in the latter case is due to an increase in the rates of the side reactions rather than a large decrease in the rate of carbonylation. These results contrast with the homogeneous system, where ethanol carbonylation was reported (218) to be considerably slower (18 times) than with methanol. Nefedov et al. (219) found that Fe,O, impregnated on RhNaX had a promoting effect on ethanol carbonylation selectivity as was the case for methanol (215). Christensen et al. (208) were unable to carbonylate 2-propanol using RhNaX, but few details were given. Russian workers (220),however, showed that the carbonylation rate of higher alcohols could be markedly increased over RhNaX, by increasing the CO pressure. C. ETHYLENE Nefedov et al. (221)have reported that ethylene can be selectively hydrocarboxylated to form propyl propionate over NaX which had been exchanged with group VIII metal ions. The activity and selectivity decreased along the series Rh >> Pd > Ni > Co. In autoclave experiments they found that Rh, NaX (1% Rh) at 250°C and 60 atm CO pressure gave 100% conversion of ethylene in the presence of n-propanol and n-propyl iodide with a selectivity of 98.7% to propyl propionate, i.e., CH,=CH,

+ CO + HOCH,CH,CH,

1

c H , c H ,coocH,cH,cH,

n-Propyl iodide has a promoter function as for carbonylation and is regenerated at the end of the catalytic cycle.

VI.

Hydroformylation

Hydroformylation is now a well-established process in the chemical industry. The reaction involves the addition of CO and H, to an olefin yielding normal and iso-aldehydes :

NONACID CATALYSIS WITH ZEOLITES

RCH=CH,

+ CO + H,

-+

47

RCH2CH,CH0

+

CHO

I

RCHCH,

Industrially the straight chain isomer is generally the most desired product and hence the normal/iso product ratio obtained for a given catalyst is of importance. Further, the hydrogenation activities of catalysts vary considerably such that alcohols can in some cases be obtained in a single step (222).The first catalysts developed for this reaction were based on cobalt carbonyl and later cobalt carbonyl phosphine complexes. However, more recently attention has been focused on the intrinsically much more active rhodium catalysts (222,223). A simplified mechanism for (223)cobalt- and rhodium-catalyzed hydroformylation has been proposed which involves the following steps: (1) reaction of a neutral hydride carbonyl complex with an olefin to form an alkyl complex ; ( 2 ) insertion reaction with CO to form an acyl complex ; (3) reaction with H, to form aldehyde and regenerate the hydride carbonyl complex. The selectivity toward straight-chain aldehydes can be increased, for example, by incorporating bulky electron-donating ligands (e.g., PR,) into the metal coordination sphere. These products are also favored by low reaction temperatures and high CO partial pressures (222). The existing commercial hydroformylation processes are carried out using homogeneous catalysts (222, 223) with the associated disadvantages of catalyst product separation, metal deposition, and catalyst regeneration. This has prompted interest in methods of heterogenizing these homogeneous catalysts. Metal carbonyl complexes have been supported on various types of polymers, SiO,, functionalized S O , , carbon, and A1,0, (224). By comparison, there has been relatively little interest in the use of zeolites as supports. This is probably due to the obvious problems involved in incorporating a bulky metal carbonyl and/or phosphine complex into the relatively small pores of a zeolite. However, more recently there have been two publications (225,226)in which metal carbonyl clusters appear to have been formed in situ and hence encapsulated into zeolites, resulting in active hydroformylation catalysts.

A. COBALT ZEOLITES The study carried out by Centola et al. (225) was apparently stimulated by a patent (227)in which cobalt zeolites X and Y were claimed as catalysts

48

1. E. MAXWELL

i'i

f

0 0

is0

1

2

3

4

5

6 7 RUN TIME, hr

FIG.20. Rate of propylene hydroformylation with CoNaCaA catalyst as a function of run time. (Reproduced from Ref. 225 with permission from the authors and Lu Chimica P L'lndzistria.)

for the liquid-phase hydroformylation of C, and higher olefins. Centola et al. (225) prepared three types of catalyst by exchanging zeolites NaA, NaCaA, and NaX with a 1 N aqueous solution of cobalt(I1) nitrate, resulting in cobalt levels of 1-12 wt.%. Propylene hydroformylation was carried out in a continuous-flow reactor. The reaction was carried out in the gas phase in the pressure range 100-400 atm, and care was taken to avoid propylene condensation. The catalysts were simply pretreated by heating under vacuum at 45OoC,after which the feed (propylene/CO/H,) was introduced. Typically, an induction period was observed, followed by a transition phase of apparent high activity, and finally a steady state as shown in Fig. 20. The induction period was attributed to the reduction of Coz+ to Coo and formation of carbonyl complexes. This seems very plausible but no direct evidence of carbonyl formation was given. The transition period was considered to be due to a homogeneous catalyzed reaction caused by the removal of small amounts of metal carbonyl from the zeolite. Some typical results obtained for various Co zeolites under steady-state conditions are shown in Table V. Clearly, CoNaA and CoNaCaA display quite good propylene hydroformylation activity, whereas CoNaX is relatively inactive. In all cases the selectivity to total aldehydes was 99%. the other products being isobutyl and n-butyl alcohol. Cobalt losses from the catalyst after 60-h operation did not exceed 3%. This is perhaps not too unexpected under gas-phase reaction conditions. The proposed in situ formation of cobalt carbonyl complexes, which are encapsulated in the zeolite cavities, does seem to be plausible. The performance of the CoNaA and CoNaCaA zeolites in a gas-phase reaction was considered to offer a number of advantages over the existing homogeneous catalyst processes. However, it was noted that the specific activity of these heterogeneous catalysts (in moles aldehyde per moles

49

NONACID CATALYSIS WITH ZEOLITES

TABLE V Propylene Hydroformylation Data Obtuined Using Various Cobalt Zeolites"

Catalyst

fwt. %)

level

Temp. ("C)

Pressure (aW

Aldehyde yield

Normal/iso aldehyde

CoNaX CoNaA CoNaCaA

8.0 6.4 6.4

220 210 210

200 310 260

1.o 21.0 21.0

1.23 1.12 1.37

Conditions: propylene, 10 m o l x ; CO, 45 mol % ; H,, 45 mol Defined as moles aldehyde per 100 moles of propylene.

x;flow rate, 20 Nl/hr.

cobalt per hour) was lower than for the homogeneous analogs. This fact, together with the poor activity of CoNaX, suggests that a considerable percentage of the cobalt remains in regions of the zeolite framework (e.g., sodalite cages) which are inaccessible to reactant molecules. A recent patent (to UOP) (228) claims that active hydroformylation catalysts can be prepared by reacting an aluminated zeolite (prepared from a hydrosol) with HCo(CO), vapor. The HCo(CO), complex apparently reacts with surface hydroxyl groups releasing hydrogen and yielding a surface-bound cobalt carbonyl complex. The catalyst so formed is claimed to hydroformylate higher olefins to aldehydes and alcohols at 120°C and 240 atm pressure.

B. RHODIUM ZEOLITES A RhNaY (1 wt.% Rh) catalyst was prepared by ion exchange using an aqueous solution of [Rh(NH,),]CI, (226, 229), followed by treatment with a CO/H, ( l / l ) mixture at 130°C and 80 atm pressure. The catalyst so formed was observed to have good activity and high total aldehyde selectivity (-95%) for the liquid-phase hydroformylation of hexene-1. Some typical results are shown in Table VI, which indicate that the normal/iso aldehyde product ratio is similar to that obtained with homogeneous rhodium carbony1 catalysts (223). The catalyst was examined (226) before and after the pretreatment with CO/H, by means of infrared (IR) spectroscopy. IR bands characteristic of terminal and bridging carbonyl groups associated with rhodium carbonyl clusters were observed following this pretreatment of the catalyst. However, these spectra were apparently not identical to those of the known rhodium clusters, Rh,(CO),, or Rh,(CO),, . The differences were associated with the positions of the bands due to bridging carbonyl groups. It is conceivable that

50

I . E. MAXWELL

Pressure (atm)

Hexene-1 conversion (mol yo)

Normal, iso aldehyde ratio

> 80 > 90 > 95

0.8 I .o

50 80 100

1.1

Liquid-phase, autoclave experiments. Conditions: catalyst conc., 1.2 mg (at. R h ) - ' liter- I ; temperature, 80°C; run time, 3 h r ; solvent, hexane. a

I,

an alternative rhodium catalyst cluster is formed as a result of the spatial restrictions of the zeolite Y supercage (13 A diam.). The authors also claimed that the zeolite catalyst exhibited unusually high selectivity to dialdehydes (60%) in the hydroformylation of 1,5-hexadiene. Unfortunately, comparative data obtained under similar conditions with homogeneous rhodium carbonyl catalysts were not presented. A patent (230) to Atlantic Richfield Co. claims that hydride platinum group metal carbonyl complexes such as CIRh(PPh,), supported on zeolites, for example, N a y , are suitable catalysts for the hydroformylation of low molecular weight olefins. However, since the bulky metal complex cannot diffuse into the inner pores of the zeolite it must simply be adsorbed on the external surface of the support. This is consistent with the rather poor catalyst stability which was attributed to leaching of the active species from the support.

VII.

Methanation

The synthesis of methane from synthesis gas (CO/H,) is of considerable importance in the production of substitute natural gas (SNG) i.e., 3Hz

$-

CO -+ CH,

+ HZO

In fact, with the expected trend toward upgrading coal and the depletion of United States natural gas reserves, methanation should further increase in importance. The methanation reaction is thermodynamically favorable even at high temperatures and pressures and there are a variety of metals which catalyze this reaction. Vannice (231) compared the turnover numbers for a variety of Al,O,-supported group VIII metals and showed that the rate of

NONACID CATALYSIS WITH ZEOLITES

51

methane formation could be correlated with the heat of adsorption of CO (see Fig. 21). For these metals the rate of methanation increased with decreasing heat of CO adsorption, giving the following order of activity : Ru > Fe > Ni > Co > Rh > Pd > Pt > Ir. Although nickel is not the most active metal, for reasons of cost and stability, commercial catalysts are based on this metal. These catalysts are in general supported on alumina and contain relatively high metal loadings. There is, however, interest in the development of more thermally stable, sulfur resistant, and possibly even regenerable catalysts in the future. It may be for these reasons that zeolite-based methanation catalysts have recently attracted more interest. A. PALLADIUM ZEOLITES Vannice (232) measured turnover numbers for methanation on a variety of well-characterized palladium catalysts. The supported catalysts were all more active than unsupported palladium (Pd black) and PdHY was intermediate between Pd/SiO, and Pd/Al,O, in specific activity (see Table VII). As is evident from the data there is no obvious correlation betwen particle size and turnover number. It was therefore suggested that the enhanced activity of various supported catalysts was due to a metal-support interaction. Figureas et al. (105) also found good evidence for a support effect during benzene hydrogenation studies. In this case the palladium zeolite

CO heat of adsorption (kcallrnol)

FIG.21. Correlation between methanation activity and AHa for CO. (Reproduced from Ref. 231 with permission from the author.)

52

1. E. MAXWELL

TABLE VI1 Comparison oj'A4erhanaiionTurnovrr-Numbers f o r Various Palladium Catalysts

Catalyst Pd/AI,O,

PdHY Pd/SiO, Pd black

Pd loading (wt. %)

2 9.5 2 0.5 4.75 4.15 -

Turnover number (sec-' x lo3, 275-C)

Average particle size

12 10 7.4 5.9 0.32 0.26 0.15

48 I20 82 31 28 46 2100

(8)

catalysts were more active than palladium supported on silica and alumina. It was proposed that the palladium turnover number increased with increasing electron-acceptor properties of the support. A similar mechanism may be applicable to palladium-supported methanation catalysts (232). B. NICKELZEOLITES The methanation activity of a series of NiCaY catalysts was recently studied by Bhatia eta(. (233).These authors found that the turnover numbers increased with increasing metal loading, whereas the average particle size remained constant and they attributed this result to increased support acidity and availability of Ni'. Unfortunately, the degree of reduction of Ni2+ to Nio does not appear to have been measured, which might also explain the results obtained. Elliott and Lunsford (234)more recently measured the methanation activity of NiNaY (2 wt. % Ni) and found this to be considerably less active than Ni/AI,O,. The turnover numbers obtained by various workers are compared in Table VIII. It is apparent that the nickel/ zeolite catalysts are significantly less active than the Ni/AI,O, catalysts. As indicated in Table VIII this does not appear to be a particle-size effect. Elliott et al. (234)have proposed that the alkali metal ions (in this case Na') of the zeolite may be responsible for this decreased activity. However, it is noteworthy that the relatively large average nickel particle size (140 A) means that most of the nickel is on the external surfaces of the zeolite crystals where a relatively low concentration of alkali metal ions might normally be expected.

53

NONACID CATALYSIS WITH ZEOLITES

TABLE VIII Comparison of Methanation Activities f o r Zeolite and Alumina Supporied Nickel Catalysls

Catalyst NiCaY NiNaY Ni/AI,O,

Nickel loading (wt. %) 6.8 2 2 5

Average particle size

(A) 7.8 140 360

-

Reaction temp. ("C)

Turnover number (sec-' x lo3)

Reference

300 280 280 275

7.68 5.0 230 32

233 234 234 23 I

C. RUTHENIUM ZEOLITES Elliott et al. (234) very recently carefully prepared and thoroughly characterized a variety of ruthenium methanation catalysts supported on zeolite Y. The purpose of the study was to prepare the catalysts in such a way that the metal remained finely dispersed within the zeolite cavities. For such a catalyst, a significant metal-support interaction might be expected to occur ( 6 , 2 3 ) and thereby induce a change in catalytic behavior. The methanation turnover numbers for RuNaY and RuCaY are compared with Ru/A1,0, in Table IX. The results show that the very high ruthenium dispersion achieved for the zeolite catalysts does not appear to have had very much effect on the methanation specific activity. However, the RuY and RuCaY catalysts were more stable during the methanation reaction than Ru/Al,O, . The deactivation process was attributed to the formation of excess surface carbon via dissociation of CO.

TABLE IX Comparison of Initial Methanution Activities f o r Zeolite und Aluminu Supported Ruthenium Cutulysts

Catalyst RuNaY RuCaY Ru/A1,0,

Ruthenium loading (wt. %) 0.5 2 0.5 0.5

Average particle size

(A)

Turnover number (sec-' x lo3, 280°C)

9.8 10 9.5 43

9.4 31.5 15.9 19.8

54

1. E. MAXWELL

The same workers (234) also studied the methanation behavior of bimetallic clusters of Ru/Ni and Ru/Cu in zeolite Y. Such clusters can be formed by metals, such as ruthenium and copper which are immiscible as bulk metals (235, 236). The turnover numbers versus bimetallic cluster composition are shown in Fig. 22. Dilution of ruthenium with copper clearly causes a marked decrease in specific activity. This decrease in activity is also accompanied by a decrease in methanation selectivity. This was attributed to an inhibiting effect of copper on the ruthenium hydrogenolysis activity. The addition of nickel to ruthenium has a less pronounced effect on the methanation activity. This is hardly surprising since nickel is also intrinsically active for methanation. However, dilution of ruthenium with nickel does result in a marked increase in catalyst stability. A catalyst of composition 0.5% Ru, 2% NiY was more stable than those prepared from the pure metals. This improved stability was attributed to an improved balance between the rates of dissociation of CO and hydrogenation of surface carbon, thereby preventing the formation of excess surface carbon. The data presented indicated that a similar improvement in stability was obtained for 0.5% Ru, 2% Ni on A1,0,, which demonstrates that this effect is not support sensitive. Of particular interest was the fact that the bimetallic cluster catalysts, i.e., RuNiY and RuCuY, had considerably better metal dispersions than the pure NiY and CuY catalysts. Further, the zeolite-supported bimetallic catalysts were more resistant to sintering during methanation than those supported on alumina. Particle-size measurement indicated, however, that most of the bimetallic clusters were too large to be located inside the zeolite pores. Gupta et al. (237)studied the effect of in siru ?-irradiation on the methanation activity of RuNaX (1.8 wt.% Ru). An enhancement in activity was found on y-irradiation of the catalyst. This enhanced activity was attributed to an increase in the rate of hydrogenation of surface carbon. No comparative data were presented for other support materials. Unfortunately, there have been no published studies on the poison resistance of zeolite-based methanation catalysts. NiY catalysts containing Cr,O, , for example, have been shown (238)to exhibit much improved sulfurpoisoning resistance under ethane hydrogenolysis conditions. This potential advantage of zeolites would seem worthy of further study. In addition, bimetallic cluster catalysts might offer improved activity without the disadvantage of high rate of deactivation due to the formation of excessive surface carbon. There is some evidence that the zeolite support enhances the stability of these small metal clusters. The recently discovered high silica zeolites (239) of the type ZSM-S/ZSM-ll might be of interest as support materials, due to their high hydrothermal resistance.

NONACID CATALYSIS WITH ZEOLITES

TpL 6 :

’

z L 4

:

.

E 2

Ru/Ni :

55

H

Ru/Cu

8

0 100 80

60 40 20 Atom Ru

%

0

FIG.22. Turnover number versus atom percent ruthenium in (a) ruthenium-nickel zeolites and (b) ruthenium-copper zeolites. (Reproduced from Ref. 234 with permission from the authors.)

VIII.

Conversion of Synthesis Gas to Hydrocarbons

It is beyond the scope of this review to cover the rapidly expanding topic of synthesis gas conversion to hydrocarbons in any great detail. However, it would be an omission not to mention a number of aspects which are at least related to nonacid catalysis. The current world shortage of cheap crude oil has stimulated intense research activity to develop commercially viable processes to convert coal-derived synthesis gas (CO/H,) into liquid fuels and petrochemical feedstocks. The synthesis of hydrocarbons from synthesis gas via the Fischer-Tropsch (FT) (240) reaction has been known for some 50 years and has been successfully applied in both Germany and South Africa. However, this catalytic reaction occurs by a well-defined chain growth mechanism which has the disadvantage of yielding a very broad product distribution. The Schulz-Flory polymerization kinetics, which accurately describe this process, show that the maximum attainable selectivity to a product in the gasoline range is only 48%. Zeolites offer enormous potential as catalysts in this area since the limitations of Schulz-Flory kinetics can be overcome by utilizing the shape-selective properties of these support materials. Initial studies using cobalt-, nickel-, and iron-modified zeolites X and Y (241, 242) were, however, not particularly encouraging with relatively poor activities, selectivities, and stabilities. This situation has now changed dramatically with the discovery by Mobil Oil Corporation of a new series of synthetic high-silica zeolites. The so-called ZSM-5 zeolite (in the H form) is capable of converting methanol quantitatively to hydrocarbons and water (239), i.e., xCH,OH

-+

(CH,),

+ xH,O

56

1. E. MAXWELL

where the hydrocarbons are in the range C, to C,, and the gasoline fraction (rich in aromatics) has a high research octane number (90-100). This novel acid catalysis is not limited to methanol as reactant, as is demonstrated by the fact that a wide variety of oxygenate, olefin, and paraffin feedstocks can also be upgraded to a product boiling in the gasoline range ( 2 4 2 ~ )The . constrained structure of the zeolite (as shown in Fig. 23) is directly responsible for the shape-selective process which yields a narrow range of product molecular weights. Another important feature of this zeolite catalyst is the relatively low rate of coke formation. The reasons for this improved stability do not seem to be completely understood as yet, but are very likely to be related to the specific zeolite structure which probably inhibits the formation of coke precursor molecules. Derouane has recently described this behavior (242b) as an example of restricted transition-state selectivity. Clearly, with this development in mind, it was an obvious step to combine a CO reducing function with such a zeolite and thereby produce a single-step shape-selective catalyst for synthesis gas conversion to gasoline. This has been achieved in a number of ways, for example by impregnating the active elements of conventional FT or methanol synthesis catalysts into ZSM-5 (e.g., Fe, Ru, T h o , , HfO,, Zn, Zn/Cr, and ZrO,) (243-247). The latter oxide component is claimed (246)to result in a catalyst with much improved resistance to sulfur poisoning. In general, these bifunctional catalysts result in much improved aromatics production and a very substantial reduction in heavier (C :,) hydrocarbons than would be obtained for the CO reduction catalyst alone. This component need not necessarily be incorporated into the zeolite itself, since physical mixtures of CO-reducing catalysts and zeolites have also been shown to be effective for selective synthesis gas conversion (243-247). In addition, this physical mixture may also include a water-gas shift catalyst (248, 249).

FIG. 23. A schematic diagram showing the structure of zeolite ZSM-5. (Reprinted with permission from Ref. 239. Copyright 1976 American Chemical Society.)

NONACID CATALYSIS WITH ZEOLITES

57

Jacobs has described ( 2 4 9 ~these ) two different approaches in terms of secondary (physical mixtures) and primary (FT function in zeolite matrix) effects. In the former case the results obtained can be quite well understood in terms of the separate behavior of each component. However, in the latter case the results may be different since the primary FT products are formed inside the spatially restricting pores of the zeolite. In a recent communication (250), deviations from Schulz-Flory kinetics were observed for a RuNaY synthesis gas conversion catalyst (see Fig. 24). A comparative catalyst, prepared by impregnating silica with ruthenium, i.e., Ru/SiO,, and tested under the same conditions, yielded a product distribution which gave a good fit to Schulz-Flory kinetics. The sharp decrease in chain growth probability for C:, products over RuNaY is perhaps surprising for such a relatively large-pore zeolite. Further studies (251-253) on this system indicated that there was a correlation between the ruthenium particle size in the zeolite and the product distribution.

FIG. 24. Coke yield versus shape selectivity of paraffin conversion for various zeolite catalysts. (Reproduced from Ref. 266 with permission from the authors.)

58

1. E. MAXWELL

This work is of particular interest if indeed an alternative mechanism exists for zeolite catalysts, other than pore-restricted shape selectivity, which may be employed to avoid broad product distributions in synthesis gas conversion reactions. In this respect it is of interest to note that in a recent communication (254), non-Schulz-Flory kinetics were reported for a series of cobalt hydrocarbon synthesis catalysts. The catalysts were prepared by depositing cobalt carbonyl on aluminas of different pore diameters and surface areas. A correlation was found between pore diameter and product distribution. Shape-selective effects for supports with pore radii in the range 65-3000 (i.e., outside the configurational range of diffusivities) do not seem to be very probable. However, a metal particle-size effect of the type proposed for RuNaY (250-253) may be a possibility. It should, however, be mentioned that in a study carried out by King ( 2 5 4 ~where ) ruthenium FT catalysts were prepared on a variety of supports, including zeolites X and Y, no correlation was observed between product chain length and metal particle size. Further work in this area will be necessary to confirm these proposed particle-size effects. The resistance of such small metal particles to sintering is, of course, of crucial importance to any practical application of such catalysts. Other recent examples of shape-selective FT catalysts include Cd-vapor reduced CoA zeolite (2546) and ion-carbonyl complexes incorporated into zeolite Y ( 2 5 4 ~ ) . Clearly, this field of research is in its infancy and the high level of current activity should lead to considerable further development.

IX.

Miscellaneous Reactions

A. WATER-GAS SHIFT The water-gas shift reaction, i.e., H,O

+ CO

Z?

CO,

+ H,

is of importance in adjusting a coal-derived synthesis gas to a product of more preferable composition. More active and more stable catalysts would enable increased productivities per reactor. In this respect a recent communication (255), which presented data indicating that RuX and RuY were more active for water-gas shift than conventional copper-based catalysts, is of interest. Infrared spectroscopy and temperature programmed reduction studies indicated that the active species was a complex within the zeolite pores of the form [Ru(NH,),(OH),(CO),]"~ ( n < 3) where the ruthenium is in the oxidation state Ru(1) or Ru(I1) or both. An interesting aspect was that

NONACID CATALYSIS WITH ZEOLITES

59

there is apparently no known homogeneous equivalent of this complex which is stable up to 250°C. However, it seems doubtful as to whether this heterogeneous catalyst would be sufficiently stable under the more severe commercial operating conditions.

B. KOLBEL-ENGELHARDT REACTION The less familiar Kolbel-Engelhardt reaction, i.e., 4CO

+ 2H20 F? CH, + 3C0,

is considered to be a two-step reaction involving a water-gas shift followed by hydrogenation of CO to produce methane and/or higher hydrocarbons, i.e., CO

+ H 2 0 + CO, + H,,

CO

+ 3H2 @ CH4 + H 2 0

Conventional catalysts are based on iron/iron oxide mixtures (256). Recently, Lunsford and co-workers (257) have shown that reduced RhNaY is also an active catalyst for this reaction at 329°C. A high selectivity to methane was observed and there was evidence of a two-step reaction scheme as shown above. Unfortunately, no comparative data were presented under similar conditions for a conventional catalyst. C . WATERSPLITTING

There is considerable interest in developing an economical thermochemical reaction cycle that would effect the decomposition of water into hydrogen and oxygen (i.e., water splitting). Such a cycle could in principle provide a method of manufacturing hydrogen, a nonpolluting fuel, from an unlimited source. Most processes proposed to date, however, involve a large number of reaction steps and require the use of highly corrosive chemicals (258). A few papers have appeared recently wherein water splitting using various cation-exchanged forms of zeolites has been demonstrated. The reaction schemes, in general, can be written as follows: 2M(H20):' M("-')+

+ 220- &2M'"-

')+

+ 2ZO-H" + )O,

+ ZO-H+ 9 M(H,o):+ + zo- + +H,

(1)

(2)

It appears that quite a number of metal ions, M"', can effect reaction (l), which Jacobs et al. (259) have termed zeolite autoreduction. Kasai et al. (260) have proposed that for reaction (2) to proceed, the reduction potential

60

1. E. MAXWELL

for the half-cell reaction, i.e., M"+

+

e-

+M("-l)+

must be more negative than the proton reduction half-cell reaction, i.e., H+

+ e-+fHz

(E" = -0.414V)

In this way these authors were able to explain the fact that although reaction (1) involving Cu2+/Cu+ ( E o = +0.158 V) proceeded for Cu2+-exchanged mordenite, reaction (2) could apparently not be accomplished. Further, for the systems Cr3+/Cr2+(Eo = -0.41 V) and I n 3 + / I n 2 +(Eo = -0.49 V) both reactions (1) and (2) were shown to occur. To facilitate reaction (2), ion exchange was carried out with H mordenite, thus increasing the concentration of 2 0 - H + . Repeatable two-step water thermolysis cycles were apparently achieved in both cases. Jacobs et af. (261) have shown that photolysis of water can be achieved using AgY zeolite. Irradiation using sunlight on water vapor-saturated AgY yields oxygen, and thermal treatment (600 C ) of the reduced silver zeolite restores the Ag+ cations with the evolution of hydrogen, i.e., 2Ag'

+ 2ZO' + H z O Ago

2Ag0

---+

+ ZO-H+ S

A g '

+ 2 Z O - H + + $0,

+ ZO- + i H ,

The system was, however, unstable due to silver metal sintering and dehydroxylation of the zeolite. Both these reactions are of course favored by the relatively high temperature required for the oxidative desorption of hydrogen. Interestingly, the above reaction is nonallowed according to the half-cell potential criteria of Kasai el a/. (260) (i.e., EoAg+(aq)/Ago= +0.7996 V). Leutwyler and Schumacher (262) have, however, pointed out that the standard redox potential where both species are hydrated is negative [i.e., EoAg+(aq)/Ago(aq) = - 1.8 V]. The real situation in the zeolite is probably somewhere intermediate between these two extremes and could thus satisfy the previously described criteria. The above workers also calculated that the photolytic oxygen-producing step in silver zeolites must involve two photochemical steps since the photon energy available from sunlight is insufficient for a realistic single-step process. Similar problems with regard to reversibility of the water-splitting cycle, such as silver agglomeration were also found by these workers. Kuznicki et uf. (263) have shown that zeolite 3A exchanged with Ti3+ is also capable of photolytic splitting of water. The ion-exchanged zeolite was shown to yield hydrogen in water under illumination with visible light. ESR studies indicated that the active entity for the hydrogen production was a higher valency state tetrahedral titanium(1V) oxygen complex with an

NONACID CATALYSIS WITH ZEOLITES

61

unpaired electron shared equally by the oxygen atoms. This is rather surprising since in all the previously discussed systems the reduced form of the metal ion is the active entity in the hydrogen production step. The second step, required to complete the water-splitting cycle, normally accomplished by heating with the evolution of oxygen, could not be demonstrated for the titanium-exchanged zeolite A system. Further, Ti3-+-exchanged zeolites X and Y were found to be inactive for photolytic hydrogen production from water. This is perhaps somewhat unexpected since these zeolites are also known to form titanium oxygen complexes (264) containing an unpaired electron. Kuznicki et al., however, infer in their paper (263)that the photoformed radical in TiA contains dissociatively adsorbed oxygen, whereas for TiY there is good evidence that the radical is a nondissociatively adsorbed superoxide species (264). Such a major difference in the modes of oxygen adsorption could possibly explain the observed results. In conclusion, the chromium and indium water thermolysis systems described (260)and patented (265)by Union Carbide workers do appear to be the most promising zeolite-based schemes in terms of providing a stable process cycle. The simplification of a two-step process provided by the zeolites is clearly advantageous compared to the alternative many-step processes. As explained by Kasai et d . (260)this simplification is achieved with zeolites as a result of the large entropy change (> 120 e.u.) associated with the endothermic first step. Such a large entropy change is attributed to the large number of water molecules involved in reaction (1). Alternative watersplitting schemes in general require many more steps in order to achieve the net increase in entropy (40 e.u.) associated with this reaction, at reasonable temperatures (258).Further research on this catalytic application of zeolites would seem worthwhile in view of the promising results so far obtained and the projected economics. Thermochemical production of hydrogen has been projected to compete with coal gasification and electrolytic processes in the latter part of this decade (258).

X.

General Conclusions and Future Prospects

A. ACTIVITY AND SELECTIVITY The often quite remarkable activities and selectivitites achieved using various cation-exchanged zeolites can be best illustrated by using highlights from the foregoing discussion. Selective oxidation of ethylene to acetaldehyde can be achieved using I’d2’, Cu2+-Y (46, 47) and as yet no other heterogeneous Wacker-type catalyst has been reported. A novel n-hexane

62

I . E. MAXWELL

dehydrocyclization catalyst can be prepared by reducing tellurium which has been incorporated into NaX (127). RhNaY catalyzes the dimerization of ethylene under extremely mild conditions to selectively yield n-butene (dimer) products (140). CrNaY is a very active catalyst (140) for the polymerization of ethylene yielding a high-density unbranched polyethylene product. NiNaY selectively catalyzes the trimerization of acetylene to benzene (159, 160) rather than the tetramer, as occurs when simple nickel salts are used in homogeneous catalysis. Small-pore alkali metal ion zeolites such as KA selectively dimerize cyclopropene to tricyclohexane (166), whereas larger pore zeolites (e.g., NaX, N a y ) and solid acid catalysts result exclusively in polymerization. Transition metal ion-exchanged zeolites catalyze the dimerization of butadiene (1 72-18]) over a broad temperature range, but in constrast to homogeneous catalysts 4-vinylcyclohexene is the sole product over zeolites. NiNaX dimerizes n-butene (185, 186) with high selectivity to octane products. Carbonylation of methanol to methyl acetate can be achieved with very high specific activity and selectivity using RhNaY (209,210) in the presence of an iodide promoter. The specific activity of this zeolite catalyst exceeds that of any other heterogeneous carbonylation catalyst to date. RuNaY is an active catalyst for synthesis gas conversion (250) and yields a product distribution which markedly deviates from Schulz-Flory kinetics as is normally observed for the more conventional Fischer-Tropsch catalysts. The relatively high activities of these catalysts can in most cases be attributed to the high dispersions of the active species. These are normally incorporated as cations via an ion-exchange process and thus remain bound onto the extensive inner surface of the zeolites by electrostatic forces. The selectivities observed, for example, in oligomerization reactions where, in general, dimers are formed in preference to higher oligomers, may be a direct consequence of the spatial limitations imposed on transition-state complexes within the small zeolite cavities.

B. STABILITY It is perhaps not sufficiently widely recognized that for the successful development of a new industrial catalyst, good stability may often be just as critical as activity and selectivity. Further, in order to achieve economically acceptable space time yields, catalysts are generally run at high work rates (gram product per gram catalyst per hour). For many processes this places a heavy heat/mass transfer load on the catalyst which often accelerates deactivation processes. Although there is, in general, only very limited information on zeolite

NONACID CATALYSIS WITH ZEOLITES

63

catalyst stability in the literature, a number of quite different decline mechanisms have been observed for the systems discussed. For example, the stability of Pd2+-Y for partial oxidation of alkenes is rather poor. This can be considerably improved (46, 47) by incorporating Cu2+ into the zeolite, which increases the rate of reoxidation of PdO back to the catalytically active Pd2+ ion. Although this is, in fact, analogous to the homogeneous Wacker catalyst system, it is nevertheless indicative of facile electron transfer reactions within the cavities of a zeolite. The deactivation of the Cu' -Y butadiene cyclodimerization catalyst was attributed to polymerization reactions catalyzed by the zeolitic acid sites. The marked improvement in stability achieved (177) by the chemisorption of NH, on these sites is consistent with this deactivation mechanism. A major problem to be anticipated with heterogeneous cobalt and rhodium hydroformylation catalysts is leaching of the active metal from the catalyst surface, particularly during liquid-phase operation, with a consequent decline in activity. The quite good stability of CoNaA and RhNaY for olefin hydroformylation (226) may be due to the encapsulation of metal carbonyl clusters in the zeolite cavities. For the RhNaY catalyst a comparative experiment using a Rh/carbon catalyst, under the same reaction conditions, confirmed the improved stability of the zeolite for this reaction. A RuNiY methanation catalyst (234) was found to have much improved stability over that obtained for NiY and RuY. This was attributed to the formation of small bimetallic clusters which provided a better balance between the rates of CO dissociation and hydrogenation of surface carbon, thereby substantially reducing the rate of formation of deactivating surface carbon. There was also evidence that improved dispersions of the bimetallic catalysts were obtained on the zeolite support compared with, for example, alumina. Moreover, the zeolite-supported bimetallic clusters did appear to be more resistant to sintering. The much-improved stability of catalysts based on ZSM-5 zeolites for synthesis gas, and methanol conversion reactions is due to markedly reduced rates of coke formation compared to many other zeolites. Rollmann and Walsh (266) have recently shown that for a wide variety of zeolites there is a good correlation between shape-selective behavior, as measured by the relative rates of conversion of n-hexane and 3-methylpentane, and the rate of coke formation (see Fig. 24). This correlation was considered to provide good evidence that intracrystalline coking is itself a shape-selective reaction. Thus, the rather constrained ZSM-5 pore structure exhibits high shape selectivity, probably via a restricted transition-state mechanism (242b), and therefore has a low rate of coke formation. Zeolite composition and crystal size, although influencing coke formation, were found to be of secondary importance. This type of information is clearly

64

I. E. MAXWELL

invaluable in the choice of zeolites for reactions where coke formation may occur. This will normally only be the case if acid sites are also present in the zeolite to catalyze the formation of coke precursors (267). There is a very definite need for more fundamental work of this type to provide a better understanding of decline behavior, which is so essential to designing zeolite catalysts with sufficient stability. SPECIES AND REACTION MECHANISMS C. ACTIVE The crystallinity of zeolites has the advantage that they lend themselves rather more than conventional heterogeneous catalysts to the study of active sites. This is exemplified in the literature where the active species in zeolites for a wide variety of reactions have been identified. There are fewer examples where the reaction mechanisms have been unambiguously defined, although they can often be inferred from studies on analogous homogeneous or heterogeneous catalysts. For example, oxygen-bridged species have been shown (18,21-26,51,65) to play an important role in transition metal ion-exchanged zeolite oxidation catalysts. However, a detailed mechanistic scheme for most of these oxidation reactions is not yet available. By contrast, a very detailed picture has been obtained (128-130) of the active site in a TeNaX dehydrocyclization catalyst which involves a telluride ion coordinated to two Na' ions inside the supercage. Further, the mechanism was shown (131) to be similar to conventional dehydrocyclization catalysts and thus involved a stepwise dehydrogenation scheme. The active sites for transition metal ion-containing zeolite Y ethylene dimerization catalysts are very likely ds metal ions, i.e., Rh' (112), Ni2' (1.39, and P d Z +(140). In the case of RhNaY three different rhodium oxidation states could be distinguished using XPS (112), and the dimerization activity was found to correlate well with the concentration of Rh' species. This is a particularly interesting demonstration of the potential that this technique would seem to offer in defining active oxidation states in zeolites. A number of transition metal ion-exchange zeolites are active for acetylene trimerization (159, 160), and the criterion for activity appears to be an even, partially filled d-orbital, i.e., d H(Ni", Co'), d" (Fez'), d4 (Cr"). This has led to the suggestion that the mechanism must involve a complex in which there is simultaneous coordination of two acetylene molecules to the transition metal ion. The active oxidation state for CuNaY butadiene cyclodimerization catalysts has been unambiguously defined as monovalent copper (172-180). The d" electronic configuration of Cu' is consistent with the fact that isoelectronic complexes of Nio and Pdo are active homogeneous catalysts for this reaction. The almost quantitative cyclodimerization selec-

NONACID CATALYSIS WITH ZEOLITES

65

tivity to 4-vinylcyclohexene exhibited by copper zeolite catalysts is indicative of a reaction mechanism which proceeds via a o-allyl, n-allylbutadiene complex. Single-crystal X-ray diffraction studies ( I 79) have shown that copper ions migrate toward sites I1 and I11 in the supercages on adsorption of butadine into faujasite. The favorable location and unsaturated coordination geometry of the cations at site 111 led to the proposal that these were the active sites. The active species for rhodium zeolite carbonylation catalysts (203, 210, 211) is very likely a four-coordinate d8 complex of the type [Rh'L,]+Zwhere the ligands, L, could be oxygen anions of the zeolite framework and/or reactant molecules (e.g., methanol). Significantly, maximum specific activities are obtained for low metal loadings (- 0.6 wt.% Rh) (210), suggesting that these sites are on average relatively far apart ( - 2 5 A). There is good evidence that the reaction mechanism on the zeolite catalysts is very similar (203,210,211,213)to that found for homogeneous catalysts. For hydroformylation over cobalt and rhodium zeolites the active species have not been defined. However, in the case of RhNaY the in sizu formation of a rhodium carbonyl cluster has been identified (226) by infrared spectroscopy. Interestingly, this cluster appears to be different from known compounds such as Rh,(C0),2 and Rh,(CO),,. This does suggest that alternative carbonyl clusters may possibly be formed in zeolites due to the spatial restrictions of the intracrystalline cavities. The mechanism of hydroformylation in these zeolites is probably similar to that known for homogeneous catalysis. A complex of the type [Ru(NH,),(OH),(CO),]"+ (n < 3) has been proposed (255) as the active species for catalysis of the water-gas shift reaction by RuX and RuY zeolites. Interestingly, there does not appear to be a homogeneous analog to this complex. The active sites for catalyzed water splitting over zeolites involve both metal cations and protons sites in the zeolite pores. It has been proposed (260)that suitable zeolites must contain metal ions with a standard reduction potential (i.e., M"+/M("- 'I) which is more negative than the proton reduction potential (i.e., - 0.414 V). Furthermore, the redox cycle involving this metal ion must be reversible in the zeolite. Redox couples involving Cr3+/Cr2+ and In3+/In2+satisfy this criterion and accomplish water thermolysis with good cycle repeatability. The two-step mechanism first involves the dehydration and reduction of the metal ions with water splitting to form protons on the surface of the zeolite and the evolution of oxygen. The second step involves the reoxidation and rehydration of the metal ions with simultaneous deprotonation of the zeolite surface and the evolution of hydrogen. From the foregoing discussion it is evident that in terms of active species and reaction mechanism there is often a close parallel between ion-exchanged

66

I . E. MAXWELL

zeolites and homogeneous catalysts. This is particularly true for reactions such as alkene partial oxidation, alkene oligomerization, acetylene and butadiene coupling reactions, carbonylation, and hydroformylation. In a sense, the zeolite can be regarded as a large, rigid anion, balancing the positive charge of the metal ion or complex, where the oxide anions of the zeolite framework may or may not participate in the metal ion coordination sphere. There is some evidence to suggest that the zeolite framework stabilizes such metal ions in unusual coordination geometries and oxidation states and in addition results in a very high degree of metal ion dispersion. The thermal stability of most zeolites is such that high-temperature exothermic regeneration procedures (e.g., polymer or coke burn-off) are often quite feasible.

D. FUTURE PROSPECTS To gain some insight as to how the subject of nonacid zeolite catalysis has advanced in recent years, it is instructive to read, for example, the sections on olefin oligomerization and carbonylation in the 1968 zeolite catalysis review by Venuto and Landis (268). The activities, selectivities, and stabilities of the zeolite catalysts (usually in the acid form) used for these reactions were extremely poor. By comparison, as has been discussed, a number of metal ion-exchanged zeolites are now known to catalyze these reactions under mild conditions with high selectivities and good stabilities. There is every reason to believe that this trend will continue, particularly if cross-fertilization is maintained between the disciplines of zeolite and homogeneous catalysis. A large number of publications on zeolite catalysis are unfortunately of only limited value in that the catalysts are often not well characterized, or the reaction conditions are not well defined. In addition, there is often no indication of catalyst stability. Studies in which zeolite catalysts are compared, under similar conditions, with conventional catalysts (where these exist) are invaluable in ascertaining the role of the zeolite support itself. I t is encouraging to observe that there does seem to be an increase toward this comparative approach. To the author’s knowledge, there are at present no major industrial processes which could be strictly defined as nonacid catalysis that make use of zeolite-based catalysts. This is in contrast to acid catalysis where zeolites continue to make an impact. Technically, a number of zeolite-based catalysts for reactions, such as Wacker chemistry and olefin or diolefin oligomerization reactions, appear to be quite attractive, and it is almost certainly economic factors that have limited further development. The discovery of a new family of shape selective zeolites by Mobil(239) has now extended the range of pore sizes and thus the accessible range of configurational diffusion. To date the majority of studies have been carried out

NONACID CATALYSIS WITH ZEOLITES

67

using the more familiar large-pore zeolites such as X and Y. Comparative studies using zeolites with different pore structures and dimensions should provide a better insight into shape-selective and stability behavior. Measurements of both diffusion and reaction rates for these types of catalysts would be invaluable in elucidating the mechanisms by which shape selectivity occurs in such systems. Zeolites also lend themselves particularly well to reactions where the catalytically active species are cationic. Under these circumstances there is a strong electrostatic interaction between the active entity and the support which minimizes activity loss via leaching processes. The following areas would, in the author’s opinion, seem worthy of further study in the field of zeolite nonacid catalysis. A systematic investigation of shape selectivity in olefin and diolefin oligomerization reactions, by incorporating a common active component such a s rhodium into a variety of zeolite structures would provide useful additional information. There are preliminary indications (226) that unusual metal carbonyl cluster formation may occur in zeolite cages. This would also seem to be particularly interesting and has obvious relevance not only to hydroformylation but also to synthesis gas conversion reactions in general. The zeolite-based heterogeneous methanol carbonylation catalysts exhibit exceptional specific activity (209, 210), but comparison with the homogeneous system indicates that this might be further improved. The encapsulation of an effective halide promoter in the zeolite cavities of such a catalyst would remove the need to recycle these corrosive species in existing processes. There does seem to be some quite good evidence for the existence of small electron-deficient metal clusters in zeolites (101-105), which may be related to their increased resistance to sulfur poisoning. Further studies are, however, necessary in order to provide a more detailed understanding in this area. Zeolites also appear to be suitable supports for the formation of small bimetallic clusters between components which are immiscible as bulk metals. The Ru/Ni zeolite Y methanation catalyst is an interesting example (234) where the properties of a single active metal component could be advantageously modified. The initial results obtained (258-263) using trivalent ion-exchanged zeolites as catalysts for thermolytic water splitting are quite encouraging. The simple two-step cycle, the good stability, and noncorrosive properties of the zeolite are all positive aspects. Further research toward a zeolite-based system which could operate at even lower temperatures would seem worth pursuing. The rapidly growing world concern over future oil supplies has led to considerable research activity into the use of synthesis gas or synthesis gas-derived molecules, such as methanol, as future feedstock materials for

68

I. E. MAXWELL

both the oil and chemical industries. The unique shape-selective properties of zeolites will be utilized to constrain the chain-growth-controlled (SchulzFlory) kinetics of conventional Fischer-Tropsch catalysts, which leads to undesirable broad product distributions. Results of initial studies in this area using FT catalyst/ZSM-5 physical mixtures, modified ZSM-5 (243-249), and even larger pore zeolites such as RuNaY (250) are quite promising. Physical mixtures of zeolites with conventional catalysts will compete with single tailor-made zeolite catalysts. A variety of catalysts will probably emerge, to meet the different product needs of the oil and chemical industries. However, it now seems certain that the zeolites will play a central role in the catalysis of these new processes, which should give an enormous stimulus to zeolite research in general. ACKNOWLEDGMENTS I would like to express my sincere gratitude to the following colleagues who have read the initial manuscript and offered valuable suggestions: Dr. R. S. Downing, Dr. A. L. Farragher, Mr. A . G. T. G . Kortbeek, Dr. J. C. Platteeuw, Dr. M. F. M . Post and Dr. G . T. Pott. In addition I am very grateful to Miss E. Breekland who provided considerable assistance with the literature and patent searches. The support for the concept of this review from Professor W. M . H. Sachtier is also greatly appreciated. REFERENCES I . Rabo, J. A. (ed.), Am. Chem. Soc. Monogr. 171 (1976). 2. Jacobs, P. A,, “Carboniogenic Activity of Zeolites.” Elsevier, Amsterdam, 1977. 3. Venuto, P. B., Catal. Ory. Synth. Con(, 6th p. 67 (1977). 4. Uytterhoeven, J. B., Puog. Colloid. Polym. Sci. 65, 223 (1978). 5. Haynes, H. W., Caul. Reu. Sci. Eng. 17, 273 (1978). 6. Gallezot, P., C u d . Rev. Sci. Eny. 20, 121 (1979). 7. Rudham, R., and Stockwell, A., Spec. Period. Rep. Ctrtul. I , 87 (1977). 7u. Naccache, C., Proc. Int. Con$ Ze/J/ifP,S,5th. London p. 592 (1980). 8. Breck, D. W., “Zeolite Molecular Sieves.” Wiley, New York, 1974. 9. Seff, K.. Acc. C h ~ mRrs. . 9, 121 (1976). 10. Olson, D. H., Mikovsky, R. J., Shipman, G. F., and Dempsey, E., J . Cutol. 24, 161 (1 972). 11. Maxwell, I . E., de Boer, J. J., and Downing, R. S., J . Cutti/.61, 493 (1980). 12. Smith, J. V., Puoc. Int. Conf: Zeolites. 5th. Naples p. 194 (1980). 13. Weisz, P. B., Chenitech August, 498 (1973). 14. Chen, N . Y., and Weisz, P. B., C h m . En$. Prog. S.ynip. Srr. 63, 86 (1967). 15. Gorring, R. L., J . Catrrl. 31, 13 (1973). /&i. Post, M . F . M., results to be published. 16. Boreskov, G. K., Bobrov, N. N., Maksimov. N . G., Anufrienko, V. F., Ione, K. G., and Shestakova, N . A,, Dokl. Aknrl. Nuuk SSSR 201,887 (1971). 17. lone, K. G . , Bobrov, N. N., Boreskov, K. G.. and Vostrikova, L. A,. Dokl. Aktirl. Nuzrk SSSR 210, 388 (1973). 18. Boreskov, G. K., Proc. I n f . Congr. Cutal., 5th 2,981 (1973).

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NONACID CATALYSIS WITH ZEOLITES

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217. Nefedov, B. K., Dzhaparidze, R. V., and Eidus, Ya. T., Izv. Akad. Nauk SSSR, Ser. Khim. 1422 (1977). 218. Hortkjaer, J., and Jsrgensen, J. C., J . Mol. Catal. 4, 199 (1978). 219. Nefedov, B. K., Dzhaparidze, R. V., and Marnaev, 0. G., Izv. Akad. Nauk SSSR, Ser. Khim. 1657 (1978). 220. Nefedov, B . K., Sergeeva, N. S., and Krasnova, L. L., Izv. Akad. Nauk SSSR, Ser. Khim. 614 (1977). 221. Nefedov, B. K., and Kozhukhareva, V. N., Izv. Akad. Nauk SSSR, Ser. Khim. 2279 (1977). 222. Gates, B. C., Katzer, J. R., and Schuit, G . C. A., “Chemistry of Catalytic Processes,” p. 140. McGraw-Hill, New York, 1979. 223. Marko, L., Aspects Homogeneous Catal. 2, 3 (1974). 224. Scurrell, M. S . , Spec. Period. Rep. Catal. 2, 215 (1978). 225. Centola, P., Terzaghi, G., Del Rosso, R., and Pasquon, I., Chim. Ind. (Milan) 54, 775 (1972). 226. Mantovani, E., Palladino, N., and Zanobi, A,, J . M u / . Catal. 3, 285 (1977/1978). 227. US Patent 3,352,924 (1967). 228. US Patent 4,070,403 (1978). 229. German Offenlegungsschrift 2,804,307 (1978). 230. US Patent 3,940,447 (1976). 231. Vannice, M. A,, J . Catal. 37,462 (1975). 232. Vannice, M. A,, J . Cutal. 40, 129 (1975). 233. Bhatia, S . , Mathews, J. F., and Bakhshi, N. N., Acta Phys. Chem. 24, 83 (1978). 234. Elliott, D. J., and Lunsford, J. H., J . Catal. 57, 11 (1979). 235. Sinfelt, J. H . , Arc. Chem. Res. 10, 15 (1977). 236. Sinfelt, J. H . , J . Catal. 29, 308 (1973). 237. Gupta, N. M., Kamble, V. S., and Iyer, R. M., Radiat. Phys. Chem. 12, 143 (1978). 238. Lawson, J. D., and Rase, H. F., Ind. Eng. Chem. Prod. Res. Dev. 9, 317 (1970). 239. Meisel, S . L., McCullough, J. P., Lenchthaler, C. H., and Weisz, P. B., Chem. Tech. 2, 86 (1976). 240. Denny, P. J . , and Whan, D. A.. Spec. Period. Rep. Catal. 2, 46 (1978). 241. Abdulahad. I., and Relek, M., Erdol Kohle, Erdgas, Petrochem. Brennstofl-Chem. 25, 187 (1972). 242. Lapidus, A . L., Isakov, Ya. I., Guseva, I . V., Minachev, Kh. M., and Eidus, Ya. T., Izv. Akad. Nauk SSSR, Ser. Khim. 1441 (1974). 242a. US Patents 3,827,968; 3,756,942; 3,894,103-7; 3,907,915; 4,046,825; 3,960,978. 242b. Derouane, E. G., in “Catalysis by Zeolites” (B. lmelik et a/., eds.), p. 5. Elsevier, Amsterdam, 1980. 243. Chang, C. D., Lang, W. H., and Silvestri, A. J., J. Catal. 56, 268 (1979). 244. Caesar, P . D., Brennan, J. A., Ganvood, W. E., and Ciric, J., J . Cural. 56, 274 (1979). 245. British Patent 1,495,794 (1975). 246. British Patent 1,489,357 (1976). 247. US Patent 4,086,262 (1978). 248. Netherlands Patent Application, 7,711,350 (1977). 249. Netherlands Patent Application, 7,804,899 (1978). 249a. Jacobs, P. A,, in “Catalysis by Zeolites” (B. Imelik et a/., eds.), p. 293. Elsevier, Amsterdam, 1980. 250. Nijs, H . H . , Jacobs, P. A., and Uytterhoeven, J. B., J . Chem. Soc. Chem. Commun. in press (1980).

76

I. E. MAXWELL

2.51. Nijs. H. H., Jacobs, P. A,, and Uytterhoeven. J. B., J . C/wm. Sot,. C'iwm. Coninrun.

1095 (1979). 252. Verdonck, J. J., Jacobs, P. A , , Genet, M., and Poncelet, G., J . Chern. SOC.Furuday Trun.s. I. 76, 403 ( 1 980). 253. Nijs. H . H., Jacobs, P. A,, Verdonck, J . J . , and Uyttcrhoeven, J . B., Pro(,. h r . COH/. Zcolrrcs, 5th, Nuplrs p. 633 (1980). 254. Vanhove, D., Makambo, P., and Blanchard, M . , J . C ' h c w . SOC.Circm. Cornmutr. 605 (1979). 2 5 4 ~ .King, D. L., J . Caul. 51, 386 ( 1 978). 254h. Fraenkel, D., and Gates, B. C., J . A m . Ckrrn. Sor,. 102, 2478 (1980). 2 5 k . Ballivet-Tkatchenko, D.. Condurier. G., and Mozzanega, H., in "Catalysis by Zcolites" (B. lmelik r~ [ I / . % eds.), p. 309. Elsevier, Amsterdam, 1980. 255. Verdonck, J. J . . Jacobs, P. A,, and Uytterhoeven, J. B., J . Chc,m. Soc.. Chcm. Cbmmwr. 181 (1979). 256. Kiilbel, H.. and Hammer, H., Z . Elekfrockrrn. 64, 224 (1960). 257. Niwa, M., lizuka, T., and Lunsford, J. H., J . Chrrn. Soc. C ' h ~ r n Conrmun. . 684 (1979). 258. Chao, R. E., Ind. Eny. Chern. Prod. R1.s. DCT. 13, 94 (1974). 259. Jacobs, P. A . , de Wilde, W.. Schoonheydt, R. A., Uytterhoeven, J. B., and Beyer, H., J . ClicJm.So(,.Furuclrry Trans. 172, 1221 (1976). 260. Kasai, P. H., and Bishop, R. J., J . P / ~ . YChem. . 81, 1527 (1977). 261. Jacobs, P. A . , Uytterhoeven, J . B., and Beyer, H. K., J . C'hom. SOL..Clr~rn.Commun. 128 (1977). 262. Leutwyler, S., and Schumacher, E., Chirniu 13, 475 (1977). 263. Kurnicki, S. M., and Eyring, E. M . , J. A m . Chem. Soc. 100, 6790 (1978). 264. Ono, Y . , Suzuki, K., and Keii, T., J. Phys. Ciirm. 78, 218 (1974). 265. US Patent 3,963,830 (1976). 266. Rollmann, L. D., and Walsh, D. E., J . Curd. 56, 139 (1979). 267. Walsh, D. E., and Rollmann, L. D., J . Card. 56, 195 (1979). 26N. Venuto, P. B.. and Landis, P. S., A&. CuruI. 18, 259 (1968).

ADVANCES IN CAT4LYSIS. V O L U M E 31

Characterization and Reactivity of Mononuclear Oxygen Species on Oxide Surfaces M . CHE Labortrtoiw de Chinzir lies Solides. ER 133. C N R S UniurrsifP de Paris VI Paris. France AND

A . J . TENCH Chemisrry Diuision Atomic Energy Research Establishment Harwell. Oxford.ihire, United Kingdom

I . Introduction . . . . . . . . . . . . . . . I1. Characterization of 0 - by EPR and Optical Spectroscopy . . . . . . . . . . . . . . . A . ThegTensor . B. The Shape of the EPR Signal . . . . . . . . . C . The Hyperfine Tensor . . . . . . . . . . . D . The Superhyperfine Tensor . . . . . . . . . E . The Influence of the Surface Crystal Field on the g Tensor . F . Optical Properties of the 0 - Ion . . . . . . . . I11. The Formation and Stability of 0 - Ions . . . . . . . A . 0 - Species Formed by Adsorption . . . . . . . B . 0 - Species Formed by Ionizing Radiation . . . . . C . EPR Invisible 0- Species Formed by Heat Pretreatment or Adsorption . . . . . . . . . . . . . IV . Aggregate 0 - Species . . . . . . . . . . . . A . The Dimer Species (n = 2) . . . . . . . . . . B . The Trimer Species (n = 3) . . . . . . . . . . V. Reactivity of the 0 - Ion . . . . . . . . . . . . A . Simple Inorganic Molecules . . . . . . . . . B . Organic Molecules . . . . . . . . . . . . . . . . . . . . . . . C . Exchange Reactions V1. The Surface Oxide Ion 0;; . . . . . . . . . . . A . Electron Donor Properties . . . . . . . . . . B . Spectroscopic Studies . . . . . . . . . . . C. Chemical Reactivity . . . . . . . . . . .

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77 Copyrighi h 1982 by Academic Press. lnc . All rights of reproduction in any form reserved .

ISBN n-12-007831.7

M. CHE AND A. J . TENCH

78

VI1. Characterization and Reactivity of M=O . . Appendix. The EPR Parameters of 0 - Ions. References.

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Introduction

Oxides have been important catalysts for many years, but it is only fairly recently that attention has been directed toward the various forms of oxygen that may exist on the surface. Early solid-state studies of irradiated salts led to the discovery that paramagnetic oxygen species could be trapped in the salt matrix, and these species were identified and characterized. A little later, in the field of surface chemistry, kinetic studies on the conductivity of semiconductors during oxygen adsorption indicated that charged oxygen species were formed on the surface. It was at this time, in the mid-l960s, that the electron paramagnetic resonance (EPR) technique had evolved sufficiently to be applied to the identification of paramagnetic oxygen species on the surface, and there has been a rapid growth in this field. The EPR of adsorbed species was reviewed eight years ago by Lunsford ( I ) , and since then the literature has increased considerably and new areas have emerged. In particular, there have been a considerable number of papers on the characterization of the 0- ion and its reactivity. This interest has developed because of the role 0- might play in oxidation reactions at the surface. The most commonly used oxidizing agent is gaseous 0, , but it is generally believed that this oxygen must be converted to either 0 - or 0’on the surface and possibly even incorporated into the lattice before it appears in the oxidation products following the general scheme O2 ads.-

02-ads.-

0-

I I I

I L---)

Oxidation products

For this reason, the present review is restricted to mononuclear oxygen species, such as the 0- ion, but it has also been broadened in scope to include the nonparamagnetic surface oxide ion O2-. Recent spectroscopic measurements have shown that 0’- has unusual properties when it is in a situation of low coordination on the surface. Other mononuclear forms of ions are known in the gas phase, but oxygen such as charged O f and 0 2 + in view of the high ionization potential of the 0 atom (2) given below, these species are unlikely to be stabilized on the surface :

+ e-, o++ 0 2 ++ e-, O+O+

IP, = 13. 2 e V

IP, = 35.12 eV

MONONUCLEAR OXYGEN ON OXIDE SURFACES

79

Neither is the oxygen atom likely to be stabilized on the surface since even in the gas phase it undergoes a highly exothermic reaction, with electrons to yield 0 - [see Cotton and Wilkinson (41, O,,,+e+O&,,

A H = -142kJmol-'

The 0 - species on the surface will be further stabilized by coulombic interactions with the lattice cations. For the sake of completeness, the review also includes a brief section on oxygen ions which are closely bonded to transition metal ions (M=O) and which may play an important part in selective oxidation [Weiss el al. (41. The coverage of the mononuclear oxygen species is restricted to those papers where there is direct evidence on the nature of the oxygen species concerned. Many of the results discussed in this review have been obtained using EPR spectroscopy, and the application of this technique to surface studies has previously been reviewed by Lunsford (5). Measurement of the EPR parameters (the 9, hyperfine, and superhyperfine tensors) gives information (6) which at the simplest level indicates the presence of a paramagnetic species, but which can be sufficiently detailed to identify and characterize the adsorbed species. Moreover, measurements of the variation of the intensity of the signal with time can give kinetic data on reactions occurring at the surface. Although most of the EPR measurements have been carried out at low temperatures, there is evidence from the examples discussed later that some of the species are stable at much higher temperatures and may play an important role in catalytic reactions.

II.

Characterization of 0 -by EPR and Optical Spectroscopy

A. THEg TENSOR In the early work, a number of authors reported the presence of the 0 species on the surface, based only on the g tensor and supported by indirect evidence. Not all of this early work has been covered in this review. It was not until recently that the existence of 0 - was firmly established by measurement of the g and hyperfine (hf) tensors, together with its chemical reactivity. All the measured parameters are now in agreement either with theory or what was known from solid-state physics (0- trapped in the bulk of various salts) or gas-phase reactions. A theoretical analysis of the g tensor for 0 - trapped in alkali-halide single crystals has been given to second order in the case of axial (7) and orthorhombic symmetry (8). Recently, refined theoretical treatments have been

80

M . CHE AND A. J. TENCH

proposed for both types of symmetry by Nistor et a/. ( 9 , I U ) . In these two cases (8), it is predicted that two components of the g tensor should exceed ge (free-spin g value) and the third should be less. This can be seen by using a simplified treatment ( I I ) and considering the energy diagram given in Fig. lc for 0 - (2p:2p:2pi). To first order, the g components are calculated to be, for orthorhombic symmetry,

where A > 0 is the spin-orbit coupling constant of oxygen. If p, and p, are degenerate (Fig. lb), the g components become, for axial symmetry, 911 = s z z

91 = ~

= g,

(4)

x = x gyy

=

Ye

+ (2’IAE)

(5)

This simplified treatment, with 1> 0, leads to one value close to ge and two values exceeding 9,. In the case of cubic symmetry (Fig. la), the p,, py, and pz orbitals of the 0-ions would be degenerate, so one is dealing with a 2P state. For zero magnetic field, the *Pstate will be split by spin-orbit coupling into two states in which the orbital ( L )and spin (S) angular momenta couple (Fig. 2). For small magnetic fields (PH << 1, which would apply to normal EPR, since 111 135 cm-’), the g value in each level is given by the Lande g factor

-

gJ=1+

(a

J(J

+ 1) + S(S + 1) - L(L + 1) 2 J ( J + 1)

(bl

(C 1

FIG. I . The energy-level diagram for the 0- ion: (a) the free ion, (b) in an axial crystal field, and (c) in an orthorhombic crystal field.

MONONUCLEAR OXYGEN ON OXIDE SURFACES

81

J = L + S =2 ! ( 2 - told degenerate]

2P

( L - fold degenerate]

FIG.2. The effect of spin-orbit coupling on the energy levels of the 0-ion in a cubic crystal field.

which is 4/3 for the J = 312 level and 2/3 for the J = 1/2 level (12). It is probable that the combination of short spin-lattice relaxation times and heterogeneity of sites would make the corresponding 0- ions very difficult to see. An alternative possibility is that the 2P state in cubic symmetry could undergo a Jahn-Teller distortion to give a site of low symmetry. In reasonable agreement with theory, EPR signals have been reported for 0- adsorbed on MgO either in orthorhombic ( 1 3 , 1 4 ) or, more usually, axial (15) symmetry. A typical spectrum for 0- in axial symmetry adsorbed on MgO is given in Fig. 3a [Wong and Lunsford (16)].Further proof of the identity of this species comes from measurements of its reactivity (13-15) and from the use of 170-enrichedN,O (16) to give an hf interaction. In Fig. 3b, the hyperfine splitting (hfs) clearly indicates the presence of only one oxygen nucleus (2nZ + 1 = 6 lines, with n = 1, I = 5/2 for I7O, associated with the g,, and gL components). Since the first reports of 0- on MgO, these ions have been seen on many other surfaces. A full list of the systems and spectroscopic parameters is given in the Appendix. For the sake of comparison, the 0- species trapped in various matrices have been included, and within each section the order is given with increasing g L or g 3 values. Inspection of the table shows that for 0 - on surfaces, nearly all of the g , , components exceed 9,. Such values cannot be explained by the theory developed on the basis of a purely ionic model for 0 - (7, 8). As will be seen later, the detection of superhyperfine (shf) interactions indicates that 0- can be covalently bonded to the matrix, and the theoretical calculations need to be modified to take into account this covalency (13, which is only rarely considered in the case of 0- ions (18). From the g tensor, a spin density of about unity is found in the pz orbital pointing perpendicular to the MgO surface (16) (Fig. 4). It is then reasonable to expect a privileged shf interaction with one Mg2+ ion along the pz axis corresponding to g i l . Thus, if this

82

M. CHE AND A . J . TENCH

f gL = 2.0L2

cg,, =2.0013

< 91 10OG

Lg,,

FIG.3. (a) The EPR spectrum for 0- on MgO; (b) the EPR spectrum for ”0- on MgO,

FIG.4. The 0 - ion on the surface of MgO: 0 , Mg”; 0 , 0 2 ;-p, orbital symbol, 0 - .

MONONUCLEAR OXYGEN ON OXIDE SURFACES

83

model is valid for other systems, it is reasonable to anticipate that the spin-orbit coupling constant of the cation leading to shf structure will influence the gll value of 0-.This is seen for 0- stabilized on M g 2 + ,Mo", and W6+ cations with gI1values of 2.0013 (16), 2.006 (17), and 2.012 (19), respectively, where the spin-orbit coupling constant is known to increase fromabout 12cm-' (Mg2+)(20)to 1030cm-' ( M o 5 + ) ( 2 ~ ) a n d 5 0 0 0 c m - ' (W" ) (22). A reverse situation has been found for molybdenyl ions, whose g-tensor components have been found to depend markedly on the spinorbit coupling constant of the surrounding ligands (23). There is also evidence that the 0- ion may be observed in orthorhombic sites. Early work on the N,O/MgO system by Tench et ul. (14) showed that an 0- spectrum was formed which varied in complexity with the starting material and thermal activation temperature. The spectrum was assigned to two types of ions with the following g tensors: g1 = 2.0016, 9, = 2.0297, g3 = 2.0505, and g1 = 2.0016, g, = 2.0193, g, = 2.0472. The tetragonal symmetry was assumed to be removed by a nearby hydroxyl group which was known to be adjacent to the adsorption site. These 0- ions react in the expected way with a wide range of adsorbed molecules including H,, 0, , CO,, and CH,OH. The 0- ions in orthorhombic symmetry have also been observed in MgO single crystals with g1 = 2.00328, g2 = 2.03668, and g3 = 2.03960 ( 2 4 ~ ) . The only case where the g values are not in reasonable agreement with theory is the spectrum attributed to the 0 - ion adsorbed on chromium oxide supported on silica (24b).The spectrum is isotropic at 90 K with giso= 2.028 and axially symmetric at 4.2 K with gL = 2.030 and g I l= 2.024. The 0- ion is described as the tetraoxyanion [CrOJ or as a covalently bonded anion radical [(02-),Cr6+0-]. This 0- ion does not fall into the general pattern of 0- ions reported in the Appendix since gL > g l lwith g l llargely exceeding ge. For this system, the explanation given above to account for large gll values and based on the spin-orbit coupling constant is not valid, since for Cr5+a value of 380 cm-' is reported (21). Although the 0- ion shows the same type of reactivity as other 0 - ions (24h), there is no "0 data to strengthen the assignment to 0-.

B. THESHAPEOF

THE

EPR SIGNAL

The shape of the axially symmetric EPR signal of 0-adsorbed on surfaces is unusual, and the extraction of the g tensor from the powder spectrum is not straightforward. The case for axially symmetric EPR signals has been recently reviewed (25) and a typical well-resolved spectrum is shown in Fig. 5. There are three turning points (at zero slope) associated with three peaks of which the high- and low-field peaks (A and C) have been shown

84

M . CHE AND A. J . TENCH

A

B

C

FIG.5 . An EPR spectrum with an axially symmetric g tensor.

(26) to correspond to a good approximation in well-resolved spectra to yL and ( J , , . The third peak B (Fig. 5) is always associated with A, i.e., g I . Therefore, the assignments of the gL and g , , peaks in Fig. 3 are correct; however, their intensities and widths are unusual. For adsorbed 0- species, the perpendicular feature is broader and less intense than the g , , peak in contrast to the expected shape (Fig. 5). A similar type of spectrum is also observed for 0- trapped in hydroxide glasses (27). Comparison of X- and Q-band spectra for 0- in 10 M NaOH/H, "0 alkaline ice glass (28) shows that the width of the perpendicular component of 0- is about four times larger at Q band than at X band. The anomalous line shape is also found for 0- adsorbed on MoOJSiO, catalysts for which the heterogeneity of Mo sites has been shown by X- and Q-band studies (29). I t is seen from Eq. ( 5 ) that variations in AE values will give rise to a distribution of gi sites which will contribute to broadening the g L line but not the g,, line, which does not depend on AE [see Eq. (4)]. To support this argument it is interesting to ) . this material, the 0 - species is thought consider polycrystalline ZnO ( 3 0 ~ In

85

MONONUCLEAR OXYGEN ON OXIDE SURFACES

to be in the bulk rather than at the surface. In the bulk of a solid, the probability of having different sites is less than on surfaces or in solutions which have been frozen. The gi feature is not broadened and a normal spectrum for an axial g tensor is observed. There are, however, probably other mechanisms such as anisotropic relaxation processes which may also contribute to the anomalous shape of 0-,but the heterogeneity of sites is believed to be an important factor. C. THEHYPERFINE TENSOR

There are only two systems, MgO (16) and MoO,/SiO, ( 3 / ) ,for which the 1 7 0 hf tensors for surface 0- have been measured, and seven for the bulk 0- species. The results are summarized in Table I. Spin densities on 0- can be calculated, but the results depend largely on the method used. For MgO ( / 6 ) , a value of 1 was obtained from the g tensor; the hyperfine tensor gives 0.83-0.92, using a one-electron wave function, depending on the choice of sign which was taken for the hfs A , whereas a many-electron wave function gives a slightly higher value of 1.01. Using a different approach, Schlick and Kevan (28) found a total spin density of 0.69 on 0- in NaOH/H2170glass, indicating that a considerable amount of spin is delocalized onto the surrounding matrix. Along the same line, Symons (34), discussing the case of adsorbed species, suggested that 0- could better be described by Oi-, i.e., 0- in interaction with a nearby 0'- ion, since there is some ambiguity in the choice of the sign of the hf

-

TABLE I The g and "0 Hyperfine Tensors for 0-in Dlferent Environments Tensor

Surface : MoO,/SiO, MgO Bulk : NaOH ice Mg o SrCI, K1 KBr KCl RbCl

+

"

P P

2.002 2.0013

2.019 2.042

2.019 2.042

96 103.2

'? 19.5

? 19.5

31

FS

2.002 2.0032 2.0032 1.9733 I .9508 1.9475 1.9154

2.070 2.0389 2.0650 2.3023 2.2346 2.2217 2.2344

2.070 2.0389 2.0650 2.3023 2.4416 2.4524 2.5330

99 104.5 104.6 81.4 76.6 79 75

< 20 14.5 18.5 17.5 < 26.8 21.1 <32.1

<20 14.5 18.5 17.5 28.7 29.5 <44.6

28 32 33 7 8 8 8

sc SC SC SC SC

sc

P, powder; FS, frozen solution; SC, single crystal

I6

86

M. CHE AND A . J. TENCH

constant. However, an experiment performed with an enriched l 7 0 MgO surface showed that interaction of 0- with I 7 O 2 - ions was not measurable (35).In the Mo03/Si0, (31) system, the hf tensor is not completely resolved and the hfs is 7% less than for MgO (16). If the ratio of the s to p character of the orbital for 0- remains essentially constant, it is possible to conclude that the 7% decrease in hfs represents a similar decrease in ionic character for 0 - on the MoOJSiO, system. The hf tensor is thus important since it can confirm the delocalization of the spin density onto the matrix.

D. THESUPERHYPERFINE TENSOR The shf tensors have been measured for 0- adsorbed on the surface of a number of systems and also for 0- trapped in the bulk of single crystals (Table 11). We have seen that for 0- adsorbed on MgO, the electron was not delocalized on nearby 02-ions. However, for the systems in Table I1 the observation of a shf tensor is an indication that not all the spin density is localized on the 0- species. The number of shf lines observed indicates that in all cases o f O - adsorbed on surfaces, the interaction takes place only with a single cation nucleus. The shf tensor has been used to obtain the amount of spin delocalization on the cation leading to the shf structure. For 0- stabilized on Mo03/Si0, (17) (Fig. 6 ) , a value of 1.9% was calculated from the shf tensor for the spin density in the dz2orbital of molybdenum, which can be compared with the value of 7% for the same system which was derived by Ben Taarit and Lunsford (31)based on the comparison of the ( 1 7 0 ) A values for MoOj/SiO, (31) and MgO (16).

I I Ill I I FIG. 6 . The 0 - ion on MoO,/SiO, showing the shf interaction with the Mo ion. [Figure according to Kolosev rt al.

(In.]

TABLE I1 The g and Superhyperfine Tensors f o r 0- in Different Encironments Tensor

~~~

~~~~~

Surface".h: Group A MoO,/SIO, MoO,/AI,O, V,O,/SiO, WO,/SiO, Group B As,O,iSiO, A1,0, iSi0, P,O,:SiO, B,O,/SiO, Bulk : Be0 MgO SiO, 3Ca,(P04),, CaF, KBr KCI

2.012

2.020 2.024 2.026 2.026

2.020 2.024 2.026 2.026

2.0055 2.003 2.0076 2.002

2.0055 2.012 2.013 2.012

2.0055

2.0026 2.0032 2.003 2.0012 1.987 1.98 1

2.0155 2.0385 2.007 2.0516 2.226 2.258

2.0155 2.0385 2.029 2.0516 2.226 2.258

2.006 2.008 ?

9

2.019 2.038

7.6 3.4

8

?

8

14

14

14.5

?

?

Mo Al? V W

48.5 7 48.2

48.5 8 49.8

48.5

AS

13

15

?

1

?

Al P B

4.0 0.3 1.7 1.8

Be Mg Si 2F K K

49.8

d

4.0 6.9 3.5 3.9

4.0 0.3 1.7 1.8

0- from 0, or N,O adsorption; matrices of Group A lead to 0 - ,with the same parameters, by y irradiation in ~ U C U O(40).

' Group B, 0 - from y irradiation in oacuo. ' Parameters: A Parameters: A:

=

2.84 MHz, B = 0.86 MHz (for significance of A and B, see original reference). 4.855 MHz, A ; = 5.492 MHz, A: = 0.655 MHz (for significance ofA:, A ; , and A: see original reference).

=

17 36 37,38

39 40 39 40

19

41 42 43 44 45a 4.562

88

M. CHE AND A. J. TENCH

The observation of the shf interaction shows that a cation of the lattice is the stabilization site for adsorbed 0 - .

E. THEINFLUENCE OF THE SURFACE CRYSTAL FIELD ON

THE g

TENSOR

Inspection of Eqs. (1)-(5) for the g tensor, together with the evidence of delocalization of some electron spin onto a nearby cation, leads us to expect that the charge of the adsorption site may have an important effect on the g tensor. In Fig. 7, the g1 value of adsorbed 0- has been plotted against the charge of the stabilizing cation. We have also included the lowest value obtained so far for n = 1 (NaOH/H2l70system) (28),although it is not an adsorbed 0- species. The plot shows that there is some trend, with gL decreasing with increasing cation charge. A similar trend is found for 0; ions adsorbed on surfaces ( I ) . The crystal field stabilization energy has been calculated for an 0 - ion chemisorbed on third-row transition-metal oxides and values up to 70 kcal/ mol can be obtained for certain geometric arrangements (4%).

2.080 gL

11 y;

H I H , ‘‘0 (281

2.0’70

2.060

2.050 -

2.0LO

-

2.030

-

2.020

-

1

0

1

I

t

2

3

L

5

6

Cation charge

FIG.7. The variation of yL of the 0- ion with the charge ( n ) of the stabilizing cation. (Numbers in parentheses are reference numbers.)

-

MONONUCLEAR OXYGEN ON OXIDE SURFACES

89

F. OPTICAL PROPERTIES OF THE 0 - ION Parallel optical and EPR experiments have confirmed that an optical absorption band at 2.3 eV can be attributed to V- centers in single-crystal MgO (46, 47). It seems likely that this optical band is associated with a range of V- centers (Fig. 8), i.e., not only the 0 - ion adjacent to a cation vacancy, but also similar defects with nearby impurity ions. Many of the properties of these 0- centers are accounted for in terms of the simple crystal field splitting of the energy levels described in Fig. 1. However, attempts to attribute the band at 2.3 eV to a crystal field transition have failed because the use of typical values of the spin-orbit coupling term for the 0- ion to calculate the optical transition energy leads to large discrepancies from the experimental value. Recently, this has been resolved (48) using the polaron model of Schirmer et al. (49),where the transition is considered as due to the hole on one oxygen transferring to an adjacent oxygen sufficiently quickly to prevent the lattice from relaxing to a new equilibrium configuration. The transition energy is then determined by the self-trapping distortion produced by the hole, i.e., it is assumed that the polarization energy does follow the hole. There is very little optical data on the 0 - ion on the surface, mainly because many of the supports are colored and the ion is often unstable. However, MgO provides a substrate with good reflectance properties (see Section VI,B), and the reflectance spectrum of the surface, after N,O has been reacted with electrons trapped at surface centers, shows a broad band centered at 2.0 eV (50). This band is destroyed by the adsorption of oxygen to give a new band attributed to 0;.The formation of 0 - and its subsequent reaction with oxygen to give 0; were confirmed by parallel EPR experiments (50).The optical band at 2.0 eV is at energy lower than that observed in the bulk.

(a)

(b)

FIG.8. Models for (a) the V and (b) the Vo centers in the bulk oxide.

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M. CHE A N D A. J. TENCH

111.

The Formation and Stability of 0 - Ions

The 0- ions on surfaces can conveniently be divided into three main classes. First, those formed by adsorption; second, those formed by ionizing radiation; and third, those invisible to EPR and formed by heat treatment or adsorption.

FORMED BY ADSORPTION A. 0- SPECIES

The early work was complicated by the difficulty of making identifications based on the g tensor alone. A spectrum observed by Kazansky and coworkers (51) on TiO, was assigned at first to 0- on the basis of a full analysis of they tensor, but this signal was later shown to be due to 0; ions (52).They also reported the existence of a species on V,O,/SiO, which appears to have the correct characteristics for 0-.These authors showed that adsorption of 0, on reduced V,05/Si0, led to a complicated spectrum first thought to be due to 0; (53)with g1 = 2.025, g2 = 2.01 1, g3 = 2.005, and A , = 14.1 G, A , = 6.9 G, A , = 6 G (from the cation); but was then reassigned to a mixture of 0; and 0- ions (54). It was shown that 0- ions could be formed alone by contacting the reduced catalyst at 300°C with 0, and pumping off excess oxygen. At 77 K an eight-line spectrum with g = 2.025 and A = 14 G was observed, indicating that the 0- was adsorbed close to a vanadium ion. Attempts to confirm that the species was 0- using 7Oenriched 0, have not been successful owing to the presence of 1 7 0 and "V nuclei, both with I # 0 (55). This species was stable in 0, up to 300°C and was found to be very reactive toward H, and CH,, even at 77 K (54). The magnetic parameters, however, did not agree with theory, and it was later recognized (38)that the observed lines corresponded to perpendicular features and that it was possible from the spectrum to estimate that g1 > g , ,. Thus the parameters were corrected and slightly modified for gL to give gi = 2.026, A , = 14 G with gL > gli and A , 3 A , , (38).It should be noted that the 0 - species could also be obtained from N 2 0 adsorption ( 5 4 ) .This is important since N,O is known to be an efficient electron scavenger in radiolysis (56) to form N ,Oions. These ions, stable only in certain conditions (57, 58), normally dissociate into N, and 0-. Because of the absence of side reactions, the adsorption of N,O has been the principal method of forming 0- on MgO (13, 15), MoO,/SiO, (38), MoS,/SiO, (59,WO,/SiO, (19), and in ZnO (30a).The adsorption of oxygen has also been used to produce 0- ions on V,O,/SiO,, as seen earlier (38), and more recently it has been reported for MoO,/Al,O, (36) Pt/AI,O, (60), and Cr/SiO, (241).

MONONUCLEAR OXYGEN ON OXIDE SURFACES

91

Among all the systems, two [ZnO (30a) and MoO,/AI,O, (36)]are exceptional because they are thought to stabilize 0- species which are located close to, but not on, the surface. Attempts to form 170-from either N, 7O or I7O, in these systems lead just to the normal l6O- spectrum. In the case of ZnO (30u), it has been proposed that the electron affinity of an 0- ion in an oxide ion vacancy would be greater than that of an 0- ion adjacent to a zinc ion vacancy, and thus the 0- formed at an oxide vacancy in the surface would release a hole and become an 02-ion. In the case of MOO,/ A1,0, (36), the migration of the hole via a “hopping mechanism” to an oxide ion adjacent to an A13+ ion was proposed. We should note then, if this is correct, that a “hole affinity concept” can be envisaged. The hole affinity of a bulk 0,- ion close to a zinc ion vacancy must be larger than that of an 0,- ion on the surface. The reaction of 0- with 0, and CO at 77 K (30a) was used to suggest that the 0- species formed are near the surface, and in practice the difference in hole affinity is not large. The situation for the MoO,/A1,0, system is more complicated. The authors report that 0is observed in Q band after heat treatment of the sample, with a hfs of 3.4 G from Al3’. However, it is clear from their published data that 0; stabilized on A13+ is also still present after heating, and this species is known to show a hfs of 3.6 G from A13+ (61). Therefore, the assignment of the hfs to 0on A13+ seems ambiguous and the stabilization of 0- on the MOO, can also be envisaged. This would bring into line the g1 values and exchange properties of MOO, on Al,O, (36) and SiO, (31). Attempts to produce 0- from N,O adsorption have failed with this system, and it is not clear whether 0; is a necessary intermediate in the formation of 0- (36). In the case of V,0,/Si02, there is evidence that 0- can be formed from 0; ions (54). Apart from adsorption of 0, and N,O, it has been suggested that 0 atoms produced from a microwave discharge give 0- in small concentrations in zeolite 4A. The reported EPR spectrum overlaps with that of other species 0; and possibly 0; (62); the assignment is doubtful because it is based on a poorly resolved g tensor and the expected reactivity was not observed. Radicals in these systems are usually in the form of covalently bonded species, such as -AI-O’, -Si-0-O’, and -Si-0-0-O‘, some of which have been investigated recently and give different parameters (63,64). More recently, Katzer et al. (60) have reported on the formation of oxygen species on Pt/Al,O, catalysts. The observed spectrum was analyzed in termsofasignalwithg, = 2.141,g2 = 2.010, andg, = 1.963,andassigned to 0; adsorbed on Pt2+ or Pt4+, together with a one- or possibly a twocomponent signal (g = 2.005 or g1 = 2.008 and g,, = 2.005) assigned to 0associated with Pt. These assignments are based solely on the y-tensor components extracted from a complex spectrum. More experimental work should be carried out on this system, particularly with N,O- and ”0-

92

M . CHE AND A. J . TENCH

enriched gases (O,, N,O), in order to confirm the assignment. Some features seem to be in line with previous work on adsorbed 0; and 0- ( I ) , but it is difficult to understand the absence of any shf structure due to lg5Pt(I = 1/2) in the reported spectra, especially since shf structure was detected in the spectra of the paramagnetic platinum before oxygen adsorption, as well as in earlier work on the adsorption of NO on Pt/AI,O, catalysts (651). Shiotani e l ul. (65b)have reported the formation of several oxygen species on titanium-supported surfaces. Two of these are produced by UV irradiation of the titanium-porous Vycor glass system in the presence of N,O or 0, at 77 K and one is believed to be an 0 - ion. Their g tensors are not complete and there is no further data from l 7 0-labeled 0, or N,O or on their reactivity to support the assignment to 0- ions.

B. 0 - SPECIES FORMED BY IONIZING RADIATION Another method of producing 0 - used in solid-state physics is the use of ionizing irradiation. In this case there is no need for 0, or N,O adsorption. The holes and electrons which are formed by y irradiation normally recombine very rapidly unless there is a trap available. In oxides with lattice defects, such as vacancies or impurity ions, the electrons and holes can be captured by the defects to form paramagnetic centers persisting after y irradiation (66). The centers can be of various kinds, such as reducedimpurity cations, F centers, and trapped-hole centers, of which the last corresponds to adsorbed 0- species. This approach has been used (39, 40) where the electron traps were in the form of supported ions on silica gel leading to surface-hole trapped (0-)centers (Table 11). In the case of trivalent cations (A13 , B 3 +), which are formally negatively charged compared with Si4+,hole centers of “bridge” structure are formed, +

I

I

o2-

I

where Me = Al, B, As. In the case of penta- and hexavalent cations (P5’,V5+, Mo6+, W6+), which are positively charged compared with Si4+,terminal hole centers are formed according to

MONONUCLEAR OXYGEN ON OXIDE SURFACES

93

where Me = P, V, Mo, W with n = 6 or 5 and the box represents a formal cation vacancy in the coordination sphere. In the latter case, the vacancy is believed to have a stabilizing effect on the oxygen hole center at the surface (40). The EPR parameters of these bridge and terminal oxygen hole centers were found to be very close to those of the same bulk defects and 0- adsorbed species, respectively (Table 11). The reactivity of the hole centers at the surface was similar to that of 0- adsorbed species (39, 40). Another class of 0- species can be formed by UV irradiation and they are less easy to characterize in view of the short lifetime (T < sec). The transition metal ion-oxygen pair can absorb energy to give a charge transfer complex (67) which rapidly disappears when the irradiation is stopped. However, in the presence of gases the trapped hole on the oxygen is able to react chemically as an 0- ion. The reaction of these ions can be detected from the nature of the reaction products (40). Ultraviolet irradiation of silica gel with Co2+, Ni2+, V5+, and Mo6+ in the presence of various gases (H,, NH,, CH,OH) has been studied by EPR, and detection of the radicals allows the following reactions to be proposed : h! Mell+OZ- + {Me(n-11+0* (8) 3 )

0;

+ RH-+OH; + R'

(9)

where 0, represents the surface oxygen hole center and R = H' (68, 69), NH; (70), CH,OH' (70) is the radical which can be detected by EPR. Under these conditions, Eq. (8) is displaced to the right-hand side. Irradiation in the presence of an alkane and 0, leads to the production of peroxy radicals (40), 0;

+ RH + OL-+OH; + ROO;

(10)

Kubokawa et al. (71a) reported that UV irradiation of porous Vycor glass at 77 K in the presence of oxygen gave an EPR signal with g values of 2.0102 and 2.0030 which was assigned to 0-. Its formation was thought to follow reaction Eq. (8), but in that case the 0- species were stable at 77 K. On raising the temperature to around 170 K, the 0- signal disappeared, while a signal identified as 0, appeared according to the reaction 0-

+02-+o;

(1 1)

94

M. CHE AND A . J . TENCH

These results parallel those found on the V,O,/SiO, system where 0was formed by adsorption (72). The photoreactivity of transition-metal oxides deposited onto PVG or silica supports has been investigated by several authors (71h-71e) and is discussed in more detail in Section V1,C under reactivity of the lattice oxygen ions. Apart from their short lifetimes, another important reason for not observing the 0- species is dipolar interaction. As explained later (Section IV), if the distance between one 0- species and another paramagnetic center (0-or a metal ion) decreases, the linewidth may increase to the extent that the 0- species become “EPR invisible.” In this category of EPR invisible 0- species, there are not only the short-lived 0- produced according to Eq. (S), but also the long-lived 0- formed by heat pretreatment or adsorption, as discussed in the following section.

c.

EPR

INVISIBLE

0- SPECIES FORMED BY HEATPRETREATMENT OR ADSORPTION

It has been suggested by Boudart rt al. (73) that thermal activation of MgO without any irradiation can generate 0- ions. A weak EPR signal was reported under these conditions and assigned to a model consisting of a triangular array of three 0- ions (V, center) on a MgO( 11 1) surface plane. Although no EPR signal due to isolated 0- ions has been reported in thermally activated MgO, Martens rt ul. (74) have followed the formation of 0- ions indirectly by measuring the hydrogen and oxygen release during the thermal decomposition of Mg(OH), to form the oxide. While the gases evolved consisted mostly of water, small amounts of hydrogen and atomic oxygen were also detected. They suggested that 0- ions were formed in the bulk or on the surface by a process such as OH- . . . - H O + 2 0 -

+ H,

(12)

in concentrations 104-105 times larger than that of V, centers. This led Praliaud r t uf. (75)to try to detect such ions by a chemical reaction involving dissolution of MgO in a redox system (Fe2’/Fe3+). They found that an oxidizing species was present after thermal activation of MgO. The nature of the oxidizing species is uncertain, but the results are consistent with the presence of either (0- 0-)pairs located in the bulk and on the surface, or with existence of coordinatively unsaturated 0’- ions on the surface. Another chemical method has been proposed by Bielanski and Najbar (76) to determine the average electrical charge transferred from the adsorbent (NiO, COO, MnO) to the adsorbed oxygen atoms. In this method, it is assumed that the formation of charged chemisorbed oxygen is ac1 . .

MONONUCLEAR OXYGEN ON OXIDE SURFACES

95

companied by the promotion of an appropriate number of cations in the oxide crystal lattice to a higher oxidation state. The number of cations oxidized during chemisorption was measured analytically, and the charge of the adsorbed oxygen species could be deduced. It was found that for high surface-area nickel oxide, the majority of oxygen was adsorbed as 0- at room temperature. T o our knowledge, there has not been any convincing spectroscopic evidence for 0- on NiO. It is possible that dipolar interaction between paramagnetic centers (0- . . . 0-, 0 - . . . N i 2 + , or 0- . . . Ni3+) has prevented their observation or that the nonparamagnetic species 0:is formed on the surface.

IV. Aggregate 0 - Species Apart from the 0 - species whose characteristics have been summarized above, there are reports of 0- species which can combine to form molecular ions 0;- or can be found in arrays of n o - which occupy near-normal lattice positions. It is thus important, in view of the reactivity of 0 - , to consider the properties of these aggregate 0- species in more detail. A. THEDIMER SPECIES (n = 2)

If two 0- ions are independent and separated by a distance of R A, then the linewidth of the EPR spectrum governed by the dipolar interaction is of the order P/R3 where is the Bohr magneton. Thus, for a linewidth of 43 G, the distance R will be about 6 A. If the two 0- ions are coupled, the dimer species formed can be expected to exist either as a ground-state singlet or triplet, depending on its nature. In the alkaline-earth oxides, bulk Vo centers have been reported; these centers consist of a cation vacancy with two hole centers, one trapped either side of the vacancy. Their spectroscopic behavior is characterized as an axial S = 1 center and can be described by the spin Hamiltonian, with the principal axis oriented along the (100) direction. The EPR parameters are summarized in Table 111, together with those of the corresponding monomer center V - . The powder EPR spectrum of V o centers appears as two doublets centered at about the gi and g , , values, respectively. The linear configuration of the Vo center is 0- cation vacancy 0- (Fig. S), and assuming that the fine-structure splitting D is due entirely to a dipole-

96

M. CHE AND A. J . TENCH TABLE 111 EPR Parameters for V - and V o Centers in Alkaline-Eurrh Oxide Single Crystals

~~~~

~

~

Lattice

Defect

g I1

Sl

MgO

VVO

2.0033 2.0033

2.0386 2.0395

VVO

2.0021 2.0021

2.0697 2.0733

V-

2.001 3 2.0012

2.0705 2.0748

CaO

SrO

VO

D(10-4cm-')

212.61 j227.33 G } 114.08 { 122.05 G } 127.05 { 135.99 G}

Ref. 77 77 78 78 79 79

dipole interaction between the two positive holes, the formula (80)

can be used to calculate the distance R between the holes, where p, is the Bohr magneton and h and c have their usual meaning (79). The results are given in Table IV, together with the spacing a of the lattices. These distortions are larger than expected and indicate that the simple model involving only dipoleedipole interactions may not be adequate. The results for the bulk centers in the alkaline-earth oxides show that, in principle, EPR spectra could be obtained for a pair of 0- ions separated by R on the surface. However, there are no examples of well-resolved finestructure spectra for such adsorbed species. The main reasons for this are expected to be as follows: From the above formula, one would anticipate the spacings between the lines in the parallel and parpendicular directions to increase with decreasing distance R. It is therefore likely that with the heterogeneity of sites that has been discussed previously for isolated 0species, the EPR spectrum would be difficult to observe. Electrostatic repulsion will favor the collinear orientation of the two trapped holes on opposite sides of the positive ion vacancy, as observed from the relaxation TABLE IV Reluxation o j V o Centers in Alkaline-Eurth Oxide Singlr Crysruls Lattice

Mgo CaO SrO

R

(A)

4.99 6.16 5.95

0

(A)

4.21 4.80 5.16

Relaxation

-- 28 -

18 15

(2))

MONONUCLEAR OXYGEN ON OXIDE SURFACES

97

of the lattices (Table IV). From Fig. 9, it is obvious that a totally different situation is expected if a molecular ion 0;- is formed, since the two added electrons to 0, would fill up the TI*antibonding orbitals forming a diamagnetic species. The transition from the paramagnetic pair of 0- to the diamagnetic molecular ion 0;- will occur at a critical distance between the oxygen nuclei in the range from about 5 A (Table IV) to about 1S O A, the distance observed generally for peroxide salts (81). Such a transition will in turn depend on the sites available at the surface (plane, edges, dislocations, etc.) and their relative distances. These two distinct situations have not been clearly distinguished, and the evidence for 0;- on surfaces comes mainly from a combination of volumetric and/or EPR measurements (82-86). However, it is ambiguous to assign to 0:- the difference between the number of oxygen molecules adsorbed and the number of spins (generally identified as adsorbed 0 ; ), since there are several forms of nonparamagnetic adsorbed oxygen which may be formed (0'- is an example). In the peroxide salts, the 0;- groups give voo bands in the range 1093-1054 cm-' for the alkali-metal group (from Li,Oz to Rb,02) or 1088-1061 cm-' for the alkaline-earth group (from MgO, to BaO,) (81). There are few reports on the 0;- species in other matrices; in KC1 single crystals doped with 0, an absorption peak at 260 nm accompanied by a strong fluorescence at 395 nm (87) was attributed 10 Oi-. This species can also be synthetized by

02

022 -

FIG.9. The energy-level diagram showing the occupied molecular orbitals for O 2 and O i - .

98

M. CHE AND A. J . TENCH

matrix reactions of alkali-metal atoms with oxygen molecules at very low temperatures in argon (88). The presence of two oxygen nuclei in the I R spectra (88) is clearly demonstrated using isotopically labeled oxygen molecules.

B. THETRIMER SPECIES (iz = 3) This species was suggested by Symons (34) and can be looked at i n two distinct ways: as either a molecular ion 0;- or an array of three distinct 0- ions. There are only few reports on O:-. This species has been thought to be produced by UV irradiation of TiO, in the presence of oxygen and is characterized by gL = 2.001 and y,, = 2.008 (89); both 0, and 0; are produced at the same time. Use of "0-enriched oxygen did not lead to any additional hf structure apart from that typical of "0; adsorbed on TiO,, which had been reported earlier (52). There was however a decrease in the intensity of the signal assigned to O:-, as expected from the use of "0-enriched oxygen. The assignment to 0:- could not be made directly from the spectroscopic evidence but was based on the identification of a product (0-0-CO-) which was thought to be formed according to the reaction 0:-

+ co+02- + o-o-co-

(15)

A similar EPR spectrum had also been observed by UV irradiation of TiOz (gi = 2.001, g,l = 2.008) (90)or TiO, supported on SiO, (yI = 2.002, yil = 2.01 1) (38).Initially assigned to 0- species (38, 90), the EPR spectrum was later recognized to be due to 0; (91). This peculiar O;, with unusual spectroscopic parameters, was also reported on SiO, (92) and V,O,/SiO, (72). In the latter system, analysis of both X- and Q-band spectra, together with the use of "0-enriched oxygen, led to the following parameters: g1 = 2.007, y, = 2.002, g 3 = 1.998 with A , = A , 2 0 G and A , = 78 G. The hf structure is consistent with two equivalent oxygen nuclei (72). Comparison ofthis data with that ofthe postulated 0:- shows that it is no&in fact possible to separate the hf structure due to * 70-enriched 0:- from that of 0; with parameters A 3 = 77 G, g 3 = 2.003 (89). It is therefore probable that 0;- is not formed, and the spectra should be reassigned to an anomalous 0; species. Calculations for the 0:- species have been performed using the CNDO/SP method (93): 0:- is a 21-electron radical and the twenty-first electron enters a o bonding orbital whose energy tends to favor a linear geometry (9446) (Fig. 10). However, several other geometries have been assumed and calculations show that in any case the values (whether corresponding to the central or the terminal oxygens) for the isotropic or

99

MONONUCLEAR OXYGEN ON OXIDE SURFACES

n-

1

L

FIG.10. Correlation diagram for AB, molecules. [Figure according to Atkins and Symons (94a).1

anisotropic parts of the hyperfine tensor lie below 400 and 40 G, respectively (93).Thus, the explanation that the hf structure is not observed because of a very large isotropic interaction and the polycrystallinity of the sample (89) is not valid. Other reports of the 0;- species are also on TiO, in the form of anatase either hydrated or dehydrated and irradiated by X rays (Y5) or UV rays (96). Unfortunately, the assignment was based on the earlier one made for TiO, (89)and is probably incorrect. It has been reported that the so-called 0;- is the only species obtained by irradiation of pure N,O (96) adsorbed onto TiO,; this is in contrast to previous work (89), and in this case our reassignment to 0; could be directly tested using N,”O. The formation of an 0;- species has been suggested (34, 35, SY) via the following reaction, 0&4

+ 0;

-+

0:-

(16 4

If 0;- is formed by interaction with oxide ions of the lattice, then the spectra of 0; on an oxide lattice where the surface oxide ions are enriched with ”0 should be different from the normal. This experiment has been carried

100

M. CHE AND A. J . TENCH

out with MgO enriched with "0 and no difference was seen (35). There is therefore no concrete evidence for the existence of the 0;- ion. Surface 0- species have also been reported to form triangular arrays of three 0 - stabilized on a (1 11) plane of MgO and are characterized by an EPR spectrum with g = 2.0030 (73). However, the assignment is questionable since these species did not react with H, even under rather drastic conditions (1 atm, 298 K, 24 hr), in contrast to what is observed for the isolated V center on iron-doped MgO (97). In addition, they react reversibly with 0, and H,O to give other V-type centers, in contrast again with the isolated V - center which reacts with 0, to give 0; on MgO (50). This is also in sharp contrast with the behavior of the adsorbed 0 - species that have been described earlier. Furthermore, no calculation has been made to show that the observed spectrum is consistent with the proposed model, and in view of the earlier evidence from Vo centers, it is not clear if exchange narrowing will be operative at the distance of 2.97 A separating the three 0- ions in the array. The only reported analog of this species is the trigonal center composed of a similar array of three F centers. As expected from the geometry of this aggregate, an anisotropic EPR spectrum was observed with gl = 2.0039 and g , , = 2.0036 in MgO (98a). Likewise, one would anticipate an anisotropic spectrum for an array of three 0- ions, which is not observed. To our knowledge, the characteristics of such centers have not been reproduced and one of the reasons may be their low concentration (maximum of about 10'' spins g-'). The formation of 0:- in larger concentrations has been reinvestigated by Indovina and Cordischi (98b) in the case of CaO. After activation in VLICUO in the 1000-1273 K temperature range, a 2g-value EPR signal is produced (gl = 2.002, g, = 2.010) on adsorption of oxygen. After adsorption of natural oxygen on CaO, previously exchanged with I7O,, the above signal showed some changes, suggesting that the paramagnetic species contains oxygen atoms. The signal is tentatively assigned to a 0:-surface species thought to be formed from low coordinated surface-oxide ions 0;;. The 0:- center consists of a triangular array of 0- species thought to be placed either on (1 11) surface microplanes or simply on the corners of oxide particles. The 0;- species is believed to be formed according to the following reactions :

+ 02,,,-+o,+ 0,

0;;

30,

-+

0:

~

( 16b)

(16c)

The same center is formed when chlorine is adsorbed on CaO. Although the EPR signal of 0:- is now in better agreement with expectation, it is not clear why an additional g component at 2.028 is detected on I7O en-

MONONUCLEAR OXYGEN ON OXIDE SURFACES

101

riched CaO rather than the 29 EPR signal on natural CaO. Further work needs to be done in order to clarify the origin of these signals. From the data available so far, there is as yet little evidence that oligomers with n > 2 can be formed on surfaces. This is in line with studies of the oxygen species trapped in the bulk of various matrices where only Vo centers (pairs of V - centers) have been observed. It is perhaps worth recalling that one of the safest ways of identifying surface oxygen species has been through careful comparison of identical species trapped in the bulk of single crystals and use of 70-enriched gases.

V.

Reactivity of the 0 - Ion

In this section the reactivity of the adsorbed 0 - ion is discussed for some elementary reactions on oxide surfaces for which there is direct spectroscopic evidence of the participation of the ion. Evidence from the gas phase shows that free 0 - ion reacts with H , , the lower hydrocarbons, and also with oxygen ; both associative and hydrogen abstraction reactions are known to occur (99, 100).

A. SIMPLE INORGANICMOLECULES The thermal stability of 0 - on MgO depends on the presence of N,O and decreases by an order of magnitude if excess N,O is removed. This indicates that some complex such as N,O- or N,O; is probably formed (101). Exchange studies with I7O-/Nt6O support this contention and Naccache and Che (102) have reported an EPR signal showing hf interaction with two nitrogen atoms (aw= 5.7 and 41.5 G) which was attributed to NzO+,but these parameters would also be consistent with N,O;. The reaction of H, with 0 - on MgO leads to immediate destruction of the 0- ion, even at 77 K (13, 14). In this case, the trapped-electron signal reappears and more 0 - ions can be formed by subsequent adsorption of N,O. This was confirmed by other workers (15), and it appears that the cycle can be repeated a number of times although some loss of intensity occurs. A reaction of the following type is probably involved, 0;

+ 0:- + H2-+20H, + e -

(17)

where the electron then becomes trapped at a vacancy (subscript “s” indicates surface ions). Adsorption of 0, destroys the 0 - ion giving a new EPR signal, with an orthorhombic g tensor, which has been attributed to 0; on the surface (13, 14, 102, 103). In the early work of Lunsford and

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M. CHE AND A. J . TENCH

co-workers (1.9, it was believed that adsorption of H,, CO, CO,, or additional N,O on 0- ions could promote the reaction 30- ‘ 0 ;

+ 2e-

(18)

as shown by the appearance of an EPR spectrum assigned to 0 ; .However, the signal of the trapped electrons did not reappear. These results, due in fact to oxygen impurity contained in the gases used, show that oxygen is far more reactive than H,, CO, CO,, or N,O, and that the preceding reaction does not occur. When freed from oxygen impurity, gases such as CO and CO, also appear to undergo a simple associative reaction to form the corresponding radical ions CO; (102, 103) and CO; (14). Surprisingly, SO, does not react to give SO; (104). Similar ions have been reported by Warman (105) for the gas-phase reaction of 0- with CO, to form C O ; . The reactivity of 0- formed on silica-supported oxides of Ti, V, Mo, and W has been investigated by Kazansky and co-workers (38). The H, molecule reacts with the 0- ion even at 170 K, and the reactivity is found to vary in the order V ) M o ) W. Only about half the sites react at this temperature for WO,/SiO, , indicating an inhomogeneity of sites on the surface. The 0- ion on V,O,/SiO, also reacts with 0, at 77 K to form 0; with g , , = 2.015 and g1 = 2.003 (106).The 0- signal can be recovered by warming and evacuation. A kinetic study (107) showed that the reaction with H,, D,, and CO was first order in the 0- concentration and the gas pressure; activation energies of 1 and 1.3 kcal mol-’ for H, and D,, respectively, were reported for 0on V20,/Si02. Reaction of 0- with CO on MoO,/SiO, gave CO,, which was stable even at 300 K, whereas on V,O,/SiO, and WO,/SiO,, the COY ion was not seen but CO, could be detected in the gas phase. The measured rate constants on the surface for the 0- reactions were about seven to eight orders of magnitude smaller than obtained for the free 0 - in the gas phase (1UO), and this is attributed to the covalency of the bond between the 0radicals and the surface. The low-temperature oxidation of CO has been studied in detail for MoO,/SiO, catalysts by Kazusaka and Lunsford (I08a). It was found that the 0- ion is not an intermediate and that the reaction in fact appears to involve molecular complexes of N,O and CO on clusters of Mo” and M o 3 + ions. The successive reactions of 0 - ,with CO to form CO; and then oxygen to form CO;, have also been studied on various systems: MgO (108b), Cr/SiO, (93), and MoOJSiO, ( 1 0 8 ~ )The . COT ions observed by EPR are an interesting example of species formed by secondary reactions (6).

MONONUCLEAR OXYGEN O N OXIDE SURFACES

B.

ORGANIC

103

MOLECULES

The reactivity of 0 with simple alkanes and alkenes has been studied with both MgO and supported metal oxides as substrates. 1. Alkanes Alkanes react with the 0- ion on the MgO surface to destroy the EPR signal, but no new EPR signals are observed. Aika and Lunsford (109)have carried out a stoichiometric study which involved the EPR detection of the 0- ion reaction with the alkanes (from methane to n-butane) and gas chromatographic analysis of the reaction products. The surface 0 - ions reacted with the alkanes at a rate governed by the transport of the hydrocarbon with a stoichiometry of 1 : 1. At low desorption temperatures the main product was the corresponding alkene, the amount of alkene desorbed increasing with temperature, and at high temperatures CO, and CH, were formed. No alkyl radicals were observed but the overall reaction involves hydrogen abstraction similar to that reported for the gas phase (100, 110a). It was suggested that the alkyl radicals formed react with surface 0 2 -ions to give surface alkoxides or undergo a second hydrogen abstraction by (02-), to give a small amount of ethylene at low temperature-the unpaired electrons being trapped in nearby anion vacancies, e.g., C,H; C,H;

+ (OZ-),-+(OC,H;), + e+ (Oz-), C,H, + e- + (OH-), -+

(19)

(20)

and later undergoing radical coupling reactions. Decomposition of the alkoxide leads to the formation of ethylene at higher temperatures. An interesting comparison can be made with the work of Yun et al. (liOb), who report the oxidation of alkanes and alkenes by N,O over UV-irradiated MgO. The active oxygen species were also thought to be 0ions, but in the reaction with alkanes most of the alkene formation occurred at room temperature without raising the temperature to desorb the products. An alternative mechanism involving dehydrogenation of alkyl radicals by 0- ions was proposed, since the 0- ions formed by UV irradiation would probably be less stabilized and in higher concentration than those found in Lunsford’s work. The reaction products with alkenes were similar to those obtained by Aika and Lunsford ( I 14), which are described in the next section, but with a larger fraction of CH,. The 0- ions formed on the supported metal catalysts V,O,/SiO,, MoO,/SiO,, and WO,/SiO, also react with alkanes such as methane. A

104

M. CHE AND A. J. TENCH

kinetic analysis was carried out for CH, (107), and the rate constant was found to be four times higher than for H, on the MoO,/SiO, system. The rate constants for similar reactions carried out in the gas phase are quite different and indicate large effects due to the surface. The 0 - ion can also be formed by irradiation of V5+ or P5+ on SiO, or by irradiation of TiO, (111a). Exposure of the irradiated samples to CH, and C,H, led to destruction of the 0- EPR signal. A product analysis, similar to that described previously, led to the suggestion that a stable alkoxide species was formed on the surface via an hydrogen abstraction process. When 0- ions react with CH,, the main desorbed products were CH, and CO. Reaction of 0 with ethane did not lead to the formation of ethylene in this system. The suggested mechanisms were complicated and are not discussed further. Work by Yang and Lunsford ( I l l b ) on the reaction between N,O or 0, and ethane over MoOJSiO, has used a single pass reactor to obtain kinetic data on the formation of C2H4 and CO,. N,O was shown to be much more effective in forming C,H, , particularly at lower temperatures. It was proposed that the N,O formed 0-ions on the surface which reacted with ethane to form C,H, on sites composed of pairs of Mo ions. It was suggested that 0-ions might also be formed thermally on the catalyst surface and these could then be responsible for hydrogen abstraction reactions (see Section VII).

2. Reaction with Alkenes The reaction of ethylene with the 0- ion has been of considerable interest because of its relevance to the formation of ethylene oxide. Adsorption of ethylene on the 0-/MgO system leads to a completely new EPR signal (103) characterized by an isotropic g value of 2.0056 and an hf structure, suggesting isotropic interaction with two pairs of protons with aH of 4 and 57 G, respectively. The existence of a large hf interaction with two protons was confirmed using C,D, (102) and the spectrum assigned to the species CH,CH,O- . Because of difficulty in understanding the possible structures, isotopic labeling of the ethylene with 13C was used to probe the spin density on the carbon atoms (112). The assignment was reconsidered (113) and finally the 13C hf interaction from one atom was found to be strongly anisotropic with A,, = 77 G and A , = 15 G ; the interaction with the other was too small to be detected. The proton interactions were then reanalyzed in terms of a large interaction from a pair of protons and the small interaction arising from a nearby hydroxyl group. This is consistent with a species such as CH,=C- . . . HO-, formed via successive abstraction reactions in which

MONONUCLEAR OXYGEN ON OXIDE SURFACES

105

the H,C=cH radical is not detected, H,C=CH, (H,C=CH),

+ 0;

-+

(H,C=CH),

+ 0:- +.(H,C=C-

+ OH;

(21)

HO-),

(22)

It is interesting to note that reaction of trapped electrons on MgO with radicals. Higher ethylene oxide leads to ring opening to give H,c-CH,Oalkenes destroy the 0- signal, but no new signals are formed with propylene and only very weak signals are observed for 1-butene and 1 :3-butadiene, which probably form radicals analogous to that formed with ethylene (114). A detailed study of the oxidation of alkenes by 0- on MgO at 300 K indicated a stoichiometry of one alkene reacted for each 0- ion (114). With all three alkenes, the initial reaction appears to be the abstraction of a hydrogen atom by the 0- ion in line with the gas-phase data (100). The reaction of ethylene and propylene with 0- gave no gaseous products at 25"C, but heating the sample above 450°C gave mainly methane. Reaction of 1-butene with 0- gives butadiene as the main product on thermal desorption, and the formation of alkoxide ions was proposed as the intermediate step. The reaction of ethylene is assumed to go through the intermediate H,C=C'- . . . HO- which reacts further with surface oxide ions to form carboxylate ions in Eq. (23),

H,c=C-

*

*

HO-

+ 02--

p

H~C--C~0

+ 3e-

(23)

which can then decompose to give CH, and COi-. A similar mechanism is proposed for the reaction with propylene. This work supports the idea that the 0- ion could be the active form of oxygen in the catalytic oxidative dehydrogenation of butene to butadiene (115). Ethylene also reacts with the 0-ion formed on various supported oxides to give rather different radical species. For MoO,/SiO, and W0,/Si02, the initial reaction of the 0- ion is additive rather than the abstraction reactions studied on MgO. On Mo0,/Si02 (116), a new EPR spectrum is formed at 110 K with hyperfine interactions of 21.5 and 31 G from two pairs of protons. This is reminiscent of the radical formed by the reaction of ethylene oxide with trapped electrons on MgO and has been attributed to cH,CH,O-. After warming to 150 K, this spectrum is replaced by a different one consisting of a doublet of triplets with splitting constants of 64 and 21.5 G, respectively, which is attributed to the CH,cH radical thought to be an intermediate in the reaction on MgO. The formation of CH,CH,O- on the MoO,/SiO, system has been confirmed using deuterium

106

M. CHE AND A. J. TENCH

and I3C-labeled ethylene (117); carbon hf interactions of 39 and 13 G are observed from the a and fl carbon atoms, respectively. Surprisingly, in this work the radical was stable for several days at 373 K, but there was no evidence for the CH,cH radical. This illustrates the variation to be expected between different preparations of the same type or samples. The 0- ion was also found to react with propene and 1-butene (117) to give methyl radicals and CH,CH,O-, indicating that a cracking process is also possible. It is interesting to note that on the V,O,/SiO, surface, no radicals were observed after reaction with 0-. It is not clear why these differences in stability or in the nature of the intermediates formed on the surface exist. On the WO,/SiO, system, the reaction of 0 - with ethylene at 77 K leads to a different signal with five hfcomponents and a splitting constant of 25.5 G (f9).Warming to 90 K leads to the formation of the radical previously identified as CH,CH. The initial species was identified as an 0- adduct to form a n-type complex such as a cyclic (CH,CH,O-) species and this assignment was supported by theoretical calculations. No studies using I3C were carried out to confirm the equivalence of the carbon atoms. In view of the fact that the hyperfine constants from the bent species CH,CH,O- can change considerably with temperature and probably also with geometry on the surface, this latter species should also be considered as a possibility.

C . EXCHANGE REACTIONS Since the 0- ion readily forms 0; on adsorption of oxygen in a process which is often easily reversible, it has been suggested (106, 118) that this process could provide a suitable pathway for exchange reactions to take place. The oxygen species on the V,05/Si0, system have been well characterized, and after reduction in CO at 500°C homonuclear exchange of oxygen is observed to occur rapidly, whereas at the same time (at low 0, pressure), EPR shows that both 0- and 0; are present (fU6, f18).Addition of H, leads to a sharp drop in the rate of exchange and the EPR signal for 0disappears, whereas the intensity of the signal from 0; remains unchanged, indicating a mechanism such as the following, (0-),

+ (O& * ~0-0-0), * (0.'), + ((Idg

(24)

The apparent activation energy for this reaction has been measured to be in the range 0.9 - 1 .O kcal mol- '. This reaction was confirmed by EPR studies using 0, labeled with I7O ( f l y ) when 'Oh: on the surface was found to be replaced by I7O;. The rate constant measured for the exchange was 0.9 x cm3 mol-' at 300 K and only about 70';., of the 0- sites were found to react. The relatively slow rate of reaction compared to the reaction

MONONUCLEAR OXYGEN ON OXIDE SURFACES

107

of 0 - with molecules such as H, and CO presumably reflects the relatively long lifetime of the ( O ; ) 5ions. A similar kind of exchange process is thought to occur after activation by UV of the TiOJSiO, system (90, Y l , 120). After UV or y irradiation of MgO, the activity toward H,-D, exchange and 0-p H, conversion is increased (Y6, 121, 122a) by a factor of 2 or 3. It has been suggested that V, centers (0- ions next to a cation vacancy on the surface) are responsible for the enhancement of these reactions by a mechanism (122a) such as (25)

A possible alternative can be proposed in view of the high reactivity of the 0 - ion toward H,. In this situation the majority of the V, centers would be expected to form an OH - ion next to a cation vacancy and the hydroxyl ion may then act as the active center.

VI.

The Surface Oxide Ion 0;;

In a discussion of the types of oxygen species at surfaces the 0’- ion is often neglected, and yet there are many such ions present on the surface of oxides; there is also evidence that they can be formed as the final stage of some adsorption processes (122b, 123, 124). It is these ions at the surface that play a major role in determining the surface properties of the material. The reason for this omission is simple; the surface oxide ion is difficult to study because it normally cannot be easily distinguished from the other oxide ions of the lattice. Measurements of the activity, either for exchange (125, 126) or for catalytic reactions (127), give some indication of the heterogeneity of the surface-oxide ion, e.g., changes in coordination, but give little evidence as to the characteristics of the species involved and in some cases exchange with the bulk oxide ions occurs rapidly (128, 129). A few moments’ thought about the nature of the surface of an oxide leads to the conclusion that the surface oxide ion should have quite different properties than the bulk lattice ions. For example, consider a simple cubic oxide such as MO with a sodium chloride structure where each ion is sixfold coordinated; if this is cleaved along a (100) plane, then the coordination of the ions in this plane is reduced from six- to fivefold. This new surface will not be ideal, and ions of still lower coordination will also be present where higher index planes are exposed at the surface. However, for MgO prepared by thermal decomposition of the hydroxide or carbonate, evidence from electron microscopy (130) indicates that these have high index planes that

108

M. CHE AND A . J . TENCH

are better regarded as “mean” planes derived from the addition or removal of repetitive units, i.e., unit cells to (100) faces. It seems likely that “actual” (as distinct from “mean”) (1 1 1) planes are not present on a macroscale, although microarrays may exist. Theoretical work by Tasker (131) confirms that this type of surface, which is charged and possesses a dipole moment perpendicular to the surface, can only be stabilized by substantial reconstruction. Overall, for the simple cubic oxides, it is possible that the lower coordination ions are most likely to be associated with imperfections in the low-index surface planes. It is useful at this point to define more precisely what is meant by ions of low coordination. The normal coordination of the ion in the bulk lattice is defined by the number of nearest neighbor ions of the opposite sign, and for the octahedrally coordinated alkaline-earth oxides this will be six ; ions of lower coordination than this will be designated by the subscript LC to denote low coordination. For example, the surface oxide ions on an MgO( 100) plane can be described as 02; with coordination 5, or O:;, and more generally, L can take the values 5,4, 3, etc., for different coordinations as shown in Fig. 1 1. The use of the subscript LC is recommended in preference to CUS (coordinatively unsaturated) because of its simplicity. During the last few years there has been increasing evidence that ions in positions of low coordination are linked with unusual electron donor

I

. 0 . 0 . 0 .

io

. .

o/

FIG. 1 I . Representation of a surface (100) plane of MgO showing surface imperfections such as steps, kinks, corners, etc., which provide sites for ions of low coordination.

MONONUCLEAR OXYGEN O N OXIDE SURFACES

109

properties and this has been supported by spectroscopic evidence. In the following sections, we have mainly concentrated on the work with the alkaline-earth oxides because they are largely ionic and are good model systems which have been studied in some detail. They are also free from the complicating effects of variable-valence transition metal ions. In these systems it is very unlikely that the effects described arise from impurities, because a wide range of material source and added impurities have been investigated by a number of different workers.

A. ELECTRON DONOR PROPERTIES Nitrobenzene (NB) and related compounds were used by Tench and Nelson (132) to study the electron donor properties of an ionic oxide such as MgO. Adsorption of NB in vucuo onto a MgO powder prepared by outgassing at 900°C in V ~ C U Ogave a strong EPR signal due to NB- ions with gL = 2.006 showing hfs from the proton (4 G) and an anisotropic nitrogen splitting (2.5, 2.5, and 2.8 G). Measurements of the spin concentration indicated that a relatively high concentration of NB- ions had been formed equivalent to about 0.5% of the surface oxide ions. It was suggested that 0;; ions were likely to be the source of the electrons, since transition metal ion impurities could be eliminated because of their low concentration. This work was extended by Che et al. (133) who used nitrocompounds and tetracyanoethylene (TCNE) to study the electron donor properties of MgO as a function of activation temperature (Fig. 12) and found two maxima in the concentration at about 220 and 700°C. The first was attributed to OHgroups acting as donors and the second to 02; ions. Cordischi et al. (1340) reported broadly similar results but attributed the electron donor center at low temperature of activation to a 0:; . . . OH- pair where the oxide ion was in a position of low coordination. These ideas have been taken further in a study by Coluccia et af. (1346) where the formation of NB- ions was found to be much higher on MgO smoke that had been etched by water vapor than on the normal smoke. These results are interpreted in terms of a considerable increase in the number of 0;; ions after etching of the regular cubic particles. In all these studies, one of the reactions involved is thought to be the donation of an electron from a surface ion of low coordination (0:;). The reported concentration of donors varies from about 0.2 to 2% of the surface ions measured for samples activated at 700°C. In none of the work described above has it been possible to see an EPR signal from the corresponding hole or 0- center formed when the 02; gives up an electron. This may be because the electron is delocalized over several ions giving a very broad

110

M . CHE AND A . J . TENCH

T ("C)

FIG.12. Radical-forming ability of MgO with TCNE as a function of activation temperature. [Figure according to Che et a/. (/33).]

signal, but this question has not yet been resolved satisfactorily. In a more extensive study (135) Cordischi and Indovina have compared the electron donor properties of several oxides (CaO, MgO, ZnO, A1,0,, and SO,/ A1,0,) using the formation of NB- as a probe. The electron donor power of the oxide surfaces after high-temperature activation decreases from CaO (100 arbitrary units) to MgO (18) to silica--alumina (< 1) in the same way as the basic strength (i.e., Lewis basicity) of the surface. The ability to transfer a single electron appears to parallel the ability to donate an electron lone pair (i.e., the Lewis strength). However, in many systems several different donor sites contribute to the overall electron-donating ability of the surface. In later work (986) Indovina and Cordischi use the concentration of 0;(see Section IV,B) on the surface as a measure of electron donor sites and find a broad correlation with surface base strength. Similar studies on A1,0, and TiO, using electron acceptors with electron affinities ranging from 1.77 to 2.84 eV (136) indicate that there is a range of strengths among the electron donor sites on the surface. The coordination of surface oxide ions has been explored further by following the adsorption of halogens on MgO (83). All the halogens are strongly adsorbed, and chlorine reacts to form chloride ions with no evidence

MONONUCLEAR OXYGEN O N OXIDE SURFACES

111

for the formation of Cl,. Heating to 300°C leads to the evolution of oxygen which corresponds in amount to the replacement of about 20% of the surface oxide ions and analysis for chloride showed that a stoichiometric quantity of chloride ions is formed, 2(C12),

+ 2(02-)\-+4(cI-),+

(26)

The reactivity of bromine with the surface is similar to that of chlorine, but it reacts with only about 10% of the surface oxide ions; and in contrast, iodine reacts to only a small extent and the iodide concentration corresponds to reaction with about 1% of the surface oxide ions. These results can be understood in terms of heats of formation of the solid oxide and halides (137) (viz., MgO, 143 kcal mol-'; MgCl,, 153 kcal mol-'; MgBr,, 125 kcal mol- I ; and MgI,, 87 kcal mol- '). The liberation of gaseous oxygen from the surface by bromine indicates that some M g 2 + 0 2 -ion pairs must be destabilized with respect to the normal bulk material by at least 18 kcal mol- A likely interpretation of these results is that iodine can only react with those 0;; ions in very low, possibly threefold coordination, whereas bromine reacts with 0: in four- and threefold coordination, and chlorine is able to react with a wide range of surface oxide ions. The origin of these effects lies in the changed coulombic and polarization effects present at the surface when compared to the bulk, and these will be discussed in more detail in the next section. The intermediate form of the oxygen after electron transfer is not known for certain, but after adsorption of chlorine a new EPR peak appears with g values of 2.0099 and 2.0020, together with a strong reflectance peak at 430 nm (3 eV). This latter peak decays as oxygen is evolved from the sample during thermal treatment, whereas the EPR signal decays more slowly. It seems probable that charge-deficient groups of oxygen ions of which the simplest would be O i - , may represent the intermediate form of the donor ions. Indovina and Cordischi (5%) have tentatively assigned an EPR signal observed under similar conditions to an 0:- species which has been formed from 0;; on the surface (see Section IV,B). Although several oxides, other than the alkaline earth oxides, form radical anions when electron acceptors are adsorbed, there is little evidence on the source of these electrons [except where transition metal ions are known to be present (133, 138) and where the resulting hole is localized]. In an early study of A1,0, (139),the A1-0- species was suggested as the electron donor site after activation at high temperatures. In general, it has been accepted that the donor is a hydroxyl ion on the surface (136,140) for alumina, titania, silica, and zirconia-titania. In a recent study by Druon et ul. (141), the EPR signal which appears when TCNE is adsorbed on type X and Y zeolites was analyzed as arising from two types of paramagnetic centers. These

'.

112

M. CHE AND A. J . TENCH

centers always appear together; one is clearly the TCNE- radical ion, the nature of the second center is less clear. It is located away from the first center and probably associated with an aluminium ion. B. SPECTROSCOPIC STUDIES

In the previous section we summarized the chemical evidence that oxide ions in a state of low coordination can act as electron donors. At the same time, spectroscopic evidence has been accumulated which shows that highly dispersed alkaline-earth oxides have optical absorption bands that are not present in the pure single crystal. This is surprising at first because the energy required for electronic excitation of bulk MgO corresponds to a frequency in the vacuum ultraviolet. In order to understand this we must look at the absorption process more closely. The absorption of light close to the fundamental absorption-band edge of an oxide leads to the excitation of an electron in the oxide ion followed by a charge-transfer process to create an exciton (an electron-hole pair) which is essentially free to migrate through the lattice, MZ+(nso)O2-(2po)-+ M2+(ns") OZ-(2p53s') Mz+(nso)O2-(2pS3s')+ M+(ns') 0-(2ps3s")

(27) (28)

These processes give rise to the electronic absorption bands of lowest energy observed in the pure undamaged single crystals which occur at 7.68 eV for MgO and 6.8 eV for CaO (142). Defects within the crystal structure are associated with optical absorption bands at reduced energies [for example, the anion vacancy band in the alkali halides (143)] because of the lower Madelung potential. The energy is still absorbed by the processes described in Eqs. (27) and (28), but the exciton is now bound to a defect and is equivalent to an excited electronic state of the defect. These ideas can be extended to the surface, In this case, the ions experience a reduced Madelung potential because of their lower coordination that leads to absorption of light at energies lower than the band edge. The various states of coordination will be associated with surface excitons at different energies. If the oxide ion were completely isolated from the lattice, then the optical absorption process would correspond to a simple electronic excitation of the oxide ion [Eq. (27)]. However, the experimental results (Table V) indicate that the energy required for this process decreases as the atomic number of the cation increases. This is consistent with the idea that a chargetransfer process is involved [Eq. (28)]; such a process may involve more than one cation because of delocalization of the electron but can be regarded as increasingly localized as the coordination of the ions is reduced.

MONONUCLEAR OXYGEN ON OXIDE SURFACES

113

TABLE V Absorption, Excitation, and Emission Bands of the Alkaline-Earth Oxide Powders

Oxide

Absorption (eVY

Excitation (ev)”

Emission (eVY

MgO CaO SrO BaO

5.70; 4.58 5.52; 4.40 4.64; 3.96 3.60; 3.22

>5.40; 4.52 > 5.40;4.40 4.43; 3.94 3.70

3.18 3.06 2.64 2.67

‘ Refs. 146. 147. Ref. 149. Ref. 149.

In principle, the optical absorptions in this region could also be associated with point defects on the surface either with or without trapped electrons and holes. However, the properties of the charged defects have been studied extensively, and the trapped charges can be thermally annealed at temperatures far lower than the normal preparation temperatures of these samples. In addition, they are characterized by optical and EPR spectra which are not observed in these samples. Contributions from point defects with no trapped charges cannot easily be eliminated. In fact, such a surface vacancy or divacancy would represent a localized state on a low-index surface associated with 4-coordinated ions. Intrinsic surface exciton absorption can be seen in high-surface area MgO, provided that the surface is thoroughly outgassed to remove surface carbonate and hydroxide. This was first reported by Nelson and Hale (144) who found that the reflectance spectra of high-surface area MgO, CaO, and SrO in uacua showed a strong fluorescence which could be quenched by oxygen to give a number of bands in the range 5.7-3.9 eV; these were ascribed to surface states. This fluorescence was confirmed for MgO by Tench and Pott (145); they observed the emission spectrum from the surface as a doublet with maxima about 400 and 440 nm. At this point, two separate approaches to the spectroscopic work were adopted which led to very similar results. The first expanded the earlier reflectance data and the second exploited the photoluminescent properties of the system. This latter method is more sensitive, since often well-resolved peaks appear in the spectrum instead of shoulders, but it is generally more limited in scope because a radiative decay of the excited state must occur for it to be observed. The reflectance spectra of the alkaline-earth oxides were examined in detail by Zecchina et al. (146, 147), but in order to do this they had to use an overpressure of 133 nm-2 (1 Torr) of oxygen to suppress the fluorescence. In spite of the large variety of possible configurations to be expected in a

114

M . CHE AND A . J . TENCH

high surface-area oxide, the spectra show only a few bands for each oxide (Table V). These bands are fairly broad (Fig. 13) and each band can be considered as representing a particular local surface coordination of the oxide ions. The bands in the spectra are changed by the adsorption of gases: those of lowest frequency corresponding to ions in lowest coordination are most affected. Oxygen has no effect on the MgO spectrum apart from a reversible quenching, but with CaO a new peak develops at 23,500 c m - ' . The nature of this species is not clear, but it probably corresponds to adsorbed 0; ; at the same time, the lowest frequency band decreases in intensity. This erosion of the bands becomes progressively more marked for SrO and BaO and the changes are irreversible. Nitrous oxide shows a similar effect with all the oxides; the bands are displaced reversibly toward the higher frequencies. Carbon dioxide and water vapor radically change the spectrum because of their strong chemical interaction with the surface, but even in this situation the bands at the lowest frequency are the first to disappear. The adsorbing gases convert low-coordination surface ions to ions of higher coordination, and the higher Madelung potential arising from this increase in coordination displaces the surface absorption toward higher frequency. A parallel approach has been to study the photoluminescent spectra of the alkaline-earth oxides (148-150). The absorption of light corresponds to the

iii VJ' cm-' FIG. 13. Reflectance spectrum of MgO after outgassing at various temperatures and exposure to 133 nmz (1 Torr) oxygen at room temperature to quench fluorescence: solid curve, outgassed at 773 K ; dotted curve, outgassed at 923 K ; dashed curve, outgassed at 1073 K . [Figure according to Zecchina er ril. (146).]

MONONUCLEAR OXYGEN ON OXIDE SURFACES

115

formation of an exciton as in the reflectance spectrum. In the absence of a quenching gas such as oxygen, the exciton may undergo radiative decay and light is emitted (Fig. 14) at a lower frequency (the Stokes shift); the difference in energy is lost to the vibration modes of the crystal M'O-

+

M 2 + 0 2 - + hv

(29)

The emitting site need not be the original site of absorption of light since the exciton may hop to other sites of low coordination, but the lower the coordination the less likely this is to occur. The photoluminescent measurements can give information about both the absorbing and the emitting sites; the excitation spectra correspond to the absorption bands obtained from the reflectance data (but only, of course, where the absorption leads to a subsequent emission) without the complication of added gases such as oxygen, and the emission spectra give information about the emitting site. The formation of the oxide surface can be followed by starting with hydroxide and dehydrating at successively higher temperatures in uacuo (148). For MgO, an emission excited at 365 nm (3.4 eV) and centered at 480 nm (2.6 eV) reaches a maximum between 200 and 300°C and has been attributed to surface hydroxyl groups. Two other emissions excited at 230 nm (5.4 eV) and 274 nm (4.5 eV) and centered close to 400 nm (3.1 eV) grow in at higher temperatures reaching a maximum between 800 and 1000°C (Fig. 15). These latter emissions are associated with the presence of 0;; ions. The behavior of these two types of centers with adsorbed gases is very different; the one linked to the OH- group is only slightly affected by 0, and CO, , whereas the emission linked to 0;; is immediately quenched by oxygen and destroyed by CO;, . It seems likely that the absorption/emission process linked to the OH- may be related to the center responsible for the electron donation at low temperatures. The absorption bands measured in the reflectance spectra and the corresponding

Exci tat ion / Absorpt lo n

I-f Emission

FIG.14. A simplified energy-level diagram of the processes involved in excitation/absorption of light followed by emission. The curved arrow ,I,. represents nonradiative decay leading to the Stokes shift.

116

M. CHE AND A. J . TENCH eV

I

6.0 5.0 1

.

200

1

3.0

L.0 '

1

I

1

1

1

'

1

I

300 LOO 500 W o v e l e n g t h (nml

I

60

FIG. 15. Photoluminescent spectra of MgO at 300 K after pretreatment at 800°C: (a) excitation spectrum of 390-nm luminescence; (b) emission excited at 230 nm; and (c) emission excited at 274 nm. [Figure according to Coluccia et a/. (149).]

bands in the excitation spectra match well for all the alkaline-earth oxides (Table V). Two peaks are observed in both absorption and excitation spectra and in general the agreement is very good, indicating that the overpressure of oxygen has no effect on the reflectance data. The excitation data could not be obtained for the highest energies because of low light intensities. The position of maximum intensity is given for the emission peak, but the shape changes with excitation energy and there is clearly more than one component involved. Some further evidence on the nature of the emitting sites was obtained by adsorbing hydrogen on SrO. The bands on the low-frequency side of the excitation spectrum are preferentially eroded by hydrogen adsorption, whereas oxygen quenches all the spectrum by interfering with the emission process. Since hydrogen may well be expected to interact with 0;; (Section VI,C), this adds weight to the argument that the sites responsible for the absorption of light are cation deficient or have a local excess of oxide ions (e.g., 0;;).Typical measured lifetimes are in microseconds. From a comparison of the optical absorption and excitation data for the oxides (Table V), it is clear that the energy decreases with increasing cation size along the series Mg to Ba. The bulk exciton transitions of these oxides also decrease in a similar manner (Table VI). It is possible to make a semiquantitative calculation of the intrinsic surface energy states using the approach of Levine and Mark (151) where the ions in an ideal surface are considered to be equivalent to bulk ions except for their reduced Madelung

MONONUCLEAR OXYGEN ON OXIDE SURFACES

117

constant. The ratio t of surface (E,) to bulk (Eb)band gaps is given by = (Y

-PIN

- PI

(30)

where y = c,/cbis the ratio of surface to bulk Madelung constant. The term p is given by ,u = 0.0347r(I - A)/C,Z

(31)

where I is the ionization potential (in electron volts) of the cation, A the electron affinity (in electron volts) of the anion (Zand A refer to the free-atom values), r the lattice parameter, and 2 the charge on the ions. Values for y and p are available and since the alkaline earth oxides are not likely to be fully ionic, particularly at low coordination on the surface, the calculation has been carried out for both Z = 1 and 2. Calculated values for the surface excitons on the surface (100) plane and on higher index planes are shown in Table V1. Comparison of these calculated exciton transitions with the experimental data in Table V shows that the main features of the results are reproduced. The energies for the (100) surface (5-coordinated ions) are only slightly shifted from the bulk, whereas those transitions corresponding to the higher index planes are much closer to the experimental data. On an atomic scale this means that ions whose coordination numbers are 4 and 3 are involved in the observed transitions, whereas 5-coordinated ions at the surface will absorb at higher energies closer to the bulk band edge. This theoretical treatment is approximate since it considers only an ideal surface and assumes that the electron affinity and ionization potential are constant for the different planes. In fact, the evidence already presented on electron transfer in Section VI,A indicates that the ionization potential varies with the coordination of the ion. Garrone et al. (154) have analyzed the Levine-Mark approach in more detail. They have reinterpreted earlier spectra (146, 147) and suggested that there is evidence for three absorption bands in MgO, CaO, and SrO rather than two, and that these do correspond to ions in 5-, 4-, and 3-coordinations. The highest energy exciton behaves similarly to the free exciton in the bulk, whereas the other two are typical of bound excitons and obey the Mollwoh e y relation (155, 156). Analysis of data using Eq. (30) showed that the ionic model with 2 = 2 accounts satisfactorily for the energies of the highest energy exciton in MgO, CaO, and SrO if it is assumed that 5-coordinated ions on (100) faces are involved. However, for the lower energy excitons the Levine-Mark theory no longer holds because of the increased covalent character of the bonds due to the reduced coordination. For 4-coordinated sites this can be analyzed in terms of a modified Levine-Mark theory as consistent with excitation, either of ions with a reduced charge on a smooth

TABLE VI Alkaline-Earth 0.wide Surface Esciton Leaels ( e V ) Calculated Using the Leuine and Mark Approach"

I?

SrO E' = 6.1

CaO

MgO = 7.8

Eh

=

7.0

?I

BaO = 4.4

Surface plane

Coordination of the ion

z = l

2 = 2

z=1

z=2

Z = l

z=2

z=1

2 = 2

(100) (110) (210) (211)

5 4 4 3

7.4 6.4 5.5 3.7

7.2 5.6 4.2 1.6

6.6 5.8 5.0 3.5

6.4 5.9 3.8 1.4

6.4 5.6 4.8 3.5

6.2 4.8 3.6 1.2

4.2 3.7 3.2 2.3

4.0 3.1 2.3 0.7

'I

Ref. 151.

* For comparison with the surface exciton, the energy of the first bulk excitonic transition has been used where possible; otherwise, the band-gap energy has been used. Ref. 142. Ref. 133.

MONONUCLEAR OXYGEN ON OXIDE SURFACES

119

(1 10) o r (210) face or of ions with a normal charge on a step or edge where a new y value corresponding to the Madelung potential a t that site has to be computed. Overall, the latter model is to be preferred since, for example, the morphology of the MgO crystallites, indicated by electron microscopy, is that of cubes with steps, etc., rather than prisms with (1 10) faces. If ions exist on the surface in various degrees of coordination, then the relative population of ions in each state of coordination should be strongly dependent on the surface topography of the particle and should change if the crystal imperfections change. For MgO, a comparison has been carried out (150) using MgO smoke, which is formed in nearly perfect small cubes (500-1000 A). This material would be expected to have a much smaller than normal concentration of 02,and this is confirmed by the photoluminescence spectrum (Fig. 16), which shows a much smaller than usual excitation peak at 274 nm. Electron micrographs show that etching in water vapor leads to erosion of corners and edges of the crystallites, together with the creation of

-

6.0

LO

5.0

eV

c

v)

.-

. I -

C

3

> L

0 c L

.-

n

L

Q

>

c .-

ln C 01

c

-C

I 1

I

nm

300

FIG. 16. Excitation spectra in uucuo of MgO smoke taken with 400-nm luminescence: (a) sample outgassed at 1200 K and (b) sample in contact with water vapor and then outgassed at 1200 K . [Figure according to Coluccia e t a / . (150).]

120

M. CHE AND

A . J . TENCH

pits in the surface. The photoluminescence spectra show a sharp increase in the peak at 274 nm (Fig. 16) which is attributed to 05; ions in the lowest, i.e., threefold coordination, confirming the general predictions of the model. Apart from the alkaline-earth oxides, some photoluminescence data have also been obtained with high surface area T h o , (157). Emissions at 2.82 and 2.78 eV are ascribed to vacancies in the surface and those at 2.52 and 1.97 eV to surface hydroxyl groups. The authors have not considered the possibility of 02, on the surface, but the dehydration process which is postulated as leading to the formation of vacancies could equally well give rise to 02, ions in the surface which would act as electron donors rather than acceptors. The excitation energy for these centers is about 5 eV, which is just less than the band gap of 5.85 eV (154). Some preliminary photoluminescent work with alumina indicates that it may be possible to study the more covalent oxides, and this work is then closely related to that on the short-lived 0 - ions described in Section IIl,B, where a covalent transition metal-oxygen complex can act as an electron-hole trap. It seems likely that it will not be easily possible to study the semiconducting oxides by this technique because in these systems nonradiative mechanisms usually predominate. In addition to these oxides, porous Vycor glass (PVG) has been reported to have a photoluminescent spectrum (158) which can be quenched by 0, and NH, (159). Excitation at 250 nm leads to an emission at 410 nm. This absorption of light well below the expected band gap (8.1 eV) has been interpreted in terms of ions in sites of low coordination, as discussed for the alkaline-earth oxides. The quenching processes have been attributed to the formation of adsorption complexes between the quencher molecules and the PVG surface. For oxygen at 300 K this is thought to consist of t w o processes : reversible weak adsorption complexes and the irreversible formation of (O;)5.The irreversible quenching with NH, was attributed to the adsorption of NH, on Lewis acid sites. Photoenhancement of the yield of propene from isopropoxide on PVG ( 1 6 0 ~has ) been attributed to 0 - centers formed by charge-transfer processes at low-coordination sites. The photoluminescence of lattice oxide ions of transition-metal oxides mixed or supported on conventional carriers has also been reported (160h). The luminescence is shown to occur from 0x0 complexes (MOJ- ( M = V, Mo, W, Cr) in which the transition-metal ion exists in a high oxidation state with a do electronic configuration. Since the d orbitals of the transition-metal ion are not occupied and therefore the d --d transitions impossible, So -+S, charge-transfer electronic transitions occur in the 0x0 complexes upon absorption of light. The result is that an electron is transferred from a filled molecular orbital localized mainly on the Oz--anions to a d orbital of the transition-metal ion. This leads to the formation of an excited singlet electronic state S, with two unpaired electrons, in which the total electron spin,

MONONUCLEAR OXYGEN ON OXIDE SURFACES

121

as in the So ground state, is equal to zero. However, there is a low-lying excited triplet state T I which is populated from S , as a result of nonradiative intersystem crossing. The deactivation of the TI excited state produces phosphorescence. Since TI +So transitions are forbidden, the lifetime of the T, state is fairly long. The excitation spectra coincides with the position of the charge-transfer spectra as measured by diffuse reflectance spectroscopy. In the case of supported transition-metal oxides, quenching of phosphorescence occurs on adsorption indicating that the luminescence centers are located at the surface, in contrast to mixed transition-metal oxides where these centers are in the bulk (1606). In many cases (716,r, 160b,c,d), the quenching is reversible at room temperature and probably occurs via the formation of a weak complex. It has been suggested that the transition-metal ion should be in a tetrahedral arrangement (1606) or exhibit a M=O-type bond (71b, 1 6 0 ~ for ) phosphorescence to occur. In some systems, for instance V,05 supported on PVG (71h), the phosphorescence spectrum can exhibit a resolved vibrational fine structure with the 0 -+ 3 transition the strongest, indicating that the V 5 + = 0 2 - bond will become longer in the excited states. In conclusion, there is now good evidence that oxide ions of unusually low coordination, 02; and 0:; for the alkaline-earth oxides, are present on the surface and their chemical and spectroscopic properties have been characterized. It is clear that they can act as electron transfer sites when the normal surface oxide ions (0;; for alkaline-earth oxides) do not, and their possible role in catalytic reactions will be discussed in the next section.

C . CHEMICAL REACTIVITY In contrast to the halogens, there are the weak electron acceptors such a 0, with an electron affinity of 0.44 eV (161) in the gas phase. Admission of oxygen to thermally activated MgO in uucuo gives a very weak EPR signal. which has been attributed to 0; (135, 162). Preadsorption of hydrogen leads to a much stronger signal in which several y,, components are visible and can be clearly identified as arising from 0; in more than one environment. A similar signal is obtained from CaO without the need for hydrogen adsorption, whereas no signal is obtained with A1,0, (163). These reactions in the absence of hydrogen can be interpreted if a small number of 0;; can act as very ready electron donors. However, no precautions were taken against photoeffects occurring at the surface, and it is known that light in the wavelength range 250-360 nm does cause photosorption of oxygen and the formation of adsorbed species (164). I t seems likely that the reactivity

122

M. CHE AND A. J. TENCH

of the oxide toward 0, without added gases might also be enhanced by very small amounts of residual hydrogen as OH- groups which are not removed during the activation process. The influence of hydrogen in substantially increasing the yield of 0; is not clear-again, photoeffects can be even more important unless the experiment is carried out in the dark. Cordischi et al. (165) have suggested that preadsorption of hydrogen increases the yield of 0; by adsorbing onto pairs of 0 , sites, created on the surface during thermal activation in uacuo, leading to O&-H+ species. These can then donate an electron when oxygen is added to give 0; adsorbed on a nearby cation and forming O,-H+. This last species is equivalent to an OH& radical and should be observable by EPR, but has not been reported. The adsorption of hydrogen on the MgO surface has been studied by Coluccia and Tench ( 1 6 6 ~ )At . low temperatures, the adsorption is largely molecular, and the photoluminescence spectra show that both 0:; and 0:; ions are involved, Infrared evidence (166b)shows that the room-temperature adsorption involves heterolytic dissociation [Eq. (32)] and is associated with 0;; ions,

In the dark, this reaction is likely to involve ion pairs of the lowest coordination, e.g., M:ZO&, whereas in the presence of UV irradiation other sites are involved, such as Mi: 0:;. The presence of such ion pairs on the model surface can be seen in Fig. 11. The coordination of the oxygen ion appears to be the dominant feature in determining whether the adsorption is molecular or dissociative. Similar ion sites are probably to be involved in the H,/D, equilibration reaction, and special configurations such as a triangular array of 0:; ions may play an important role. The increased yield of 0; after adsorption of hydrogen is most likely to arise from electron donation by the MEZ-H- species, the hydrogen either forms more surface hydroxyl groups or is evolved as a gas. Ito et al. (166c) have interpreted the thermal desorption spectra of adsorbed hydrogen using a similar model involving heterolytic dissociation of hydrogen. Coluccia et af. ( 1 6 6 4 have extended the 1R studies to cover CaO and SrO. For all of these oxides, the same mechanism is thought to apply. The isomerization of butenes over MgO is a catalytic reaction which has been studied by a number of workers and is thought to occur via an anionic mechanism which involves basic sites on the surface. Baird and Lunsford (167) were able to show a correlation between the concentration of electron donors on the surface and the 1-butene isomerization rate, and they suggested that the formation of a carbanion occurred on 0,- ions located on corner

MONONUCLEAR OXYGEN ON OXIDE SURFACES

123

sites of the cubic lattice. This is in line with the ideas proposed above for hydrogen adsorption and involves the intermediate formation of H+ . . .02; on the surface. Possibly, the most reactive sites are not involved in the catalytic reaction since the proton would not be easily removed under these circumstances, but an ion pair such as MiZO:, could well be the active site. Garrone et al. (168)have shown that the sensitization of MgO to electron donation by preadsorption extends to propene, butene, and acetylene. Ultraviolet reflectance measurements show bands characteristic of the carbanions, and the protons are assumed to react with 02; to form OH,. The addition of oxygen then leads to an electron transfer from the carbanion to form 0; ,whereas the radical then formed can oxidize or dimerize. They suggest that in this way 0; can be formed without the need for electron transfer from the solid to form a preexisting radical. However, it is clear that the oxide surface is involved and without the presence of 0:; the reaction will not proceed. The function of the 0;; would be to abstract a proton from the adsorbed molecule; alternatively an electron could be donated to the hydrocarbon molecule and then the 0, would abstract a hydrogen to form OH,. The photoactivation of lattice oxygen has been studied for V,O, supported on PVG (7Jb) and silica ( 7 1 4 . UV irradiation of the V,O,/PVG or V,O,/ silica systems at 25°C in the presence of CO molecules was found to lead to the formation of CO,. The quantum yields for the production of CO, closely followed the shape of the excitation band in the photoemission with a pronounced peak at about 340 nm. The CO was thought to form a complex with an excited V 5 + = 0 2 - species in the surface, leading to loss of oxygen ions from the lattice. Anpo ef al. (716) suggest that elongation of the internuclear distance in the excited state, together with the formation of the charge-transfer excited state (V4+-O-), would account for the easy photoformation of C O , in this system. In a more general study by Anpo et ai. (7Ze), it has been found that photoreduction with CO of metal oxides supported on PVG is confined to the oxides having metal-oxygen double-bond character, such as V,O,, MOO,, and CrO, . Their photoreducibility decreases in that order, ie., with decreasing lifetime of the charge-transfer excited triplet states determined from the phosphorescence decay curves. Anpo et al. (71c) have also studied the photoinduced reaction of C,H, on MOO, supported on PVG to form C2H, and 2-butene. The dependence of yield on the excitation wavelength was in good agreement with the photoluminescence excitation peak. The photoinduced reaction was thought to be closely associated with the charge-transfer excited triplet state of MOO,/ PVG. Fenin et al. (160d) have reported the oxidative dehydrogenation of 1-butene and related compounds over V,O,/MgO over the temperature range 350-550°C. Photoluminescent studies before and after adsorption of the I-butene indicated that a complex was formed with the V=O bond,

124

M . CHE AND A . J . TENCH

which quenched the luminescence. In related work, lwasawa and Ogasawara ( 1 6 0 ~found ) that the photoluminescence associated with the M o 6 + = 0 2 in fixed MoOJSiO, catalysts was almost completely destroyed by 10 Torr of cis-2-butene, indicating the formation of a surface complex.

VII.

Characterization and Reactivity of M =O

In contrast to the alkaline-earth oxides described in Section VI, the oxidation state of the metal ion in catalysts involving transition metal oxides can be easily varied, leading to the possibility of electron transfer from the cation and also of different kinds of metal-oxygen bonds. The main features of these systems are outlined below. In such catalytic oxides, the oxidation of a hydrocarbon can often be described using the early redox model suggested by Mars and van Krevelen (169) for the oxidation of napthalene. In this model, the hydrocarbon reacts with surface oxide ions to form oxidized products, leaving a reduced catalyst which is reoxidized to its initial state by reaction with gas-phase oxygen. The incorporation of the oxide ions into the oxidized products depends on their environment, and correlations have been made between the strength of the metal-oxygen bonds and their activity in the oxidation of several types of molecules (170, 171). It has been suggested that catalysts containing oxygen doubly bonded to metal cations would selectively oxidize alkenes (172). This has led a number of authors to characterize such bonds and to study their reactivity. The IR technique has been a powerful tool, since the M=O double bonds are known to give IR absorption bands in the 900-1100-cm-' region depending on the metal ion (172, 1734,6), whereas metal-oxygen single bonds (M-0-M) vibrate in the 800-900-cm- region giving broad absorption bands (174). When oxygen is less strongly bonded to the surface of the catalysts, a more rapid oxidation of the hydrocarbon takes place; on the other hand, too weak a bonding results in a more complete and less selective oxidation. The strength of the bonding between the metal ion and oxygen can be measured by the frequency of the 1R wavelength absorbed. In a systematic study of molybdates Trifiro ef a/. (175, 176) have observed a strong absorption at 940-970 cm- for iron, cobalt, and manganese molybdates, and at 920-940 cm- for bismuth molybdate. In the latter system, the lower frequency indicates more labile oxygen, and this catalyst is known to be active in mild oxidation ( 1 7 7 ~ ) .For molybdates of calcium, lead, and thallium, which are totally nonselective, no metal-oxygen double bonds are

'

'

MONONUCLEAR OXYGEN ON OXIDE SURFACES

125

detected by IR. This indicates that selective oxidation is related to a particular region of bond strength (177b).However, this approach seems now to be an oversimplification as recently discussed by Higgins and Hayden (127). The EPR technique has also been largely used to detect either the cation ofthe M(”-l)+=O pair (M(”-1)+is . V4+, Cr5+,M o 5 +,W 5 + ,. . .) obtained by reduction of M”’=O (6)or the oxygen species 0 - formed by y irradiation in vucuo of the double bond [Eq. (S)](40).In the first case, for samples reduced by heat treatment, the g tensor of the EPR signal observed at room temperature is axial and corresponds to a C,, symmetry because of the dominant role played by the M=O double bond. The case of molybdenyl compounds, including molybdenum catalysts, has been recently examined by Che et ul. (23) who measured the covalency of the M o 5 + = 0 bond. They also showed that the metal cation could have a vacant coordination position in the trans position to the “yl” oxygen which could be occupied on adsorption. This is in agreement with Van Reijen (179) who also proposed that the short M=O bond would be directed toward the interior of the catalysts ; whereas the long M-0 bond in the trans position would be directed toward the exterior and could be involved in the catalytic reaction. In contrast to this, Weiss et al. (4) proposed classical organic reaction mechanisms involving exposed double bonds, both of the M=O entities and of organic molecules, to explain the selectivity of catalysts in mild oxidation reactions. The characterization of chromyl, vanadyl, and molybdenyl catalysts by EPR is now well documented (23, 29, 178, 180, 181). Both EPR (182,183,184) and photoluminescence (71b, 160b-d) have been applied to the characterization of the charge-transfer complex obtained by y or UV irradiation of the M=O pair, refer to reaction 8. Both Krenzke and Keulks (185) and Yang and Lunsford (1116) have suggested that the charge transfer could also be obtained by thermal activation. However in these cases, the 0 - ion formed cannot be seen by EPR, because of either short lifetime or dipolar interaction, as explained in Section 111. Evidence for such ions comes from their photoluminescence spectra, their chemical reactivity (Section II1,C) and from theoretical calculations. For example, calculations by Surratt and Kunz (186) and Klein (187) show that the 0- ion is stable in NiO and COO, provided it is associated with cation vacancies. The existence of these species on the surface would explain the observed dissociative chemisorption of hydrogen. These results also suggest that 0 - ions may be formed more readily from oxide ions on unsupported metal oxides where cation vacancies exist on the surface. According to this model the role of gaseous 0, is to replenish the oxygen ions which are removed during the reaction (I116). From the preceding discussion it is clear that where transition metal ions are involved, there is substantial evidence that charge-transfer processes of

126

M. CHE AND A. J. TENCH

the kind shown below occur quite easily,

under a variety of conditions and the process may be considered as an equilibrium. From electrostatic arguments alone we can see that charge transfer to form M("-')+0- will lead to a weakening of the M-0 bond. From the ideas developed in the previous section, it is clear that this weakeningof the bond is likely to be further favored if the "yl" oxygen is in a position of low coordination. In fact, the process 0'-

+

0-

+ e-

(34)

becomes energetically more favorable as the coordination of the ion decreases because of the decreased Madelung contribution. If we now relate this to catalytic reactions, it seems evident that both increase of temperature and decrease in coordination of the oxygen in the MO bond will favor the formation of 0-.It is interesting to consider that in a catalyst operating under real conditions, there may be no clear distinction between 0- and 0 2 -ions on the surface at the catalytically active sites. This idea has not been suggested previously, but it may prove to be a useful approach in considering the mechanisms of catalytic reactions.

Appendix.

The EPR Parameters of 0 - Ions'

Tensor Matrix

91

92

93

Reference

2.0155 2.0193 2.0385 2.0386 2.0389 2.0516 2.0650 2.0697 2.0705 2.0806 2.226

41 188 189 77 32 44 33 78 79 9, I0 4%

Single crystals

Axial symmetry: Be0 ZnO MgO Ca,(PO,),, @F, SrCI, CaO SrO NaCl K Br

2.0026 2.0024 2.0032 2.0033 2.0032 2.0012 2.0032 2.0021 2.00 I3 I .9976 1.987

2.0155 2.0193 2.0385 2.0386 2.0389 2.0516 2.0650 2.0697 2.0705 2.0806 2.226

127

MONONUCLEAR OXYGEN ON OXIDE SURFACES

Tensor Matrix KCI RbI NaI KI Orthorhomhic symmetry: ZnO

Mg o NaCl KBr KCI RhCl

91

92

Y3

1.981 1.9733 1.9769 1.9733

2.258 2.2888 2.2931 2.3023

2.258 2.2888 2.2931 2.3023

2.0028 2.0038 2.0038 2.00328 1.9976 1.9508 1.9475 1.9154

2.0173 2.0180 2.0182 2.03668 2.0542 2.2346 2.2217 2.2344

2.0183 2.0191 2.021 7 2.03960 2.1214 2.4416 2.4524 2.5330

Reference 45u

7 7 7 188

I90 190 24u

I91 8 8 8

Powders or frozen solutions Axial symmetry: As,O,/SiO, MoO,/SiO,

ZnO

Mo0,/A1,03 WO, /SiO, V,O ,/SiO, Cr/SiO, A1203

MgO Aqueous alkali metal hydroxides NaOH + ice Orthorhombic symmetry: AI20,/Si0, P,O,/SiO, SiO, B,O,/SiO, MgO

"

2.0055 2.002 2.006 2.006 2.0026 2.003

2.0055 2.019 2.020 2.020 2.021 2.021

2.003

2.023

2.008 2.012 2.024 1.988 2.0013 2.002

2.024 2.026 2.026 2.030 2.033 2.042 2.063-2.070

2.0055 2.019 2.020 2.020 2.021 2.021 (site 1) 2.023 (site 2) 2.024 2.026 2.026 2.030 2.033 2.042 2.063-2.070

2.002

2.070

2.070

2.003 2.0076 2.003 2.002 2.0016 2.0016

2.012 2.013 2.007 2.012 2.0193 2.0297

?

40

2.019 2.029 2.038 2.0472 2.0505

39 43 40 14

?

3Y 31 17 59 30u 306

30b 36 19

37,38 24b 192 16 193

28

14

The order of the 0 - species listed in each section is given with increasing gL or g3 values.

128

M. CHE AND A. J . TENCH NOTEADDEDI N PROOF

An oxidized surface state model of vanadium oxides and its application to catalysis have been discussed by A. Anderson [ J . Solid Stare Chem. 42, 263 (1982)]. The author concludes that 02-ions, in the form of V=O surface groups, are responsible for the catalytic oxidation of hydrocarbons. ACKNOWL~DGMENTS The authors acknowledge facilities provided by A. E. R. E. Harwell during the writing of this review, and M. Che acknowledges a European Exchange Fellowship and a Vacation Associate appointment at Marwell. They are also very grateful to Mr. K. Y. Che for his help in collecting some references. and to Dr. A. Szuszkiewicz and Dr. A. E. Hughes for discussions. The authors wish to dedicate this review to the memory of Jiiri Kukk, Estonian Professor of Chemistry, who died in a Soviet labor camp on March 27, 1981 at the age of 40. REFERENCES I . Lunsford, J. H., Cutul. Rev. 8, 135 (1973). 2. Kaye, G . W. C., and Laby, T. H., “Tables of Physical and Chemical Constants” 14th ed., p. 268. Longman, London, 1973. 3. Cotton, F. A,, and Wilkinson, G., in “Advanced Inorganic Chemistry,” 3rd ed., p. 58. Wiley, New York, 1972. 4 . Weiss, F., Marion, J., Metzger, J., and Cognion, J. M., Kincf. Kutul. 14,45 (1973). 5 . Lunsford, J . H., Ado. Curul. 22, 265 (1972). 6. Che, M., in “Magnetic Resonance in Colloid and Interface Science,” (J. P. Fraissard and H. A. Resing, eds.), p. 79. Reidel, Dordrecht, 1980. 7. Brailsford, J. R., Morton, J . R., and Vannotti, L. E., J. Chrm. Phys. 49, 2237 (1968). 8. Brailsford, J. R., and Morton, J . R., J. Chem. Phys. 51,4794 (1969). 9. Nistor, S. V., and Stoicescu, G . , Rru. Roum. Phys. 16, 515 (1971). 10. Nistor, S. V., and Ursu, I., Rev. Roum. Phys. 16, 495 (1971). I f . Iyengar, R. D., Codell, M., Karrd, J. S., and Turkevich, J., J. Am. Chew. Soc. 88, 5055 (1966). 12. Kuhn, H. C., in “Atomic Spectra,” p. 198. Longmans. London, 1962. 13. Tench, A. J., and Lawson, T., Chew. Phys. Lett. 7,459 (1970). 14. Tench, A. J., Lawson, T., and Kibblewhite, J . F. J., J. Chem. Soc. Furuduy Truns. 168, I169 (1972). 15. Williamson, W. B., Lunsford, J. H., and Naccache, C., Chwiz. Phys. Lrrt. 9, 33 (1971). 16. Wong, N. B., and Lunsford, J. H., J . Chrm. Phys. 55, 3007 (1971). 17. Kolosov, A. K., Shvets, V. A,, and Kazansky, V. B., Chrm. Phys. Lrrt. 34, 360 (1975). 18. Griscom, D. L.. Taylor, P. C., Ware, D. A., and Bray, P. J., J. Chem. Phys. 48,5158 (I 968). 19. Shvets, V. A., Sapozhnikov, V. B., Chuvylkin, N. D., and Kazansky, V. B., J. Cutul. 52, 459 ( I 978). 20. Kemp, J . C . , Cheng, T. C., Izen, E. H., and Modine, F. A., Phys. Rru. 179, 818 (1969). 21. Dunn,T. M., Trans. FuruduySoc. 57, 1441 (1961). 22. Shimizu, T., J. Phys. Soc. .//,? 23,848 I. (1967). 23. Che, M., Fournier, M., and Launay, J. P., J . Chrm. Plzys. 71, 1954 (1979). 24u. O’Mard, W. C., and Wertz, J. E., J . Mugn. Reson. 8, 366 (1972). 24h. Lipatkina, N. I., Chuvylkin, N. D., Shvets, V. A,. and Kazansky, V. B., Kinrt. Kurd. 19, 1561 (1978).

MONONUCLEAR OXYGEN ON OXIDE SURFACES

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

52. 53. 54. 55. 56. 57. 58. 59. 60. 61. 62. 63. 64.

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ADVANCES IN CATALYSIS. VOLUME 31

Sulfur Poisoning of Metals C . H . BARTHOLOMEW

P . K . AGRAWAL

Department of’ Chemical Engineering Brigham Young University Provo. Uiah

School of Chemical Engineering Georgia Institute of Technology Atlanta. Georgia

J . R . KATZER Centerfor Catalytic Science and Technologj Department of Chemical Engineering University o j Delaware Newark . Delaware

I . Introduction . . . . . . . . . . . . . . 11. The Nature of Metal-Sulfur Bonds . . . . . . . A . Sulfur Bonding Chemistry . . . . . . . . . B. Thermodynamics of Bulk Sulfides . . . . . . . I11. Sulfur Adsorption on Metals . . . . . . . . . A . Surface Structures . . . . . . . . . . . B . Adsorption Mechanismsand Kinetics . . . . . C . Adsorption Stoichiometries . . . . . . . . D . Stability of Sulfur Adsorbed on Metal Surfaces . . . IV . Effects of Sulfur on Adsorption of Other Molecules . . . A . Nickel Catalysts . . . . . . . . . . . . B. Other Metals . . . . . . . . . . . . V . Effects of Sulfur on Catalytic Activity and Selectivity Properties of Metals . . . . . . . . . . . . A . Experimental Considerations . . . . . . . . B . CO Hydrogenation Reactions . . . . . . . . C . Steam Reforming . . . . . . . . . . . D . Ammonia Synthesis . . . . . . . . . . . E . Hydrogenation, Dehydrogenation. and Hydrogenolysis of Organic Compounds . . . . . . . . . . VI . Regeneration . . . . . . . . . . . . . . VII . Conclusions and Recommendations . . . . . . . A . Conclusions . . . . . . . . . . . . . B. Recommendations . . . . . . . . . . . References . . . . . . . . . . . . . . .

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I.

Introduction

Catalyst poisoning is one of the most severe problems associated with the commercial application of catalysts. It is a phenomenon whose global behavior is studied extensively in industrial laboratories to allow adequate prediction of commercial catalyst life and commercial behavior. Yet, a quantitative understanding of the intrinsic rates and mechanisms of catalyst poisoning is generally lacking, partly because of the complexity of poisoning processes and partly because of the lack of sufficiently careful studies of these processes. Probably the most severe poisoning encountered in catalytic systems is that induced by sulfur on metal catalysts. With the increasing application of supported metal catalysts and their potential large-scale applications for synthesis gas (CO H2)conversion to fuels and chemicals, sulfur poisoning is becoming increasingly more important, particularly since synthesis gas produced by coal gasification contains significant amounts of sulfurcontaining compounds such as H,S and COS. Sulfur apparently bonds so strongly to metal surfaces that marked activity reduction occurs at extremely low gas-phase concentrations of sulfur-containing compounds. In commercial practice the life of supported metal catalysts may be reduced to only a few months or weeks in the presence of only ppm quantities of sulfur contaminants in the feed. Because of the essentially irreversible adsorption of sulfur compounds on metals, regeneration is usually impossible or impractical. In spite of its obvious practical importance, sulfur poisoning has received only moderate attention in the literature. In fact, the most recent comprehensive review of the literature dealing with poisoning of metals was by Maxted in 1951 ( I ) . More recently, Madon and Shaw (2) reviewed the pre1970 literature describing the effects of sulfur in Fischer-Tropsch synthesis and methanation. During the last decade a great deal of fundamental information has accumulated regarding the interaction of sulfur with surfaces, and recent investigations involving metal catalysts have led to an enhanced understanding of the role of sulfur in catalyst poisoning. This review is an attempt to integrate available information on the interaction of sulfur with metal surfaces with that of these recent poisoning studies to provide a more complete picture of' sulfur poisoning and of the mechanism thereof. In this review the discussion focuses on several fundamental questions :

+

0 How strong are surface metal-sulfur bonds and why are they as strong as they are? Under what conditions of temperature and partial pressure of the sulfur-containing compound does less than complete coverage of the surface occur?

SULFUR POISONING OF METALS

137

0 What is the mechanistic nature of sulfur adsorption on metal surfaces and what are the surface structures formed? 0 What is the mechanism of sulfur poisoning of metals? Is it due to a geometrical blocking or does it involve electronic effects that may propagate many atomic distances away from the adsorption site? 0 Is the mechanism of sulfur poisoning a function of the metal, of the reaction and/or of reaction conditions; can the dependence of the system on any of these variables be predicted a priori? 0 What are the rates of sulfur adsorption and of sulfur poisoning and can they be predicted? Can catalyst life in commercial catalyst applications be predicted based on poisoning mechanisms and models? 0 What are the mechanisms of sulfur removal from surfaces and can these mechanisms be used in the development of regeneration techniques for sulfur-poisoned catalysts? 0 How does sulfur poisoning fit more generally into the overall picture of the catalytic behavior of metals?

In the ensuing discussion we consider the fundamental characteristics of sulfur adsorption on metals, the effects of adsorbed sulfur on the adsorption of other gaseous species, the effects of sulfur on catalytic activity and selectivity, and regeneration of sulfur-poisoned catalysts. The discussion of reaction studies emphasizes CO hydrogenation but covers other reactions including steam reforming, ammonia synthesis, hydrocarbon hydrogenation and hydrogenolysis. In addition, practical experimental problems are considered. Indeed, this review will stress the importance of experimental techniques which provide definitive, fundamental information regarding sulfur adsorption and poisoning. Finally, recommendations for further work are included.

II.

The Nature of Metal-Sulfur Bonds

A. SULFURBONDINGCHEMISTRY The electronic structure of sulfur is 3s23p4; therefore, bonding involves s and p orbitals. Bond formation with transition elements involves mainly d orbitals. Maxted ( 1 ) suggested that sulfur compounds chemisorb on transition metals by forming bonds in which previously unshared electrons in the sulfur atom are donated into the d orbitals of the metal. In a recent

138

C. H. BARTHOLOMEW ET AL.

study of chemical bonding of sulfur with metals of groups IB, IIB, and IIIA using electron spectroscopy (ESCA), it was shown that in sulfides of Cu, Ag, Zn, Cd, Ga, and In, the d states of the metal are significantly involved in the chemical bond, and this bonding has a significant effect on the energies of the sulfur orbitals that are involved (3). For instance, for Cu and Ag the bonding involves the 3d and the 4d metal orbitals and predominantly 3p sulfur orbitals; for Zn and Cd, the 3d and the 4d metal orbitals interact more strongly with 3s sulfur orbitals. Similar information is not available for other metals. Quantum chemical calculations have been made for sulfur bonded to small Ni clusters (4). The calculations for NiS suggest that the Ni-S (T bond is composed of a Ni 4s orbital paired with a single occupied S 3p orbital. Calculations for Ni,S and Ni,S suggest similar bonding. Capehart and Rhodin ( 5 ) have suggested on the basis of low-energy electron diffraction (LEED) and angular resolved photoemission spectroscopy that nickel 3d orbitals are not involved in bonding of sulfur to the Ni(ll1) surface. Instead, the Ni-S bond in the Ni,S structure involves Ni 4s and S 3s orbitals. For Ni,S the S-Ni bond is stabilized by interaction between S 3s and Ni 3d orbitals. The calculated bond length for NiS clusters is 0.191 nm, which is substantially shorter than that for bulk nickel sulfide (4). For Ni,S clusters the calculated bond length of 0.221 nm is in good agreement with a value of 0.216 ? 0.016 nm obtained from LEED for S on a Ni(100) surface (6). The calculated bond strength ( 4 ) for NiS is 320 kJ/mol, which is in reasonable agreement with current experimental estimates (7) of 341 -t 15 kJ/mol. The bond strength for Ni,S is 518 kJ/mol(4), indicating that bridge bonding is much more favorable than bonding directly over a single Ni atom. Because sulfur has a lower electronegativity than oxygen (8),metal sulfides tend to be much more covalent than metal oxides. Sulfides frequently occur as nonstoichiometric phases, some with semimetallic behavior. Only the more electropositive elements-alkali and alkaline earth metals-form sulfides which are mainly ionic. The ionic character of oxygen compounds as compared with sulfur compounds is illustrated by the difference in boiling point of H,O (373 K) and H,S (212 K), indicating lack of bond polarization for H,S. Auger electron spectroscopic (AES) studies show that the valence electron densities around the Ni atom change very little upon sulfide formation, whereas marked electron transfer away from the Ni atom occurs upon oxide formation (9-11). That metal-sulfur bonds are nonionic for transition metals is also suggested by recent theoretical studies of Ni,S clusters (4, 12), showing a net charge transfer between nickel and sulfur to be between 0.4 and 0.6 electrons per sulfur atom. Thus, theoretical calculations have considerable potential in quantifying the nature of metal-sulfur bonds.

SULFUR POISONING OF METALS

139

B. THERMODYNAMICS OF BULKSULFIDES Due to their nonionic nature, metal sulfides are often nonstoichiometric rather than stoichiometric compounds in the classical sense. For most systems (Cr-S, Ni-S, Fe-S, Co-S, etc.), many phases occur, and each phase exists over a limited range of sulfur pressure and temperature. The nickelsulfur system, for example, shows many phase transitions and bulk phases at high temperatures and large concentrations of H,S (13-16). Data from Rosenqvist (13) in Fig. 1 show that at temperatures of catalytic interest (673-773 K), nickel exists in the metallic state only at PH2s/PH2 values below 10-3-10-4. At higher H,S concentrations, the most stable bulk sulfide

FIG.1. Plot of H,S/H, versus IjT for the nickel-sulfur system (13).

140

C. H. BARTHOLOMEW ET A L

NiJ, is formed. Thus in commercial processes under reducing conditions involving high-sulfur-content feeds, e.g., steam reforming, H2S concentrations may be high enough to cause bulk sulfide formation of the Ni catalyst, whereas typical levels of H,S encountered in processing low-sulfur feeds are in the ppm range, and under these conditions bulk metal sulfides are not stable and do not form. Therefore, the poisoning observed under these low-sulfur concentration conditions must be interpreted in terms of sulfur adsorption on the surface of the metal (surface sulfide formation). The thermodynamic properties of bulk metal sulfides are fairly well documented (13-20). Multiple phases similar to those of Ni are observed for Fe, Co, and other metal-sulfur systems of catalytic interest. Figure 2 shows 0

300

600

900

1 ; 1 1 1 1 1 1 1 1 1

1200

1500

I I 1 1 1 1

I800

2100

I l l l l l

0-

-100

r-

-200-

RuSp

-300

5 E

-

-400-

\

7 1

-10001

I l l / I ( I I I I I I l l I l l I I I I I I

0

300

600

900

T (K)

1200

1500

BOO

2100

FIG.2. Standard free energy of formation of various sulfides (kJ/mol S,) as a function of temperature. To change basis to H,S subtract AG for formation of HIS (also shown above). (Ref. 19.)

141

SULFUR POISONING OF METALS

the temperature dependence of the standard free energy of formation of various bulk sulfides. Table I gives the free energy of formation of bulk metal sulfides at 300 and 600 K (spanning the typical temperature range for catalytic reactions). Table 11 gives thermodynamic properties for sulfides of Fe, Co, and Ni. The values for the thermodynamic properties (Tables I and 11) are per sulfur atom and based on the following stoichiometry: (X/.Y)M

+ H2S * (1/y)M,S, + H,

(1)

where AGO is defined by

In the typical range of reaction temperature, the free energy of formation for a given metal sulfide is insensitive to temperature (Table I). However, large differences exist among these metals in terms of the stability of their bulk sulfides. Metals more commonly used as catalysts, such as Pt, Ni, Ru, Rh, Ag, Fe, Co, and Ir, have lower free energies of formation of their bulk sulfides, indicating that relatively large gas-phase H2S concentrations are required for stable bulk sulfides to exist. Some other metals, such as Cr, Mn, Mo, Re, Ti, V, W, Zn, and Zr, which may be used either in alloy form or in oxide form as a catalyst or as a support, have much higher free energies of

AG; (kJ/g atom)

AG;' (kJ/g atom) Element Ag Cd

co Cr cu Fe Hf Ir Mn Mo Nb

Sulfide

300 K

600 K

Element

-1.3 -117 -65.3 - 39.3 - 54.0

-5.0 - 105 -55.3 - 39.8 -47.7 -108 -52.3 - 57.3 - 123 - 18.4 - I80 - 76.8 -70.3 - I16

Ni

-110

-46.5 - 58.2 - 131 -31.4 -185 -85.4 -77.4 -124

Pt Re Rh Ru Ta Ti V

w

Zn Zr

Sulfide Ni,S2 Ni,S, Ni,S2-, PtS PtS, ReS, Rh,S RuS, TaS, TIS, "2%

WS, ZnS Zr2,

300K

600K

-65.7 -55.7 - 54.0 - 39.8 + 10.9 -69.1 -3.1 - 28 -7.1 -80.8 - 278 - 69.1 - 168 - 123

-56.1 -48.6 -58.2 - 26.4 +23.0 -97.5 - 10.9 -

- 2.5 -84.1 - 266 -60.7 - 152 -115

Values computed from the data provided in Refs. 13 and 1 9 : basis is per mole of H,S [seeEq.(l)].TochangetoS, basisaddAG"forH, + 0.5S2 = H,SforwhichAG" = -90.64 + 0.049T kJ/mol H,S.

142

C. H. BARTHOLOMEW ET AL.

TABLE I1 Molar Free Energies, Entropies, und H e m of Formation of Iron, Cobalt, and Nickel Su@des at 1000 K (727 C)u.h

Compound FeS FeS, Co33

c0,s.Y

COS Co,S,

cos,

Ni,S, (lowtemperature form) Ni,S, (hightemperature form) Ni,S, NiS NiS,

AG (kJ)’

A S (J/K)’

AH (kJ)

SIOOoK (JiK)”

- 55.9 - 5.4 - 39.7 -41.6 -39.1 - 26.2 - 5.2 -43.3

- 3.4 - 42 +0.7 -33.9 -21.3 - 29.8 -48.2 - 32.2

-59.3 -41.3 -39.0 - 75.5 -60.5 - 56.2 -53.5 -75.5

149 37.7 175 I26 131 105

-48.0

-5.2

-53.0

183

- 38.6 - 33.5 + 1.2

-24.2 -22.6 -39.8

-62.8 -56.1 -38.5

131

~~

70.5

156

I43 80 ~

~

~

~

~~~~

~

Adapted from Ref. I S . * Reference states: solid metals, hydrogen sulfide, and hydrogen at 1 atm. To obtain the free energy and heat of formation from S, (gas), -40.9 and -90.3 kJ, respectively, must be added for each atom of sulfur in the compound. The entropies of the reference states at 1000 K have been calculated from their entropies at 298 K, and their specific heats between 298 and 1000 K, as given by Kubaschewski et al. (20). The following values for S,,,,,, were used: Fe = 67.0, Co = 66.1, Ni = 67.8, H,S = 252, H, = 166. The accuracy of the calculated entropy values is estimated to 4.2 J/K per atom of sulfur in the compound.

formation of their bulk sulfides. This suggests that sulfur poisoning of one metal may be reduced by combining it with another metal having a higher free energy of formation of the bulk sulfide. Observations supporting this hypothesis have been reported (22-24). For example, nickel catalysts containing Cr, Mn, and Mo oxides have been reported to be more resistant to sulfur poisoning (21-23). The presence of chromium or other oxides in Ni-containing catalysts are said to provide adsorption sites for sulfur atoms, thus preventing or delaying the adsorption of sulfur on Ni (23, 24). However, this explanation is unlikely in those instances where the stability of the oxide exceeds that of the metal sulfide. For example, chromium oxide is much more stable (AG” = - 1046 kJ/mol) than chromium sulfide (AGO = - 109 kJ/mol). In the case of the ZnO-Cr,O, methanol synthesis catalyst, it is proposed that Cr,O, induces reversibility to sulfur poisoning (22),but definitive information to support this hypothesis is absent. On the other hand, in methanol synthesis on Cu-ZnO catalysts under similar conditions, sulfur causes essentially irreversible poisoning (25),

SULFUR POISONING OF METALS

143

suggesting that ZnO does not provide protection from sulfur poisoning even though ZnS is more stable than Cu,S (AGO = -168 kJ/mol for ZnS, AG" = -46.5 kJ/mol for Cu2S). Whether such bulk gettering effects are transient or steady state is unclear. The mechanisms of such effects and whether they are due to bulk or surface phenomena is also unknown. Thus it is clear that bulk thermodynamic information used in a cursory way does not help in rationalizing observed sulfur-poisoning behavior. Furthermore, since most metals of catalytic interest do not form stable bulk sulfides under typical reaction conditions, the observed severe poisoning by sulfur suggests that surface rather than bulk thermodynamics may be required.

111.

Sulfur Adsorption on Metals

A number of sulfur compounds known to poison catalysts include H,S, CS,, MeSH, Et2SH, Me,& EtS, thiophene, COS, SO,, and SO,. These sulfur compounds have unshared electron pairs which can lead to very strong chemisorption on the metal surface. Under reducing conditions, essentially independent of the starting compound, the adsorption on the surface is typically dissociative, leaving a reduced sulfur atom strongly bonded to the surface. Thus results obtained with H2S adsorption have general applicability and constitute the majority of the literature results. In order to be able to interpret quantitatively the extent and nature of poisoning by sulfur, it is essential to know the structure and bonding of sulfur to metal atoms at the surface. Thus, surface structures are considered first in this section followed by adsorption mechanisms and stoichiometries and finally stability of surface metal sulfides. Structural information regarding the geometry of the surface and of the adsorbed sulfur layer is best obtained from single-crystal studies using LEED. The discussion of structure begins with single-crystal planes and then proceeds to other metal configurations including foils and supported metals. Since the sulfur-nickel system has been studied most extensively, it will be considered in the greatest detail. A . SURFACE STRUCTURES 1. Nickel

A number of recent surface studies (5, 6, 26-50) have provided useful information regarding structures of sulfur adsorbed on various faces of Ni (Table 111). During the initial stages of sulfur adsorption on clean Ni single-

TABLE Sulfur Adsorption on

Crystal face

Technique used

Reference

Remarks

LEED", AES". work function measurements LEED

34

Adsorbate forms two structures: ('(2 X 2). /J(2 X 2)

42

Angle-resolved UPS' LEED

44

Angle-resolved UPS

41

LEED. AES

43

LEED

45

ELEED*

40

LEED

29

LEED

6

Presence of two structures: ((2 x 2) a n d p ( 2 x 2) Sulfur adsorption induces surface states in Ni p(2 x 2) and ,jT+ , / T R 30" structures; two-dimensional surface sulfide Adsorption effect on S is identical on the three planes in terms of' inducing two extra levels in Ni at comparable energies Surface segregation of S from hulk ; (,(2 x 2) structure observed for both S and C Electronic surface resonance band structure of S-covered surface can be described by 2-D free-electron model Distance of S atom from plane of Ni surface is d, = 0.13 nm 4 2 x 2) overlayers of S on Ni(100) correspond to S atom in the fourfold sites at d, = 0.13 +_ 0.01 nm Sulfur atoms reside in highcoordination sites; all nearestneighbor Ni-S bond lengths are less than those of stable bulk compounds Sulfur adsorbs in hollow sites 0.13 nm above the Ni surface; interaction between S and Na during coadsorption-Na bonds to S rather than Ni, but electron transfer is small, < 1 electron/Na atom For the same adatom sites, coordination for binding to Ni atoms may change depending on S coverage 4 2 x 2) S forms in a bridge-type Ni,S structure, S is adsorbed on the top of the Ni surface layer Ni ~Sbond lengths are same for p(2 x 2) and 4 2 x 2) structures

42

Dynamic LEED

31

Analysis of INS' ( 3 3 ) LEED (29a 33,469, and UPS data INS, work function measurements, LEED LEED

32

33

30

" Low energy electron diffraction. Auger electron spectroscopy. Ultraviolet photoemissions spectroscopy. 144

111 N i Single-Crystal Planes

Crystal face

Technique used LEED

Reference

Remarks

35,36

High-coordination sites preferred; reorientation of surface Ni atoms to (100) structure upon S adsorption on Ni(lI1); formation of S layer on the top of reoriented Ni layer Preliminary adsorption sites are the highest coordination sites on each surface; reconstruction into a surface compound-2-D sulfide, consisting of S and Ni (both in ionic form) in the reconstructed layer; 2-D sulfide more stable than 3-D sulfide Faceting of stepped Ni surface, with C as a n adsorbate; step coalescence is inhibited by S adsorption. Strong Ni interaction with C and S With S bonded to three Ni atoms, Ni-Ni bond length increases by 0.04 nm; Ni-Ni interactions weaken due to longer Ni-Ni distance Segregation of S to 0 = 0.3 lowers Ni LEED intensity. N o additional spots due to sulfur. Interpreted a s random sulfur For 0.18 < 0 < 0.3 S adsorbs in I-D chains along the [ITO] direction only detectable by RHEED. For 0.3 < 0 < 0.48 'sulfur adsorbs in a (5 o x 2) structure Sulfur adsorbs as p(2 x 2) a t 0.2 < 0, < 0.25, ( J ~ X,,IT) R 30 at 0.25 < 0, < 0.4, ((20 x 2) at 0.4 < 0, < 0.5 Heating or exposure to H,S of a c(2 x 2) structure introduces a new binding state in the subsurface Study of the detailed electronic structure of c(2 x 2) and p(2 x 2) overlayers shows important differences in the S 3p levels. Sulfur adsorption causes changes in the Ni 3d bond

LEED, radioactive tracer "S

26

Stepped surface, (1 10) and (1 1 1) planes

LEED, AES

38

Organometallic Ni-S metal cluster

X-ray measurements, LCAO MO theory

37

LEED, AES

46

RHEED', LEED

47

(111)

LEED, AES, flash desorption

49

(100)

Angle-resolved UPS. L E E D

sou

( 100)

Angle-resolved UPS

5Ob

~

Elastic low energy electron diffraction. Ion neutralization spectroscopy. Reflection high energy electron diffraction. 145

146

C. H. BARTHOLOMEW ET AL.

(001)

(110)

Ibl

(111)

2 36?004 130?01

L-2 17?%1

FIG. 3. Location of adsorbed atomic sulfur on the (IIO), (IOO), and (111) planes of Ni determined from LEED intensity measurements. Solid circles represent sulfur atoms; open circles, nickel atoms. Dimensions are given in angstroms (Ref. 6 ) ; “a” denotes top view, “b” side view.

crystal surfaces (IOO), (I lo), and (1 1 I), sulfur atoms reside in high-coordination sites, i.e., the atomic hollows of the surface (6,26-32,40) (see Fig. 3). For example, on a Ni( 100) surface, sulfur is adsorbed in an ordered p ( 2 x 2) overlayer, bonded to four Ni atoms up to 0 = 0.25’ and from 8 = 0.25 to 0.5 with an ordered 4 2 x 2) overlayer with each sulfur atom bonded to two Ni atoms. The same two structures are obtained from adsorption of thiophene, H,S, n-propyl mercaptan, and dimethyl sulfide (36b).Figure 4 shows the arrangement of sulfur atoms on a Ni(100) surface for p ( 2 x 2 ) and c(2 x 2 ) structures. For the c(2 x 2 ) structure, the distance between sulfur atoms and the plane through the center of the surface Ni atoms (Fig. 4c) is 0.13 & 0.01 nm (6,29-31,40). The distance of 0.1 33 nm calculated for small clusters of Ni atoms using ab initio calculations ( 4 ) is in remarkably good agreement with the distances obtained from LEED intensities measurements. The adsorption of sulfur on Ni(ll0) and Ni(ll1) is more complicated (26, 30, 33-36, 41, 42, 4 4 ) . For example, during initial stages of sulfur adsorption (42, 49) on clean Ni(ll1) at room temperature, a p ( 2 x 2) structure is observed at 8 < 0.25 (Fig. 5a). At slightly higher coverages (0.25 < 0 < 0.33), the structure changes to a f i x f i R 30” pattern, with each sulfur atom bonded to three Ni atoms in three-atom coordination sites (Fig. 5b). There is also evidence from a recent RHEED study (47) that onedimensional chains having linear disorder (and which therefore produce weak, diffuse LEED patterns) are segregated on a Ni(ll1) surface at 8 = 0.2-0.3 (see Fig. 6). Higher coverages result in essentially complete satura-

’

B is the ratio of the number of sulfur atoms per square centimeter to the number of surface metal atoms per square centimeter.

147

SULFUR POISONING OF METALS

0 25 MONOLAYER

. . . . . . \. . \. ,. 05

MONOLAYER \o

> C '

. ./ . ./ . . '. '. .' . ./ . . 0

0

3

. . . . . P ( 2 X 2 ) [Cb = 41

(a 1

0

\o

/*

I '

0

0

,/

c(2x 2 1

[Cb = 23

(b)

FIG.4. Schematic views of sulfur bonding on Ni (100) surface for (a) top view at 0.25 monolayer coverage p(Z x 2) with sulfur in a fourfold site; (b) top view at 0.50 monolayer coverage c(2 x 2) with sulfur in fourfold site; and (c) side view of 0.5 monolayer c(2 x 2) structure (Ref. 32).

FIG.5. Preliminary 'states of sulfur on Ni(l1 I ) at progressively higher coverages and room temperature: (a) p(2 x 2), (b) f i x flR 30",and (c) p(5 x 5). Reprinted from Ref. 27 by courtesy of Marcel Dekker, Inc.

148

C. H. BARTHOLOMEW ET AL.

FIG.6 . One-dimensionally disordered sulfur chains adsorbed on Ni(ll1) at B = 0.22. Solid circles represent adsorbed sulfur atoms; open circleb, nickel atoms in the adsorbent surface layer (Ref. 47).

tion of the surface; the last stage corresponds to a compact arrangement of sulfur atoms with a coincidence lattice in relation to the substrate atoms (surface Ni atoms) having a p ( 5 x 5), ( 5 J3 x 2), or 4 2 0 x 2) structure ( 4 2 , 4 7 , 4 9 )(see Figs. 5c, 7, and 8). In summary, sulfur atoms adsorb in three stages, but only the first two involve binding of all the sulfur atoms in threecoordinate Ni sites (42). The phenomenon of surface reconstruction or faceting is important in catalysis, particularly in regard to surface-structure-sensitive reactions. The question of whether sulfur adsorption induces reconstruction of nickel surfaces is interesting but also controversial, due in part to the complications in assigning surface structures to LEED patterns. Perdereau and Oudar (26) originally interpreted their data for high-temperature adsorption of H,S on Ni(ll1) in terms of a reconstruction leading to a true, bidimensional, ionic

FIG.7. Coincidence ( 5 43 x 2) structure for sulfur adsorbed on Ni(lI1) at 0 = 0.40. Solid circles represent adsorbed sulfur atoms; open circles, nickcl atoms in the adsorbent surface layer (Ref. 47).

149

SULFUR POISONING O F METALS

FIG. 8. Coincidence 420 :< 2) structure for sulfur adsorbed on Ni(ll1) at 0 = 0.50. Solid circles represent adsorbed sulfur atoms, open circles, nickel atoms in the adsorbent surface layer (Ref. 49).

Ni,S layer. However, diffraction data from two later studies (36, 49) and work function measurements (48) caused them to revise this hypothesis (48). McCarroll et al. (35--36) studied the structures formed during sulfur adsorption on a Ni(ll1) plane at temperatures up to 700 K. Adsorption at room temperature produced successively p ( 2 x 2), x R 20", and hexagonal patterns, in agreement with the observations by Perdereau and co-workers (26, 42). At higher temperatures (700 K), the LEED pattern observed was interpreted to be a distorted 4 2 x 2) structure resulting from reorientation of the Ni( 1 1 1) face to a ( 1 00) surface layer on which adsorbed sulfur had formed a c(2 x 2)-unit mesh. It was suggested that the sulfur atoms are adsorbed on the top of the reoriented Ni( 100)layer (Fig. 9). Furthermore. only surface Ni atoms were thought to undergo reorientation, the bulk maintaining its (1 11) geometry. Presumably, a strong Ni-S interaction

3 3

FIG.9. The Ni(100) 4 2 x 2) surface structure which can exist on Ni(ll1) (Ref. 36n): (0)Ni atom position; ( 0 )S atom adsorbed on top.

150

C. H . BARTHOLOMEW ET AL

weakens the bonding between the surface and the next-lower Ni layers, permitting rearrangements of the surface Ni atoms. Indeed, the weakening of Ni-Ni bonds in cluster complexes by sulfur ligands has been shown to occur (37). Nevertheless, Erley and Wagner (49) discarded the hypothesis of reconstruction, interpreting their data in terms of a p ( 2 x 20) coincidence overlayer of sulfur atoms adsorbed on unperturbed Ni( 1 11). This point of view appears to be consistent with the work of Delescluse and Masson (47). However, in a very recent study Masson and co-workers (48b) obtained evidence that certain crystalline planes of nickel are unstable when covered with sulfur to near saturation. Specifically the (810) plane was observed to decompose into (410) and (100) facets ; moreover, the phenomenon was reversible, i.e., as the adsorbed sulfur was desorbed the surface returned to the (8 10) structure. It has been suggested on the basis of theoretical calculations that a reconstruction of both electronic states and the geometry of the surface layer may result from adatoms interactions (51); a change in the structure of surface metal atoms may also cause a shift in adatom equilibrium positions. Work function measurements on sulfur-covered surfaces (48)provide evidence of a weak electronic transfer from the metal to the adsorbed sulfur atom, although the dominant character of the bond is covalent. Thus, it should not be surprising to observe “electronic effects” as well as “geometric effects” in catalytic properties arising from sulfur poisoning. Regardless of the nature of geometric or electronic changes in the surface metal layer, it is clear that more significant modifications occur in the structure of the sulfur layer with increasing coverage. The change from fourto twofold coordinated sulfur bonding as a function of increasing sulfur coverage on a Ni( 100) surface implies that the adatom-adatom interactions are strong enough that subsequent adsorption requires modification of the bonding of preadsorbed atoms (32).The fact that the sulfur must modify its bonding to the nickel to allow further sulfur adsorption also implies that with progressive coverage of the surface, strong restraints are put on further adsorption of any sort. This is indeed borne out by data presented in Section IV. Since poisoning by sulfur usually involves saturation coverage, the surface structures under these conditions are of great interest. Apparently the sulfursaturated surface of all three low-index faces of Ni contains about one sulfur atom for every two surface Ni atoms and has a somewhat distorted c(2 x 2) overlayer structure. Thus it may be concluded that in general the surface structures at saturation coverage of nickel by sulfur are similar for different single-crystal surfaces ; moreover, the Ni-S bond lengths found on all surfaces are smaller than those occurring in stable Ni-S bulk compounds

SULFUR POISONING OF METALS

151

( 4 , 5 , 3 1 ) . This suggests that adsorbed sulfur is bonded to Ni more strongly than sulfur in a bulk Ni sulfide. Indeed, both experimental and theoretical evidence provided in Section II1,D support this conclusion.

2 . Other Metals Sulfur adsorption studies of single-crystal faces of other metals and in particular of Ag (52-57), Cu (58-65), Fe (66-73), Mo (74-76), Ru (77-79), Pd (80), Pt (81-90), and Pb (91) indicate surface metal sulfides as having structures qualitatively similar to those of Ni. In all the metals studied, there are two steps involved in sulfur adsorption: (1) adsorption on highcoordination metal sites, and (2) formation of a 2-D surface sulfide. For example, on the Pt( 100) surface three main structures have been reported (K5,86):p ( 5 x 20), p(2 x 2), and c(2 x 2 ) in the order of increasing sulfur coverage. The same LEED patterns are observed at room temperature and at temperatures to 573 K. The c(2 x 2) structure corresponds to a saturation coverage for which the ratio of adsorbed sulfur atoms to surface Pt atoms ( 6 ) is 0.5. The density of sulfur atoms per square centimeter of Pt surface is 6.5 x 1014 (85). Increased exposure of the Pt(l00) surface to elemental sulfur at room temperature leads to further adsorption of sulfur, but this adsorption is reported to be nondissociative since the sticking coefficient of S, does not decrease appreciably, and also because heating the surface to elevated temperatures (573 K), results in the desorption of S, (85). Heegemann et al. (85) speculated that two distinct adsorption states occur Pt(100) : dissociative adsorption of S, (a states) and nondissociative adsorption ( b states). Similar observations were made by Contractor and La1 (90) for H,S and SO, adsorption on a polycrystalline Pt surface at 353 K. They (90) suggested that sulfur atoms are either adsorbed with one sulfur atom strongly bonded to two Pt atoms or with one sulfur atom weakly bonded to only one Pt atom ; however, no direct evidence for these two kinds of bonding was provided. Seidl and Bechtold (82) also reported that two states of sulfur adsorption on Pt foil exist at temperatures between 543 and 653 K and at S, pressures between 5 x lo-’ and 7.6 x l o p 6 Torr. For sulfur adsorption on a Pt( 110) surface, Bonze1 and Ku (830) reported four structures occurring with successively increasing coverage : c(2 x 6), p ( 2 x 3), p(4 x 3), and c(2 x 4). Prior to the appearance of each of these four structures, they reported transition structures characterized by illdefined spot positions in the LEED patterns. Similar observations were reported by Berthier et al. (84, 89); they noted that at saturation coverage c(2 x 4) and p(4 x 4) structures coexist. The structures observed during initial stages of sulfur adsorption were identical, however. During sulfur

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adsorption on Pt(l1 l), four stages were observed in the formation of a saturated 2-D surface sulfide: p(l x l), y(2 x 2), ( $ x $) R 30", and a hexagonal compact structure (84, 85). Figure 10 shows those four stages in the formation of a 2-D sulfide. At saturation coverage, the sulfur conceng S/cm2, in good agreement with the values tration is 42 i 4 x reported for several faces of Ni (Table IV). Restructuring of the metal surface by adsorbed sulfur appears to be a general phenomenon (58,63,65,70,71,7#, 81,92,93).One observation is of particular interest here. It was observed that the Pt(ll1) surface reorients to the (100) plane in the presence of H,S (81). McCarroll(94) reported, on the other hand, that when Ca' or Na' ions were added to a clean Pt(l00) surface, it caused a reorientation of surface Pt atoms to a (1 11) plane. He suggested that the switching between (100) and (1 11) planes may be the result of the cationic/anionic role played by Na' or Ca+/S, respectively, and that this might possibly explain the role of promoters in catalysis. The reorientation of the metal surface is not unique to sulfur, sodium, or calcium only. It has been observed, for example, that the Ni( 1 1 1) surface reorients to the (100) surface in the presence of ethylene or benzene (35). Somorjai (93) has I000 I

0

I

I

10

20

-*

I

I

1

30

40

50

s(10g/cm2)

FIG.10. Surface structures of sulfur on Pt(1 I I) as a function of coverage and temperature (Ref. 84).

SULFUR POISONING OF METALS

153

suggested that the presence of small amounts of impurities may either promote or retard phase transformations that are accompanied by marked changes of atomic structure : “If the adsorbed impurity changes the surface free energy of the various crystal planes by different amounts, it can induce the rearrangement of the surface structure to form crystal planes that have lower surface free energy in the presence of the adsorbed impurity than the cyrstal planes that bound the clean solid.” Thus, two effects, induced by sulfur adsorption, may be observed: (1) inaccessibility of active surface sites due to geometric blockage by sulfur, and (2) changes in catalytic activity due to changes in the structure of the catalytic surface. The second effect is generally attributed to “structure-sensitive’’ reactions (95). Accordingly, in the case of a reaction network involving both “structure-sensitive’’ and “structure-insensitive’’ reactions, significant changes in the selectivity may be expected upon sulfur adsorption.

B. ADSORPTION MECHANISMS AND KINETICS The interaction of H,S, organic sulfides, and other sulfur compounds may involve a number of consecutive steps including reversible molecular adsorption of the sulfur compound, its dissociation, reorientation or reconstruction of the metal surface, formation of a 2-D surface sulfide, and at still higher H2S/H2 ratios, formation of a three-dimensional (3-D) (bulk) metal sulfide. Kinetic information about these processes may generally be helpful in elucidating the adsorption mechanism. Unfortunately, such quantitative kinetic information is not adequately available, with one exception, formation of bulk sulfides (9, 96). It has been shown from corrosion studies (97, 98) that bulk sulfide formation involves metal cation diffusion through the sulfide layer to the surface with the formation of a new metal sulfide layer on the outer surface beyond the original metal. Apparently the formation of multilayer sulfides occurs slowly at room temperature and at PHZs= 1 atm, cation diffusion through the sulfide layer controlling the rate (9, 96). Presumably this step would also be rate limiting at higher temperatures and at H,S concentrations as low as 10-100 ppm. However, in most catalytic processes H2S concentrations are below those needed for bulk sulfide formation. Only a few kinetic studies of the rate of sulfur adsorption on metals have been made. They reveal that rates of adsorption of H2S on metals are generally very rapid, the high sticking probability suggesting no barrier to adsorption and dissociation until saturation is approached. In the case of Pt and Cu (83, 92), two adsorption regimes are observed: (1) at 8 < 0.25-0.3, the adsorption of sulfur occurs with a high sticking coefficient ( z 1 .O); and

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(2) at 0 > 0.3, sulfur adsorbs slowly due to a decreased sticking coefficient (<0.03). The sticking coefficient for H,S on a Ni surface containing 0.35 of a monolayer of sulfur (ds = 0.35) during reaction of H, with CO at 661 K was estimated to be 5 x (99,100). Under similar reaction conditions but at lower sulfur coverages the sticking coefficient for H,S would be logically higher, suggesting that H2S adsorbs rapidly and quantitatively on Ni catalysts during methanation. In similar sulfur-poisoning studies of CO hydrogenation catalyzed by Co, Fe, and Ru surfaces at 663 K , Agrawal(101) reported H,S adsorption rates similar to those obtained for Ni (99). Sulfur adsorption on previously carbon-deactivated Co and Fe catalysts for which the bulk was carburized, and the surface was covered with multilayer graphitic carbon deposits also exhibited a large sticking coefficient (101). Thus, the presence of surface carbon apparently did not inhibit sulfur adsorption in the same manner as preadsorbed sulfur. The initial sticking coefficient of H,S adsorption on Fe( 100) at room temperature was reported to be 0.2, and decreased in accordance with (1 - OjU,) where 0, represents the saturation coverage (73). Similar behavior was reported at higher temperatures; e.g., the initial sticking coefficient for H,S adsorption on Fe(100) at 873 K was estimated to be about 0.7 (73). A number of workers (96, 102-115) have examined the nature of equilibrium adsorption of inorganic and organic sulfur compounds on nickel foils and films and supported catalysts employing a variety of techniques, e.g., infrared spectroscopy, volumetric, gravimetric, and magnetic measurements. Most of the previous work (96, 102-115) suggests that H,S chemisorbs dissociatively on Ni surfaces even below room temperature, although there is disagreement regarding the number of surface Ni atoms involved per sulfur atom. Saleh et al. (96) suggested a three-site mechanism for H,S chemisorption in the temperature range of 193-373 K,

-

H H,Si,,

+ -Ni-Ni-Ni

-+

S

H

I I. -Ni-Ni-Ni

I

(3)

whereas Den Besten and Selwood (104) inferred from magnetic measurements at 273-393 K that H,S forms four bonds upon adsorption on the Ni surface: H

H,Si,, + -Ni-Ni-Ni-Ni

-+

S

I / \ -Ni-Ni-Ni-Ni

H

I

(4)

In a more recent magnetic study of Ni/SiO,, Ng and Martin (110, 111) also suggested that at room temperature, H,S adsorption involves four surface Ni atoms per H,S adsorbed; at a PHZsof 1 atm deeper layers of Ni are subsequently, slowly attacked.

SULFUR POISONING OF METALS

155

Rostrup-Nielsen (106a) suggested a one-site mechanism for H,S adsorption at higher temperatures (823-91 8 K), H,S,,,

+ Ni

+

Ni-S

+ H,,,,

(5)

based on the value of 1.1 obtained for the exponent n in a Langmuir fit to adsorption isotherms over a broad rangeofsulfur coverage (0.13 < OS c 0.5):

From desorption data at 725 K and high sulfur coverages, (0 > 0.5), Oliphant et al. (112) obtained Langmuir exponents of 2.9 and 2.7 for Ni powder and 3% Ni/AI,O, , respectively, suggesting a three-site mechanism. The data of Oliphant et al. ( 112) are considered to be more reliable than those of Rostrup-Nielsen (106a), since their desorption isotherms, each determined from a single sample, evidence considerably less scatter than RostrupNielsen’s adsorption isotherm, each point of which corresponds to a different sample. Accordingly, the three-site mechanism is favored, at least at high temperature (675-875 K). However, the limitations involved in the application of the Langmuir isotherm should be recognized. For example, the adsorption of H,S may involve two different kinds of sites for S and H, respectively. In a recent study Rostrup-Nielsen et a/. (106b) showed that their hightemperature data for H,S adsorption on supported nickel, the data of McCarty and Wise (114, 115) for Ni/Al,O,, and the data ofoliphant et a/. for Ni/AI,O, were all very well fitted by a Temlkin-type expression of the form: PH2sIPH2 = exp[AH;(l - d ) R T - ASo/R]

(6b)

where AH: = -289 kJ/mol, ASo = - 19 J/K, and a = 0.69. As in classical Temkin theory, this model predicts a linear increase in AHo (decrease in heat of adsorption) with increasing coverage, but differs from the classical theory in that the entropy is independent of coverage. The decrease in the heat of adsorption with increasing coverage suggests the importance of surface heterogeneity in sulfur adsorption on supported polycrystalline nickel. The constant entropy suggests (106h) the possibility of subsurface sulfur adsorption (in addition to the adsorbed surface monolayer) consistent with recent observations by Weeks and Plummer of a subsurface species formed in addition to the c(2 x 2) layer on a Ni(100) surface (50a). SH surface species have been suggested (96, 104, 109) as intermediates in the dissociation of H,S on metals. Such a possibility is supported by the observation that at increasing sulfur coverages, the dissociated hydrogen is gradually desorbed (96, 104). Exchange experiments involving deuterium provide clear evidence for the existence of two types of hydrogen atoms on

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the surface (96). Saleh et al. (96) suggest the possibility that hydrogen atoms close to (but not necessarily bonded to) sulfur atoms are slowly exchanged, whereas hydrogen atoms freed from the interactions with sulfur adatoms are readily exchanged by deuterium. No direct evidence is available for the presence of SH surface species on Ni; however, recent data from a secondary ion mass spectrometry (SIMS) study (116) in which HS', NiH,S+, and Ni,HS+ species were detected at 120 K but not at room temperature suggests that such species may be present on the Ni surface at low temperatures, but not at room temperature where complete dissociation may occur. Recent spectroscopic evidence (78) indicates such an SH species to be present on Ru during H,S adsorption. Adsorptions of CS, (102, 117), SO, (118), (CH3)$ (104, 119, 120), (C,H,),S (105), n-propyl mercaptan (120, 121), thiophene (120), and isopropyl, n-butyl, isobutyl, and tert-butyl mercaptans (121) on Ni have also been investigated. It is speculated that CS, adsorbs dissociatively on Ni surfaces at room temperature (117). The interaction of CS, with Ni is limited to the surface at 193 K, but above 298 K bulk sulfidation is observed (102).Sulfur dioxide chemisorbs rapidly and irreversibly on Ni at 193 K, but extensive incorporation into the bulk is not observed below 373 K. Methyl mercaptan is adsorbed dissociatively on Ni at room temperature (103, 108) accompanied by evolution of hydrogen, methane, and dimethyl sulfide (103).In contrast, dimethyl sulfide is associatively adsorbed at 298 K, but at higher temperatures (> 500 K) rapid dissociation is observed accompanied by the evolution of CH,, C,H,, and H, (104). Moreover, X-ray photon spectroscopy (XPS) results for (CH,),S adsorption at room temperature indicate the presence of a CH3S-type adspecies (119). Mercaptans involving longer chain alkyl groups (108, 120-122) are generally believed to adsorb at room temperature as mercaptide structures through dissociation of the H-S bond; under these conditions the C-C and C-S bonds remain intact. Heating to higher temperatures (> 350 K), however, decomposes the mercaptide with formation of the corresponding olefin and a sulfided surface (121). Evidence for mercaptide structures is obtained from infrared observation of the C-S stretching band (121, 122). Other catalytic metals, e g , Fe (69,117, 118, 123-125), Pt (126),Pd (127), W (126), and Cu (128),adsorb sulfur compounds dissociatively in much the same way asNi. Lead, however, adsorbsH,S (129),SO, (118),and CS, (117) molecularly and almost reversibly below 373 K. No incorporation into the bulk is observed up to 573 K. In the case of H,S adsorption on Pt, less than monolayer sulfur coverage occurs at 193 K and the process is partially reversible (126). Adsorption of CS, and SO, on Pt is weak and completely reversible between 193 and 523 K. No sulfur incorporation into the Pt bulk was observed at temperatures to 523 K (126). Auger electron spectroscopy

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157

(AES) studies confirm that sulfur is not incorporated into the bulk of Pt in 0.54 kPa of H,S at temperatures to 550 K or in the presence of 0.54 kPa of SO, at temperatures to 625 K (130).This is in contrast to the observation by Tsai et al. (131) of sulfur incorporation into the Pt bulk at 473 K in the presence of 50 ppm SO, during reduction of NO by NH, . The presence of a reaction environment in the later case (131) is apparently responsible for the markedly different behavior of Pt with respect to SO,. Most of the discussion up to this point has centered on the adsorption of gaseous sulfur compounds on metals, although several of the previously mentioned studies involved segregation of sulfur from the bulk metals at high temperature (43, 46, 53). The development of segregation models and segregation isotherms which relate the equilibrium distribution of impurity on the surface and in the bulk are less advanced than for adsorption. The quantitative determination of segregation isotherms requires knowledge of surface energies, which are affected by the presence of adsorbed species such as sulfur (48a, 53). Nevertheless, since the changes in surface free energy due to adsorbed sulfur are apparently relatively small (53), the determination of segregation isotherms is apparently feasible, as illustrated by data for the S-Ag( 110) system (53). A few studies have been made of the adsorption of sulfur on metal oxides (100, 101, 112, 232-140). These can be classified as (1) nonreducible oxides which are used typically as catalyst supports and (2) reducible oxides which may be present originally as an oxide catalyst, or which may form during the reaction. Sulfiding of metal oxides by H,S is an interesting, though relatively undeveloped area of research. A number of metals owe their corrosion resistance to protective oxide films ;in addition, many catalytic reactions may occur on a passivated surface oxide layer with reduced metal lying underneath. Adsorption of sulfur on A1,0, (100, 101, 112, 132-134), on SiO, (100, 135), on ZrO, (ZOO), and on Al,O,-supported Mo and Co-Mo (23, 113, 136) has been studied; these results will be considered only briefly. Adsorption of H,S on irreducible oxide supports at high temperatures is typically small in comparison to that on metals. For example, at 725 K the quantity of H,S adsorbed on Al,03 is about 4% of that adsorbed on 3% Ni/Al,O, (112). Only negligible amounts of H,S were found to adsorb on Al,O, (100, l01), SiO, (ZOO), and ZrO, (100) at 653-663 K in the H,S concentration range of 13-100 ppb. However, at lower temperatures (e.g., 500-625 K), the amount of H,S adsorbed on Al,O, was found to be quite significant compared to that on the metal (140). Studies of sulfur adsorption on single-crystal planes of the oxides of Ni (137),Co (138), and Pb (91, 139) provide evidence that sulfidation does not occur by direct substitution of sulfur for oxygen since the oxide and sulfide structures may be different. Indeed, it has been reported (100, 101) that

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C. H. BARTHOLOMEW ET AL.

sulfur contamination of metal films (Ni, Co, Fe, and Ru) during transfer through the ambient can be minimized by payivating the surface of these films by a thin oxide layer. The oxide layer appears to play a role in the nucleation of surface sulfide. It has been suggested, for example, that during sulfidation of COO (loo), the reaction proceeds via (1) formation of an epitaxial metallic film on the oxide, and (2) sulfur adsorption on the metallic surface (138). However, in sulfidation of NiO (137) and PbO single-crystal planes (Yl), the possibility of oxygen and sulfur atoms adsorbed together on the same surface layer appears more likely. Thus, a clear distinction can be made between reducible oxides and irreducible oxides. Sulfur adsorption on reducible oxides (such as NiO, COO, and PbO) is significant and appears to involve a removal of surface oxygen somewhere in the intermediate stage of adsorption, whereas because of their irreducibility, typical catalyst supports (such as A1,03, SO,, and ZrO,) adsorb relatively little sulfur. It is also clear that our knowledge of the sulfidation of metal oxides is sparse; furthermore, a more detailed understanding of metal sulfidation will probably be necessary before significant progress is made in understanding the sulfidation of metal oxides. To summarize, previously published information regarding the adsorption of sulfur-containing compounds on metal foils and films and supported metals provides a qualitative and useful understanding of the mechanism of sulfur adsorption. Some of the most important observations are:

(1) irrespective of the form of the metal (foil, film, or supported metal), sulfur compounds chemisorb dissociatively on the metal surface in the range of typical reaction temperatures (473- 673 K), to form surface sulfides; ( 2 ) the adsorption process involves several metal sites for every molecule of sulfur compound adsorbed: and (3) the rate of sulfur adsorption is rapid, having a sticking coefficient approaching 1.O at low sulfur coverage. Future studies should emphasize the further application of spectroscopic techniques, AES, LEED, electron energy loss spectroscopy (EELS), and laser Raman spectroscopy to provide a more quantitative understanding of the interaction of sulfur compounds with metal surfaces. C. ADSORPTION STOICHIOMETRIES 1. Nickel

a. Single-Crystul Nickel. Values of 0 (i.e., the number of sulfur atoms adsorbed per surface metal atom) obtained for various crystal planes of Ni at saturation coverage by sulfur are listed in the last column of Table IV (26).

159

SULFUR POISONING OF METALS

For Ni(l1 l), Ni(100), Ni(1 lo), and Ni(210) planes, 0 ranges from 0.48 to 1.09. Thus it appears that the value of H at saturation coverage increases with the increasing openess of the crystal plane, whereas the sulfur concentration (per square centimeter) is essentially independent of crystal plane (Table IV). In other words, these variations in the maximum value of 6, arise from the differences in Ni atom concentration on the various surface planes. It should be emphasized, however, that in those cases where reconstruction occurs, the 2-D nickel sulfide or sulfur overlayer thus formed may have the same sulfur-to-nickel atom ratio independent of the starting plane. Therefore, a better way to evaluate 0 would be to consider the structure of the surface Ni layer at saturation coverage rather than the structure of the clean Ni surface before adsorption of any sulfur. If all surfaces of Ni are assumed to reorient to a (100) configuration at saturation coverage, then the values of 0 in Table IV for Ni(l1 I), (loo), (1 lo), and (210), and polycrystalline surfaces would be 0.54, 0.50, 0.51, 0.49, and 0.51, respectively. Thus, a value of Om,, of 0.5 appears to be universal. This method of calculating 0 for various planes on the basis of the density of the (100) plane is reasonable; it is supported, in fact, by spectroscopic observations from several recent studies (9-11, 41, 116). For example, Bordoli et al. (116) observed the same SIMS intensity ratios, N,Sf/Ni,+, for H,S adsorption on two different crystallographic planes of nickel, namely Ni( 100) and Ni(ll1). Nguyen and Cinti (41) found by AES study of Ni(100), (llO), and (111) planes that the peak ratio S(152 eV)/Ni (62 eV) was 0.95 on all three planes at saturation coverage, independent of the original plane. Colby ( 9 ) and Windawi and Katzer (10, 11) also reported the peak ratio S (152 eV)/Ni (62 eV) to be 1.0 on polycrystalline Ni films at saturation coverage. Since SIMS intensities for Ni,+ and Auger yields for Ni (62 eV) are directly proportional to the number of surface Ni atoms, the concept of surface reorientation to the same 2-D sulfide regardless of starting plane is supported by these observations. TABLE 1V Sulfur Adsorprion Densities on Vurious Crystul Fuccs of Nickel"

Crystal fdCe

"

Data from Ref. 26.

Sulfur conc. at saturation (g S/cmZ)

Number of S atoms/cm2 ( X loi5)

Number of N i atoms/cmz (x

101~)

S atoms per surface Ni atoms

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C . H . BARTHOLOMEW ET AL.

b. Polycrystalline Ni. Although sulfur adsorption stoichiometries at saturation on metal single-crystal planes are well defined, they are more difficult to measure and are less well defined in polycrystalline and supportedmetal systems. Indeed, there is apparently only fair agreement as to the stoichiometry of H,S adsorption on polycrystalline and supported nickel. The reasons for this are primarily twofold: (1) The saturation stoichiometry apparently depends upon PHZsabove about 0.1 ppm (112, 113, 140); and ( 2 ) it varies with temperature (140). Based upon adsorption studies at high PHZs (10 ppm to 1 atm) on polycrystalline or supported Ni, reported S/Ni, values2 range from 0.25 to 0.33 at room temperature (96, 104, 111) to as high as 0.75-1.0 at 725-875 K (106, 112, 113). On the other hand, Fitzharris et al. (99, 100) observed a S/Ni, value of 0.5 on poorly dispersed Ni/Al,O, at 673 K and 100 ppb H,S/H2. A priori, it is reasonable to expect that at most about 0.5-0.6 atoms of sulfur chemisorb per nickel atom on a clean nickel surface in view of the relative areas of about 0.1 1 and 0.060 (nm)’/atom for S and Ni atoms, respectively. These apparent discrepancies can be resolved as follows. First, the values at lower temperatures (S/Ni, = 0.25 and 0.33) (96, 104, I l l ) are smaller because hydrogen atoms from the dissociative chemisorption of H,S remain adsorbed on the nickel surface blocking sites for further sulfur adsorption. At higher temperatures hydrogen desorbs allowing sulfur atoms to cover most or all of the nickel sites, and thus higher S/Ni, ratios (e.g., 0.6-0.7) are observed. The agreement of the S/Ni, ratio of 0.5 reported by Fitzharris et al. (99, 100) with the values obtained for single-crystal planes is expected because: (1) the measurements were carried out at very low H,S concentration (< 100 ppb in H2), (2) the nickel surfaces of the large crystallites were composed of predominately low-index planes, and (3) the measurements were carried out at sufficiently high temperature (673 K) that H, adsorption was reversible. The near-unity values of S/Ni, observed by Oliphant et al. ( I 12) for both well-dispersed supported and for unsupported Ni obtained by desorption after saturation at 25-30 ppm H,S/H, could be explained by: (1) surface reconstruction at higher H,S concentrations (10-25 ppm) to a Ni surface sulfide layer having a stoichiometry of NiS, (2) the presence of high-index Ni planes which would lead to higher S/Ni, values (Table IV), and (3) variation from 1.0 in the H/Ni, chemisorption stoichiometry. The facts that the Ni powder studied by Oliphant et al. (112) was poorly dispersed and that the separately determined H/Ni, ratio was The S/Ni, ratio is defined as the number of sulfur atoms adsorbed per nickel surface atom measured by H, chemisorption. If H/Ni, = 1, then S/Ni, and B are equivalent.

SULFUR POISONING OF METALS

161

1.0 (141) appear to rule out explanations (2) and (3). Thus surface reconstruction leading to new surface phases at higher H,S concentrations is favored. Evidence that S/Ni, at saturation coverage varies with H,S concentration was obtained in a study by Erekson and Bartholomew (140)of H,S poisoning of unsupported nickel during methanation at 523-673 K and PH2s/PH2 values of 0.2-30 ppm. At 523 K, S/Ni, values were found to increase from 0.5 to 1.3 as the H2S concentration was increased from 0.2 to 1.O ppm. These data were obtained at PHzs /PHI values lower than those estimated from bulk thermodynamic data (13) to be necessary for bulk sulfide formation. Thus these results suggest that S/Ni, values increase with increasing PHzs/PH2. Moreover, they suggest that Ni-S surfaces vary from well-defined structures (0 = 0.5) at low H2S concentrations to reconstructed surfaces of Ni,S, and NiS stoichiometry at intermediate conditions to multilayer sulfides at high H S concentrations. Although surface structures consistent with Ni,S, and NiS stoichiometries and multilayer surface sulfides have not been reported, all of the previous LEED and AES studies [with the exception of the study by Weeks and Plummer (50a)l were conducted at partial pressures of H,S significantly lower or temperatures significantly higher than in the previously cited studies of polycrystalline and supported Ni in which S/Ni, ratios of 0.7-1.0 were observed (106, 112, 113). Thus there is a clear need for spectroscopic characterization of single-crystal and polycrystalline surfaces during or following = 0.1-25 ppm, exposure to H,S under comparable conditions (PH2s/PH, total pressure of 1 atm, 575-800 K), especially since these are conditions most relevant to typical catalystic processes in which sulfur poisoning is a problem.

,

2. Other Metals Sulfur adsorption stoichiometries at saturation coverage for single-crystal surfaces of Pt (84,85), Fe (72, 142), Mo (75,143,144), Ag (53-56), and Cu (58,65),and for polycrystalline metal surfaces of Pt (14.9, Fe (IOZ),Co (ZOZ), and Ru (101) have been reported. The general features observed for these metals are similar to those observed for Ni; accordingly, only the more interesting observations will be discussed. Berthier et al. (84) studied sulfur adsorption on Pt(100), (1 lo), and (1 11) planes, using LEED, AES, and H,35S radioactive tracer techniques. The stable saturated 2-D sulfide phases on the (loo), (110), and (111) planes contained 6.6 x 1014,6.0 x 1014,and 5.5 x 1014 S atoms/cm2, respectively. Although their results suggest surface reconstruction, this question was not resolved in the discussion of results. If in a manner similar to that for Ni, we calculate the site density of surface Pt atoms for the Pt(100) plane, 1.3 x 10l5 Pt atoms/cmZ (146), and use it as the basis for calculating 6 for saturation

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C. H. BARTHOLOMEW ET AL.

coverage, we obtain values of 0.51, 0.46, and 0.42 for the (loo), (1 lo), and (1 1 1) planes, respectively. These values of 8 are in sufficiently close agreement to suggest that surfaces of Pt (like Ni) are reconstructed in the presence of adsorbed sulfur to form 2-D sulfides of almost the same structure irrespective of the starting phase. When Berthier et al. (84) exposed the saturated surface of Pt(ll1) to 0.1 Torr H,S at 473 K, more sulfur was adsorbed than at higher temperatures giving B,,, = 0.59 ; however, heating the surface to 623 K resulted in desorption of a portion of the adsorbed sulfur, and a stable coverage of 6, = 0.42 was obtained. These workers considered 8 = 0.42 to represent a saturated layer (stable at higher temperatures), and assumed that the amount of sulfur adsorbed in excess of 8 = 0.42 was due to a weakly adsorbed sulfur. Similar observations were reported by Heegemann et al. (85). They observed the value of 8 at saturation coverage to be 0.5 and 0.38 for the (100) and (1 11) planes, respectively, when S, was adsorbed on these planes. Saturated layers of Pt(ll1) and (100) adsorbed more sulfur at room temperature, giving values of 8 of 1.12 and 1.08, respectively. Evacuation at room temperature reduced the value of 8 to 0.92 for both surfaces, indicating some fraction of S, to be in a physisorbed state. When both these surfaces were heated to temperatures above 575 K, values of 0 = 0.5 and 0.38 were obtained for the (100) and (11 1) planes, respectively; further continued heating to 723 K resulted in no further reduction in the value of 8. The results reported by Berthier et al. (84)and Heegemann et al. (85)are in good agreement. Both groups indicate that more sulfur can be adsorbed at lower temperatures and higher gas-phase sulfur concentrations, although a portion of the sulfur is in a weaker bonding state. This is also supported by the results of Contractor and La1 (90) and Seidel and Bechtold (82). Bechtold (145)reported that the surface of polycrystalline Pt films was saturated with sulfur, and that between 673 and 723 K a value for 8 of 0.5 1 was stable over long periods of time. Thus it can be concluded that on Pt surfaces, at higher temperatures a saturated surface layer contains about one S atom per two surface Pt atoms [based on Pt(l00)l. The value of 6, at saturation coverage has also been reported to be 0.5 for Fe (72, 101, 142), Cu (58), Mo (75, 143, 144), Ru ( I U l ) , and Co (101). Cabane-Brouty (54-56)reported the sulfur adsorption density to be constant at a value of 7.1 x 1014 S atomslcm’ for saturated Ag(100), (1 lo), and (1 11) planes. Based on the density of the Ag(100) surface, a value of 0 = 0.59 is obtained for the saturation coverage of all three planes. Thus, based on the above observations, it is possible to generalize the stoichiometry of sulfur adsorption at saturation coverage for various metals. A value of 0 = 0.5 appears to have a general applicability in representing the saturation coverage of the metals of catalytic interest at high reaction

SULFUR POISONING OF METALS

163

temperatures (>673 K) and low H,S concentrations ( ~ 0 . 1ppm). At higher concentrations (>0.1 ppm) and lower temperatures (473-673 K), values of 8 = 0.5-1.0 are possible. Clearly, more work is needed in this area to define the structure of reoriented or reconstructed surface phases for metals other than Ni. Another question relevant to catalytic chemists is how much the sulfur adsorption stoichiometry is affected by the presence of other molecules in the gas phase or on the surface of the metal, i.e., how sulfur adsorbs under actual reaction conditions. Very little information is available in this area of great practical importance. Agrawal et al. (147) studied sulfur poisoning of Co/Al,O, in CO hydrogenation at 1 atm and 663 K using 13-100 ppb H,S. The catalyst was poisoned by one of the following three techniques: (1) preadsorption of H2S to achieve a saturated surface layer before introducing any CO, (2) introducing CO and H,S in H, at the same time so that coadsorption of CO and H,S was concurrent, and (3) introducing CO/H, first at reaction temperature to obtain a catalyst deactivated 100-fold by multilayer graphitic deposits on the surface and formation of bulk carbide. The value of 8 at saturation coverage by sulfur was 0.5 in all three cases, suggesting that the interaction between S and Co is too strong to be influenced by the presence of CO in the gas phase, or carbon on the surface, at least under these conditions. On the other hand, markedly different behavior was observed during reaction for H,S adsorption on Ni (140) and for SO, adsorption on Pt (81, 131). Erekson and Bartholomew (140) observed that during methanation of CO on polycrystalline Ni, the number of sulfur atoms required to deactivate one nickel atom decreased from 1.0 to 0.23 as the mol % of CO (H, diluent; PHzs= 0.2 ppm) was increased from 2 to 20%. During those experiments H,S and COS appeared in the exit stream when the CO concentration was 5% or above. At 20 mol % CO only 1/4 of the reactant H,S (and apparently none of the COS)adsorbed on the catalyst, suggesting that adsorbed CO or carbon inhibits H,S adsorption. No sulfur was incorporated into bulk Pt at temperatures to 523 K (126, 148) in the presence of SO, alone, but incorporation of sulfur into the Pt bulk at 473 in the presence of SO, during reduction of NO by NH, was reported (131). Similarly, Berthier et al. (84) reported no reorientation or reconstruction of the Pt( 1 1 1) surface when exposed to H,S at elevated temperatures, but Schmidt and Luss (81) observed reorientation of Pt(ll1) surface to Pt(100) by H,S in the presence of HCN. The different results discussed above for Co, Ni, and Pt illustrate the complexity of sulfur adsorption under reaction conditions and show that any generalizations or extrapolations from nonreacting to a reacting environment may be dangerous. Clearly, more work is needed to define how various

164

C. H. BARTHOLOMEW ET AL.

sulfur compounds adsorb in competition with other molecules present either in the gas phase or on the surface.

D. STABILITY OF SULFUR ADSORBED ON METALSURFACES The fact that metal catalysts are poisoned for all practical purposes irreversibly by sulfur compounds at concentrations well below those necessary for bulk metal sulfide formation suggests that 2-D surface sulfides are significantly more stable than bulk metal sulfides. In principle, the stability of surface sulfides (like those of bulk metal sulfides) can be expressed in terms of AG" and can be determined according to Eq. (2) from experimental values of PHls/PH2at equilibrium and at fixed values of temperature and 13.Since Eq. (2) involves a direct proportionality between AGO and absolute temperature, the stability of surface sulfides is quite temperature sensitive. This is illustrated by the large variations in AGO with temperature shown in Table V. The necessity of obtaining accurate values of AGO in order to estimate H,S levels at which the surface sulfide is stable can be illustrated by the following calculation. At 600 K, 51.8 ppb H,S in H, is needed to form a stable surface sulfide for AGO = - 84 kJ/mol, whereas only 0.78 ppb H,S in H2 is sufficient to form a stable surface sulfide for AGO = - 105 kJ/mol. In practice, thermodynamic data such as heats and free energies of formation of surface sulfides are very difficult to obtain, mainly because quantitative measurements must be made of adsorbed and gas-phase sulfur at extremely low concentrations. Only during the last decade has the developTABLE V The Rario P H I S / P H us2a Funcrion o j Temperuiure for Various Vulues qf Free Energy o j Formution" PHZSIPH>

Temperature (K) 300 400 500 600 700 800 900 I000 ~

~~

Based on Eq. (2).

AG" = -63 kJ/mol

AG'' = -84 kJ/mol

AC = - 105

1.18 x 6.36 x 2.77 x 3.43 x 2.07 x 7.98 x 2.27 x 5.26 x

2.68 x 1.18 x 1.81 x 5.18 x 5.69 x 3.43 x 1.39 x 4.25 x

6.1 x 2.18 x 1.17 x 7.8 x 1.56 x 1.48 x 8.48 x 3.43 x

lo-" lo-' lo-'

lo-6 10-5 lo-'

10-4 10-4

lo-" 10-9 lo-' 10-7 10-5

lo-'

'

kJ/mol lo-'' lo-'' lo-'' lo-* lo-' 10-7 10-6

AG" = - I26 kJ/mol

1.39 x 4.0 x 7.69 x 1.18 x 4.3 x 6.36 x 5.18 x 2.77 x

lo-"

lo-" 10-l' 10-~ lo-* lo-'

SULFUR POISONING OF METALS

165

ment of sophisticated analytical and spectroscopic techniques enabled such measurements to be carried out accurately and conveniently. Accordingly, previously reported values of AGO and AH for surface metal sulfides are relatively few. These data, nevertheless, confirm the extreme stability of adsorbed sulfur on metals.

1. Nickel A number of previous studies (29a, 96, 99-115, 149) provide evidence showing surface nickel-sulfur bonds to be substantially more stable than bulk nickel-surface bonds. For example, Goddard et al. (149) estimated the enthalpy of formation for sulfur adsorbed on Ni using ab initio calculations to be - 240 kJ/mol compared to values of - 85.4 kJ/mol for bulk NiS (150), and -75.3 kJ/mol for bulk Ni3S2(13). From LEED measurements, Demuth et al. (29a) found the length of the surface Ni-S bond (0.218 nm) in single crystals to be smaller than the Ni-S bond lengths in bulk sulfides (0.238 nm for NiS and 0.228 nm for Ni,S2), these data indicating a decrease in bond length and hence a more stable bond on the surface. More quantitative measures of the stability of adsorbed sulfur on nickel were obtained by several groups of workers from isotherms for H2S adsorption on Ni (26, 106b, 112-115). For example, in the most definitive investigations Bartholomew et al. (112, I 1 3 ) , McCarty et al. (114, 115), and Alstrup et al. (106b) obtained heats of H,S adsorption on polycrystalline Ni powders and supported Ni from desorption isotherms, adsorption isosteres, and adsorption isobars respectively. Remarkably good agreement is evident among their three sets of data listed in Table VI. Since according to these data the heats of adsorption for H,S on various forms of nickel may range from 130-160 kJ/mol (at 8 = 0.7-0.9), the enthalpy of adsorption is 55-85 kJ/mol more exothermic than the enthalpy of formation of Ni,S2 of 75 kJ/mol(l3). The remarkable stability of adsorbed sulfur is further demonstrated in Fig. 11 in which most of the previous equilibrium adsorption data for nickel (26, 106, 112-115) are represented in a single plot of log(PH,,/P~,) versus reciprocal temperature. The solid line corresponds to the equilibrium data reported by Rosenqvist (13) for formation of Ni3S, (in the temperature range 675-810 K). Based on the equation AGO

=

RTln(PH,,/PH,)

=

AH - T A S

(7)

the slope of this line is AH/R(2.303), where AH = - 75 kJ/mol, and the intercept is - ASlR(2.303). The dashed lines in Fig. 11 represent equilibrium lines for chemisorbed sulfur with heats of adsorption of 85, 125, and 165 kJ/mol, assuming formation ofNi,S2 (i.e., the same intercept as the solid line for Ni3S2).Thus, according to these data, the enthalpy of adsorption of

166

480 725--775

990- I 140 773-1023

c. H .

5% Ni AI,O, Polycrystal 14x NiiAl,O, 167" Ni/AI,O, Sponge 5% Ni AI,O, Ni( 100) Ni/Mg AI,O,

BARTHOLOMEW E'r A L

136" 134' 159' I50 144d I50* 164" I47",'

Ads. isostere Desorption

114, 115

Ads. isostcre

114, 115 114, I15 114, 115 106h

Ads. isobars

75 (Ni,S,)

I I2 I13 I12

Heat of adsorption per mole of H,S. Heat of bulk formation per mole of H,S (Ref. 13) ' 0 > 0.7. H 2 0.70. "AH = -289 ( I - 0.690). "

H2S on nickel is 50- 100 kJ/mol more exothermic than the enthalpy of formation of Ni,S2, depending upon temperature and coverage. It is also apparent that the absolute value of the heat increases with decreasing coverage and that the equilibrium partial pressure of H2S increases with increasing temperature. Considering the different experimental approaches [i.e., radioactive tracers (26),desorption isotherms (112,113) and adsorption isosteres (114, 115)] and obvious experimental problems and/or limitations /PH2 with increasing temperature (26) and high sulfur [e.g., a decreasing PHzs coverages (112, 113)],the data in Fig. 1 1 present a remarkably consistent picture. From equilibrium adsorption isotherms or AG' values at different temperatures, it is possible to estimate the extent to which H2S adsorption on Ni is reversible at various temperatures and concentrations. For example, the isotherms of Oliphant et ul. (112) in Fig. 12 reveal that in the temperature range of 725-775 K, less than saturation coverage (i.e., reversible adsorption) occurs on the Ni surface only at concentrations less than 1-2 ppm H,S. From the data in Fig. 11 it is also possible to estimate the equilibrium partial pressure of H2S at any given temperature for fractional coverages ranging from 0.5 to 0.8. For instance, at 725 K and 0 = 0.5, values of In other words, half coverage is PH2s/PHz range from about lo-* to obtained at 1-10 ppb H,S, a concentration range at the lower limit of our present analytical capability! At the same temperature (725 K), almost

167

SULFUR POISONING OF METALS

2

I

I

I

0

-2 c

N

'a \ v)

N

-4

CL=

u

P rn o -6

-8

-10 RECIPROCAL TEMPERATURE

X

lOOO(K-')

FIG. 1 1, Equilibrium partial pressure of H,S versus reciprocal temperature for the nickel sulfur system. (Values of A H f based on 1 mol of H,S.) Ref.

Symbol" Unsupported Ni 3% Ni/AI,O, 14% Ni/AI,03 16% Ni, 0.5% Pt/AI,O, 10% Ni/MgO-Al,O, Ni Foil 5% Ni/AI,O, Ni sponge Ni foil ~~~~~

~

~

Open symbols, 0 9 = 0.80-0.9. a

112 112 113 112 106 26 114 114 114

= 0.50-0.6;

closed symbols,

168

C. H. BARTHOLOMEW ET AL.

I .2

I

I

I 100% Ni'POWDER

=

3% N i / A I 2 O 3

1.0

E, 0.9

\ v) u)

J

,

I. I -

-

10% N I / 10% C0/A1203

,.

n

0.8

1

15%N1 / O 5 % P t / A 1 2 0 3

5 0.7

t

0

v

w 0.6 Y

2 0.5 3

a: 0.4 3

5 0.3 3

cn 0.2

i

O.I 00

4

8

12

16

20

24

28

32

ti, S CONCENTRATION ( p p m )

FIG.12. H,S desorption isotherms for nickel and ruthenium catalysts at 725 K normalized to H, uptake (Ref. 112).

complete coverage (0 = 0.9) obtains at values of P H 2 s / P H 2of 10-8-10-6 (0.01- 1 ppm) or, in other words, at H2Sconcentrations encountered in many catalytic processes after the gas has been processed to remove sulfur compounds. At lower temperatures, e.g., 500-600 K, there is no practical equilibrium concentration of H2Sthat will not result in irreversible poisoning at monolayer coverage by sulfur ! All of the data in Fig. 11 were obtained using H,S in pure H,. The stability of the nickel-sulfur bond may be affected by the presence of other gas-phase or surface-phase species during reaction (140). Nevertheless, it appears that in H,-rich reaction environments, its stability is about the same as in pure H,. For example, Fitzharris et al. (99, 100) observed during CO hydrogenation on Ni/AI,03 at 661 K (H,/CO = 25) that only 13 ppb H,S caused the formation of a saturated sulfide on Ni surface. From their data an upper limit for the free energy of formation of the surface nickel sulfide at 661 K, of - 100 kJ/mol was obtained in good agreement with - AG" values of 100-125 kJ/mol obtained in pure H, for the same range of temperature (106,112,113, 151) (see Table VII).

169

SULFUR POISONING OF METALS

TABLE V11 Free Energies a/ Formation oJ'Surface Sulfides Compared to Free EnergieJ of Formation of Bulk Metal SidJides Sample form

Metal Ni

Ag

On A1,0, Powder On AI,O, Foil On AI,O,/MgO Foil Several forms on Al,O, (100) (1 10)

66 1 725 725 823 823-918 973 I173 660-1400 600-1000 673

( - AG);;.h

( - AGY; (kJ/mol)

T (K)

Ref.

(kJ/mol)

> 100

99 54.4 (Ni,S,) 112 112 52.1 (Ni,S,) 117-125' 112, 113 52.1 (Ni,S,) 1100 151 49.0 (Ni,S,) 106 99.6 46.5 (Ni,S,) 74.9 26 44.2 (Ni,S,) 90.4 26 37.2 (Ni,S,) 107.2 - 0.012 T' 114, 115 30-60 154.2 - 0.034 Td 114, 115 35.2 57 5.9, 18.7 (Ag,S)

Ref. 13 13 13

13 13 13 13 13 19, 153

(111)

Cu

Foil

I I03

Fe

OnA1,0, Foil Powder OnA1,0, On A1,0, (0001) Two forms OnAI,O, Powder

663 1123 725 663 663

Ru

Co

"

70.3

60

> 100

1000

650-1400 663 702

101

62.0 (Cu,S) 2.5 (CuS) 56.9 (FeS) 55.2 (FeS)

99.2 154 100.7 + 0.011 Td 152 > 100 I01 28 (Ru,S) 114 114, 156 At 300 K 101 I14 127.9 - 0.023 Td 114, 156 > 100 101 52.7 (Co,S,) 110.2 - 0.013 P 152

19 13 13

19

13

Free energy of formation is computed using the following equations: (x/y)M

+ H,S,

* Free energy of formation

y-'M,S,

+ H,,

AG'

=

RTln(P,,,/P,,).

of bulk sulfide is at the same temperature at which results are

reported for surface sulfide. 0 = 0.9. d e = 0.7. 0 = 1.0.

2. Other Metals Only a limited amount of thermodynamic data are available for 2-D sulfides of metals other than nickel. Data reported for Ag (57), Co (152), Cu (253), Fe (152, 154), Mo (153, 1.55), Ru (156), and Pt (84, 85) indicate that the heats of sulfur adsorption are generally 20-40% larger than the heats of formation of the most stable bulk sulfides. Indeed, Benard et al. (153)have shown a linear correlation between the heats of adsorption of sulfur and the

I70

C. H. BARTHOLOMEW ET AL.

heats of formation for the bulk sulfides (see Fig. 13). From these data (Fig. 13) and recently obtained data from McCarty and Wise (115, 152, 156) it appears that the bond strength of sulfur on various metals decreases in approximately the following order: Cr > Ni > Mo > Co > Ru > Pt > Fe > Cu > Ag. From the data in Fig. 13 for Ag, Cu, and Pt it appears that the heat of sulfur adsorption increases with increasing atomic roughness of the surface. In the (3.2 x case of Ag, the saturation coverage is reached at the same PHZs/PH2 on all three low-index planes [(loo), ( I I O ) , and (ill)]. However,

c u (1OO)o O P t (1 1 1 1 60-

ii

3A g ( 1 1 0 ) 40Ag(l00)

2 02Ag(l11) -AH, ( $ M t H2S=$M, S, +H, 1

0

I

1

I

I

I

I

FIG. 13. Heats of H,S adsorption (0 = O.5)versus heats of bulk sulfide formation from H,S. Ni and Ru data from Ref. 114 ; other metals from Ref. 153. Linear least-squares fit : slope = 1.323, intercept, 31.77, r = 0.95. To convert data to S, basis. add A H , for H,, + 0.5 S,(,, --* H,S@,,where AHr = - (82.4 + O.O089T), T = 300-700 K .

SULFUR POISONING OF METALS

171

subsaturation coverages vary considerably for the three planes at any fixed value of PH2s/PH2 (see Fig. 14). For example, at 673 K and PHrS/PH2 = lop3, values of 6' are approximately 0.3,0.6, and 0.1 for the (loo), (1 lo), and (1 11) planes, respectively. Moreover, the magnitude of the heats of adsorption for these three faces appears to correlate with these coverages at subsaturation (see Fig. 13 and 15). These are very interesting findings. Unfortunately, similar detailed attempts to study sulfur adsorption on various single-crystal faces on other metals have not been reported. Table VII compares the free energies of formation of surface sulfides on various metal surfaces with those of corresponding bulk sulfides. With the exception of Ag and Cu [which are less sensitive to sulfur adsorption due to their lower (-AG)s], it is clear that the free energies of formation of surface sulfides are at least 40 kJ/mol more stable than the corresponding bulk metal sulfides. Hence it is generally true for all metals that the surface sulfides are considerably more stable than the bulk sulfides. As demonstrated in the case of H,S on Ni, thermodynamic data such as AGO and AHadare of great value in understanding the extent and reversibility of sulfur poisioning. Clearly, more thermodynamic data for adsorption of sulfur are needed for important metals of catalytic interest such as Co, Fe, Mo, Pd, Pt, Re, Rh, Ru, etc. Since sulfur adsorption on Co, Pd, Rh, and Ru

FIG.14. Adsorption isotherms for H,S on (IOO), ( I lo), and ( I 11) planes of Ag at 675 K (Ref. 57).

172

C. H. BARTHOLOMEW ET AL.

175

I50

-g

h

125

\ 3 1 Y

'0

r"

a

100

75

0

25 50 75 I00 DEGREE OF COVERAGE (Yo)

FIG. ats of adsorption of sulfur on (IOO), ( I lo), and (111) planes of Ag at 67 (Ref. 57). To determine heat of adsorption of H,S subtract AH, for H,,,, + 0.5Sz,,) = H,S(,, for which AH = -(82.4 + O.O089T), T = 300-700 K.

has received little attention, the investigation of these important catalytic metals should be a high priority.

IV.

Effects of Sulfur on Adsorption of Other Molecules

Since one of the necessary steps in a heterogeneous reaction is the adsorption of one or more of the reactants, investigation of the effects of adsorbed sulfur on the adsorption of other molecules can reveal a great deal about the poisoning process. Sulfur adsorption on a metal surface may alter its adsorption characteristics for various reactant molecules, either by blocking the active surface sites and thus making them inaccessible to the adsorbing molecules (geometric effects) or by structural changes caused by strong metal-sulfur interactions (electronic effects). It is also possible that sulfur will have direct chemical interactions with the adsorbed molecules instead of having indirect interactions through metal atoms. The effects of sulfur poisoning on CO and H2 adsorption on metals are of particular interest, not only because these molecules participate in numerous reactions,

173

SULFUR POISONING OF METALS

but also because they are used as selective titrants to measure metal surface areas. A. NICKELCATALYSTS 1. Efsects of’ Preadsorbed Sulfur on H, Adsorption Results of previous investigations (23, 110, 111, 113, 141, 157-165) show that hydrogen adsorption on nickel at room temperature is lowered by preadsorbed sulfur. Moreover, the fraction by which hydrogen adsorption is reduced in polycrystalline and supported nickel catalysts is generally proportional to the mean fractional coverage of sulfur. This is illustrated by data in Fig. 16 from Bartholomew and co-workers (112, 113, 141, 157-162). This plot of fractional H, adsorption (H, uptake at 300 K of catalysts presulfided in 5, 10, or 25 ppm H,S at 725 K divided by initial H, uptake) versus mean sulfur coverage (in molecules H,S adsorbed per molecule of surface nickel), suggests a linear relationship between H, uptake and sulfur coverage. Interestingly, the intercept at zero H, coverage (saturation sulfur coverage) is H,S/Ni, = 0.75, in excellent agreement with the adsorption stoichiometry reported by Oliphant et al. (112) for adsorption of H,S at 725 K.

0

0.5

I .o

H2S / Nis

FIG.16. Fractional H, uptake versus H,S coverage in molecules of H,S per nickel surface atom: 0,Ni powder (INCO) (Ref. 161); 0, 14-16% Ni/A1,0, Washcoat/Monolith (Ref. 159); A, 3, 8, 14, and 23% Ni/A1,0, (Ref. 160); 0, 14%; Ni/AI,O, (Ref. 113); @, 3% Ni/Al,O, (Ref. 158); 0. 3% Ni/AI,O, (Ref. 112). Open symbols, 5 or 10 ppm H,S; closed symbols, 25 ppm H,S; 8.50 ppm H,S.

174

C. H. BARTHOLOMEW ET AL

A linear correlation between the decrease in volume of H, adsorbed ( VH2) and the volume of H,S adsorbed (VHZs) is also demonstrated by the

data of Ng and Martin (111) for Ni/SiO, shown in Fig. 17. The ratio of the intercepts for the abcissa and ordinate is VHZs1 VH2 = 20130 or 213. Assuming one hydrogen atom adsorbs per nickel site, the value of H,S/Ni, is 1/3 for adsorption of H,S at 300 K, in good agreement with the earlier work of Saleh et af. (96) and of Den Besten and Selwood (104). Values of a(H,),the change in saturation magnetization per adsorbed molecule of hydrogen, are apparently constant over most of the range of sulfur coverage, suggesting that the adsorption properties of the free nickel sites are not strongly influenced by the sites on which sulfur is adsorbed. In other words, adsorbed sulfur appears to poison H, adsorption in supported nickel by a simple blocking mechanism. Hence the studies of Ng and Martin and Bartholomew et al. establish that H, adsorption at 298 K on partially sulfur-poisoned catalysts provides an accurate measure of the unpoisoned nickel surface.

I

I

\

I

10

20

VHZs (ml NTP/g N i l

FIG. 17. Volume of adsorbed H, at room temperature and ca. 5 Torr and g ( H Z )versus the volume of preadsorbed H,S (adapted from Ref. 111). B. M. denotes Bohr magnetons.

SULFUR POISONING OF METALS

175

In terms of the quantitative effects of sulfur on H, adsorption, the stoichiometric data of Bartholomew ct ul. are probably more <useful than those of Ng and Martin since catalysts from Bartholomew’s work were presulfided in 5-25 ppm H,S/H, at 725 K, a temperature and concentration range more representative of poisoning processes. Moreover, a t 725 K and 5-25 ppm H,S/H,, the surface is saturated with sulfur atoms, whereas at 300 K the adsorption of H,S results in strongly adsorbed sulfur and hydrogen, both of which can prevent further H, adsorption at the same temperature. The studies mentioned to this point (111, 162) involved H, adsorption at 298 K. Pannell et ul. (141) investigated the effects of preadsorbed sulfur on H, adsorption at 475 K, their data indicating a significantly smaller reduction in H, adsorption at 475 K, i.e., 10% compared to 33% at 298 K, for the same presulfided sample. The increased hydrogen uptake at 475 K on the sulfided catalyst is in agreement with work by Griffith and co-workers (166, 167). Since only 5- 10% of the adsorbed sulfur can be removed in pure H, at 725 K (106, 112), it is unlikely that the increased hydrogen uptake is a result of removing adsorbed sulfur. However, this observation does suggest a different mechanism of H, adsorption on presulfided nickel at 475 K than at 298 K. Perhaps H, is adsorbed weakly on the sulfided portion of the surface at the higher temperature. Regardless of the mechanism, on the basis of its behavior, H, adsorption at 475 K is not a good indicator of the unpoisoned nickel surface area. Studies of single-crystal nickel (163-165) show that the decrease in H, adsorption is not linear over the full range of sulfur coverage. For example, Rendulic and Winkler (163) observed a linear decrease in the initial sticking coefficient of hydrogen on Ni with increasing sulfur coverage below Q, = 0.3 (see Fig. 18). However, above 19, = 0.3 the slope decreases and the curve deviates to higher coverages. At maximum sulfur coverage (0, = 0.5), the sticking coefficient is virtually zero. Kiskinova and Goodman (164) observed a similar relationship between the sticking coefficient of hydrogen on Ni(100) and sulfur coverage. The linear decrease in the sticking coefficient at low coverage (0, < 0.2), equivalent to poisoning of about four nickel atoms per adsorbed sulfur atom, was attributed to an electronic interaction. At QS > 0.2 the number of nickel atoms poisoned per adsorbed sulfur atom decreased significantly with increasing coverage, presumably due to changes in the surface structure and repulsions between sulfur atoms. In view of the nonlinear decrease in H, adsorption with sulfur coverage in the single-crystal studies, it is interesting that such well-defined linear decreases with mean sulfur coverage are observed in the studies of polycrystalline and supported nickel (111, 162). However, since the partially poisoned catalysts are thought (111, 162) to consist of completely poisoned

176

C. H. BARTHOLOMEW ET AL.

MEAN SULFUR COVERAGE (Ns/"I)

FIG. 18. Variation of the initial sticking coefficient of hydrogen on polycrystalline nickel as a function of sulfur coverage (adapted from Ref. 163).

and unpoisoned phases having sticking coefficients of either 1 or 0, a linear decrease in H, proportional to the fraction of catalyst completely poisoned is expected. How does preadsorbed sulfur prevent hydrogen adsorption? The studies of single-crystal Ni (163-165) suggest that electronic effects may be important at low coverage since one atom of sulfur can prevent hydrogen adsorption on four or more Ni atoms. At intermediate coverages desorption states of H, are shifted to lower energies (164, 165) suggesting that sulfur weakens the bonds between Ni and hydrogen. Johnson and Madix (165) obtained evidence of both electronic and geometric effects of adsorbed sulfur on the desorption kinetics of H, on Ni(100). Both activation energies and frequency factors were substantially reduced for hydrogen recombination on the 4 2 x 2 ) surface relative to the clean surface. The lower activation energy is suggestive of electronic changes due to sulfur whereas the significantly lower (by 6 orders of magnitude) frequency factor appears to result from severe steric hindrance to hydrogen recombination. Thus, at complete coverage a simple blocking or geometrical effect is adequate to explain the poisoning since the ensembles of Ni required for dissociation of H, are no longer accessible at the surface. 2. Effects of' Preadsorbed Sdjur on CO Adsorption

The effects of sulfur poisoning on the adsorption of CO on nickel (46,49, 102, 108, 111, 116, 120, 121, 141, 157-160, 162, 164, 165, 168-172) are very complex, the nature of the adsorbed species and the adsorption

177

SULFUR POlSONING OF METALS

stoichiometry varying considerably with changes in pressure, temperature, and sulfur coverage. For example, IR data (120) show that bridged bonding is diminished and subcarbonyl bonding enhanced by sulfur. Moreover, there are significant differences in the behavior of supported and unsupported Ni (141, 162). In fact, in the case of supported nickel catalysts very significant increases in CO adsorption are observed at 190-298 K and moderate pressures (10-60 kPa) after treatment with H,S (157-160, 162), apparently as a result of sulfur-catalyzed Ni(CO), formation (162, 273, 174). However, at low CO pressures (less than 0.1-0.5 kPa), sulfur diminishes room-temperature CO adsorption on single-crystal Ni (46, 49, 164, 165), shifts the desorption states to lower energy, and completely poisons the highest energy state ( p z ) at 0 r 0.3 (46, 49, 164) on Ni(ll1) or (100) and in the form of a p ( 2 x 2) diminishes or as a c(2 x 2) eliminates fourfold coordinate a-CO on Ni(100) (165). On the other hand, under high-temperature reaction conditions representative of methanation (475-675 K, 100 kPa, Pco = 5-10 kPa, Hz/CO = 3), dissociative adsorption of CO occurs on a completely

c

0.2

0

0.2 0.4 0.6 0.8 1.0

2

FIG.19. SIMS intensity ratio for CO-containing species versus sulfur coverage in terms of Ni,S+/Ni: for Ni(100) and Ni(l11) (Ref. 116).

178

C. H. BARTHOLOMEW ET A L .

poisoned Ni/AI,O, catalyst, the extent of which is about 50% of that observed for an unpoisoned catalyst (172). The results of a SIMS study at low CO pressures by Bordoli et ul. (116) provide evidence that effects of sulfur on CO adsorption vary with surface structure and with the nature of the adsorbed species (see Fig. 19). On Ni(100) the concentration of bridged CO (proportional to the Ni,CO+/Ni,+ intensity) decreases in almost linear fashion with increasing sulfur coverage (in terms of Ni,S+/Ni,+). Although the concentration of the linear species (NiCO+/Ni+)is constant over a wide range of sulfur coverage, it then decreases rapidly near saturation coverage. Apparently no species is adsorbed at saturation coverage. The changes in these species appear to correlate well with decreases in the fll and p2 states on Ni(100) observed by Kiskinova and Goodman (164) as a function of O,, although the latter group (164) observed a significant amount of adsorbed (presumably non-fl) CO even at 8 = 0.5. Johnson and Madix (165), however, observed both and f12 species on the x 2) sulfur overlayer on Ni(100). Thus, there are obviously significant c(2 discrepancies among the results reported (116, 164, 165) for Ni (100). On Ni( 1 11) the amount of bridged species (in Fig. 19) decreases in exponential fashion to zero with increasing 0, , whereas the coverage of the linear species drops rapidly to one-half of its original value but is not further decreased with increasing 0, , suggesting that linear species can adsorb on a completely poisoned Ni( 1 1 1) surface. The effects of sulfur on CO adsorption at moderate pressures and low temperatures are illustrated in Tables VIII-X and Figs. 20-21. Table VIII shows H, adsorption uptakes and ratios of adsorbed CO molecules to hydrogen atoms determined by Pannell and co-workers (141) for an unsupported nickel powder (INCO) before and after exposure to 5 ppm H,S in H, at 5 5 5 K for 6 h. As expected, the H, uptake is reduced significantly by the presulfiding treatment. Irreversible and total CO adsorption on the freshcatalyst are equalat 190 K, whereas total CO adsorption is two times the irreversible at 300 K and four times that at the lower temperature. These data suggest that nickel carbonyl or subcarbonyl-type intermediates [e.g., Ni(CO), or Ni(CO),] are formed at 300 K but not at 190 K. Upon poisoning with H,S the (CO/H)irr,,,,,ib,, decreases 35-40% at both temperatures, indicating that adsorbed sulfur diminishes the amount of strongly adsorbed CO. Nevertheless, the total adsorption at 300 K doubles after poisoning, the CO/H stoichiometry consistent with formation of Ni(CO),. In other words, the formation of nickel tetracarbonyl is catalyzed by the presence of adsorbed sulfur. The formation of nickel carbonyl at temperatures above 273 K and pressures above 0.1 kPa and its promotion by adsorbed sulfur have been reported by other workers (121, 162, 173, 174). The effects of sulfur poisoning on CO adsorption are even more complex in the case of supported catalysts. Again, there are two different kinds of

179

SULFUR POISONING OF METALS

TABLE V l l l H , Adsorption Uptakes and C O J H Ratios for Unsupported Nickel (ONCO) before and ufter Pre.dfidiny"

CO/H' H, uptakeh (Pmolig)

CO adsorption temperature (K)

Total

Irreversible

Fresh

4.52

Poisoned'

3.0 I

190 300 190 300

0.54 2.09 0.35 3.85

0.54 1.16 0.35 0.70

Sample

From Ref. 141. Total H, uptake measured at 300 K and 13-53 kPa. Molecules of CO adsorbed at the temperature indicated and 13--53kPa per hydrogen atom adsorbed at 300 K. Irreversible CO uptake is the difference between the total uptake and that measured after evacuation for 30 min at the temperature of adsorption. Exposed to 5 ppm H,S for 6 h at 555 K and a space velocity of 2000 hr-'. a

TABLE 1X H 2 and C O Adsorption Uptakes and C O / H Ratiosfor 3% NiIA120, before und ajrer Presulfiding"."

Catalyst Fresh Presulfided Fresh Presulfided Fresh Presulfided after 2 runs Fresh Presulfided (run 1) Presulfided (run 2)

H, uptake' &mol/g)

Temperature (K) of CO adsorption

CO uptaked trmolig)

CO/H"

21 14 30.6 23.2' 30.6 23.2' 8.0' 41.5 0 0

I90 I90 273 213 300 300

80 374 121 950 109 315

1.9 13.4 2.0 20.4 1.8 6.8

273 273 273

316 563 35

3.8 6.8"

~~

From Ref. 162. Presulfided in 10 or 25 ppm H,S/H, at 725 K at a space velocity of 2000 hr-' for 6-12 hr. Total hydrogen uptake at 300 K and 13-53 kPa. lrreversible CO uptake at temperature shown and 13-53 kPa. Irreversible uptake is the difference between total uptake and that after evacuation at the adsorption temperature for 30 min, also corrected for adsorption on A120, Molecules of CO adsorbed at the temperature indicated per H atoms adsorbed at 300 K. The same sample was tested before and after poisoning at both 273 and 300 K. The value of 23.2 pmol/g was determined after poisoning but before CO adsorption at either temperature. After CO adsorption measurements of the poisoned catalysts, the H, uptake was 8.0; the sample changed from black to gray, suggesting a significant loss of Ni via formation of the tetracarbonyl. Molecules of CO adsorbed at 273 K per atom H adsorbed on the fresh sample. "

180

C. H. BARTHOLOMEW ET AL.

TABLE X Infrared Bands for CO Adsorption on NilSiO, :Assignments and the EfJtcts of Sulfur Poi.7oning Thereon“ Infrared band (cm- ’) I960

Structural assignment

Effect of sulfur poisoning

Bridged or multicenter: 0

I1 / \

Progressively, irreversibly poisoned with increasing sulfur coverage. Band is not observed at high S coverages

C

Ni 2030-2050

Ni

Linear: 0 C

’.

I

Ni 2050-2090

Subcarbonyl : 0 0 0

c c c \I/ Ni

a

Progressively, irreversibly poisoned with increasing sulfur coverage with concomitant growth of band at 2056-2090 cm- The band is completely shifted at high sulfur coverages Band increases in intensity with increasing amounts of preadsorbed sulfur at low sulfur coverages and subsequently decreases at higher coverages. This band is not completely diminished at high sulfur coverages. Can be removed, however, by evacuation for 30 min. Can be regenerated in part with oxygen treatment

0

0

c c \/ Ni

From Ref. 120.

behavior depending upon the adsorption pressure of CO. That is, at low pressures (less than 0.1 kPa) CO adsorption (both total and irreversible) is decreased (110, 111, 169, 170); however at moderate pressures (13-53 kPa) both total and irreversible CO uptakes are dramatically increased (157-160, 162)as illustrated by data for Ni/Al,O, catalysts in Table IX and Fig. 20. For example, decreases in H, uptake ranging from 20-100% are evident in Table IX for 3% Ni/Al,O,. At the same time, quantities of “irreversibly” adsorbed CO increase dramatically from 300 to 800% depending upon sulfur coverage and temperature. Assuming again that H/Ni, = 1, the adsorption stoichiometries are consistent with CO/Ni, values ranging from 7 to 20. These values of CO/Ni, are consistent with the formation of subcarbonyl species [Ni(CO), ,x = 2-31 and of Ni(C0)4,which is able to migrate to the gas phase and/or the support (depending upon the adsorption temperature) thereby exposing new nickel sites for adsorption of CO.

181

SULFUR POISONING OF METALS

I-

0

I 0

I

I

I

I

100

200

300

400

J

PRESSURE(Torr) FIG.20. Effects of H I S on CO adsorption on 8% Ni/AI,O, (open symbols denote fresh sample; closed symbols denote poisoned sample) (Ref. 162).

The hypothesis of tetracarbonyl and subcarbonyl species formation on sulfur-poisoned nickel catalysts finds considerable support from previous IR studies (102,120, 121,168-170). Indeed, this topic is discussed in some detail by Rochester and Terrell(120), who associate various modes of CO adsorption with observed IR bands consistent with data from their poisoning experiments at 300 K on a 7.6% Ni/SiO, and previous IR studies. The important IR bands, their molecular assignments, and the observed effects of sulfur poisoning are summarized in Table X. Accordingly, the band intensities at 1960 and 2030-2050 cm- ' are progressively diminished and the intensity of the band at 2065-2090 first increases, then decreases with increasing sulfur coverage. These observations are interpreted by the authors as (1) inhibition of bridged or multicenter adsorption of CO and (2) a weakening of the linear CO complex accompanied by addition of a second and/or third CO molecule to each surface nickel atom to form subcarbonyl complexes. The volatile tetracarbonyl is also formed and observed in the vapor phase at 2055 cm- especially if the samples are heated. The assignment by Rochester and Terrell(l20) of the band at 1960 cm-' to strongly held, bridge-bonded CO is reasonable since sulfiding of the surface could greatly reduce the availability of suitable clusters or adjacent pairs acting as multiple adsorption sites. It is also consistent with the observation of Pannell and co-workers (141) that strongly adsorbed CO is reduced by preadsorbed sulfur. The results of Rochester and Terrell are also in good

',

182

C. H. BARTHOLOMEW ET AL.

1 9 4 6 cm-'

0.1

0

60

H,S

80

EXPOSURE(^^^

x

1 0 7

10')

FIG.21. Effect of H2S on CO preadsorbed on Ni/Al,O, catalyst (CO exposure, 7.1 x l o b 4 mol). (Adapted from Ref. 169.)

agreement with those of Crell et al. (168) who reported that IR peaks for CO at 2060 and 1950 cm-' were absent after treatment with H2S; however, CO adsorption on the presulfided surface led to the appearance of a new peak at 2095 cm-'. Garland (102)likewise indicated that IR bands at 2045 and 1960 were progressively diminished by increasing doses of CS, . Moreover, CO adsorption on heavily sulfided catalysts produced a single band at 2084-2094 cm-', which could be completely removed by a 30-min evacuation at room temperature. More recently, Rewick and Wise (169, 170) observed an increase in band intensity at 2075 cm-' and a corresponding decrease at 1946 cm-' for CO adsorption on a catalyst exposed to progressively larger doses of H2S. However, at sufficiently high coverages of sulfur, the intensity of the peak at 2075 cm- ' passed through a maximum and then decreased with increasing coverage (see Fig. 21). The elimination of the bridge-bonded species (1946

SULFUR POISONING OF METALS

183

cm- ') on a presulfided catalyst, they hypothesized, was due to a restructuring of the nickel surface to a 2-D nickel sulfide, causing the disappearance of neighboring Ni sites. In this configuration the surface Ni ions could bind CO molecules weakly and in a linear, single-bonded fashion only. The facts that sulfur poisoning of CO adsorption is less inhibited and tetracarbonyl formation is enhanced on supported Ni catalysts compared to polycrystalline or single-crystal nickel suggest that either support or metal crystallite-size effects may play a role in determining the CO adsorption properties. Hegedus and McCabe ( I 75) speculate that SiO, and A1,0, supports create low-coordination nickel sites which more readily adsorb linear and subcarbonyl CO. This is consistent with experimental observation of Bartholomew and Pannell(176) showing that CO/H adsorption ratios for well-dispersed Ni/Al,O, and Ni/SiO, increase with increasing Ni dispersion. However, Bartholomew and Pannell conclude that this observation may be explained by either crystallite-size or support effects. Moreover, most of the available evidence favors an electronic interaction of the nickel crystallites with the support (177) to explain the changes in CO adsorption behavior and the associated greater resistance to poisoning of CO adsorption by supported Ni. The dramatic increase in irreversible CO adsorption on presulfided supported nickel catalysts at moderate pressures (162) has significant, practical implications in regard to the use of CO chemisorption to measure nickel dispersion. For example, it is often desirable to determine nickel surface areas for catalysts used in a process where sulfur impurities are present in the reactants. Substantial differences in the measurements of nickel surface area by H, or CO adsorption are possible depending upon the catalyst history and choice of adsorption conditions. In view of the ease with which catalysts may be poisoned by sulfur contaminants at extremely low concentrations in almost any catalytic process, and since large CO uptakes may be observed on supported Ni not necessarily representative of the unpoisoned nickel surface area, the use of CO adsorption to measure nickel surface areas is highly questionable under almost any circumstance. The fact that the nature of CO adspecies on nickel is considerably modified by the presence of adsorbed sulfur has important implications for reactions involving CO as a reactant, particularly for synthesis reactions. Although high coverages of sulfur generally cause complete loss of activity, small amounts of sulfur may bring about desirable changes in selectivity, assuming the linear or subcarbonyl form of CO rather than a bridged form is the desired intermediate. The relationship of the above-discussed effects of sulfur poisoning on CO adsorption to changes in selectivity in synthesis reactions is discussed in Section V on reaction studies.

184

C. H. BARTHOLOMEW ET AL.

3. Ejfects of Sulfir on the Adsorption of Other Molecules Since sulfur poisoning is known to modify selectivity properties in a number of hydrocarbon reactions, investigation of the effects of preadsorbed sulfur on the adsorption of organic molecules on metal catalysts may provide the basis for understanding this behavior. Only two such studies involving Ni have been reported (111, 171). Ng and Martin (111) investigated acetylene and benzene adsorption on presulfided Ni/SiO, , their volumetric, adsorption, and saturation magnetization data revealing a more complex behavior than that of either H, or CO. Both acetylene and benzene adsorb on presulfided Ni/SiO, at room temperature, these results suggesting interactions of the hydrocarbons with deeper nickel layers. Perhaps the most significant observation from this study is that preadsorbed H,S inhibits the cracking of hydrocarbons on nickel. In a study of temperature-programmed desorption of methanol on Ni(100) Johnson and Madix ( 1 7 1 ) found that the distribution of products from desorption was significantly modified by the presence of sulfur. For example, on the clean Ni(100) surface methanol was cracked to form H, and CO, whereas on the Ni(100) c(2 x 2) S surface, principally formaldehyde was formed. Their results, like those of Ng and Martin ( I l l ) , suggested that preadsorbed H,S inhibits the cracking of hydrocarbons such as methanol on nickel. The adsorption of 0, is apparently affected by the presence of adsorbed sulfur; for example, 0, adsorbs strongly on clean Ni(ll1) at room temperature but only weakly on the presulfided surface (1 78). However, the same behavior is not observed at high-temperature reaction conditions. For example, Neff and Kitching (121) reported that a sulfided nickel surface reacted readily with oxygen at 353 K. Colby (9)also observed similar results, namely that atmospheric pressure 0, treatment of sulfur-poisoned nickel at 573 K resulted in bulk nickel oxide formation. The adsorption of oxygen on sulfur-poisoned metal surfaces has significant importance from the viewpoint of regeneration of the sulfur-poisoned catalysts. This subject will be dealt with in more detail later, as we consider the regeneration of sulfurpoisoned catalysts. That only two investigations (111, 171) so far have considered the effects of sulfur poisoning on the adsorption of organic molecules and only three studies (9, 121, 178) poisoning of the adsorption of 0, on nickel, strongly suggests needs for further research. A number of important molecules are suggested by considering some of the important reactions in which nickel catalysts are poisoned by sulfur, e.g., ethylene (hydrogenation), methane (steam reforming), and cyclohexane (hydrocracking). In addition, the study of poisoning of H, and CO adsorption by different sulfur compounds is of

SULFUR POISONING OF METALS

185

high priority, especially since only one investigation (120) has considered such effects.

B. OTHERMETALS Relatively few studies (76, 79, 83, 88, 96, 101, 125, 157-161, 179-191) have examined the consequences of sulfur poisoning on the adsorption of molecules on metals other than Ni. Those systems having received the greatest attention include CO on Fe (179-182), CO on Pt (83a, 88, 183-185), and H, and CO on Ru (79, 144, 158, 187, 188). Effects of sulfur on H, adsorption on Co (157) and Mo (76); on CO adsorption on Co (101), Au, Os, Re, Rh (187), and Zn (189); on 0, adsorption on Pt (83); NO, benzene, and acetylene adsorption on the Pt( 100) surface (88); formaldehyde adsorption on the Ru(lI0) surface (190); and alkyl aromatic adsorption on Pt (191) have also been reported. Generally, it appears that adsorptions of hydrogen, oxygen, and CO are prevented by adsorbed sulfur. For example, Bonze1 and Ku (83a) observed that partially sulfurized Pt( 1 10) surfaces were characterized by several new CO binding states whose energies decreased with increasing sulfur coverage. At low coverages (0 < 0.25), their data were consistent with each sulfur atom blocking two CO chemisorption sites ; at higher coverages the slope deviated to higher coverages than would be allowed by a simple c(2 x 2) structure, suggesting phase changes at the surface. LEED data indicated that these phase changes may result form repulsive interactions between S atoms at higher coverages. Similar observations were reported by Keleman et al. (185) for CO adsorption on sulfurcovered Pt(ll1) and (100). Bartholomew and co-workers (157-161) observed that hydrogen uptakes for nickel bimetallics and ruthenium were generally decreased proportional to sulfur coverage, whereas Schwartz (79) reported that sulfur blocks both dissociation sites for H, adsorption and recombination sites for H, desorption on Ru(100). Evidence that adsorbed sulfur atoms rather than sulfur molecules are responsible for blocking the adsorption of hydrogen was obtained by Kikuchi and Ishizuka (76). They observed that a clean Mo surface, after preadsorption of S, vapor at low temperature, could be saturated with Hz at room temperature. However, when the sulfided surface was heated to 773 K in vacuum, only a very small amount of H, could be adsorbed. A few previous studies (83a, 181,183,185,187) have examined the effects of preadsorbed sulfur on the nature of adsorbed CO species or CO adsorption states on metals. Rhodin and Brucker (181) investigated variations in CO bonding on clean and partially sulfur-deactivated a-Fe( 100) surfaces.

186

C . H . BARTHOLOMEW ET A L .

Bonding of adsorbed CO to clean Fe was found to occur according to the Blyholder model (192), with complete dissociation at 300 K. On a sulfursaturated (8 = 0.5) Fe surface, no stretching in the CO bond was observed, apparently due to reduction in forward- and back-bonding interactions. Agrawal(101) investigated the effects of sulfur on dissociative adsorption of CO on Co/AI,O,. Three Co surfaces with different sulfur coverages (0 = 0.5, 0.3, and 0.15) were exposed to a gas mixture (10% C0/900/;:H,) at 1 atm total pressure and at 673 K for 2 hr. After the treatment, the surfaces were examined by using AES. No carbon was present on the Co surfaces with 0 = 0.5 and 0.3, whereas multilayer graphitic carbon deposits were observed on the Co surface with small amounts of sulfur (8 = 0.15). This result is significant in that severe carbon deactivation of Fischer-Tropsch catalysts can be prevented by small quantities of sulfur present on the surface. Since adsorption and dissociation of CO are necessary intermediate steps in the carburization of catalysts, the above results suggest that dissociative chemisorption of CO on Co is inhibited by the presence of sulfur. Similar results were obtained by Goodman and Kiskinova (193) for dissociative adsorption of CO on Ni( 100). Guerra (187) reported shifts in the M-CO bond to lower frequencies due to poisoning of Ir and 0 s by H,S. Exposure of Rh to H,S enhanced the bond attributed to two CO molecules per metal site and eliminated bridged Rh,-CO species. Thus sulfur appears to affect the adsorption states for CO on some group VIII metals in a manner similar to nickel. As in the case of Ni it appears that the effects of sulfur on the adsorption of organic molecules on other metals are more complex than for either H, or CO. For example, formaldehyde (HCHO), which dissociates at 80 K on a clean Ru(ll0) surface (190),adsorbs molecularly on the sulfur-poisoned surface (78). Similarly, Cosyns et al. (191) suggest that sulfiding of the platinum enhances the adsorption of the alkylaromatics compounds. Since preadsorbed sulfur generally blocks the adsorption of other molecules, it would only be logical to expect that it would also prevent the adsorption of H,S or S, . Previously discussed studies of sticking coefficients (73, 83, 92, 99, 101) and H,S adsorption on metals (57, 106, 112-115) provide evidence that the sticking coefficient and heat of adsorption for H,S or S, decrease with increasing coverage. Thus, rates and strengths of sulfur adsorption on sulfur-saturated metal surfaces are clearly lower than those on a clean metal surface. To summarize, previous studies have established the qualitative effects and to a lesser extent quantitative effects of preadsorbed sulfur on the adsorption of other molecules, particularly CO and H,. Unfortunately, in most of the previous work quantitative relationships between the coverage of sulfur and decrease in adsorption or changes in adsorption states for a given adsorbate were not obtained. Thus, determination of the quantitative effects

SULFUR POISONING OF METALS

187

of sulfur on adsorption of reactants on metals, particularly those having received little attention such as Co and Rh, is a potentially fruitful area for further research.

V. Effects of Sulfur on Catalytic Activity and Selectivity Properties of Metals

Because sulfur adsorbs very strongly on metals and prevents or modifies the further adsorption of reactant molecules, its presence on a catalyst surface usually effects substantial or complete loss of activity in many important reactions, particularly in hydrogenation reactions. Where the reaction network leads to two or more products, adsorbed sulfur can markedly affect the selectivity by reducing the rate of one of the reactions more than the other(s). In a few reaction systems these changes in selectivity are desirable; however, in many others they are not. The purpose of this section is to critically review work dealing with the effects of sulfur poisoning on catalytic activity and selectivity in several important reaction systems, to quantify rates and extents of sulfur poisoning, and to identify poisoning mechanisms and sulfur tolerances of important catalysts where possible. It is also important to address the complexities of investigating sulfur poisoning and to review experimental techniques which have been proven to provide accurate, unambiguous data. Consistent with the overall theme of this review and with the availability of literature, hydrogenation and synthesis reactions on supported metals, particularly nickel, are emphasized.

A. EXPERIMENTAL CONSIDERATIONS The experimental techniques utilized in the study of sulfur poisoning are markedly more critical to obtaining quantitative, basic data than in any other type of reaction study. In most previous studies of sulfur-poisoned catalysts, rate data were affected by experimental complexities to a sufficient extent that a basic understanding of poisoning rates and mechanisms was not possible. Most of these difficulties occurred because of: (1) strong nonuniform adsorption of sulfur on the catalyst; (2) choice of an improper catalyst form (i.e., pellets rather than powder or supported rather than unsupported metal) ; (3) choice of an improper reactor type and reactor configuration;

188

C. H. BARTHOLOMEW ET AL

(4) choice of inappropriate materials of construction which reacted with sulfur-containing compounds ; and ( 5 ) choice of improper experimental conditions so that heat and mass transport limitations disguised the intrinsic effects of sulfur poisoning. These problems were most severely aggravated by the combined use of a porous supported catalyst in a fixed-bed reactor. For example, in a fixed bed the extremely strong adsorption of sulfur on metal catalysts results in a sharp concentration gradient of poison which moves slowly like a wave front through the catalyst bed. The catalyst behind the wave front is severely poisoned, whereas that ahead of the wave front is essentially poison free (194, 195). In addition to this problem there is a poison wave front moving radially into each porous catalyst particle severely poisoning an increasingly thick peripheral region surrounding an unpoisoned core (113, 194). This nonuniform sulfur distribution in bed and particle makes fundamental interpretation of the sulfur-poisoning process extremely difficult or impossible. However, it is possible to infer effects of poisoning from comparison of fresh and completely poisoned catalysts, but great care must be taken to ensure that the catalyst is poisoned under equilibrium conditions. It is also important to avoid PHzJPHI values large enough to cause bulk metal sulfide formation, since this introduces additional and usually undesirable complications. Unfortunately, much of the previous work is suspect for this very reason. The choice of materials of construction for controlled-atmosphere studies is also crucial. Most metals adsorb, desorb, and react with sulfurcontaining compounds depending upon the experimental conditions. For /PH2values ( < 1- 10 ppm), adsorption, desorp experiments at very low PHls tion, and even generation of H,S from metal- and glass-containing systems result in serious problems, particularly at reaction temperatures (196). Quartz and teflon are the preferred materials of construction for studies at low H,S concentrations (99, 100, 114, 140, 196, 197). To enable quantitative determination of rates of sulfur deactivation, of extents of sulfur deactivation at very low gas-phase sulfur concentrations, of true dynamic equilibrium between gas-phase sulfur concentration and metal surface, and/or of the amount of sulfur adsorbed on the surface, the following requirements must be satisfied in the design of experimental apparatus : (1) The reactor must be gradientless with respect to reactant(s) and to gas-phase sulfur concentration so that all parts of the catalyst are exposed to the same gas-phase sulfur concentration. Even a shallow bed of catalyst operated at very low conversion of reactant(s) is not necessarily gradientless

SULFUR POISONING OF METALS

189

with respect to sulfur because of the very strong, nonuniform adsorption of sulfur on metals. (2) All surface metal atoms must be exposed to the same gas-phase concentrations of sulfur-containing compound and reactants. Because of the very strong adsorption of sulfur, metal atoms at a short distance into the pore of a porous catalyst will not experience the same concentration of sulfur as the metal atoms near the pore mouth or on the exterior surface of the catalyst. (3) The reactor should neither adsorb nor desorb sulfur in quantities which are significant compared to adsorption on the catalyst nor in such a manner as to decrease or increase the gas-phase concentration. (4) The feed-gas composition must be carefully controlled and gasphase sulfur compound concentrations as low as 10 ppb or less must be reproducibly prepared. ( 5 ) On-line analytical instruments must be capable of quantitatively measuring gas-phase sulfur concentrations as low as 5 ppb and of quantitatively measuring reactants and products for the reaction under study for a broad range of concentrations. In essentially no previous studies have all these requirements been met (in most cases none of them has). An example of a system which most nearly meets these requirements is a quartz continuous-flow stirred-tank reactor (CFSTR) (99-101,140,196,197) with catalyst configurations in which all surface metal atoms are on the exterior surface of the support. It satisfies the relevant requirements listed above and allows investigation over a broad range of both product and reactant concentrations. Furthermore, true poisoning rates can be measured directly, without requiring assumption of a model for the poisoning. The amount of sulfur adsorbed can be directly determined as a function of time and gasphase H,S concentration, and the catalytic activity of the metal can be measured as a function of sulfur on the surface.

B. CO HYDROGENATION REACTIONS Sulfur poisoning is a key problem in hydrocarbon synthesis from coalderived synthesis gas. The most important hydrocarbon synthesis reactions include methanation, Fischer-Tropsch synthesis, and methanol synthesis, which occur typically on nickel, iron, or cobalt, and ZnO-Cu catalysts, respectively. Madon and Shaw (2) reviewed much of the early work dealing with effects of sulfur in Fischer-Tropsch synthesis. Only the most important conclusions of their review will be summarized here.

190

C. H . BARTHOLOMEW ET AL.

1. Ejects of Sulfur on Activity and Selectivity

From the pre-1975 literature (2, 198) it is apparent that gas-phase sulfurcontaining compounds, especially H,S, when present in only ppm quantities for a sufficient period of time cause order of magnitude losses in global activity for methanation and hydrocarbon synthesis over Co, Fe, and Ni catalysts. For example, Fig. 22 shows the loss of activity with increasing amounts of adsorbed sulfur for hydrocarbon production on a fused-iron catalyst (199). These data show how a commercial packed-bed reactor system behaves with sulfur poisoning, but they suffer from the problems discussed above. The observed activity reduction is the response to a material balance on sulfur fed to the reactor and a poison wave moving through the reactor. There are also undoubtedly severe poisoning nonuniformities in the catalyst

I .o

>-

k > -

0.8

I-

0

a w

z

0.6

>

-1

0.4

0.2

I

0

I

I

1

I

0 .I 0.2 0.3 0.4 SULFUR FEED TO CATALYST, mg S / g Fe

FIG.22. Poisoning of reduced fused iron catalyst by H,S: T = 535 K, HJCO = 1, P = 2.16 MPa. Sulfur concentration in feed (mg S/m3): (0) 6.9, ( 0 )23.0, (m) 69.0. Reprinted with permission from Ref. 199. Copyright 1963, 1964 American Chemical Society.

SULFUR POISONING OF METALS

191

particles; these probably explain the limited extent of sulfur deactivation observed. Alkali and oxide promoters such as K,O and A120, apparently improve the resistance of Fischer-Tropsch catalysts to poisoning. For example, Anderson et al. (200) found that Al,O,-promoted Fe evidenced somewhat greater resistance to sulfur poisoning than MgO-promoted Fe. Moreover, the addition of alkali improved su@r resistance (i.e., decreased the rate of deactivation) by a factor of 10. This effect may be due to the scavenging of sulfur by the alkali metal to form a highly stable sulfide, thereby extending the active life of the transition metal. If so, the sulficr tolerance (i.e., the steadystate activity) should be no higher with alkali than without. This, however, has not yet been proved or disproved. Selectivity properties of Fischer-Tropsch catalysts can be significantly altered by addition of relatively small quantities of sulfur during preparation or in the form of a gas-phase pretreatment (2). However, the changes are reasonably complex, varying with catalyst composition and the form and extent of treatment with sulfur and the details of the entire experimental procedure. For example, Herington and Woodward (201) found that a small amount of sulfur added as gas-phase H2S caused a marked increase in the yield of liquid hydrocarbons and a decrease of gaseous hydrocarbons produced over a Co/ThO,/kieselguhr catalyst. They attributed this effect to the preferential poisoning of hydrogenation activity leading to formation of longer hydrocarbon chains via polymerization. However, Schultz and co-workers (202) found that initial additions of small amounts of sulfur to an iron catalyst (up to 1 mg S/g Fe) prevented wax formation while increasing CI to C, hydrocarbon formation; the capacity to form liquid-range hydrocarbons was unchanged. However, further additions of sulfur to the catalyst caused a progressive reduction in the molecular weight of hydrocarbon products until only gaseous products were observed. Anderson and coworkers (200) also found that the changes in selectivity due to sulfur depended on particle size, the smaller catalyst particles producing more wax upon the initial addition of sulfur consistent with the results of Herington and Woodward (201). The effects of sulfur on product distribution also varied for nitrided, carbided, and freshly reduced iron catalysts (199). That the results of these early studies (199-202) were undoubtedly influenced by mass-transfer limitations and nonuniform sulfur concentrations, explains why contradictory trends were observed with particles of different size. Nevertheless, one or two general principles are suggested. For example, it is generally believed that olefins are a primary product of hydrocarbon synthesis over transition metals, that these may react to form higher molecular weight products but that they are rapidly hydrogenated to paraffins over an unpoisoned catalyst. Small amounts of sulfur apparently markedly

192

C . H. BARTHOLOMEW ET AL.

FIG.23. CO hydrogenation at 523 K on (1) 5% Ni/ZrO,; (2) 2% Ni/AI,O,; (3) Raney Ni (Ref. 204).

reduce the hydrogenation of the olefins formed and may also reduce the rate of chain termination leading to a generally observed increase in the formation of C , + hydrocarbons for integral reactor operation. Larger amounts of sulfur produce severe deactivation, for which condition good selectivity data are not available from the early work. The effects of sulfur poisoning on activity/selectivity properties of Ni, Co, Fe, and Ru in methanation were more recently investigated by Dalla Betta et al. (203, 204), Bartholomew et al. (23, 113, 140, 172, 194), Katzer and co-workers (99-101, 147, 205-208), Rostrup-Nielsen and Pedersen (209), Wentrcek et al. (195), and Goodman and Kiskinova (193). Dalla Betta and co-workers determined steady-state activities of Ni/A1,0,, Ni/ZrO,, and Raney Ni catalysts (powder form) in a tubular packed-bed reactor in the presence of 10 ppm H2S. At 523 K the activities of Ni/ZrO, and Raney Ni decreased about two orders of magnitude, whereas the activity of Ni/Al,O, was lowered three orders of magnitude during 24-hr exposure to 10 ppm H,S in a CO-H, feedgas (see Fig. 23). The loss of activity was about an order of magnitude less for samples pretreated 24 hr in a sulfur-free reactant mixture at 623 K (similar results were obtained for Ru) and two orders of magnitude greater if the reactant mixture contained water vapor in

193

SULFUR POISONlNG OF METALS

€

Ni/A1p3

10' : -D

r

lo2

-

-

lo3

loQ

105

.I! D

A

E

G

REACTION CONDITIONS

FIG.24. CO hydrogenation at 673 K on 2% Ni/A1,0, with H,O in the reactant stream. See Table XI for description of conditions. Arrows give rate in water-free system (Ref. 204).

addition to H,S (see Fig. 24 and Table XI). However, the presence of water vapor alone reduced the rate only about fivefold. The reasons why preconditioning the catalyst in a reaction mixture results in a higher steady-state activity in the presence of H,S are not clear, although it was speculated to be due to inhibition of H,S adsorption by surface carbon formed during the pretreatment (204). This speculation, however, has not yet been confirmed TABLE XI Reaction Conditions at Steadv-State for Runs in Fig. 24" Reactant composition (atm) Reactions conditionsh

A B C D E F G 'I

T (K)

HZO

H 2s

H,

co

He

523 523 673 673 673 673 673

0 0.133 0.133 0.133 0.133 0 0

0 0 0

0.5938 0.515

0.1562 0.135 0.135 0.135 0.135 0.1562 0.5938

0.250 0.217 0.217 0.217 0.217 0.250 0.250

10-5

0.515 0.515

0 0 lo-s

0.515 0.5938 0.5938

From Ref. 204.

* The conditions were maintained for 24 hr before the steady-state reaction rate was determined.

194

C. H. BARTHOLOMEW ET AL.

(205).The deleterious effect of water vapor was speculated to be due to its inhibition of carbon formation freeing the metal surface for interaction by H,S. Thus, sulfur poisoning of nickel at high temperature (above 673 K) may be more representative of a carbon-fouled surface, whereas at low temperatures it may be more characteristic of the clean metal surface. Again, this needs to be confirmed by direct measurements of carbon and sulfur adsorption. For Ni/A1,03 and Ni/ZrO, the extent of sulfur deactivation was about fiftyfold at 673 K ; at 523 K the extent of deactivation was about 1000-fold. However, for Raney Ni the extent of sulfur deactivation was tenfold higher at 673 K than at 523 K ; this difference in behavior also needs confirmation and explanation. In addition to significantly lowering methanation activity, sulfur poisoning significantly affects product distribution. Dalla Betta et af. (203) found that methane yield was significantly decreased (e.g., from 100 to 58% for 2% Ni/Al,O, at 673 K) and Cz+ hydrocarbon yields were substantially increased (from 0 to 42%) after poisoning 24 hr with 10 ppm H,S. Based on these results, these workers concluded that sulfur poisons the ability of the surface to hydrogenate carbon atoms more severely than the ability to form carbon-carbon bonds. This selectivity shift is basically in agreement with the postulate of Herington and Woodward (201) that H,S adsorbs on sites TABLE XI1 Effect of Purtiul Presulfidiny (Fixed Bed) on Actiuity/Selectivify Properties of Ni/AI,O, and CoIAl,O, Catalysts at 500 K , 1 atmn.h

Sample 20% Co/AI,O,* Fresh Presulfidcd 23% Ni/A1,03g Fresh Presulfidcd

H* uptake ”/, CO (prnolig) conversion

43.5 34.5 283 1 13

Turnover number” x 103

% Yield‘ ___

(NCH4

(NCH,

CH,

CO,

Cz+

CO

CH,

fresh

5.2 4.3

84 76

2 5

14 19

3.7 3.9

3.1 2.9

0.94

7.2 3.1

78 85

5 5

17

4.3 4.6

3.4 3.3

0.97

10

)

poisoned”/ )

From Ref. 157. Reactant mixture of 1% CO, 4% H,, 95% N,. ‘ Percent of converted CO appearing as a CH,, CO,. or C , , product. Molecules converted or produced per catalytic site per second. Turnover number for poisoned catalyst based on H, adsorption after presulfiding divided by that for the fresh catalyst based on uptake before presulfiding. Sample in the form of 0.32-cm spheres Sample in powdered form. a

’

195

SULFUR POISONING OF METALS

where H, otherwise adsorbs, reducing hydrogenation activity more than the rate of chain growth. Based on the results of Dalla Betta and co-workers, it is clear that the steady-state activity of a completely sulfur-poisoned Ni or Ru methanation catalyst is 102-104 times lower than that of the fresh catalyst. However, a typical industrial methanation process would more probably involve a catalyst only partly poisoned by sulfur. Bartholomew and co-workers (23, 113, 157) attempted to assess how sulfur poisoning of only a portion of the catalyst would affect its activity/selectivity properties in fixed-bed and fluidized-bed reactors. Data in Table XI1 show the effects on specific activity and product distribution of partially presulfided Co/A120, and Ni/A120, catalysts in a fixed bed. Catalysts were presulfided with 10 ppm H2S at 725 K, and reaction was carried out with sulfur-free feedgas. Corresponding data are listed in Table XI11 for catalysts partially presulfided and then studied in a fluidized-bed reactor under the same conditions. The decrease in H, uptake TABLE XI11 Specific Activity Data" before and after Exposureb tn 10 p p m H,S of Alumina-Supported Ni, Co, and Ni Bimetallics in Powder Form'

Catalyst

CH4 turnover number

% co

% CH4

conversion

yieldd

x

103

(sec-')

Poisoned site activity ratio'

14% Ni

Fresh Poisoned 20% NiLCo (10% Ni, 10% C O ) Fresh Poisoned 20% c o Fresh Poisoned 30.7% Ni-Mo (10.8% Ni. 19.8% Mo) Fresh Poisoned

9.85 6.28

83.6 14.2

6.9 2.8

8.60 8.44

85.6 89.7

16.1 i 3.4

5.81 3.25

56.9 65.2

10.3

10.3 6.1

12 64

3.5

6.8 2. I

0.40

0.83

0.34

0.34

At 523 K , 140 kPa, and space velocities of about 40,000-100,000 hr-' in a gas mixture containing 1% CO, 4% H,. 95% N , . Exposure to 10 pprn H,S in a fluidized bed over a period of several hours sufficient to poison about 50% of the surface. ' From Ref. 157. Methane yield is the fraction of converted CO which is transformed to methane. Turnover number for the poisoned catalyst divided by that for the fresh catalyst.

196

C. H . BARTHOLOMEW ET AL.

is presumably a measure of the fraction of metal covered by sulfur. Much of the metal surface probably contained no sulfur and was responsible for the majority of the activity. For the supported Co and Ni catalysts in Table XI1 it appears that CO conversion is decreased roughly proportional to the decrease in H, uptake. Moreover, the turnover numbers and product distributions for the partly poisoned catalyst are not changed significantly within experimental accuracy (estimated at & 10%). These data suggest that in a fixed-bed reactor a portion of the bed is poisoned completely and the activity/selectivity properties of the unpoisoned portion are apparently unaffected. Unfortunately, such data provide no information on the relationship between the fraction of the surface covered with sulfur and catalytic activity, the kinetics of sulfur poisoning or the sulfur-poisoning mechanism. This illustrates the complexity of investigating sulfur poisoning in a fixed bed and the limitations on interpretation of such data. Furthermore, if the reaction is run at high conversion where diffusional limitations are important, the effects of poisoning will be completely masked until most of the bed has been poisoned and breakthrough occurs (113). This is evident from the work of Baird et al. (210) showing that the decrease in CO conversion from adding a given percentage of sulfur to a fixed bed of Raney nickel decreases as the size of the bed is increased, and no effect is observed if the bed is large enough. Under such conditions conclusions about sulfur tolerance of catalysts, etc., are totally erroneous. The data of Bartholomew et al. (113,157) in Table XI11 suggest that the sulfur is apparently more uniformly distributed in a fluidized bed, causing significant changes in activity and methane yield. A factor of two decrease in turnover number is observed for both Ni and Co. In the case of Ni-Co there is apparently a synergistic interaction resulting in greater resistance to sulfur poisoning, as the turnover number is only decreased by 20% due to sulfur poisoning. Bartholomew and co-workers also measured the loss of catalytic activity with time of Ni and Co bimetallics (157, 1 9 4 , Ni-molybdenum oxide (23, 113), and borided Ni and Co catalysts (161) during methanation in the presence of 10 ppm H,S. Typical activity versus time plots are shown in Figs. 25 and 26. Activity is defined as the ratio of the mass-based rate of methane production at any time t divided by the initial rate. The activitytime curves are generally characteristic of exponential decay; some catalysts decay more slowly than others, but all catalysts suffer at least two orders of magnitude loss in activity within a period of 100-150 hr. Accordingly, it does not appear that other metals or metal oxides in conjunction with Ni significantly change the sulfur tolerance defined in terms of steady-state activity of Ni. These materials can, however, influence the rate at which the

197

SULFUR POISONING OF METALS

10

__T_

I

I

I

I

I

09

\

08

5 >

0.7

-

+ 0.6

c3

w 0.5

N

2 0.4 -J

u

2003

A

0.2

,lo% Ni , IO%Co

0.I 0 0

10

20

30

40 TIME(hr1

50

60

70

FIG.25. Activity versus time for alumina-supported catalysts during methanation at 525 K, 1 atm in 10 ppm H,S (H,/CO = 99) (Ref. 194).

steady-state activity of a Ni catalyst is reached, some materials apparently prolonging the deactivation process. The relative rates of deactivation (i.e., sulfur resistances) for these different catalysts are discussed further in the following subsection. Although metals or even promoted metals have very low sulfur tolerances in synthesis reactions, other materials, such as metal oxides, nitrides, borides, and sulfides, may have greater tolerance to sulfur poisoning because of their potential ability to resist sulfidation (18).The extremely low steadystate activities of Co, Ni, and Ru metals in a sulfur-contaminated stream actually correspond to the activities of the sulfided metal surfaces. However, if more active sulfides could be found, their activity/selectivity properties would be presumably quite stable in a reducing, H,S-containing environment. This is, in fact, the basis for the recent development of “sulfur active” synthesis catalysts (211-215), which are reported to maintain stable activity/ selectivity properties in methanation and Fischer-Tropsch synthesis at H,S levels of 1% or greater. Happel and Hnatow (214), for example, reported in a recent patent that rare-earth and actinide-metal-promoted molybdenum oxide catalysts are reasonably active for methanation in the presence of 1-3% H,S. None of these patents, however, have reported intrinsic activities

198

C . H . BARTHOLOMEW ET AL.

SURFACE AREA NORMALIZED TIME (molecules H,S/Ni site) FIG.26. Activity versus surface area normalized time, for alumina-supported catalysts a t 525 K, 1 atm (10 ppm H,S, H,/CO = 99) (Ref. 194).

or steady-state activities in the presence of sulfur. In other words, the sulfur tolerance of these materials has not really been quantitatively demonstrated. The studies mentioned to this point provided measures of sulfur tolerance and effects of sulfur on selectivity of Fe, Co, Ni, and Ru in CO hydrogenation. Several recent studies (99-101, 147, 172, 193, 19.5, 205-209) provide additional insight into the mechanism of poisoning in CO hydrogenation on these metals. In a gravimetric study of CO adsorption/desorption during methanation, Bartholomew and Gardner (172) inferred that adsorption of CO occurred on sulfur-poisoned Ni/Al,O, at approximately the same rate and in about half the quantity observed for the fresh catalyst. The reactivity of surface species was drastically affected; instead of the principally hydrogenatable or volatile surface species observed on unpoisoned Ni/AI,O,, a substantial portion ( - 80%) of the surface species on the sulfur-poisoned catalyst could not be hydrogenated (see Table XIV). Bartholomew and Gardner postulated that sulfur poisons the dissociative adsorption of H, to a greater extent than CO, preventing the reaction of atomic hydrogen with the adsorbed carbon species formed upon dissociative adsorption of CO. Wentrcek et ul. (195) obtained similar results in pulse reactor experiments with fresh and sulfur-poisoned Ni/AI,O, . By pulsing CO over 15 mg of the

199

SULFUR POISONING OF METALS

TABLE XIV Gravimetric (Adsorption und Desorption) Measurements of'Surface Species Formed during Methanation on Fresh and Sulfur-Poisoned 14% Ni/Al,O, at 675 K"

Sample

Adsorbed species" (mg/g catalyst)

Percent reactive'

Percent inactived

Fresh 50% Poisoned 100% Poisoned

11.6 8.9 6.2

78 57

22 43 81

19

' From Ref. 172.

' CO adsorbed at 675 K, 1 atm, H,/CO

= 3.

' Percent of adsorbed species desorbed in N, and H, streams at 675 K, 1 atm. Presumably active C, CO, CH,, H,O. Percent not removable by either N, or H, at 675 K, 1 aim; presumably inactive carbon.

fresh catalyst at 553 K, they deposited 1.05 x mol of surface carbon (as indicated by CO, produced), all of which was easily removed by reaction with hydrogen pulses to form methane. Exposing the catalyst to progressively greater amounts of H,S caused the amount of carbon deposited to decrease somewhat and the reactivity of the deposited carbon to decrease substantially (Table XV). After exposure of the catalyst to sufficient H,S (3.2 x mol) to poison about 20-30% of the catalyst surface (sulfur adsorption capacity of 15 x mol), the amount of surface carbon deposited was 40% of that

TABLE XV Deactivation of Surface Curbon by Hydrogen Sulfide on 25% NiIAl,O, CaralystaSb

H,S exposure (mol x 10')

Deposited'

Convertedd

Fractional deactivation of surface carbon (%)

0 0.4 1.2 2.4 3.2

1.05 0.95 0.93 0.73 0.42

1.11 0.53 0.32 0.15 0.06

0 44 66 80 86

Surface carbon (mol x 10')

* From Ref. 195.

15.2 x g of NijA1,0, (25 wt. %). From pulse of CO (4.86 x lo-' mol) at 553 K. Surface carbon converted to CH, by exposure to hydrogen pulses.

200

C. H. BARTHOLOMEW ET AL.

observed for the fresh catalyst; only 14% of this carbon could be removed with H, . After preexposure to saturation sulfur coverage, no carbon deposition was detected. These results suggest that sulfur not only poisons the dissociative adsorption of CO but also prevents hydrogenation of surface carbon to methane. Wentrcek et al. suggested on the basis of their pulse reaction and LEED data that adsorption of sulfur causes a transformation of the reactive surface carbon to a less reactive polymeric carbon form. Rostrup-Nielsen and Pedersen (209) recently studied sulfur poisoning of supported nickel catalysts in both methanation and Boudouard reactions by means of gravimetric and differential packed-bed reactor experiments. In their gravimetric experiments a synthesis mixture (H,/CO/He = 5/7/3) containing 1-2 ppm H2S was passed over a catalyst pellet of 13% Ni/A1,03MgO at 673 K and 1 atm. The rates of Boudouard and methanation reactions were determined from weight increases and exit methane concentrations respectively. In the presence of 2 ppm H,S a factor of 20 decrease was observed in both methanation and Boudouard rates over a period of 30-60 min. However, the selectivity or ratio of the rates for Boudouard and methanation reactions was constant with time. From these results the authors concluded that the methanation and Boudouard reactions involve the same intermediate, carbon, and that sulfur blocks the sites for the formation of this intermediate. A number of important differences are apparent from comparing the results of Rostrup-Nielsen and Pedersen (209), Bartholomew and Gardner (172, 216), and Wentrcek et al. (195). Bartholomew and Gardner (172) and Wentrcek et al. (195) report a moderate decrease in the amount of CO adsorbed and/or carbon formed and a relatively large decrease in methane production with increasing sulfur content, i.e., a wide variation in selectivity with increasing sulfur content, whereas the data of Rostrup-Nielsen and Pedersen (209)suggest a constant selectivity with sulfur coverage. Bartholomew and Gardner observed a measureable quantity of dissociative CO adsorption on completely sulfided Ni/A1,03. Moreover, Gardner and Bartholomew (216) observed approximately the same rates of carbon formation in the Boudouard reaction on fresh and presulfided 14% Ni/AI,O, over a wide range of temperatures. Nevertheless, Wentrcek et al. (195) observed no carbon formation on their completely presulfided Ni/AI,O, . These apparent differences may be explained by differences in experimental conditions, differences in catalyst composition, and by transport limitations in the work of Rostrup-Nielsen and Pedersen (209). For example, the unsteady-state experiments of Bartholomew and Gardner (172) and of Wentrcek et al. (195) were conducted on presulfided catalysts in powder from at 575 K (H,/CO = 3), whereas the steady-state experiments of Rostrup-Nielsen and Pedersen were carried out in situ at 675 K (H,/CO =

SULFUR POISONING OF METALS

20 1

0.7) using a catalyst pellet. Bartholomew et al. (140, 194) have shown that the rate and extent of poisoning of nickel catalysts in methanation are affected by temperature and gas-phase composition, particularly H2/CO ratio. In the work of Rostrup-Nielsen and Pedersen there was undoubtedly strong, nonuniform eggshell deposition (113, 194) of sulfur as well as mass and heat transport limitations resulting from the rapid, exothermic methanation reaction in the catalyst pellet. The reaction was probably limited mainly to the exterior of the pellet where the deposited sulfur was concentrated. Under these high-temperature conditions (675 K), carbon deposition itself may have caused a significant loss of methanation activity (216-219), although there was no evidence of such deactivation during the first 25 min of reaction in the absence of sulfur. Thus the gravimetric results of RostrupNielsen and Pedersen do not necessarily reflect the effects of sulfur poisoning alone, since they may have been coupled with other deactivation phenomena and heat/mass transport influences. It is also not clear how the unsteady-state results of Bartholomew and Gardner and Wentrcek et al. apply to the steadystate activity/selectivity properties of nickel in the presence of sulfur. Differential reaction studies of presulfided catalysts reported by RostrupNielsen and Pederson (209) were conducted in the apparent absence of such complications as transport and carbon deposition. The sulfidation of 25% Ni/Al,O, was carried out at 575,775, or 1075 K with H2S/H2varied between 8.4 x to obtain a broad range of sulfur coverages. Their and 5.7 x results, plotted as the log of normalized activity versus 1 - S/So where S and Soare sulfur coverage and saturation sulfur coverage, respectively, show that activity decreases very significantly with increasing sulfur coverage (Fig. 27). The slope of the straight-line fit of the data of -4.3 was interpreted by the authors to suggest that an ensemble of four nickel atoms is required for reaction. They also concluded that since the strong, nonlinear deactivation was accompanied by no change in activation energy, sulfur poisoning in methanation involves site blockage. Most of the previous studies mentioned to this point did not meet the experimental requirements outlined in Section V,A and involved H2S concentrations in the ppm range. In view of the previously discussed adsorption studies showing that reversible adsorption of H2S occurs only at ppb levels under typical methanation conditions, it is reasonable to expect that full saturation coverage of sulfur occurred on the catalysts used in these previous studies; hence their steady-state activities should not vary significantly for a given metal such as Ni. To allow quantitative measurements of the rates of catalyst deactivation, of the steady-state activity in the presence of ppb concentrations of H2S, and to determine the catalytic activity as a function of surface coverage by sulfur, Katzer and co-workers (99-101,147, 205-208), used a reactor and catalyst configuration which satisfied all the

202

C. H. BARTHOLOMEW ET AL.

I 0.001

1 0.5

\" 0.2 l-S/SO

FIG.27. Methanation activity of presulfided catalysts at 300"C, 1 atm, 0.05 g catalyst (0.3-0.5 mm). Sulfidation temperature: filled circles, 800°C; half-filled circles, 500°C; unfilled circles, 300°C (Ref. 209).

requirements of Section V,A. An all-quartz gradientless internal recycle reactor (99, 196, 197) was used to study sulfur poisoning of Ni (99, 100), Co (101, 147, 205, 206), Ru (207), and Fe (101, 208) at H,S concentrations as low as 10 ppb. The catalysts used consisted of annular cylinders or flat plates of fused A1,0, with metal deposited on the external surface by vacuum evaporation or by impregnation and spherical beads with metal deposited only on an external peripheral shell less than 5 pm thick. Operation of the system was such that intrinsic kinetic and rate data were obtained, unaffected by heat and mass transfer effects. Sulfur distributions were uniform. Measurements were made both in the absence and presence of sulfur. The turnover numbers and activation energies for methanation over all four metals freshly reduced and under sulfur-free conditions (Table XVI) were in good agreement with values reported for supported metal catalysts (220).At 673 K, Ni and Ru exhibited only very slow losses in activity apparently due to slow carbon deposition, whereas Co and Fe underwent rapid, severe carbon deactivation after maintaining their fresh catalyst activity for a few hours. After rapid deactivation the final steady-state activity, which was about 100-fold lower than the activity of the fresh catalyst, was approached slowly; this activity region was referred to as the lower pseudosteady state. Likewise, the fresh catalyst behavior was referred to as the upper pseudo-

203

SULFUR POISONING OF METALS

TABLE XVI Methanation Activity and Activation Energy for Methanation over Ni, Co, Fe, and Ru" Methanation turnover numbers (sec- ') 548 K

'I

673 K

Catalyst

1%CO in H,

4XCO in H,

I%CO in H,

4%CO in H,

Ni/AI,O, CO/AI,O,~ Fe/AI,O, Ru/AI,O,

0.012 0.10 0.047 0.020

0.22 0.05

0.70 15.0 2.7 4.5

1.30 -

-

0.005

Activation energy (kJ/mol) 100 1 I7 100 1 I3

1.o

From Refs. 99-101 and 147.

* The turnover numbers are for the upper pseudo-steady-state region TABLE XVII Methanation Activation Energies over Ni, Co, Fe, and Ru Catalysts'

Catalyst Ni/AI,O, Co/AI,O, Fe/AI,O, Ru/AI,O, a

Fresh catalyst

Aged catalyst

Sulfurpoisoned catalyst

100 1 I7 100 113

I00 67 75 113

100 67 113

From Refs. 99-101 and 147. Energies in kJ/mol.

steady state. AES analysis of Fe and Co deactivated to the lower pseudosteady state showed that activity loss was due to bulk carburization of the metal and buildup of graphitic deposits. AES analysis of aged Ni and Ru showed no bulk carburization; graphitic carbon was present on the surface of aged Ni, whereas aged Ru contained no surface carbon. Activation energies for methanation over Ni and Ru were the same for both the fresh and the aged catalysts (Table XVII). In contrast, activation energies for methanation over aged Co and Fe were lower by 50 and 25 kJ/mol, respectively as compared to fresh Co and Fe (Table XVII). In the case of Co the CO partial pressure dependence changed from a negative order in the upper pseudosteady state to a positive-order dependence in the lower pseudosteady state (Table XVIII). The dependence on H, partial pressure was positive one-half order in both upper and lower pseudosteady states. For Fe the kinetic behavior was not investigated.

204

C. H. BARTHOLOMEW ET AL.

TABLE XVlll Kinetic Behavior of Co/AI,O, in D?fferent Regions ?f' Methunutioii Activity" ~~~~~-

NCH,at 673 K E,

~~

Carbon deactivated

Sulfur poisoned

15.0 sec-l 117 kJ/mol

0.IOsec I 67 kJ/mol

0.001 sec67 kJ;mol

-0.25 0.5 "

~~~

Fresh catalyst

0.3 1.0

0.3-1.0

0.5

-

'

From Refs. 101, 147.

Having established the kinetic behavior of Ni, Co, Fe, and Ru under sulfur-free conditions, observed differences could then be attributed to sulfur deactivation. To further clarify the role of carbon and sulfur in deactivation, two different methods of sulfur poisoning were employed : (1) in situ deactivation in which H,S, CO, and H, were fed to the freshly reduced catalyst directly, and the transient activity behavior was followed ;

Z(o-41 0

I

I I I I I I I i I i I i J 1000 2000 3000 4000 5000 6000 7000 REACTION TIME (rnin)

FIG.28. In siru poisoning of Ni/AI,O, plate catalyst. Reaction conditions: 100 kPa, 661 K, 13 ppb H,S, (2) 62 ppb H,S, (3) 95 ppb H,S (Refs. 99, 100).

42,C0/96%, H,; (1)

205

SULFUR POISONING OF METALS

and (2) presulfiding with an H,S/H, mixture reaction temperature until equilibrium sulfur coverage was established prior to introducing CO into the feed gas and determining activity-versus-time behavior. The transient methanation activity behavior of Ni, Co, Fe, and Ru during in situ sulfur poisoning was more or less identical in all cases. Hence, only the results for Ni will be discussed in detail; those for Co, Fe, and Ru will be discussed only where significant differences were observed. Figures 28 and 29 show the transient methanation activity of a Ni/AI,O, flat-plate catalyst and the gas-phase H,S concentration profile, respectively. The presence of just 13-ppb H,S caused about a 200-fold loss in steady-state methanation activity. Increasing the H,S level to 62 ppb resulted in an additional tenfold activity loss; an increase to 95 ppb lowered the activity further. However, increasing the H,S level above 95 ppm did not cause a significant additional decrease in activity (Fig. 30) and decreasing the H,S level from 95 to about 15 ppb reversibly restored the activity level originally observed at this latter concentration level, thereby demonstrating that sulfur adsorption and poisoning by sulfur are reversible, and that a truly dynamic

0

1000

2000

3000

4000

5000

6000

7000

REACTION TIME (mid

FIG.29. Transient H,S concentration profiles during in siru poisoning of Ni/A1,0,. Reaction conditions are the same as those in Fig. 28 (Refs. 99, 100).

206

C . H. BARTHOLOMEW ET A1

0

3000

6000

9000

12,000

REACTION TIME ( m i d FIG.30. In situ poisoning of Ni/AI,O, pellet. Reaction conditions: 100 kPa, 661 K, 4% C0/96%H,; (1) 12 ppb H2S, (2) 52 ppb H,S, (3) 73 ppb H2S, (4) 93 ppb H2S, ( 5 ) 15 ppb H2S (Refs. 99, 100).

equilibrium exists between the gas-phase H,S and the catalyst surface. To our knowledge this is the first time such a dynamic adsorption equilibrium under reaction conditions has been shown. Comparison of the methanation activity in Fig. 28 with the H2S concentration transients in Fig. 29 indicates a prolonged adsorption of H,S during the 13-ppb deactivation with the H,S level reaching a steady state at about the same time as the methanation activity. The difference between the sulfur fed to the reactor and that in the effluent represents the sulfur that was adsorbed on the Ni surface, since the alumina support and the quartz reactor did not adsorb sulfur. Time integration of the amount of sulfur adsorbed, obtained from the effluent H,S concentration data, gives the total amount of sulfur adsorbed at equilibrium with the gas-phase H2S concentration. At 661 K and 13-ppb H2S in H,, sufficient sulfur was adsorbed to give a sulfur-to-surface-Ni atom ratio of 0.5, where the number of surface Ni atoms was determined by H, chemisorption. This S/Ni, ratio corresponds to saturation coverage of single-crystal Ni with sulfur (Section 111). Figure 29 also shows that increasing the H,S level above 13 ppb results in little, if any, further adsorption of H,S by Ni, although the methanation activity continues to decline (Fig. 28). Apparently, the amount and rate of sulfur adsorption on a Ni surface equilibrated with 13-ppb H,S in H, is too

SULFUR POISONING OF METALS

207

small to be detectable in the H2S concentration profiles, consistent with the observation above that already at 13-ppb H2S, the Ni surface contained a saturation layer of sulfur (measurable to within i-5%). AES studies showed the presence of a saturated sulfur layer on all metals thus confirming the results of the sulfur adsorption studies in the reactor; the ratio of the sulfurto-metal peak remained constant (within $.5%) even though H2S concentrations were varied between 13 and 100 ppb during different runs, thereby indicating no significant further adsorption. This indicates that the kinetic probe is more sensitive than the surface probe. It should be emphasized that the surface probe (AES) showed that sulfur was present only on the surface and not in the subsurface regions. AES studies of in situ sulfur-poisoned Ni and Ru showed the complete absence of carbon on the surface or in the near-surface region; for Co, on the other hand, a fraction of a monolayer of carbon was present but there was no bulk carburization. For Co which was carbon deactivated prior to sulfur poisoning, surface carbon and bulk carburization were both observed (205). The presulfiding pretreatment involved contacting the catalyst with an H2S/H, mixture (without CO) at 661 K until the feed and effluent H,S concentrations were equal, indicating adsorption equilibrium, following which CO was introduced. Figure 31 shows typical activity-versus-time

Iw Z-

200 400 600 800 1000 REACTION TIME (rnin) FIG.31. Transient methanation activity behavior of prepoisoned and sulfur-saturated Ni/A1,0,. Prepoisoning conditions: 100 kPa, 661 K, 13 ppb H,S in H, for 50 hr. Reaction conditions: 100 kPa, 661 K , 4% C0/96% H, with 13 ppb H,S (Refs. 99, 100).

0

208

C. H. BARTHOLOMEW ET AL.

behavior for presulfided Ni/A1,0, upon introduction of CO. The initial activity was reduced over 100-fold by presulfiding at 13-ppb H2S, and it remained essentially unchanged over the next 15 hr. This activity is essentially the same as that for in situ poisoning runs at the same H,S concentration. The activity behavior of presulfided Co, Fe, and Ru was similar to that of Ni, and the activity reductions were the same as those observed during iiz situ poisoning. AES showed that with the exception of Fe, no carbon was observed on presulfided catalysts even after 700 hr of steady-state operation with a sulfur-containing feed. These results strongly suggest that sulfur poisoning is due primarily to blockage of active metal sites by adsorbed sulfur. Figure 32 shows the steady-state methanation activity of Ni, Co, Fe, and Ru relative to the fresh, unpoisoned surface activity as a function of gasphase H,S concentration. These data indicate that Ni, Co, Fe, and Ru are all similarly extremely sensitive to poisoning by very low H,S concentrations (15-100 ppb). These results predict that, in principle, poisoning studies of

3

10-5

2

0

30 45 60 75 90 105 H2S CONCENTRATION (ppb)

15

Co (A), Fe ([I), FIG.32. Relative steady-state methanation activity profiles for Ni ,).( and Ru (0) as a function of gas-phase H,S concentration. Reaction conditions: 100 kPa, 663 K, 1% C0/99%,H, for Co, Fe, and Ru; 100 kPa, 661 K, 4:4, C0/96x, H, for Ni (Ref. 101).

209

SULFUR POISONING OF METALS

these metals at H2S levels in the ppm range should show no sensitivity to H2S concentration since saturation and complete activity loss are reached at ppb levels. Nevertheless, this may not be true in practice, as Erekson and Bartholomew (140) report changes in poisoning rates and adsorption stoichiometries with variations in PHlS/PH2 values from 0.2 to 10 ppm, possibly as a result of surface restructuring. From Fig. 29 it is possible to obtain the fractional sulfur coverage relative to saturation coverage as a function of time by piecewise integration of the sulfur-adsorbed-versus-time data. Because sulfur distribution was uniform and the rate of poisoning relative to the reaction time constant was slow, a direct relationship between the methanation rate and the fractional sulfur coverage could be established. Figure 33 shows that the nickel-catalyzed methanation rate (relative to fresh catalyst) is directly proportional to the square of the fraction of the unpoisoned surface, (1 - 6j6,)'. Similar behavior was observed for Co. These results suggest that (1) the poisoning by

SQUARE

OF THE UNSULFIDED SURFACE FRACTION

Methanation turnover number for Ni/AI,O, as a function of (1 conditions: 100 kPa, 661 K , 404 C0/96% H 2 (Refs. 99, 100). FIG. 33.

-

0,)'. Reaction

210

C. H . BARTHOLOMEW ET AL.

sulfur is primarily a geometric effect in which sulfur blocks two active metal sites per adsorbed sulfur atom, thus making them inaccessible for the reaction; and (2) the rate-controlling step in the methanation requires two surface sites. That the activation energies and rate concentration dependences for methanation over sulfur-poisoned Ni and Ru are the same as those observed under sulfur-free conditions (99-100,140,147,209) provides further evidence that sulfur poisoning of Ni and Ru involves geometric blockage of active metal sites. However, the observations that (i) the activation energies for methanation over sulfur-poisoned and carbon-deactivated Co are both 50 kJ/mol lower than that for fresh Co and (ii) the rate dependence on CO partial pressure is positive order for both carbon-deactivated Co and sulfur-poisoned Co, suggest that the deactivations of Co by sulfur and carbon are similar and may involve electronic as well as geometric effects. The reaction studies discussed to this point involved supported nickel catalysts. There are few examples in the literature of reaction/poisoning studies involving single-crystal surfaces and only one thus far involving CO hydrogenation on nickel. In their investigation of the effects of sulfur on the methanation activity of Ni( IOO), Goodman and Kiskinova (193) observed 2.5 orders of magnitude loss of activity at Qs < 0.2 with no further decrease in activity up to saturation coverage (0, = 0.5). From this they concluded that a p ( 2 x 2) S structure deactivates the nickel surface for methane production. At temperatures below 600 K the activation energies for the fresh and poisoned surface were the same within experimental error, in agreement with studies of supported and polycrystalline nickel (99, 140, 209). At temperatures greater than 600 K, carbon deposition was found to lower activity at Q, = 0.05; however, no surface carbon was observed at saturation sulfur coverage (0, = 0.5). Goodman and Kiskinova found a correlation between loss of methanation activity and the loss of adsorption capacity for CO and H, with increasing sulfur coverage. Their observed large decrease in CO and H, adsorption with increasing sulfur coverage at 0, < 0.2 was consistent with the deactivation of 4-10 nickel sites per adsorbed sulfur atom, which they interpreted as an electronic effect. To summarize this discussion of sulfur poisoning in CO hydrogenation, several recent studies provide limited fundamental information regarding the poisoning of transition metals. Much of the previously reported data were masked by transport effects and nonuniform sulfur poisoning. However, all available data suggest that transition metals are extremely sensitive to sulfur poisoning. The more quantitative studies show that severe deactivation due to formation of monolayer surface metal sulfides occurs at ppb H2S concentrations, levels of H2S orders of magnitude lower than those required to

SULFUR POISONING OF METALS

21 1

form stable bulk sulfides. For Ni there is general consensus that the mechanism of sulfur poisoning involves site blockage due to sulfur adsorption and that electronic effects are not important at high sulfur coverages typical of almost all reaction systems. There is evidence that electronic factors may also play a role in sulfur poisoning of metals including Ni, Co, Fe, especially at low coverages, although site blockage by sulfur is still the more important factor. Although carbon deactivation can be an important factor under sulfur-free reaction conditions, its role in sulfur poisoning may vary depending upon the metal and poisoning conditions, and requires further clarification. Sulfur is known to affect the selectivity in Fischer-Tropsch synthesis; therefore, poisoning studies may be helpful in elucidating the reaction mechanism and in controlling selectivity. A more fundamental understanding of sulfur poisoning and of sulfur interactions with metals can ultimately provide the key to development of catalysts with significantly improved sulfur tolerance, although some improvements will undoubtedly continue to come through trial-and-error empiricism.

2 . Deactivation Rutes Loss of catalyst activity due to poisoning is an important consideration in reactor design. Because of the extreme sensitivity of methanation and hydrocarbon synthesis catalysts to poisoning by sulfur compounds, it is also essential to obtain accurate deactivation rates or, alternatively, deactivation models in order to estimate accurately catalyst life in a given process. If the rates are determined under well-defined conditions (i.e., in the absence of heat and mass transport influences and concentration gradients), they also provide a basis for better understanding the mechanisms of poisoning. These objectives have apparently stimulated several recent investigations of sulfur poisoning in CO hydrogenation (23, 99-101, 113, 194, 195, 210, 221-224). Several of these studies were not definitive because of heat/mass transport influences, concentration gradients, and reactor contamination (210, 223, 224). Some of the results have fairly specific applications and serve more than anything to illustrate the complexity of modeling the poisoning process in an industrial reactor (221,222);other studies have fairly broad applications and will serve in this review as examples of how the subject of modeling sulfur poisoning might be approached (23,99-101,113,194,195). It should become apparent t o the reader in the ensuing discussion that the modeling of sulfur poisoning in high-surface-area, porous catalysts, especially in a fixed bed, is very difficult because most sulfur compounds (e.g., H2S) adsorb strongly and nonuniformly at the entrance to the bed and on the exterior shell of a catalyst particle. One can avoid these problems in basic studies by using a mixed-flow

212

C. H . BARTHOLOMEW ET A L .

reactor and nonporous catalysts as in the work of Katzer and co-workers discussed above ; however, in typical industrial processes these complexities usually have to be dealt with directly. It should be emphasized here that even though sulfur tolerances (steadystate activities in the presence of sulfur impurities), of promoted and unpromoted nickel catalysts are extremely low, their su/fur resistances (rates of activity loss in the presence of sulfur impurities), can vary greatly with catalyst configuration, composition, and support. Thus it may be possible to extend significantly nickel catalyst life in the presence of sulfur poisons through addition of promoters or use of novel supports (23,99,100,113,114, 161, 194, 225-227). This is in fact the basis of two recent patents (225,226). Rates of deactivation of Ni and Ni bimetallic catalysts as a result of poisoning by 10-ppm H,S during methanation were investigated in a series of studies by Bartholomew and co-workers (23, 113, 161, 194). Effects of catalyst composition and geometry, gas composition and reaction temperature on the rate of deactivation were considered. Deactivation rates were found to be relatively insensitive to temperature and quite sensitive to gas and catalyst composition (194). In fact, the rates of deactivation were 2-3 times more rapid in a H,-rich mixture (H, j C 0 = 99), compared to a normal synthesis ( H 2 / C 0 = 3-4) mixture. Activity-versus-time curves shown in Fig. 25 for alumina-supported Ni and Ni bimetallic catalysts show two significant facts: (1) the exponential decay for each of the curves is characteristic of nonuniform pore-mouth poisoning, and (2) the rate at which activity declines varies considerably with metal loading, surface area, and composition. Because of large differences in metal surface area (i.e., sulfur capacity), catalysts cannot be compared directly unless these differences are taken into account. There are basically two ways to do this ;(1) for monometallic catalysts normalize time in terms of sulfur coverage or the number of H,S molecules passed over the catalysts per active metal site (161,194), and (2) for mono- or bimetallic catalysts compare values of the deactivation rate constant calculated from a poisoning model (113, 195). Figure 26 shows activity data from Fig. 25 plotted versus surface-area normalized time (the number of H2S molecules passed through the sample per initial number of metal sites). The deactivation curves in Fig. 26 for 3 and 20% Ni and for 20% Co and are very nearly coincidental. The curve for Ni-Co shows slightly higher activity compared to Ni, suggesting a synergestic effect. The activity of 20% Ni-MoO,/Al,O,, however, is significantly higher than Ni over a wide range of surface-area normalized time. In the case of Ni-Moo,, the Mo acts as a sulfur sink, thereby reducing the loss of activity due to adsorption of sulfur on Ni sites (23, 113). The adsorption of sulfur on a metallic catalyst can be modeled as irrevers-

SULFUR POISONING OF METALS

213

ible adsorption or as an ion exchange process, for which models have already been developed (228-232). For uniform plug flow the continuity equation for transport of contaminant species through a catalyst bed is an/&

+ V(an/i?x) = -kns

(8)

where n is the gas-phase concentration of poison, V is the molar average velocity, x is the distance along the bed, s is the density of sites which can adsorb sulfur, and k is the rate constant for sulfur adsorption. A balance on the site density is given by ds/dt

=

-kns

(9)

Solutions to these equations presented by Wise et al. (195, 233) for conditions where pore diffusion and external mass transfer resistance can be neglected are ln[n,/n, - 11 = In[exp(ks,l/V) - I ] - kni[t - ( L / V ) ]

(10)

ln[so/s - 11 = In[exp(kni(t - L / V ) ) - 11 - ks,(L/V)

(11)

where n, and nL are the inlet and outlet concentrations of poison and so and s are the total number of sites initially and at any given time, respectively. From experimental measurements of the poison concentration at the outlet of the catalyst bed (i.e., the breakthrough curve), it is possible to obtain values of k from a plot of In(n,/n, - 1) against ( t - V / L ) .The slope of the resulting line is n,k and the intercept is a function of so. Diffusional resistances, both external and pore diffusion, can be included by adding the appropriate terms. The resultant system of equations is soluble only by numerical techniques. However, pore diffusional limitations vanish at low Thiele modulus and external mass transfer limitations are reduced by high Reynolds number. This means that it may be possible to apply the model to very small particles of a porous catalyst at low poison concentrations, although in general it applies strictly to nonporous catalyst particles only where film diffusion resistance of the poison may be neglected. The application of the model of Wise et a/. (195, 233) to determining deactivation rate constants encounters rather serious experimental limitations in that at low poison concentrations and space velocities, the breakthrough curves are very slow to appear, and the accurate measurement of sulfur concentrations in the ppm or ppb range is difficult. If activity decline of the catalyst for the reaction of interest could be related to the loss of sites by poisoning, a more direct measurement of deactivation rate would be realized. Bartholomew and co-workers (113, 140, 161) extended this model by expressing the rate of deactivation in terms of normalized activity a : daldt

=

-kna

(12)

214

C. H . BARTHOLOMEW ET AL

by assuming that catalytic activity was proportional to the number of remaining sites, a = s/so. This relationship would only be true under very specific reaction conditions and concentrations of H,S (140). Upon substitution of a = s/so into Eq. (1 1) and assuming that exp[kni(t - xi?')] >> 1, a simplified relationship between activity and time is obtained : ln[(1 - a)/a] = knit - (kni

+ ks,)(L/V)

(1 3) According to this relationship a plot of ln[(1 - a)/a] versus t should result in a straight line with slope n,k which enables calculation of deactivation rate constants from activity-versus-time curves like those in Figs. 25 and 26 if activity is proportional to the number of unpoisoned sites. Deactivation parameters obtained by plotting ln[(1 - a)/a)] versus time are listed in Table XIX for a number of nickel and nickel bimetallic catalysts. The fact that these plots were generally linear confirms that these data are fitted well by this deactivation model. These data, which include initial site densities for sulfur adsorption, deactivation rate constants, and breakthrough times for poisoning by 1-ppm H,S at a space velocity of 3000 hr- ' provide meaningful comparisons of sulfur resistance and catalyst life for both unsupported and supported catalysts. Table XIX shows that the TABLE XIX Deuctivation Rate Constants j o r H , S Poisoning''

Catalyst H,/CO = 99 Pure Ni Raney Ni Ni-B Ni-Ru-B HJCO = 4 14% Ni/AI,O, 25% Ni/l% Ir/Al,O,g 10% Ni/20% MoO,/AI,O,

H, uptakeb So' (pmolig) (pmolig)

Deactivation' rate constant Catalyst k x lo3 life' S / N i d (hr- ppm- 1 (days)

'

'

5 267 103 188

6.2 2150 2660 3745

0.62 4.02 12.9 9.96

60 12 9.5 9.0

2.6 235 247 232

182

213 216 693

0.59

5 24 2.3

89 53 289

72

-

4.8

From Refs. 113 and 161. Measured at 298 K . Site density for sulfur adsorption from Wise-Bartholomew model. Sulfur molecules adsorbed per nickel surface site measured by H, adsorption, calculated assuming that 1 ppm = (1.5-2.7) x I O l 9 sites/g depending upon the bulk density. ' Determined from Wise-Bartholomew model; du/dt = - kun at 525 K. 130 kPa, and 10 ppm H,S. In 1 ppm H,S at 3000 h r - ' . Ref. 195. "

215

SULFUR POISONING OF METALS

catalysts with the largest site densities and smallest deactivation rate constants are longest lived. For example, the largest site density and the smallest deactivation rate constant are observed for the 10% Ni/20"/, Mo/Al,O,, having a breakthrough time of 289 days for a feed containing 1 ppm H,S, about a factor of three larger than the Ni/Al,O,. Comparison of the site densities from Table XIX with metal areas determined from H, adsorption provides important insights into the nature of H,S adsorption on these catalysts. For example, the sulfur site density of 213 pmol/g compared to the metal site density of 182 pmol/g (from H, adsorption) for 14% Ni/AI,O, is equivalent to S/Ni, = 0.6, in reasonable agreement with the earlier discussed studies (Section II1,C) which show values of 0.5-0.8 and consistent with the value 0.6 determined for pure unsupported Ni. However, in the case of a typical molybdenum-containing catalyst, e.g., 10% Ni/20% Mo/Al,O,, the sulfur site density and H, uptake are 693 and 72 pmol/g, respectively (S/Ni, = 4.8), providing evidence that a considerable amount of sulfur adsorbs on molybdenum oxide sites which do not adsorb H,; a similar behavior is also observed for Raney Ni and nickel-boride catalysts. The above studies show that oxide and metal promoters influence the rate at which nickel deactivates in the presence of sulfur. Metal-support inter-

10-4

0

1

I

1

1

2000 4000 6000 REACTION T I M E (min)

8000

FIG.34. Sulfur deactivation of Ni/AI,O, (pellets) (0) and Ni/ZrO, (pellets) (0) catalysts in the presence of low concentrations of H,S; T = 388"C, 4% CO, 96% H,,PHZs = 55 ppb at steady state (Ref. 99).

216

C. H. BARTHOLOMEW ET AL.

actions may also significantly influence the poisoning rate of nickel as indicated by the activity versus time data in Fig. 34; Ni/ZrO, maintains its activity for methane production over a significantly longer period of time than Ni/Al,O, (100). Since H,S was shown not to adsorb on the ZrO, support (ZOO), this is apparently an effect of an interaction between nickel and ZrO, affecting the adsorption of H,S. In view of these results it would be interesting to examine the sulfur resistance of nickel on other supports such as TiO, and MgO and as a function of metal crystallite size.

C. STEAM REFORMING Conventional nickel reforming catalysts are very sensitive to poisoning by sulfur compounds, suffering significant losses in activity at concentrations greater than 1-20 mg S/m3 (0.5 to 15 ppm by volume) (198,234-238). Since organic sulfides are present in conventional reforming feedstocks such as natural gas and naptha at levels as high as 300-500 ppm and at much higher levels in heavy oils and resids under consideration as potential feedstocks, desulfurization of the feed is essential prior to reaction over conventional catalysts. Organic sulfur compounds remaining in the feed after desulfurization are readily hydrogenated to H,S under typical reforming conditions; thus it is sufficient to consider poisoning by H,S. Adsorption studies of H,S on nickel (summarized in Fig. 11) indicate increasing reversibility of adsorbed sulfur with increasing temperature. Since the reforming process is carried out at high temperatures (6751125 K), poisoning by sulfur should be somewhat reversible at ppm levels; indeed the reversibility was demonstrated by Morita and Inoue (234, 235), who found that nickel catalysts regained initial activity after removal of sulfur compounds in the feed. At any given temperature, there was a threshold concentration below which no detectable poisoning was observed. Sulfur threshold concentrations determined by Morita and Inoue (234, 235) and other workers (198, 236) for outlet temperatures of 1075- 1175 K are listed in Table XX and compared with equilibrium adsorption data from Fig. 11 for 0, = 0.5. The agreement among experimental values and of experimental values with predicted values from Fig. 11 at the same temperature is generally good, suggesting that the reversibility of sulfur poisoning during reforming is a predictable consequence of equilibrium adsorption of H,S. Accordingly, the poisoning effect of sulfur in reforming may be ascribed to blocking of the nickel surface. This hypothesis is confirmed by the data of Rostrup-Nielsen (237, 238) listed in Table XXI,showing the effects of sulfur poisoning on the specific activity of 25 wt. % Ni/MgOAl,O, in steam reforming of ethane at 775 K.

217

SULFUR POISONING OF METALS

TABLE XX Suljiu TCIreshold Lrurls" in Swum Reforming Catalyst exit temp. (K) Catalyst

1075

14-25% Ni (ICI) 4.404 Ni 26% Ni Various Ni catalysts

0.7 1.4 3 7

3

1125

1175

Ref.

18

236 234,235

3.5 ~

~

11 14

-

21

198

7

16

Fig. 11

Comments Full-scale reforming Isothermal bed Experimental concentrations based on equilibrium gas composition for CO + 3H, = CH, + H,O PH2s/PH2 values from Fig. 11 based on extrapolation of equilibrium adsorption data at 8, = 0.5 for temperature shown

~

* Threshold concentrations below which there is no measureable loss of activity at the temperature shown in ppm (VHIS/VCH4);0.7 ppm ( V / V ) = 1 mg S/m3 (STP).

TABLE XXI Influence o f Suljur Poisoning on Spec& Acticity in Strum RrJbrming of'Ethunr on 25% Ni/AI,O, MgOa.b Sulfur content (wt. PPm) 80 239 360 398 61 5 805 'I

"

Sulfur coverage <0.1

0.30 0.45 0.49 0.76 1 .oo

Reaction rate (mol/g hr) x 10

Reaction rate (mol/m* Ni hr) x lo3

2.41 0.66 0.53 0.59 0.38 <0.01

120 62 69 64 56 -

773 K, 31 atm, H,O/C = 4. Data from Refs. 237 and 238

The fact that the specific activities based on remaining nickel surface area are reasonably constant over a wide range of sulfur coverage is evidence that chemisorbed sulfur poisons by blocking the metal surface for adsorption of reactants. At a sulfur coverage of 1.0, the rate is lowered by more than two orders of magnitude. Thus the actual tolerance of conventional nickel catalysts to sulfur poisoning during steam reforming at 775 K is very low.

218

C. H . BARTHOLOMEW ET AL.

Although the above-discussed studies have defined sulfur-poisoning tolerances for conventional nickel catalysts used in steam reforming of natural gas and naptha, they have not considered in sufficient detail the kinetics of poisoning at above-threshold concentrations nor the effects of catalyst and/or gas compositions on rate of deactivation and tolerance level. Nor is there any previous report on the effects of sulfur on product distribution (i.e., relative rates of production of H,, CO, CH,) in steam reforming of hydrocarbons. A recently announced Japanese process (239)for high-temperature steam reforming of high-sulfur, heavy-hydrocarbon feedstocks (such as resids) based on Ni/Ca aluminate catalysts demonstrates the potential for developing sulfur-tolerant catalysts. Certainly, steam reforming of such highsulfur, high-coking feedstocks will necessitate development of a new class of steam-reforming catalysts. The study of their sulfur tolerance properties should constitute an important, fruitful area of research. Because of large temperature gradients in steam reforming and the significant reversibility of H,S adsorption at high temperatures representative of the outlet conditions (1075-1 125 K), the equilibrium distribution of sulfur in the catalyst bed and particles is of interest. Rostrup-Nielsen (237b), has demonstrated how equilibrium sulfur-coverage profiles through a steam reforming reactor can be estimated for different inlet sulfur concentrations (Fig. 35). His calculations assume that chemisorption equilibrium is readily established and depends only upon PH2s/PH2 and temperature. Axial profiles of hydrogen flow and temperature are assumed to be linear (see legend, Fig. 35). Since it is evident that poisoning will influence the sulfur, hydrogen, and temperature profiles over a period of time, the results of RostrupNielsen may be only approximately representative for a limited time span. Qualitatively, however, they illustrate how overall sulfur coverages as a function of position in a fixed bed can be determined for different sulfur concentrations at any given time. Since sulfur apparently poisons by blocking the sites, it should be possible to link the sulfur profiles to changes in activity as a function of concentration, thereby enabling conversion-versus-time behavior to be modeled. In addition, the calculations of Rostrup-Nielsen (237) consider the effects of pore diffusion, the results suggesting that equilibrium coverage is attained rapidly at the external surface of the catalyst pellets in the entire bed; thus, the adsorption front moving gradually through the bed is not as sharp as it would be in the case of methanation at lower temperatures. This means that a large, but short-term, increase in the inlet sulfur concentration in the feed could significantly upset the entire process by causing a significant increase in the coverage of the external pellet layer throughout the reactor bed. It also means that accumulation in the interior of the pellet is a slow

219

SULFUR POISONING OF METALS

0 05 0 02 0 01 O 0

, L

O

I

n

rJ

i."n

-7 I M I I.

I ,

FIG.35. Sulfur coverage of the nickel surface of a reforming catalyst in a typical naphthabased ammonia plant as a function of distance through the reactor and inlet sulfur concentration (ppm) (Ref. 237b). Assumed profiles of temperature and hydrogen flows: 0 m: 4 m: 11 m:

773 K, 3.19 x kmol HJkg naphtha feed 948 K, 7.77 x lo-* kmol HJkg naphtha feed 1068 K, 1.25 x lo-' kmol H,/kg naphtha feed

process. Indeed, Rostrup-Nielsen estimates overall inlet coverages for typical catalysts in naptha reforming for an ammonia plant to be 10 and 70% of equilibrium coverage after 100 and 1000 hr, respectively. The work of Rostrup-Nielsen is very informative, but it also raises a number of important questions. How can more realistic temperature and concentration profiles through the reactor be incorporated into a reactor deactivation model? Could experimental measurements be performed to determine how sulfur is actually distributed in the catalyst pellets and in the bed and how this distribution changes as a function of time at various H,S concentrations? Would it be worthwhile to consider a modification of the model by Wise and co-workers (195,233)for steam reforming in which pore

220

C. H . BARTHOLOMEW ET AL.

diffusional resistance is considered so as to account for the slow deactivation of the pellet interior?

D. AMMONIA SYNTHESIS Poisoning of iron catalysts during ammonia synthesis by sulfur compounds has received relatively little attention (154, 240-244). Nevertheless, the previous work provides information on the poisoning mechanism and interesting examples of how oxide promoters may influence the sulfur poisoning behavior of a catalytic metal. Conventional ammonia synthesis catalysts contain 3-4 wt. % A1203 and 0.5-1 wt. % K 2 0 , both promoters being present largely at thecatalyst surface. Indeed, it has been estimated that more than 90% of the surface is composed of these two promoters and less than 10% of metallic iron (245). Bulatnikov et al. (%0) studied the effects of promoters on sulfur resistance of iron by measuring the amount of radioactive H2S adsorbed on iron catalysts promoted with A1203 and/or K,O. They reported irreversible deactivation of Fe promoted with A1,0,, A1,03 + K 2 0 , or K 2 0 after 0.8, 1.5, and 5 monolayers of sulfur had been adsorbed. In other words, the presence of K,O was responsible for increasing sulfur adsorption capacity, although it was not clear upon which portion of the surface sulfur had adsorbed. It was also reported that A120, was necessary to prevent volatilization of K 2 0 . Similar results were obtained by Zubova and co-workers (244) in a study of ammonia synthesis activity of iron catalysts containing varying amounts of Al,03, K 2 0 , and CaO promoters. The catalysts were exposed during reduction (673-773 K) to H,S in concentrations of 0.001-1.0%. Potassium oxide-promoted catalysts were reported to be significantly more resistant, and A1203-promotedcatalysts slightly more resistant, to sulfur than unpromoted iron. Alumina was speculated to be responsible for maintaining thermal stability in the presence of low-temperature-melting sulfides such as FeS, K2S, etc., but did not itself adsorb significant quantities of sulfur. Potassium oxide was thought to preferentially adsorb sulfur to form K2S and KHS compounds, thereby delaying formation of iron sulfides and loss of activity even at sulfur contents as high as 1.2%. This gettering effect of the alkali promoter, however, was not observed in two cases: (1) at very low H2S concentrations [e.g., at 0.001% (10 ppm)] or (2) if the sulfur was present in the reducing gas during the preparation of the catalyst. In the former case all catalysts were poisoned after only 0.25 wt. % sulfur had accumulated on the catalyst. Thus, promoted ammonia synthesis catalysts may nut be more sulfur resistant than unpromoted iron at industrially relevant sulfur

SULFUR POISONING OF METALS

22 1

concentrations. In the latter case FeS was reported to be easily formed and the catalyst completely deactivated by just 0.6 wt. % sulfur. This emphasizes the necessity of' carefully controlling sulfur impurity levels during catalyst preparation and reduction. Brill and Tauster (241) investigated unpromoted and Al,O,-promoted iron, finding that the effects of poisoning on activity were similar for the two catalysts, although larger quantities of poison were required to achieve the same effect in the promoted catalyst. Indeed, the unpromoted iron was poisoned completely by sufficient sulfur to cover only 1/6 of the iron surface, whereas the promoted catalysts was poisoned completely by one monolayer of sulfur. The authors postulated that the presence of an Fe-A1 alloy at the surface might account for the enhanced adsorption of sulfur, since A1,03 present at the surface would adsorb only small quantities. Other possible explanations not considered by the author include differences in H,S concentration in the sulfiding treatments and in metal crystallite size between the two catalysts and the possibility of bulk iron sulfide formation at the entrance to the catalyst bed. Three other observations from the work of Brill and Tauster (241) are noteworthy. The loss of activity of both catalysts after exposure to a pulse of poison was very slow (on the order of 10-100 hr), suggesting that the sulfur was preferentially adsorbed at the entrance to the bed and then diffused slowly through the bed until the poison was uniformly distributed. The activation energy for unpoisoned and partially poisoned catalysts (both promoted and unpromoted) was the same (96 kJ/mol), suggesting that the poisoning involves a blocking of iron sites, rather than a modification of the electronic properties or Fermi level of the metal. Moreover, a linear dependence between the rate constant and the square of the concentration of unpoisoned surface was observed (Fig. 36), suggesting that two iron sites were poisoned by each adsorbed sulfur in the promoted catalyst. The effect of H2S poisoning (10-30 ppm) on singly (A1,0,) and doubly (A1,03 and K,O) promoted iron was investigated by Krabetz and Peters (242). They found that the overall activation energy did not vary with degree of poisoning for either kind of catalyst. However, doubly promoted, KzOcontaining catalysts were found to lose significantly less activity than singly promoted catalysts at comparable coverages of sulfur. This is illustrated by the data in Fig. 37 showing that normalized activity (the rate constant of the poisoned catalyst normalized with respect to that for the fresh catalyst) for the singly promoted catalysts drops rapidly to zero at a sulfur coverage (0,) of 3%, whereas the doubly promoted catalysts maintain 20% activity at 0, > 30%. It should be emphasized that extensive poisoning at such low sulfur coverages is consistent with the fact that only a small fraction of the surface in promoted catalysts consists of iron metal. Krabetz and

222

C. H. BARTHOLOMEW ET AL.

0.10

X

0.05

0.0I

0

0.5

1.00

S2 FIG.36. Dependence of the reaction rate constant on the square of the unpoisoned surface in ammonia synthesis (Ref. 241).

Peters (242) also observed that under identical conditions of poisoning, potassium-promoted catalysts adsorbed approximately twice as much sulfur as the singly promoted catalyst, suggesting as in the Russian studies (240, 244) that K 2 0 increased the adsorption, possibly forming K2S. Brill and co-workers (243) investigated the poisoning of unpromoted and A120,-promoted iron over a range of temperature and H,S concentrations. At high temperatures (623-673 K) and high concentrations of H2S, formation of bulk FeS was observed. Although the bulk layers of sulfur could be removed by synthesis gas at high temperature, the surface layers could not be removed at temperatures up to 893 K. After reduction of the bulk layer or following exposure to 1.6 ppm H2S at 673-773 K, the concentration of sulfur on the surface was 0.4-0.5 mg/m2 of the iron surface corresponding to nearly monolayer coverage (i.e., 0, = 0.8-1.0) for both promoted and unpromoted iron. Nevertheless, only 0.2 mg S/m2 of iron surface (0, = 0.4)

223

SULFUR POISONING OF METALS

loo L

o m Y \ L

m Y

50

0 0 II

\

la

‘x

0

I 10

20

30

FIG.37. Normalized activity as a function of sulfur concentrations: x , singly promoted catalysts; 0 ,doubly promoted catalyst (Ref. 242).

was sufficient for complete deactivation of either catalyst. Sintering (loss of BET surface area) of both catalysts was reported to be enhanced by sulfur poisoning, although it was observed to a lesser extent in the Al,O,-promoted catalysts, thus explaining its greater resistance to sulfur poisoning. This conclusion is supported by recent experiments of Grabke et al. (142) involving the sintering of porous iron pellets in H, and H,S/H, atmospheres. They interpreted their results to suggest that sulfur decreases the surface energy of Fe and thereby increases the surface self-diffusion coefficient of iron atoms. Increases in the self-diffusion coefficients of other metals due to adsorbed sulfur have also been reported (48b). Probably the most definitive study of iron-sulfur interactions was performed by Grabke and co-workers (154). Their approach to obtaining thermodynamic data was novel. Indeed, the effects of sulfur on nitriding and carbiding of iron were used as the basis for determining the adsorption isotherm of sulfur on iron at 1325 K. These results show that less than complete values below saturation of the surface with H,S occurs only at PHIs/PH2 2x at 1325 K. In other words, extremely high temperatures are required before reversible adsorption occurs. In addition, the results provide insight into ammonia synthesis and CO/H, synthesis processes. For example, in the nitriding of iron, the results can be interpreted to mean that dissociation of N, is the rate-controlling step for which two sites must be free of sulfur

224

C . H. BARTHOLOMEW ET A L.

(see Fig. 38). In other words, sulfur apparently poisons the adsorption of N,, which is also the rate-controlling step in ammonia synthesis. By analogy, one might postulate that sulfur poisons the carbiding of iron catalysts and the dissociation of CO, also thought to be the rate-determining step in CO/H, synthesis at low temperature. The above studies provide a qualitative, and to a limited degree, a quantitative understanding of poisoning in ammonia synthesis. However, the role of promoters in activity maintenance is not well understood. In most of the studies cited, the working state of the catalysts, i.e., surface composition and surface chemistry, was not specified. The most serious difficulty was the ill-defined nature of the sulfur distribution on the different surface phases and throughout the catalysts or catalyst bed. The need is apparent for careful, basic investigations of adsorption and reaction in which the sulfur is uniformly distributed on the catalyst surface, accompanied by quantitative determination of its coverage and its chemical state using surface techniques. I .o

0.8

0.6 0

L

\

L

0.4

0.2

0

0

0.2

0.4 0.6 2 ( I - 8/8,)

0.8

I .o

FIG.38. Normalized activity for nitriding of Fe versus the square of the fraction of unpoisoned surface (Ref. 154).

SULFUR POISONING OF METALS

225

E. HYDROGENATION, DEHYDROGENATION, AND HYDROGENOLYSIS OF ORGANIC COMPOUNDS

Maxted ( I ) reviewed pre-1950 studies of sulfur poisoning of hydrogenation catalysts. A number of more recent studies have reported the effects of sulfur poisoning on activity and selectivity of platinum and (to a lesser extent) nickel catalysts in hydrogenation (191, 246-257), dehydrogenation (251, 258-266), and hydrogenolysis (88, 110, 171, 246, 251, 252, 255, 267) of organic compounds. Generally, sulfur compounds lower the activity of platinum (and nickel) catalysts for these reactions. However, the extent of poisoning may vary for each of these different reactions and may depend upon (1) the organic compounds undergoing reaction, (2) the kinds of sulfur compounds, (3) the conditions under which catalysts are poisoned, and (4) the properties of the catalyst. Therefore, in order to systematically address these complications, the previous work will be discussed in terms of three important concepts originally introduced by Maxted (1): toxicity, poisoning susceptibility, and selective poisoning. 1. Toxicity This term refers to the extent of inhibition of a poison for a given reaction and catalyst. According to Maxted ( I ) , toxicity depends upon electronic configuration, molecular structure and size, and length of stay on the surface. Shielded structures such as SO:- are less toxic than unshielded structures such as H,S for poisoning of a given metal. Thus the toxicity of sulfur compounds diminishes as the oxidation state increases. In general, the toxicity per gram atom of sulfur increases with molecular weight and size of the poison and its mean lifetime on the surface, i.e., the stability of the adsorbed sulfur compound. These generalizations are most applicable to hydrogenation reactions carried out at relatively low temperatures (e.g., 300450 K). Since at higher temperatures many organic and other sulfur compounds decompose on metal surfaces to form 2-D sulfides, toxicity would be independent of the starting compound under these severer conditions. There are, however, several examples of low-temperature hydrogenation reactions in which the concept of toxicity finds application ( I , 248,250,254, 268). In their investigation of the hydrogenation of crotonic acid o n platinum and nickel at 300 and 433 K, respectively, Maxted and and Evans (268)found the relative toxicities to increase in the following order: H,S, CS,, thiophene, and cysteine. This is a clear example of toxicity increasing with increasing molecular size. Curtis and Baker (248) found that dimethyl disulfide, t-butyl mercaptan, and n-propyl sulfide were 1250, 100, and 5 times as deleterious as thiophene in poisoning the hydrogenation of 1-hexene at 300 K on platinum,

226

C. H . BARTHOLOMEW ET AL.

suggesting that molecular structure plays an important role. Koppova et al. (250) found that 30-400/, less thiophene and diheptl sulfide (by weight) than heptyl mercaptan were required for complete poisoning of Ni/kieselguhr in the liquid-phase hydrogenation of 1, 2, 3, or 4-octyne. A somewhat more quantitative study of toxicity than those previously mentioned was undertaken by Gonzalez-Tejuca and co-workers (254), who determined the stoichiometry of adsorption for various poisons on Pt/A1,0, and Pt/SiO, catalysts during hydrogenation of olefins and benzene at relatively low temperatures. H,S and COS were found to deactivate one platinum site per molecule adsorbed, whereas CS, generally deactivated two platinum sites per adsorbed molecule. Titration with CS, to determine the number of active sites for hydrogenation was recommended over H, chemisorption since more consistent turnover numbers were obtained using CS,, apparently because H, adsorbs on sites which are not active for olefin hydrogenation. Not only was CS, found to deactivate more sites per molecule than the other poisons in olefin hydrogenation, it was also quite strongly adsorbed at 273-293 K. Indeed, the order of increasing reversibility was H,S E CS, < COS < thiophene. In fact, under the conditions studied thiophene was so easily desorbed that it was not an effective poison for Pt. In the case of nickel, however, thiophene was a very effective poison in liquid-phase hydrogenations of 1-alkenes at 300 K (257). Indeed, Holah and co-workers (257) observed that thiophene was strongly and nonuniformly adsorbed on nickel boride and Raney nickel catalysts. On the other hand, the poisoning curves for n-butanethiol were more typical of a weakly adsorbed poison distributed uniformly on the catalyst surface. Despite its being strongly adsorbed, thiophene was not capable of poisoning more than 75% of the sites in nickel boride or Raney nickel catalysts even at high poison concentrations, apparently because the geometric requirements for thiophene restricted easy access to a portion of the catalyst, although this portion was accessible to reactants. This illustrates how diffusional effects in micropores may influence the toxicity of a sulfur compound because of its molecular size and structure. Accordingly, the role of pore structure must be carefully considered in poisoning studies involving high-surface-area catalysts.

2. Poisoning Susceptibility This term refers to the sensitivity of a catalyst to poisoning under specified conditions. Two other terms are typically used to describe poisoning susceptibility. Poisoning resistance is the degree to which a catalyst resists deactivation, i.e., a catalyst which deactivates slowly is more resistant to poisoning than one that deactivates rapidly. Poisoning tolerance is defined typically as either the ultimate amount of poison a catalyst can adsorb and

SULFUR POISONING OF METALS

221

maintain a specified activity or the degree to which it maintains activity after it has adsorbed a steady-state amount of poison under specified reaction conditions. Sulfur resistance depends upon (1) catalyst composition, ( 2 ) physical properties of the catalyst such as surface area and pore structure, and ( 3 ) reaction conditions, most particularly, temperature and poison concentration. Sulfur tolerance could depend on all of the above except physical properties of the catalyst since it is an intrinsic property of the metal surface. Very few studies have considered the effects of catalyst composition and properties on sulfur resistance or tolerance in hydrogenation reactions. Maxted and Evans (268) measured poisoning coefficients (slopes of activity versus poison concentration plots) for nickel and platinum catalysts during hydrogenation of crotonic acid, and found that the values for nickel poisoned by H,S, CS,, thiophene, and cysteine were 200-300 times greater than those for platinum. In other words, nickel was 200-300 times less tolerant than platinum to sulfur poisoning under these conditions. That nickel boride and Raney nickel were strongly and irreversibly poisoned by thiophene during low-temperature hydrogenation of alkenes (257),whereas platinum catalysts are not (254), provides further evidence that platinum is more sulfur tolerant than nickel for hydrogenation reactions. This behavior could, in fact, be anticipated on the basis of the large heat of adsorption of H,S on Ni relative to that for Pt (see Fig. 13). There is some evidence that promoters and supports influence sulfur tolerance in hydrogenation and dehydrogenation reactions. For example, Landau et al. (253) found dealuminized Pd/zeolite to be a factor of 10 more sulfur tolerant than untreated Pd/zeolite. This was attributed to the higher acidity of the dealuminized zeolite and a stronger interaction of the metal with proton centers. Pt-Re/Al,O, is apparently less sulfur tolerant than Pt/AI,O, in dehydrogenation of cyclohexane and hexane and dehydroisomerization of methylcyclopentane because Pt-Re adsorbs sulfur more irreversibly under reaction conditions and is thus more sensitive to poisoning by additional reversible adsorption of sulfur compounds (266). Reaction conditions, especially temperature, play an important role in determining the sensitivity of catalysts to poisons. Generally, an increase in temperature will favor more reversible adsorption and thereby increase sulfur tolerance for a given catalyst. For example, CS, irreversibly poisons olefin hydrogenation on Pt at 273 K (254); however, at 323 K, 20% of the unpoisoned activity can be recovered and at 373 K almost 80% of the unpoisoned activity can be recovered upon heating in H, for 15-20 hr. Nevertheless, the actual steady-state activity in the presence of contaminants may not change appreciably with increasing temperature. This depends upon the degree of reversibility with respect to the adsorption of poison. For example, the poisoning coefficient for crotonic acid hydrogenation in the presence of

228

C. H. BARTHOLOMEW ET AL.

thiophene only decreased about 10% when the hydrogenation temperature was increased from 288 to 323 K even though a fivefold increase in reaction rate occurred (1). It should also be emphasized that the effects of temperature on sulfur tolerance can be complicated by changes in adsorption mechanism (e.g., dissociative rather than associative) and in stoichiometry at sufficiently high temperatures (see Sections III,B and C). 3. Selective Poisoning In a catalytic process involving single sites, a poison may preferentially interact with the most active sites. This has been referred to as “selective poisoning” (17.5) but will be more specifically classified here as “site-selectiw poisoning.” In a catalytic process involving more than one reaction, a poison may suppress the activity of one reaction more than another, leading to a change in product distribution. This has been called “reaction selectivity” ( I 7.5), but is more specifically denoted reac.tion-selL~ctivepoisoning in this review. Reaction-selective poisoning may be beneficial as in the case of the selective poisoning of the hydrogenolysis (coking) reactions in reforming (251,255,258-266). Somorjai (93)has proposed that facile reactions such as hydrogenation (or dehydrogenation) should be less affected by sulfur poisoning than demanding reactions (95) such as hydrogenolysis because the sulfur can, by restructuring of the surface, effect the deactivation of more than one to two surface sites for the structure-sensitive reaction. There are a number of interesting examples of how sulfur influences selectivity in hydrogenation (191, 246, 248, 249, 252, 2.56, 2.57), dehydrogenation (251, 2.52, 2.58), and hydrogenolysis (246, 251, 252) reactions. For example, Pines and co-workers (246) found that a small amount of sulfur prevented undesirable hydrogenation during hydrogenolysis of primary alcohols on Ni/kieselguhr. At very low sulfur contents the main reaction was the dehydroxymethylation, whereas highly poisoned Ni-catalyzed reductive dehydroxylation only : R H + CH, RCH,

+ H,O

+ H20

In the hydrogenation of olefins such as 3-3-dimethyl- 1-butene and cyclohexene on Ni/kieselguhr, the addition of sulfur caused hydroisomerization to 2-3-dimethylbutane and methylcyclopentane, respectively (247). Maurel ef al. (251) reported that CO, and H,S were “nonselective poisons” in the hydrogenation of benzene and benzene-D, exchange on platinum. That is, they did not affect the ratio of exchange/hydrogenation rates. However, elemental sulfur produced by the reaction of SO, and H,S was a “selective poison,” causing the ratios of the rates for exchange/hydro-

SULFUR POISONING OF METALS

229

genation to increase by a factor of two. They concluded that H,S and SO, are adsorbed indifferently on all catalytic sites, whereas elemental sulfur adsorbs on sites responsible for hydrogenolysis. The role of sulfur in reforming on Pt/Al,O, and Pt-Re/Al,O, was investigated by Menon and Prasad (266). They found that sulfur preadsorbed at 773 K suppressed all of the representative metal-catalyzed reforming reactions including hexane and cyclohexane dehydrogenation, dehydroisomerization of methylcyclopentane, dehydrocyclization of benzene, and hydrocracking. On Pt-Re, the extent of hydrocracking was reduced to a much greater extent than these other reactions, explaining why this catalyst produces less coke than Pt/A1,0,. That sulfur is responsible for suppressing hydrocracking of organic molecules on Pt is consistent with the work of Fischer and Kelemen (88) showing that bonding of benzene on Pt(100) is sufficiently modified by preadsorbed sulfur to enable an increasing fraction of the adsorbed benzene to desorb at elevated temperatures rather than to dehydrogenate. Sterba and Haensel (252) reviewed the effects of sulfur in reforming reactions, specifically hydrogenation, dehydrogenation, and isomerization reactions of cyclohexane on Pt/AI,O,. They noted that low levels of sulfur (50 ppm) cause striking changes in activity and selectivity; for example, in experiments at high space velocity the addition of 50 ppm sulfur decreased overall conversion of cyclohexane by a factor of 10, whereas conversion to benzene decreased from 97 to 0% and conversion to methylcyclopentane and methylcyclopentene increased from 0.9 and 0% to 44 and 35%, respectively. Thus the presence of only 50 ppm sulfur changed the entire course of the reaction. The examples discussed above illustrate some interesting qualitative features of hydrogenation; probably the most interesting feature is that of selective (sometimes beneficial) poisoning. Unfortunately, few of the previous studies provide quantitative measures of toxicity, sulfur tolerance, and selective poisoning. Thus the need for additional work in which these areas are quantitatively investigated is clear. Specifically, measurements of the intrinsic sulfur tolerance of different metals for different poisons in several important hydrogenation, dehydrogenation, and hydrogenolysis reactions over a range of temperature would be highly desirable.

VI.

Regeneration

Relatively few studies of regeneration of sulfur-poisoned catalysts have been reported in the literature ( I , 9, 11, 83, 92, 178, 194, 237, 269-274). Attempts to regenerate poisoned catalysts have been made using oxygenlair

230

C. H . BARTHOLOMEW ET AL.

( 9 , I I , 83, 92, 178, 194, 269, 272, 273), steam (237, 270), H, (9, 11, 112, 194, 274), and inorganic oxidizing agents ( I ) . Most of the studies have been made on Ni (9, 11, 178, 194, 237, 269-271, 273, 274), with only few exceptions, eg., Cu (92), Pt (83), and Mo (272). In the only study of regeneration using steam (237,270),it was found that up to 80% removal of surface sulfur from Ni steam reforming catalysts either unpromoted or promoted with Mg and Ca could be achieved at 973 K. The effluent gas analysis showed the presence of SO, and H,S, and the following reaction pattern was suggested: Ni-S

+ H,O

H,S

+ 2H,O

-+

--+

NiO

+ H,S

(15)

SO,

+ 3H,

(16)

Catalyst regeneration using steam was insigificant at temperatures below 873 K. Moreover, catalysts promoted with Na and K could not be effectively regenerated with the same treatment, presumably because of strong bonding between sulfur and alkali metals. Treatment with steam and air resulted in the formation of sulfates, which were subsequently reduced back to sulfide upon reduction with H2. These observations are consistent with those made by Bartholomew et ul. (194) during reduction of Ni/Al,O, catalysts which had been previously poisoned by H,S and regenerated in air. Although the steam regeneration above 973 K successfully removes adsorbed sulfur, use of such high temperatures (900-1000 K) results in severe sintering of commercial high-surface-area catalysts (237, 238, 275278). Accordingly, the investigation of the thermal stability of Ni bimetallics (277) and of Ni on different supports (276-278) in steam or H,/steam environments, may be a potentially fruitful area of research. Since adsorbed sulfur apparently increases the self-diffusion coefficients of metals (48b, 142), studies of sintering in the absence and presence of sulfur would also be worthwhile. Efforts to regenerate sulfur-poisoned Ni catalysts using H, ( 9 , I I , 112,194) have been largely unsuccessful. One patent (274) proposes using H2 to regenerate 15-60"/, Ni/Al,O, at 700-1 100 K, although reportedly no sulfur is removed by this treatment. In view of the fact that regeneration cannot be achieved without removing sulfur from the surface, this claim (274) is questionable. Studies using AES (9, 11) indicate that negligible sulfur is removed by H, (PH2z kPa; total pressure of 100 kPa) up to 873 K. Studies of H,S desorption from sulfided nickel surfaces indicate that the elution curves are exponential and that only 5-10% of the adsorbed sulfur can be removed in flowing hydrogen at 725 K over a period of 2-3 days (112). The very high stability of surface metal sulfides as compared to bulk metal

23 1

SULFUR POISONING OF METALS

sulfides (Table VII) suggests that if the H, gas passing over the poisoned catalyst is in thermodynamic equilibrium with the poisoned surface, the effluent gas would contain less than 5 ppb H,S at 700 K. If a Ni catalyst were treated with 1 ppm H,S at 700 K, the regeneration under identical conditions (H, flow rate) would take 200 times longer than the poisoning process. Thus regeneration using H, appears to be impractical. Recent investigations (9, I I , 8 3 , 92, 178,272) of regeneration by 0, have resulted in modest success. Table XXII lists studies made of Ni (11, 178), Cu (92), Pt (83), and Mo (272). These studies indicate that sulfur can be removed as SO, by low-pressure (Pol = 3 x 10- 9-7 x 10- kPa) oxidation at temperatures ranging from ambient to nearly 1600 K. Significant differences in the kinetics of sulfur removal observed in these studies are attributed to different reaction mechanisms being predominant at different temperatures. On the other hand, attempts to remove sulfur from polycrystalline Ni surface by atmospheric pressure oxidation (Po, = 10-760 Torr) at temperatures from 300 to 800 K were unsuccessful (9). Instead, nickel oxide layers were formed on top of the sulfur layer. Treatment of the oxidized catalyst with H, at 700 K for 2 hr reduced the Ni oxide completely to the metal; at the same time it caused the return of the sulfur layer to the surface. Colby ( 9 ) explained the failure to regenerate sulfur-poisoned Ni catalysts by the high-pressure 0, treatment in terms of two competing events on the surface: (1) oxidation of the sulfur to SO,, and (2) oxidation of the Ni to nickel oxide. At low oxygen pressures a monolayer of nickel oxide forms slowly, and the sulfur is oxidized to SO, and removed at a comparable rate. At the higher oxygen pressures, the formation of nickel oxide is markedly faster, resulting in the rapid growth of nickel oxide which rapidly buries the sulfide; in this instance conversion to SO, and sulfur removal does not occur.

'

TABLE XXIl Regeneration of Sulfur-Poisoned CatalyJts Using 0 , Treatment Poz (Torr)"

Metal(p1ane) Cu(l10) Mo(polycrysta1line) Ni(l1l) Ni(polycrysta1line) Pt(ll0) I'

760 Tom

=

103 kPa.

1.4 x lo-'

I 3 2 1.4

-

7 x

x 10-8 - 2 x 10-6 x 10-8 - 3 x x - 2 x x 1 0 - 8 - 7 x 10-6

Temperature (K)

Reference

880-1100 975 -1575 300 500 -875 440 -670

Y2 272 178 I1 83b

232

C. H . BARTHOLOMEW ET AL.

Recent work (11) shows that the competition between sulfur removal and nickel oxide formation is pressure and temperature dependent. The work of Colby may explain the difficulties encountered by Bartholomew et al. (194) in their attempt to regenerate Ni/A1,0, with 0, at near atmospheric pressures ; their observed recovery in catalyst activity after 0, treatment and subsequent low-temperature (525 K) reduction may have been due to only partial segregation of sulfur to the surface, with some sulfur still present in the bulk Ni/NiO layers. However, reduction with H, at 725 K resulted in complete loss of activity ( I N ) , apparently due to the complete segregation of sulfur to the nickel surface. The mechanism of sulfur removal by oxygen at low pressure is likewise explained by the work of Colby (9). A more complete understanding of the mechanism is obtained from the studies listed in Table XXII. The proposed mechanisms from these studies have several common features. The regeneration scheme involves basically three steps: (1) nucleation of sites to adsorb 0 , ,(2) reaction between adsorbed sulfur and adsorbed oxygen species via a Langmuir-Hinshelwood-type mechanism at these defect sites, and (3) growth of oxygen islands around these sites or “holes” in the sulfur structure. Again, differences may be observed depending on the temperature. At low temperatures, surface diffusion is considered to be slow, and sulfur is immobile. Hence, the sulfur-oxygen reaction takes place only at boundaries between the sulfur and the oxygen regions. Under these conditions, sulfur and oxygen coexist on the surface. On the other hand, at high temperatures, rapid surface diffusion makes it difficult for sulfur and oxygen to coexist on the surface. Thus, two different kinetic regimes may be observed, depending on the temperature. In fact, such an interpretation has been considered by several researchers (83, 92, 178, 272). Kawai et ul. (272) have suggested that at lower temperatures ( - 975 K), the regeneration proceeds via the Eley-Rideal mechanism in which a reaction between gaseous oxygen and surface sulfur to form SO, is predominant. Holloway and Hudson (178) on the other hand, suggest that surface defect sites physisorb molecular oxygen which then reacts with sulfur to form SO,. These discrepancies in the nucleation cycle of the regeneration scheme cannot be resolved on the basis of a limited number of sets of kinetic data alone. Additional kinetic studies, which involve a concerted application of surface-sensitive spectroscopic techniques, should prove helpful in understanding the regeneration mechanism. Future efforts should be directed toward development of more practical regeneration schemes, i.e., at higher pressures (atmospheric and higher) and at temperatures low enough to minimize sintering of commercial catalysts. Attainment of these goals remains a challenging area for future researchers.

SULFUR POISONING OF METALS

VII.

233

Conclusions and Recommendations

A. CONCLUSIONS 0 Metal-sulfur bonds of adsorbed sulfur are significantly stronger than metal-sulfur bonds in bulk metal sulfides. Heats of H,S adsorption on metals are SO-lOO% larger than heats of formation of the bulk sulfides. PHZs/PH2 ratios for gas-phase H,S in equilibrium with completely covered Co, Fe, Ni metal surfaces are typically 10-7-10-'0 at moderate temperatures (525-725 K). In the case of more noble metals, higher P H Z S / P H 2 ratios would be expected for monolayer coverage under the same conditions. 0 In the case of well-defined, single-crystal metal surfaces, a wealth of fundamental information is available regarding the structures and stoichiometries of adsorbed sulfur. Unfortunately, such information is generally not available for polycrystalline and supported metals. 0 Adsorption of H,S on metals occurs dissociatively over a wide range of temperatures and coverages. At low coverages of sulfur on metals, sulfur resides in the high-coordination sites, each sulfur atom bonded to three or four metal atoms; however, as the sulfur coverage is increased, the surface restructures, and in some instances may incorporate metal atoms into the surface sulfur layer. The strength of the metal-sulfur bonding is reduced somewhat by restructuring. 0 At moderate temperatures a surface metal sulfide is formed independent of the nature of the original sulfur-containing compound. 0 Adsorbed sulfur reduces the capacity of metals for adsorption of other molecules. H, adsorption capacity is reduced in approximately direct proportion to the fraction of surface covered by sulfur. At low pressures CO adsorption is decreased by adsorbed sulfur. On the other hand, adsorbed sulfur can cause increased CO adsorption under conditions where metal carbonyls or subcarbonyls are stable. 0 Because of the strength of the sulfur-metal bond, sulfur adsorption is frequently very nonuniform with respect to catalyst bed and particle, making fundamental interpretation of poisoning phenomena very difficult or impossible. 0 Metal catalysts employed in CO hydrogenation processes are generally extremely sensitive to poisoning by sulfur, losing three to four orders of magnitude specific activity at parts per billion levels. 0 Product selectivity in CO hydrogenation is generally shifted by sulfur poisoning in the direction of higher molecular weight products. 0 In CO hydrogenation, NH, synthesis, and steam reforming the data

234

C. H . BARTHOLOMEW ET AL.

strongly suggest that the poisoning involves primarily a geometric blocking of active sites by adsorbed sulfur. 0 At low temperatures characteristic of conditions for hydrogenation of organic compounds, sulfur adsorption may occur associatively, in which case the toxicity of the poison is dependent upon molecular size, structure, and strength of adsorption. 0 Sulfur resistance, i.e., the rate of deactivation, can be a function of catalyst properties (e.g., composition, physical and chemical properties of the support, and promoter type) and reaction conditions (particularly temperature and sulfur concentration). Sulfur tolerance, on the other hand, is independent of the physical properties of the catalysts and is rather an intrinsic property of the metal. It is mainly a function of the strength of the metal-sulfur bond. 0 Because of strong metal-sulfur bonds, regeneration under reducing conditions is impractical; use of oxidative conditions is a much more promising approach.

B. RECOMMENDATIONS The discussion of previous work in this review has focused on several fundamental questions, all of which were specifically addressed to some extent. Nevertheless, there are many basic questions that remain unanswered. These relate to the strength, structure, and stoichiometry of bonding to metal surfaces, to the role of sulfur in affecting the adsorption of other species, to the effects of sulfur on the rates and mechanisms of surface-catalyzed reactions, and to the role of supports, promoters, compound formation, and reaction conditions in determining the sensitivity of metals to sulfur poisoning. The need for expanded efforts at the fundamental level is apparent. Such efforts have the greatest potential for providing the insights necessary for development of practical solutions to poisoning problems such as regeneration schemes and more sulfur-tolerant catalysts. Further surface science studies are necessary to better define the interactions between sulfur and metals, but these should focus on conditions more representative of those found in actual poisoning processes and should attempt to bridge the gap between the information obtained in clean and in reacting systems. Because of the strong, nonuniform nature of sulfur adsorption and potential heatimass transport disguises in pellet-catalyst/reactor systems, the results obtained therein are often difficult or impossible to interpret; thus extreme care must be exercised in the choice of reactor design, catalyst form, and reaction conditions in sulfur poisoning studies so as to avoid these difficulties.

SULFUR POISONING OF METALS

235

Finally, engineering modeling studies for prediction of deactivation behavior are needed which consider quantitatively all of the chemical and rate processes involved in the commercial process.

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Decay in Methanol Synthesis and Fischer-Tropsch Catalysis.” Final Report to ERDA, PERC-0060-8, Sept. 30, 1977. 234. Morita, S., and Inoue, T., I n t . Chem. Eng. 5, 180 (1965). 235. Morita. S., and Inoue, T., Kog. Kugak. Zussh. 68,659 (1965). 236. Bridger, G. W., and Chinchen, G . C., in “Catalyst Handbook.” Wolf, London, 1970. 237. (a) Rostrup-Nielsen, J. R., “Steam Reforming Catalysts.” Teknisk Forlag. Copenhagen, 1975; (b) Rostrup-Nielsen, J. R., Paper presented at NATO ASI, Lagos, Portugal, May, 1981. 238. Rostrup-Nielsen, J . R., J. C u r d 31, 173 (1973). 239. Tomita, T., and Kitagawa, M., Eur. Meet. Chem. Eng., Frankfurt June 23 (1976). 240. Bulatnikov, Yu. I . , Apel’baum, L. O., and Temkin, M. I., Zh. Fiz. Khim. 32, 2717 (1958). 241. Brill, R., and Tauster, S., Ber. Bunsenges. Phys. Chem. 67, 390 (1963). 242. Krabetz. R., and Peters, C., Ber. Bunscnges. P l i ~ s Chem. . 67, 522 (1963). 243. Brill, R., Schaefer, H., and Zimmermann. G., Ber. Bunsenges. Phvs. Chem. 72, 1218 (1968). 244. Zubova, I . E., Rabina, P.D., Pavlova, N. Z., Kuznetsov, L. D., Chudinov, M. G., and Shang, L. K., Kinrt. Carol. 15, 11 18 (1974). 245. “Catalyst Handbook,” 1). 138. Wolf, London, 1970. 246. Pines, H., Shamaiengar, M., and Postl, W. S., J Am. Chem. Soc. 77, 5099 (1955). 247. Pines, H., Marechal, J., and Postl, W. S., J. Am. Chem. SOL.77, 6390 (1955). 248. Curtis, J. L. S., and Baker, M. O., A n d . Chrm. 42, 278 (1970). 249. Kitayama, Y., and Hayakawa. M., Chem. Lett. (Chem. Soc. Jpn.) 181 (1973). 250. Koppova, A,, Zapletal, V., RbiiEka, V., and Soukup, J., Coll. Czech. Chem. Commun. 38, 2472 (1973). 251. Maurel, R., Leclercq, G., and Barbier, J.. J . Curd. 37, 324 (1975). 252. Sterba, M. J . , and Haensel, V., Ind. Eng. Chem. Prod. Res. Dev. 15, 2 (1976). 253. Landau, M. V . ,Kruglikov, V. Ya., Goncharova, N. V., Konoval’chikov, 0.D., Chukin, G. D., Smirnov, B. V., and Malevich, V. I., Kiner. Cutul. 17, 1104 (1976). 254. Gonzalez-Tejuca, L., Aika, K., Namba, S., and Turkevich, J., J. Phys. Chem. 81, 1399 (1977). 255. Fuentes, S., and Figueras, F., J. Chem. SOL.Furuday Trans. 174, 174 (1978). 256. Cosyns, J., Franck, J. P., and Gil, J. M., C. R . Acud. Sci. Paris 288, 85 (1979). 257. Holah, D. G . , Hoodless, 1. M., Hughes, A. N., and Sedor, L., J . Cutal. 60, 148 (1979). 258. Chang, Y. C., and Kalechits, I. V., K’o Hsueh T’ung Pa0 15,478 (1958); [Chem. Ahstr. 53, 10718g (1959)l. 259. Lin, T . Y., and Cheng, Y. C., Wu Hun Ta Hsueh, Tzu Jan Ico Hsiieh Hsiieh Puo 5, 32 (1959); [Chem. Ahstr. 54, 55987e (196O)l. 260. Minachev, Kh. M., and Kondrat’ev, D. A,, Izv. Akud. Nuuk S S S R Otdel. Khim. Nuuk 877 (1961); [Chem. Absrr. 55,27145d (1961)l. 261. Minachev, Kh. M., and [sagulyants, G. V., Proc. Int. Congr. Cutul., 3rd, Amsterdam p. 308 (1965). 262. Maslyanskii, G. N., Zharkov, B. B., Rubinov, A. Z., and Klimenko, J. M., Kinet. Kutul. 12, 1060 (1971). 263. Pfefferle, W . C., Prepr. A m . Chem. SOC.Div. Pet. Chem. 15(1), A21 (1970). 264. Hayes, J . C., Mitsche, R. T., Pollitzer, E. L.. and Homeier, E. H., Prepr. Nat. Meet., 167th Am. Chem. SOC.,Los Angeles (1974). 265. Sivasanker, S., and Ramaswany, A. V., J . Cural. 37, 553 (1975). 266. Menon, P. G., and Prasad, J., Proc. Int. Congr. Catat., 6th London p. 1061 (1976).

242

C. H. BARTHOLOMEW ET AL.

267. Smith, R. L., Naro, P. A,, and Silvestry, A. J., J . Curd. 20, 359 (1971). 268. Maxted. E. B., and Evans, H. C., J . Chrm. Soc. 603 (1937). 269. Riesz, C. H., Dirksen, H. A., and Kirkpatrick, W. J., "Sulfur Poisoning of Nickel Catalysts." Institute of Gas Technology Research Bulletin No. 10, September, 195I . 270. Rostrup-Nielsen, J . R., J . Catal. 21, 171 (1971). 271. Pieters, W. J., Freel, J., and Anderson, R. B., U.S. Patent 3,674,707, 1972. 272. Kawai, T., Kunimori, K., Kondow, T., Onishi, T., and Tamaru, K.. J . Clicm. Sot,. Fnruduy Trctns. 172, 833 (1976). 273. Schoofs, R. J., Nordhausen, L. J., and Dugdale, L. A.. U.S. Patent 4,026,821, 1977. 274. Dobashi, H. H., U S . Patent 4,065,484, 1977. 275. Williams, A,, Butler, G. A,, and Hammonds, J . , J . Crrtul. 24, 352 (1972). 276. Pannell, R. B., Bartholomew, C. H., and Fowler, R. W., Prepr. Am. Chem. Soi,. Div. Pet. Chenz. 22 (4), 1331 (1977). 277. Bartholomew, C. H., Pannell, R. B., and Fowler, R. W., J . Catal., in press (1982). 278. Sorenson, W. L., and Bartholomew, C. H., in preparation (1982).

ADVANCES IN CATALYSIS. VOLUME 31

Methano I Synthesis K . KLIER Department of Chemistry and Center.for SurJuce and Coatings Research Lehiyh Uniiiersiiy Bethlehem. Pennsvlcaniu

1. Introduction . . . . . . . . . . . . . . . . I1 . Catalyst Selection . . . . . . . . . . . . . . 111. The Activity of Pure Copper Metdl . . . . . . . . . IV . The Activity of Pure Zinc Oxide . . . . . . . . . . V . The Binary Copper-Zinc Oxide System . . . . . . . . A . Preparation of the Cu/ZnO Binary Catalysts . . . . . B. Physical Characteristics of the Cu/ZnO Catalysts . . . . C . Particle Size and Morphology . . . . . . . . . D . Surface Analysis by XPS-Auger Spectroscopy . . . . . E. Surface Areas of the Individual Catalyst Components . . . F . Chemisorption of Carbon Monoxide and Hydrogen . . . G . Activity Patterns . . . . . . . . . . . . . H . Reaction Kinetics in the Presence of Carbon Dioxide . . . I . The Effect of Carbon Dioxide on Selectivity . . . . . J . Relative Hydrogenation Rates of Carbon Monoxide, Hydrocarbons, and Oxygenates . . . . . . . . . VI . Other Binary Catalysts . . . . . . . . . . . . . A . Copper-Based Catalysts . . . . . . . . . . . B . Transition-Metal-Based Catalysts . . . . . . . . VII . Ternary and Quaternary Catalysts . . . . . . . . . . VIII . Mechanisms . . . . . . . . . . . . . . . . A . Possible Reaction Pathways . . . . . . . . . . B . Methods Used in Mechanistic Studies of Methanol Synthesis C . Evidence for Adsorbed Carbon Monoxide and Hydrogen . . D . Evidence for Adsorbed Formate and Methoxide . . . . E . Surface Complexes Formed by Coadsorption of the Reactants F . Mechanism of Methanol Synthesis at Low Temperatures and Pressures . . . . . . . . . . . . . . References . . . . . . . . . . . . . . . .

I.

.

.

.

. 251

243

. . 254 . . 257 . . 251 . . 258 . . 258 . . 261 . . 266 . . 261 . . 268 . . 271 . . 274 . . 284 . . 285 . . 287 . . 287 . . 289 . . 291 . . 296 . . 291 . . 299 . . 302 . . 303 . . 304 .

.

. 308 . 310

Introduction

Methanol synthesis is a process of major industrial importance consisting of hydrogenation of carbon monoxide or of carbon dioxide according to the equations CO

+ 2H. + CH. OH 243

.

Copyright .i. 1982 ..by . Academic Prccs Inc . All rjghrs of reproducrion in any form reserved . ISBN O - I ~ - O O-7~ U ~ I

244

K . KLlER

where AH,,,,

= - 100.46 kJ/mol

CO,

and AG,,,,

=

+45.36 kJ/mol:

+ 3H, *CH,OH + H 2 0

(2)

where AH,,,, = -61.59 kJ/mol and AG,,,, = +61.80 kJ/mol. Methyl alcohol is one of the least thermodynamically probable products of CO and CO, hydrogenations, since the formation of higher alcohols and hydrocarbons is accompanied with more negative free-energy change, as illustrated in Fig. 1. Dimethylether (DME) may also be produced by the dehydration 2CH,OH

where AH,,,,

=

* CH30CH, + H,O

-20.59 kJ/mol and AG,,,,

(3)

= - 10.71 kJ/mol.

GlBBS FREE ENERGY OF CARBON MONOXIDE HYDROGENATION TO:

ARAFFINS t WATER PARAFFINS t C o p

"i

7 I

2

3

4

5

6

7

8

9

10

NUMBER OF CARBON ATOMS IN THE HYDROGENATED PRODUCT

FIG. I . Gibbs free energies A G at 600 K (kcal/mol of carbon) in the product alcohol or hydrocarbon of the reactions:

+ +

nCO 2nH, +CH,(CH,),-,OH (2n - 1)CO + ( W + 1)HZ +CH,(CH2),-,OH 1)H2 +CH,(CHz),-,CH, K O + (2n 2nCO + ( n + 1)H2 dCH,(CH,),-,CH,

+ (n - I)H,O + (n - 1)COz + nHZO + nCO,

(0) (+)

(0)

(C,)

From D. R. Stull, E. F. Westrum, and G. C. Sinke. "The Chemical Thermodynamics of Organic Compounds." Wiley, New York, 1969. 1 cal = 4.184 J .

METHANOL SYNTHESIS

245

To avoid the formation of DME, higher alcohols, and hydrocarbons, methanol synthesis requires a selective catalyst that rapidly executes the hydrogenation of CO or CO, to methyl alcohol but at the same time does not allow the reaction to proceed any further downhill on the thermodynamic free-energy scale. Based on general experience with catalysts in other hydrogenation reactions, this would appear a formidable if not an impossible task. However, after the development involving some six decades, industry now does have a catalyst which selectively steers the hydrogenation of the synthesis gas, CO CO, + H, in various proportions, to almost pure methanol. The activity, and consequently, the pressure and temperature range at which the methanol synthesis catalysts operate have undergone dramatic changes in the last two decades, which culminated in the commercialization of the so-called low-pressure process operating at pressures less than 100 atm. The low-pressure catalysts invariably contain copper and a mixture of oxides such as ZnO-Al,O, or ZnO-Cr,O,. Oxides V,O,, V 2 0 3 ,Tho,, and others have also been claimed as effective components of the low-pressure copper-based catalysts. The credit for the first invention of the copper-based catalysts goes to M. Patart (I), who has patented many inorganic oxides, salts, and metals, including copper, for methanol synthesis at pressures 150-200 atm and temperatures 300-600°C. Following Patart's work, several investigators noted the curious fact that oxides and copper separately had relatively low activity, whereas mixtures of an oxide and copper increased the yield considerably. It was this mutual promotion effect to which many conflicting reports on the activity of copper-based catalysts could be traced. Realization came early that preparation of copper catalysts is an extremely important factor in determining their activity for methanol. In one report, copper catalysts prepared by precipitation by alkali hydroxide were active, while those precipitated by ammonium hydroxide were inactive, and suspicion was raised that alkaline impurities may be responsible for the activity of copper catalysts (2). Methanol decomposition also gave conflicting results in the early research. Sabatier and co-workers, who first employed methanol decomposition as a guide for finding active methanol synthesis catalysts, found many catalysts, including copper, that decomposed methanol (3, 4), but the attempts to synthesize methanol by the reverse process gave very poor results with copper catalysts (5). The question of the activity of copper and of the promotion effect of zinc oxide was further investigated by Frolich and his co-workers (6, 7), who prepared the whole range of binary compositions CuO/ZnO from OjlOO to lOOj0 by coprecipitation of hydroxides from nitrate solutions by ammonia, calcination at 220°C, and reduction by methanol at 200-220°C. The key results of their methanol decomposition

+

246

K . KLlEK

(6) and methanol synthesis (7) study are summarized in Fig. 2. At 360°C and 204 atm, zinc oxide was quite a good catalyst for both the decomposition and the synthesis, while copper, which was present in the form of metal, as Frolich et al. showed later (8),displayed some activity for the decomposition but none for the synthesis. The maximum rate of the synthesis and decomposition reactions appeared at catalyst compositions near CuO/ZnO = 30170, and Frolich et a/. (8) engaged in a study to elucidate the nature of this

CUO 100

80 COMPOSITION

60

40

20

0

OF C A T A L Y S T (MOL % )

FIG.2. Comparison of methanol decomposition at 360’C and atmospheric pressure with synthesis at 350°C and 204 atm using the same catalysts (7). [Reprinted with permission from Ind. Eng. Che~n.20, 1327 (1928). Copyright (1928) American Chemical Society.]

METHANOL SYNTHESIS

247

effect by X-ray determinations of the lattice constants of copper and zinc oxide over the whole compositional range. Indeed, lattice parameters of both the copper and the zinc oxide were found to depend on the catalyst composition. The lattice extension of copper was attributed to alpha brass formation upon partial reduction of zine oxide, and an attempt was made to correlate the lattice constant of copper with the decomposition rate of methanol to methyl formate. Furthermore, the decomposition rate of methanol to carbon monoxide was found to correlate with the changes of lattice constant of zinc oxide. Although such correlations did not establish the cause of the promotion in the absence of surfacearea measurements and of correlations of specific activities, the changes of lattice parameters determined by Frolich et al. are real and indicate for the first time that the interaction of catalyst components can result in observable changes of bulk properties of the individual phases. Frolich et af. did not offer an interpretation of the observed changes in lattice parameters of zinc oxide. Yet these changes accompany the formation of an active catalyst, and much of this review will be devoted to the origin, physicochemical nature, and catalytic activity of the active phase in the zinc oxide-copper catalysts. The patterns of methanol decomposition rates were qualitatively confirmed by Kostelitz and Huttig (9) who concluded in an extensive study of the whole Cu/ZnO compositional range that a “synergic promotion” took place in the zinc-oxide-rich catalysts which also displayed the highest activity. The term synergic promotion was coined by Mittasch to describe effects due to chemical and physical interactions of the catalyst components, as opposed to “structural promotion,” associated with increasing or maintaining the surface area of the active substance without changing the activity of unit area. The synergic promotion was described by Kostelitz and Hiittig in the following words : “Es miissen sich also an den Kupfer-Zinkoxyd-Phasengrenzflachen spezifisch wirkende Kraftfelder herausgebilt haben, denen ein wesentlicher Anteil an der katalytischen Wirksamkeit zugeschrieben werden muss,” i.e., “There must be specifically acting force fields built up at the copper-zinc oxide interfaces, to which a substantial part of the catalytic activity must be attributed.” The thread of evidence and thought that an active catalyst is formed by the coexistence and through the physicochemical interactions of two or more separate solid-phase components, has once again appeared in the early catalytic literature. With the experimental tools of the first half of the century, however, it was not possible to analyze what the “spezifisch wirkende Kraftfelder” were and how they were related to catalysis. It must be noted that despite these interesting findings and interpretative

248

K . KLIER

developments, it was not demonstrated until much later that the copperzinc oxide catalysts would work at low pressures, and industry used instead a catalyst based on zinc oxide and chromia (also termed “zinc chromite”) in a high-pressure process commercialized by BASF in 1923 (5). As late as 1955, the copper-zinc oxide-chromia catalyst was considered by Natta as impractical because of its poor resistance to thermal shocks, concomitant poor reproducibility, and sensitivity to chemical poisons (10). The development of the high-pressure process, which included thorough kinetic studies, was reviewed by Natta (10). In the 1950s methanol synthesis appeared to be a mature process and zinc chromite a very satisfactory catalyst. The decade of 1960s brought a dramatic industrial breakthrough, however. Pioneered by ICI, soon followed by other companies, low-pressure methanol synthesis processes emerged, all utilizing ternary catalysts based on copper, zinc oxide, and another oxide such as chromia or alumina. Although the basic combination copper-zinc oxide had been known to be an active methanol synthesis catalyst for a considerable length of time, it was the more systematic and careful, although apparently still quite empirical, preparation and testing of the mixed catalysts that led to the new process in which the synthesis was carried out at pressures of 50-100 atm and temperatures of 220-250°C. Tables I and 11 summarize catalyst compositions, conditions, and methanol yields for these various new low-pressure catalysts. Compared to the older high-pressure processes which ran typically at 200 atm and 350”C,the low-pressure conditions offered such economic and operational advantages that virtually all the world’s methanol plants built after 1967 were the low-pressure plants. Along with the commercial success of the new low-pressure process, there emerged a renewed research activity aiming at the scientific understanding of the function of the copper-containing mixed-oxide catalysts. Furthermore, various reports summarized in a recent review by Denny and Whan (26) indicated that carbon dioxide has a significant effect on the synthesis rates, and Andrew recently reported (27) that the dependence of the synthesis rate on carbon dioxide concentration has a maximum around 1% C 0 2 ,a result derived from tests involving the commercial ICI catalyst Cu/ZnO/Al,O, . It appears probable, and will be substantiated below, that the effects of CO, and the “synergic promotion” of one catalyst component by another are but two facets of the formation of an active catalyst through the interactions among its components and the surrounding gas phase. It is the purpose of this article to review the recent observations and interpretations concerning the distribution of phases and elements, the physicochemical state of the catalyst components, the mode of activation of the reactants, and to show that a consistent picture is emerging that elucidates the function of the copper-based catalysts in the synthesis mechanism.

TABLE I

CujZnIAI Oxide Catalysts Used in the Synthesis of Methanol Composition" (wt. % 1

12: 62: 25 23 :46 :30 24: 38: 38 35:45 :20 53: 21: 6 60: 22: 8

2 2 3 2 1 1 1

66: 17: 17

2 3 3 1

C

1

64: 32:4

a

Reactantsb

CuO: ZnO: A1,0,.

Space velocity (hr-')

(kg liter-' h r - ' )

Company

so

10,000 10,000 20,000 12,000

3.290 2.086 2.5 0.7

50 50 100 50 50 70 50

40,000 9600 10,000 10,000 2Ood 10,000

0.5 0.5 0.3 0.9 4.75

BASF BASF CCI 1c1 Academic ICI 1CI ICI Academic Academic DuPont Academic

Temp. (-C)

Pressure (atm)

230 230 240 226 250 250 250 226 250 300 275 250

200 100

Yield

' 1, H, + CO + CO, ;2, H, + CO + CO, + C H I ; 3, CO + H, ; N, is sometimes used as a diluent. ' SNM-1 catalyst. Moles per hour.

Reference

12

I3 I4 15 16 I7 18 19

TABLE I1 CulZnlCr Oxide Catalysts Used in the Synthesis of‘ Methanol

Composition” (wt. %) I 1 :70: 19 15 :48 :37 31:38:5

33:31:36

40: 10:50 40: 40: 20 60:30: 10

Reactantsh

Temp. (‘C)

Pressure (atm)

Space velocity (hr-’)

Yield (kg liter-’ h r - ’ )

3 3 3 4 3 3 I 2 2 1

250 270 230 230 250 300 260 250 250 250

145 50 50 150 150 100 40 80 100

4000 10,000 10,000 10,000 10,000 10,000 10,000 6000 10,000 9800

1.95‘ 0.755 1.275 1.1 2.2 0.48‘ 0.26 0.77 2.28

Company

Reference

Power-Gas Corp. Jap. Gas-Chem. Co. BASF BASF Academic Academic

20

T, HFA

23 24

ICI ICI Metall-Gesellschalt

” CuO :ZnO : Cr,O,.

’ 1, H, + CO + CO, ; 2, H, + CO + CO, + CH, ; 3. CO + H, ; 4, CO + H, + 0, ;N, Kilograms per kilogram per hour.

is sometimes used as a diluent

21

22 17

25

METHANOL SYNTHESIS

II.

25 1

Catalyst Selection

Methanol synthesis from carbon monoxide formally consists of an attachment of three hydrogen atoms onto the carbon end and of one hydrogen atom onto the oxygen end of the CO molecule without the cleavage of the carbon-oxygen bond. The two bonding 71 orbitals of carbon monoxide, the lone pair (50) orbital or carbon, and the o orbitals of two hydrogen molecules are utilized to make three C-H sigma bonds, one 0-H sigma bond, and one additional lone pair orbital on the oxygen atom, schematically. H H-H

Ic~olH

H

\

H-c-Q-H /

H

The carbon-oxygen sigma (30) and the lone-pair oxygen (40) orbitals remain intact except for a change in electron repulsion. These features of the reaction impose the following requirements on the catalyst : 1. The catalyst must not cleave the carbon-oxygen sigma bond (- 360 kJ/mol). 2. The catalyst must activate carbon monoxide molecule so that hydrogenation can occur on both ends of the molecule. 3. The catalyst must be a fairly good hydrogenation catalyst which activates hydrogen molecules in a manner suitable for the above reaction. From the pattern of bonds broken and formed, this appears to require splitting of hydrogen molecules on the catalyst surface at some stage of the reaction. Although all three requirements must be satisfied by a good methanol catalyst, there is evidence that activation of carbon monoxide is more difficult than that of hydrogen. It is therefore instructive first to examine the ability of various catalysts to activate CO by nondissociative chemisorption. For metals there exists a clear-cut relationship between their position in the periodic table and their ability to chemisorb CO dissociatively. A division line has been established for ambient temperature by Broden et af. (28) between metals that chemisorb CO nondissociatively and those which have at least some crystal faces that split the CO molecule into surface carbon and oxygen. Table I11 shows a section of the periodic table wherein elements on the right-hand side of the borderline adsorb CO nondissociatively and vice versa. Broden et al. have also demonstrated on the basis of UV photoelectron spectroscopic measurements that the n-40 energy separation is affected by the bonding of undissociated carbon monoxide to the metal surface: the (71-40) energy gap increases with increasing strength of the carbon-to-metal

252

K . KLlER

TABLE I l l CO Chemisorption on Transition Metals"

VI A Cr

Mo W AMBIENT TEMPERATURE

\SYNTHESIS

I

TEMPERATURES

200-300°C

Dividing lines separate metals on the left which split CO from those on the right which adsorb CO nondissociatively. Broden et al. (1 976).

'

bond as is apparent from Table IV. The differences of the (71-4a) energy gap for metals on the two sides of the room-temperature borderline are only tenths of electron volt (10 kJ) however, and it is entirely expectable that the borderline will depend on temperature. At high temperatures, cracking of carbon monoxide will become more probable, and so the borderline will move to the right. Although no measurements similar to those by Broden et al. have been reported for elevated temperatures, the chemical behavior of various metal catalysts permits the conclusion that the division line between nondissociative and dissociative CO chemisorption lies, at 25O-30O0C, between nickel and copper, rhodium and palladium, and osmium and iridium. This high-temperature division line is also indicated in Table 111. At still higher temperatures, even copper and palladium will dissociate carbon monoxide. It must be emphasized that nondissociative chemisorption of carbon monoxide, which appears to be a necessary condition for its hydrogenation to methanol, is not a sufJicient prerequisite. For example, nickel (29) or gold

Weakening of the C-0 Mo,Fe 3.50

>

W 3.20

>

TABLE 1V Bond by Back-Donation Metal + CO (2~")'

Ru 3.15

>

Ni 3.08

>

Pd 2.90

>

Ir 2.75

>

Pt 2.60

Consequences: E(1x)increases more than E(4u) and A(1rr-4u) increases with back-bonding. The average values of A for different transition metals are in electron volts. The scatter ofA for different crystal faces is 0.2 eV (0.026 eV = 300 K ) . From Ref. 28.

METHANOL SYNTHESIS

253

(30) chemisorb CO nondissociatively at ambient temperature but they are not methanol synthesis catalysts. Of the metals on the right-hand side of the high-temperature borderline in Table 111, copper (31), palladium, platinum, and iridium (32) have been reported to be selective methanol synthesis catalysts, while certain forms of rhodium have been found catalysts of intermediate activity and selectivity for methanol (33,34).All these metals have the common property that they lie close to the high-temperature borderline in Table 111. Moreover, metals guiding the synthesis selectively to methanol are to the right, but not far to the right, of this boundary. This indicates that good methanol catalysts chemisorb CO with moderate strength, which is sufficient to perturb this molecule to enable it to react with hydrogen but insufficient to break it or any of the intermediates into fragments. Aside from depending on the position of the metal in the periodic table as outlined above, the catalytic activity of each individual metal further depends on its physical and chemical state. Therefore, while the nondissociative chemisorption of CO serves as a crude guiding principle for the selection of a catalyst that will potentially synthesize methanol, the optimum performance of the catalyst results from maximizing and stabilizing its form that is most active, and in the case of methanol also most selective, under the desired reaction conditions. Factors that have contributed to the development of the low-pressure methanol catalyst have been the stability of metal dispersion ; the choice of oxide “support”; the concentration of impurities such as alkali-, sulfur-, and chlorine-containing compounds ; preparation variables giving rise to different precursor compounds; calcination and reduction regimes ; etc. After a complex preparation procedure, pinpointing the “active form” of a catalyst appears to be a difficult task. Experimental efforts to find the active component of a given catalyst often spur controversies, and the history of the copper-based methanol catalyst is a prime example of a collection of “puzzling results,” claims and denials, and disagreements among the various workers in the field. The early investigations dating back to Sabatier, Patart, Frolich, and Huttig were outlined in Section I. After the successful development of the low-pressure process, debate has concentrated on the active component of the Cu/ZnO/Al,O, catalyst. Detailed description and analysis of this system are offered in the following text. The palladium, platinum, iridium, and rhodium catalysts have not been considered as of this writing for commercial applications because their activity was found to be significantly lower than that of the modern copperbased catalysts. While the relation between the position of metals in the periodic table and their ability to synthesize methanol has been established, a similar pattern is lacking for oxides (35). It can be anticipated that the requirements 1-3, listed earlier, must also be satisfied by oxide catalysts, but so far attempts

254

K . KLlER

have not been made to link the methanol-synthesis activity of oxides with their physicochemical properties such as basicity or acidity, metal -oxygen bond strength, energy gap and semiconductivity, work function and electron affinity, etc. There does appear to be consensus among the workers in the field that surface acidity of an oxide leads to the formation of dimethyl ether by a reaction consecutive to methanol synthesis.

111.

The Activity of Pure Copper Metal

The early conflicting reports on the activity of pure copper metal could not be reconciled without the simultaneous or concurrent measurements of activity, surface area, and surface composition. Moreover, it became evident that it is important to use unsupported copper as the reference material to avoid support -metal interactions that may influence the catalytic properties of the latter. In order to resolve the issue of the activity of pure copper metal, the following experiment has been carried out in the author's laboratory (36): Copper hydroxy nitrate [Cu,(OH),NO,] was precipitated from a solution of copper nitrate by a dropwise addition of sodium carbonate, carefully washed and stepwise calcined at temperatures not exceeding 350 C. X-ray diffraction showed the product of calcination to be crystalline CuO (tenorite) and Auger surface analysis was able to detect only copper and oxygen. Careful search for Na, N, Fe, C1, and S revealed that these elements were absent to within the detection limits of the technique, i.e., 0.02-1 at. % on the surface. The CuO catalyst was then loaded into a reactor in which many active catalysts were tested, reduced in a stream of 2% H, in N, at atmospheric pressure and 250'C, and subject to a standard methanol synthesis test by passing the synthesis gas CO/CO,/H, = 24/6/70 at 75 atm and 250°C at the rate 15 liters (STP)/hr over 16 g of the copper catalyst. The effluent was analyzed by an on-line gas chromatograph, and the carbon conversion to methanol was found to be less than 0.001% under the outlined conditions. No dimethyl ether was detected. However, small amounts of water (0.005%) and methane (0.036%) were found in the effluent gas. The catalyst was removed under nitrogen and subjected to X-ray diffraction, BET surface area measurement, and Auger analysis. X-ray diffraction determined copper to be in its normal fcc crystalline state with the cubic lattice constant that of bulk copper metal (36). The BET surface area was found to be 0.43 m2/g, so that the 16 g of the tested catalyst exposed an area of 6.88 m2 to the synthesis gas. A scanning electron micrograph of this copper specimen is shown in Fig. 3. The analysis of argon isotherms indicated that the metal was not

METHANOL SYNTHESIS

255

FIG. 3. Scanning electron micrograph of copper used in methanol synthesis study (40). [Reprinted with permission from J . Cutal. 57, 339 (1979). Copyright (1979) Academic Press, New York.]

microporous (37). A sample of the used copper catalyst was subject, after a brief exposure to air, to Auger analysis, which found no other elements than copper and oxygen on the catalyst surface. The catalyst was then reduced in the Auger spectrometer at 250”C, and only copper was found in its surface. This set of experiments establishes that pure copper metal, free of surface impurities, yields less than l o - * kg of methanol per square meter of the catalyst per hour under the standard conditions outlined above. Such a yield is far below the specific activity of “supported” copper-based catalysts, e.g., 3.63 x kg CH,OH m-’ hr-’ for the Cu/ZnO = 30/70 catalyst (39), and shows that copper metal is a very poor catalyst for methanol synthesis at 75 atm at 250°C.

FIG.4. Transmission electron micrographs of pure zinc oxide prepared by calcination of Zn,(OH)6(C0,), (40). [Reprinted with permission from J . Cutaf. 57, 339 (1979). Copyright (1979) Academic Press, New York.]

METHANOL SYNTHESIS

IV.

257

The Activity of Pure Zinc Oxide

In a similar experiment, 2.5 g of zinc oxide prepared by precipitation from zinc nitrate solution by sodium carbonate, calcination, and attempted reduction under similar conditions as previously employed, gave a catalyst of surface area of 40 m2/g, which yielded less than kg of methanol per square meter of the catalyst per hour under the standard conditions used for the testing of the copper catalyst. The zinc oxide was in its wurtzitic crystal modification as in most laboratory as well as industrial catalysts, was free of surface impurities, and had a morphology shown in Fig. 4. Details of the pore structure of this catalyst are given in reference (38). The experiment outlined above demonstrates that like metallic copper, pure zinc oxide is a very poor methanol catalyst at 75 atm and 250°C. This finding contrasts with the relatively high activity of zinc oxide at pressures exceeding 200 atm and temperatures above 350°C. Further, since the activity of zinc oxide has been found to be higher than that of other oxides such as alumina or chromia, no single-component catalyst, oxide, or metal is known at the present time that would effectively catalyze methanol synthesis at low temperatures and pressures, i.e., below 250°C and 100 atm. It is the synergic promotion in multicomponent catalysts that brings about greatly enhanced activity at low temperatures.

V.

The Binary Copper-Zinc Oxide System

After the introduction of the low-pressure processes using the Cu/ZnO/ A1,0, and Cu/ZnO/Cr,O, catalysts summarized in Tables I and 11, the question arose as to what causes the high activity of these catalysts compared to that of their individual components. Herman el al. (39) investigated the relation of the activity of the single components and their binary, as well as ternary, combinations to various physical and chemical characteristics of these catalysts. They have established that the low-temperature-low pressure activity resides in the binary copper-zinc oxide system and have optimized the catalyst composition for maximum methanol yields. The catalysts were characterized in all stages of preparation and before as well as after use by X-ray diffraction, electron microscopy, Auger/ESCA spectroscopy, optical spectroscopy, BET surface-area measurements, and chemisorption measurements (38-43). The principal findings are summarized in the following sections.

258

K. KLIER OF THE Cu/ZnO BINARY CATALYSTS A. PREPARATION

The catalysts were coprecipitated from CujZn nitrate solutions by sodium carbonate, keeping the initial nitrate concentration constant and varying the Cu/Zn ratio from O j l O O to lOOj0. In the zinc-rich compositional range hydroxy carbonates were precipitated, while in the copper-rich range the main product was copper hydroxy nitrate Cu,(OH),NO, . Upon calcination in air at 350"C, these precursor compounds turned into a heterophase mixture of tenorite CuO and wurtzite ZnO, and upon following reduction in 2% H,/N, gas, the cupric oxide was converted to copper and the zinc oxide remained unreduced in its wurtzite modification. Although the precipitate precursors did not influence the phase composition of the calcined CuO/ZnO or the reduced CujZnO catalysts, their composition determined the interdispersion and morphology of the final catalyst. A very intimate mixture of small copper and zinc oxide particles was obtained by calcination and reduction of the (Cu, Zn),(OH),CO, precursor in which copper and zinc initially occupy the same type of lattice sites and are interdispersed on atomic scale. It was this preparation that led to the formation of the most active binary catalyst (39). PHYSICAL CHARACTERISTICS OF THE CuiZnO CATALYSTS

B.

A quantitative X-ray diffraction study of the Cu/ZnO catalysts (41) has revealed that varying amounts of copper appear in an X-ray amorphous r

1

0

10

7

-

1

I

I

I

I

I

I

I

60

70

80

90

100

1 0

0 8

z -

I

0 6

\

u

0 4

0 2 0 0 20

30

40 %

50

Copper

Content

FIG.5 Comparisons of the crystalline Cuo (0) and ZnO (0) components in the reduced as deterCu/ZnO catalysts, I,, with those in the reduced mechanically mixed composites, IMM, mined by X-ray powder diffraction; I is the integrated intensity. Shaded area shows the fraction of amorphous copper (41). [Reprinted with permission from .IPhys. . Chern. 83, 31 18 (1979) Copyright (1979) American Chemical Society.]

259

METHANOL SYNTHESIS

TABLE V BET Argon Surface Areas of the CujZnO Catalysts" Catalyst composition CuO/ZnO 0/100 2/98 l0/90 20/80 30/70

Surface area (m2/g) of Reduced

Catalyst composition CuO/ZnO

Used

Surface area (mZ/g)of Reduced

40/60

25.2 28.9 27.0 30.0 37.1

50jS0

67/33 1 00/0

6.3 I .4

Used 13.5 10.3 8.6 -

I

" From Ref. 39.

form, as demonstrated in Fig. 5. It is evident that the amorphous state of copper is induced by the presence of zinc oxide, since all copper metal is crystalline where zinc oxide is absent. Since the largest amounts of amorphous copper were found in the most active catalysts, it became important to further identify the location and state of this component of the catalyst. The BET surface areas are summarized in Table V. It is seen from the data in Fig. 5 and Table V that no dramatic increase of the total catalyst surface area occurs in samples with large amounts of amorphous copper, contrary to what would be expected if amorphous copper were dispersed as very small particles. The answer to the question of the location of amorphous copper was provided by a combination of structural and elemental analyses of the Cu/ZnO binaries in the scanning transmission electron microscope (STEM) (40). Significant amounts of copper were found in the zinc oxide phase of the catalyst; these are compared with the amount of X-ray amorphous copper in Table VI to show that most if not all the amorphous copper is located in the zinc oxide crystallites. Furthermore, this amorphous copper is not accumulated on the zinc oxide crystallite surface because the amount of oxygen and zinc determined by surface analysis in ESCA/Auger spectrometer is that expected from a uniform interdispersion of 0,Zn, and Cu throughout the sample. Hence the combined STEM, ESCA/Auger, and surface-area measurements indicate that the amorphous copper is dispersed, or dissolved, in the zinc oxide phase. This result was corroborated by optical studies and further refined by chemisorption measurements discussed below. The electronic interaction between the catalyst components is best exemplified by its color and optical spectra. For example, the very active binary catalyst Cu/ZnO = 30/70 has a pitch black color and although it is composed of crystallographically identifiable copper and zinc oxide, its optical spectrum is not a superposition of the spectrum of copper metal and zinc oxide, but rather comprises a very intense continuous absorption band in the visible part of the spectrum that contains no trace of the characteristic

260

K . KLIER

TABLE Vf Extent of Cu/ZnO Solid Solution Formation in the Reduced Coprecipitated Catalyst, us Determined by X-Ray Powder Dzzruction und Electron Microscopy Cu PA)in solid solution Nominal catalyst composition"

X-ray diffraction"

Electron microscopy'

Cu (%) in solid solution Nominal catalyst composition"

X-ray diffractionh

Electron microscopy'

30170 40/60 50150 67/33 80/20

16.8 F 1.5 12.6 f 1.5 13.0 f 2.0 15.4 f 3.0 14.2

15.1 k 3.7 5.0 f 3.7 8.7 f 3.6 11.6 5 2.8

~~~~~~~~~~

2/98 5/95 10/90

15/85 20/80

2.0 2.5 4.8 8.1 10.5

zinc oxide absorption edge at 25,800 cm-' in the near UV. There is an absorption edge in the near IR which appears at frequencies lower than those of the copper metal d-hump (41). This spectrum was attributed by Klier and co-workers (39,41) to the copper solution in the zinc oxide, and a physical model was put forward which is depicted in Fig. 6. The disappearance of the zinc oxide absorption edge was explained as originating from the overlap

36 ICuI B A N 0

BAND OF ZnO

36 Iznl BAND

I__,

M - 0 Distance

___,

M - 0 Oistance

FIG.6. A diagram showing the band spectra of (a) ZnO and (b) Cu/ZnO solid solution (39). [Reprinted with permission from J . Cutal. 56,407 (1979). Copyright (1979) Academic Press, New York.]

METHANOL SYNTHESIS

26 1

of the copper solute spd band with the valence-band edge of zinc oxide, the appearance of the near-IR edge was proposed to be due to the transition between the copper dsp band and the conduction band of the zinc oxide, and the high intensity of the transition was attributed to the high degree of allow that also is characteristic of the UV band-to-band transition in the pure zinc oxide. The copper solution in the zinc oxide characterized by the outlined analytical andphysical methods was found to exist only after mild reduction of the calcined catalyst. Before reduction, the solubility of CuO in ZnO is limited to 4-6% (44,45) and after more severe reduction, the optical spectra begin to resemble a superposition of those of pure copper metal and zinc oxide. Hence the black solute phase is metastable and does not appear to be the final product of reduction. For this reason, the dispersed copper species were assigned the valence state + 1 ; Bulko rt al. (41) visualized these copper species not as isolated Cu' ions but rather as electron-deficient copper atoms with strong electronic overlap with the host zinc oxide lattice, particularly with neighboring oxygens whose orbitals dominate in the valence band of zinc oxide.

C. PARTICLE SIZEAND MORPHOLOGY Using the preparation method described above, the particle sizes of the crystalline components of the Cu/ZnO catalyst were strongly determined by the precursor composition and structure. Figure 7 shows that small particles of CuO, copper, and zinc oxide formed by calcination and reduction of hydroxy carbonate precursors at concentrations Cu/ZnO = O / 100 to 30/70, while large particles of all components result from calcination and reduction of precursors that contain copper hydroxy nitrate at compositions Cu/ ZnO = 40/60 to lOOj0. The rather sudden change of particle size between compositions Cu/ZnO = 30/70 and 40/60 is accompanied by interesting changes of morphology which are demonstrated in the electron micrographs of the Cu/ZnO = 30/70 (Fig. 8) and 67/33 (Fig. 9) catalysts (40). In the zinc-rich region, zinc oxide forms intertwined lacework with the hexagonal crystal axis parallel to its longer dimension, while in the copper-rich composites, zinc oxide crystallizes as well-shaped hexagonal platelet crystals with their major crystal axis perpendicular to the platelet plane. The hexagonal platelet morphology of zinc oxide is believed to originate during the thermal decomposition of hydroxy carbonates and hydroxy nitrates of zinc and copper. Since the hydroxy carbonates that are rich in zinc decompose at slightly lower temperatures than hydroxy nitrate of copper, the zinc oxide crystals can epitaxially nucleate on the pseudohexagonal (001) plane of

262

K. KLIER

50

40

10

0

10

20

30 k Copper

40

50

60

70

C a n tent

FIG. 7. The average particle sizes of copper (O), its CuO precursor (O), and zinc oxide ( A ) determined from X-ray diffraction line broadening (41). [Reprinted with permission from J . Phys. Chem. 83, 3118 (1979). Copyright (1979) American Chemical Society.]

Cu,(OH),NO,, giving rise to the hexagonal ZnO crystals. In subsequent stages of calcination the copper hydroxy nitrate decomposes to form separate particles of CuO. Further interesting details are also apparent from the electron microscope investigations of Mehta et al. (40). Figure 8 shows that copper particles are in intimate contact with zinc oxide in the 30/70 catalyst. Moreover, these particles are in epitaxial registry with the zinc oxide crystallites, as demonstrated by the selected area diffraction pattern in Fig. 10. The spot pattern indicates that the whole zinc oxide lacework is single-crystal oriented with its hexagonal axis parallel to the edge of the particle assembly marked in Fig. 8. Furthermore, the copper crystallites are in definite crystallographic orientation with respect to the zinc oxide and to each other. The Cu/ZnO = 30/70 catalyst is thus composed of zinc oxide exposing its (1010) prism planes with the copper particles being spherical caps facing the underlying zinc oxide with their (211) planes. In addition to these structural features, a large amount of copper (15% in the 30/70 catalyst) was found to be located in the zinc oxide phase by 6-nm spot analyses in STEM (40). This amount coincides, within analytical error, with the amount of amorphous copper determined by X-ray diffraction as shown in Table VI.

FIG.8. (a) Transmission electron micrograph of a Cu/ZnO = 30/70 binary catalyst (40); 60 8, copper spheres are placed on crystalline zinc oxide network. (b) Dark field image of the copper crystallites in the area shown in the bright field image (a) obtained using the [TI I] reflection of copper. [Adapted with permission from J . Curd 57, 339 (1979). Copyright (1979) Academic Press, New York.]

FIG.9. (a) Transmission electron micrograph of Cu/ZnO = 67/33 catalyst. The hexagonal platelet particles are zinc oxide exposing its basal plane (0001) and the dark irregularly shaped particles are metallic copper (40). (b) Convergent beam diffraction pattern in STEM from the hexagonal ZnO particle shown in (a) above demonstrating that the sixfold (0001) crystal axis is perpendicular to the plane of the platelet. [Adapted with permission from J . Cuful. 57, 339 (1979). Copyright (1979) Academic Press, New York.]

I

i

8

FIG. 10. (a) Selected area diffraction pattern of the reduced Cu/ZnO = 30/70 catalyst from an area shown in the upper part of Fig. 8; (b) schematic diffraction pattern showing the spots from ZnO ( 0 )and Cu (u).The circles with the solid lines and the circles with dashed lines are the more intense diffraction rings from ZnO and Cu, respectively (40).[Adapted with permission from J . Catal. 57, 339 (1979). Copyright (1979) Academic Press, New York.] 265

266

K . KLIER

0 ZINC, @

COPPER,

0 OXYGEN

FIG.11. A view of copper and zinc oxide crystallites with dispersed copper ions in the binary Cu/ZnO catalyst = 30/70 derived from diffraction and characteristic X-ray emission analysis in TEM and STEM.

The structural and analytical findings concerning the Cu/ZnO = 30/70 catalyst are summarily represented in Fig, 1 1. It is seen that this catalyst has several parts t o be considered as potential seats of catalytic activity: the copper crystallite surface, the zinc crystallite surface, the copper solute in zinc oxide, and finally the interface between the copper and zinc oxide particles. To resolve which constituent of this catalyst is mainly responsible for its activity, it is helpful to examine the structural and compositional features of the second morphologic range at compositions Cu/ZnO = 40/60 and at higher copper concentrations. The ZnO particles are hexagonal platelets exposing their (0001) basal planes as major faces, and copper is in the form of large irregular blunted particles. The copper and zinc oxide particles are well separated and there appears little contact between the two phases (see Fig. 9). However, the zinc oxide particles again contain significant amounts of dispersed copper (40). Since Cu/ZnO = 67/33 also is an active catalyst, the copper solute appears to be a species of prime importance, and the contact between copper and zinc oxide crystals seem to play at most a secondary role in the formation of the active catalyst. This issue will be discussed further in conjunction with the chemisorption and activity measurements.

D. SURFACE ANALYSIS BY XPS-AUGERSPECTROSCOPY Elemental surface composition of the same Cu/ZnO catalysts as those investigated by other methods and tested for methanol synthesis was determined by X-ray photoelectron (XPS) and Auger spectroscopy and reported by Herman rt ul. (39).The catalysts show surface concentrations of Cu, Zn, and 0 that roughly correspond to their nominal elemental compositions. Given the semiquantitative nature of electron spectroscopic analysis and the

METHANOL SYNTHESIS

267

effect of particle shape on the effective analytical cross sections for the individual elements, subtle changes of concentration ratios cannot be determined reliably. However, it is evident that there is no dramatic redistribution of elements between the bulk and the surface of the catalyst particles such as might be expected to occur if copper were to form overlayers on the zinc oxide surface or if zinc oxide were to form a “coating” on the copper particles. In principle, Auger/XPS analysis can be used for the determination of valence states of the catalyst components in various stages of reduction and use in the synthesis reaction. These refined analyses are in progress in the author’s laboratory and their results will be reported in future communications. At this time the analysis indicates that no cupric ions or zinc metal atoms are present at any stage of reduction or synthesis at temperatures including and below 250°C. The possible surface contaminations were carefully followed by AugerXPS analysis. Similarly, as with the copper catalyst described earlier (Section 111) the Cu/ZnO binaries were free from alkali metals, iron, chlorine, and sulfur, and contained only small amounts of carbon after the use in catalytic reactor (39).The latter result indicates that reactants, intermediates, and the product are adsorbed with moderate strength, a feature that is desirable for all efficient catalysts. E.

SURFACE AREASOF THE INDIVIDUAL CATALYST COMPONENTS

Many researchers have argued that the activity of the low-temperature methanol catalysts is associated solely with the copper surface (27, 46, 47). Although the experiment using unsupported copper (Section 111) indicates otherwise, it is still desirable to determine the copper surface area in the supported catalysts. Among the several methods offered in the literature, low-temperature oxygen chemisorption has been most widely used. Most investigators agree that oxygen irreversibly chemisorbed on copper at temperatures lower than 140 K occupies an area of around 0.16 nm2 per oxygen atom, based on 0.14 nm2 for argon area (48), which corresponds approximately to a formula Cu,O. The method has been used to measure surface areas of pure copper (48, 49), and of copper in the Cu/ZnO/Al,O, catalysts (27, 50). Parris and Klier (43) used oxygen chemisorption to determine copper areas in the binary Cu/ZnO catalysts. Oxygen chemisorbed irreversibly at 78 K was found to cover an area of 15.2 A2 per oxygen atom on pure copper, and this value was also assumed to hold for copper surface in the Cu/ZnO composites. No irreversible oxygen was found on pure zinc oxide. The zinc oxide area was then calculated as the difference between the total

268

K. KLIER

TABLE VII Area of Copper Metal and Zinc Oxide in the Binury CuJZnO Catalysis"

Catalyst composition Cu/ZnO

o/ 100 10190

20/80 30/70 "

Surface area (m2/gcatalyst)

Cu

ZnO

Total

Catalyst composition Cu/ZnO

-

24.1 30.1 27.6 23.3

24.1 35.0 36.5 39.3

40160 50/50 67/33 lOOj0

4.9 8.9 16.0

Surface area (m2/gcatalyst) Cu

ZnO

Total

6.5 5.7 3.9 0.62

8.5 5.4 3.5

15.0 11.1 7.4 0.62

-

Parris and Klier ( 4 3 ) .

BET area (Table V) and the copper area from oxygen chemisorption. Table VII summarizes the copper and zinc oxide areas so determined for the whole compositional range. The oxygen chemisorption method suffers from the uncertainty that some oxygen may be adsorbed on the copper solute and on defects in the zinc oxide surface that are formed only in the presence of copper. There is indirect evidence from a comparative study of carbon monoxide and oxygen chemisorption, however, that this is not the case and that oxygen titrates only the copper metal surface.

F. CHEMISORPTION OF CARBON MONOXIDE AND HYDROGEN Quantitative and qualitative changes in chemisorption of the reactants in methanol synthesis occur as a consequence of the chemical and physical interactions of the components of the copper -zinc oxide binary catalysts. Parris and Klier (43) have found that irreversible chemisorption of carbon monoxide is induced in the copper-zinc oxide catalysts, while pure copper chemisorbs CO only reversibly and pure zinc oxide does not chemisorb this gas at all at ambient temperature. The CO chemisorption isotherms are shown in Fig. 12, and the variations of total CO adsorption at saturation and its irreversible portion with the Cu/ZnO ratio are displayed in Fig. 13. The irreversible portion was defined as one which could not be removed by 10 min pumping at Torr at room temperature. The weakly adsorbed CO, given by the difference between the total and irreversible CO adsorption, correlated linearly with the amount of irreversibly chemisorbed oxygen, as demonstrated in Fig. 14. The most straightforward interpretation of this correlation is that both irreversible oxygen and reversible CO adsorb on the copper metal surface, The stoichiometry is approximately CO : 0 = 1 :2 , a ratio obtained for pure copper, over the whole compositional range of the

CO CHEMISORPTION ON C u / Z n O

CATALYSTS

(D

0 8 x

t-m

3 a $ 6

.2 V

cn

-

0

w m

4

LL

m 0 n a 0 V LL

2

0

t-

z

0 2

5

20

0

60

40

80

100

120

PRESSURE (Torr)

FIG.12. Carbon monoxide chemisorption isotherms at 25°C on the binary Cu/ZnO catalysts. The labels at the individual isotherms denote the molar composition Cu/ZnO (43). 1-

I

I

I

I

I

I

(D

0

4

x

I-

m -J >

a $a 0

3

LL

0

N

E

(L

w

a

n w m

2

LL 0 0 I n

a

0

u

1

LL

0

I/

IRREVERSIBLE IRREVEPSIBLE

J

fn w J 0

I

0

10

20

30

40 cu Cu t Z n

50

60

70

80

90

100

x 100

FIG. 13. The dependence of the carbon monoxide saturation adsorption (total) and irreversible adsorption (irreversible) on the Cu/ZnO ratio in the binary copper-zinc oxide catalysts (43).

270

K . KLIER 4 REVERSIBLE CO VS IRREVERSIBLE 0 2 CHEMISORPTION ON C u /

(D

ZnO

CATALYSTS

0

;3VI

a 0

L u

E

\

2 -

A

0

-I 1 w

0

I 0 2 (IRREVERSIBLE)

2

3 MOLII~

4

CATALYST x 1 0 6

FIG. 14. The relation between the amounts of weakly chemisorbed carbon monoxide and irreversibly chemisorbed oxygen, indicating that these two adsorbates are a measure of copper metal surface area (43).

Cu/ZnO catalysts. Furthermore, sites other than copper metal must be available for the irreversible CO, and it was suggested by Parris and Klier that these are the copper solute sites in the zinc oxide surface. The plot of the surface coverage by irreversible CO against the concentration of copper solute shown in Fig. 15 has a linear portion for each morphology of zinc oxide. Therefore, if the irreversibly bound CO molecules titrate the surface copper solute sites, the basal (0001) ZnO planes tend to accumulate more of the solute copper atoms than the prism (lOT0) planes. Since the basal plane is formally electrostatically charged while the prism plane is not, the accumulation of copper is consistent with the notion that the solute copper species are also charged and tend to neutralize the excess charges of the basal planes of zinc oxide. These copper ions are then considered to be the active centers for the irreversibly chemisorbed carbon monoxide. The catalysts containing the hexagonal zinc oxide platelets are more complex than the catalysts with prism morphology for two reasons: first, in addition to the basal planes, the platelets also expose the prism faces, and their chemisorption activity will be a superposition of the two; and second, the basal planes may be composed of oxygen- and zinc-rich patches as has often been reported for hexagonal crystals (51,52). Despite this complexity, however, catalysts with hexagonal-platelet zinc oxide morphology may be of practical value in applications where low-surface-area catalysts are desired, such as in liquid-

27 1

METHANOL SYNTHESIS

TITRATION OF Cu SOLUTE SITES BY CO 0

n w m E E n

2

a

3

m 0-

$ e

> *

w -

E L

I

3

% 2 I - $

3 I a 5

O f

0

2

4

6

CONCENTRATION OF Cu SOLUTE SITES ON Z n O SURFACE, [G-ATOM PER GRAM OF CATALYST ( X I O ~ ) I

FIG.15. The dependence of the amount of irreversibly chemisorbed carbon monoxide on the concentration of amorphous copper found by X-ray diffraction (Fig. 5) and by STEM (Table VI). The (lOT0) prism planes of zinc oxide are exposed in the Cu/ZnO = l0/90, 20/80, and 30/70 catalysts and the (0001) basal planes of ZnO are exposed in the Cu/ZnO = 40/60, 50/50, and 67/33 catalysts. The different slopes for the two morphologies indicate slightly different concentrations of copper atoms on the prism and basal surfaces (43).

phase methanol synthesis (53).Of these composites, the Cu/ZnO = 67/33 of surface area of only 7-9 m2/g proved quite an effective irreversible sorbent for CO and, as is shown below, a catalyst of the highest specific activity.

G. ACTIVITY PATTERNS All the binary Cu/ZnO catalysts were found highly selective toward methanol without DME, methane, or higher alcohols and hydrocarbons detected in the product by sensitive gas chromatographic methods (39). Several of the composites were also found to be very active when subjected to a standard test with synthesis gas CO/CO,/H, = 24/6/70 at gas hourly space velocity of 5000 hr- pressure 75 atm, and temperature 250°C. The activities, expressed as carbon conversions and yields, are summarized in Table VIII. The end members of the series, pure copper and pure zinc oxide, were inactive under these testing conditions, and maximum activity was obtained for the composition Cu/ZnO = 30/70. The yields per unit weight, per unit area of the catalyst or the individual components, turnover rates per site titratable by irreversible oxygen and by irreversible carbon monoxide, are graphically

272

K . KLlER TABLE VIII The Catalytic Testing Results for the CulZnO Systems" Yield

Catalyst compositionh CuO/ZnO/M,O,

o/ lOOj0 2/98/0 10/90/0 20/80/0 30/70/0 40/60/0 50/50/0 6713310

1 oo/o/o 60/30/10' 60/30/ 1Od 60/30/ 10'

Carbon conversion 0 0.7 1 .o

10.2 51.1 9.6 11.3 21.8 0 40.0 17.0 47.0

kg liter-

hr-

kg kg- ' hr0 0.03 0.02 0.24 I .35 0.18 0.20 0.41 0 I .52 0.58 1.01

0 0.02 0.02 0.22 1.10 0.21 0.25 0.48 0 0.95 0.45 1.17

'

k g m - 2 hr-' ( x lo5) 0 0.10 0.70 0.80 3.63 1.33 1.94 4.76 0 5.82 1.73 6.47

'

Obtained at 250"C, 75 atm, and GHSV = 5000 hr- with a synthesis gas of H2/CO/C0, = 70/24/6 ~ 0 1 %From . Ref. 39. " Weight percent as the oxides. M = AI, prepared from the acetates. M = Al, prepared from the nitrates, tested at 100 atm. ' M = Cr. tested at 100 atm.

represented in Fig. 16. Aside from showing that the active catalyst indeed requires the simultaneous presence of zinc oxide and copper, these graphs indicate that no single component surface is responsible for the activity. For example, should only copper metal or only zinc oxide be the active component, the activities per unit area of copper or zinc oxide would be independent of the composition. The reasons for the lack of constancy of the specific activities and turnover frequencies defined and presented in various ways in Fig. 16 must be sought, therefore, in the interactions of the catalyst components that give rise to the active catalyst. Several observations have already indicated that the amorphous copper solute induces irreversible chemisorption of carbon monoxide, and it therefore behooves us to correlate the methanol yields with the surface concentration of irreversibly chemisorbed CO. The best understanding of the synthesis pattern is provided with the help of the representation in Fig. 17. The turnover rate, defined as the number of CO molecules converted to methanol in 1 sec on one site titratable by irreversibly bound carbon monoxide, depends smoothly on the concentration of these sites for each zinc

273

METHANOL SYNTHESIS

.c

8 IL

0'05 0

L

rV

:

0

"N

O LD LD-

OO

.";'IJ//q n "

0

20 40 60 80 ATOMIC PERCENT OF COPPER IN Cu/ZnO BINARY CATALYSTS

100

FIG.16. Catalytic activity expressed as yield of methanol per unit weight, per unit surface area of the Cu/ZnO catalyst, per unit surface areas of the components Cu and ZnO, and as turnover numbers over copper metal sites titratable by irreversible adsorption of oxygen and over Cu/ZnO solute titratable by irreversible adsorption of carbon monoxide: (a) g MeOH/g catalyst hr- ; (b) g MeOH/m2 catalyst hr- ; (c) g MeOH/mZ Cu hr- ' ; (d) g MeOH/mZZnO hr- ;(e) turnover rate per oxygen site; (f) turnover rate per carbon monoxide site.

'

oxide morphology, i.e., for the prism morphology in the compositional range Cu/ZnO = l0/90 to 30/70, and for the hexagonal platelet morphology for Cu/ZnO = 40/60 to 67/33. At the same time, the turnover rate varies as the second to third power of the surface site concentration, indicating that three to four sites are involved in the conversion of each CO molecule. This is possible if one site is utilized for activation of carbon monoxide and several other sites for activation of hydrogen. Since the sites titrated by irreversibly adsorbed CO are fairly diluted ( 1 6 x 1017/m2ZnO for the prism morphology and I1 x 10'*/m2 ZnO for the basal morphology; see Fig. 17), the activated reaction components must have a significant lateral mobility to reach the site of reactive encounter.

274

K. KLlER

AMOUNT OF IRREVERSIBLY ADSORBED

co

(rnol/rn'

Z n o ) x 106

FIG.17. The dependence of turnover rates of methanol synthesis (at 250"C, 75 atm) over sites binding carbon monoxide irreversibly upon the concentration of these sites on ZnO surface (43).

To summarize the qualitative findings, the methanol synthesis activity in the binary Cu/ZnO catalysts appears to be linked to sites that also irreversibly chemisorb CO and not to sites that adsorb CO reversibly. Since irreversible adsorption of CO follows linearly the concentration of amorphous copper in zinc oxide, these sites are likely to be that part of the copper solute that is present on the zinc oxide surface. No correlation of the catalyst activity and the copper metal surface area, titrated by reversible form of CO or by oxygen, could be found in the binary Cu/ZnO catalysts (43). In contrast with this result, it has been claimed that the synthesis activity is proportional to copper metal area in copper-chromia ( 4 3 , copper-zinc aluminate ( 2 3 , and copper-zinc oxide-alumina (46) catalysts. In these latter communications (27,46,47),the amount of amorphous copper has not been determined, and obviously there is much room for further research to confirm one or another set of results and interpretations. However, in view of the lack of activity of pure copper metal quoted earlier, it is unlikely that the synthesis activity is simply proportional to the copper metal surface area in any of the low-temperature methanol-synthesis catalysts.

H. REACTION KINETICS IN

THE

PRESENCE OF CARBON DIOXIDE

It has been known for some time that a small amount of carbon dioxide in the CO/H, synthesis gas acts as a promoter or prevents catalyst deactivation

METHANOL SYNTHESIS

275

(27, 39). This effect was attributed to the ability of CO, to keep the catalyst in its intermediate oxidation state such as Cu(1) (39), or to prevent reduction of zinc oxide with concomitant formation of brass (27),or to the reversal of the Boudouart reaction

2co=co, + c

(4)

that would cause carbon fouling of the catalyst. However, by Auger spectroscopy Simmons (54) found very little carbon and no carbidic species on the Cu/ZnO catalyst exposed to the C0,-free CO/H, synthesis gas and Chatikavanij ( 5 5 ) did not detect any traces of methane, which would be expected to form by hydrogenation of surface carbon, when running the synthesis in CO/H, = 30/70 synthesis gas. Hence, the Boudouart reaction is an unlikely cause of catalyst deactivation in the C0,-free synthesis gas. That the C 0 2 enhancement of the synthesis rate is a true chemical promotion effect of CO, and not a kinetic effect that would be produced by an addition of any inert has been established by replacing CO, by Ar, whereupon the CO conversion rate decreased by a factor 4.5 (56). In pursuing the effects of CO, further, it was found that there is an optimum in the COJCO ratio at which the synthesis runs at maximum rate (56, 27). At high concentrations, CO, has a gradually retarding effect on the synthesis (27, 56). The influence of the CO,/CO ratio on the synthesis has been embedded in kinetic equations only recently. The early kinetic equations for the highpressure ZnO/Cr,O, catalyst did not contain a C0,-dependent term at all, perhaps because the effects of CO, were not significant when zinc chromite catalysts were used; Natta et al. (57) proposed the rate equation for the ZnO/Cr,O, catalyst at temperatures 300-360°C as follows : r=

2 2 "?COPCOYH,PHz

(A +

B?COpCO

-

?CH30HPCH30H/Keq

+ CYH2PHz + DYCH,0HPCH30H)3

(5)

where r is the rate of methanol synthesis; y i is the fugacity coefficient of species i; pi is the partial pressure of species i ; K,, is the equilibrium constant of carbon monoxide hydrogenation to methanol; and A, B, C, and D are empirical constants. Equation (5) was derived under the assumption that the rate-controlling step of the synthesis is the trimolecular reaction of carbon monoxide with two hydrogen molecules in the adsorbed phase. Following Natta's publication of his data and kinetic treatment, a number of empirical as well as model-derived rate equations for methanol synthesis have been proposed, some of which have demonstrated that Natta's original data could be fitted with several kinetic models. These developments were summarized by Denny and Whan (26)up to 1977.

276

K . KLIER

When the synthesis was carried out with carbon dioxide-rich synthesis gas, however, the comparison of measured rates with those calculated using kinetic equation ( 5 ) or its modifications (58) showed a substantial disparity (59), and it was realized that carbon dioxide-dependent terms had to be incorporated in order for the kinetic equations to be useful for process design. This was accomplished by Bakemeier et al. (60)who obtained the following rate equation, once again for the ZnO/Cr,O, catalyst:

In equation ( 6 ) the quantities A , E , n, m , D, and F are semiempirical parameters. It can be seen that this rate equation predicts that the methanol yield would decrease as the CO, partial pressure is increased and would drop to zero for synthesis gas that contains carbon dioxide only. Hence Eq. (6) describes a process in which carbon dioxide is a retardant but not at all a reactant or a promoter. In 1973, Leonov et al. (61) introduced a kinetic rate equation for the lowpressure copper-zinc oxide-alumina catalyst for temperatures between 220 and 260°C. The rate equation was proposed to be (7 ) where k is the rate constant for the forward reaction and K , , is the equilibrium constant. Similar to the early kinetic studies with the high-pressure ZnO/ Cr,O, catalysts, there are no C0,-dependent terms in Eq. (7) for the lowpressure synthesis. Denny and Whan (26) also reviewed various contrasting reports on the effects of CO, on the synthesis and emphasized that any complete kinetic expression should include a term involving the partial pressure of CO,. A rate equation that does contain an empirical C0,-dependent term for the copper-zinc oxide-alumina catalysts has been presented in 1980 by Andrew (27) in the form r

0.2-0.6

= kPC0

0 7

PH;@co,

Although the function (Dco2 was not explicitly determined, it was reported that the rate of methanol synthesis reached a maximum at the C0,:CO partial pressure ratio around 0.01-0.03 and decreased as the CO, pressure was further increased. It was also indicated that the methanol synthesis rate would decline at very small concentrations of CO,. Behavior like this was

METHANOL SYNTHESIS

277

also reported to be the property of the binary Cu/ZnO catalysts (39), and hence it seems established that an optimum concentration of CO,, or a ratio of CO, and CO concentrations, exists at which the low-pressure synthesis runs at its maximum rate. A feature common to all kinetic equations summarized above is that the forward rate is proportional to a power of the product pco . p i 2 . When the reaction is carried out a low conversions, the reverse reaction is negligible, and the dependence of the synthesis rate on the product p c o . p & , such as illustrated in Fig. 18 for the Cu/ZnO/Cr,O, catalyst, is strong evidence that several sites are involved in one molecular synthesis step. Klier et al. (56)presented a kinetic model based on previous observations of the physical and chemical characteristics of the binary Cu/ZnO catalysts and put forward equations that quantitatively account for the CO,/CO dependence of the synthesis rates. Their model was required to accommodate the following observations : 1. The increase of methanol conversion rates from the CO/H, synthesis gas upon additions of 1-2% CO, can also be obtained by additions of hundredths of percent of oxygen. 2. A catalyst that has been reduced by hydrogen at 250 C can be further reduced by CO with the appearance of CO, . When this CO-reduced catalyst is exposed to CO,, also at 250'C, it is reoxidized and CO, is converted to CO. The CO/CO, equilibria are established in the time frame of hours (42). 3 . The activity of the catalyst increases as a high power of the concentration of the amorphous copper, as is seen from Fig. 17, and does not depend monotonously on the surface area of crystalline copper particles.

-e W

0.I

0 01

2 H2 (9) X CO (g) (Torr3)

FIG. 18. Methanol synthesis rate as a function of the product of concentrations [HZ(g)lzx [CO(g)] over a Cu/ZnOCr,O, catalyst at 250 C (114). [Reprinted with permission from Bull. C / w n . SOC.Jpn. 38, 1727 (1965). Copyright (1965) Chemical Society, Japan.]

278

K . KLIER

4. Carbon dioxide is adsorbed quite strongly. The reactant adsorption strengths are estimated to decrease in the order CO, > CO > H, . Based on these observations, Klier et al. (56) formulated the model for the synthesis as follows : (i) The catalyst can exist in a reduced state A,,d and in an oxidized state A,, . The oxidized state is active and the reduced state is inactive (cf. observation 1 above). The proportion of A,, and A r e d is controlled by the ratio of CO, and CO in the synthesis gas (cf. observation 2 ) . (ii) Several active centers A,, are involved in each reaction step. These centers may be identified with copper solute species in zinc oxide (cf. observation 3). (iii) All three components of the synthesis gas CO, H,, and CO, react in the adsorbed layer; CO, competes for active sites with at least one of the reactants CO and H,. The adsorption strengths are in the order CO, > CO > H, (cf. observation 4).

The redox reaction that controls the concentration of the active form A,, can be written as A,,

+ CO(& *Ared + COAd

(9)

which together with the requirement that the reduced and oxidized forms of the catalyst sum up to a constant A , = Ared A,,, yields for the concentration of active centers

+

Klier et a/ . investigated several cases of kinetics in which methanol is formed by a surface reaction between CO and hydrogen adsorbed on the A,, sites competitively or noncompetitively with CO on the A,, sites and hydrogen elsewhere on the surface; in each case CO, effects sub (i) and (iii) above were taken into account. In addition, it was found empirically that a small amount of CO, is hydrogenated to methanol at a rate that linearly depended on partial pressure of CO, . All kinetic equations that successfully described the CO, effects had the general form

279

METHANOL SYNTHESIS

where F is a linear function of pressures of H,, CO, H,O, and CH,OH; the exponent m is close to 3; the exponent n ranges from 1 to 3; and A,, is determined by the ratio pco,/pco as shown in Eq. (10). The term Key is an equilibrium constant, defined by partial pressures, for the reaction (1) and Klq for the reaction (2). These constants are related to fugacity ratios K , and KI and to the true equilibrium constants K , and K ; as Key = K,/K,; K & = K;,/K.:,, and their pressure-temperature dependences are given by

(5,W K , = 3.27 x lo-', exp(11,678/T),

K ; = 3.826 x lo-" exp(6851/T)

K ; = 1 - A,P,

KI, = (1 - A,P)(I - A 2 P )

A , = 1.95 x 10-4exp(1703/T),

A, = 4.24 x

(12)

exp(l107/T)

where absolute temperature T is in degrees Kelvin and the total pressure P is in atmospheres. The dependence of carbon conversion on the CO,/CO ratio and on temperature is shown in Fig. 19. Experimental rates are marked by circular (250"C), triangular (235"C), and square (225"'C)symbols, and the theoretical kinetic curves using the integrated form of the above kinetic equation (1 I ) with m = 3, n = 3, and F = (1 K H g H 2 Kc,pco) are represented as solid lines passing through the experimental points. With this form of F the extent of adsorption of H,O and CH,OH is neglected; the effect of water and methanol is, however, taken into account by kinetic terms for the reverse reaction, but not for product desorption in the forward reaction. The set of constants with which best fits were obtained is given in Table IX for three cases : case I , as above, describing competition of CO, H, ,and CO, for sites A,,; case I1 wherein CO are adsorbed on A,, and H , competes with CO, for different sites; and case 111 wherein CO, competes with both CO and hydrogen, but CO and H, are adsorbed on different sites. The statistical factors in the last column of Table IX show that all three models are nearly equally satisfactory. Thus the kinetic models appear to be insensitive to a specific mechanism by which CO and hydrogen utilize the catalyst sites in the synthesis. However, the promoting effect of CO, in creating the A,, sites and the strong retardation of the synthesis by the adsorption of this gas had to be taken into account in all three cases. The kinetics case I may be regarded as a convolution of the original Natta kinetics with the general form describing the CO, effects in Eq. (1 I). It must be emphasized that Natta's kinetic model (57) was tailored for catalysts used at higher temperatures, while the model of Klier r t al. (56) aimed primarily at a quantitative description of the low-temperature dependence

+

+

280

K . KLIER

80

A

0

z a

+ w

70

f

60

I

z 0 v)

50

K

w

>

5

40

0

z

0 30

m K

a

0

c

z w O K w a

20

(

I

10,

0 0

10

20

PERCENT CO2 IN SYNGAS,

30 C O + C O ~ = ~ O

FIG.19. The dependence of methanol synthesis rates at 225-250°C. 75 atm total pressure. and gas hourly space velocity 5000 hr-' upon the ratio of concentrations of C 0 2 and CO. The total hydrogen-to-carbon ratio in the feed gas was 7: 3 for all COJCO ratios. Experimental data are marked as open symbols and theoretical dependences, described by the kinetic model in text with the values of constants as in Tdble IX. as full curves. The closed symbok describes an equivalent of conversion rate when carbon dioxide in the mixture CO/C02/H, = 24/6/70 was replaced by argon ( 5 6 ) .[Reprinted with permission from J . Curd., 74,343(1982). Copyright (1982) Academic Press, New York.]

of the synthesis on carbon dioxide pressure. Klier rt u1. considered adsorption of CO, dominant over that of methanol and incorporated the methanolretarding terms only in the reverse reaction rates. This constitutes a major difference from Natta's treatment, and evidently the model can be refined. Klier r t a/. further scrutinized their model to determine whether the temperature dependences of the various adsorption equilibrium constants would yield reasonable values of adsorption enthalpies and entropies. These thermodynamic entities, calculated from data in Table IX, are summarized

TABLE IX Vulues of Constants Usrdjor- the Construction of Theoretical Dependences of' S t u d y - S t u t e Curhori Concersioris to Methanol upon ihe Percentage x of CO, in the Synthesis Gas of Composition C O J C O I H , = x/(30 - x)/70"

I

I1

111

225 235 250 225 235 250 225 235 250

12.52 8.58 5.00 12.52 8.58 5.00 12.52 8.58 5.00

1.77 1.40 1 .oo

1.77 1.40 1 .oo

1.77 1.40 1 .oo

39.62 21.52 9.00 98.51 60.0 30.0 19.8 10.8 4.5

158.2 125.4 90.0 18.1 16.1 13.8 79.1 62.7 45.0

1.064 1.253 1.584 0.898 1.135 1.584 0.088 0.120 0. I95

2.18( -4) 2.70( -4) 3.75( - 4) 2.18( -4) 2.70( - 4) 3.75( -4) 2.1 8( -4) 2.70( - 4) 3.75( -4)

9.034( - 3) 5.409( - 3) 2.625( - 3) 9.034( -3) 5.409(-3) 2.625( - 3) 9.034( - 3) 5.409( - 3) 2.625( - 3)

9.237( - 5) 6.573( - 5) 4.095( - 5) 9.237( -5) 6.573( -5) 4.095( - 5) 9.237( - 5) 6.573( - 5) 4.095( - 5 )

0.0012

0.0016

0.0015

A total pressure of 75 atm was maintained at all conversions (56). Rate constants per gram of catalysr per hour for rates expressed in moles of methanol per hour. Equilibrium constants for reactions CO(y) + 2H,(g) S CH,OH(g) ( K c q )and CO,(g) + 3H,(g) S CH,OH(g) H,O(g) ( K & ) ; K,, and K:, are defined on p. 279. Sum of the squares of differences between theoretical and observed conversions for C0,-containing synthesis gas mixtures divided by the number of measurements n.

+

282

K . KLIER

in Table X along with the activation energies for the synthesis of methanol from CO (E,) and from CO, (Ei).In all cases the adsorption heats were found to be in the expected order -AH(CO,) > -AH(CO) > -AH(H,). Furthermore, the adsorption entropies had very reasonable values : AS(C0,) corresponded to an entire (cases I and 111) or very substantial (case 11) loss of translational and rotational entropy of gaseous CO,, 55.4 cal mol- deg- ', while adsorbed CO and H, retained a significant portion of their gas-phase entropy (56). Thus the model that fits well the CO, dependence of the synthesis rate also provides a sound and attractive physicochemical picture of the synthesis: whereas CO and H, are bonded with intermediate strength and retain some mobility in the adsorbed phase, CO, appears relatively strongly bound and immobile. The behavior of CO and H, is favorable to the reaction in which carbon monoxide and hydrogen, adsorbed on different active centers, must meet in a reactive collision to form the product. At the same time, the strong immobile CO, adsorption points to the retarding effect of this gas at high CO,/CO ratios, which is more pronounced and produces flatter plateaus on the methanol yield versus CO, concentration as the temperature decreases from 250 to 225'C. The retarding effect of CO, by its adsorption can also be seen from the residence times calculated from sorption entropies : at 235°C the CO, residence time, approximately 1 sec, is comparable with methanol turnover time, 11 sec, whereas the CO and H, residence times were estimated to have values between lo-' and sec (56).The latter finding justifies a posteriori the assumption of the model that the CO and H, adsorption equilibria are rapidly established. The long residence time of CO, indicates that the adsorption of this gas may not be at equilibrium at low temperatures, a fact that is reflected by observed slow attainment of steady state when going from low to high temperatures in C0,-rich synthesis gas (55, 56). Evidently, the Cu/ZnO catalyst investigated in this work has a very good inherent activity for methanol but can be further improved for performance

'

TABLE X Vulues of Activution Energies of' Methunol Synthesis from Curhon Monoxide, E,(k), und,from Curbon Dioxide, E,(k' ), und Adsorption Enthulpies AH and Entropies AS Derived from the Kinetic Model Utilizing Constunts in Tuhle I X Case I AE&) AE,(k') AH(C0) AH(H,)

Case 11

Case 111

11.74 11.28

16.53 11.28

8.23 11.28 -19.0

-19.0

-19.0

-11.8

-11.8

- 11.8

Aff(C0,) AS(C0) AS(H,) AS(C0,)

Case 1

Case I1

Case I l l

-30.7 -33.1 -22.6 -54.3

-24.6 -33.1 -22.6 -40.3

-30.7 -33.1 -22.6 -55.7

~~

a 1 cal = 4.184 J ; from Ref. 56. Energies and enthalpies are in kcal/mol and entropies in cal mol- deg- ',

'

283

METHANOL SYNTHESIS

in C0,-rich synthesis gas by modifiers that weaken the adsorption energy of CO, . A selection of such modifiers will be determined by the nature of sites that strongly adsorb CO,. Such sites are more likely located on the zinc oxide surface than on copper metal surface and may be associated with the basicity of zinc oxide. Therefore it appears desirable to diminish the basicity of zinc oxide to achieve weaker bonding of CO, . A specific model of basic sites for CO:, that block sites that activate CO is shown below.

-0

P C U - 0 lZn-"\

l4

Cu site for carbon monoxide chemisorption

Adjacent oxygen site for carbon dioxide chemisorption

Once again the site involved in CO hydrogenation requires initimate interdispersion of copper and zinc oxide, which is consistent with the earlier conclusion that copper solute in zinc oxide is the catalytically active component. In the kinetic treatment presented above it was necessary to postulate the active form of the catalyst to be its oxidized form A,,, but it was not necessary to specify further what A,, and Ared are chemically. In order to present a consistent picture of the catalyst formation and deactivation, Klier et al. (56) divided the redox equilibria between Ared,CO, ,A,, , and CO into a reversible and an irreversible class, the latter giving rise to permanent deactivation of the catalyst. The reversible redox equilibria can be schematically represented as CO,

+ no+ 2e- sco + 02-,

and also CO,

CU+

ZnZ++ e- S Z ~ + (13)

+ e- =CU,O,

+ H-Cu+ e C 0 + HO-Cu'

where 0, is a symbol for oxygen vacancy and Cu; for dispersed copper. The reduction of Cu' is considered predominant over that of Z n 2 + a t temperatures below 250°C. The irreversible processes are segregation of copper, Cu: +Cu metal, and reduction to zinc and brass formation, Zn+

+ e-

--f

Zn,

Zn

+ nCu

+

Cu,Zn

(14)

both of which have been observed at high temperatures (27, 39, 62). Natta, however, noted that the reduction of zinc oxide is prevented by small amounts of water and carbon dioxide always present in the synthesis gas

284

K . KLIER

(lo),and so the most serious deactivation mechanism appears to be segregation of copper.

DIOXIDE ON SELECTIVITY 1. THEEFFECTOF CARBON Aside from the synthesis pattern shown in Fig. 19, carbon dioxide was found to cause amounts of methane to be formed at concentrations of CO, exceeding 10% in (CO COJH, = 30/70 synthesis gas (55, 56). Methane production increased with increasing concentration of CO, and with de-

+

0

10 20 PERCENT C02 IN SYNGAS

30

FIG.20. The steady-state concentrations of carbon dioxide (top), water (center), and methane (bottom) as functions of the CO&O ratio in the feed gas at 225-250°C, 75 atm total pressure, 250C, (-.-V-.-), 235°C. (-U-)225°C. Conversions to rnethand GHSV 5000 hr-' (56): (-O-) anol are given in Fig. 19. From Ref. 55.

METHANOL SYNTHESIS

285

creasing temperature, as shown in Fig. 20 together with the steady-state concentrations of water and carbon dioxide. The latter behavior indicates that methane is formed from adsorbed CO, . The pathway to methane must involve a direct hydrogenation of CO, and not a stepwise conversion of CO, to CO followed by hydrogenation of CO to methane, since in a CO-rich synthesis gas, methane is not a product.

J. RELATIVE HYDROGENATION RATESOF CARBON MONOXIDE, AND OXYGENATES HYDROCARBONS,

One of the useful methods of catalyst characterization is to compare rates of related reactions under similar pressure, temperature, and flow conditions. Furthermore, when carbon monoxide, another substrate, and hydrogen are allowed in contact with the catalyst simultaneously, their relative conversions reflect the extent to which these reactions components compete for the same active sites. Using the Cu/ZnO = 30/70 catalyst, the substrates carbon monoxide, 1-hexene, propionic acid, acrolein, propanaldehyde, benzene, toluene, and ethylbenzene were hydrogenated in vapor phase and in the presence of carbon dioxide under the pressure, temperature, and flow conditions at which CO is hydrogenated to methanol (63). The results are summarized in Table XI, which shows that the hydrogenation rates were substrate dependent in the order olefin > aldehyde > CO +CH,OH > carboxylic acid > ethylbenzene w toluene > benzene. A noteworthy observation is that methanol synthesis is a faster reaction than the aromatic ring hydrogenation over the Cu/ZnO catalyst, a pattern just opposite to that over methanation catalysts such as nickel (64). Vedage and Klier (63) further showed that carbon monoxide and other hydrogenable substrates compete for the same active sites. Their results, summarized in Table XII, demonstrate that all hydrogenable substrates suppress to varying degrees the rate of methanol synthesis from CO, and the presence of CO hinders to varying degrees the rate of hydrogenation of other substrates. Particularly interesting is the observation that the aromatic ring hydrogenation is completely stopped by the presence of CO, while methanol synthesis proceeds at a considerable rate in the presence of aromatics. The relative hydrogenation rates of ethylbenzene, toluene, and benzene indicate that the aromatics are activated as n-donor complexes, the substituted benzenes being better donors than benzene. In view of the proposals identifying the active centers as Cu(1) species in the ZnO surface, it is beneficial to point out that Cu(I)/CO (65) and Cu(I)/olefin (66) complexes exist and that the Cu(1) species, free from the 4s screening electrons and being

286

K . KLIER TABLE XI Relative Hydrogenation R a m oJ Carbon Monoxide, Hydrocarbons, and Oxygenates over the CujZnO Catalysta

Conversionb Substrate Benzene Toluene Ethylbenzene Propionic acid Propanaldehyde I-Hexene Carbon monoxide

(%I 16.6 25 29 63 99 100

55

Turnover rate' (sec- ')

4.95 x 6.27 x 6.24 x 1.48 x 4.23 x 4.33 x 2.96 x

10-3 10-3 10-3

lo-' lo-' 10-1

lo-'

Substrate feed rate (molihr)

Ionization potential (eV)

0.0068 0.0057 0.0049 0.0080 0.29 0.30 0.147

9.24 8.82 8.76

a Cu/ZnO = 30/70 catalyst described in Ref. 39. The hydrogenation products were saturated hydrocarbons in the case of aromatics and the olefin, primary alcohols in the case of oxygenates, and methanol in the case of CO. From Ref. 63. Conversion measured at 250°C, 75 atm, over 2.45 g catalyst using a feed of 10.5 liters (STP) of hydrogen per hour (0.43 mol H,/hr) and that of the substrate indicated in the fourth column of this table. ' The turnover rate was calculated from the measured conversion rates assuming 2 x lo'* sites/m2of ZnO on the Cu/ZnO catalyst having 23.3 m2 ZnO surface per gram of catalyst ( 4 3 ) . The turnover rate is defined as number of conversions per site per second per hydrogen molecule consumed. Hence the substrate turnover is two (three) times smaller in the case where the substrate is hydrogenated by two (three) hydrogen molecules.

'

TABLE XI1 Hydrogenation Raies of Carbon Monoxide, Hydrocarbons, and Oxygenates When CO PIus an Additional Substrate and Hydrogen Were Admitted to the CulZnO Catalyst' Amounts of substituents

MeOH

Other product

Reaction mixtureh

(%)

(%I

COICO,/H, 1-Hexene/CO/CO, /H2 PropanaIdehyde/CO/CO,/H, Acrolein/CO/CO,/H, Propionic acid/CO/CO,/H, Ethyl benzene/CO/CO,/H, Toluene/CO/CO,/H, Benzene/CO/CO, /H,

24/6/70 4.4/22.9/5.8/66.8 3.9/23/5.8/67.2 1.5/23.6/5.9/69.0 1.3/23.7/6.0/69.0 0.8/23.8/6.0/69.4 0.9/23.8/6.0/69.3 I .1/23.7/6.0/69.2

55 25-42" 32-51" 13.6" 0.6" 31.5" 34.0" 30.5"

100 n-Hexane

99 n-Propanol 99 n-Propanol 63 n-Propanol Traces of toluene 0 0

See Table XI footnotes. CO/CO,/H, = 24/6/70 mixture was used at 250"C, 75 atm, and total flow rate of 15 liters/hr over 2.45 g of the Cu/ZnO = 30/70 catalyst. The substrate was admitted from a liquid feed valve to produce concentrations indicated in the first column. a

* The

METHANOL SYNTHESIS

287

good “two-way acceptors,” are better suited for binding carbon monoxide and unsaturated organic compounds than any other valence states of copper (67). VI.

Other Binary Catalysts

There are several metal-oxide combinations other than the Cu/ZnO that have been reported active in methanol synthesis at low temperatures and pressures. These can be divided into catalysts containing copper and an oxide and catalysts containing a transition metal and an oxide.

A. COPPER-BASED CATALYSTS One of the recently studied catalysts is the copper-thoria composite made by oxidation and subsequent reduction of copper-thorium intermetallics such as Th,Cu, ThCu,, ThCu,,, , and ThCu, . Baglin et al. (68) found that air oxidation caused the alloy elements to segregate into a heterophase thoria, copper, and copper oxide mixture and that subsequent reduction in hydrogen or in the synthesis gas produced a catalyst containing crystalline copper and thoria. Catalysts with surface areas in the range 25-50 m2/g were made in this manner. Baglin et al. also stated, quoting Herman et al. (39), that the absence of diffraction lines of copper oxides would not preclude the presence of Cu(1) or Cu(I1) in the reduced catalysts and noted that further characterization of the copper-thoria catalyst may prove illuminating. The results of testing of this catalyst at 22O--30O0C,60 atm, H,/CO ratio equal to 16/1, and space velocity 31,000 hr-’ are summarized in Table XIII. The reported methanol conversions were high, and the yields at 260°C were comparable to those of the Cu/ZnO catalysts tested at 25OoC, 75 atm, H,/CO ratio 3/1, and space velocity 5000 hr-’: the Cu/ThO, = 5/l catalyst yielded 0.995 kg methanol per liter of catalyst per hour (68),while the Cu/ZnO = 30/70 catalyst yielded 1.10 kg methanol per liter of catalyst per hour (39) under the stated conditions. Table XI11 also shows products other than methanol, i.e., methane, ethylene, dimethylether, CO,, and water that were formed over the Cu/ThO, catalyst. The selectivity of the Cu/ThO, catalyst for methanol was good but not as high as that of the Cu/ZnO catalyst, the main side product being methane. A major difference appears to be that high conversions are obtained with the Cu/ThO, catalyst in a C0,-free synthesis gas, while the presence of CO, was necessary for maximizing the activity of the Cu/ZnO catalysts. The activity of the Cu/ThO, catalyst was found to decrease at

28 8

K . KLIER

TABLE XI11 Gaseous Product Analysis as a Function of Temperature from the Synthesis of Methanol over ThCu,,,, Preaciiuuted in Air" Mol % excluding H, Temp. ("C)

CO

220 240 260 270 280 300

83.52 67.22 55.93 52.26 49.42 54.44

CH,

CO,

0.24 0.08 0.60 0.28 1.24 1.07 1.99 1.58 2.85 2.46 5.49 3.65

C,H,

0.06 0.08 0.13 0.17

H,O

CH,OH

0.56 1.11 1.91

16.16 31.89 41.70 43.54 44.02 34.04

(CH,),O

Yield of CH,OH (kg/liter catalyst hr-')b 0.386 0.761 0.995 1.039 1.050

0.31

0.8 12

At 400°C; 16H,: CO; 60 atm pressure; 31,000 gas hourly space velocity; according to Baglin et al. (68). Calculated by the present author.

2-2.5%/day in a test involving 161 hr on stream. The catalyst performance was compared with that of a commercial Cu/ZnO/Al,O, catalyst (United Catalysts, Inc., C79-4) under identical reaction conditions and the yields were found to be 6.7 times higher over the Cu/ThO, catalyst. The activity of the Cu/ThO, catalyst could not be correlated with the total BET surface areas or with areas determined by chemisorption of carbon monoxide. The authors also stated that no standard method to measure the surface area of copper catalysts exists, an opinion based on a report by Scholten and Konvalinka (69)but which is at variance with those of Andrew (27),Parris and Klier (43), and literature cited in these two latter reports. It is evident that the Cu/ThO, catalysts are good prospects and due for a detailed characterization study. However, the claim that they are significantly more active than commercial Cu/ZnO/Al,O, catalysts is not substantiated because the best industrial catalysts based on this composition give yields higher than those reported for Cu/ThO, under comparable conditions (compare Tables XI11 and I). Aside from the recently described Cu/ThO, catalysts, copper on chromia and copper on silica have been reported to catalyze methanol synthesis at low temperatures and pressures in various communications that are neither patents nor refereed publications. I t is not feasible to critically review statements unsupported by published data or verifiable examples. However, physical and chemical interactions similar to those documented in the copper-zinc oxide catalysts are possible in several copper-metal oxide systems and the active form of copper may be stabilized by oxides of zinc, thorium, chromium, silicon, and many other elements. At the same time it is doubtful that more active and selective binary copper-based catalysts than

METHANOL SYNTHESIS

289

the optimized Cu/ZnO composites will be found, since the components of the latter seem to be ideally suited for the mutual synergic promotion required for the low-temperature methanol synthesis.

B. TRANSITION-METAL-BASED CATALYSTS Among the transition metals, Pd, Pt, Ir, and Rh in various forms have been reported active in methanol synthesis (32-34), as noted in Section 11. Palladium, platinum, and iridium metals were supported on silica, and it has recently been suggested that palladium is present in its valence state Pd(I1) which is the active form of the catalyst (70). Under the synthesis conditions the Pd(I1) ions could not survive the highly reducing atmosphere of the CO/H, synthesis gas, and so this valence state would have to be induced by the presence of silicon dioxide. Should this be a general case, silica would not act merely as an inert support, and the silica-supported transition metals would have to be considered binary catalysts whose active state is formed by a “support-metal interaction.” A dramatic effect of the support has been reported by Ichikawa (33,71,72) for rhodium catalysts prepared by deposition and thermal decomposition of multinuclear rhodium carbonyls on various oxides. When zinc oxide or magnesium oxide were used as supports, methanol was the predominant product. On the other hand, rhodium supported on titania or lanthana gave ethanol as major product, and rhodium on silica or y-alumina yielded hydrocarbons. These experiments were carried out at subatmospheric pressure and the yields of methanol were very low. In addition, methane was a significant by-product, as demonstrated in Fig. 21. In a subsequent communication Ichikawa reported higher yields and selectivities when the synthesis was carried out at higher pressures (72). Methanol and methane were also formed over a catalyst comprising rhodium and lanthanum hexaboride (73). In order to compare the activity of Rh/LaB, with those of Rh/ZnO and Rh/MgO, published by Ichikawa, the former was prepared in the author’s laboratory by a deposition of 0.1 g of Rh,(CO),, from tetrahydrofuran (THF) solution onto 20 g of LaB, support (14.5 m’/g), followed by evaporation of T H F and thermal decomposition of the rhodium carbonyl at 180°C in nitrogen. The resulting catalyst was tested at 75 atm in the temperature range 200-275°C with the synthesis gas CO/CO,/H, = 24/6/70; methanol and methane were the only carbon-containing products with yields given, along with those over the Pd/SiO,, Rh/MgO, and Rh/Fe/SiO, catalysts noted in Table XIV. A comparison of these yields with those over the Cu/ZnO catalysts demonstrates that the rhodium-based catalysts are inferior to the copper-zinc oxide catalysts on the basis of both the catalyst total volume

290

K . KLIER rnmol hr-’ 0

0.1

c

0

c

0

E

9

-

0.01

0)

O

n

0.001

FIG.21. Arrhenius plots of methanol and methane formation over the pyrolyzed catalyst on MgO prepared from Rh,(CO),, 0.1 1 g (0.15 mmol) and MgO 20 g, 0.75% Rh wt. dispersion. The reactor volume was 320 ml; CO : H, = 200 Torr: 450 Torr (33). [Reprinted with permission from Bull. Chem. SOC.Jpn. 51, 2268 (1978). Copyright 1978 Chemical Society, Japan.]

and the weight of the rhodium metal. Although the reaction conditions are not quite compatible, the results indicate that the activity of the palladiumsilica catalysts are intermediate between Cu/ZnO and Rh/MgO or Rh/LaB, . The silica-supported platinum and iridium catalysts had activity lower than the Pd/SiO, catalysts (32). In the effort to modify the catalytic properties of rhodium, Wilson and co-workers prepared Rh/Fe/SiO, and other two-metal-silica combinations (74); methanol yields over a Rh/Fe/SiO, catalyst shown in Table XIV were higher than those over the Rh/MgO, Rh/ZnO, and Rh/LaB, catalysts prepared from cluster carbonyls, but the selectivity of the latter toward methanol was better than that of bimetallic rhodium-iron catalysts. In the absence of a detailed characterization of the supported rhodium, it is premature but nevertheless interesting to speculate about which property of the oxide “support” induces methanol synthesis activity in small rhodium particles. The effects that might be involved here are acid-base or electron transfer interactions between the support and the rhodium particles, including the change of the effective valence state of rhodium. Basic oxides such as ZnO or MgO will tend to donate electron pairs to the rhodium

29 1

METHANOL SYNTHESIS

TABLE XIV Supported Transition-Metal Catalysts,for Methunol: A Comparison of Activities

Catalyst composition

Reactants

Temp. ( C)

Pressure (atm)

Yield (kg liter-' h r f * )

Reference

Pd/SiO, Rh/MgO

H, H,

+ CO + CO

350 250

100 0.85

H,

+ CO

250

75

H,

+ CO

300

68

0.23" 0.0002" 0.07" 0.00015' 0.05' 0.004 0.15" 0.0044 0.1 9d 0.02" 0.06'

32 33 33 33 33 73 73 73 73 34, 74 34, 74

Rh/LaB,

Rh/Fe/SiO,

Methanol yield in kgiliter catalyst hr-'. Yield of methanol in kg/kg Rh hr-'. Methane yield in kgiliter catalyst hr- ', Methane yield in kg/kg R h hr-'. CH,OH i C,H,OH + C, chemicals and oils 'Yield of CH, "

"

+

particles to fill some of its d levels and possibly produce effective electron configurations of the metal closer to one on the right-hand side in the periodic table, i.e., to palladium. Such an electron transfer or polarization will have the effect of suppressing carbon monoxide dissociation on the metal and will amount to crossing the division line between hydrocarbon and methanol synthesis catalysts in Table 111 from the left to the right (75). In order for the basic support-metal interactions outlined above to be effective in modifying the chemical properties of the metal, it is imperative that the metal particles be very small. Although the catalyst preparation by decomposition of carbonyls is designed to produce very small metal clusters, there is a suspicion that these clusters may agglomerate during use. Therefore, any rational explanation of support effects in the rhodium catalysis must await a detailed characterization of the metal dispersion.

VII. Ternary and Quaternary Catalysts Ternary compositions Cu/ZnO/Al,O, and Cu/ZnO/Cr,O, are currently the most important industrial catalysts. The solid-state chemistry of these composites is often complex, but there is no evidence that addition of alumina or chromia to the binary Cu/ZnO systems causes significant synergic

292

K . KLIER

promotion. Other third components, however, are known to alter the selectivity of the Cu/ZnO catalysts and in some cases act as retardants through a great number of possible solid-state reactions which perturb the delicate balance of copper in zinc oxide that is necessary for the catalyst formation. In this section we shall concentrate on the “structural promotion” by alumina and by various products of its interaction with zinc oxide and copper. There are many preparations and different ratios of copper, zinc oxide, and alumina, or chromia that result in successful catalysts, as is apparent from Tables I and 11 and the literature quoted therein. Two types of these “formulations” were characterized better than others, and these are discussed in some detail. ICI catalysts contain deliberate additions of crystalline zinc aluminate spinel ZnAl,O,. Andrew (27) stressed the importance of this refractory oxide in stabilizing the catalyst for long-term performance. The ICI catalyst was reported to be a mixture of microcrystalline copper of average particle size between 5 and 7 nm, microcrystalline zinc aluminate, and perhaps also zinc oxide. One of the functions of zinc aluminate was suggested to be to maintain the copper particles in fine dispersion, and data from ICI plants were analyzed under the assumption that the surface of the copper crystallites was the catalytically active part of the composite (27). The ICI Cu/ZnAl,O,/ ZnO catalysts exhibited dependence on the CO,/CO concentration ratio similar to the binary Cu/ZnO catalysts described by Klier et al. (56). The catalysts were severely deactivated in C0,-free synthesis gas, an effect attributed by Andrew to the reduction of zinc oxide with the concomitant alloying of zinc with copper. This would be the more severe mode of deactivation, grouped among the irreversible processes [Eq. (1 3)]. In principle, the CO,/CO dependence of the synthesis rates over the ICI catalysts can be described by the same formal kinetics as presented for the Cu/ZnO binaries on p. 279. Both sets of data require the reduced form of the catalyst ArCd to be inactive, despite the disagreement on what is the chemical nature of Ared at low temperatures: Andrew’s proposal is that Ared is brass and A , , is metallic copper, while the Klier cz a/. basic concept is that ArCdis metallic copper and A,, is monovalent copper in zinc oxide. It appears likely that the same set of solid-state reactions underlies the observed C0,jCO dependences over the ternary Cu/ZnAI,O,/ZnO and the binary Cu/ZnO catalysts, no matter which interpretation of the chemical nature of the redox process involved in the formation of A,, will eventually prevail. Another type of preparation of the Cu/ZnO/Al,O, catalyst involves coprecipitation of the three components and the resulting catalyst contains neither spinel nor crystalline alumina. Fischer r’f 01. (76) reported that a ternary catalyst of the nominal composition Cu/ZnO/Al,O, = 60/35/5 mol% contained only crystalline copper and zinc oxide, and all the alumina

METHANOL SYNTHESIS

293

was in X-ray amorphous form. The chemical composition and the amounts of X-ray crystalline matter are given in Table XV. The error given for the crystallinity, f lo%, was too large to determine whether any part of zinc oxide or copper was X-ray amorphous, but the chemical Cu/ZnO ratio (1.7) was substantially greater than the average X-ray crystallinity Cu/ZnO ratio (1.38). It is interesting to note that were all ZnO crystalline as in the binary Cu/ZnO catalysts (see Fig. 5), the comparison of the chemical and crystalline Cu/ZnO ratio would indicate 19% of the copper to be X-ray amorphous. Consequently, the mole ratio of X-ray amorphous copper to zinc oxide would be 0.4, exceeding that found by Bulko et LEI. (41) in the binary Cu/ZnO catalysts. Further evidence that part of copper is X-ray amorphous in the ternary catalysts follows from the work of Herman et al. (77). These authors prepared a Cu/ZnO/Al,O, = 60/30/10 catalyst by coprecipitation of a precursor from the solution of Cu2+, Zn2+, and A13+ acetates by sodium carbonate, washing, calcination at 350"C, and reduction at conditions identical with those employed for the preparation of Cu/ZnO composites. The resulting catalyst consisted of particles shown in Fig. 22a. The large hexagonal particles had a structure of a single-crystal wurtzite zinc oxide oriented with its sixfold axis perpendicular to the hexagonal plane outline of the crystal, as documented by the diffraction pattern in Fig. 22b. The analysis in the STEM showed that this type of particle contained 8 at.% Cu, 50 at.% Zn, and 42 at.% Al. Separate large copper crystallites, showing as dark areas in the electron micrograph of Fig. 22a, were also constituents of the catalyst, but their surface area was estimated to be negligible compared to that of the porous ZnO platelets. Given the earlier result that largeparticle crystalline copper does not catalyze methanol synthesis, the activity of this catalyst resides in the "micromonolith" porous particle of ZnO that contains all three elements Cu, Zn, and Al. Interestingly enough the pore walls, which are perpendicular to the platelet major dimension, expose the same prism plane of zinc oxide as the most active binary catalyst Cu/ZnO = 30/70, and very likely the chemical and physical nature of the active surface TABLE XV Chemical Composition and X-ray Crysrailiniry of CulZnlAl Oxide Caralysi" Catalyst compositionb (Cu/Zn/Al) (at. %) determined (1)

chemically, 60/35/5

(2) X-ray diffraction (floyo),58;42,0

According to Fischer rt ai. (76). Technical low-pressure methanol catalyst prepared by coprecipitation and reduced at 240 C. 'I

"

FIG.22. (a) Transmission electron micrograph and (b) selected area diffraction pattern of a single hexagonal porous particle of the Cu/ZnO/Al,O, catalyst prepared from acetates (77). The Cu/Zn/Al ratio exclusive of crystalline copper, which appears as dark crystallites,was equal to 8/50/42.

METHANOL SYNTHESIS

295

of the two catalysts is identical. Herman et a/. (77) suggested that the role of amorphous alumina is to “glue” the zinc oxide microcrystals doped with amorphous copper solute around the 5-nm pores. The specific activity for methanol of this Cu/ZnO/Al,O, catalyst was nearly equal to that of the binary Cu/ZnO = 30/70 catalyst, so that the effect of alumina appears to be structural promotion only. A different view on the role of amorphous alumina was taken by Fischer et al. (76). These authors noted that copper in the ternary Cu/ZnO/Al,O, catalysts displayed paracrystallinity, i.e., a defect crystalline state in which long-range order may be perturbed by “endotactic” inclusions-randomly distributed species whose dimensions have a strong misfit with the host lattice of the copper crystallites. Fischer et uf. argued that these endotactic inclusions are small A1,0, clusters because no crystalline alumina could be found in the catalyst. The effect of inclusions was suggested to be indirect : the lattice disorder would cause some of the surface copper atoms to be displaced from their lattice positions, which then bind carbon monoxide with a different strength than the lattice atoms. Quantum mechanical calculations have predicted that the atoms displaced only 10% from their lattice positions would bind CO more weakly than atoms in ideal lattice positions, and Fischer et al. further argued that this weaker form of CO, because of its greater mobility, would be one that is most rapidly hydrogenated to methanol. This concept is in contrast to the observations of Parris and Klier (43)who found an irreversibly bonded type of carbon monoxide on the more active of the binary catalysts (see Fig. 13) and concluded that the hydrogenable form is the more strongly adsorbed CO on the solute copper centers in zinc oxide. In neither case is the catalyst activity expected to be proportional to the copper metal surface area: the density of defects induced by inclusions will depend more on the nature and concentration of the inclusions, and the density of copper solute atoms will depend more on chemical interactions between the copper and the host lattice than simply on the copper surface area. A comparison of activities of the Cu/ZnO catalysts and the Cu/ZnO/Al,O, catalyst prepared by Herman ef al. (77) indicates that the presence of alumina did not bring about a dramatic change of specific activity. It is therefore concluded that paracrystallinity due to alumina inclusions was not a significant determinant of activity of the latter catalyst. In summary, several roles were attributed to alumina in the Cu/ZnO/Al,O, catalysts : prevention of copper particle sintering through the formation of zinc aluminate, induction of surface defects by endotactic inclusion of alumina clusters in copper, and that of stabilization of highly dispersed Cu/ZnO binary catalyst. It is possible that all these effects take place in varying degrees in the commercial catalysts, but there is no evidence at this time that the specific activity of the present commercial catalysts Cu/ZnO/

296

K . KLIER

A1,0, catalysts originates from interactions other than those existing in the binary system Cu/ZnO. The structural promotion by alumina is a very significant factor in the formulation of industrial catalysts, however, as it imparts chemical and mechanical stability required for a long-lived catalyst in large-scale reactors. It is noteworthy that components other than alumina often have detrimental chemical and physical effects on the catalyst. For example, Herman et ul. (77) reported that addition of ceria to the Cu/ZnO catalyst lowered methanol conversion by a factor of 5, despite the presence of a large concentration of microparticulate copper metal. This effect was explained by the ability of ceria to drive copper from the active state in zinc oxide solution to inactive metallic copper. Chromia, which had been used as a component of catalysts for methanol for a considerable period of time, is a suitable structural promoter, but some preparations result in an increase of concentration of side products such as higher alcohols (39),dimethyl ether (47), or even hydrocarbons. There are other components such as alkali, transition metals cobalt and rhodium, oxides of iron, manganese, vanadium, and rare earths, which have at one time or another been used in combination with copper and zinc oxide. The effects of alkali and cobalt metal are to lower the selectivity to methanol with a significant coproduction of hydrocarbons and higher alcohols. Recently a quaternary catalyst has been patented (78-80), based on a combination of copper, cobalt, transition-metal oxide, and alkali oxide that produced from the synthesis gas a mixture of C,-C, alcohols with the major product being ethanol. This catalyst contains components of methanol catalysts and components of Fischer-Tropsch or homologation catalysts. The process is being evaluated as of the time of this writing. The reported product composition and yields over this and other recently reported catalysts that have produced significant amounts of low alcohols have been summarized by Klier (75). The claimed activities of the copper-cobalttransition oxide-alkali catalysts are higher than those of the rhodium-based catalysts. The understanding of the function of these catalysts must await detailed characterization of the structural, electronic, and morphologic properties of their components, similar to the one achieved with the simpler Cu/ZnO methanol catalysts.

VIII.

Mechanisms

Kinetic experiments performed by Natta et a/. (57) with a ZnO/Cr,O, catalyst and by Klier et ul. (56) with the Cu/ZnO = 30/70 catalyst, as well as the high-power dependence of the synthesis rate on the concentration of

297

METHANOL SYNTHESIS

active sites measured by irreversible CO chemisorption (Fig. 17) and on reactant pressures (Fig. 1S), indicate that each molecular synthesis step requires several sites. Natta stressed that kinetic equations, based on the hypothesis that the surface reaction is bimolecular, did not fit the observed rate dependences on the partial pressures of the reaction components. Only with the hypothesis of a trimolecular surface reaction involving one CO and two H, adsorbed molecules was it possible to interpret the rates observed over the ZnO/Cr,O, catalyst quantitatively. Later developments have shown that several kinetic models will fit Natta's original data, but a common feature to all was that the synthesis rate depended on positive powers of partial pressure of hydrogen and carbon monoxide, the dependence on hydrogen being stronger than that on carbon monoxide. Klier et al. (56) also treated the synthesis over the low-temperature Cu/ZnO catalyst as a trimolecular surface reaction and obtained a successful kinetic model. Mechanistic interpretation of the kinetic models may be as follows. Equilibria of carbon monoxide and molecular hydrogen between the gaseous phase and adsorbate are rapidly established. While the CO and H, molecules undergo dynamic exchange between the gas and adsorbed phase and among the surface sites, a small fraction of each, proportional to their surface concentrations, is activated for a reactive collision. The activation may involve splitting of hydrogen molecules into atoms, and partially hydrogenated species may include formyl HCO, formaldehyde HCHO, hydroxycarbene HCOH, hydroxymethyl CH,OH, and methoxyl CH,O. With the participation of oxygen atoms of the solid oxide, of carbon dioxide, or water, formate HCOO- is also a possible intermediate. Once the reactive collision has occurred, the subsequent partial reactions are irreversible and rapid. For this reason, the kinetically significant intermediates are difficult to detect during the synthesis by either physical or chemical methods, and consequently the synthesis mechanism has not been firmly established. Yet it is very important to understand some fundamental features of the synthesis pathway because knowledge of mechanistic details will allow the control of selectivity to side products, where desirable, and pave the way for the design of catalysts for alkylations with the synthesis gas and other related reactions.

A. POSSIBLE REACTION PATHWAYS The mechanisms proposed in the past may be divided into three basic groups : carbon-down, represented by the sequence ' 0 '

II

cII M

H

H,C$I&

I

M

H,

WH

C--GxH II

M

2H-C-HII

M

H

M

+ CH,OH

(I)

29 8

K . KLlER

oxygen-down, represented by the sequence H

H

CH,

H

H: M

+ CH,OH

and side-on hydrogenation

Here M is a symbol for a site involved in the activation of carbon monoxide. It is helpful to summarize briefly the thermodynamic energies of the various intermediates in pathways 1-111. Figure 23 shows enthalpies and Gibbs free energies at 250°C which would be required for noncatalyzed reactions that involve a particular set of intermediates. Pathway I1 is drawn in Fig. 23 for potassium hydroxide, formate, methoxide, and hydride because of the lack of available thermodynamic data for the corresponding compounds of zinc and copper. Nevertheless, the relative values of AG and AH illustrate the points to be made. The role of the catalyst for Scheme I would be to lower the 200 kJ/mol thermodynamic barrier by suitable bonding of hydroxycarbene HCOH, while pathway I1 would require destabilization of the surface formate and methoxide. It is also seen that the removal of surface methoxide by hydrolysis would proceed with a lower thermodynamic barrier than by hydrogenation. Pathway 111 would seem to be a feasible one because the free energy of formaldehyde lies only 17 kJ/mol above that of methanol. However, formaldehyde has never been observed as intermediate in methanol synthesis over catalysts reviewed here. On the other hand, formaldehyde was frequentiy found to be the product of methanol decomposition (9, 81-83). It is therefore likely that large kinetic barriers exist for the synthesis of formaldehyde that offset the thermodynamic feasibility of pathway 111. It is evident that entirely different requirements would be imposed on catalysts for pathways 1,II, or 111 : route I would require stabilization of hydroxycarbene, route I1 destabilization of formate and methoxide, and route 111 lowering of the kinetic barrier for formaldehyde synthesis. In the effort to detect some of the intermediates listed in Fig. 23, various workers in the field resorted to experimental

‘“t

299

METHANOL SYNTHESIS

0

co+ti2

‘\,CHjOH

I

FIG. 23. Standard enthalpy (dashed lines) and Gibbs free-energy (full lines) changes for intermediates in methanol synthesis. (a) Pathway I: hydroxycarbene route CO(g)

+ HAg) +HCOH(g),

HCOWg)

+ H2k) +CH,OH(g)

(b) Pathway 11: formate-methoxide route

CO(y) + KOH(s) + HCOOK(s) HCOOK(s) 2H,(g) + CH,OK(s) + H,O(g) CH,OK(s) HZ(g)+ CH,OH(g) KH(s) H20(g)+ CH,OH(g) KOH(s) CH,OK(s)

+ + +

+ +

(c) Pathway 111: formaldehyde route COk)

+ Hzk)

-+

HzCO(g),

H,CO(Q)

+ Hz(g) +CH,OH(g).

methods, outlined below, which gave partial insight into the methanol synthesis mechanisms.

STUDIES B. METHODSUSEDIN MECHANISTIC OF METHANOL SYNTHESIS Chemical composition of side products of methanol synthesis and decomposition has often been reported but rarely used for the resolution of

300

K . KLlER

mechanism. The most frequent side products in the synthesis, when run over less than optimally selective catalysts, are dimethylether, higher alcohols, methane, and methyl formate. While dimethylether and higher a.lcohols are clearly products of secondary reactions, methane and methyl formate may be formed directly from the surface intermediates containing one carbon atom. The presence of methane is diagnostic of the C-0 bond cleavage, as discussed in Section V,I. Methyl formate may be formed by a number of reactions ; among others, thermodynamic driving forces exist in favor of esterification HCOOH(g)+ CH3OH(g)*HCOOCH3(g)

where A H 2 5 0 0 C = -9.54 kJ/mol, AG,,,:, proportionation of formaldehyde

+ H,O(g)

(15 )

= - 13.39 kJ/mol, and of dis-

ZHCHO(g)* HCOOCH,(g)

(16)

where AHzso., = - 116.44 kJ/mol, AG,,,., = -46.77 kJ/mol. If related reactions take place on the catalyst surface, then perhaps methyl formate could be formed from surface formate, methoxide, or formyl residues. It is interesting to note that in an early study of methanol decomposition over the Cu/ZnO catalysts, Frolich et al. (6)found methyl formate and formaldehyde to be the predominant products over the copper-rich composites, including copper metal, whereas carbon monoxide was the dominant product over the zinc-oxide-rich catalysts. While the decomposition rates to carbon monoxide paralleled those of methanol synthesis rates (6), there was no correlation between the decomposition of methanol to methyl formate and the synthesis. Since methyl formate was also formed over pure copper, the oxygen of an oxide does not appear to be necessary for the occurrence of methyl formate. Consequently the presence of methyl formate, either in the synthesis or in the decomposition of methanol is not unambiguous evidence for pathway I1 with surface formate being an intermediate. Coadsorption of reactants and subsequent thermal decomposition of the surface complexes formed have been used to resolve the mechanisms in several studies (81-84). Mutual enhancement of the adsorbed amounts of the reactants is indicative of their interaction, and if the adsorption of separately admitted components is negligible, the stoichiometry of the adsorbed complex can be determined. Further evidence for the formation of an adsorbed complex, employed in a mechanistic study of methanol synthesis over ZnO (84), is obtained by thermal decomposition of the adsorbed complex : if the reactants appear simultaneously at one temperature upon thermal desorption from a coadsorbed layer, but if each reactant adsorbed separately gives a thermal desorption peak at a different temperature, the existence, although not necessarily the structure or com-

301

METHANOL SYNTHESIS

position, of the adsorbed complex has been clearly established. Zxotope methods can also be used for the determination of certain features of the methanol synthesis mechanism. If the carbon monoxide is initially mixed from l3C''O and 12C'80 and if the methanol product contains no '3CH3180Hand '2CH,'60H, it can be concluded that (1) C-0 bond does not rupture and (2) the synthesis does not proceed via formate with the participation of oxygen of an oxide or other sources such as CO, or water. Clearly, a rapid shift reaction or exchange reactions of the catalyst oxygen with CO would scramble 13C and "0 between CO, CO,, and water, and would preclude the above result. In one experiment hydrogen in the synthesis gas was rapidly replaced by deuterium, and CH,OD was found to be predominant product (85). Such an isotope composition of the product would prove that the slow hydrogenation is that of a CH,O intermediate, if it were proven simultaneously that CH,OH methanol does not rapidly exchange its hydroxylic hydrogen with deuterium. In the opposite case the predominance of CH,OD would be the result of an exchange reaction and not of the synthesis. Perhaps the most powerful mechanistic tool is the trapping qf' intermediates by reactive chemicals. For example, dimethyl- and diethylsulfate were used for trapping of surface methoxides and formates (86). Unsaturated

ZnO,022449

0l;m N - 0

NNN

0'

3200

2800

2400

2000

1800

1600

1400

1200

1000

crn-'

FIG.24. Infrared spectra of methanol chemisorbed on zinc oxide at 200°C: (.--.) background; (-) CH,OH; (...) CD,OD. The bands at 2930 and 2830 cm-' are assigned to the CH3 asymmetric and symmetric stretching vibrations of surface methoxide ions. Bands at 2870, 1571, and 1367 cm-' are assigned to the C-H stretching, 0-C-0 antisymmetric symmetric stretching vibrations of surface formate ions. Further stretching, and 0-C-0 details of assignments are given in Ref. 88. [Reprinted with permission from Trans. Furuduy SOC.67, 3585 (1971). Copyright (1971) Chemical Society, London.]

302

K . KLlER

hydrocarbons and heteroatomic compounds may be considered for trapping of hydroxycarbenes as illustrated by the many examples given by E. 0. Fischer (87).In methanol synthesis, however, most of the unsaturated compounds will undergo a competitive hydrogenation (63),and their reactions with the fleeting intermediates may not be readily observable. Finally, if conditions can be found under which the kinetically significant intermediates appear in sufficiently large surface concentrations, various surfiice spectroscopic methods may be used for their detection. As an example, IR spectra of surface methoxide and formate are shown in Fig. 24 (88). Spectra of surface hydroxyls and hydrides (89-92), as well as of alkyl groups (93), carbonates, and carbonyls (94) have been well described. The vibrational spectra on small flat surfaces can also be determined by electron energy-loss spectroscopy (EELS). Semiquantitative surface analysis can be obtained from Auger electron (AES) or X-ray photoelectron (XPS) spectroscopy. Information on which orbitals of the adsorbed reactants, intermediates, and products are involved in bonding with the catalyst surface may be gained from an analysis of UV photoelectron spectroscopic (UPS) data. All of the methods cited in this section have been used with varying degrees of success for the detection of intermediates in methanol synthesis and decomposition. Additional mechanistic conclusions have been made on the basis of the change of product composition or of reaction rate upon a deliberate modification of the catalyst by bulk or surface dopants, and from analogies with the chemical behavior of metal-organic complexes. C. EVIDENCE FOR ADSORBED CARBON MONOXIDE AND HYDROGEN Hydrogen is weakly adsorbed on copper metal, on zinc oxide, and on the copper-zinc oxide catalysts (43). However, a variable portion of the adsorbate is irreversibly bound. On zinc oxide, IR bands at 3400-3500 and 1475-1708 cm-I were attributed to the stretching vibrations of OH and ZnH groups, respectively, and various workers agree that the OH and ZnH groups originate from heteropolar dissociation of hydrogen molecules on adjacent zinc and oxygen surface sites (89-92). Hydrogen chemisorption on copper is nondissociative at temperatures below 300°C because the H,/D, conversion does not occur at these conditions (95). Carbon monoxide adsorption on ZnO has been found to depend strongly on the pretreatment of the oxide. Kortiim and Knehr (94) detected a weakly bound CO with at least one rotational degree of freedom showing rotationalvibrational bands with maxima at 2188 and 2120 cm-' and a strongly bound complex similar to bidentate carbonate. In addition, well-degassed ZnO specimens adsorbed CO with the IR frequency within the carbonyl

METHANOL SYNTHESIS

303

range, 1800-2100 cm- I . The sites for this carbonyl-type adsorbate were suggested to be interstitial Zn' ions, and carbon monoxide was assumed to be bonded carbon-down in this adsorption mode. The carbonyl-like CO was highly reactive toward oxidation. No evidence for surface formate can be found in the work of Kortum and Knehr on hydrogen-pretreated ZnO samples exposed to CO. Ostrovskii et al. (96) studied carbon monoxide adsorption on a ternary Cu/Zn/Al oxide catalyst by volumetric methods and came to the conclusion that surface carbonate was formed on oxygenpretreated catalyst but that both the rate and the extent of CO adsorption were low on oxygen-depleted catalyst. The chemisorption studies of Parris and Klier (43) using the Cu/ZnO catalyst have been mentioned earlier. Carbon monoxide was irreversibly bonded at room temperature to the surface of the binary catalysts that were also active in methanol synthesis; however, this irreversible adsorbate could be desorbed as CO, which indicates that it was not a surface carbonate but rather a strongly bonded carbonyl-type CO. Infrared studies of this chemisorbate are lacking and it would be very desirable to determine the structure of this surface species. Carbon monoxide chemisorption on copper metal is weak and for the most part reversible, and has been studied by a number of techniques (97-103).

D. EVIDENCE FOR ADSORBED FORMATE AND METHOXIDE Surface formate and methoxide bands have been detected by 1R spectroscopy in methanol decomposition over zinc oxide (88) and in the shift reaction over the Cu/ZnO catalyst (104). The surface methoxide, which is characterized by the CH symmetric stretch at 2840 cm- specific for the CH,O group, was also detected in IR spectra of methanol chemisorbed on iron, nickel, and cobalt (105), magnesia (106), and alumina (107-112). Bowker and Madix (82) used UPS, XPS, and thermal desorption techniques to identify products of methanol adsorption on the Cu(ll0) crystal face and determined that methanol orbitals which have the highest electron density on the oxygen atom were most affected by bonding to the surface. Their findings are consistent with half monolayer coverage of methoxy species originating from methanol dissociation on the clean Cu( 1 10) surface at 270 K. It is evident that surface methoxide is a fairly common product of the interaction of methanol with solids which are methanol catalysts as well as with many other solids that are not catalysts for methanol. To provide complete information, however, the fate of hydrogen from the methanol hydroxyl should also be determined. This hydrogen is expected to end up

304

K . KLIER

in an M-H bond on metals and in an O-H bond on oxides. The M-H and O-H stretching vibrations, so far unreported in methanol decomposition studies, should be readily detectable in IR spectra. Methoxide and formate were also found by chemical trapping with SO,(CH,), and SO,(C,H,), of surface complexes formed in methanol synthesis over the BASF ZnO/Cr,O, catalyst (86) according to the reactions R,SO, R,SO,

+2 C H p + 2HC00-

-+

2CH,OR

+ SO:-

(17)

-+

2HCOOR

+ SO:-

(18)

where R is methyl or ethyl. Both the methoxide and the formate were found when the synthesis was run in a flow system at 250°C 50 atm, and H,/CO = 2, with the CH,O-/HCOO- ratio being equal to 7.9; 0.5% of methyl formate was also detected in the methanol product.

E. SURFACE COMPLEXES FORMED BY COADSORPTION OF THE REACTANTS

Many studies of simultaneous adsorption of hydrogen or water and CO or CO, have been carried out on the high-pressure methanol synthesis catalysts based on zinc oxide and one or several other oxides, but only three investigations (104, 113, 114) dealt with catalysts containing copper, and two of these were made in reference to the mechanism of the low-temperature shift reaction. The interaction of adsorbed carbon monoxide and hydrogen was examined by IR spectroscopy and other methods. A mutual enhancement of CO and H, adsorption on zinc oxide was reported in an early study by Nagarjunan et al. (115) to occur at 0 and 25°C and interpreted as the formation of the complex =CHOH, i.e., hydroxycarbene, in the adsorbed phase. However, a more recent 1R investigation of Boccuzzi et af. (116) provided evidence for interaction of adsorbed carbon monoxide and hydrogen of a different kind than attributable to the formation of hydroxycarbene. At room temperature, carbon monoxide admitted to zinc oxide with preadsorbed hydrogen was found to perturb both the Zn-H (1700 cm-') and O-H (3520 cm-') bands originating from hydrogen chemisorption. While the OH frequencies were shifted upward, new discrete lower frequency bands of ZnH at 16801690 and 1660-1670 cm-' were formed, the relative intensities of which depended on CO pressure, as illustrated in Fig. 25. These new bands cannot be assigned to any expected bands of hydroxycarbene. The carbon monoxide adsorption was accompanied by a carbonyl-type band at 2180 cm- and was attributed to carbon-down species on the cations since it was displaceable by pyridine. A model explaining the more pronounced effect of CO on

METHANOL SYNTHESIS

305

FIG. 25. Effect of increasing CO pressure on the Zn-H stretching bands on zinc oxide (PH2= 100 Torr) (116). [Reprinted with permission from J . Cutul. 51, 160 (1978). Copyright (1978) Academic Press, New York.]

the ZnH and less pronounced effect on the OH band, proposed by Boccuzzi et al. (116), is illustrated in Fig. 26. No IR bands of surface formates, methoxides, or carbonates appeared in the spectra of ZnO with hydrogen and CO coadsorbed at room temperature. More surprisingly, no surface complexes were detected upon coadsorption of CO and H, at temperatures

(C)

FIG.26. Model for coadsorption of carbon monoxide and hydrogen explaining the infrared shifts in Fig. 25 (116): (a) clean surface; (b) hydrogen-covered surface; (c) hydrogen- and carbon monoxide-covered surface. [Reprinted with permission from J . Cutul. 51, 160 (1978). Copyright (1978) Academic Press, New York.]

306

K . KLIER

up to 300 C and pressures of 1 Torr followed by programmed thermal desorption in the study of Bowker et ul. (84). These authors concluded that no stable surface formate is formed on ZnO from carbon monoxide and hydrogen, and their conclusion must be held valid within the temperature and pressure range outlined above. However, Deluzarche rt a/. (86) employed pressures up to 50 atm at 250'C and were able to trap chemically both surface formate and methoxide when a ZnO/Cr,O, catalyst was exposed to hydrogen and carbon monoxide either sequentially or simultaneously in the course of methanol synthesis. With an H,/CO = 2 ratio in a continuous run, Deluzarche et al. found the ratio of surface methoxide to formate CH,O-/HCOO- = 7.9, corresponding to an average formula CH,,,,O,,, The amounts of coadsorbed carbon monoxide and hydrogen were also measured during methanol synthesis over the Cu/ZnO/Cr,O, catalysts (114) at atmospheric pressure and 250°C with a widely varying ratio of gaseous components. The adsorption of each reactant was enhanced in the presence of the other. The principal result of this study is represented in Fig. 27. The authors noted that the average composition of the adsorbate, H, :CO = 1.5: 1, corresponded to the stoichiometric formula CH,O and concluded that the synthesis rate is limited by the hydrogenation of this complex. A closer inspection of Fig. 27 indicates that the H,:CO ratio on the surface varied from 1 : 1 to 2: 1 as the gaseous concentration of hydrogen increased, corresponding to stoichiometric formulas CH,O to CH,OH. Thus several intermediates may be responsible for the mutual enhancement of the carbon monoxide and hydrogen adsorption during the synthesis

0.0I

0.I

I

10

I00

1000

H,(g)/CO(g)

FIG.27. Relationship between H,(g)/CO(g) and H,(a)/CO(a), the ratios of hydrogen and carbon monoxide concentrations in the gas (9) and adsorbed ( a ) phase during methanol synthesis (0and 0 ) and decomposition ( x ) over Cu/ZnO/Cr,O, catalyst at 250°C (114). [Reprinted with permission from Bull. Chem. SOC.Jpn. 38, 1727 (1965). Copyright (1965) Chemical Society, Japan.]

307

METHANOL SYNTHESIS

Kinetic evidence for synergic adsorption of ccrrhon rnono.\-i& cind n ~ w r on the low-temperature shift catalyst Cu/ZnO/Fe,O, was obtained by van Henvijnen and deJong (113), and IR spectra of surface formate were detected on several oxide catalysts, including CuO/MgO, at temperatures as low as 20°C and pressures of 20 Torr, as reported by Davydov r t a/. (104). Decomposition of the surface formate to CO, and H, occurred at 100- 150°C over the Cu/MgO catalyst and at 250-3OO'C over the MgO catalyst, and the promotion effect of copper was attributed to the formation and decomposition of a labile surface formate (HCOO),Cu. Ueno et al. (117) have shown earlier that surface formates are formed on zinc oxide, from CO and H,O as well as from CO, and H,, and hence an associative mechanism of the shift and reverse-shift reaction, involving formate intermediate, is believed to operate on many oxide catalysts. Curhon dioside und h)3drogen also interact with the formation of surface formate. This was documented for ZnO by the IR investigation of Ueno et NI. (117) and, less directly, by coadsorption-thermal decomposition study (84). Surface complex was formed from CO, with H, at temperatures above 18OoC,which decomposed at 300°C with the evolution of carbon monoxide and hydrogen at the ratio CO :H r 1 : 1. When carbon dioxide and hydrogen were adsorbed separately, the CO, and H, desorption temperatures were different, indicating conclusively that a surface complex was formed from CO, and H,. A complex with the same decomposition temperature was obtained upon adsorption of formaldehyde and methanol. Based upon the observed stoichiometry of decomposition products and upon earlier reported IR spectra of CO, H, coadsorbates, this complex was identified as surface formate. Table XVI compares the thermal decomposition peak temperatures and activation energies, product composition, and surface

+

TABLE XVI Adsorbed Species, Coverages, und Desorption Activurion Energies"

Adsorbed species C02 H Formate CO, H, H,CO CH,OH

+

"

"

Major TPRS products

Peak temperature (K)

E: (kJ mol-I)

Maximum total coverage observed (Cm')

390, 420, 510, 560 420,465, 540

109, 117, 142, 153 117, 130, 150

6 x 10''

-157 160 160

3 x 1013 10" lo1"

- 580 570

580

From Ref. 84. Calculated assuming a preexponential factor of 1013 sec- '

1013

308

K. KLIER

+

coverage for the surface formate from CO, H , , formaldehyde, and methanol with the desorption temperatures for carbon dioxide and hydrogen. The coadsorption and IR spectroscopic studies cited in this section indicate that surface formate is readily formed from CO, and H, or from CO and H,O vapors. It appears that the reaction of CO and H, to formate is considerably more difficult. Moreover, while both formate and methoxide have been detected in methanol decomposition over ZnO and in the synthesis over the ZnO/Cr,O, catalysts, no positive evidence of these species has so far been gathered for the copper-based catalysts.

F. MECHANISM OF METHANOL SYNTHESIS AT Low TEMPERATURES AND PRESSURES I t has been shown conclusively that methanol can be decomposed over various catalysts to carbon monoxide and hydrogen via surface methoxide and formate, and therefore a reversed path must exist for the synthesis. This has been confirmed by the trapping experiments that detected formate and methoxide during methanol synthesis. However, there persists the principal question whether more effective mechanisms may operate at low temperatures and pressures. Since the associative formate mechanism of the synthesis from carbon monoxide requires a source of oxygen additional to that of CO, be it water, surface OH groups, or CO, through the reverseshift reaction, isotopic labeling of the feed gas may provide a clue to the mechanism: if methanol from a mixture of 3 C ' h 0 and 1 2 C s 0contains only unscrambled molecules '3CH,'60H and '2CH,'80H, the formate mechanism and any other mechanism leading to isotope exchange

.-

+

~ 1 3 ~ 1 8+0 ~ 1 2 ~ 1 6 0 ~- 1 2 ~ 1 8 0 ~ 1 3 ~ 1 6 0

where R is any substituent including the catalyst oxygen, must be ruled out. Such an experiment has in fact been performed with the use of the Rh/TiO, catalyst ( I 18), and only isotopically unscrambled methanol molecules have been found. Thus there is evidence, at least on one catalyst, for a synthesis mechanism that does not proceed via surface formate. Bowker et al. (84)also proposed a low-temperature mechanism, as a minor pathway in addition to the formate mechanism, wherein carbon monoxide and hydrogen react via formyl, formaldehyde, and methoxide to methanol over zinc oxide. The current mechanistic findings and interpretations are summarized in Fig. 28. Evidence for all pathways indicated in Fig. 28 exists but the reaction conditions and catalysts differ. It is likely that palladium, platinum, and

METHANOL SYNTHESIS

309

FIG.28

iridium and copper metal will behave similarly as the Rh/TiO, catalyst, and therefore they were assigned the pathway bypassing the formate intermediate. It is not known at the present time whether methanol synthesis over the Cu/ZnO catalysts runs the course similar to that over the preciousmetal catalyst or over the ZnO catalysts. In the latter case the specificity of the promotion effect of copper would require an explanation analogous to that proposed for the function of the CuO/MgO-shift catalyst, namely that ionic state of copper in the oxide matrix accelerates not only the decomposition but also hydrogenation of the surface formate. The specific composition of surface intermediates on the Cu/ZnO catalysts has yet to be determined. A mechanism involving carbon-down CO and formyl, which flips over to be hydrogenated to methoxide and methanol, has been proposed by Kung (119) and by Bowker et al. (84) for the low-temperature synthesis on ZnO. This mechanism would also bypass the formate route and involves a crossing of pathway I into pathway I1 prior to methanol formation. Further mechanistic probing into the synthesis pathways over the copperbased catalysts appears very desirable. In particular, trapping, isotope, and spectroscopic studies may reveal further mechanistic details. It is important to conduct such investigations under the synthesis conditions, since at other conditions intermediates irrelevant to the synthesis may become prominent. Despite our limited knowledge of the mechanism, however, significant progress has been made in the characterization of the active component of the copper-based catalysts, in the modeling of the CO/CO, dependence of the synthesis rate, based on sound mechanistic features, and in the understanding of the physicochemical concepts according to which selective methanol catalysts may be chosen and further developed.

3 10

K . KLIER

REFERENCES 1. Patart, M., French Patent 540,343 (Aug. 1921). 2. Audibert, E., and Raineau, A,, h d . Eng. Chem. 20, 1105 (1928). 3. Sabatier, P., and Senderens, J. B., Ann. Chim. Plzys. 4,418 (1905). 4. Sabatier, P., and Mailhe, A., C . R. 146, 1376 (1908). 5 . Lormand, C., Ind. Eny. Chem. 17,430 (1925). 6 . Frolich Per, K., Fenske, M. R., and Quiggle, D., Ind. Eny. Chem. 20, 694 (1928). 7. Frolich Per, K., Fenske, M. R., Taylor, P. S.. and Southwich, C. A., Jr., Intl. Eng. Chem. 20, 1327 (1 928). 8. Frolich Per, K., Davidson, R. L., and Fenske, M. R., Ind. Eny. Chem. 21, 109 (1929). 9. Kostelitz, O., and Huttig, G . F., KolloidZ. 67, 265 (1934). 10. Natta, G . , Ccrtulysis 3, 349 (1955). 11. Brocker, F. J., Marosi, L., Schroder, W., and Schwarzmann, M., German Patent 2,056,612 (May 31, 1972); assigned to Badische Anilin- & Soda-Fabrik AG. 12. Catalysts and Chemicals, Inc., German Patent 1,965,007 (Oct. 15, 1970). 13. Collins, B. M., German Patent 2,302,658 (Aug. 2, 1973); assigned to Imperial Chemical Ind., Ltd. 14. Shishkov, D. S.. and Kasabova, N. A., Dokl. Bulg. Akad. Nauk 27, 73 (1974). 15. French Patent 1,489.682 (July 21, 1967); assigned to Imperial Chemical Ind., Ltd. 16. French Patent 2,037,567 (Dec. 21, 1970); assigned to Imperial Chemical Ind., Ltd. 17. Kotera, Y., Oba, M.. Ogawa, K., Shimomura, K., and Uchida, H.. in “Preparation of Catalysts” (B. Delmon, P. A. Jacobs, and G. Poncelet, eds.), pp. 489-597. Elsevier, Amsterdam, 1976. 18. Stiles, A. B.,German Patent 2,320,192(0ct. 25,1973); assigned to E. I. duPont de Nemours. 19. Lender, Yu. V., Tsybina, E. N., Popov, I . G., Pirozhenko, L. F., and Petrishcheva, G. S., Khim. Prom. (Moscow) 49,899 (1973). 20. Oliver, R. B., German Patent 1,229,990 (Dec. 8, 1966); assigned to Power-Gas Corp., Ltd. 21. Eguchi, T., Yamamoto, T., Yamauchi, S., Kuraishi, M., and Asakawa, K., U.S. Patent 3,256,208 (June 14, 1966); assigned to Japan Gas-Chemical Co. 22. Brocker, F. J.. German Patent 2,116,949 (Oct. 19, 1972); assigned to Badische Anilin- & Soda-Fabrik A.G. 23. Jorgensen, M. H., and Rushede, K., German Patent 2,016,596 (Dec. 3, 1970); assigned to Topsoe, Haldor Frederik Axel. 24. Davies, P.. and Snowdon, F. F., U.S. Patent 3,326,956 (June 20, 1967); assigned to Imperial Chemical Ind., Ltd. 25. Baron, G., Bechtholdt, H., Bratzler, K., Leibgolt, H., and Ehrland, E., German Patent 1,300,917 (Aug. 14, 1969); assigned to Metallgesellschaft AG. 26. Denny, P. J., and Whan, D. A,, Catalysis London 2, 46 (1978). 27. Andrew, S. P. S., Post Conyr. Symp. Int. Congr. Catal., 7th, Oscrka, Paper 12 (July 1980). 28. Broden, G . , Rhodin, T. N., Bruckner, C., Benbow, R.. and Hurych, 2.. Surf: Sci. 59,593 (1976). 29. Klier, K., Zettlemoyer, A. C., and Leidheiser, H., Jr., J . Chrm. P h p . 52, 589 (1970). 30. Hayward, D. 0.. and Trapnell, B. M. W., “Chemisorption,” 2nd ed. Butterworths, London, 1964, p. 23 1. 31. The catalytic activity of copper metal is reviewed on pp. - . 32. Poutsma, M. L., Elek, L. F., Ibarbia, P. A,, Risch, A. P., and Rabo. J. A,, J . Ctrttrl. 52, 151 (1978).

31 1

METHANOL SYNTHESIS

33. Ichikawa, M., Bull. Chem. Soc. Jpn. 51,2268 (1978). 34. Ellgen, P. C . , and Bhasin, M. M., U S . Patents 4,014.913 (March 1977); 4,096,164 (June 1978); and 4,162,262 (July 1979). 35. Early reports ( I , 2) indicate that the activity of oxides follows the pattern ZnO MnO, Cr,O, > BeO, Ce,O,, UO,, ZrO,, oxides of AI, Si, Mo, W. V, Ti, Th, Mg, Ca, Ba, and Sr being inactive. 36. Klier, K., to be published (1982). 37. Bulko, J. B., Ph.D. dissertation, Lehigh University, 1980 (Univ. Microfilms Int., Order No. 8102497). 38. Bulko, J. B., Herman, R. G., and Klier, K.. J . Am. Ceram. Soc. in press (1 982). 39. Herman, R. G., Klier, K., Simmons, G. W., Finn, B. P., Bulko, J. B., and Kobylinski, T. P., J . Cutal. 56,407 (1979). 40. Mehta, S., Simmons. G. W.. Klier, K., and Herman, R. G., J . Catul. 57, 339 (1979). 41. Bulko, J. B., Herman, R. G., Klier, K., and Simmons, G. W.. J. Phys. Chem. 83, 31 18 (1979). 42. Parris, G. E., Ph.D. Dissertation, Lehigh University, 1981. 43. Parris, G. E., and Klicr, K., to be published (1982). 44. Chapple. F. H., and Stone, F. S., Proc. Br. Cerum. Soc. I, 45 (1964). 45. Schiavello, M., Pepe, F., and DeRossi, S., J . Phvs. Chem. 92, 109 (1974). 46. Kochloefl, K., Proc. h i t . Congr. Catal., 7th. Tokyo 1980 p. 486 (1981). 47. Notari, B., discussion t o paper A 32, Proc. Int. Congr. Cat& 7th, Tokyo fY80 p. 487 (1981). 48. Zettlemoyer, A. C., Yu. Y. F., Chessick, J. J., and Healey, F. H., J . Phys. Chem. 61, 1319 (1957). 49. Kington, G. L., and Holmes, J. M., Trans. Faruduy Soc. 49,425 (1953). 50. Vasilevich, A. A., Shpiro, G . P., Alekseev, A. M., Semenova, T. A,, Markina, M. I., Vasil’eva, T. A., and Budkina, 0. G., Kinet. Cutal. (English ed.) 16(6), 1363 (1975). 51. Yao, Y.-F. Y., J . Phys. Chem. 69, 3930 (1965). 52. Kummer, J. T., and Yao, Y.-F. Y.. Cun. J . Chem. 45,421 (1967). 53. Sherwin, M., and Blum, D., “Liquid Phase Methanol.” Report to the Electric Power Research Inst., EPRI AF-202, Research Project 317-1 (August 1976). 54. Simmons, G. W. et a/., to be published. 55. Chatikavanij, V.. M. S. Thesis, Lehigh University, 1980. 56. Klier. K . , Chatikavanij, V., Herman, R. G . ,and Simmons, G . W., J . Cutul. 74,343 (1982). 57. Natta, G., Pino, P., Mazzanti, G., and Pasquon, I., Chim. Ind. 35, 705 (1953). 58. Pasquon, I . , and Dente, M., J . Cutal. 1, 108 (1962). 59. Wermann, J., Lukas, K., and Gelbin, D., Z . Phys. Chem. 225,234 (1964). 60. Bakemeier, H., Laurer, P. R., and Schroder, W., Chem. Eng. Prog. Symp. Ser. 66(98), 1 (1970). 61. Leonov, V. E., Karavaev, M. M., Tsybina, E. N., and Petrishcheva, G. S., Kinet. Cutul. (English ed.) 14, 848 (1973). 62. van Herwijnen, T.. and deJong, W. A,, J . Card. 34,209 (1974). 63. Vedage, G., and Klier, K., J . C u r d in press (1982). 64. Vannice, M. A., J . Cutul. 37, 449 (1975); reports turnover rate of 3.2 x lo-, sec-’ at 275°C for hydrogenation of CO to methane over nickel catalysts, while turnover rate for benzene hydrogenation can be estimated at 0.77 x 10- sec- at 66°C and a minimum of0.46 sec- at 275°C from data given in the report by Danes, V., Cabicar, J . , Grubner, O., Klier, K., and Jiru, P., Chem. Listy 50, 1049 (1956); see also Ref. 63. 65. Cotton, F . A., and Marks, T. J . , J . Am. Chtw. Soc. 92, 51 14 (1970). 66. Herberhold, M., “Metal n-Complexes,” pp. 227-232. Elsevier, Amsterdam, 1972. ,

’

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67. Cotton, F. A,, and Wilkinson, G . , “Advanced Inorganic Chemistry,” 3rd ed. Wiley, New York, 1972. 68. Baglin, E. G., Atkinson, G. B., and Nicks, L. J., Ind. Eng. Chem. 20,87 (1981). 69. Scholten, J. F. F., and Konvalinka, J. A,, Trans. Faruday Soc. 65,2465 (1969). 70. Ponec, V., Proc. Int. Congr. Cutul., 7th,Tokyo 1980 p. 486 (1981). 71. Ichikawa, M., Bull. Chem. Soc. Jpn. 51, 2273 (1978). 72. Ichikawa, M., and Shikakura, K., Proc. Int. Congr. Catul., 7th, Tokyo 1980 pp. 925-940 (1 98 I). 73. Klier, K., and Herman, R. G., unpublished results. 74. Ellgen, P. C., Bartley, W. J., Bhasin, M. M., and Wilson, T. P., Adv. Chem. 178,147 (1979). 75. Klier, K., Proc. Conf Catal. Organ. Synth., 8th pp. 75-89 (1981). 76. Fischer, A., Hosemann, R., Vogel, W., Koutecky, J., Pohl, J., and Rrilek, M., Proc. Inr. Congr. Cutul., 7th,Tokyo 1980 p. 341 (1981). 77. Herman, R. G.,Simmons, G. W., and Klier, K., Proc. Int. Congr. Cutul., 7th, Tokyo 1980 p. 475 (1981). 78. Sugier, A,, and Freund, E., US.Patent 4,122,110 (Oct. 1978). 79. Sugier, A,, and Freund, E., Fr. Patent 2,369,234 (May 1978). 80. Sugier, A,, and Freund, E., Ger. Offen. 2,748,097 (May 1978). 81. Wachs, I. E., and Madix, R. J., J . Catal. 53, 208 (1978). 82. Bowker, M., and Madix, R. J . , Surf Sci.95, 190 (1980). 83. Madix, R. J., Adv. Cutal. 29, 1 (1980). 84. Bowker, M., Houghton, H., and Waugh, K. C., J. Chem. Soc. Faraduy Trans. I, 77,3023 (1 98 I ) . 85. Borowitz, J. L., J. Catal. 13, 106 (1969). 86. Deluzarche, A,, Kieffer, R., and Muth, A,. Tetrahedron Lrtt. 38, 3357 (1977). 87. Fischer. E. O., Adu. Organomet. Chem. 14, I (1976). 88. Ueno. A,, Onishi, T., and Tamaru, K., Trans. Furaduy Soc. 67,3585 (1971). 89. Eischens, R. P., Pliskin, W. A,, and Low, M. J. D., J. Cutul. 1, 180 (1962). 90. Chang, C. C., Dixon, L. T., and Kokes, R. J., J . Phys. Chem. 77, 2634 (1973). 91. Dent, A. L., and Kokes, R. J., J . Phys. Chem. 73, 3772 (1969). 92. Boccuzzi, F., Borello, E., Zecchina, A,, Bossi, A,, and Camia, M., J . Catal. 51, 150 (1978). 93. Carrizosa, I., Munuera, G., and Castafiar, S., J . Catul. 49, 265 (1977). 94. Kortum, G., and Knehr, H., Z. Phys. Chem. N.F. 89, 194 (1974). 95. Mikovsky, R. J., Boudart, M., and Taylor, H. S., J . Am. Chem. Soc. 76, 3814 (1954). 96. Ostrovskii, V. E., Dyatlov, A. A,, and Ogneva, T. P., Kinet. Catal. (English ed.) 19, 410 (1978). 97. Tracy, J . C.,J . Chem. Phys. 56, 2748 (1972). 98. Pritchard, J . , Catterick, T., and Gupta, R. K., Surf: Sci. 53, 1 (1975). 99. Doyen, G., and Ertl, G., Surf. Sci. 43, 197 (1974). 100. Papp, H., and Pritchard, J., Surf: Sci. 53, 371 (1975). 101. Alexander, C. S., and Pritchard, J., J . Chem. Soc. Furuduy Truns. I, 68, 202 (1972). 102. Thurston, E. F. W., Trans. Faruday Soc. 64,2181 (1968). 103. Smith, A. W., and Quets, J. M., J. Catal. 4, 163 (1965). 104. Davydov, A. A,, Boreskov, G. K., Yur’eva, T. M., and Rubene, N. A,, Dokl. Akud. Nuuk USSR 236, 1402 (1977). 105. Blyholder, G., and Wyatt, W. V., J . Phys. Chem. 70, 1745 (1966). 106. Kagel, R. O.,and Greenler, R. G., J . Chem. Phys. 49, 1638 (1968). 107. Soma, Y., Onishi, T., and Tamaru, K., Trans. Furaduy Soc. 65, 2215 (1969); Greenler, R. G., J. Chem. Phys. 37,2094 (1962).

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313

108. Treibmann, D., and Simon, A,, Ber. Bunsenges. Phys. Chem. 70,562 (1966). 109. Arai, H., Take, J., Saito, Y . ,and Yoneda, Y . ,J. Cutul. 9, 146 (1967). f10. Kagel, R. O., J . Phys. Chem. 71,844 ( 1 967). f I f . Deo, A. V., and Dalla Lana, I. G., J . Phys. Chem. 73, 716 (1969). 112. Hertl, W., and Guenca, A. M., J . Phys. Chem. 77, 1120 (1973). 113. Herwijnen, T., van, and deJong, W. A,, J . Catul. 63,83 (1980). 114. Tsuchiya, S., and Shiba, T., Bull. Chem. Soc. Jpn. 38, 1726 (1965). 115. Nagarjunan, T. S., Sastri, M . V. C., and Kuriacose, J . C., J . Cutul. 2,223 (1963). f16. Boccuzzi, F., Garrone, E., Zecchina, A,, Bossi, A., and Camia, M., J . Cutul. 51,160 (1978). 117. Ueno, A,, Onishi, T., and Tamaru, K., Trans. Furuduy Sor. 66, 756 (1970). f l 8 . Takeuchi, A,, and Katzer, J. R.. J . Phys. Chem. 85,937 (1981). 119. Kung, H. H., Cutul. Reu. 22, 235 (1980).

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Author Index Numbers in parentheses are reference numbers and indicate that an author’s work is referred to although his name is not cited in the text. Numbers in italics show the page on which the complete reference is listed

A Albo, S. 87(36), 91(36), 127(36). 129 Abdulahad, I., 55(241), 75 Aben, P. C., 21(117), 72 Aberdam, D., 143(45), 144(45), 236 Ablow, C. M., 151(53), 157(53), 161(53), 188 (195), 192(195), 198(195), 199(195), 200 (195), 21 1(195), 212(195), 213(195), 214 (195), 219(195), 236,239 Abragam, A., 96(80), 130 Abraham, M. M., 89(47), 96(77, 78, 79), 126 (77, 78, 79), 129, 130 Adams, C . J., 157(132),238 Adams, E. O., 213(229), 240 Agafonov, A. V., 19(103),20(103), 67(103), 71 Agievskii, D. A., 19(103), 20(103), 67(103), 71 Agrawal, P. K., 154,1157(10 1, 13 I), 116 1 ( 10I), 162(101), 163, 165(101), 185(101), 186, 189(101), 192(147, 205, 206, 207, 208), 194(205), 198(101, 147, 205, 206, 207, 208), 201(101, 147, 205, 206, 207, 208), 202(101,147,205,206,207,208),203(10l, 147), 204(101, 147), 204(101, 147), 208 (lot), 210(147), 211(101), 237,238,240 Aika, K., 103, 105(114), 131, 225(254), 227 (254), 241 Aizawa, T., 153(97), 237 Akimoto, M., 125(178), 133 Alcock, C. B., 140(20), 235 Alderson, T., 28(142), 72 Alekseev, A. M., 267(50), 311 Alexander, C. S., 303(101), 312 Alsdorf, E., 14(37), 69 Alstrup, I., 154(106b), 155(106b), 165, 166 (106b), 237 Altshuler, 0. V., 14(35, 36, 37, 44), 69 Andersen, T., 97(87), 130 Anderson, R. B., 185(179), 191, 190(199), 191, 229(271), 239,240,242

Andersson, A,, 128, 128 Andersson, S., 143(31a, 31b), 146(31a, 31b). 235,236 Andrews, L., 98(88), 130 Andrews, S. P. S., 248(27), 267(27), 274(27), 275(27), 276, 283(27), 288, 292, 311 Anpo, M., 93(71a), 94(71b, 71c, 71e), 103 (IIOb), 104(111b), 120(158, 159, 160a), 121(71b, 71e, 160a), 123, 125(71b), 130, 131, 132 Antoshin, G. V., 28(145), 72 Anufrienko, V. F., 11(16), 12(21), 13(16), 16 (16, 21), 64(21), 68, 69 Apel’baum, L. O., 220(240), 222(240), 241 Aptekar, E. L., 97(85), 130 Arai, H., 15(47), 16(51), 17(51), 21(121, 122), 22(122), 61(47), 63(47), 64(51),69,70, 72, 303(109), 313 Argano, E. S., 185(183), 239 Arieti, A,, 97(86), 130 Arpe, H. J., 16(48), 38(48), 69 Artemov, A. V., 18(67), 70 Asakawa, K., 250(21), 310 Ashmore, P. G., 107(128), 132 Asrieva, V. D., 19(103), 20(103), 67(103), 71 Atanasova, V. D., 29(148), 73 Atkins, P. W., 98(94a), 130 Atkinson, G. B., 287(68), 288(68), 312 Atwood, K., 21 1(223), 240 Audibert, E., 245(2), 310 Auzins, P., 126(189), 133 Avaev, V. I., 20(113), 71 Avetisyan, R. V., 24(138, 139), 25(139), 72

B Baglin, E. G., 287, 288, 312 Baird, M. J., 122,133, 196,211(210,222), 240 Bakemeier, H., 276, 311 Baker, M. O., 225, 228(248), 241

315

316

AUTHOR INDEX

Bakhshi, N . N., 52(233), 53(233), 75 Balistreri, S., 125(182), 133 Ballivet-Tkatchenko, D., 58(254c), 76 Baptista, J. L., 97(87), I30 Barbier, J . , 225(251), 228(251), 241 Barbouth, N., 169(153), 238 Barin, I., 139(16), 140(16), 235 Baron, G., 250(25), 310 Barraclough, C. G., 124(174), 133 Barret, P., 153(98), 237 Bartholomew, C. H., 142(23), 154(112, 113), 155(112), 157(23, 112, 113, 140), 160(112, 113, 140), 161(112, 113), 168(112, 140), 173(23,112,113,141,161b,162),175(112, 141, 162), 176(141, 162), 177(141, 162). 178(141, 162), 179(141, 162), 180(157a, 157b, 162), 181(141, 162), 183(162, 176, 177), 185(157b), 188(113, 194, 197), 189 (140, 197), 192, 193, 195, 196, 197(194), 198(194), 200, 201, 209(140), 210(140), 211(23, 113, 194), 212(23, 113, 194), 213 (140), 214(140), 229(194), 230, 232, 235, 237,238,239,240,242 Bartley, W. J . , 290(74), 291(74), 312 Barton, A,, 109(134b), 132 Battrel, C. F., 156(119), 238 Baudoing, R., 143(45), 144(45), 236 Bayer, J . , 185(179), 239 Bechtold, E., 151(85), 152(85), 161(85, 145), 162, 169(85), 237, 238 Bechtold, G., 151, 162, 237 Bechtholdt, H., 250(25), 310 Becker, G . E., 143(33), 144(33), 146(33), 236 Belozerov, A. L., 19(103), 20(103), 67(103), 71 Benard, J., 151(57, 67), 169, 171(57), 172(57), 186(57), 236,238 Benbow, R., 251(28), 252(28), 311 Ben Taarit, Y., 84(30a), 85(31), 86(31), 90 (30a), 91(30a, 31), 101(101), 102(108b), 104(112, 113), 127(30a, 31). 129, 131 Benziger, J., 185(182), 239 Bernard, J., 143(27), 146(27), 147(27), 235 Berthier, B., 151(66), 236 Berthier, Y., 151(75, 80), 161(75), 162(75), 169(84, 153, 155), 237,238 Besenjei, G., 20(107), 71 Besoukhanova, T., 30( 159), 62(159), 64(159), 73 Beyer, H. K., 13, 59(259), 60(261), 67(259, 261), 69, 76

Bhasin, M . M., 253(34), 289(34), 290(74), 291 (34, 74), 311,312 Bhatia, S., 52, 53(233), 75 Biberian, J . P., 151(71), 152(71), 236 Bielanski, A., 94(76), 107(123), 130, 131 Bill, H., 85(33), 126(33), 129 Bird, C . W., 30(161), 73 Bishop, R. J . , 59(260), 60(260), 61(260), 67 (260), 76 Blake, W. B. J . , 87(42), 129 Blakely, J., 143(38), 145(38), 236 Blanchard, M., 58(254), 76 Blandamer, M. J., 84(27) 127(193), 128, 133 Bleaney, B., 96(80), 130 Blinov, B. B., 19(91,92), 93, 56(91,92,93), 71 Blum, D., 271(53), 311 Blunt, F. J., 97(81), 130 Blyholder, G., 154(105, 108), 156(108), 165 (105, log), 176(108), 186, 303(105), 237, 239,313 Boatner. L A , , 96(78), 126(78), 130 Bobrov, N. N., 11(16), 12(21, 22, 23), 13(16, 23), 16(16, 17, 19, 21, 22, 23), 18(69), 53 ( 2 3 ) , 64(21,22, 23), 68,69, 70 Boccuzzi, F., 122(166b, 166d), 133, 302(92), 304, 305,312,313 Bogdanovic, B., 32(170), 73 Bohart, G. S., 213(229), 240 Bohme, D. K., 103(110a), 131 Bonzel, H. P., 151, 152(92), 153(92), 185, 186 (92), 229(92), 230(92), 231(83b, 92). 232 (92), 237 Bordoli, R. S., 156(116), 159, 176(116), 177 (1 16), 178,238 Borello, E., 124(173b), 133, 302(92), 312 Boreskov,G. K., 11, 12, 13(16), 16(16, 17, 18, 19, 21, 22), 18(64), 64(18, 21, 22), 68, 69, 70, 107(125), 131,303(104),304(104), 307 (104), 313 Borowitz, J . L., 301(85), 312 Bosacek, V., 29(154), 73 Bossi. A , , 304(116), 305(116), 313 Bottoms, W. R., 185(189), 239 Boudart, M., 16(25, 26), 19(95), 56(95), 64(25, 26), 69, 71,94, 130, 140(18), 153(95), 197 (18), 235,237, 302(95), 312 Bowen, D. O., 154(105), 165(105), 237 Bowker, M., 298(82), 300(82, 84), 303, 306, 307(84), 309,312 Bradley, E. B . , 212(227), 240

AUTHOR INDEX

Brailsford, J. R., 79(7, 8), 80(8), 81(7, 8), 85 (7, 8), 127(7, 8), 128 Bratzler, K., 250(25), 310 Bray, P. J., 81(18), 128 Breck, D. W., 3,68 Brennan, J. A,, 56(244), 68(244), 75 Breysse, M., 120(157), 132 Bridger, G. W . , 216(236), 217(236), 241 Brill, R., 220(241, 243), 221,222(241), 241 Broden, G., 251,252(28), 311 Brocker, F. J., 249(1 l), 250(22), 310 Brown, D. J., 20(116), 72 Brucker, C. F., 185,239 Bruckner, C., 251(28), 252(28), 311 Buckley, D. H., 151(69), 156(69), 236 Budkina, 0. G., 267(50), 3 1I Bulko, J. B., 255(37, 39), 257(38, 39, 41), 258 (39, 41), 259(39), 260(39, 41), 261, 262 (41), 266(39), 267(39), 271(39), 275(39), 283(39), 287(39), 293, 296(39), 311 Bulatnikov, Yu. I., 220,222(240), 241 Burdshanadze, M. N., 22(124), 72 Burkkhardt, I., 14(37), 6Y Bunvell, R. L., 157(135), 238 Butler, G. A,, 230(275), 242 Bystrikov, A. V., 91(63), 129

C Cabane-Brouty, F., 151(53, 54a, 54b, 55, 56, 57, 64), 157(53), 161(53, 54a, 55, 56), 169 (57), 171(57), 172(57), 186(57), 236 Cabicar, J., 285(64), 312 Caesar, P. D., 56(244), 68(244), 75 Cagle, G. W., 154(108), 156(108), 165(108), l76( log), 237 Camia, M., 304(116), 305(116), 313 Candlin, J . P., 32(169), 73 Capehart, T. W., 138, 143(5), 151(5), 235 Carrizosa, I., 302(93), 312 Carroll, W. M., 20(114, 115), 21, 71, 72 Castafiar, S., 302(93), 312 Catterick, T., 303(98), 312 Centola, P.,47(225), 48, 75, 123(173b, 175), 133 Chang, C. C., 302(90), 312 Chang, C. D., 56(243), 68(243), 75 Chang, Y . C., 225(258, 259), 228(258, 259), 241

317

Chao, J. L., 143(46), 145(46), 157(46), 176(46), I77(46), 236 Chao, R. E., 59(258), 61(258), 67(258), 76 Chapple, F. H., 261(44), 311 Chatikavanij, V., 275 (55, 56), 277(56), 278 (56), 279(56), 280(56), 281(56), 282(55, 56), 283(56), 284(55, 56), 292(56), 296 (56), 297(56), 309(56), 311 Chauvin, Y., 24(133), 25(133), 28(133), 72 Che, M., 79(6), 83(23, 25), 84(29), 90(52, 55), 94(72), 97(82), 98(52), 101, 102(6, 102, 108c), 106(119), 109, 110, 111(133, 138), 125,128, 129,130, 131, 132,133 Chemagina, V. P., 38(190), 74 Chen, N. Y . , 10(14), 13,68,69 Chen, Y., 89(46, 47), 96(77, 78, 79), 126(77, 78, 79), 129, 130 Cheng, T. C., 83(20), I28 Cherkashin, A. E., 127(30b), 129 Chessick, J. J., 267(48), 311 Chester, A. W., 13(31), 69 Chinchen, G. C., 216(236), 217(236), 241 Christensen, B., 42, 43, 44(21 I), 45(208, 21 I), 46, 65(21 I), 74 Chudinov, M. G., 220(244), 222(244), 241 Chukin, G. D., 19(102, 103), 20(102, 103), 67 (102, 103), 71, 225(253), 227(253), 241 Chung, K. S., 161(141), 173(141, 158), 175 (141), 176(141, 158), 177(158), 178(141), 179(l41), 180( 158), 181( l41), 185(158), 238,239 Chuvylkin, N. D., 83(19, 24b), 87(19), 90(19, 24b), 98(93), 99(93), 102(93), 105(116), 106(19), 127(19, 24b), 128, 130, 131 Cid, R., 23(126), 72 Cirnino, A,, 109(134a), 132 Cinti, R. C., 143(41, 44), 144(41, 44), 146(41, 44), 159(41), 236 Ciric, J., 56(244), 68(244), 75 Clarke, A,, 20(106), 71 Clarke, J. S., 38, 74 Claudel, B., 120(157), 132 Codell, M., 80(11), 128 Cognion, J. M., 79(4), 125(4), 128 Colby, S. A,, 138(9), 153(9), 159, 184(9), 229 (9), 230(9), 231, 232, 235 Collins, B. M., 249(13), 310 Colson, J . C., 153(98), 157(137, 1381, 158(137, 138), 237, 238 Coluccia, S., 94(75), 109, 113(149), 1l4(148,

318

AUTHOR INDEX

149, 150). 115(148), 116, 119, 122, 124 (173b). 130. 132, 133 Compton, W. D., 112(143), 132 Condurier, G., 58(254c), 76 Contractor. A . Q., 151(90), 162, 237 Cotdischi, D., 97(86), 100(98b), 109, 110(98b, 135), 1 1 I , 122, 130, 131, 132, 133 C o m a , A,, 23, 72 Cornet, D., 16(52), 70 Cosyns, J., 185(191), 186, 225(191, 256), 226 (256), 228(191), 239, 241 Cotton, F. A., 79, 128, 285(65), 287(67), 312 Coughlan, B., 20(114, 115). 21, 71, 72 Cowan, D. L., 87(42), 129 Craddock, J . H., 39(196), 40(196), 41(196), 42 (196, 207), 45(196), 74 Cramer, R., 25(141a), 72 Crell, W., 176(168), 181(168), 182, 230 Csuros, Z . , 20(107), 71 Curtis, J. L. S., 225, 228(248), 241 Cusumano, J. A., 140(18), 197(18),235

D Dadyburgor, D. B., 125(177b), 133 Dahl, L. F., 143(37), 145(37), 236 Dainton, F. S., 90(57), 129 Dalla Betta, R. A., 16(26), 19(95), 56(95), 64 (26),6Y, 71. 192, 193(204), 194,212(225), 240 Dalla Lana, I. G., 303(11 I), 313 Danes, V., 285(64), 312 Darabont, A,, 127(191), 133 Datka, J., 19(104), 20(104), 67(104), 71 Davidova, N . , 19(98), 56(98), 71 Davidson, R. L., 246(8), 310 Davies, P.. 250(24), 310 Davydov, A. A., 12(22,23), 13(23), 16(22,23), 20(110), 53(23), 64(22, 23), 69, 71, 303 (304), 304(104), 307,313 Davydova, L. P., 18(64), 70 Deane,A. M.,94(75),113(149), 114(148, 149), 11 5( l48), 1 16(l49), 130, 132 de Boer, J. J., 8(1l), 33(179), 35(179), 62(179), 64(179), 65(179), 68, 73 de Boer, N . H., 124(170), 133 de Jong, W. A., 283(62), 304(113), 307, 312, 313 Delbouille, A. J., 94(73), 130 Delescluse, P., 143(47), 145(47), 146(47), 148 (47), 150, 236

Delgass, W. N., 16(25), 64(25), 69 Del Rosso, R.,47(225), 48(225), 75 Deluzarche, A., 301(86), 304(86), 306,312 Delzer, G., 196(210), 21 1(210), 240 Dempsey, E., 8(10), 23( 129,130), 64( 129,130). 68, 72 Demuth, J. E., 138(6), 143(6, 29a, 29b, 34), 144(6,34), 146(6,29a, 29b, 34), 165(29a), 235, 236 Den Besten, I. E., 154, 155(104), 156(104), 160 (104), 165(104), 174, 237 Denny, P. J., 55(240), 75, 248(26), 276,310 Dent, A. L., 302(91), 312 Dente, M., 276(58). 311 D e o , A . V . , 303(111),313 Derks, F.. 17(58), 70 de Rosset, A. J., 157(132),238 De Rossi, S . , 261(45), 311 Derouane, E. G., 56(242b), 63(242b), 75, 94 ( 7 3 , 130 de Siebenthal, J. M., 85(33), 126(33), 129 de Wilde, W., 59(259), 67(259), 76 Dillard, J. G., 156(119), 238 Dimitrov, C . , 22( 125), 72 Dimitrova, R. P., 22(125), 72 Dini, P., 19(100), 56(100), 71 Dirksen, H. A., 229(269), 230(269). 242 Dixon, L. T., 302(90), 312 Dobashi, H. H., 229(274), 230(274), 242 Doering, W., 35(183), 73 Domange, J . L.. 151(59, 61, 62, 63), 152(63), 236 Domasheveskaya, E. P., 138(3), 235 Doneo, D., 19(100), 56(100), 71 Downing, R. S., 8(1 I), 33(177, 178). 62( 177, 178, 179), 63(177), 64(177), 178, 179), 65 (179), 68, 73 Doyen, G., 303(99), 312 Drent, E., 33(175), 62(175), 64(175), 73 Druon, C., 1 1 I , 132 Druzhkov, V. N., 29(153), 73 Dubinsky, Yu. G., 24(132), 72 Dudzik, Z., l7(6l), 70 Dufaux, M., 1 Il(l38), 132 Dugdale, L. A., 229(273), 230(273), 242 Duke, C. B.. 143(40), 144(40), 146(40), 236 Dumas, P., 157(137, 138), 158(137, 138), 238 Dunn, T. M., 83(21), 128 Dyatlov, A. A,, 303(96), 312 Dzhaparidze, R.V., 45(214,216,217),46(215, 219), 74, 75

AUTHOR INDEX

E Ebisawa, M., 25(140), 26(140), 28(140), 29 (140), 62(140), 64(140), 72 Echigoya, E., 125(178), 133 Edmonds, T., 143(35, 36a, 36b), 145(35), 146 (35, 36b), 149(35, 36a, 36b), 152(35), 236 Eguchi, T., 250(21), 310 Ehrland, E., 250(25), 310 Eidus,Ya. T.,24(134, 135, 136, 137, 138. 139), 25(139), 28(144), 29(155, 156, 157), 38 (189, 190, 191, 192, 193), 40(199, 202, 203), 41(205), 42(203, 205), 43(203), 44 (203), 45(217), 64(135), 65(203), 72, 73, 74, 75 Eischens, R. P., 302(89), 312 Ekedahl, E., 213(230), 240 Eley, D. D., 91(61), 107(122a), 129, 131, 253 (32), 289(32), 290(32), 291(32), 311 Ellgen, P. C., 253(34), 289(34), 290(74), 291 (34, 74). 311, 312 Elliott, D. J., 52(234), 53, 54(234), 55(234), 63 (234), 67(234), 75, 139(15), 140(15), 235 Epishina, G. P., 22(124), 72 Epishina, Z. V., 22(123), 72 Erekson, E. J., 157(140), 160(140), 161, 163, 168(140), 188(140, 197), 189(140, 197). 193(140), 201(140), 202(197), 209(140), 210(140), 213(140), 214(140), 238,239 Erley, W., 143(49), 145(49), 146(49), 148(49), 149(49), 150, 176(49), 177(49), 236 Ermakov, Yu. I., 29(153), 73 Ershov, N . I., 24(134, 135, 136, 137), 29(155, 156, 157), 64(135), 72, 73 Ertl, G., 303(99), 312 Esumi, K., 110(136), 111(136), I32 Evans, E. L.,140(20), 235 Evans, H. C . , 225(268), 227,242 Eyring, E. M . , 60(263), 61(263), 67(263), 76

F Faure, L., 120(157), 132 Fehsenfield, F. C., 103(1 IOa), 131 Fenin, V . A., 120(16Ob), 121(160b), 125 (160b, d), 132 Fenske, M. R., 245(6, 7), 246(6, 7, 8), 300(6), 310 Fenin, V . A . , 121(160d), 123, 132 Ficalora, P. J . , 156(125), 185(125), 238

319

Figueras, F., 20,51,67(105), 71,225(255), 228 (2551, Finn, B. P., 255(39), 257(39), 258(39), 259 (39), 260(39), 266(39), 267(39), 271(39), 275(39), 283(39), 287(39), 296(39), 311 Finstrom, C. G., 157(132), 235 Firth, J . G., 13(33),69 Fischer, A,, 292, 293(76), 295, 312 Fischer, E. O., 302, 312 Fischer, T. E., 151(77, 86, 87, 88), 185(185), 225(88), 237, 239 Fisher, G. B., 143(32), 144(32), 146(32), 147 (32), 150(32), 151(78), 156(78), 185(19O), 186(78, 190), 236, 237,239 Fitzharris, W. D., 154(99, IOO), 157(100), 158 (100). 160, 165(99, IOO), 168, 186(99), 188 (99, 100, 196), 189(196), 192(99, IOO), 198 (99, 100). 201(99, loo), 202(99, 100, 196), 203(99, IOO), 204(99, IOO), 205(99, IOO), 206(99, IOO), 207(99, IOO), 208(99), 209 (IOO), 210(99, IOO), 211(99, IOO), 212(99, IOO), 214(99), 216(100), 236, 237 Flentge, D. R., 16(56), 70 Flockhart, B. D., 18(68), 70 Foley, J. M., 157(131), 163(131), 238 Forney, A . J., 21 1(221), 240 Forni, L., 36(186), 37(187), 62(186), 74 Fournier, M . , 83(23), 125(23), 128 Fowler, R. W., 142(23), 154(112, 113), 155 (112), 157(23, 112, 113), 160(112, 113), 161(112, 113), 165(112), 166(112, 113), 167(112, 113), 168(112, 113), 173(23, 112, 113), 175(112), 186(112, 113), 188(113), 192(23, 113), 195(23, I13), 196(23, 113), 201(113), 211(23, 113), 212(23, 113), 230 (276, 277), 235,237,242 Fraenkel. D., 58(254b), 76 Franck, J . P., 185(191), 186(191), 225(19l, 256), 226(256), 228(191), 239,241 Franck-Neumann, M., 35(183), 73 Freed, J. H., 92(65b), 129 Freel, J . , 229(271), 242 Freeman, (3. B., 143(46), 145(46), 157(46), 176 (46), 177(46), 236 Freund, E., 296(78, 79, go), 312 Freund, F., 94(74), 130 Friebele, E. J . , 91(64), 129 Frolich Per, K., 245(6, 7), 246, 300, 310 Frolov, A . M., 97(85), 130 Fueki, K., 153(97), 237 Fuentes, S., 225(255), 228(255), 241

320

AUTHOR INDEX

Fujitsu, H., 14(40), 69 Fujiwara, T., 111(140), 132 Funabiki, T., 94(71d), l23(71d), 130 Furuyama, M., 156(124), 238

G Galich, P. N., 19(94), 56(94), 71 Galinski, A . A,. 19(94), 56(94), 71 Gallad, D., 126(188), 127(188), 133 Gallezot, P., 2(6), 18(6), 19(104), 20(104), 30 (160), 53(6), 62(160),64(160),67(104),68, 71, 73

Galli, A,, 127(192), 133 Garanin, V. I . , 19(77,78,81,82,84,86,87,88, 89, 90, 91, 92, 96), 28(144, 145, 146), 56 (96), 70, 71, 72, 73 Gardner, D. C., 200,201(216), 239,240 Garland, C. W., 154(102), 156(102), 165(102), 176(102), 181(102), 182, 237 Garrone, E., 117, 120(154), 123, 132, 133,304 (116), 305(116), 313 Garten, R. L., 16(25, 26), 64(25, 26), 69 Ganvood, W. E., 56(244), 68(244), 75 Gates, B. C., 40(200), 47(222), 58(254b), 74, 75. 76

Gautheir, Y . , 143(45), 144(45),236 Gelbin, D., 276(59), 311 Genet, M., 57(252), 58(252), 76 Gentry, S. J., 14, 6 9 Gentsch, H., 94(74), 130 Germain, J. E., 107(124), 131 Ghiotti, G., 122(166b, 166d), 133 Ghorbel, A , , 92(65a), 129 Gibson, M . A . , 105(115), 131 Gikis. B. J., 213(233), 219(233), 240 Gil, J. M., l85(191), 186(191), 225(191, 256), 226(256), 228( l91), 239, 241 Giordano, N., 19(100), 56(100), 71 Glascock, H. H., 118(153), 132 Glass, R. 157(133),238 Glueckaut, E., 213(232), 240 Goddard, W. A , , 165,235,238 Gogol, N. A., 19(97), 56(97), 71 Gomez, R., 20(105), 51(105), 67(105), 71 Goncharova,N. V., 19(102,103),20(102,103), 67(102, 103), 71, 225(253), 227(253), 241 Gonzalez-Elipe, A. R., 99(95, 96), 102(108c), 107(96), 125(184), 130, 131, 133

w.,

Gonzalez-Tejuca, L., 225(254), 227(254), 241 Goodman, D. W . , 186. 198(193), 210,239 Goodman, W . D., 173(164), 175, 176(164), 177(164), 178(164), 239 Gorring, R. L., IO(l5), 68 Goryachev, A. A , , 19(81, 82), 70 Could, R., 157(130), 163(148), 238 Grabke, H. J., 161(142), 162(142), 169(154), 185(186), 220(154), 223, 224(154), 230 (142), 238, 239 Grabowski, E . , 16(53), 70 Greenler, R. G., 303(106), 313 Griffith, R. H., 175,239 Griffiiths,J . H. E., 126(189), 133 Grimley, T B., 138(12), 235 Griscom, D. L., 81(18), 91(64), 128, 129 Gritscov, A . M . , 93(67, 70), 130 Griva, A. P., 106(118), 1 3 / Grubner, O., 285(64), 312 Gryaznov, 2 . V., 19(85), 20(108, i l l ) , 22 (123, 124), 70, 71, 72 Guenca, A . M., 303(112), 313 Guenin, M., 120(157), 132 Guerra, C. R., 185(187), 186, 23Y GUptd, N . M . , 54, 75 Gupta, R . K . , 303(98), 312 Guptd, S. K . , 20(116), 72 Guseinov, A . D., 38(188), 74 Guseva, I. V., 55(242), 75 Gutyrya, V . S., 19(94), 56(94), 71

H Haber, J., 107(123), 131 Haensel, V., 225(252), 228(252), 229(252), 241 Hagstrum, H. D., 143(33), 144(33), 146(33), 236

Hale, J . W . , 113, 132 Hall, W. K., 87(36), 91(36), 127(36), 129 Halliburton, L. H., 87(42), 129 Halpern, B., 107(124), I31 Hammer, H., 59(256), Hammonds, J., 230(275), 242 Happel, J.. 197, 240 Hara, N., 25(140), 26(140), 28(140), 29(140, 147), 42(209), 44(209), 62(140, 209). 64 (140). 67(209), 72, 73, 74 Hardt, P., 32(170), 73 Harkins, C. G., 107(121), 131

321

AUTHOR INDEX Harman, R. E., 20(116), 72 Hartzell, F. D., 13(31),69 Hasegawa, T., 31(164), 73 Hassan, S. M., 30(163), 73 Hasselmann, D., 35(183), 73 Hatada, K., 18(66), 70 Hausberger, A . L., 21 1(223), 241) Hawkins, R. J., 240 Hayakawa, M., 225(249), 228(249), 241 Hayata, S., 14(38,43), 15(43), 16(49, 50), 69 Hayden, P., 107(127), 125, 131 Hayek, K., 151(85), 152(85), 161(85), 162(85), 169(85), 237 Hayes, J. C . , 225(264), 228(264), 241 Haynes, H. W., 2(5), 18(5), 68 Haynes, W. P., 211(221, 222), 240 Hayward, D. O., 253(30), 311 Healey, F. H., 267(48), 311 Heegmann, W., 151, 152(85), 161(85), 162, 169(85), 237 Hegedus, L. L., 183, 228(175), 239 Heimbach, P., 32(167, 170), 35(167), 73 Heinz, R. E., 212(227), 240 Hemidy, J. F., 16(52), 70, 106(117), 131 Henderson, B., 89(47), 92(66), 100(98a), 129, 130, 131 Hendra, P. J . , 97(81), 130 Hennebert, P., 16(52), 70 Hensley, I18(153), 132 Herberhold, M., 285(66), 312 Herington, E. F. G., 191, 194, 240 Herman, R. G., 255(39, 40), 257, 258(39,41), 259(39, 40). 260139, 40,411, 261(40,41), 262(40, 41), 263(40), 264(40), 265(40), 266, 267(39), 271(39), 275(39, 56), 277 (56), 278(56), 279(56), 280(56), 283(39, 56), 284(56), 286(39), 287, 289(73), 291 (73), 292(56), 293, 294(77), 295, 296, 297 (56), 311, 312 Hermerschmidt, D., 91(62), 129 Hershman, A., 39(196), 40(196), 41(196), 42 (207), 43(196), 45(196), 74 Hertl, W., 303(112), 313 Herve, A., 87(41), 126(41, 188), 127(188), 129, I33 Higgens, R., 107(127), 125, 131 Hightower, J . W., 105(115), 131 Hill, S. G., 175(167), 239 Himmele, W., 39(195), 74

Hjortkjaer, T., 43(212), 44(212), 74 Hnatow, M. A,, 197,240 Hoang, C. Y., 29(156), 73 Hobbs, A . P., 196(210), 21 1(210), 240 Hobert, H., 156(122), 176(168), l81(168), 182 (168),238,239 Hockey, J . A,, 107(128), 132 Hoefs, E. V., 107(129), 132 Hofer, L. J. E., 185(179), 191(202), 239,240 Hohenschutz, H., 39(195), 74 Holah, D. G., 225(257), 226, 227(257), 228 (257), 241 Holland, H. B., 13(33), 69 Holloway,P. H., 184(178). 229(178),230(178), 23 I (178), 232,239 Holmes, J. M . , 267(49), 311 Holzwarth, N., 145(50b), 236 Homeier, E. H., 225(264), 228(264), 241 Hoodless, I. M., 225(257), 256(257), 227(257), 228(257), 241 Hopper, J . R., 240 Hortkjaer, J., 46(218), 75 Hosaka, H., 11 1(140), 132 Hosemann, R., 292(76), 293(76), 295(76), 312 Houghton, H., 300(84), 306(84), 307(84), 309 (84), 312 Howe, R. F., 18(72), 70, 87(36), 91(36), 125 (182), 127(36), 129,133 Huber, M., 151(70, 71, 80, 89), 152(70, 71), 236,237 Hucknall, D. J., 124(177a), 133 Hudson, J. B., 184(178), 229(178), 230(178), 23 I (178), 232, 239 Hiittig, G. F., 247(9), 298(9), 310 Hughes, A. E., 92(66), 130 Hughes, A. N., 225(257), 226(257), 227(257), 228(257), 241 Hurych, Z., 251(28), 252(28), 311 Huss, A,, 157(130), 163(148), 238 Hwang, J.-T., 98(88), 130

I Ibarbia, P. A,, 253(32), 289(32), 290(32), 291 (32), 311 Ichikawa, M., 253(33), 289(33, 71, 72), 290 (33), 291(33), 311,312 lizuka, T., 59(257), 76 Ikeda, S., 156(124), 238

322

AUTHOR INDEX

Ikeda, Y . , 14(40),6 9 Imai, H., 31(164), 73 Imanaka, T., 20(112), 26(112), 27(112), 64 (1 12). 71 Imelik, B., 19(104), 20(104), 30(159, 160), 62 (159, 160), 64(159, 160), 67(104), 71, 109 (133), 110(133), lll(133. 139), 132, 143 (28), 146(28), 235 Indovina, V., 94(73), 97(86), 100(98b), 109 (134a), 110(98b, 135), 111, 122(165), 130, 131, 132, 133 Inoue, T., 216, 217(234, 235), 241 Invernizzi, R., 36(186), 37(187), 62(186), 74 lone. K. G., 1 1 , 12(21, 22, 23), 13(16, 23), 16 (16, 17, 19, 21, 22, 23), 18(64,69), 19(99), 20(109, I I O ) , 53(23), 56(99), 64(21, 22, 23), 68, 69, 70, 71 Isagulyants, G. V., 24(132), 72, 225(261), 228 (261), 241 Isakov, Ya, I., 24(138, 139), 25(139), 38(189, 190, 191, 192, 193), 55(242), 72, 74, 75 Isakova, T. A., 19(77, 78), 70 Isakson, W. E., 213(233), 219(233), 240 Ishida, N., 20(112), 26(112), 27(112), 64(112), 71 Ishizuka, K., 151(76), 185,237 Ito, T., 122. 133 h e y , H. F.. 1 17( 155), 132 Iwamoto, M . , 107(122b), 131 Iwasawa, Y., 121(160c), 125(160c), 132 Iyengar, R. D., 80(1 I), 128 Iyer, R.M . , 54, 75 Izen, E. H., 83(20), 125

J Jackson, S. D., 185(184), 239 Jacobs, P. A., 2(2), 4(2), 5(2), 12(24), 13(27), 16(24), 18(2), 57(249a, 250,251,252,253), 58(250, 251, 252, 253, 25% 59, 60, 62 (250), 64(24), 65(255). 67(259, 261), 68 (250), 68, 69, 75, 76 James, W. H.. 32(169), 73 Jarrell, M . S., 40(200), 74 Jarvi, G. A., 188(194). 192(194), 197(194), 198 (1941, 201(194, 217), 211(194). 212(194): 229(194), 230(194), 232(194), 239, 240 Jenner, E. L., 28(142), 72 Jensen, V. W., 43(212), 44(212), 74

Jepsen, D. W., 138(6), 143(6,29a, 29b), 144(6), 146(6, 29a, 29b), 151(72), 161(72). 162 (72), 165(29a),235, 236 Jewur, S. S., l25( 177b), 133 Jiru, P., 124(175), 133, 285(64), 312 Jitsumatsu, T., 14(39),69 Johnson,S., 173(165), 175(165),176, 177(165), 178, 184, 225(171), 239 Johnstone, P. D., 87(44), 126(44), 129 Jolly, P. W., 32(167), 35(167), 73 Jona, F., 151(72). 161(72), 162(72), 236 Jmrgensen, J. C., 46(218), 75 Jorgensen, M . H., 250(23), 310 Joyner, R. W., 151(58, 91), 152(58), 157(91), 158(91), 161(58), 162(58), 236, 237

Kagel, R. O., 303(106, 1 lo), 313 Kalechits, I. V . , 225(258), 228(258), 241 Kaliaguine, S. L., 131 Kalyaguine, S. L., 87(39), 92(39), 93(39), 127 (39), 129 Kamble, V. S., 54(237), 75 Kamneva, A. I., 18(67), 70 Kanazirev, V., 19(98), 56(98), 71 Kanzansky, V. B., 87(39, 40). 92(39, 40), 93 (39,40), 125(40), 127(39,40), 129 Kapustin, M . A,, 19(90,91,92), 56(90,91,92), 71 Karavaev, M . M., 276(61), 311 Karhlamov, V. V . , 14(41, 42), 69 Karn, F. S., 190(199), 191(199,200,202), 240 Karra, I. S., 80(1 I), 128 Kasabava, N. A,, 249(14), 310 Kasai, P. H., 59,60,61(260), 67(260), 76 Kato, A,, 14(37, 38,43), 15(43), 16(49, 50), 69 Katzer, J. R., 47(222), 75,90(60), 91, I Z Y , 138 (10, I I ) , 143(39), 154(99), 157(130, 131), 159, 160(99), 163(131, 147), 165(99), 168 (99, 151), 186(99), 188(99, 196), 189(99, 196), 192(99, 147,205,206,207,208), 194 (205), 198(99, 147, 205, 206, 207, 208), 201,202(99, 147, 196,205, 206,207,208), 203(99, 147), 204(99, 147). 205(99), 206 (99), 207(99), 208(99), 210(99, 147). 21 1 (99), 212(99), 214(99), 229(1 I), 230(1 I). 231(11), 232(11),235,236,237,238,239, 308(118), 313

AUTHOR INDEX

Kaufherr, N.. 121(164), 13-1 Kawai, T., 229(272), 230(2’72),231(272), 232, 242 Kaye, G. W. C., 78(2), 128 Kaye, R. L., 35(183), 73 Kazansky, V . B., 24(132, 134, 135), 29(148, 149). 64(135), 72, 73, 81(17), 83(17, 19, 24b), 86(17), 87(17, 19, 37, 38), 90(24b), 93(67, 68, 69, 70), 98(38, 90, 91, 92, 93), 99(93), 102, 104(107), 105(107), 106(19, 106, 118), 107(90, 91, 120), 120( 160b), I 2 1 ( 160b, 160d). 123( 160d), 1 25( 160b,d, 183), 127(17, 19, 24b, 37, 38, 59), 128, 129, 130, 131, 132, 133 Kazusaka, A., 102, / 3 / Keii, T., 18(66), 61(264), 70, 76 Keim, W., 32(170), 73 Kelemen, S. R., 151(77, 86., 87, 88), 185, 225 (88), 237. 239 Kelley, R. E., 190(199), 191(199), 240 Kemball, C., 153(96), 154(96, 103), 155(96), 156(96, 103), 160(96), 165(96, 103), 174 (96), I85(96), 237 Kemp, J . C., 83(20), I28 Keulks, G. W., 107(129), 125, 132, 133 Kevan, L., 84(27), 85, 88(28), 127(88), 128 Khalif, V. A,, 97(85), 130 Kharatishvili, N. G., 19(89), 71 Kharlamov, V. V., 19(77, 78, 81, 82, 84, 86, 87, 88, 89, 90, 91, 92), 70, 71 Kherd, S. S., 197(212), 240 Khodakov, Yu. S., 19(83), 70 Khulbe, K. C., 157(136), 238 Kibblewhite, J. F. J., 81(14), 83(14), 86(35), 97(83), 99(35), 100(35), 101(14), 102(14), 127(14), 128, 129, 130 Kieffer, R., 301(86), 304(86), 306(86), 312 Kikuchi, T., 151(76), 185,237 King, D. L., 58(254a), 76 Kington, G . L., 267(49), 311 Kirkpatrick, W. J., 139(14), 140(14), 229(269), 230(269), 235,242 Kiryushkin, S. G., 38(188), 74 Kishi, K., 151(91), 156(124), 157(91, 139), 158 (91), 185(180), 237,238,239 Kiskinova, M . , 173(164), 175, 176(164), 177 (164), 178(164), 186, 198(193), 210, 239 Kitagawa, M., 218(239), 241 Kitayama, Y., 225(249), 228(249), 241

323

Kitching, S. C . , 156(121), 176(121), 178(121), 181(121), 184, 238 Kittel, C., 161(146), 238 Klein, D. L., 125, 133 Klier, K., 252(29), 254(36), 255(39, 40), 257 (38, 39, 40, 41, 43), 258(39, 41), 259(39, 40), 260(39, 40, 41), 261(40, 41), 262(40, 41), 263(40), 264(40), 265(40), 266(39,40), 267, 268, 269(43), 270(43), 271(39, 43), 274(43), 275(39, 56), 277, 278, 279, 280 (56), 281(56), 282(56), 283, 284(56), 285 (63, 64), 286(43, 63), 287(39), 289(73), 291(73, 75), 292, 293(41, 77), 295(77), 296, 297, 302(43,63), 303(43), 311,312 Klimenko, J. M., 225(262), 228(262), 241 Klueva, N. V., 18(69), 70 Knache, O., 139(16), 140(16), 235 Knappe, B., 176(168), 181(168), 182(168), 239 Knehr, H., 302(94), 312 Knight, C. B., 211(223), 240 Kobylinski, T. P., 255(39), 257(39), 258(39), 259(39), 266(39), 267(39), 271(39), 275 (39), 283(39), 287(39), 296(39), 311 Kochloefl, K., 267(46), 274(46), 311 Kodratoff, Y., 111(139), 132 Kolbel, H., 59(256), 76 Koidl, P., 89(49), 129 Kokes, R. J . , 302(90, 91). 312 Kollrack, R., 88(45b), 129 Kolodieva, Y. V., 20( 11 l), 71 Kolosov, A. K., 81(17), 83(17), 86, 87(17), 90 (59), 127(17, 59), 128, 129 Komarov, V. S., 18(70), 70 Kon, M . Ya., 102(106), 106(106), 131 Kondow, T., 229(272), 230(272), 231(272), 232 (272), 242 Kondrat’ev, D. A,, 255(260), 228(260), 241 Konoval’chikov, 0. D., 19(103), 20(103), 67 (103), 225(253), 227(253), 71, 241 Konovalchikov, 0. P., 19(102), 20(102), 67 (102), 71 Konvalinka, J . A ,, 288(69), 312 Koopman, P.. 17(59), 70 Koppova, A,, 225(250), 226,241 Kortiim, G., 302(94), 312 Kostelitz, O., 247(9), 298(9), 310 Kotani, K., 153(97), 237 Kotera, Y . , 249(17), 250(17), 310 Koutecky, J . , 292(76), 293(76), 295(76), 312

324

AUTHOR INDEX

Kovalenko, L. I., 24(132), 72 Kozhukhareva, V. N., 46(221), 75 Kozlov, G. A., 87(39), 92(39), 93(39), 127(39), 129 Krabetz, R., 220(242), 221, 222,223(242), 241 Krasnova, L. L., 46(220), 75 Krauss, H. L., 29(151, 152), 73 Krdvtsov, V. E., 150(51), 236 Krenzke, L. D., 125, 133 Kroner, W., 32(170), 73 Kruerke, U.,29, 33(174), 62(174), 64(174), 73 Kruglikov, V. Ya., 19(102, 103), 20(102, 103), 67(102, 103), 71, 225(253), 227(253), 241 Krupennikova, A. Yu., 20(112), 26(112), 27 (1 12), 64( I12), 71 Krylov, 0. V., 14(36, 37), 69,97(85), 130 Krymova, V. V., 22(124), 72 Krzywicki, A., 40(198), 41(204), 42(198), 74 Krzyzanowski, S., 97(83), 130 Ku, R., 151, 185,231(83b), 237 Kubaschewski, O., 140(20), 235 Kubo, T., 11(20), 13(20), 15(47), 16(20), 21, 22(122), 61(47), 63(47), 69, 72 Kubokawa, Y., 93, 94(71b, 71c, 71e), 103 (110b), 104(111b), 120(158, 159, 160a), 121(71b, 71c, 71e, 160a), 123(71b, 71c, 71e), 125(71b), 130,131, 132 Kubota, T., 15(46), 61(46), 63(46), 69 Kuhn, H. C . , 81(12), 128 Kumada, F., 15(46), 61(46), 63(46), 69 Kummer, J. T., 270(52), 311 Kung, H. H., 309,3/3 Kunimori, K., 229(272), 230(272), 231(272), 232(272), 242 Kunugi, T., 11(20), 13(20), 16(20), 21(121, 122), 22(122), 35(182), 69, 72, 73 Kunz, A. B., 125, 133 Kuraishi, M . , 250(21), 310 Kuriacose, J. C., 304(115), 313 Kuwada, T., 18(71), 70 Kuznetsov, L. D., 220(244), 222(244), 241 Kuznetsov, 0. I., 30(163), 38(188), 73, 74 Kuznetsov, P. N., 12(21), 16(21), 64(21), 69 Kuznicki, S. M . , 60, 61, 67(263), 76

L Laby, T. H., 78(2), I28 Lai, H., 151(90), 162,237

Lambert, R. M . , 151(52), 236 Lambertin, M . , 153(98), 237 Landau, M . V., 19(102, 103), 20(102, 103), 66, 67(102, 103), 71, 76, 225(253), 227, 241 Lang, W. H., 23(128), 56(243), 64(128), 68 (243), 72, 75 Lapidus, A . L., 24(138), 28(144, 145, 146), 38 (189, 190, 191, 192, 193), 55(242), 72, 73, 74, 75 Laramore, G. E., 143(40), 144(40), 146(40), 236 Lauer, P. R., 276(60), 311 Launay, J. P., 83(23), 125(23), 128 Lawson, J. D., 54(238), 75 Lawson, T., 81(13, 14), 83(14), 90(13), 101(13, 14), 102(14), 127(14), 128 Lebedev, Ya. S., 84(26), 128 Leclercq, G., 225(251), 228(251), 241 Lefebore, G., 24( 133), 25( 133), 28( 133), 72 Legg, K . O., 151(72), 161(72), 162(72), 236 Leibgolt, H., 250(25), 320 Leidheiser, H., Jr., 252(29), 311 Leland, T. W., 107(121), 131 Lenchthaler, C. H., 55(239), 56(239), 66(239), 75 Lender, Yu.V., 249(19), 310 Leonov, V. E., 276,311 Leutvein, K . , 127(190), 133 Leutwyler, S., 60, 67(262), 76 Levine, J. D., 116, 118(151), 132 Levy, R. B., 140(18), 197(18), 235 Lewis, J., 124(174), 133 Lichtman, V . A,, 143(37), 145(37), 236 Lidow, D. B., 185(189), 239 Liebsch, A., 145(50b), 236 Lin, T. Y., 225(259), 228(259), 241 Lindey, R. V., 28(142), 72 Lipdri, N. O., 143(40), 144(40), 146(40), 236 Lipatkina, N. I., 83(24b), 90(24b), 98(93), 99 (93), 102(93, 107). 104(107), 327(24b), 128, 130, 131 Llyodlangston, J., 20(106), 71 Lofthouse, M . G., 113(146), 114(146), 117 (146). 132 Lopez Agudo, A , , 23(126), 72 Lormand, C., 245(5), 248(5), 276(5), 310 Louis, C., 102(108c), 131 Low, M . J. D., 302(89), 312 Ludwig, G. W., 87(44), 126(44), 129

AUTHOR INDEX

Lukas, K., 276(59), 311 Lukas, J., 31(166), 62(166), 73 Lunsford, J. H., 16(54, 55), 17, 18(72), 52, 53 (234), 54(234), 55(234), 56, 59(257), 63 (2341, 67(234), 70, 75, 76, 78, 79, 81, 84 (30a), 85(16, 311, 86, 88(1), 90(15, 30a), 91(30a, 31), 100(97), lOl(15, 101), 102, 103, 104, 105(114), 122, 125, 127(16, 30a, 31), 128, 129, 131, 133 Luss, D., 151(81), 152(81), 163(81), 163, 237 Luz, Z., 85(32), 126(32), 129

M McAteer, J. C., 84(29), 125(29), 129 McCabe, R. W., 183,228(175), 239 McCann, W., 20(114, Il5),21, 71, 72 McCarroll, J. J., 143(35, 36a, 36b). 145(35), 146(35, 36b), 149, 152, 236,237 McCarty, J. G., 154(114, 115), 155, 165, 166 (114, 115), 167(114), 169(152, 156), 170, 170(56), 186(114, 115), 188(114, 195), 192 (195), 198(195), 200(195), 211(195), 212 (114, 195), 213(195, 233), 214(195), 219 (195, 233), 237,238,239.240 McCullough, J . P., 55(239), 56(239), 66(239), 75 McKee, C. S., 151(58), 152(58), 161(58), 162 (58), 236 Mackenzie, J. R., 97(81), 130 McRae, E. G., 143(45), 144(45), 236 Madey, T. E., 185(190), 186(190),239 Madix, R. J., 173(165), 175(165), 176, 177 (165), 178, 184, 185(182), 225(171), 239, 298(81, 82, 83), 300(81, 82, 83), 303, 312 Madon, R. J., 136, 190(2), 235 Maffeo, B., 87(41), 126(41), 129 Mahishima, S., 124(171), 133 Mahoney, F., 18(65), 64(65), 70 Mailhe, A,, 245(4), 310 Makambo, P., 58(254), 76 Maksimov, N. G., 11(16), 12(21), 13(16), 16 (16, 21), 64(21), 68, 69 Malashevich, L. N., 18(70), 70 Malevich, V. I., 19(102), 20(102), 67(102), 7 / , 225(253), 227(253), 241 Mal'shukov, A. G., 150(51), 236 Mal'tser, V. V., 28(144, 145, 146), 39(193), 72, 73, 74

325

Malyshev, A. G . ,87(43), 127(43), 129 Mamaev, 0. G., 46(219), 75 Mann, R. S., 157(136), 238 Manogue, W. H., 154(99), 157(131), 160(99), 163(131, 147), 165(99), 168(99), 186(99), 188(99), 189(99), 192(99, 147, 205), 206, 207,208), 198(99,147,205,206,207,208), 201(99, 147, 205, 206, 207, 208), 202(99, 147, 205, 206, 207, 208), 203(99, 147), 204(99, 147), 205(99), 206(99), 207(99), 208(99), 210(99, 147), 21 1(99), 212(99), 214(99), 237,238 Mantovani, E., 47(226), 49(226), 63(226), 65 (226), 67(226), 75 Marcus, P. M . , 138(6), 143(6, 29a, 29b), 144 (6). 146(6, 29a, 29b), 151(72), 161(72), 162(72), 165(29a), 235, 236 Marechal, J., 225(247), 22812471, 241 Margot, E., 151(67), 169(153), 236,238 Marion, J., 79(4), 125(4), 128 Mark, P., 116, 118(151), I32 Markina, M. I., 267(50), 311 Marko, L., 47(223), 49(223), 75 Marks, T. J., 285(65), 312 Marosi, L., 249(1 l), 310 Mars. P., 17, 70, 124, 133 Marsh, J. D. F., 175(166), 239 Marshakova, N., 138(3), 235 Martens, R., 94, 130 Martin, G., 143(28), 146(28), 235 Martin, G. A,, 154, 165(1 lo), 173(1 lo), 180 ( I lo), 225(1 lo), 237 Maschenko, A. I., 90(51), 129 Maslyanskii, G. N., 225(262), 228(262), 241 Massardier, J., 19(104), 20(104), 67(104), 71 Masson, A-, 143(47), 145(47), 146(47), 148 (47), 150, 236 Mathe, T., 20(107), 71 Mathews, J. T., 52(233), 53(233), 75 Mathieu, M., 30(159), 62(159), 64(159), 73 Mathieu, M. V., 16(53), 70 Matsumura, Y . , 94(71d), 123(71d), 130 Matsuura, I., 124(173a), 133 Maurel, R., 225(251), 228,241 Maxted,E. B., 136, 137,225(268), 227,229(1), 235,242 Maxwell, 1. E., 8(11), 33, 35(179), 62(175, 177, 178), 63(177), 64(175, 177, 178, 179), 65 (179). 68, 73

326

AUTHOR INDEX

Mazzanti, G.. 275(57). 279(57), 311 Meguro, K . , 110(136), lll(136, 140). 132 Mchta. S . . 255(40), 257(40), 259(40), 260(40). 26 1(40), 262, 263(40), 264(40), 265(40), 266(40), 31 1 Meiscl, S . L.. 55(239), 56(239), 66(239). 75 Meister, K . H., I5I(X5), 152(85), 161(8S), 162 ( 8 5 ) , 169(85), 237 Menon, P. G., 255(266), 227(266), 228(266), 229, 242 Mericaudeau, P., 90(52), 96(65a), 98(52, 89), 99(89), 102(108b), 129, 130, 131 Metzger, J., 79(4), 125(4), 128 Miale, J . N., 23, 62(127), 72 Mikhaleva, I . M., 22(123), 72 Mikheikin, I. D., 90(51), 129 Mikovsky, R. J., 8(10), 23(128, 129, 130), 64 (128, 129, 130), 68, 72, 302(95), 312 Milliams, D. E., 177(173), 178(173), 239 Mills, K . C., 140(17), 235 Minachev, H., 19(98), 56(98), 71 Minachev, Kh. M., 14, 19(77, 78, 79, 81, 82, 83, 84, 86, 87, 88, 89, 90, 91, 92, 96), 20 (113), 24(138, 139), 25(139), 28(144, 145, 146), 38(189, 190, 191, 192, 193), 52(242), 69, 70, 71, 74, 75, 255(260,261), 228(260, 261). 241 Mine, R. S . , 19(99). 56(99), 71 Mirra, C. Z . , 122(166b), 133 Mishra, S. P., 90(58), 12Y Mitchell, P. C. H., 124(176), 133 Mitsche, R. T., 225(264), 228(264), 241 Mizokoshi, Y., 103(110b), 104(111b), 131 Mochida, I., 14, 15(43), 16,69 Modine, F. A , , 83(20), 128 Moeller, A. D., 201(218), 240 Moison, J. M . , 151(59), 169(155), 236, 238 Mollan, P. A. F., 18(68), 70 Mollwo, E., 117(156), 132 Monnier, J. R., 107(129), 132 Montelatici, S . , 19(100), 56(100), 71 Montgolfier, Ph., 102(108b), 131 Montgomery, P. D., 40(197), 42(206), 74 Moody, A. F., 107(130), 132 Moriai, N., 122(166c), 133 Morita, S., 216. 217(234, 235), 241 Moro, G., 92(65b), 129 Morterrd, C . , 122(166d), 133

Morton, J. R., 79(7, X), 80(8), 81(7, 8). 85 (7. 8). 127(7, 8). 128 Moya, F., 151(64), 236 Mozzancga, H.. 58(254c), 76 Mudge, W . A ,. 142(21), 235 MUnUerd, G . , 99(95, 96), 107(96). 130, 131, 302(93), 3/17 Murchison. C . B., 197(213), 240 Murdick. D. A., 197(213), 240 Muth, A,, 301(86), 304(86), 306(86), 312 Muzykantov, V. S.,107(125), 131

N Naccache, C., 2(7a), 68, 81(15), 83(25), 90(15, 52), 94(52, 72), 101, 102(15, 102, 103, 108b), 104(103, 112), lll(133, 138, 139), 128, 129, 130, 131, 132 Nagarjunan, T. S . , 304,313 Nagata, J., 29(147), 73 Najbar, M., 94(76). 130 Namba, S . , 19(99), 56(99), 71, 225(254), 227 (254), ,741 Narayanan, S.,20( I 14, I IS), 2 I , 71, 72 Naro, P. A., 225(267), 241 Naskidashvili, N . N., 19(85), 70 Natta, G.. 248. 275, 279, 284(10), 310, 311 Neeley, V . I . , 132 Nefedov, B. K., 35(181), 40, 41(203, 205). 42, 43(203), 44(203), 45,46,62(181), 65(203), 73, 74, 75 Nefedov, V . I . , 138(3), 235 Neff, L. D.. 156(121), 176(121), 178(121), 181 (121), 184, 238 Neinska, Y . , 19(98), 56(98), 71 Nelson, R . L., 109, 113, 132 Nesterov, V. K . , 19(83), 70 Newling, W. B. S . , 175(166), 239 Ng, C. F., 154, 160(111), 165(111). 173(111), 174, 175(111), 176(111). 180(111), 184, 23 7 Nguyen,T. T. A., 143(41,44), 144(41,44), 146 (41, 44), 159(41), 236 Nieuwenhuijse, B., 17(59), 70 Nijs, H. H., 57(250,25 I , 253), SS(250.253). 62 (250), 68(250), 75, 76 Nikiska, V . V., 98(90, 91, 92), 102(106), 106 (106, 118), 107(90.91, 120). 130, 13/

327

AUTHOR INDEX

Nistor,S. V., 80, 126(9,10), 127(191), 128,133 Niwa, M . , 59(257), 76 Noda, S., 94(71d), 123(71d), 130 Nordhausen, L. J., 229(273). 230(273), 242 Norgett, M. J., 89(48), 129 Notari, B., 267(47), 274(47), 296(47), 311 Novozhilov, A. I., 87(43), 127(43), 129 Novruzov, T. A., 19(96), 56(96), 71 Nyholm, R. S., 124(174), 133

0 Oba, M . , 249(17), 250(17), 310 Oberkirch, W., 32(170), 73 Occhiuzzi, M., 97(86), 122(165), 130, 133 Ogasawara, S.,121(160c), 125(160c), 132 Ogawa, K., 249(17), 250(17), 310 Ogneva, T. P., 303(96), 312 Okamoto, Y., 18(71), 20(112), 26(112), 27 (1 12), 64(112), 70, 71 Oliphant, J. L., 154(112), 155, 157(112), 160, 161(112), 165(112), 166(112), 167(112), 168(112), 173, 175(112), 186(112), 230 (1 12), 237 Oliver, R. B., 250(20), 310 Olson,D. H., 8(10), 23(129,130),64(129,130), 68, 72 O’Mara, W. C., 83(24a), 127(24a), 128 O,Neill, C . S . , 177(174), 178(174), 239 O’Neill, P., 90(57), 129 Onishi, T., 229(272), 230(272), 23 l(272), 232 (272), 242, 301(88), 302(88), 303(88, 107), 307(117), 312,313 Ono, Y., 18(66), 61(264), 70, 76 Orikasa, Y., 42(209), 43(210), 44(209), 62 (209, 210), 65(210), 67(209, 210), 74 Orlova, L. B., 20(1 lo), 71 Orton, J . W., 126(189), 133 Ostrovskii, V. E., 303,312 Ota, K., 153(97), 237 Oudar, J., 143(26,48a, 48b), 145(26), 146(26), 148, 151(55, 56, 57, 60,61, 62.63, 65, 66, 67, 68, 70, 74, 75, 84, 89), 152(63, 65, 70, 74), 157(48a), 158(26), 159(26), 161(55, 56, 65, 75, 84), 162(75, 84), 165(26), 166 (26), 169(57, 84, 153), 171(57), 172(57), 186(57), 223(48b), 230(48b), 235, 236, 237,238

Ozaki, A., 107(126), 131

P Paetow, H., 13, 69 Palladino, N., 47(226), 49(226), 63(226), 65 (226), 67(226), 75 Panchenkov, G. M., 30(163), 38(188), 73, 74 Pannell, R. B., 154(112), 155(112), 157(112), l60(112), 16l(l12, 141), 165(112), 166 (112), 167(112), 168(112), 173(112. 141, 160, 162), 175, 176(141, 160, 162), 177 (141, 160, 162), 178, 179(141, 162), 180 (160, 162), 181, 183(162, 176, 177), 185 (160), 186(112), 230(112, 276, 277), 237, 238,239,242 Pannetier, G . ,40(198), 41(204), 42(198), 74 Panov, G. I., 107(125), 131 Pantages, P., 196(210), 21 1(210), 240 Papp, H., 303(100), 312 Paranosenkov, V. P., 20(108, 11I), 71 Parkes, D. A., lOl(99, loo), 102(100), 103 (IOO), 105(100), 131 Parris, G. E., 257(42, 43), 267, 268, 269(43), 270(43), 27 1(43), 274(43), 277(42), 286 (43), 288(43), 295, 302(43), 303(43), 311 Pasquet, D., 111(141), I32 Pasquon, I., 47(225), 48(225), 75, 124(172, 175), 133,275(57), 276(58), 279(57), 311 Patart, M . , 245,310 Pathak, A. P., 89(48), 129 Patzelova, V., 29(154), 73 Paulik, F. E., 39(196), 40(196), 41(196), 43 (196), 45(196), 74 Pauling, L., 138(8), 235 Paulitschke, W., 161(142), 162(142), 185(186), 223(142), 230(142), 238,239 Pavlova, N. Z., 220(244), 222(244), 241 Pearce, J. R., 17, 70 Pecora, L. M . , 156(125), 185(125), 238 Perdereau. M., 143(42), 144(42), 146(42), 148 (42), 151(84), 161(84), 162(84), 169(84), 192, 198(209), 200, 201, 210(209). 236, 237,240 Penchev, V., 19(98), 56(98), 71 Pendry, J . B., 143(31a), 146(31a), 235 Pennline, H. W., 21 1(222), 240 Pepe, F., 261(45), 311

328

AUTHOR INDEX

Peralta, L , 151(75, 80), 161(75), 162(75), 237 Perdereau, M . , 143(26), 145(26), 146(26), 148, 158(26). 165(26), 166(26), 167(36),235 Perry,J. H., 165(150),238 Pershin, A. N., 125(183), 133 Peters, C . , 220(242), 221, 222, 223(242), 241 Peterson, E. M.. 169(154), 220(154), 223(154), 224( 154), 238 Petrino, P., 151(64),236 Petrishcheva, G. S., 249(19). 276(61), 310,311 Petro, J . , 20(107), 71 Pfefferle, W. C . , 225(263), 228(263), 241 Picardi, R., 87(41), 126(41), 129 Pichat, P.. 30(159, 160), 62(159, 160), 64(159, 160), 73 Pichler, H., 190(198), 216(198), 217(198), 240 Pickert, P. E., 19(101), 20(101), 67(101), 71 Pieters, W. J., 229(271), 242 Piken, A. G., 192(203), 194(203),212(225), 240 Pines, H., 225(246, 247), 228,241 Pink, R. C . , 18(68), 70, I 1 1(I39), I32 Pino, P., 275(57), 279(57), 311 Pirozhenko, L. F., 249(19), 310 Pismennaya, A. V., 18(70), 70 Piterskii, L. N., 18(67), 70 Pitkethly, R. C . , 143(35, 36a, 36b), 145(35), 146(35, 36b), 149(35, 36b), 152(35), 236 Platteeuw, J. C., 21(117), 72 Pliskin, W. A,, 302(89), 312 Plummer, E. W., 145(50a, 50b), 155(50a), 161, 236 Pohl, J., 292(76), 293(76), 295(76), 312 Pollitzer, E. L., 225(264), 228(264), 241 Poncelet, G., 57(252), 58(252), 76 Ponec, V., 289(71), 312 Popov, 1. G., 249(19), 310 Popovskii. V. V., 18(64), 70 Post, M . F. M . , 10(15a), 68 Postl, W . S., 225(246,247), 228(246, 247). 241 Pott, G . T., 113, 132 Pour, V., 154(109), 155(109), 165(109), 237 Poustma, M . L., 29(158), 73,253(32), 289(32), 290(32), 291(32), 311 Prahaud, H . , 94, 130 Prasad, J.. 255(266), 227(266), 228(266), 229, 242 Primet, M., 16(53), 19(104), 20(104, 105), 51 (105), 67(104, IOS), 70, 71 Pritchard, J., 177(173), 178(173), 239, 303(98, 100, I O l ) , 312

Przhevalskaya, L. K., 29(149, 150), 73 Puzitskii, K . V., 24(134, 135, 136, 137). 64 (1 35), 72

Q Quets, J. M., 303(103), 313 Quiggle, D., 245(6), 246(6), 300(6), 310

R Rabina, P. D., 220(244), 222(244), 241 Rabo, J. A., 2(1), 18(l), 19(101), 20(101), 29 (158), 67(101), 68, 71, 73, 253(32), 289 (32), 290(32), 291(32), 311 Radchenko, E. D., 19(103), 20(103), 67(103), 71 Radtsig, V . A., 91(63), 129 Raineau, A , 245(2), 310 Rakowski-DuBois, M. C . , 197(215), 240 Ralek, M . , 292(76), 293(76), 295(76), 312 Ramaswany, A. V . , 225(265), 228(265), 241 Randhava, S. S., 185(183), 239 Rappe, A . K., 165(149), 238 Rase, H . F . , 54(238), 75 Regner, A., 154(109), 155(109), 165(109), 237 Rehmat, A ., 185(183), 239 Reik, H. G., 89(49), 129 Reimlinger, H., 33(174), 62(174), 64(174), 73 Relek, M . , 55(241), 75 Rendulic, K . D., 173(163), 175, 176(163), 239 Pepina, V. V., 20(109), 71 Rewick, R. T., 176(169, 170), 180(169, 170), 181(169, 170). 181(170), 182, 239 Reynolds, R . W., 96(78), 126(78), 1311 Rhein, R. A,, 38, 74 Rhodin, T. N., 138, 143(5, 34), 144(34), 146 (34), 151(5), 185, 235, 236, 239, 251(28), 252(28), 311 Richardson, J. T., 142(24), 154(107), 165(107), 235,237 Riekert, L., 13, 69 Riesz, C . H., 229(269), 230(269), 242 Risch, A. P., 253(32), 289(32), 290(32), 291 (32), 311 Rius, G., 87(41), 126(41), 129 Roberts, M . W., 151(58), 153(96), 154(96, 103), 155(96), 156(96, 103, 123), 157(91, 139), 158(91), 160(96), 161(58), 162(58),

AUTHOR INDEX 165(96, 103), 174(96), 185(96, 180), 236, 237,238,239 Robinson, K. K . , 42(207), 74 Rochester, C. H., 156(120), 176(120), 177 (120), 180(120), 181, 185(120), 238 Roen, S., 154(106b), 155(106b), 165(106b), 166(106b), 237 Rollmann, L. D . , 57(266), 63, 64(267), 76 Romannikov, V. N., 18(69), 20(109, IIO), 70, 71 Rosenqvist,T., 139, 140(13), 141(13), 142(13), 161(13), 165, 166(13), 169(13), 235 Rosolovskaya, E. N., 19(80), 70 Ross, J. R. H., 156(123), 157(133), 238 Rostrup-Nielsen, J . R., 154(106b), 155, 165 (106a, 106b), 192, 198(209), 200, 201, 210(209), 216, 217(238), 218, 219(237b), 229(270), 230(238, 270), 240,241,242 Rosynek, M . P., 157( 134), 238 Roth, J. F., 39(196), 40,41(196), 42(196,207), 45(196), 74 Rozengart, M. I., 24(132), 72 Rozentuller, B. V., 97(85), 130 Rubene, N. A , , 303(304), 304(104), 307(104), 313 Rubinov, A. Z., 225(262), 228(262), 241 Rubio, 0. J., 96(79), 126(79), 130 Ruckenstein, E., 125(177b), f 3 3 Rudajevova, A , , 154(l09), 155(l09), 1 6 3l09), 23 7 Rudakova, L. N., 38(189, 190, 191, 192, 193), 74 Rudham, R., 2(7), 13(34), 14(34, 45), 18(7, 65). 64(65), 68,69, 70 Riishede, K., 250(23), 310 RBiiEka, V., 225(250), 226(250), 241 Ryashentseva, M . A , , 20(113), 71

S Sabatier, P., 245(3, 4). 310 Sachtler, W. M., 124(170), I33 Saito, V . , 124(171), 133 Saito, Y., 303(109), 313 Sakai, T., 35( I82), 73 Saleh, J . M., 153(96), 154, 155(96), 156, 160 (96), 163(126), 165(96, 103), 174, 185(96), 237,238 Salmon, G. A , , 90(57), 129 Samoilovich, M. 1.. 87(43), 1?7(43), 129

329

Sancier, K. M . , 213(233), 219(233), 240 Sander, W., 87(45a), 126(45a), 127(45a), 129 Sanders, M . K., 13(34), 14(34, 45), 69 Santier, C., 87(41), 126(41), 129 Sapozhnikov, V. B., 83(19), 87(19), 90(19), 105(116), 106(19), 127(19), 128, 131 Sargint, G . A , , 143(46), 145(46), 157(46), 176 (46), 177(46), 236 Sarichev, M . E., 90(53), 129 Sasaki, Y., 35(182), 73 Sastri, M . V . C . , 304(115), 313 Savchenko, B. M., 19(83), 70 Schaefer, H., 220(243). 241 Schechter, S., 213(233), 240 Schehi, R. R., 21 l(221, 222), 240 Schiavello, M., 261(45), 311 Schipperijn, A . J., 31(166), 62(166), 73 Schirmer, 0. F., 89, 129 Schlick, S., 84(27), 85, 88(28), 127(28), 128 Schmidt, H., 29(152), 73 Schmidt, L. D . , 151(81), 152(81), 163(81), 163, 237 Schnabel, K. Kh., 14(37), 69 Schneider, J., 127(190), 133 Schoenberg, A , , 85(32), 126(32), 129 Scholten, J. F. F., 288(69), 312 Schoofs, R. J., 229(273), 230(273), 242 Schoonheydt, R. A., 12(24), 16(24), 59(259), 64(24), 67(259), 69, 76 Schroder, W., 276(60), 311 Schroder, W., 249(1 l), 310 Schuit, G. C. A., 47(222), 75, 90(60), 91(60), 129 Schultz, J. F., 191, 240 Schultz, R. G., 40(197), 42(206), 74 Schumacher, E., 60, 67(262), 76 Schwaha, A. K., 151(52), 236 Schwarzmann, M., 249(1 I), 310 Schwarz, J. A,, 151(79), 185,237, 239 Scurrell, M. S., 32(171), 40, 42(208), 43(201, 211), 44, 45(201, 208, 211), 46(208), 47 (224), 65(21 I , 213), 73, 74, 75 Sedor, L., 225(257), 226(257), 227(257), 228 (257), 241 Seff, K., 6, 7(9), 68 Segall, B., 87(44), 126(44), 129 Segall, R. L., 114(150), 119(150), 132 Seidl, G., 151, 162, 237 Seiyama, T., 14(38, 39,43), 16(49, 50), 69, 107 (122b), 131

330

AUTHOR INDEX

Sekino, T., 122(166c), 133 Seleznev, V. A,, 14(44), 18(67), 69, 70 Selwood, P. W., 154, 155(104), 156(104), 160 (104), 165(104), 174, 237 Semenova, T. A., 267(50), 311 Senderens, J. B., 245(3), 310 Sergeeva, N. S., 40(199,203), 41(203,205), 42 (203, 205), 43(203), 44(203), 46(220), 65 (203), 74 Sergeywa, N . S., 35(181), 62(181), 73 Sergushin, N. P., 138(3), 235 Shakhnovskaya, 0. L., 19(80), 70 Shamaiengar, M., 225(246), 228(246), 241 Shang, L. K., 220(244), 222(244), 241 Shang, W. W., 107(121), 131 Shapiro, G . P., 267(50), 311 Shatynski,S. R., 140(19), 141(19), 169(19),235 Shaw, H., 136(2), 190(2), 235 Shelef, M., 192(203, 204), 193(204), 194(204), 212(225), 240 Shelimov, B. N., 87(39), 90(55), 92(39), 93(39, 68, 691, 98(72, 90,91, 92), 102(106), 106 (106, 118, 119), 107(90, 91, 120), 125 (183), 127(39), 129,130,131,133 Sherwin, M., 271(53), 311 Shestakova, N. A., 11(16), 13(16), 16(16), 68 Shiba, T., 304(114), 306(114), 313 Shields, L., 84(27), 127(193), 128, 133 Shikakura, K., 289(72), 312 Shimazaki, Y., 29(147), 73 Shimizu, T., 83(22), 128 Shimomura, K., 249(17), 250(17), 310 Shiotani, M., 92, 129 Shipman, G. F., 8(10), 23(130), 64(130),68, 72 Shishkov, D. S., 249(14), 310 Shliomenzon, N. L., 19(97), 56(97), 71 Shoemaker, C. F., 156(119), 238 Shomaker, V., 19(101), 20(101), 67(101), 71 Shostakovskii, M. F., 45(216), 74 Shpiro, E. S., 28(145), 72 Shuets, V . A,, 29(148, 149, 150), 73 Shuklov, A. D., 93(69), 130 Shultz, J. F., 190(199), 191(199, 200). 240 Shulman, J. H., 112(143), 132 Shutkina, E. M . , 40(202, 203), 41(203), 42 (203), 43(203), 44(203), 65(203), 74 Shvets, V. A,, 81(17), 83(17, 19, 24b, 37, 38), 86(17), 87(17, 19, 37, 38), 90(19, 24b, 38, 53, 54, 59), 93(67, 70), 98(38, 93), 99 (93), 102(38, 93, 106, 107), 104(107), 105(116),

106(19, 106, 118), 120(160b), 121(160b, 160d), 123(160d), 125(160b,d), 127(17, 19,24b, 38, 59), 128,129, 130, 131, 132 Sibley, W. A,, 89(46), 129 Sickafus, E. N., 143(43), 144(43), 157(43), 236 Sillen, L. G., 213(230), 240 Silvestri, A. J., 23(128, 129, 131), 56(243), 64 (128, 129, 131), 68(243), 72, 75 Simmons, G. W., 255(39,40), 257(39,40j, 258 (39), 259(39, 40), 260(40), 261(40), 262 (40), 263(40), 264(40), 265(40), 266(39, 40), 267(39), 271(39), 275(39, 56), 277 (56), 283(39, 56), 284(56), 287(39), 257 (41), 258(41), 260(41), 261(41), 262(41), 275, 292(56), 293(41, 77), 294(77), 295 (77), 296(39, 56, 77), 297(56), 311,312 Simon, A., 303(108), 313 Sinfelt, J. H., 54(235, 236), 75 Sivasanker, S., 225(265), 228(265), 241 Slaugh, L. H., 197(211), 240 Smir, S. A,, 107(120), 131 Smirov, B. V., /9(102, 103), 20(102, 103), 67 (102, 103), 71,225(253), 227(253), 241 Smirnov, V. S., 22(124), 72 Smith, A. W., 303(103), 313 Smith, J . V., 8, 68 Smith, R. L., 23(131), 64(131), 72, 225(267), 242 Snowdon, F.F., 250(24), 310 Sokolskii, D. V., 19(97), 56(97), 71 Soma, K., 35(182), 73 Soma, Y., 303(107), 313 Somorjai, G. A., 152(93), 228,237 Sorenson, W. L., 230(278), 242 Soria, J., 99(95, 96), 107(96), 130, 131 Sorokina, A. K., 46(215), 74 Soukup, J., 225(250), 226(250), 241 Southwich, C. A., Jr., 245(7), 246(7), 310 Spencer, N. D., 151(52), 236 Spiridonov, K. N., 97(85), I30 Srinivasan, S. R., 169(154), 185(186), 220 (154), 223(154), 224(154), 238,239 Stach, H., 29(151), 73 Stapelbroek, M., 91(64), 129 Steffgen, F. W., 196(210), 211(210), 240 Steijns, M., 17(57, 58, 59,60), 70 Stein, K. C., 185(179), 239 Steinbrunn, A., 157(137, 138). 158(137, 138), 238 Steinriicke, E., 32(170), 73

AUTHOR INDEX

Stencel, J. M., 212(227), 240 Sterba, M. J., 225(252), 228(252), 229,241 Stiles, A. B., 249(18), 310 Stockwell, A,, 2(7), 18(7), 68 Stoicescu, G., 80(9), 126(9), 128 Stone,F. S., 113(146, 147), 114(146), 117(146, 147, 154), 120(154), 123(168), 132, 133, 261(44), 311 Stoneham, A. M., 89(48), 129 Stouthamer, B., 21(117), 72 Stowell, D. E., 173(159), 176(159), 177(159), 180(159), 185(159),239 Strakey, J. P., 21 1(222), 240 Strey, F. L., 157(134), 238 Sugier, A , , 296(78, 79, 80), 312 Summers, J. V., 18(65), 64(65), 70 Surin, S. A,, 87(39), 92(39), 93(39, 68, 69), 98 (92), 127(39), 129, 130 Surratt, G. T., 125, 133 Suss, J. T., 85(32), 126(32), 129 Suzuki, K., 61(264), 76 Svejda, P. P., 91(62), 129 Sykes, K. W., 177(173), 178(173), 239 Symons, M. C. R.,84(27), 85, 90(58), 98, 99 (34), 104(113), 127(193), 128, 129, 130, 131, I33 Szapiro, S., 85(32), 126(32), 129

T Tabourier, P., 11 1(141), 132 Tagiev, D. B., 19(81, 82, 84), 70 Tagiyev, D. B., 14(41, 42), 19(86, 87), 69, 71 Takahashi, N., 42(209), 43(210), 44(209), 62 (209, 210), 65(210), 67(209, 210), 74 Take, J., 303(109), 313 Takeuchi, A,, 308(118), 313 Tamaru, K., 229(272), 230(272), 231(272), 232 (272), 242, 301(88), 302(88), 303(88, 107), 307,312,313 Tamura, N., 125(181), 133 Tanahashi, I., 94(71b, 71c), 71e), 121(71b, 71e), 123(71b, 71c, 71e), 125(71b), 130 Tanaka, K., 32(170), 73 Tarama, K., 125(181), 133 Tasker, P. W., 108, 132 Tauber, G., 161(142), 162(142), 223(142), 230 (142), 238 Tauster, S., 220(241), 221, 222(241), 241

331

Taylor, H. S., 302(95), 312 Taylor, P. C., 81(18), 128 Taylor, P. S., 245(7), 246(7), 310 Teichner, S. J., 92(65a), 129 Temkin, M. I . , 220(240), 222(240), 241 Tench, A. J., 81(13, 14), 83, 84(29), 86(35), 89 (50), 90(13, 35, 52), 94(75), 97(82, 83), 98 (52), lOO(35, 50), lOl(13, 14), 102(14), 104(112, 113), 106(117), 109, 113, 114 (148, 149, 150), 115(148), 116(149), 119 (150), 121(164), 122, 125(29), 127(14), 128,129, 130, 131, 132, 133 Teranishi, S., 20(112), 26(112), 27(112), 64 (112), 71, 125(181), 133 Terekhov, V. A., 138(3), 235 Terletskikh, A., 11(19), 16(19), 69 Terrell, R. J., 156(120), 176(120), 177(120), 180(120), 181, 185(120), 238 Terzaghi, G., 47(225), 48(225), 75 Thapliyal, H. V., 143(38), 145(38), 236 Thomas, C. L., 142(22,25), 235 Thomas, W. J., 20(106), 71 Thomson, S. J., 185(184), 239 Thurston, E. F. W., 303(102), 313 Tohver, H. T., 89(47),96(79), 126(79), 129,130 Tokuda, T., 122(166c), 133 Tominaga, H., 11(20), 13(20), 15(46, 47), 16 (20, 51), 17(51), 21(121, 122), 22(122), 35 (182), 61(46, 47), 63(46, 47), 64(51), 69, 70, 73 Tomita, T., 218(239), 241 Tong, S. Y., 143(30), 144(30), 146(30), 235 Tonner, B., 145(50b), 236 Topchieva, K. V., 19(80), 70 Tracy, J. C., 303(97),’312 Trapnell, B. M. W., 253(30), 311 Treibmann, D., 303(108), 313 Trifiro, F., 124, 133 Trindle, C., 98(88), 130 Tsai, J., 157, 163(13I), 238 Tsinober, L. I., 87(43), 127(43), 129 Tsitovskaya, I. L., 14(35, 36, 37, 44), 69 Tsitsishvili, G. V., 19(85), 20(11 I), 70, 71 Tsuchiya, J., 16(51), 17(51), 64(51), 70 Tsuchiya, S., 304(114), 306(114), 313 Tsuruya, S., 18(71), 70 Tsybina, E. N., 249(19), 276(61), 310, 311 Tungler, A., 20(107), 71 Turkevitch, J., 19(99), 56(99), 71, 80(11), 128 Tvaruzkova, Z., 29(154), 73

332

AUTHOR INDEX

U Uchida, H., 31(164), 73, 249(17), 250(17), 310 Ueno, A,, 301(88), 302(88), 303(88), 307, 312, 313 Ugai, Y. A., 138(3), 235 Uken, A. H., 173(161a, 161b), 239 Unruh, W. P., 96(77), I26(77), 130 Upton, T. H., 165(149), 238 Ursu, I., 80(10), 126(10), 128 Ushida, Y . , 25(140), 26(140), 28(140), 29 (140), 62(140), 64(140), 72 Uytterhoeven, J. B., 2(4), 12(24), 13(27), 16 (24), 18(24), 33(176), 57(250, 251, 253), 58(250, 251, 253, 255), 59(259), 60(261), 62(176, 250), 64(24, 176), 65(255), 67 (259, 261), 68(250), 68,69, 73, 75, 76

V Vahrenkamp, H., 143(37), 145(37), 236 Vandamme, L. J., 12(24), 13(27), 16(24), 64 (24),69 van Henvijnen, T., 283(62), 304(113), 307, 312,313 Van Hooff, J. H. C., 90(60), 91(60), 129 Vanhove, D., 58(254), 76 Van Hove, M., 143(30), 144(30), 146(30), 235 van Krevelen, D. W., 124, 133 van Langen, S. A. J., 33(177, 178), 62(177, 178), 63(177), 64(177, 178, 179), 65(179), 73 van Mao, L., 36(186), 37(187), 62(186), 74 Vannice, M. A,, 50, 51, 52(232), 53(231), 75, 202(220), 240,285(64), 312 Vannotti, L. E., 79(7), 81(7), 85(7), 127(7), 128 Van Reijen, L. L., 125(179, 180). 133 Vasil’eva, T. A,, 267(50), 311 Vasilevich, A. A,, 267(50), 311 Vedage, G., 285(63), 286(63), 302(63), 312 Vedrine, J. C., 30(160), 62(160), 64(160), 83 (25), 98(89), 99(89), 102(108b), 73, 128, 130,131 Verdonck, J. J., 57(252, 253), 58(252, 253, 255), 65(225), 76 Venuto, P. B., 2(3), 18(3), 66,68, 76 Verloop, A., 17(58), 70 Vermeulen, T., 21 3(228, 231), 240 Vogel, W., 292(76), 293(76), 295(76), 312 Vollhardt, K. P. C . , 30(162), 73

Volodin, A. M., 127(30b), 129 von, E., 35(183), 73 von Kutepow, N., 39(195), 74 Vorotinzev, V. M., 90(54), 129 Vorotyntsev, V. M., 87(37), 127(37), I29 Vostrikova, L. A., 11(17), 16(17), 68 Vickerman, J. C., 156(116), 159(116), 176 (116), 177(116), 178(116), 238 Viefnaus, H., 161(142), 162(142), 223(142), 230(142), 238 Vinogradova, 0. M., 14(44), 69

W Wachs, I. E., 298(81), 300(81), 312 Wachrenier, J., 11 1(141), 132 Wagner, H., I43(49), 145(49), 146(49), 148 (49), 150, 176(49), 177(49), 236 Walch, S. P., 165(149), 235, 238 Walker, W. C., 118(142), 132 Walsh, A. D., 98(94b), 130 Walsh, D. E., 57(266), 63, 64(267), 76 Walter, D., 32(170), 73 Walters, A. B., 94(73), 130 Warble, C. E., 107(130), 132 Ware, D. A., 81(18), 128 Warman, J. M., 90(56), 102, 129, I31 Watanabe, M., 151(73), 154(73), 186(73), 237 Waugh, K. C., 300(84), 306(84), 307(84), 309 (84), 312 Weatherbee, G. D., 188(194), 192(194), 197 (194), 198(194), 201(194), 21 1(194), 212 (194), 229(194), 230(194), 232(194), 239, 240 Webb, G . , 185(184), 239 Weeks, R. A,, 91(64), I29 Weeks, S. P., 145(50a), 155(50a), 161, 236 Weinberger, D. T.. 196(210), 211(210), 240 Weiss, F., 79, 125, 128 Weissermel, K., 16(48), 38(48), 69 Weisz, P. B., 9, 10(14), 13, 23, 55(239), 56 (239), 62(127), 66(239), 68,69, 72, 75 Wentreck,P. R., 154(114), 155(114), 165(114), 166(114), 167(114), 186(114), 118(1 14), 212(114), 213(233), 219(233), 237, 240 Wentrcek, P. W., 188(195), 192,198,199(195), 200, 211(195), 212(195), 213(195), 214 (195), 219(195), 239 Werlen-Ruze, B., 151(65), 152(65), 161(65), 236

333

AUTHOR INDEX

Wermann, J., 276(59), 311 Wertz, J. E., 83(24a), 87(42), 126(189), 127 (24a), 128, 129,133 Whan, D. A,, 55(240), 75,248(26), 276,310 Whited, R. C., 118(142), 132 Wilke, G., 32(167, 170), 35(167), 73 Wilkinson, G., 287(67), 312 Williams, A , , 230(275), 242 Williamson, W. B., 16, 70, 81(15), 90(15), 101 (15), 102(15), 128 Wilson, J. M., 161(143, 144), 162(143, 144), 238 Windawi, H., 138(10, II), 143(39), 159, 229 (II), 230(11), 231(11), 232(11), 235,236 Windhorst, K. A., 16(54, 55). 70 Winkler, A., I73(163), 175, 176(163), 239 Wise, H., 151(153), 154(114, 115), 157(53), 161(53), 165,166(114, 115), 167(114), 169 (152, 156), 170, 170(156), 176(169), 180 (169), 181(169), 182, 186(114), 188(114), l92( 19S), 198(l9S), 199(l9S), 200( 195), 211(195), 212(114, 195, 226), 213, 214 (195), 219,236,237,238,239,240 Witkinson, G., 79, 128 Wolstenholme, J., 156(116), 1591116), 176 (116), 177(116), 178(116), 238 Wong, N. B., 81, 84(30a), 85( 16), 86( 16), 90 (30a), 91(30a), 127(16, 30a), 128,129 Wood, B., 212(226), 240 Wood, B. J., 154(114), 155, 165,166(114), 167 (114), 186(114), 188(114), 212(114), 213 (233), 219(233), 237,240 Woodbury, H. H., 87(44), 126(44), 129 Woodward, L. A,, 191, 194,240 Wragg, R. D., 107(128), 132 Wyatt, W. V., 303(105), 313

Y Yamamoto, T., 250(21), 310 Yamashiro, T., 15(47), 61(47), 63(47), 69

Yamauchi, S., 250(21), 310 Yamazoe, N., 107(122b), 131 Yang, F. S., 104, 125,131 Yao, Y.-F. Y., 270(51, 52), 311 Yashima, T., 25, 26(140), 28, 29, 42(209), 43 (210), 44(209), 62(209, 210), 64(140), 65 (210), 67(209, 210), 72, 73, 74 Yates, J. T., Jr., 185(190), 186(190), 239 Yem, H. C., 29(155, 157), 73 Yoda, Y., 107(122b), 131 Yoneda, Y., 124(171), 133, 303(109), 313 Yoshida, S., 94(71d), 123(71d), 125(181), 130, 133 Youngblood, A. J., 21 1(222), 240 Yu, Y. F., 267(48), 311 Yun, C., 93(71a), 103, 104(111b), 120(158, 159, 160a), 121(160a), 130,131,132 Yur’eva, T. M., 303(104), 304(104), 307(104), 313

Z Zakharova, V. I., 18(67), 29(153), 70, 73 Zammitt, M. A,, 91(61), 107(122a), 129, 131 Zanobi, A., 47(226), 49(226), 63(226), 65 (226), 67(226), 75 Zapletal, V., 225(250), 226(250), 241 Zasorina, N. M., 45(216), 74 Zecchina,A., 113,114,117(146,147,154), 120 (154), 123(168), 124(173b), 132, 133, 302 (92), 304(116), 305(116), 312,313 Zelenina, M., 14(37), 69 Zettlemoyer, A. C., 252(29), 267(48), 311 Zharkov, B. B., 225(262), 228(262), 241 Zhdanov, S. P., 19(80), 70 Zimmermann, G., 220(243), 241 Zimmermann, H., 32(170), 73 Ziolek, M., 17(61), 70 Zubova, I. E., 220, 222(244), 241 Zueva, T. V., 40(203), 41(203), 42(203), 43 (203), 44(203), 65(203), 74 Zulfugarov, Z. G., 14(42), 69

Subject Index

A Absorption light, 112 bands, alkaline-earth oxides, 112-113, 117 spectra, 116 Acetylene adsorption complexes, 6-7 oligomerization, 29-30 Acid sites, 5-6, 19 Acidity, ion exchange and, zeolites, 5-6 Activation energy adsorbed formate, 307-308 exchange reactions, 106 methanation, 202-203 methanol synthesis, 282 promoted iron catalyst, 221, 223 rhodium zeolites, 43-45 Activation temperature, radical-forming ability, 109-110 Activity, methanation, 200-203 Adsorption CS2, 156 complexes, 6-7 dependence on Cu/ZnO ratio, 269 energy, weakening, Cu/ZnO binary catalyst modifiers, 283 halogens, 110-1 11 hydrogen, MgO, 122 isotherm, 171,223 mercaptans, 156 methanol synthesis, 282 N,O, 90 oxygen species formation, 90-92,94-95 rates, H2S, 153-154 sulfur on metals, see Sulfur, adsorption on metals Aggregate oxygen species, 95-101, see also Dimer species; Trimer species AI,O,, electron donor properties, 11 I

Alcohols, higher, carbonylation, 45-46 Alkali metal zeolites, 14 oligomerization, 31, 38 Alkali promoters, Fischer-Tropsch catalysts, 191 Alkaline-earth oxides absorption bands, 112-1 13, 1 I7 exciton levels, 1I8 powders, 112-113 reflectance spectra, 113-1 14 single crystals, 95-96 Alkanes, 13- 14, 103-104 Alkenes hydrogenation, sulfur toxicity, 226-227 hyperfine interaction, 104-105 mononuclear oxygen species, reactivity, 104- 106 oxidation, 14-16, 63 Alkoxide, decomposition, 103 Alumina-supported catalysts, methanation, activity, 196-198 Ammonia, synthesis, 220-224

B Benzene, hydrogenation, 2 1 Bond strength, 124-125, 138 Bonds, metal-sulfur, see Sulfur poisoning, metal-sulfur bonds Bromine, reactivity, 111 Bronsted acid sites, 5 Butadiene, oligomerization, 32-36 Butene, isomerization, 122- 123 n-Butene dimerization, 36-37 composition change, 25-26

C CS, adsorption, 156 334

SUBJECT INDEX sulfur toxicity, 226-228 CaO, trimer species, 100- 101 Calcination, Cu/ZnO binary catalyst, 261-262 Carbanion, butene isomerization, 122-123 Carbon conversion, 272,279-281 Carbonylation, 39-46 ethanol, 45-46 ethylene, 46 higher alcohols, 45-46 methanol, 39-45, see Methanol, carbonylation Catalysts activity, effects of sulfur, 187-229 binary, 287-291, see also specific catalysts composition, hydrogenation, sulfur poisoning, 227 deactivation rates, 21 1-216 metal, see specific catalysts mixed, see specific catalysts pellets, external sulfur coverage, 218-219 poisoning susceptibility, 226-228 selection, methanol synthesis, 25 1-254 sulfur-active, 197-198 sulfur poisoning regeneration, 229-232 tolerance, 227 Cation diffusion, bulk sulfide formation, 153 oxidation, chemisorption, 94-95 siting, zeolites, 6-9, 33-34 C,H, , kinetic analysis for, 104 CH,, kinetic analysis for, 104 Chabazite, alkane oxidation, 14 Charge-transfer complex, photoluminescence, 125 Chemisorption, see CO, chemisorption; H,, cheinisorption Chlorine, reactivity, 110-1 11 Chromium oxide, g tensor, 83 C-0 bond, weakening by back-donation metal, 252

co adsorbed, methanol synthesis, 302-303 adsorption catalysts other than nickel, preadsorbed sulfur, 185-186 H,S effects, 181-182 infrared bands, nickel catalyst, 180-1 82 preadsorbed sulfur effects, 177-179 chemisorption, 183,251-253,268-271

335

coadsorption, methanol synthesis, 304-308 hydrogenation, see Methanol synthesis oxidation, 10-13, 102 pressure, Zn-H stretching bands, 304-305 CO/H, , see Synthesis gas CO, effects, methanol synthesis, 274-284 hydrogenation, see Methanol synthesis photoformation, 123 selectivity effects, Cu/ZnO binary catalyst, 284-285 steady-state concentrations, 284 Cobalt zeolite, hydroformylation, 47-49 Cobalt hydrocarbon synthesis catalysts, 58 Color spectra, Cu/ZnO binary catalyst, 25926 1 Contaminant species, transport continuity equation, 213 Continuous-flow stirred-tank reactor, 189 Controlled-atmosphere studies, choice of materials for construction, 188 Copper amorphous, 259-260,270-271 average particle size, 261-262 pure, activity, 254-255 relationship between weakly and irreversibly chemisorbed O,, 270 scanning electron micrograph, 255 solid solution, 259-260 surface areas, Cu/ZnO binary catalyst, 267268 Copper catalysts activity, effects of sulfur, 187-229 binary, 287-289, see also Cu/ZnO; Zinc chromite CO adsorption, preadsorbed sulfur, 186 methanation, 195-198,202-204,208 methanol synthesis, 245-247 Copper zeolites, butadiene oligomerization, 32-35 Crotonic acid, hydrogenation, sulfur poisoning, 227 Crystal field, influence on g tensor, 88 Crystallinity, X-ray, Cu/ZnO/Al,O, , 293 Cu2'-exchanged zeolites, 11-13 Cu/ThO, binary catalyst, 287-288 Cu/Zn/Al,O, catalysts, 249 Cu/Zn/Cr,O, catalysts, 250 Cu/ZnO binary catalyst, 257-287 activity patterns, 271-274

336

SUBJECT INDEX

BET argon surface areas, 259 calcination, 261-262 catalytic testing, 272 chemisorption, 268-271 CO, effects,selectivity, 284-285 color spectra, 259-261 component comparison, 258-259 methanol synthesis, 246-247 modifiers, weakening of adsorption energy, 28 3 optical spectra, 259-261 particle size and morphology, 261-266 physical characteristics, 258-261 preparation, 258 reaction kinetics, presence of CO, ,274-284 redox reaction, 278 reduced, area diffraction pattern, 262, 265 relative hydrogenation rates, 285-287 structure, 266 surface analysis, 266-268 synthesis model, 278 Cu/ZnO/Al,O, , 292-295, 276 CuO, 11, 13,254,261-262 CuY, oxidation, 11-12, 14 Cyclodimerization, butadiene, 32-36, 64-65 Cyclohexane dehydrogenation, 14,21-22 reforming, selective poisoning, 229 Cyclopropenes, oligomerization, 31

D Deactivation rates, 21 1-216 Dehydrocyclization, 23 Dehydrogenation organic compounds, 225-229 selective poisoning effects, 228 zeolites, 21-24 Dehydrohalogenation, 41,45 Dehydroxymethylation, selective poisoning effects, 228 Desorption energy, adsorbed formated, 307308 Desorption isotherms, H,S, 166, 168 Diffusion, in zeolites, 9-10 Dimer species, 95-98 Dimerization, 24-28, 36-37 Dimethyl ether, 41, 244 Dissociation, H,S, 155-156

E Electron donor, properties, surface oxide ions, 109-1 12 Electron paramagnetic resonance charge-transfer complex, 125 exchange reactions, 106-107 invisible oxygen species, 94-95 oxygen species, characterization, see Mononuclear oxygen species, characterization parameters alkaline-earth oxide single crystals, 95-96 dimer species, 95-96 hole centers, 93 MgO, 100 oxygen species, 126-127 TiO, , 9 8 radical detection, 93 shape of axially symmetric signal, 83-85 signal, mononuclear oxygen species, reactivity, 101-102 spectrum dimer species, linewidth, 95 oxygen species, MgO, 81-82 Emission spectra, 113, 115-116 Energy activation, see Activation energy bonding orbital, trimer species, 98-99 formation, see Formation energy free, 244, 298-299 gap, (n-4a), 251-252 levels oxygen species, 80-81 molecular orbital, 97, 80 optical transition, 89 stabilization, oxygen species, 88 surface, 1 16- 1I7 Entropy, 61, 142 Erionite, alkane oxidation, 14 Ethane, steam reforming, 217 Ethanol, 45-46 Ethylbenzene, cyclodimerization, 36 Ethylene carbonylation, 46 dimerization, 24-28 hydrogenation, 20 mononuclear oxygen species, reaction with, 104-105 oligomerization, 24-29 polymerization, 29

337

SUBJECT INDEX

Exchange reactions, 106- 107 Excitation spectra, 113, 115--116, 119, 121 Exciton absorption, MgO, I13 levels, alkaline-earth oxides, 118 transitions, 116-1 17

F Fischer-Tropsch catalysts, 191 Fixed-bed reactor, catalyst choice, 188 Flujd cracking catalysts, 13 Formaldehyde, formation, methanol synthesis, 300 Formate, adsorbed, methanol synthesis, 303304, 307-308 Formation energy adsorbed sulfur stability, 164 free, surface versus bulk sulfides, 169, 171 sulfides, 140-142

G g tensor, 79-83

axially symmetric, 83-85 components, 80 EPR parameters, 126-127 MgO, 85 orthorhombic symmetry, 81, 83 oxygen species, formed by adsorption, 90 spin density, 81-82 stabilizing cation charge and, 88 surface crystal field, influence, 88 values, 85, 87 Gibbs free energy, 244, 298-299 Gradientless internal recycle reactor, 202

H H2 adsorbed, methanol synthesis, 302-303 adsorption, 122, 173-175, 178-179, 185 catalyst regeneration, 230 chemisorption, 268-271 coadsorption, methanol synthesis, 304-308 sticking coefficient, 175-176 Halogens, 1 10- 1 1 1 Hamiltonian, spin, dimer species, 95 Heat of adsorption, 51, 171-172 versus heat of formation, 165-166,169-170

Heat of formation bulk sulfides, 142 versus heat of adsorption, 165-166, 169170 Heat pretreatment, oxygen species formation, 94-95 Hexene-1, hydroformylation, 49-50 H2S adsorption rates, 153-154 chemisorption on nickel, 154- 155 concentration, 139, 205-206 desorption isotherms, 166, 168 dissociation on metals, 155-156 equilibrium, 165- 167 oxidation, 17 poisoning, deactivation rate constants, 214 sticking coefficient, 154 Hydrocarbons cracking, preadsorbed sulfur effects, I84 relative hydrogenation rates, 285-287 synthesis gas conversion, 54-58 Hydrocracking, suppression by sulfur, 229 Hydroformylation, 46-50, 63, 65 Hydrogenation alkenes, sulfur toxicity, 226-227 catalyst composition, sulfur poisoning, 227 CO, see Methanol synthesis CO, , see Methanol synthesis crotonic acid, sulfur poisoning, 227 Gibbs free energy, 244 low-temperature, sulfur toxicity, 225 olefins, sulfur poisoning, 191-192, 226 organic compounds, 225-229 rates, substrate dependence, 285-286 selective poisoning effects, 228 zeolites, 19-21 Hydrogenolysis, organic compounds, 225229 Hyperfine interaction, reaction with alkenes, 104-105 Hyperfine splitting, 81 Hyperfine structure, trimer species, 98-99 Hyperfine tensor, oxygen species, 85-86

I Infrared spectra, 124,201 Iodine, reactivity, 11 1 Ion exchange, 5-6, 12 Ionizing radiation, oxygen species formation, 92-94

338

SUBJECT INDEX

Iron catalysts, 190-191, 202-203, 208, 222223 Isobutylene, oligomerization, 37-39 Isomerization, butene, 122-123

K Kinetics, rate equations, 275-276 Kolbel-Engelhardt reaction, 59

L Langmuir isotherm, 155 Lattice parameters, Cu/ZnO, 247 Levine-Mark theory, 116-1 19 Lewis acid sites, 6 Light absorption, 112 Low-energy electron diffraction, 146,148-149

M M=O bond, 124-126 Madelung potential, 112 Mercaptans, adsorption, 156 Metals back-donation, C - 0 bond weakening, 252 catalytic activity, methanol synthesis, 253 ion-oxygen bond strength, 124-125 position in periodic table, ability to synthesize methanol, 252-253 Metal catalysts, sulfur poisoning, see Sulfur poisoning Metal ion-exchanged zeolites, 13 Metal oxides, 78-79, 89, 102, 123, 157-158, 191, see also specific oxides Metal-sulfur bonds, see Sulfur poisoning, metal-sulfur bonds Methanation, 50-55 activity, 51, 196-198 gravimetric measurements, 198-199 nickel catalysts, 52-53, 198, 212 palladium zeolites, 51-52 ruthenium zeolite, 53-55 sulfur poisoning, 192-195 turnover numbers, 52-53,202-203, 209 water vapor effect, 193-194 zeolite stability, 63 Methane, 284,289-290 Methanol carbonylation, 39-45

other metal zeolites, 45 rhodium catalysts, 39-40 rhodium zeolites, 40-45 zeolite Y, 42-43 chemisorbed, infrared spectra, 201 decomposition, 245-247 dehydrohalogenation, 41 Methanol synthesis, 190-21 1, 243-313, see also Carbon conversion activation energy, 282 adsorption, 282 catalyst selection, 251 -254 CO, effects, 274-284 CuO catalyst, test, 254 deactivation rates, 21 1-216 defined, 243-244 H, :CO ratio, 306 high-pressure process, 248 low-pressure, 245, 248, 308-310 mechanisms, 296-310 adsorbed CO, 302-303 adsorbed formate, 303-304,307-308 adsorbed H, , 302-303 adsorbed methoxide, 303-304 coadsorption, 304-308 low temperature and pressure, 308-310 methods used in studies, 299-302 reaction pathways, 297-299, 309 model, 278 rates, 272-274,277 reaction conditions, 193 Rh/LaB, binary catalyst, 289-291 sulfur poisoning, 190 versus decomposition, 246 ZnO catalyst, 246 Methoxide, adsorbed, methanol synthesis, 303-304 Methyl acetate, specific activity, 41 Methyl formate, formation, 300 MgO alkene oxidation, 104-105 butene isomerization, 122- 123 electron donor properties, 109-1 10, 123 EPR, 81-82, 100 excitation spectra, 119 exciton absorption, 113 g tensor, 85 halogen adsorption, 110-1 11 hydrogen adsorption, 122 hyperfine tensor, 85

339

SUBJECT INDEX

irradiation, exchange reactions, 107 mononuclear oxygen species, 81-82, 101 photoluminescence spectra, 11 5-1 16, 1 19120 radical-forming ability, 109- 110 reactivity, alkanes, 103 reflectance spectra, 114 spin density, 8 1-82 surface imperfections, 108 surface trimer species, 100 thermal activation, 94 MoO,/AI,O, , 91 MoO,/SiO, , 86 Molecular orbitals, energy levels, 97, 80 Molybdenyl compounds, M=O bond, 125 Mononuclear oxygen species, 12, 77-134, see also Surface oxide ion aggregate, see Aggregate oxygen species energy levels, 80-81 EPR, characterization, 79-89 orthorhombic symmetry. 81, 83 invisible, 94-95 g tensor, 79-83 UV irradiation, 93 formation, 90-95 adsorption, 90-92,94-95 heat pretreatment, 94-95 hole centers, 92-93 ionizing radiation, 92-94 hyperfine splitting, 81 hyperfine tensor, 85-86 optical spectroscopy, 79-89 radicals, 93 reactivity, 101-107 alkanes, 103-104 alkenes, 104- 106 exchange reactions, 106-107 organic molecules, 103-106 simple inorganic molecules, 101-102 spin densities, 85 spin-orbit coupling constant, 83 stability, 78-79,90-95 superfine tensor, 86-88 thermal activation, 94 Mordenite, 3, 5 , 8, 13-14, 21

N NaX, cyclodimerization, butadiene, 35 NH, ,oxidation, I6

N,O, adsorption, 90 Nickel catalysts adsorbed sulfur, 146, 148-150, 154-155 activity, methanation, 195-198, 204-208 Boudard reactions, 200-201 deactivation, 198-200, 212, 215 gravimetric measurements, 198-199 H,S chemisorption on, 154- 155 H,S desorption isotherms, 166, 168 methanation, 200-201,203 preadsorbed sulfur effects on CO adsorption, 176-183 effects on H, adsorption, 173-176 presulfided, 178-179,201-202,207 regeneration, 230-232 site densities, 214-215 sulfur adsorption effects on adsorption of other molecules, 183-185 preliminary states, 146-147 single-crystal planes, 144-145 stability, 165-169 stoichiometries, 158-161 surface structures, 143-151 sulfur bonding, schematicviews, 146-147 sulfur coverage, ammonia synthesis, 219 sulfur poisoning, 21 1, 226-227 Nickel sulfides, bond strength, 138 Nickel zeolites, 20, 24-25, 30-31, 36, 52-53 Nickel-sulfur bonds, surface versus bulk, stability, 165 Nickel-sulfur system, H,S concentration, 139 Nitriding, iron, 223-224 Nitrobenzene ions, electron donor properties, 109-110 NO adsorption complexes, 6-7 oxidation, 16-17 NO,, oxidation, 16

0 0 - ions, see Mononuclear oxygen species Olefins hydroformylation, 46-47 hydrogenation, sulfur poisoning, 191- 192, 226 oligomerization, rhodium zeolites, 25 Oligomerization acetylene, 29-30

340

SUBJECT INDEX

butadiene, 32-36 n-butene, 36-37 cyclopropenes, 31 ethylene, 24-29 isobutylene, 37-39 olefins, 25 zeolites, 24-39 Optical spectroscopy, 79-89,259-261 Oxidation active sites, 64 alkanes, 13- 14 alkenes, 14-16, 63, 104-105 catalyst regeneration. 231-232 cation, chemisorption, 94-95 CO, 10-13, I02 miscellaneous reactions, 18 sulfur-containing compounds, 17- 18 nitrogen-containing compounds, 16- 17 zeolites, 10-18 Oxygen adsorption, 184 catalyst regeneration, 231 reactivity of oxide, 121-122 species, mononuclear, see Mononuclear oxygen species Oxygenates, relative hydrogenation rates, 285-287

P Palladium zeolites, 19-20, 28, 51-52 Paracrystallinity, Cu/ZnO/Al,O, , 295 Paramagnetic oxygen species, 78-79, see also Mononuclear oxygen species Peroxide salts, dimer species, 96-97 pH, ion-exchange solution, 12 Phillips catalyst, 29 Phosphorescence spectra, transition-metal oxides, 121 Photoactivation. lattice oxygen, 123 Photoformation, CO,, 123 Photoluminescent spectra, 114-1 16, 119-121, I25 Photolysis, water, 60 Photoreduction, metal oxides, 123 Platinum catalysts adsorbed sulfur, 15 1- 152, 156-1 57 sulfur adsorption, stoichiometries, 161-163 sulfur toxicity, 226-227, 229

Platinum zeolites, dehydrogenation, 22 Polymerization, isobutylene, 38-39 Pore parameters, zeolites, 3-4,9 Presulfiding, partial, 194-1 96 Propylene hydroformylation, 48-49 oligomerization, 30-31 oxidation, 14-16 Pt/AI,O,, oxygen species, stability, 91-92

Q Quaternary catalysts, 291 -296

R Reaction-selective poisoning, 228 Reactivity, mononuclear oxygen species, see Mononuclear oxygen species, reactivity Redox reaction, Cu/ZnO binary catalyst, 278, 283 Reflectance spectra, alkaline-earth oxides, 113-1 I4 Regeneration, using steam, 230 Rh/LaB, binary catalyst, 289-291 RhNaX, carbonylation, 40-43,46 RhNaY, 42-43,49-50 Rhodium catalysts, 39-42, 291 Rhodium zeolites, 20-21 activation energy, 43-45 ethylene dimerization, 25-28 hydroformylation, 49-50 kinetics, 43-45 mechanism, 43-45 methanol carbonylation, 40-45 specific activity, 41,44-45 Rhodium zeolite X , 40-42 Rostrup-Nielsen calculations, 2 18-2 19 Ruthenium catalysts H,S desorption isotherms, 168 methanation, 202-203, 208 Ruthenium zeolites, 53-55, 57-59

S SIMS intensity ratio, CO-containing species versus sulfur coverage, 177-178 SO,, oxidation, 17-18 Scanning electron micrography, copper. 255

SUBJECT INDEX

34 I

Scanning transmission electron microscopy, model, 212-214 Cu in solid solution, 259-260 nickel single-crystal planes, 144- 145 Schulz-Flory kinetics, synthesis gas converpresence of other molecules, 163 sion, 55 reversibility, 216 Selectivity saturation coverage, 162- 163 Boudouard reactions, 200-201 stability, 164-172 Fischer-Tropsch catalysts, 191 stoichiometries, 158-164 catalytic, effects of sulfur, 187-229 surface structures, 143-153 methanation, 192- 195, 200-20 I bonding chemistry, 137-138, 146-147 partial presulfiding, 1944 196 coverage, 175-176.209, 218-219 Silica-supported oxides, mononuclear oxyeffects on adsorption of other molecules, gen species, 102 172-187 Silver zeolites, water splitting, 60 nickel catalyst, 173-1 85 Sintering, iron catalysts, enhancement by sulother catalysts, 185-187 fur poisoning, 223 hydrocracking suppression, 229 Spectroscopic studies, see also specific specpoisoning, 135-242, see also Sulfides troscopic methods catalyst regeneration, 229-232 nickel-sulfur bonding, 138 CO hydrogenation, 190 surface oxide ions, 112-121 conclusions, 233-234 Spin-orbit coupling constant, oxygen species, controlled-atmosphere studies, 188 83 deactivation, methanation, 194 Stability experiment considerations, 187-1 89 oxygen species, 88,90-95, 101 Fischer-Tropsch catalysts, 191 sulfur adsorption, 164- 172 fundamental questions, 136- 137 nickel catalysts, 165-169 iron catalysts, 190, 222-223 Steam reforming, 216-220 mechanisms, 21 1 ethane, sulfur poisoning effects, 217 metal-sulfur bonds, 137-143 high-temperature, 218 methanation, 192-195 selective poisoning effects, 229 reaction-selective, 228 sulfur coverage profiles, 218-219 recommendations, 234 sulfur threshold levels, 216-21 7 site-selective, 228 Steam regeneration, 230 susceptibility, 226-228 Sticking coefficient, 154, 175-176 preadsorbed, 177- 178, 184- 186 Structural promotion, 247 resistance, promoted iron effects, 220-221 Sulfides tolerance, 227 bond strength, 138 toxicity, 225-228 bulk Sulfumontaining compounds, oxidation, 17formation mechanism, 153 18 thermodynamics, 139- 143 Superhyperfine interaction, 81, 83,86-88 versus surface, free energies of formation, Surface oxide ion, 107-124 169, 171 chemical reactivity, 121-124 Sulfur coordination, 110-1 11 adsorbed electron donor properties, 109-1 12 location, 146 spectroscopic studies, 1 12- 121 nickel, 144-150, 154, 159 Synergic promotion, 247 platinum, structures, 151-152 Synthesis gas, conversion, 54-58 adsorption on metals, 143-172 cobalt hydrocarbon catalysts, 58 kinetics, 153-158 Schulz-Flory kinetics, 55 mechanisms, 153- 158 ZSM-5 zeolite. 55-56

342

SUBJECT INDEX

T TeNaX, dehydrogenation, 23 Temperature activation, radical-forming ability, 109-1 10 effects on sulfur toxicity, 227 oxygen regeneration, 231 Temkin theory, 155 Ternary catalysts, 29 1-296 Theophene, nickel catalyst poisoning, 226 Thermodynamics, bulk sulfides, 139- 143 Th o , , photoluminescence spectra, 120 TiO,, UV irradiation, trimer species, 98 Titanium zeolite, water splitting, 60-61 Transition metal ion-exchanged zeolites, 1 1 13 active species, 64-65 acetylene oligomerization, 29-30 hydrogenation, 19 oxidation alkane, 13-14 alkene, 14-16 hydrogen, 18 nitrogen-containing compounds, 16- 17 Transition-metal oxides, 120-121, 124-126 Transition-metal catalysts, 289-29 1 Transmission electron microscopy, 256, 263264,294 Trimer species, 98-101 Turnover numbers, 54-55,202-203,209

U UV irradiation, oxygen species, 93, 98

V V@,, lattice oxygen photoactivation, 123 Vinylcyclohexene, cyclodimerization, butadiene, 32-35 Vycor glass, porous, 120, 123

W Wacker catalyst system, 15-16 Water adsorption on metal oxides, 157-158 photolysis, 60 splitting, 59-61, 65 steady-state concentrations, 284

Water-gas shift, 58-59 Wax formation, iron catalyst, 191

X X-ray photoelectron/Auger spectroscopy, 266-267 X-ray powder diffraction, Cu in solid solution, 259-260

Y Y parameter, 14-15

Z ZSM-5 zeolite, synthesis gas conversion, 5556 Zeolites, nonacid catalysis, 1-76, see also specific zeolites acidity, 5-6 active species, 64-66 activity, 61 -62 adsorption complexes, 6-7 cation siting, 6-9 coke yield versus shape selectivity, 57 composition, 3-5 dehydrogenation, 21 -24 diffusion in, 9-10 future prospects, 66-68 hydrogenation, 19-2 1 ion-exchanged, 5-6, 1 1- 13,59-60 mordenite structure, 5 oligomerization, 24-39 oxidation reactions, 10- 18 oxygen bridge species, 12 pore parameters, 3-4 reaction mechanisms, 64-66 schematic diagram, 4 selectivity, 61-62 shape-selective, 56-58,66-67 stability, 62-64 structure, 3-5 sulfur poisoning resistance, 54 Y parameter, 14-15 Zeolite A, 3-4,6-7, 14,21, 60-61 Zeolite L, 14, 21

SUBJECT INDEX

Zeolite T, 13 Zeolite X carbonylation, methanol, 40-42 cation siting, 7 composition, 3 cyclodimerization, butadiene, 35 dehydrogenation, 22-23 hydrogenation, 21 metal ion-exchanged, 13 oligomerization, 30-31, 38 oxidation, 11-12, 14, 18 structure, 4 Zeolite Y carbonylation, methanol, 42-43 cation siting, 7 composition, 3

343

dehydrogenation, 22-23 dimerization, ethylene, 25-27 hydrogenation, 21 methanation, 53-54 oligomerization, 29-31, 38 oxidation, 11-12, 14-15, 18 polymerization, ethylene, 29 structure, 4 Zinc chromite, 248,275-277 Zn-H stretching bands, CO pressure, 304-305 ZnO average particle size, 261 -262 methanol synthesis, 246 mononuclear oxygen species, stability, 91 pure, 256-257 surfaces, 259-260, 268

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Contents of Previous Volumes Volume 1 The Heterogeneity of Catalyst Surfaces for Chemisorption HUGHS. TAYLOR Alkylation of Isoparaffins V. N. IPATIEFF AND LOUISSCHMERLING Surface Area Measurements. A New Tool for Studying Contact Catalysts P. H. EMMETT The Geometrical Factor in Catalysis R. H. GRIFFITH The Fischer-Tropsch and Related Processes for Synthesis of Hydrocarbons by Hydrogenation of Carbon Monoxide H. H. STORCH The Catalytic Activation of Hydrogen D. D. ELEY Isomerization of Alkanes HERMAN PINES The Application of X-Ray Diffraction to the Study of Solid Catalysts AND I. FANKUCHEN M. H. JELLINEK

Volume 2 The Fundamental Principles of Catalytic Activity FREDERICK SEITZ The Mechanism of the Polymerization of Alkenes LOUIS SCHMERLING AND v. N. IPATIEFF Early Studies of Multicomponent Catalysts ALWINMITTASCH Catalytic Phenomena Related to Photographic Development T. H. JAMES Catalysis and the Adsorption of Hydrogen on Metal Catalysts OTTOBEECK Hydrogen Fluoride Catalysis J . H. SIMONS Entropy of Adsorption CHARLES KEMBALL

About the Mechanism of Contact Catalysis GEORG-MARIA SCHWAB

Volume 3 Balandin’s Contribution to Heterogeneous Catalysis B. M. W. TRAPNELL Magnetism and the Structure of Catalytically Active Solids P. W. SELWOOD Catalytic Oxidation of Acetylene in Air for Oxygen Manufacture AND K. A. KRIEGER J. HENRYRUSHTON The Poisoning of Metallic Catalysts E . B . MAXTED Catalytic Cracking of Pure Hydrocarbons VLADIMIR HAENSEL Chemical Characteristics and Structure of Cracking Catalysts JR., AND A. G. OBLAD,T. H. MILLIKEN, G. A. MILLS Reaction Rates and Selectivity in Catalyst Pores AHLBORN WHEELER Nickel Sulfide Catalysts WILLIAM J. KIRKPATRICK

Volume 4 Chemical Concepts of Catalytic Cracking R. C. HANSFORD Decomposition of Hydrogen Peroxide by Catalysts in Homogeneous Aqueous Solution J. H. BAXENDALE Structure and Sintering Properties of Cracking Catalysts and Related Materials E. RIES,JR. HERMAN Acid-Base Catalysis and Molecular Structure R. P. BELL Theory of Physical Adsorption TERRELL L. HILL

345

346

CONTENTS OF PREVIOUS VOLUMES

The Role of Surface Heterogeneity in Adsorption GEORGED. HALSEY Twenty-Five Years of Synthesis of Gasoline by Catalytic Conversion of Carbon Monoxide and Hydrogen HELMUT PICHLER The Free Radical Mechanism in the Reactions of Hydrogen Peroxide JOSEPHWEISS The Specific Reactions of Iron in Some Hemoproteins PHILIPGEORGE

Volume 5 Latest Developments in Ammonia Synthesis ANDERSNIELSEN Surface Studies with the Vacuum Microbalance: Instrumentation and Low-Temperature Applications T. N. RHODIN,JR. Surface Studies with the Vacuum Microbalance: High-Temperature Reactions EARLA. GULBRANSEN The Heterogeneous Oxidation of Carbon Monoxide MORRISKATZ Contributions of Russian Scientists to Catalysis J. G. TOLPIN,G. S . JOHN,AND E. FIELD The Elucidation of Reaction Mechanisms by the Method of Intermediates in QuasiStationary Concentrations J. A. CHRISTIANSEN Iron Nitrides as Fischer-Tropsch Catalysts ROBERTB. ANDERSON Hydrogenation of Organic Compounds with Synthesis Gas MILTONORCHIN The Uses of Raney Nickel EUGENELIEBERAND FREDL. MORRITZ

Some General Aspects o Chernisorption and Catalysis TAKAOKWAN Nobel Metal-Synthetic Polymer Catalysts and Studies on the Mechanism of Their Action WILLIAM P. DUNWORTH AND F. F. NORD Interpretation of Measurement in Experimental Catalysis P. B. WEISZAND C. D. PRATER Commercial Isomerization B. L. EVERING Acidic and Basic Catalysis MARTINKILPATRICK Industrial Catalytic Cracking RODNEY V. SHANKLAND

Volume 7 The Electronic Factor in Heterogeneous Catalysis M. McD. BAKERAND G. I. JENKINS Chemisorption and Catalysis on Oxide Semiconductors AND M. BOUDART G. PARRAVANO The Compensation Effect in Heterogeneous Catalysis E. CREMER Field Emission Microscopy and Some Applications to Catalysis and Chemisorption ROBERTCOMER Adsorption on Metal Surfaces and Its Bearing on Catalysis JOSEPHA. BECKER The Application of the Theory of Semiconductors to Problems of Heterogeneous Catalysis K. HAUFFE Surface Barrier Effects in Adsorption, Illustrated by Zinc Oxide , S. ROY MORRISON Electronic Interaction between Metallic Catalysts and Chemisorbed Molecules R. SUHRMANN

Volume 6

Volume 8

Catalysis and Reaction Kinetics at Liquid Interfaces J. T. DAVIES

Current Problems of Heterogeneous Catalysis J. ARVIDHEDVALL

CONTENTS OF PREVIOUS VOLUMES Adsorption Phenomena J. H. DE BOER Activation of Molecular Hydrogen by Homogeneous Catalysts S . W. WELLERAND G. A. MILLS Catalytic Syntheses of Ketones v . I. KOMAREWSKY AND J. R. COLEY Polymerization of Olefins from Cracked Gases EDWINK. JONES Coal-Hydrogenation Vapor-Phase Catalysts E. E. DONATH The Kinetics of the Cracking of Cumene by Silica-Alumina Catalysts CHARLESD. PRATERAND RUDOLPHM. LAGO Volume 9 Proceedings of the International Congress on Catalysis, Philadelphia, Pennsylvania, 1956 Volume 10 The Infrared Spectra of Adsorbed Molecules AND W. A. PLISKIN R. P. EISCHENS The Influence of Crystal Face in Catalysis ALLANT. GWATHMEY AND ROBERTE. CUNNINGHAM The Nature of Active Centers and the Kinetics of Catalytic Dehydrogenation A. A. BALANDIN The Structure of the Active Surface of Cholinesterases and the Mechanism of Their Catalytic Action in Ester Hydrolysis F. BERGMANN Commercial Alkylation of Paraffins and Aromatics EDWINK. JONES The Reactivity of Oxide Surfaces E. R. S . WINTER The Structure and Activity of Metal-on-Silica Catalysts G. C . SCHUITAND L. L. V A N REIJEN Volume 11 The Kinetics of the Stereospecific Polymerization of a-Olefins G. NATTAAND I. PASQUON

347

Surface Potentials and Adsorption Process on Metals R. V. CULVER AND F. C. TOMPKINS Gas Reactions of Carbon P. L. WALKER,JR., FRANK RUSINKO,JR., AND L. G. AUSTIN The Catalytic Exchange of Hydrocarbons with Deuterium C. KEMBALL Immersional Heats and the Nature of Solid Surfaces J. J. CHESSICKAND A. C. ZETTLEMOYER The Catalytic Activation of Hydrogen in Homogeneous, Heterogeneous, and Biological Systems J. HALPERN Volume 12 The Wave Mechanics of the Surface Bond in Chemisorption T. B. GRIMLEY Magnetic Resonance Techniques in Catalytic Research D. E. O’REILLY Bare-Catalyzed Reactions of Hydrocarbons HERMAN PINES AND LUKE A. SCHAAP The Use of X-Ray and K-Absorption Edges in the Study of Catalytically Active Solids ROBERT A. VANNORDSTRAND The Electron Theory of Catalysis on Semiconductors TH. WOLKENSTEIN Molecular Specificity in Physical Adsorption D. J . C. YATES Volume 13 Chemisorption and Catalysis on Metallic Oxides F. S. STONE Radiation Catalysis R. COEKELBERGS, A. CRUCQ, AND A. FRENNET Polyfunctional Heterogeneous Catalysis PAULB. WEISZ A New Electron Diffraction Technique, Potentially Applicable to Research in Catalysis L. H. GERMER

348

CONTENTS OF PREVIOUS VOLUMES

The Structure and Analysis of Complex Reaction Systems JAMBWEI AND CHARLES D. PRATER Catalytic Effect in Isocyanate Reactions A. FARKAS AND G. A. MILLS

Volume 14

Quantum Conversion in Chloroplasts MELVIN CALVIN The Catalytic Decomposition of Formic Acid P. MARS, J. J. F. SCHOLTEN,AND P. ZWIETERING Application of Spectrophotometry to the Study of Catalytic Systems JR. H. P. LEFTINAND M. C. HOBSON, Hydrogenation of Pyridines and Quinolines MORRISFREIFELDER Modern Methods in Surface Kinetics.: Flash, Desorption, Field Emission Microscopy, and Ultrahigh Vacuum Techniques GERTEHRLICH Catalytic Oxidation of Hydrocarbons L. YA. MARGOLIS

Volume 15

The Atomization of Diatomic Molecules by Metals D. BRENNAN The Clean Single-Crystal-Surface Approach to Surface Reactions N. E. FARNSWORTH Adsorption Measurements during Surface Catalysis KENZl TAMARU The Mechanism of the Hydrogenation of Unsaturated Hydrocarbon on Transition Metal Catalysts G. C. BONDAND P. B. WELLS Electronic Spectroscopy of Absorbed Gas Molecules A. TERENlN The Catalysis of Isotopic Exchange in Molecular Oxygen G. K. BORESKOV

Volume 16

The Homogeneous Catalytic Isomerization of Olefins by Transition Metal Complexes MILTONORCHIN The Mechanism of Dehydration of Alcohols over Alumina Catalysts HERMAN PINESAND JOOST MANASSEN A Complex Adsorption in Hydrogen Exchange on Group VIlI Transition Metal Catalysts J. L. GARNETTAND W. A. SOLLICHBAUMGARTNER Stereochemistry and the Mechanism of Hydrogenation of Unsaturated Hydrocarbons SAMUEL SIEGEL Chemical Identification of Surface Groups H. P. BOEHM Volume 17

On the Theory of Heterogeneous Catalysis NAKAMURA JUROHORlUTI AND TAKASHI Linear Correlations of Substrate Reactivity in Heterogeneous Catalytic Reactions M. KRAUS Application of a Temperature-Programmed Desorption Technique to Catalyst Studies R. J. CVETANOVIC A N D Y. AMENOMIYA Catalytic Oxidation of Olefins R. ADAMS HERVEY H. VOGEAND CHARLES The Physical-Chemical Properties of Chromia-Alumina Catalysts CHARLES P. POOLE,JR. AND D. S. MACIVER Catalytic Activity and Acidic Property of Solid Metal Sulfates KOZO TANABE AND TSUNEICHI TAKESHITA Electrocatalysis S. SRINIVASEN, H. WROBLOWA,AND J. O’M. BOCKRIS Volume 18

Stereochemistry and Mechanism of Hydrogenation of Napthalenes in Transition Metal Catalysts and Conformational Analysis of the Products A. W. WEITKAMP

CONTENTS OF PREVIOUS VOLUMES

The Effects of Ionizing Radiation on Solid Catalysts ELLISON H. TAYLOR Organic Catalysis over Crystalline Aluminosilicates P. B. VENUTO AND P. s. LANDIS On the Transition Metal-Catalyzed Reactions of Norbornadiene and the Concept of n Complex Multicenter Processes G. N. SCHRAUZER

Volume 19

Modem State of the Multiplet Theory of Heterogeneous Catalysis A. A. BALANDIN The Polymerization of Olefins by Ziegler Catalysts AND M. N. BERGER,G. BOOCOCK, R. N. HAWARR Dynamic Methods for Characterization of Adsorptive Properties of Solid Catalysts L. POLINSKI AND L. NAPHTALI Enhanced Reactivity at Dislocations in Solids J. M. THOMAS Volume 20

349

JEANEUGENE GERMAIN AND MICHELBLANCHARD Molecular Orbital Symmetry Conservation in Transition Metal Catalysis FRANKD. MANGO Catalysis by Electron Donor- Acceptor Complexes KENZITAMARU Catalysis and Inhibition in Solutions of Synthetic Polymers and in Mieellar Solutions H. MORAWETZ Catalytic Activities of Thermal Polyanhydroa-Amino Acids DUANEL. ROHLFING AND SIDNEY W. Fox Volume 21

Kinetics of Adsorption and Desorption and the Elovich Equation AND F. c. TOMPKINS c . AHARONI Carbon Monoxide Adsorption on the Transition Metals R. R. FORD Discovery of Surface Phases by Low Energy Electron Diffraction (LEED) JOHNW. MAY Sorption, Diffusion, and Catalytic Reaction in Zeolites L. RIEKERT Adsorbed Atomic Species as Intermediates in Heterogeneous Catalysis CARLWAGNER

Chemisorptive and Catalytic Behavior of Chromia L. BURWELL, JR., GARYL. HALLER, ROBERT KATHLEEN C. TAYLOR, AND JOHNF. READ Volume 22 Correlation among Methods of Preparation of Solid Catalysts, Their Structures, and Hydrogenation and Isomerization over Zinc Catalytic Activity Oxide KIYOSHI MORIKAWA, TAKAYASU SHIRASAKI, R. J. KOKESAND A. L. DENT AND MASAHIDE OKADA Chemisorption Complexes and Their Role in Catalytic Reactions on Transition Metals Catalytic Research on Zeolites J. TURKEVICH Z. KNOR AND Y. ON0 Influence of Metal Particle Size in NickelCatalysis by Supported Metals on-Aerosil Catalysts on Surface Site DisM. BOUDART tribution, Catalytic Activity, and SelecCarbon Monoxide Oxidation and Related Reactions on a Highly Divided Nickel tivity R. VANHARDEVELD AND F. HARTOG Oxide Adsorption and Catalysis on Evaporated P. C. GRAVELLE AND S. J . TEICHNER Alloy Films Acid-Catalyzed Isomerization of Bicyclic R. L. Moss AND L. WHALLEY Olefins

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CONTENTS OF PREVIOUS VOLUMES

Heat-Flow Microcalorimetry and Its Applications to Heterogeneous Catalysis P. C. GRAVELLE Electron Spin Resonance in Catalysis JACKH. LUNSFORD

Volume 23 Metal Catalyzed Skeletal Reactions of Hydrocarbons J. R. ANDERSON Specificity in Catalytic Hydrogenolysis by Metals J . H. SINFELT The Chemisorption of Benzene P. B. MOYESAND P. B. WELLS The Electronic Theory of Photocatalytic Reactions on Semiconductors TH. WOLKENSTEIN Cycloamyloses as Catalysts DAVID W. GRIFFITHS AND MYRONL. BENDER Pi and Sigma Transition Metal Carbon Compounds as Catalysts for the Polymerization of Vinyl Monomers and Olefins D. G. H. BALLARD

Analysis of Thermal Desorption Data for Adsorption Studies MILO$ SMUTEK, SLAVOJ CERNL, AND FRANTISEK BUZEK

Volume 25 Application of Molecular Orbital Theory to Catalysis ROGERC. BAETZOLD The Stereochemistry of Hydrogenation of cQ-Unsaturated Ketones ROBERTL. AUGUSTINE Asymmetric Homogeneous Hydrogenation J. D. MORRISON,W. F. MASLER,AND M. K. NEUBERG Stereochemical Approaches to Mechanisms of Hydrocarbon Reactions on Metal Catalysts J. K. A. CLARKE AND J. J. ROONEY Specific Poisoning and Characterization of Catalytically Active Oxide Surfaces HELMUTKNOZINCER Metal-Catalyzed Oxidations of Organic Compounds in the Liquid Phase: A Mechanistic Approach ROGERA. SHELDON AND JAYK. KOCHI

Volume 24 Kinetics of Coupled Heterogeneous Catalytic Reactions L. BERANEK Catalysis for Motor Vehicle Emissions JAMESWEI The Metathesis of Unsaturated Hydrocarbons Catalyzed by Transition Metal Compounds J. C. MOLAND J. A. MOULIJN One-Component Catalysts for Polymerization of Olefins AND V. ZAKHAROV Yu. YERMAKOV The Economics of Catalytic Processes J. DEWING AND D. s. DAVIES Catalytic Reactivity of Hydrogen on Palladium and Nickel Hydride Phases W. PALCZEWSKA Laser Raman Spectroscopy and Its Application to the Study of Adsorbed Species R. P. CGQNEY, G. CURTHOYS, AND NGUYEN THETAM

Volume 26 Active Sites in Heterogeneous Catalysis G. A. SOMORJAI Surface Composition and Selectivity of Alloy Catalysts W. M. H. SACHTLER AND R. A. VAN SANTEN Mossbauer Spectroscopy Applications to Heterogeneous Catalysis AND HENIUKTOPS@E JAMBA. DUMESIC Compensation Effect in Heterogeneous Catalysis A. K. GALWEY Transition Metal-Catalyzed Reactions of Organic Halides with CO, Olefins, and Acetylenes R. F. HECK Manual of Symbols and Terminology for Physicochemical Quantities and UnitsAppendix 11 Part 11: Heterogeneous Catalysis

CONTENTS OF PREVIOUS VOLUMES

Volume 27 Electronics of Supported Catalysts GEORG-MARIA SCHWAB The Effect of a Magnetic Field on the Catalyzed Nondissociative Parahydrogen Conversion Rate P. W. SELWOOD Hysteresis and Periodic Activity Behavior in Catalytic Chemical Reaction Systems VLADIM~R HLAVA~EK AND JAROSLAV VOTRUB A Surface Acidity of Solid Catalysts H. A. BENESIAND B. H. C. WINQUIST Selective Oxidation of Propylene GEORGEW. KEULKS,L. DAVID KRENZKE, AND THOMAS N. NOTERMANN o-n Rearrangements and Their Role in Catalysis BARRYGOREWIT AND MINORUTSUTSUI Characterization of Molybdena Catalysts F. E. MASSOTH Poisoning of Automative Catalysts M. SHELEF,K. OTTO,AND N. C. OTTO

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Volume 29 Reaction Kinetics and Mechanism on Metal Single Crystal Surfaces ROBERTJ. MADIX Photoelectron Spectroscopy and Surface Chemistry M. W. ROBERTS Site Density and Entropy Criteria in Identifying Rate-Determining Steps in Solid-Catalyzed Reactions RUSSELL W. MAATMAN Organic Substituent Effects as Probes for the Mechanism of Surface Catalysis M. KRAUS Enzyme-like Synthetic Catalysts (Synzymes) G. P. ROYER Hydrogenolytic Behaviors of Asymmetric Diary lmethanes AND TADASHI KAWAI YASUOYAMAZAKI Metal-Catalyzed Cyclization Reactions of Hydrocarbons ZOLTAN PALL

Volume 30 Volume 28 Elementary Steps in the Catalytic Oxidation of Carbon Monoxide on Platinum Metals T. ENCELAND G . ERTL The Binding and Activation of Carbon Monoxide, Carbon Dioxide, and Nitric Oxide and Their Homogeneously Catalyzed Reactions AND DAN E. HENRICHARDE~SENBERG DRIKSEN

The Kinetics of Some Industrial Heterogeneous Catalytic Reactions M. I. TEMKIN Metal-Catalyzed Dehydrocyclization of Alkylaromatics SIGMUND M. CSICSERY Metalloenzyme Catalysis AND JOSEPHJ. VILLAFRANCA FRANKM. RAUSHEL

Mechanisms of Skeletal Isomerization of Hydrocarbons on Metals F. G. GAULT Tin-Antimony Oxide Catalysts FRANKJ. BERRY Selective Oxidation and Ammoxidation of Propylene by Heterogeneous Catalysis ROBERT K. GRASSELLI AND JAMES D. BURRINGTON Mechanism of Hydrocarbon Synthesis over Fischer-Tropsch Catalysts P. BILOENAND W. M. H. SACHTLER Surface Reactions and Selectivity in Electrocatalysis GEORGE P. SAKELLAROPOULOS Solvent and Structure Effects in Hydrogenation of Unsaturated Substances on Solid Catalysts AND VLASTIMIL R~~~IEKA LIBORCERVENL

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