ADVANCES IN CATALYSIS VOLUME 32
Advisory Board
M. BOUDART Stanford, California
M. CALVIN Berkeley, California
V. B...
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ADVANCES IN CATALYSIS VOLUME 32
Advisory Board
M. BOUDART Stanford, California
M. CALVIN Berkeley, California
V. B. KAZANSKY Moscow, U.S.S.R.
G . A. SOMORJAI Berkeley, California
P. H. EMMETT Portland, Oregon
A. OZAKI
G.-M. SCHWAB
Tokyo, Japan
Munich, Germany
R. UGO Milan, Italy
ADVANCES IN CATALYSIS VOLUME 32
Edited by
D.D. ELEY TIw Uniivrsity Nottinghatn. Englerncl
HERMANPINES Northit~t~.srern Uniiwsity Ei,trnsion,
Illinois
PAULB. WEISZ Mohil Rr.serirc,h crnd
DcJidoptnc,nr Cotportit ion Princeton. NeM' Jcwev
1983
ACADEMIC PRESS A Subsidiary of Harcourt Brace Jovanovich, Publishers
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COPYRIGHT @ 1983, BY ACADEMIC PRESS,INC. ALL RIGHTS RESERVED. NO PART OF THIS PUBLICATION MAY BE REPRODUCED OR TRANSMITTED IN ANY FORM O R BY ANY MEANS, ELECTRONIC OR MECHANICAL, INCLUDING PHOTOCOPY, RECORDING, OR ANY INFORMATION STORAGE AND RETRIEVAL SYSTEM, WITHOUT PERMISSION IN WRITINQ FROM THE PUBLISHER.
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United Kinadom Edition oitbhlred bv ACADEMiC PRESS, INC. ( L O N D O N ) LTD. 24/28 Oval Road, London NWI
1DX
LIBRARY OF CONGRESS CATALOG CARDNUMBER:49-7755 ISBN n-12-007832-5 PRINTED I N THE UNITED STATES O F AMERICA 83 84 85 86
9 8 7 6 5 4 3 2 1
Contents CONTKIBUK~KS. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . P u E r ~ r.t. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
iX
xi
Characterization and Reactivity of Molecular Oxygen Species on Oxide Surfaces M. CHLA N D A. J . TENCH
I. 11.
Ill. IV. V. VI. VII. VIII.
Introduction ............................................... Neutral Oxygen Species . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . characterization of Charged Dioxygen Species. . . . . . . . . . . . . . . . . . . . . . . . . . Formation and Stability of Charged Diatomic Species. . . . . . . . . . . . . . . . . . . . Oxygen Ions Containing More Than Two N Reactivity of Molecular Ions . . . . . . . . . . . . . The Relation of Mononuclear Surface Oxyg Spectroscopic and Catalysis Studies . . Comparison of Oxygen Species and Appendix A. Summary of g,, Value Appendix B. The Experimental "0 Diatomic Oxygen Species ( 0 2and ROO) . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . f Oxygen Species by Infrared Spectroscopy Appendix Reference ....................................... Note Added in Proof . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
2 3 X 36 X2 98
109 Ill 123 12x
130 134
148
Catalysis by Alloys in Hydrocarbon Reactions VLADlMlK PONEC
I. 11. 111. IV. V. VI.
Introduction. . . . . . . . . . . . . . . . . . . . . . . . . . . . . , . . . . . . . . . . . . . . . . . . . . . . . Alloys . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Particle Size Effects. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Mechanism of Hydrocarbon-Hydrogen Reactions . . . . . . . . . . . . . . . . . . . . . . . Hydrocarbon Reactionson Alloys . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Conclusions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
149
151 159 162
186 205 206
Modified Raney Nickel (MRNi) Catalyst: Heterogeneous Enantio-Differentiating (Asymmetric) Catalyst YOSHIHARU IZUMI 1. 11.
What Is MRNi? . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 215 History of Discovery and Development of MRNi.. . . . . . . . . . . . . . . . . . . . . . . 218 V
vi
CONTENTS
111. IV.
V.
v1. VII. VIII. IX. X.
Profile of MRNi in Hydrogenation.. , . . Profile of MRNi in Stereo-Differentation . . . . . . . . . . . . . . . . . . . . . . . ..................... Other Profiles . . . . . . , . . . . . . . . . . . . . . . . Surface Conditions. . , . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Mechanism of Enantio-Differentiation. .................... Characterization of Catalyst by Modify .................................. TA-NaBr-MRNi . . . Other Investigations. , . . . . . . . . . . . . . . . . . . . . . . . . . . . .................................. References . . . . . . . . .
224 229 248 249 254 262 264 267 269
Analysis of the Possible Mechanisms for a Catalytic Reaction System JOHN
I. 11.
Ill. IV.
V.
VI. VII.
HAPPEL A N D PETERH. SELLERS
..................................
274 278 General Formulas for Mechanisms and Reactions . . . . . . . . . . . . . . . . . . . . . . 283 287 A Procedure for Finding Every Direct Mechanism Systems with a Simple Overall Reaction.. . . . . . . . . . . . . . . . . . . , . . . , . . . . . . . 29 I 300 Overall Reactions with a Multiplicity Greater Than One.. . 317 Discussion . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 320 List of Symbols . . . . . 32 1 References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
.
Homogeneous Catalytic Hydrogenation of Carbon Monoxide: Ethylene Glycol and Ethanol from Synthesis Gas B. D. DOMBEK 1.
II. 111.
1V.
V.
VI. VII.
Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 326 Cobalt Catalysts.. . . . . . . . . . . . . . . . . . . . . . . . Rhodium Catalysts. . . . . . . . . . . . . . . . . . Unpromoted and Carboxylic Acid-Promoted Ruthenium Catalysts. . . Lewis Base-Promoted Ruthenium Cata Other Catalysts.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Conclusions . . . . . . . . . . . . . . . . . 408 References .......................................................... 410
Cyclodextrins and Cyclophanes as Enzyme Models IWAOTABUSHI AND YASUHISAKURODA 1. 11.
Introduction.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 417 Basic Principles of Molecular Recognition . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 420
CONTENTS 111 .
IV . V.
vii
Enhancement of Binding and Catalysis by Host Design . . . . . . . . . . . . . . . . . . 436 Enhancement of Binding and Catalysis by Guest Design . . . . . . . . . . . . . . . . . 456 Conclusion . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 461 462 References and Notes . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
AUTHOR IN1)EX
...............................................................
INDEX . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . SUBJECT OF P~iiviousVOLUMES .............................................. CONTENTS
467 494 509
This Page Intentionally Left Blank
Contributors Numbers in purentheses indicute the puges on which the authors' contributions begin.
M. CHE,Laboratoire de Chimie des Solides, ER 133, C N R S , UniversitP Pierre et Marie Curie (Paris V l ) , 75230 Paris Cedex 05, France ( 1 )
B . D. DOMBEK,Union Carbide Corporation, South Charleston, West Virginia 25303 (325) JOHN HAPPEL,Department of Chemical Engineering and Applied Chemistry, Columbia University, New York, New York 10027 (273) YOSHIHARUIZUMI, Institute for Protein Research, Osaka University, 3-2 Yamadaoka, Suita, Osaka 565, Japan (215) YASUHISAKURODA,Department of Synthetic Chemistry, Kyoto University, Kyoto 606, Japan (417) VLADIMIR PONEC,Gorlaeus Laboratoria, Rijksuniversiteit Leiden, 2300 R A Leiden, The Netherlands (149) PETERH . SELLERS,The Rockefeller University, New York, New York I0021 (273) IWAO TABUSHI,Department of Synthetic Chemistry, Kyoto University, Kyoto 606, Japun (417) A. J . TENCH.*Chemistry Division, Atomic Energy Research Establishment, Harwell, Oxfordshire OX11 ORA, United Kingdom ( 1 )
*Deceased.
ix
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Catalysis embraces a variety of fields, and it is the aim of the Editors to see that each volume of Advances in Catalysis contains articles spanning a broad spectrum of interest. The opening contribution, by M. Che and A. J . Tench, is a survey of work on adsorbed molecular species and their role in oxidation reactions. This article, together with its companion piece in Volume 31 by the same authors, will stand as a memorial to A. J . Tench, who died on March 17, 1983 from Hodgkin’s disease, two days after submitting the contribution presented here. The importance of catalysis by alloys is well recognized in the petrochemical industry. By means of alloying. dramatic changes can be achieved in the stability and selectivity of metal catalysts. The last decade has witnessed a renaissance in alloy research, and the review by V. Ponec gives a comprehensive survey of this active field. Great strides in developing highly selective enantio-differentiating (asymmetric) catalysts have been made by modifying Raney nickel. Y. Izumi. a pioneer in this area of endeavor, surveys this field. Catalytic reactions proceed through a network of intermediates that are connected by elementary reactions. To explain a catalytic reaction it is necessary to consider how steps may be combined in appropriate proportions. The article by J . Happel and P. H . Sellers reviews methods to achieve it. Hydrogenation of carbon monoxide by heterogeneous catalysts has been studied for decades: it was surveyed in Volume I of this publication. The use of homogeneous catalysts for this type of reaction is, however, of a more recent vintage, and opens new synthetic feasibilities. Conversion of carbon monoxide to two carbon atom compounds is reviewed by B. D. Dombek. Cyclodextrins, also called cycloamylases, doughnut-shaped oligosaccharides, have attracted much attention as enzyme models. Although this area of research was surveyed in Volume 23, much subsequent progress in this field through multifunctionalization of cyclodextrin necessitates a new review. This contribution was written by I . Tabushi and Y. Kuroda. active researchers in this area.
HERMANPINES xi
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ADVANCES I N CATALYSIS. VOLUME 32
Cha racte r izat io n a nd Reactivity of Molecular Oxygen Species on Oxide Surfaces M . CHE
.
Laboratoire de Chimie des Solides ER 133. CNRS UniuersirC Pierre et Marie Curie (Paris V I ) Paris. France AND
A . J . TENCH* Chemistry Division Atomic Energy Research Establishment Harwrll. Oxjordshire. United Kingdom
I . Introduction . . . . . . . . . . . . I1 . Neutral Oxygen Species . . . . . . . . . A . Triplet Oxygen . . . . . . . . . . B . Singlet Oxygen . . . . . . . . . . I11. Characterization of Charged Dioxygen Species . . A . TheO; Ion . . . . . . . . . . . B . The 0: Ion . . . . . . . . . . . C. The 0:-Ion . . . . . . . . . . . D . The 0:- Ion . . . . . . . . . . . IV . Formation and Stability of Charged Diatomic Species A . Ionic Oxides . . . . . . . . . . . B . Transition Metal Oxides . . . . . . . C. Aluminosilicates . . . . . . . . . . D . Supported Metals . . . . . . . . . E . Dioxygen Adducts . . . . . . . . . V . Oxygen Ions Containing More Than Two Nuclei . . A . The 0 ; Ion . . . . . . . . . . . B . The 0;Ion . . . . . . . . . . . VI . Reactivity of Molecular Ions . . . . . . . A . Exchange Reactions . . . . . . . . . B. Oxidation Reactions . . . . . . . . C . Photo-Induced Reactivity . . . . . . .
. . . . . . . . . . . . . . . . . . . . . .
. . . . . . . . . . . . . . . . . . . . . .
. . . . . . . . . . . . . . . . . . . . . .
. . . . . . . . . . . . . . . . . . . . . .
. 2 . 3 . 3 . 6 . 8 . 10 . 33 . 34 . 35 . 36 . 36 . 44 . 51 . 74 . 78 . 82 . 82 . 95 . 98 . 98 . 100 . 105
Deceased March 1983.
I
.
Copyright 0 1983 by Academic Press Inc . All rights of reproduction in any form reserved . ISBN 0-12-007832-5
2
M. CHE AND A. J . TENCH VII. The Relation of Mononuclear Surface Oxygen Species to Electron Spectroscopic and Catalysis Studies . . . . VIII. Comparison of Oxygen Species and Their Role in Catalytic Reactions . . . . . . . . . . . A. Charicterization. . . . . . . . . . . . B. Reactivity . . . . . . . . . . . . . C. Future Directions . . . . . . . . . . . D. Conclusions . . . . . . . . . . . . . Appendix A. Summary of gzz Values for 0;on Surfaces . Appendix B. The Experimental "0 Hyperfine Parameters (in gauss) of Diatomic Oxygen Species (0; and ROO') . . . . . . . . . Appendix C. Characterization of Oxygen Species by Infrared Spectroscopy . . . . . . References . . . . . . . . . . . . . . . Note Added in Proof. . . . . . . . . . . .
1.
. . .
109
. . . Ill . . . Ill
. . . .
.
. . . . . . .
.
116
. 121 123 123
. 128 I30 I34
.
148
Introduction
Oxidation and oxidative dehydrogenation reactions over oxide catalysts have been widely studied in recent years. The precise role of oxygen in these reactions remains elusive, but slowly a more detailed picture is emerging which suggests that both oxide ions of the lattice and oxygen species on the surface can play an important role (1,2). The surface oxygen species can conveniently be divided into two broad classes, i.e., mononuclear and molecular. The mononuclear species such as 0-, 0;;(lattice ions in low coordination), and M=O have recently been reviewed by Che and Tench (1).These are now well characterized and their role in simple and in some more complicated reactions is now better understood. Molecular oxygen species are also formed on the surface and there has been considerable progress since these were last reviewed by Lunsford (3).The characterization of the adsorbed species has improved markedly as isotopic labeling with "0 has become more widely used. Some novel forms of molecular oxygen species have been reported and, in particular, the reactivities of species such as 0; and 0; have been studied. Molecular oxygen species have been also identified as intermediates in some biological reactions and are important as oxygen adducts in natural and artificial oxygen carriers ( 4 , 5 ) . The purpose of this review is to survey the work on adsorbed molecular oxygen species and to show how recent developments point the way toward an understanding of the role that they and the mononuclear forms of oxygen may play in oxidation reactions. The coverage is restricted to those papers where there is direct evidence on the nature of the oxygen species concerned.
MOLECULAR OXYGEN SPECIES ON OXIDE SURFACES
II.
3
Neutral Oxygen Species
The ground state of the oxygen molecule is a triplet 'C,- state with two unpaired electrons, and at slightly higher energy there are two low-lying electronically excited states, the singlet 'A, and 'C,' levels (6) (Fig. 1). A. TRIPLET OXYGEN The electron paramagnetic resonance (EPR) spectrum of ground-state oxygen has been characterized both in the gas phase and in the solid state. In the gaseous state, the coupling of the spin angular momentum with the end-over-end molecular rotation angular momentum gives rise to an EPR spectrum with many lines covering more than 10 kG (7-9). In the solid state, oxygen has been observed as an impurity in solid N,, CO, Ar, and CD, (10-12). The lines are very broad with g1 = 2.02, gll 0.7, and a zero field splitting of 108 GHz (11).Kon ( 1 2 4 observed an isotope effect between the EPR spectra of l 6 0 l 6 Oand ' s O ' 8 0 ( S = 1) at temperatures below 10 K, which was explained in terms of torsional oscillation of 0, in the matrix around the equilibrium position. In similar work, 0, molecules trapped in NaClO, and KClO, single crystals show well-resolved isotope shifts (I2b). Model calculations reveal that the discrepancy between the spin Hamiltonians of the trapped and the free molecule originates from the angular librations of the trapped molecule.
-
FIG.1 . n, orbital occupancy and energies ( 6 ) of triplet and singlet dioxygen.
4
M. CHE AND A. J. TENCH
It might be expected that effects due to adsorption on the surface would be observable as a change in the EPR spectrum of gas-phase oxygen. Clarkson and Turkevich (13)have adopted this approach and used the variation in linewidth of one of the EPR lines of gas-phase oxygen to follow the adsorption of oxygen on porous Vycor glass at 78.3 K. A best overall fit was obtained using a BET plot, but at low pressures there was evidence for an initial adsorption step of higher energy. Takaishi et a/. (14)have shown that oxygen adsorbed on mordenite gave the same spectrum as that of gaseous oxygen and was freely rotating and translating above 200 K, whereas Seymour and Wood (15) reported two new EPR lines for oxygen adsorbed on carbon corresponding to physisorbed and chemisorbed oxygen. More recently, Lemke and Haneman (16)have carried out a careful study and shown that no changes were observed in the free oxygen lines when crushed silicon was exposed to oxygen in an ultra-high-vacuum (UHV) system at 100 K, although it is known to adsorb on the surface. No new lines were observed and no oxygen lines remained after the system had been evacuated, while maintaining the silicon sample at 100 K. The authors suggest that adsorption on the surface will split the ni and energy levels, and the two electrons would pair up in the lower of the two states to give a nonparamagnetic adsorbed state (Fig. 1). For weak interactions, the splitting may not be large enough for the electrons to be paired all the time and an EPR signal may be observed which decreases as the interaction increases. It should be kept in mind that the crushed silicon used in the work of Lemke and Haneman (16)is relatively low-surface-area material compared to the catalyst supports commonly used and this may account for the different results. The broadening of an EPR signal from a surface species when exposed to oxygen has been known for many years (17). Such broadening is brought about by the magnetic interaction occurring on collision of oxygen with the surface species. Because of the exchange interaction, such broadening of the EPR spectrum is commonly referred to as exchange broadening and is also known in solution (180). Busca (18b) has evaluated the literature values for infrared bands attributed to coordinated and adsorbed dioxygen species. He concludes that it is very difficult to deduce the nature of the dioxygen coordination from measurements of the frequency shift, Avo,, with respect to the stretching frequency of the free molecule. It needs to be stressed that it is also difficult to distinguish between mononuclear and molecular species from measurements of voo, and this can only be achieved by careful interpretation of experiisotopic mixtures. The absence of such experiments ments using 160/'80 very often accounts for the conflicting attributions in the literature which are discussed in later sections. Griffiths et al. ( 1 9 4 have investigated the adsorption of oxygen on cr-Fe,O,
MOLECULAR OXYGEN SPECIES ON OXIDE SURFACES
5
and observed infrared (IR) bands at 1350 and 1270 cm-', which they assigned to 0, and 0; species. In a later report, Al-Mashta et al. (19b) have reinvestigated the same system and have suggested that both bands at 1350 and 1270 cm-' should be reassigned to perturbed 0; species as discussed in Section III,A,A Davydov et a/.(20a,b)have reported low-intensity absorption bands in the IR on adsorbing oxygen on a range of high-surface-area systems, including TiO,, SnO,, V,O,/SnO,, MoO,/AI,O,, MoO,/MgO, and NiY zeolite. They suggest that bands in the range 1600-1700 cm-' can be attributed to adsorbed molecular oxygen in a neutral state, possibly as singlet oxygen, and the IR transitions being allowed, due to coupling with the lattice. Many likely impurities, such as water and oxides of carbon and nitrogen, absorb in this same region, but the authors argue that these may be eliminated because other associated IR bands are not observed. For comparison, Raman measurements on triplet oxygen gas show a band at 1555 cm-' (21) and it is not clear why such a band would be induced to move to higher energies when oxygen is adsorbed at the surface. According to Al-Mashta et al. (196),a possible explanation is the partial electron withdrawal from the antibonding orbitals of the oxygen by a polarization induced by the cations which act as the adsorption sites. In fact, this explanation is difficult to understand since the polarization arises from a purely electrostatic effect of the charge on the cation whereas the electron withdrawal from antibonding orbitals will be related to the availability of suitable empty orbitals on the cation. Eberhardt et al. ( 2 2 4 have studied the photoemission of oxygen physisorbed on graphite at 10 K. The photoemission spectra exhibit vibrational structure in the 27c band. From calculations based on Franck-Condon factors, the authors conclude that on the graphite surface the equilibrium distance of the oxygen nuclei is decreased by 0.065A relative to the gas phase. This would also be consistent with a partial electron withdrawal from oxygen antibonding orbitals into available orbitals in the graphite. Long and Ewing (22b) have reported IR evidence for the formation of bound (O,), dimers in the gas phase at 90 K characterized by two narrow bands at 1586.1 and 1596.6 cm-' superimposed on the broad collisioninduced IR spectrum of oxygen. The energy of formation of the dimer was found to be -530 2 70 cal/mol, indicating a van der Waals type complex. Dimerization of oxygen to form O4 has been reported (22c,d)on y-alumina. Magnetic susceptibility studies show a significant decrease (about 25%) in paramagnetic susceptibility of oxygen at 77 K over a small pressure change. This is taken as evidence for a dimerization equilibrium on the surface 2 0 , * 0,
with very weak bonding between the oxygen molecules. Making use of
6
M. ('HE AND
A. J . TENCH
parallel observations of 0; labeled with "0 and of gas-phase oxygen labeled with "0, Tanaka and Kazusaka (22e)postulate that 0, is the intermediate for the homomolecular oxygen exchange reaction on ZnO at 77 K. Anufrienko et ul. (22f) suggest that 0, complexes can also be formed on SnO, at temperatures between 100 and 130 K .
B. SINGLET OXYGEN The possibility that singlet ( 'Ag) oxygen could play a role in reactions at oxide surfaces has not been considered seriously, because the energy level is 22.64 kcal above that of ground-state triplet oxygen. The EPR spectrum of ('A,) oxygen in the gas phase has been investigated by Miller (241); Wilkinson and Brummer (24b) have collected rate constants for the decay and reactions of singlet oxygen in solution. Kearns (23a)has suggested that decomposition of 0; might yield singlet oxygen, and Khan (23b) has observed the ( ' A g ) 0 2 emission spectrum at 1.29 pm from the reaction of KO, with water. The reactions of singlet oxygen with organic molecules have recently been reviewed (24c and references therein) and the study of this chemistry is made possible due to the lifetime of the singlet states (since transition to the triplet ground state is forbidden). Tsyganenko et al. (24d,e) have investigated the low-temperature adsorption of oxygen on NiO and C r 2 0 3 using 160/'80 mixed isotopes to check the presence of two oxygen atoms in the surface species. They detected IR bands at 1500 cm-' on NiO and 1460 cm-' on C r z 0 3 ,which they assigned to singlet oxygen because of the closer proximity of the bands to the frequency of gas-phase singlet oxygen at 1483 cm-' (241') than to that of gas-phase triplet oxygen at 1555 cm-' (21).There are a number of factors which can influence the voo frequency of adsorbed oxygen, as discussed in Appendix C. This assignment needs to be verified using reactions specific to singlet oxygen. Recently, Slawson and Adamson have shown (25) that films of linolenic acid on silica gel undergo an autoxidation which is accompanied by a chemiluminescence. The emission spectrum contains two components and the low-energy component close to 630 nm is attributed to bimolecular reaction of two ('Ag) oxygen molecules. In a subsequent paper, Slawson el al. (26a)report that heating 2,Sdiphenylfuran in air on a silica or titanium dioxide surface results in its conversion to cis-dibenzoylethylene, which is characteristic of reaction with singlet oxygen in the homogeneous phase. A singlet oxygen quencher inhibited the reaction which was not affected by a free radical scavenger. The authors suggest that for adsorbed oxygen, the
MOLECULAR OXYGEN SPECIES ON OXIDE SURFACES
7
normal orbital degeneracy is removed and the singlet configuration may become the one of lowest energy. This is similar to the suggestion of Lemke and Haneman (16)on the state of adsorbed oxygen. It has been proposed that singlet oxygen is formed from 0; on transition metal oxides (26b) and is the active form of oxygen which interacts with olefins. Dmuchovsky et al. (26c) have proposed that singlet oxygen is important in the oxidation of benzene to maleic anhydride over vanadiamolybdena catalysts, whereas Khan (26d) has reported singlet oxygen is formed on hot tungsten filaments. Lipatkina et al. (26e) report an unusual EPR signal with g1 = 1.95 and g l l = 1.92 when oxygen is adsorbed on chromium oxide catalysts containing Cr5 ions. They suggest that a surface complex [ C r 5 + 0 2 ] is formed since the EPR spectrum is characteristic of Cr5+,indicating that the oxygen must be in the singlet state. The presence of oxygen in the complex was confirmed by "0 labeling; the change in line shape was thought to be consistent with a total hyperfine interactions of 10-15 G, which can be compared with the 140-150 G observed for 0; (see Section III,A,2). Guillory and Shiblom (26f) have used the reaction of rubrene to form an endoperoxide to detect the presence of singlet oxygen formed in a flowing gas stream over a range of catalysts. Positive results were obtained only with a lithium-tin-phosphorus catalyst, but the results were irreproducible. Munuera et al. (269) have adopted an approach based on the use of chlorinated TiO, to produce singlet oxygen on surfaces. Chloride'ions on the surface of these samples are thought to be transformed into C10- by ultraviolet (UV) irradiation in the presence of oxygen (26h). If the TiO, surface is highly hydroxylated, the irradiation also produces H,O, (26h), which further reacts with the C10- ions in the classical reaction seen in aqueous solution (26i)to form singlet oxygen: +
(CIO-),
+ H202+H20 + Cl; + ( ' A g ) 0 2
The presence of singlet oxygen was shown by a specific reaction with a sulfonic acid (269).It might be expected that surface oxygen would show the same reaction chemistry as singlet oxygen does in homogeneous media. The proposals that singlet oxygen is involved in heterogeneous catalytic reactions have not yet been explored fully. The methods used for the detection of the excited singlet state of oxygen need to be improved, and an approach based on the detection of the emission from ( 'Ag)O, as observed by Khan (23b) will present a significant advance if it can be applied to heterogeneous systems. It is clear that more quantitative work is required and in this respect the evidence that 0; can react with water to form ( ' A e ) 0 2 (23b)could have considerable mechanistic interest.
M. CHE AND A. J. TENCH
8 111.
Characterization of Charged Dioxygen Species
Several kinds of charged dioxygen species have been reported on surfaces, including O l , O , , O:-, and O ; - . All of these, with the exception of Oz-, would be expected to be paramagnetic and to give an EPR signal. In addition, the optical and IR absorption bands are known for some of these species and can also be used for characterization (Appendix C). Table I summarizes the properties of the dioxygen species relevant to this paragraph, whereas Table I1 is concerned with the thermodynamics of processes involving dioxygen species. Since no values are available for the species adsorbed on the surface, we have given the gas-phase values in Table I1 and these should be taken only as a general guide. Inspection of Table I shows that the dioxygen bond length becomes progressively larger on going from 0;to 0;- and this increase is accompanied by a decrease in the dioxygen bond strength. These facts can be explained
TABLE I Properties of’ Dioxygen Species“ ~~
~
Db ‘0-0
Species
Example
(4
O,PtF, (27) Gas Gas Gas LiO, NaO, KO2 HO2 Gas NazO, Rb,OZ BaO, H,OZ ROOR‘
1.17 (27) 1.123 (28) 1.207 (28) 1.216 (28) I .33 ( 4 ) 1.33 (29) 1.28 (30) 1.3 (3Ia) 1.34 (32) 1.49 (33a) 1.54 (336) 1.49 (30) I .49 (34)
(kcal/mol)’
(kJ/mol)’
149 (28)
623 490
I I7 (28) 95 (28)
64 (31b)
268
49
Bond ordeP 2.5 2.5 2 2 1.5 I .5 I .5 I .5 1.5 1 1.1
1.1
51 (34) 38 (35)
213 159
1.1 1.1 0.5
* References appear in parentheses.
D denotes dissociation energy. Conversion factors used are as follows: 1 eV = 23.060 kcal/mol, 1 cal = 4.184 J . Defined as N = (n - n*)/2, where n and n* are the numbers of electrons in the bonding and antibonding molecular orbitals, respectively, of the corresponding dioxygen species ‘ R = alkyl.
MOLECULAR OXYGEN SPECIES ON OXIDE SURFACES
9
TABLE I1 Entlmlpy of Processes Involving Dioxygen Species in rhe Gas Phase" AHb Process
I. 2. 3. 4. 5. 6. 7. 8. 9. 10.
11. 12. 13. 14.
(kcal/mol)
0, + e - -0; 0, + 2e- -0;0; + e- -0;0, 0; + e0, - 2 0 0; -0 + 00;- - 2 0 20; -0, + 0:0; + e - - 2 0 0, + e - -0- + 0 0, + 2e- - 2 0 0; -0 + o+ 'Zgg0, ' A g0, 'Zg- 0, ' Z l 0,
- 10.15
154.5 164.65 278.45 118 94.37 - 104.06 174.8 60.59 83.02 50.44 153.63 22.64 31.73
-
--
(kJ/mol)
Ref.
-42.47 646.4 688.9 1165.0 493.7 394.8 -435.4 731.4 253.5 347.4 21 1.0 642.8 94.7 157.9
32 36 c
37 28
d d e d 40 d
.f 6 6
All reactants and products (except those of processes 13 and 14) are assumed to be in the ground state. Conversion factors used are 1 eV = 23.060 kcal/mol, lo3 c m - ' = 2.859 kcal/mol (1 c m - ' = 1.2398.10-4eV) when original values are given in eV or cm- and I cal = 4.184 J. ' Calculated from processes I and 2. Calculated using 0 + e - +O-, A H = - 33.78 kcal/mol, as given in Ref. 38, and a thermochemical cycle, a s described by Tuck (39). Calculated from a thermochemical cycle involving processes I and 2. Calculated using 0 -0' + e - , A H = 314.08 kcal/mol, as given in Ref. 41, and a thermochemical cycle, similar to that described by Tuck (39) for 0; -0 + 0 - ; the value 153.63 kcal/mol is to be compared with the spectroscopic value of 149 kcal/mol (28).
',
by reference to the molecular orbital energy level diagram (Fig. 2 ) . In the case of O,, the gg and nu bonding orbitals are fully occupied and the two additional electrons reside in the degenerate ng antibonding orbitals, giving a bond order of 2. Removal of an antibonding electron from 0, to give 0: will increase the bond order to 2.5 and lead to a shortening of the 0-0 bond, while the formation of 0; and 0;- from 0, requires that electrons be added to the antibonding orbitals, leading to a decrease in bond order and a lengthening of the 0-0 bond. There is no data available for O:-,
M. CHE AND
10
?:E : 0
A. J. TENCH
7G r;
+
e-
0,
+e-
r::
0;
*:
FIG.2. The simplified energy level diagram for 0:. 0 , ,and 0 ; in their ground state. When a crystal field is present, the n, and nu levels are not degenerate.
but this species would be expected to have a bond order of 0.5 and a weak 0- 0 bond. Of the reactions listed in Table 11, the only process that leads to a decrease of the energy of molecular oxygen is the formation of the free superoxide ion, 0; ( - 10.15 kcal/mol). The superoxide ion would therefore be expected to be the dioxygen species most commonly formed on oxide surfaces and in fact it is the species most studied, both in the bulk of various matrices and on surfaces. The other species (0; and 0;-)are not stable in the gas phase, although they can be stabilized in the solid state (Table I) due to the additional coulombic stabilization from the lattice. Nearly all the data in the literature refers to the characterization of 0; on various surfaces and this is discussed in detail in the following sections. A. THEO; ION
By far the most commonly reported species on oxides is the 0; (superoxide) ion, which has been characterized mostly by EPR using the g, hyperfine, and superhyperfine tensors. The EPR signals have only been seen in the case of 0; adsorbed on nonparamagnetic ions since if the ion at the adsorption site is paramagnetic, there will be a strong interaction between the unpaired electrons leading to line broadening. The absence of an EPR signal does not necessarily mean that the oxygen is in a nonparamagnetic form such as 0 2 - as assumed by some authors. In these situations other techniques
MOLECULAR OXYGEN SPECIES O N OXIDE SURFACES
11
such as IR. although of much lower sensitivity, become the major source of information. The usually accepted approach is to adopt an ionic model for the superoxide ion on the surface. In this model, an electron is transferred from the surface to the oxygen to form O;, and there is an electrostatic interaction between the cation at the adsorption site and the superoxide ion. A calculation of the CJ tensor based on this model (Section III,A,I) accounts for nearly all the data from adsorbed 0; and is consistent with the evidence that the spin density on both oxygen nuclei is the same (Section III,A,2). However, there are examples of oxygen adsorbed on the surface where the g values do not fit the predictions of the ionic model (Section IV,E) and also a few cases where the spin density on the two oxygen nuclei is found to be different. In these situations it seems likely that a covalent model in which a 0 bond is formed between the cation and the adsorbed oxygen, is more relevant. These two approaches are considered in the following sections. 1.
The g Tensor
a. The Ionic Model. The 0; ion is formed by adding an electron to one of the degenerate x g orbitals of the oxygen molecule to give the electron ~ ( with a 'n ground configuration ( I aJ2( 1a,)2(2a,)2(2a,)2(3a,)2( l ~ , ) lzJ3 state. Interaction of the free ion with the matrix either in the bulk or on the surface removes the degeneracy of the highest occupied x g orbital, splitting it into two components with a separation A (Fig. 2). Kanzig and Cohen (42) have derived theoretical expressions [Eqs. (1)-(3)] for the y tensor of 0; assuming an ionic model :
where the x axis is chosen along the ng orbital containing the unpaired electron and z is along the internuclear axis. To prevent ambiguity, all the results discussed in this review are presented with this convention. The energy level separations A and E are defined in Fig. 2, and 1 is the spinorbit coupling constant of oxygen, generally assumed to be 135 cm-' (4.3~). The parameter 1 is a correction to the angular momentum about z caused by the crystal field and is normally found to be close to unity.
12
M. CHE A N D A. J . TENCH
The properties of these equations for the g tensor can be seen more clearly if they are simplified by assuming I = 1,1 < A )-A1 AlSb GaAs
GaAs Co ammonia adducts in Y zeolite Co amine adducts in Y zeolite
s*z
g,,
2.038 2.030 2.044 2.038 2.033 2.024 2.023 2.022 2.01 76 2.017
2.009 2.009 2.009 2.009 2.008 2.01 I 2.01 I 2.01 1 2.0105 2.010
2.010 2.009 ? ? 2.0279 2.0089 2.038 2.006 2.040 2.008 2.041 2.005 2.035 2.007 2.046 2.009 2.084 2.000 2.017 2.039
{ :::::
2.01 1.998
A*, (G)
A,, (G)
6.5 2.003 5.7 2.002 2.005 12 6.5 2.003 ? 2.002 ? 2.003 9.7 2.004 ? 9.6 2.0050 2.0 1.2 2.004
4.8 4.4 9 4.7 7.4 ? 6.8 6.9 1.9 1.8
g xx
? 5.4 37 32 3.04 ? 4.9 6.4
2.004 2.004 1.987 1.993 2.0041 2.006 2.004 2.002 2.004 2.006 2.000
24.5" 17.8
2.00 1.992
17.8 20
15
3.3 3.6 ?
? 4.84 3.6 3.8 4.4 25 39" 12.5
12 10
A,,
(G)
Nucleus
5.7
103 i03 103
5.1
8 5.7 6.1 15 5.9 ? 1.0
104 105a 1056
106 107
108 109
21
? 5.2 17.5 15.0 13.2 ? 3.7 4.4 2.5 10" 12.5
12.5 13
Ref.
108 108 110 56
111 112 47 113 114 51
cyco Y o
115 116
" Note g and A tensors do not have the same principal axes-consult original article
(108)and y5M003/Si0,(108,109) (Fig. 10). On the former system, formation
of the ion at 77 K gives an EPR signal with gz, = 2.017 and a resolved set of six hyperfine lines about qyy,whereas a sample that is not enriched shows no superhyperfine structure. This is consistent with adsorption at a Mo6+ site. At 300 K, a new 0; signal appears on both samples with gzz = 2.039 and a superhyperfine structure of A,, = 5.4,Ay,v = 3.6, and A,, = 5.2 G. This corresponds to adsorption at an A13+ site. On "Mo03/Mg0, there is no indication of Mo6+ sites available either at 77 K or at 300 K. The 0; ions are adsorbed on Mg2+ sites (108). The superhyperfine tensor has also been used to derive the amount of spin delocalization on the cation leading to the superhyperfine structure. In view of what has been said above on the origin of the superhyperfine interaction, the result must be handled with caution. Thus, the unpaired electron
30
M. CHE AND
A.
J. TENCH
i 10 Oe
42
1 I Ill I I 91
I I Ill I I FIG.10. The EPR spectrum of 0;ion on 9SMo0,/Si0, at 77 K showing the superhyperfine interaction with the Mo ion ( I O Y ) .
of 0; is 5% localized in the Al atom orbital on AlSb (47)and 23% localized in the Ga atom orbital on GaAs (113, 114), whereas on oxide surfaces the figures are generally smaller (118). These latter figures are consistent with the spin densities obtained from the "0 hyperfine tensor (Tables I11 and IV). In addition to the nature of the cation at the adsorption site, the superhyperfine tensor can also give information on neighboring atoms further away. For example, 0; adsorbed on MgO exhibits a superhyperfine tensor ascribed to the presence of a nearby proton, presumably as a hydroxyl group (68) and this has been confirmed by isotopic labeling with deuterium (see Section IV,A). Superhyperfine tensors indicating the presence of nearby protons have also been reported for 0;adsorbed on ferrocene deposited on porous Vycor glass (PVG) (120)and for alkylperoxy radicals supported on TiO, (90). The information which is obtained from the superhyperfine tensor is important and much effort has been aimed at obtaining this parameter.
MOLECULAR OXYGEN SPECIES ON OXIDE SURFACES
31
Where naturally occurring isotopes of the cation are not suitable, then enrichment of the surface cations with nonzero-nuclear-spin isotopes is a powerful technique, e.g., 9 5 M 0 0 3 on A1203, S i 0 2 , or MgO (108). Often, the presence of a superhyperfine interaction leads to spectra of low resolution and it is useful to increase both the intensity of the spectra and the resolution. Second or even higher derivative spectra can be used to enhance the resolution and in some cases secondary reactions (121)have been shown to increase the intensity.
4. Optical Properties An absorption band in the UV results from a transition between the o g and the 7cg orbitals of 0; (Fig. 2). For 0; in the alkali halides, this band is centered at about 5 eV ( f 2 2 , 123) and detailed measurements of the stress dependence have been carried out (69h). A yellow luminescence in alkali halides has been known for many years (124) but was not identified as originating from 0; until 1961 (122). At low temperature, the emission spectrum in both alkali halides (125)and sodalites (126)shows a number of sharp zero phonon transitions between 400 and 600 nm; the spacing of these lines corresponds to the ground-state vibrational frequency of the 0; ion. The Ag'O; complex has been observed (127a)in a matrix experiment and gives an absorption at 275 nm. Although well documented in the solid state, no optical absorption or luminescent spectra have been reported for 0; on an oxide surface even where EPR has shown the ion to be present on, for example, zeolites (127h) or the alkaline-earth oxides (128, 129). This may arise because absorption from surface oxide ions in low coordination occurs at about the same energy as the optical absorption for 0; in the oxides such as MgO (f30,131). The 0; ion is not normally expected to be active in the IR but laser Raman studies on the crystalline alkali metal superoxides have led to the assignment of frequencies between 1137 and 1164 cm-' to the 0-0 stretching vibration (21, 132). A number of 0; complexes with the alkali metals (133a, b, c) and transition metals (127a and references therein) have been studied as matrix isolated species. There are no observations of 0; on the surface using Raman spectroscopy but there are now several reports on 0; by IR spectroscopy. Davydov et al. (20h) have reported a band at 1180 cm-' on TiO,, which they assign to a molecular species such as 0;. If correct, this means that the surface must perturb the adsorbed oxygen sufficiently to make the molecular ion infrared active. This is unexpected, since the EPR data (Section III,A,2,a) show that the oxygen nuclei are equivalent in this system (75); however, it is possible that different species
32
M. CHE AND A. J . TENCH
are being observed. Conflicting results have been also obtained for oxygen adsorbed on cr-Fe,O,. Griffiths et al. (19a) observed bands at 1350 and 1270 cm-' and assigned them to adsorbed 0, and 0; species, respectively. On heating Fe,O, in oxygen, Davydov et al. (134a) observed two strong bands at 965 and 918 cm-' and assigned them, together with weaker bands at 890, 835, and 797 cm-', to the vibrations of bonds between the surface cation and oxygen produced as a result of dissociative adsorption of oxygen on cation sites of different coordination. Al-Mashta ef al. (19b) have reconsidered the case of Fe,O, and tried to rationalize the previous results on the following basis. As indicated in Appendix C, dioxygen species are known to absorb at 1550 cm-' (O,), at ca. 1150 cm-' (O;), and at ca. 800 cm-' (O:-). These wavenumbers relate to formal bond orders of 2, 1.5, and I , respectively (Table I), but intermediate situations are possible (134b).Al-Mashta et al. (19b)have suggested that such species perturbed by the strong electrical forces of the quasi-ionic solid would give bands of higher wavenumbers than the above values, due to partial electron withdrawal from antibonding orbitals. In this discussion, the possibility of backbonding from metal orbitals to the antibonding orbitals of oxygen (134c),which will tend to decrease the voo frequency of the dioxygen species, has been neglected. They have suggested that all reported bands on m-Fe2O3 between 1350 and 1250cm-' should be assigned to a perturbed 0; species, intermediate between 0, and O;, and absorption between 1100 and 900 cm-' to perturbed 0:- species, intermediate between 0;and 0;-. However, mononuclear species such as M=O also absorb in the region I 100 to 900 cm-' (Appendix C) and the assignment needs to be confirmed by '60/'80 experiments. On the basis of isotopic studies on Cr,03 (134d), Sheppard and co-workers have revised their original attribution of the bands to 0:- ions and now conclude that they are more consistent with a mononuclear species such as Fe=O. In the case of diluted MgO-Coo solid solutions, Zecchina et al. (1344 have made labeling experiments and assigned 0-0 stretching frequencies in the 1160-1015 cm-' range to adsorbed 0; superoxide ions. These results are in line with those obtained with oxygen carriers where absorption in the range 1120- 1140 cm- has been observed and assigned to coordinated molecular oxygen in agreement with the approximate representation Co(II1)-0; (134s).
'
5. Photoelectron Spectroscopy
Gopel et al. (135a)have reported ultraviolet photoemission spectra (UPS) of the interaction of 0, with the (lOT0) face of a single crystal of ZnO. Between 300 and 600 K, chemisorption of oxygen is observed on "stoichiometric" ZnO (1010) surfaces and UPS difference spectra indicate peaks at
MOLECULAR OXYGEN SPECIES ON OXIDE SURFACES
33
4,6.5, and 10.5 eV below E, (valence band edge). By comparing with EPR data and using the notation of Fig. 2, the bands were thought to arise from O,, with the peak at 6.5 eV tentatively attributed to the cglevels of 0; and the broader peak at 4.0 eV to the n,(,, levels. The electronic states of the adsorbed 0; are almost decoupled from the ZnO states, confirming that it can be regarded as a localized surface complex. It is also interesting to note that X-ray photoelectron spectroscopy (XPS) has also been applied to the study of dioxygen cobalt adducts. Burness et al. (13.56)have obtained the binding energies of the cobalt 2p electrons in the parent cobalt (11) complexes and the cobalt 2p and oxygen 1s1,2 binding energies in the dioxygen adducts and interpreted their results in terms of a formulation of the dioxygen adducts as Co(111)-0;. They were also able to measure the electron transfer from cobalt to oxygen and found a value of 86%, in good agreement with the value derived from EPR data (46b). B. THE0; ION The 0; ion has only a single electron in the ng orbitals, compared to three for the 0; ion (Fig. 2), and is isoelectronic with NO. The presence of the surface will break the degeneracy of the ng orbitals in the same way as for O;, but in this case an unoccupied molecular orbital is formed slightly higher in energy than the ng orbital since this must now contain the unpaired electron and the EPR signal is expected to have a negative g shift. In the solid state, EPR signals have been observed from a series of fluoride complexes, e g , 0; -AsF; (136);two of the g tensor components were between 1.96 and 2.00 and the third was between 1.73 and 1.76. This is in agreement with an analysis based on crystal field theory assuming an ionic model, which to first order gives gzz
= ge
- 2AlA,
gxx = g e -
21IEy
gyy
= ge
where the symbols have the meaning given in Fig. 2 and ni is assumed to contain the unpaired electron. The parameter 1is the spin-orbit coupling constant of oxygen. Certain preparations of Ti02 after heating in oxygen show complex EPR signals which have been assigned successively to 0; (137), coordinated oxygen (138), solid-state defects (139),and (TiO)3+(140).These assignments were based on an analysis of the g tensor believed to be orthorhombic. Repetition of experiments in X- and Q-bands has shown that the earlier interpretation of the Q-band spectra (137) was erroneous since two of the g tensor components were in fact hyperfine features due to 14N ( I = 1) (141).The EPR signals are formed when ammonia introduced by the pre-
34
M. CHE AND A. J. TENCH
parative method (142) is catalytically oxidized according to the Ostwald process (141): 2NH,
+ $0,- 2 N O + 3H,O
Nitric oxide can then be detected alone or in interaction with an oxide ion to form NO:- (141, 143). These processes suggest that the catalytic activity of T i 0 2 samples is likely to depend on both preparation and heat treatment (144). From the work discussed above, there is no clear-cut evidence at present that the 0;ion can exist on oxide surfaces. In this connection, Davydov (204 has investigated the adsorption of oxygen on the Mo0,/A120, and MoO,/MgO catalysts by IR spectroscopy. Broad absorption bands were observed in the 1500-1700 cm-' range which disappeared on heating to about 100°C. Davydov assigned these bands to adsorbed molecular oxygen; he explained the increase of voo on adsorption by a transfer of oxygen to its singlet state 'Ag. This explanation is doubtful since voo for the gas-phase singlet oxygen at 1483 cm-' ( 2 4 4 is lower than that for gas-phase triplet oxygen at 1555 cm-' (21). It is more likely, as indicated in Section 11, that a partial electron withdrawal from the antibonding orbitals of the oxygen molecule occurs to form an adsorbed species with some 0;character. C. T H E O ~ION -
The 0:- ion is normally referred to as the peroxide ion (33b),which should be distinguished from the covalent peroxy radical (ROO.). It has been previously treated in an earlier review by the present authors (1) as a dimer 0 - species. Although well known as a bulk peroxide (33b), this ion is difficult to characterize on the surface because it is diamagnetic and would be expected to be infrared inactive. Peroxides are associated with a broad optical absorption at about 260 nm (245a,b),which is very similar to 0;(Section III,A,4). Andersen and Baptista (146) have reported 0:-in KCl crystals, characterized by an optical absorption at 260 nm and distinguished from 0; by the absence of an EPR signal. Yao and Shelef (147) report a new EPR signal when oxygen is admitted to 12% Re/y-A120, catalyst after previous reduction in hydrogen. No EPR parameters are given but the signal is attributed to Re2+ and therefore taken as evidence that 0:- is formed on the surface. The arguments are not very convincing and the state of the oxygen on the surface is not well defined. Studies of metal-dioxygen complexes show that the peroxide-like complexes have IR bands voo in the range 800-932 cm-' (148).These data, taken together with the Raman work described below, indicate that the frequency
MOLECULAR OXYGEN SPECIES ON OXIDE SURFACES
35
range 1061-1088 cm-' (quoted on p. 97 of Ref. 1 ) previously assumed as characteristic of the peroxide ion in the alkaline-earth oxides (21)is incorrect and the bands probably arise from a superoxide impurity. Davydov et al. (149a) have reported IR observations of a negatively charged oxygen species formed when oxygen is adsorbed on chromic oxide absorbing at a frequency of 985 cm-'. The molecular nature of this species was confirmed by isotopic labeling and it was found to convert to an atomic form when the reaction temperature was increased to 200°C. The molecular species was not identified by the authors, but comparison with Raman data on bulk peroxides (149b) where the stretching frequency is near 800 cm-' suggest that the 0-0 bond is very weak and if it is a dioxygen species, it would correspond to 0;- on the surface. However, isotopic studies by Sheppard and co-workers (1344 do not confirm Davydov's original suggestion of dioxygen species and are more consistent with the presence of a mononuclear species such as Cr= 0. Conductivity and chemical methods (2) of measuring the charge on the oxygen do not distinguish between 2 0 - and O;-. It is therefore not possible to obtain direct evidence on the nature of the oxygen species and these approaches are not discussed further.
D. THEO:- ION Symons (150) has suggested that the 0- species on the surface would be better described as the species O;-, which is isoelectronic with F;. However, an experiment performed with MgO enriched in 1 7 0 showed that the interions was not measurable (151).Ben Taarit et al. (152) action of 0- with 702have observed an oxygen species on Pd(1) zeolite with g values of g1 = 2.050 and gl, = 1.99. The g values are inverted from those expected for 0, (i.e., g1 < gll) and this could be accounted for by rotational averaging. Alternatively, the authors speculate that the signal may be due to 0;- ions with the unpaired electron in a CT* orbital, since the g values would then be as expected from analogy to Cl, and F; (153). In support of this, approximate measurements of the intensity of the signals indicate that three Pd(1) ions are lost to form one oxygen species, as would be expected for O i - . At the present time, the argument is open and more evidence is needed to support the existence of 0;- on surfaces. Even in the solid state there are few examples of O i - . Pure crystals of CeO, UV-irradiated at 77 K exhibit several paramagnetic centers with orthorhombic g tensors; a typical g tensor is gxx = 2.0175, g y p = 2.0054, and gzz = 2.0317. Several possible models for these centers are proposed, one of which involves an 0;- molecular ion near a stabilizing impurity ion (154).
'
36
M. CHE AND A. J. TENCH
IV.
Formation and Stability of Charged Diatomic Species
The adsorption of oxygen on an oxide surface depends on the method of pretreatment and this can be divided into three main types :
(i) The method most generally used in studies of oxygen species is slight reduction, by thermal treatment in uucuo or in a reducing atmosphere, at a few hundred degrees Celsius. This cleans the surface and produces a slight nonstoichiometry or valence change in the metal oxide so that the surface adsorbs oxygen readily. (ii) For stoichiometric oxides, UV or y irradiation has often been used after thermal treatment to provide excess electrons at the surface. The oxygen species are produced by adding oxygen to the samples after irradiating either in uucuo or in a reducing atmosphere; alternatively, by contacting the thermally treated sample with oxygen and then irradiating it in oxygen. (iii) Lastly, the oxygen species can be produced by secondary reactions. In these cases, the thermally treated surface can be activated by pretreatment with a reactive molecule ( H 2 , CO, etc.), normally at room temperature, followed by contacting with oxygen to produce the oxygen species. The adsorption of the reactive molecule is viewed as the primary reaction, whereas the formation of the oxygen species is the secondary one (96). In the following paragraphs, in order to help the discussion on the formation of the oxygen species, the various oxide surfaces have been divided somewhat arbitrarily into groups. Most of the discussion refers to the 0; species since, although there is much evidence to show that this is not the only dioxygen species, there is essentially no direct information on the nature of the other oxygen species. A.
IONIC OXIDES
1. Alkaline-Earth Oxides and Their Solid Solutions with Transition Metal Oxides Ultraviolet or y irradiation of MgO (68,155) in uucuo or in hydrogen forms electrons trapped on the surface, which will react with oxygen to form 0;. A typical g tensor is gzz = 2.0777, g y y = 2.0089, gxx = 2.0018 with a range of gzzvalues indicating the presence of several sites on the surface. The superoxide ion is stable at room temperature for several months. The identification was confirmed by using 1 7 0 2 and the two oxygen nuclei were found to be equivalent, indicating adsorption parallel to the surface (68). Irradiation in the presence of oxygen leads to a more complex spectrum, indicating the
MOLECULAR OXYGEN SPECIES ON OXIDE SURFACES
37
presence of several oxygen species (68,156). Superhyperfine interactions can also be observed with nearby protons under some conditions (68,159). The formation of 0; by adsorption of oxygen has been reported on MgO (157)and CaO (71,158)which have been thermally activated in uucuo at high temperature. This has been confirmed by 1 7 0 2 (71) adsorption on CaO, but 0; is only formed when the CaO is made by thermal decomposition of hydrated CaCO, in uucuo after heating to 800°C and is not formed on CaO prepared from Ca(OH), in the same way. These observations can probably be best understood in terms of the influence of hydrogen-containing impurities (see next paragraph) in activating the surface. A similar effect would account for the observations with MgO. Derouane and Indovina (157) have attempted an analysis of the variations in the g tensor for the 0; on different crystal planes of MgO but the analysis does not seem justified, particularly since there is great difficulty in obtaining accurate values for gyy and gxx because of overlapping signals in a polycrystalline sample. It has been shown recently, that the alkaline-earth oxide surface can be activated for the formation of 0; by either preadsorbed gases (159) or by transition metal ions (110). Indovina and Cordischi (159) reported that exposure of a MgO surface to H,, CO, or C,H, after thermal activation followed by subsequent exposure to oxygen leads to a strong EPR signal from 0;. Preadsorption of H, (160) gave a multicomponent gzz feature ranging from 2.0895 to 2.0623, similar to that seen on irradiated samples (68, 155). The 0; signal was completely destroyed after heating at 300°C and different thermal stabilities were obtained for 0; in the different sites. Activation of the surface by H, was thought to occur via the homolytic dissociative adsorption of H, onto a pair of adjacent surface 0- ions originally formed by dehydrogenation of the surface under vacuum : H+ H, + 0,. . 0,
I
-+
Ht
I
0;; ' .Ot,
where 0;; refers to ions in low coordination on the surface ( I ) . The (0;;-H') entity then acts as the active center for electron donation to form 0; adsorbed at a nearby Mgz+ ion. However, neither the pair of 0- ions nor the active center could be seen directly by EPR, although they would be expected to be paramagnetic. A similar effect has been reported in CaO (158) and other work has shown that preadsorption of pyridine on MgO (72,161,162), CaO (70),and SrO (72,129) followed by adsorption of oxygen leads to the formation of 0; on the oxide surface. For SrO, measurements of spin concentration have shown that the 0; ions are produced by electron
38
M. CHE AND A. J . TENCH
transfer from a dipyridyl anion radical formed by adsorption of pyridine (129). Garrone et al. (163) have rationalized all these results by proposing the following mechanism for the adsorption of hydrogen and alkenes. Oxygen ions in positions of low coordination on the surface abstract protons from the adsorbed molecules to form OH, ions and a carbanion: R H + 0:; h O H , + R R - + 0, ‘ 0 ; + R . 2R. + 0, + R O O R
These carbanions react with oxygen to form O;, whereas the radical forms a bridged peroxide or dimerizes. In the case of hydrogen, adsorption leads to heterolytic dissociation at sites of low coordination on the surface (164a,b) to give(02--Hf)and(Mg2+-H-); the (Mg2+-H-)complex is then thought to act as the electron donor on oxygen adsorption, ultimately forming H, and 0; (165). This mechanism does not involve the formal transfer of electrons from the surface, although the overall reaction appears the same. A new form of adsorbed 0; has been reported by Ben Taarit et al. (63) which is formed by adsorbing oxygen at low temperature onto a MgO surface containing CO;. The resulting complex is assigned to (CO,-O,)with a g tensor of gzz = 2.040, gyy = 2.0072, gxx = 2.0015 and a hyperfine tensor obtained using I7O, which shows that the oxygen nuclei are not equivalent, with A,, = 100 and 50 G. On warming to room temperature, this species is transformed into 0; with the normal spectroscopic parameters. These hyperfine values indicate that the two oxygen atoms must be bonded end-on, at an angle to the CO,, in a peroxy-type linkage (Section III,A,2). Added transition metal ions can also induce the formation of 0; and adsorption of oxygen onto 1% Mn ions in MgO gives a poorly resolved signal centered about g = 2.007, which has been attributed to 0; (166). The Coo-MgO system has been studied in some detail, covering concentration ranges of 0.05-5 Co atoms per 100 Mg atoms (110).Two kinds of 0; ions adsorbed at Mg2+ and Co3+ sites can be identified (Fig. 1 l ) , together with some evidence for 0;. The O;-Co3+ complex is characterized by a g tensor of 2.124, unresolved, and 1.987. The large value (gzz) is high for adsorption at a 3 + cation (Fig. 3, Section III,A,I), but a superhyperfine interaction of 37, unresolved, and 17.5 G was observed and confirms that the adsorption site is a cobalt ion. It would seem that the bonding in the (Co” . . . 0;) complex is more D type than n type and the ionic model is not suitable. No I7O2 work has been reported. Adsorption of oxygen at temperatures above -70°C gives O;, thought to be adsorbed at a Mg2+
MOLECULAR OXYGEN SPECIES ON OXIDE SURFACES
39
I )
0;IMg2')
FIG. 1 1 . The EPR spectra at 77 K of 0; on COO-MgO samples. Spectra (a) and (b) were recorded after evacuation of oxygen at 298 K, (c) and (d) in the presence of a small amount of oxygen. Spectra (a) and ( c ) refer to a 0.2%COO-MgO sample, whereas spectra (b) and (d) refer to a 5", Coo-MgO sample (the cobalt concentration is expressed as Co atoms per 100 Mg atoms) (110).
site (gzz= 2.098-2.062). The (Co3+. . . 0;) complex disappears entirely at 25°C and is thought to form Oi-, although there is also an increase in the concentration of 0; adsorbed on Mg2+ ions. Zecchina et a!. (134e)have also studied oxygen adsorption on the surface of CoO/MgO solid solutions in the same composition range using IR spectroscopy. They conclude that -85% of the adsorbed oxygen is in an undissociated molecular form at 77 K, probably as 0; characterized by stretching frequencies in the range 1160-1015 cm-'. There was some indication of a bridged superoxide structure on the surface. The EPR data (110) indicates that 0; is adsorbed at both Mg2+ and Co3+ sites at fairly similar concentrations at 273 K, whereas for 0; formed at 77 K the Co3+ site predominates. This would suggest that the IR data at 77 K refer to the oxygen ion adsorbed at Co3+ rather than Mg2+. 2. Zinc and Cadmium Oxides The formation of oxygen species on ZnO has been of interest for some time (17). Thermal activation of ZnO at about 500°C in U ~ C U Ogives an EPR signal at g = 1.96 which is thought to arise from a donor species such as Zn+ ions (155,167-170). Adsorption of oxygen decreases the signal at 1.96,
40
M. CHE AND A . J. TENCH
and a new signal is formed corresponding to 0; with gzL in the range 2.052'2.042 (155, 167, 170-1740). The range of gLZvalues indicates that there are several sites on the surface which are all in reasonable agreement with adsorption at a cation charge 2 +. The signal is broadened reversibly by excess (80)indicates equivalent oxygen oxygen in the gas phase. Adsorption of 702 nuclei with a hyperfine tensor of A,, = 80, A,, = 0, and A,, = 15 G corresponding to oxygen adsorbed parallel to the surface. A report by Codell et 01. (174b) with different hyperfine splittings is inconsistent with the rest of the data and appears to be due to a misinterpretation of the EPR results. Some exchange of 0; with the lattice ions has been observed after heating at 200°C (80),which is slightly above the limit of its stability ( 180- 19OOC)on the surface as inferred from thermal desorption studies (170, 173). Oxygen adsorbed on ZnO with preadsorbed hydrogen was also identified as 0; by EPR, but it desorbed at 1 10- 120°C, indicating a marked reduction in adsorption strength for 0; on the hydrogen-preadsorbed sample (173).A similar species is seen after y or UV irradiation of ZnO (173, 174c, 175-179) and possibly also on Be0 (179) in the presence of oxygen. A signal with gZz= 2.045 has been attributed to 0; adsorbed on a ZnO (1010) face (180). Adsorption of CO or C,H, at 25°C on samples with preadsorbed oxygen does not change the EPR parameters of O;, but adsorption of NH, at - 30°C leads to an initial increase in g,, to 2.069, rising to 2.109 for excess NH, (181).The increase in gzz may reflect a change in bonding from n to o type or it may just reflect a decreased effective charge seen by the 0; ions. No work with I7O2 has been reported to check these two possibilities. From water adsorption experiments on ZnO, surfaces, it is suggested that 0; ions can be reversibly produced from 0;- ions (1744 according to the reaction 30:-
+ 2H,O
#
20;
+ 40H-
Setaka and Kwan (182) have investigated the CdO/Al,O, system and reported g values of 2.039, 2.009, and 2.002 for 0;. Although the authors did not discuss the adsorption site, the gzz value suggests that the 0; ion is adsorbed on A13+ rather than on CdZ+,in agreement also with the line shape of the signal, which indicates probable broadening due to 27AI nuclei. 3. Tin Oxide Thermal activation of SnO, at 500°C in
U ~ C U Ogives
an EPR signal at
g = 1.896 which has been attributed to donor electrons (183). Oxygen adsorption gives a complex signal (183-186). A major triplet with a g tensor of
2.024, 2.009, and 2.0036 has been assigned to 0; adsorbed at Sn4+ sites (74, 184) and both oxygen nuclei are found to be equivalent with an "0 hyperfine splitting A,, = 80.5 G (74). Two minor triplets with g = 2.028,
MOLECULAR OXYGEN SPECIES ON OXIDE SURFACES
41
2.008, 2.002 and g = 2.00, 2.00, 1.9984 have been assigned to 0- and O:, respectively, by Mizokawa and Nakamura (183).This seems unlikely in view of the fact that only the major species shows a hyperfine splitting with "0. Furthermore, the assignment of the latter triplet to 0; seems to be in disagreement with theory, which predicts gYY= ge and gx,, gzz < ge (see Section 111,B). Separation of the peaks into triplets gives ambiguous results without work at more than one frequency; for example, Meriaudeau et al. (74) find another triplet with g = 2.034, 2.004, and 1.994, the origin of which is uncertain. It is clear from measurements of adsorption and conductance (183) that 0; is not the only species formed, and nonparamagnetic oxygen species must also be formed on the surface. Parallel thermodesorption and EPR experiments indicate that 0; species adsorbed at Sn4+ ions are stable up to 150°C (170). The complex EPR spectrum of 0; adsorbed on SnO, has received another interpretation by Anufrienko et a!. (22f,187), who assign the various peaks
FIG. 12. Analysis of the spectrum of 0;stabilized on Sn2+ showing hyperfine lines due to 16% naturally abundant 115-117*119Sn(I =
wn.
42
M. CHE AND A. J . TENCH
to superhyperfine lines due to interaction of the unpaired electron with l15Sn, "'Sn, and '19Sn, all with I = 3 (all have about the same magnetic moment and a total natural abundance of 16%)(Fig. 12). The 0 , species are believed to be adsorbed either on Sn4+ ions with g , = 2.025, y 2 = 2.009, y3 = 2.0036 and A , = 27.5, A 2 = 34, A 3 = 25.5 G , or on Sn2+ ions with g , = 2.049, g2 = 2.009, g3 = 2.0028 and A , = 47, A , = 58, A 3 = 47 G, the former disappearing at lower temperatures (200°C) than the latter (300°C).
I al
50 G
F-----i
Ibl
FIG.13. The EPR spectra at 77 K of 0 ; ion on SnO, pretreated at 400 C (a) in (b) in hydrogen (188).
L'IICUO or
MOLECULAR OXYGEN SPECIES ON OXIDE SURFACES
43
Recent experiments using thermal treatment in uucuo and hydrogen (Fig. 13) and a Q-band spectrometer suggest that the hypothesis of Anufrienko et al. is not correct and that the complex spectrum is best explained by the presence of 0; stabilized on Sn4+ ions in different environments (188). 4. Thorium Oxide Thermal treatment of T h o , at 450°C in uacuo or hydrogen gives no EPR signal, but subsequent adsorption of oxygen gives a complex signal with gzz = 2.0589, 2.0540, gyy = 2.0098, 2.0092, and gxx = 2.0073, 2.0042 (189). This is consistent with the formation of 0; in two different sites on the surface. The nature of the cation at the adsorption site is not clear, but the gzz values are more consistent with a cation charge of 2+ than 4-t. The signal from 0; gradually decreases in intensity with increasing temperature of annealing and disappears at about 300°C; this process is thought to involve conversion to other oxygen species such as 0-,but there is no evidence for such species from the EPR data. T h o , is one of the cases (see alkaline-earth oxides and ZnO) where the effect of preadsorbed gases on the formation of 0; has been studied (62, 190). Preadsorption of C O leads to an EPR signal with g1 = 1.998 and gII = 1.981 (190).An additional new signal appears on adsorption of 0, at 77 K which has been attributed to two species; species A with g values of 2.019, 2.008, 2.002 and species B with g values close to the preceding values, except that gzz is between 2.088 and 2.040 (62).After warming to 298 K, just one species remains with a gzz of 2.048. Species B is also formed by adsorption of oxygen at 298 K followed by hydrogen. Adsorption of 1 7 0 2 (62) shows that species B is 0; with two equivalent oxygen nuclei ( A x x= 75 G), whereas species A appears also to be 0; but with inequivalent oxygen nuclei ( A x , = 95 and 65 G). Comparison of the EPR spectra for 0; (B) at 77 and 298 K indicates that rotation is occurring about the gzz axis at the higher temperature to give an axially symmetric g tensor while the hyperfine splitting decreases with increasing temperature. It was suggested that the bonding of the 0; with the surface varies with temperature, but it seems more likely that the effects on A,, arise from the rotation (Section III,A,2). 5. Rare-Earth Oxides, Including Scandium and Yttrium Oxides Thermal activation of either pure (191, 192) or silica-supported CeO, (45, 191) at 500°C in uacuo gives a signal at gav = 1.963 which has been attributed to Ce3+ ions or to electrons in the solid. The adsorption of oxygen on pure CeO, gives a poorly resolved signal with gll = 2.0312 and g1 = 2.0137 (192),whereas on the Si0,-supported system an orthorhombic g tensor is reported with gzz = 2.028, gvv = 2.0109, and gxx = 2.0158 (45). Adsorption
44
M. CHE AND A. J. TENCH
of 702 on either system (45, 73) indicates that both oxygens are equivalent with A,, = 75 G. This spectrum gives rise to an unusual situation because the hyperfine structure is not centered about the smallest g value as usually expected, but about the middle value of the g tensor, indicating that the nf and n: orbitals may be inverted (45). This gives rise to an inconsistency between the g and the A tensors, since both the smallest value of g and the largest value of A are expected to lie in the direction of the orbital containing the unpaired electron. In the g tensor above, the A,, component has been taken so as to correspond to the orbital containing the unpaired electron and used to label the g value of 2.0158 as gx,. The g tensor is in reasonable agreement with values obtained by Dufaux et al. (191) and by Setaka and Kwan (182)for CeO, on A l , 0 3 . A large giWvalue is found also for 0; on UO,/AI,O, (182). Probably the simple ionic model is not suitable for these systems, which have an unpaired electron in extended f orbitals. Steinberg and Eyal (193)assign an EPR signal with g = 2.010, 2.060,and 2.12to 0; on Y20,.Although the nature of the 0; species was confirmed by 170labeling experiments, the assignment to an orthorhombic y tensor would appear to be in disagreement with the line shape of the EPR signal, which gives a better fit to an axial g tensor with gIl = 2.060and g1 = 2.010. It is likely that 0 ; adsorbed at a second site with gll = 2.12 and gl = 2.010 is also involved. This would bring the results of Steinberg and Eyal in line with those of Loginov et a/. (194),who found three types of 0; ions on Y , O J , two of which, with g z z = 2.055, g,,,. = 2.007, gAx= 2.003 and gI1= 2.121, g1 = 1.9995, are in reasonable agreement with our new assignments. Loginov et ul. (1Y4) have looked at 0; on a number of rare-earth oxides which have preadsorbed gases such as H, and CO. The gZLvalue is shown to vary considerably with the type of pretreatment, from 2.035 to as high as 2.121. The large g shifts were interpreted in terms of considerable 0 bonding with the cation at the adsorption site. These ideas are not completely consistent with evidence from 0; on the supported CeO, system ( 4 9 , where the two oxygen nuclei are known to be equivalent. In the case of lanthanum oxide (La,O,), a small superhyperfine splitting was observed, confirming a lanthanum cation adsorption site. More information on 0; adsorbed on rare-earth-exchanged zeolites can be found in Section IV,C,3.
B. TRANSITION METALOXIDES I.
Titanium Oxide Thermal activation of TiO,, either as anatase or rutile, at 300-5OO'C
in uucuo or in a reducing atmosphere gives a slightly reduced solid with a paramagnetic signal which has been attributed to Ti3+ centers (137, 138,
45
MOLECULAR OXYGEN SPECIES ON OXIDE SURFACES
195, 196); thermal activation in air at 500°C produces complex EPR signals which have been assigned successively to 0; (137), coordinated oxygen (138), solid-state defects (13Y), and TiO: (140).These complex EPR signals are in fact due to adsorbed NO and NO:- species (141, 143), as discussed in Section II1,B. Exposure of the slightly reduced samples to oxygen leads to a new complex signal which has been attributed to various forms of coordinated oxygen and/or o-, 0; (138,185,197) or to 0; ions adsorbed at different surface sites (I74c, 196). The situation was clarified by the work of Naccache et al. (75), who were able to show, using "O,, that the signals observed on anatase or rutile (Fig. 14) should be attributed to 0; adsorbed at different sites on the surface, but in all cases the oxygen nuclei were equivalent (Table VI), indicating that 0; is adsorbed parallel to the surface. This has been confirmed by theoretical calculations (43b).The gzzvalues are reasonably consistent with adsorption at a Ti4+ site and the A,, values show only a small variation. In addition to the 0; signals, a symmetrical signal was observed at giso= 2.003, which was seen also by Van Hooff (185). The origin of this signal is not clear, since no hyperfine splitting could be detected using 70-enriched oxygen.
Y
:
10
Ti3'
n I
100 G
L
160170
170 170
L
I I
I
I
I
1
I
I I
I
I I
I
I
I
I
FIG.14. The EPR spectra at 77 K of 0 ; ion on reduced TiO, showing the hyperfine interaction with two equivalent oxygen nuclei (75).
46
M. CHE AND A. J . TENCH
TABLE VI Spectroscopic Constantsfor 0;Adsorbed on Thermally Activated TiO,"
TiO,
8rz
Rutile (species I unstable at 25°C) Rutile (species I1 fairly stable at 25°C) Anatase Ih (species I stable at 25°C) Anatase Ilh(speciesI 1 unstable at 25°C)
2.030 2.020 2.025 2.024
g ,,
2.008 2.009 2.009 2.009
Yxx
A,, (G)
2.004 2.003 2.003 2.003
76 72 77 77
From Ref. 75.
' Anatase I was prepared by flame hydrolysis of the chloride, whereas anatase I1 was prepared by precipitation.
In their study of oxygen adsorption on TiO, by temperature-programmed desorption and EPR, Iwamoto et al. (170)have shown that several types of 0; ions with different gzz values and thermal stabilities could be detected. They found that 0; ions with yzz values of 2.019, 2.023, and 2.026 were related to desorption temperatures of 125, 250, and 190°C, respectively, and suggested that variations in g z z values and thermal stabilities for 0; were due to differences in coordination number of the Ti4+ adsorption centers (see Section III,A,l). Similar results are obtained for thermally activated bulk or supported TiO, systems. Shvets and Kazansky (198)found that two types of 0; could be observed at 77 K on TiO, supported on silica, with yzz values of 2.026 and 2.020, their relative intensities depending on the TiO, content. They assigned the gzz value at 2.026 to 0; adsorbed on tetrahedrally coordinated titanium ions formed at low TiO, content, whereas the yzzvalue at 2.020 was related to 0; adsorbed on titanium ions in square pyramidal coordination prevailing at higher concentration. In a later study on TiO, supported on porous Vycor glass (PVG), Shiotani et al. (66) reported two triplets at 77 K with yzzvalues of 2.0237 and 2.0305, which they assigned, by analogy with the results obtained by Naccache et al. (75) for unsupported TiO,, to 0;adsorbed on anatase and rutile, respectively. Che and Naccache (199)have studied the kinetics of 0; formed on slightly reduced anatase using EPR. They found that the adsorption could be explained on the basis of different formation rates for 0; adsorbed at different sites, with zero- and first-order kinetics for the oxygen and Ti3+ concentrations, respectively. Using the same approach, Hauser (200) has extended this work and proposed different models to explain the kinetics based on the formation of O;, 0,-, and O f -ions for which activation energies around 1 kcal/mol were obtained. Nikisha et al. (201) have studied the oxygen adsorption kinetics using EPR, conductivity, and volumetric measurements.
-
N
MOLECULAR OXYGEN SPECIES ON OXIDE SURFACES
47
They concluded that the initial fast step involved the localization of electrons by oxygen, but significant amounts of 0; were formed more slowly. The amount of oxygen adsorbed exceeded the concentration of 0; by an order of magnitude or more in the case of highly reduced TiO,. Infrared studies (20b) indicate that oxygen is adsorbed in three forms: a neutral molecular form which absorbs in the range 1600-1700 cm-', a second molecular species absorbing at 1180 cm-' which is assigned to O;, and a dissociated form which is characterized by a metal-oxygen bond vibration in the range 700- 1000 cm-'. Presumably the surface perturbs the adsorbed oxygen sufficiently to make the molecule infrared active as discussed in Section III,A,4. With increasing reduction of the sample, the amount of the neutral molecular form became progressively less and most of the oxygen was adsorbed in a dissociated form. A wider range of gzz values (2.0213-2.0330) for 0; was observed for the more highly reduced sample. Calculations using extended Huckel theory (202) suggest that these changes are to be expected with an increasing degree of reduction. In the presence of oxygen atoms, 0; is not formed on rutile but subsequent exposure of the sample to molecular oxygen gives 0; with y,, = 2.019 (203). Many papers covering oxygen photoadsorption on Ti 0, have been published [see, for example, Refs. 204,205 and the references quoted therein, and also the review by Bickley (206)l.The subject is complex, but there is general agreement that the hydroxyl groups at the surface participate in the photoadsorption of oxygen by TiO, (207). Ultraviolet irradiation of TiO, in the presence of oxygen at 77 K can lead to a number of paramagnetic oxygen-containing species depending on the outgassing conditions of the solid prior to UV irradiation. 0 - (or O:-), HO,., O,, O,, and 0;- have been reported (88, 205, 208), but unambiguous assignment has proved difficult (I; see also Sections II1,D and V,A). If the sample is warmed to room temperature, only the 0; species remains visible. Meriaudeau and Vedrine (88)have used "0, labeling to study the species produced by photolysis at 77 K in oxygen on TiO, dehydrated at 450°C. A normal 0; is formed with g values of 2.021,2.009,2.001 and a hyperfine splitting of 77 G with equivalent oxygen nuclei. Two other species were observed with g = 2.014,2.009, 2.003 and gIl = 2.008, g1 = 2.001 which were attributed to 0; and O:-, respectively, but no hyperfine structure was seen. However, the species attributed to 0;- (see Section V) readily reacts with C O at 77 K to give a new species identified as 0,-O,,-CO- with a g tensor of 2.0465, 2.006, 2.001 and with A,, = 104 and 42.5 G for 0, and O,,, respectively. This is a peroxy-type radical with nonequivalent oxygen nuclei, in which all the unpaired spin resides on oxygen atoms I and I1 originating from the gas-phase molecular oxygen, and is thought to be formed by the reaction 0:-
+ co -+O& + 0-0-co
48
M. CHE A N D A. J . TENCH
Supported T i 0 , systems have also been used for photoadsorption studies. Shiotani et ul. (66) have reported that UV irradiation of the TiO,/PVG system in oxygen gave rise to various oxygen species. One species was unambiguously identified as 0; by means of 70-enriched oxygen. Measurement of the hyperfine tensor at low temperature showed that two slightly inequivalent oxygens were present in the same 0; with A,, values of 74.9 and 80.3 G at 36 K, in good agreement with earlier data obtained for anatase (82). 2.
Vanudium Oxide
Because of the superhyperfine interaction which arises when the 0; ion is formed on a cation with nonzero nuclear spin (see Section 111,A,3) vanadium pentoxide, with 100% naturally abundant "V isotope (1 = i),has been of considerable interest. However, the presence of a superhyperfine splitting has created some difficulty in the assignment of the signals. V,O, cannot be prepared with large surface area and most of the data refer to supported V,O, systems. Silica-supported V,O, is generally activated by thermal treatment at ca. 500°C in an atmosphere of oxygen followed by hydrogen. This procedure leads to the formation of tetrahedrally coordinated V4+ ions (209) and subsequent adsorption of oxygen gives a complex EPR signal. This was initially thought to be from 0; (210) or a combination of 0; ions (211) and was then reinterpreted in terms of a mixture of 0; and 0 - (106)with the following parameters for 0; : gzz = 2.023, yyy = 2.01 1, glx = 2.004 and A,, = 9.7, A,, = 6.8, A,, = 5.9 G. The superhyperfine interaction arises from 51V and confirms that the adsorption site is a vanadium cation. Calculations of the electronic structure and superhyperfine parameters indicate that the 0; ion is donor of n and B electrons to the metal ion and give reasonable agreement with the experimental values (117, 118). Spectra ob) not sufficiently well resolved to give the tained using I7O2 ( 2 1 2 ~were unpaired electron distribution between the oxygen nuclei. The thermal stability is dependent on the experimental conditions. Shvets et ul. (106) reported that the 0; ion was stable at temperatures up to 300°C in an oxygen atmosphere, whereas Yoshida et al. (212b)observed that heating for 15 min at 150°C caused a decrease of 0; by 80%. These results are to be compared with those reported by Iwamoto et ul. (170),who observed a broad desorption peak ranging from 100 to 500°C and assigned it mainly to 0; by comparison with earlier EPR results. Fricke et ul. (107) have studied the formation of 0; and 0 - on silicasupported V20,-P,0, catalysts. The 0; and 0 - formed are stabilized on vanadium ions, but the amount decreased with increasing fraction of P,O,.
MOLECULAR OXYGEN SPECIES ON OXIDE SURFACES
49
However, the maximum selectivity for butene oxidation to maleic anhydride occurs at a P/V ratio of 2/1, where the concentrations of 0- and 0; are much reduced. This was taken to indicate that the predominant role of 0; and 0- in this system is in nonselective oxidation. ZrO,, MgO, and Al,O, have also been used as supports (198,213, 214). 0; ions with g values of 2.032,2.009, and 2.003 can be formed after adsorption of oxygen at room temperature on slightly reduced V,O,/ZrO,. The gzzvalue of 2.032 coincides with that obtained for 0; adsorbed on the ZrO, support and is therefore characteristic of Zr4+ adsorption sites (198).After admission of 0, at room temperature on thermally reduced V,O,/MgO (213, EPR signals with gzZ = 2.070 and 2.080 are formed which are consistent with adsorption of 0, at Mg2+ sites; these ions are stable up to 150°C. Adsorption of oxygen at 77 K resulted in a more complex situation and gzz values were observed at 2.080 and 2.090 originating from 0; on Mg2+ sites together with a new signal with g values of 2.026, 2.009, and 2.003. This latter signal did not exhibit any superhyperfine structure from ,'V and disappeared on warming to room temperature. The authors suggested that this signal also was due to 0 , because it did not disappear on contact with H,, and that lattice 0,- ions or pairs of vanadium ions were involved as the adsorption sites. In the case of V,O,/AI,O, it was necessary to adsorb oxygen at 77 K to detect a signal at gZz= gII= 2.024 and g1 = 2.008 which disappeared on heating to room temperature. The assignment was similar to that given for the signal at 2.026 in the V,O,/MgO system (213,214). Khalif et a!. (215) have carried out adsorption, microcalorimetric, and EPR studies of oxygen chemisorption on V,05/Mg0 and V,05/Al,03 to determine heats of adsorption. The interpretation of this type of measurement is difficult because oxygen is adsorbed in more than one form. For V,O,/MgO a comparison of the adsorption isotherm for oxygen and the EPR data for 0, showed that 0, only appeared in the spectrum after adsorption of about half the oxygen, and it was assumed that the heat of adsorption of oxygen at the last adsorption point corresponded to the heat of 0, formation. This gives a value of 18-24 kcal/mol, which agrees well with the heat of 0; formation on MoO,/MgO and MoO,/AI,O, (216). The oxygen adsorbed during the first half of the isotherm was thought to be in the form of 0- ions because it reacted with C O with a heat of reaction of 60 kcal/mol, whereas the oxygen adsorbed in the second part did not react. The 0- ions are not visible by EPR because their association with paramagnetic ions leads to a strong exchange interaction. The heat of adsorption of the oxygen in the first half is 40 kcal/mol, but adsorption of other molecular forms of oxygen is thought to reduce the observed value from the expected 60 kcal/mol observed for MoO,/MgO (216). On V205/ Al,03, oxygen is adsorbed with a heat of adsorption larger than 60 kcal/mol;
50
M. CHE AND
A . J . TENCH
there is no EPR signal from oxygen species and almost no reaction with CO. It seems likely that in this case oxide ions are predominantly formed on the surface. This work is an interesting illustration of how microcalorimetric data can be used in conjunction with other techniques to obtain direct information about the thermodynamics of adsorbed species, but the identification of the adsorbed species is not always certain. 3. Chromium Oxide
Despite their importance in olefin polymerization reactions, little attention has been paid to the nature of the adsorbed oxygen species on supported chromium oxide systems. The formation of 0; ions has been reported (217) for supported chromosilicate catalysts after reduction at 500°C with carbon monoxide and subsequent exposure to oxygen. However, the values yII = 2.007 and yI = 2.004 are quite different from what would be expected for 0, on this system on the basis of the ionic model [Eq. (6), Section III,A,l]. Doi (218) has reported different 0; species for CrO,/SiO, catalysts; reduction with ammonia gave a signal with g1 = 2.020, g2 = 2.009, and g3 = 2.004 together with a line at 2.027, whereas reduction in hydrogen gave only the line at 2.027; signals attributed to 0; were also observed. Howe (219) has used a different method of preparation based on the decomposition of Cr(CO), on silica, and tentatively identified a poorly resolved EPR signal as due to 0; with yzz= 2.01 7, yyy = 2.010, and g,, = 2.010. The gzZvalues in the range 2.017-2.020 would not seem unreasonable for 0; on Cr6+, but this system is not fully understood at present and experiments using 170need to be carried out. A difficulty arises with Cr ions because several oxidation states (from + 2 to +6) can be stabilized on the surface, depending on the thermal treatment, and a range of yzzvalues is possible. The thermal stability of these 0; ions is also unusual, since their EPR signals disappear on evacuation at room temperature and can be restored by subsequent reexposure to oxygen (217,219). Shvets ef ul. (86) have recently reported both 0; and a CO; + 0, adduct on CrOJSiO,. The CO; + 0, adduct is described as CO, and has a y tensor of 2.046,2.006,2.001 with 7O hyperfine interactions corresponding to two inequivalent oxygen nuclei (98 and 42 G). This is very similar to the adduct on MgO (63) and MoO,/SiO, (87) (Table IV). On warming, the adduct decomposes, giving off CO, and forming 0; with a y tensor of 2.070, 2.006, 2.001 while the hyperfine tensor remains the same. When observed at 300 rather than 77 K, the EPR signal is isotropic with giro= 2.022 and an isotropic hyperfine interaction of 30 G is observed, indicating considerable rotational freedom on the surface. From IR studies, Davydov et al. (1494 have reported an absorption from
51
MOLECULAR OXYGEN SPECIES ON OXIDE SURFACES
oxygen at 985 cm-I on Cr,03 which was attributed to a molecular oxygen species such as O:-. Subsequent work using isotopic labeling with I6O/l ' 0 by Sheppard and co-workers (134d) has not confirmed this assignment and it is more consistent with a mononuclear species such as Cr=O. 4. Molybdenum Oxide
A number of very important reactions such as selective oxidation (220), olefin metathesis (221), and hydrodesulfurization (222) are catalyzed by systems involving molybdenum. Because of this interest, the adsorption of oxygen on supported molybdenum oxide has been investigated by many authors. Usually MOO, is supported on SiO, or Al,O,, but MgO and TiO, have also been used as supports. The molybdenum is activated for oxygen chemisorption by thermal treatment in uucuo at 500-600°C or by reduction in hydrogen to give EPR signals assigned to Mo5+(191,223).The adsorption of oxygen at 77 K on MoOJSiO, and MoO,/AI,O, was first studied by Dufaux et a/. (191) and subsequently by a number of other workers (81,82, 84, 85, 198, 213, 219, 224-228). The EPR lines from the adsorbed oxygen species are broad, particularly on Mo03/A1,0,, leading to some variation on reported g values, but generally accepted values are y1 = 2.016-2.0175, 9, = 2.0098, and y3 = 2.0042.These values are consistent with 0; adsorbed at a cation site with high charge such as Mo6+ (191).The departure of the smallest g value from that of the free electron is probably indicative of some covalent bonding rather than a purely ionic interaction (117). Krylov et d. (213) and Howe and Leith (226) have used the yzzvalue of the y tensor to show that adsorption of oxygen at 300 K on MoO,/MgO and Mo03/A1,03 leads to oxygen adsorbed at Mg2+ and A13+ cation sites, respectively. Electron transfer from one adsorption site to another (Fig. 15) was proposed by Krylov et ul. (213)where the original 0; adsorbed on Mo6+ was formed by adsorption of oxygen at 77 K. Similar evidence is also available from other work (227). Since the grr value can only give an indication of the charge at the adsorption site, it is more informative to study the superhyperfine interaction from the cation (see Section III,A,3). For this purpose, 95Mo-enriched catalysts
I
Mo6'-0
&
l
l
M o 6 L 0 - A13'
-
I
Mo6*-0-A13'
FIG. 15. Electron transfer occurring at the surface of MoO,/AI,O, according to Krylov e r a / . (213).
52
M. CHE A N D A . J . TENCH
have been prepared (109, 229) and oxygen adsorption investigated (108, 109).Che et al. (108)have reported a superhyperfine interaction with A,, = 2, A,, = 1.9, and A,, = 1 G, for 0; at 77 K and 300 K on MoO,/SiO,, arising from interaction with a Mo nucleus ( I = $). For MoO,/Al,O,, the situation was more complicated, and on warming to 300 K both a g,, value and a superhyperfine structure characteristic of adsorption at AI3+ were observed. This is clearly not a simple transfer process, since the total concentration of the different 0; ions increases on warming and 0; stabilized on Mo6+ can still be observed if the sample is recooled to 77 K. Thus, formation of 0; stabilized on A13+ has been attributed to electron transfer not from 0; adsorbed on Mo6+ but from reduced molybdenum sites not available at 77 K. The electron transfer can then be envisaged from Mob'' ions located in the bulk (Fig. 16). This is not unexpected, since it is known that molybdenum ions deposited on the support surface can migrate at moderate temperatures into the bulk of the matrix, e.g., MoO,/TiO, (230) and MoO,/SnO, ( 2 3 1 ~ )The . results obtained for oxygen adsorption on MoO,/AI,O, and MoO,/MgO suggest that a similar migration into the bulk occurs for alumina and magnesia as supports. This is a good example which demonstrates the difficulty of ascertaining the environment of the adsorbed oxygen and the complexity of the processes on the surface. "0, studies on both MoO,/SiO, and MoO,/Al,O, (82,84,85)confirm a diatomic adsorbed species; for MoO,/Al,O, the oxygen nuclei are nearly equivalent, with A , , = 77 and 80 G, but for MoO,/SiO, they are clearly different, with A , , = 72 and 85 G (see Section III,A,2). Che et d.(8.5)attributed this difference to a particular geometry at the surface, probably depending on the energy levels of the d orbitals of the Mo ion relative to those of 0; (see Section lIl,A,2). Using "0 enriched oxygen, Giamello et ul. (231b) have observed both equivalent and inequivalent oxygens in various types of adsorbed 0; depending on the Bi/Mo ratios of the bismuth molybdates supported on silica. The reason for these observations is not clearly
I
Fiti 16 (ION)
I I I
l
i
\e-
1
I II
Electron transfer occurring at thc surface ol' MoO,/AI,O, according to Che r t
ctl
MOLECULAR OXYGEN SPECIES O N OXIDE SURFACES
53
understood. Balistreri and Howe (2314 have irradiated MoO,/SiO, catalysts at 77 K with 306-nm light in the presence of both 0, and H, to form 0; and OH radicals. Warming above 77 K gives new signals which the authors tentatively assign to 0- and HO,. Irradiation in the presence of 0, enriched with 1 7 0 gave the normal hyperfine pattern for 0; but none for the OH signal, suggesting that the OH was formed from the lattice oxide ions. The 0; ion is thermally stable in oxygen for MOO, on Al,O, and MgO up to 150°C (213).Khalif et al. (216)have measured the heat of formation of 0; as 20 kcal/mol independent of support and in good agreement with work on V,O,/MgO (215), and suggest that other forms of oxygen such as Oi-, 0-, and 0,- which have higher heats of formation (viz., 60-80 kcal/mol) are also present. It is not always clear at which site the oxygen is adsorbing, but for low MOO, concentration the gzz value indicates that the adsorption site is A13+, and at higher concentration of MOO, the grr value of 2.0234 is rather larger than expected for a Mo6+ site. Akimoto and Echigoya (232) have studied the reactivity of 0; on supported MOO, in the catalytic oxidation of butadiene and this is discussed in Section VI. 5. Tungsten Oxide WO, supported on MgO or Al,O, can be activated by thermal treatment at 600°C in onc'uo or reduction in hydrogen to give weak EPR signals in the range y = 1.76-1.82 which have been attributed to W5+ (233).Adsorption of oxygen at 77 or 300 K with Al,O, as a support gives 0; stabilized at A13+ sites characterized by gzz = 2.040 and at W6+ ions with gZL= 2.019. For MgO as a support, adsorption of oxygen at 300 K leads to 0; with g,, = 2.070 characteristic of Mg2+ adsorption sites, whereas at 77 K three gII values were reported at 2.070, 2.080, and 2.026 (233).While the 2.070 and 2.080 values were indicative of Mg2+ sites by comparison with results obtained with the MoO,/MgO system (21.9, it is not clear from the discussion given by Spiridonov et nl. (233) what is the assignment for the 2.026 value. For silica-supported tungsten prepared by decomposition of various organometallics containing tungsten, Howe (219) reported an EPR signal with g values of 2.025, 2.01 I , and 2.004 after oxygen adsorption which was assigned to 0; formed on tungsten. The gzZvalue is larger than expected for adsorption at a W6+ site and may indicate stabilization of 0; on the support ions (234). The molecular nature of the species was confirmed later by Kazusaka et nl. (61). Adsorption of 170-enriched oxygen gave a resolved spectrum at 77 K, indicating that both oxygen nuclei are equivalent with A , , = 74 G (i.e., adsorption parallel to the surface). Raising the temperature leads to broadening of the hyperfine lines and an averaging of the g tensor,
54
M. CHE AND A. J . TENCH
consistent with a restricted motion on the surface (Section III,A,I,c). The 0; ions were stable at 145°C but disappeared after outgassing at 200°C. More information on 0; adsorbed in tungsten-exchanged zeolites can be found in Section IV,C,3. 6. Iron Group Oxides
There are only a few spectroscopic studies on the adsorption of oxygen on the iron group oxides (COO,NiO, FeO, and MnO) which give direct evidence on the nature of the adsorbed oxygen species. This is because these oxides are difficult to prepare with a large surface area and also they are easily oxidized or reduced to form higher oxides or metal particles. In addition, the superoxide ion 0; cannot be observed by EPR if it is adsorbed on cations which are paramagnetic (Section IV) or on superparamagnetic or ferromagnetic particles. The papers dealing with the iron group elements exchanged in zeolites will be discussed in Section IV,C,3. a. Iron Oxides. Using IR spectroscopy, Griffiths et al. (1%) have studied the adsorption of oxygen on Fe,O, previously degassed at room temperature. After adsorption at 350°C two bands were observed at 1350 and 1270 cm-' and assigned to adsorbed 0, and O;, respectively. These assignments have been criticized by Davydov et al. (235), who suggested that the two bands were due to carbonate-carboxylic species. The problem has been reconsidered in a later study by Al-Mashta et al. (19b),and a number of bands have been observed and classified into A-bands ( I 350- 1250 cm- ') and B-bands (1100-900 cm-I), which were assigned to perturbed 0; and 0;- species, respectively, by comparison with absorption frequencies of model dioxygen compounds. However, subsequent isotopic labeling experiments have shown that the type B-bands should be reassigned to mononuclear oxygen groups of the general type Fe=O (1344, in agreement with the expected pattern (Appendix C). The frequencies of the A-bands lie between those of superoxides and gas-phase oxygen (Appendix C), making the original assignment of perturbed 0; (intermediate between 0, and 0;) reasonable. The chemisorption of oxygen on FeO prepared by decomposition of the oxalate has been investigated by Dyrek (236).This author reported a change in the EPR spectrum of Fez+ at g 3.0 into that of Fe3+ at g 2.0 as chemisorption proceeds. No oxygen EPR signal is observed and this was interpreted to mean that oxygen is chemisorbed as 0,-, in agreement with results obtained using the iodometric analysis described by Bielanski and Najbar (237). It is difficult to assess the validity of this latter method since we know of no example where it has been checked for oxides in which there was also independent spectroscopic evidence for the existence of 0 - or 0;
-
-
MOLECULAR OXYGEN SPECIES ON OXIDE SURFACES
55
adsorbed on the surface. By the nature of the method, the results are likely to be ambiguous except where the oxygen is adsorbed either as 0’- or 0; ions, as the only species. On the iron oxides, there is no E P R evidence for the electron transfer between Fez+ ions and oxygen, although it is well known in biological systems such as hemoglobin (238).However, it has been reported for some inorganic systems; for example, Imai and Habgood (239)have shown that, in Y-type zeolites, the formation of 0; was increased by adding small amounts of Fez+ ions which could act as electron donors: Fez+ +Fe’+
+
P-
Similar results have been obtained by Ismailov et al. (240) in Y-type zeolites containing iron impurities (Section IV,C,3). The formation of 0; on the ferrocene/porous Vycor glass system has been observed by Vanderspurt et al. (120) with a y tensor of 2.0300, 2.0100, 2.0020 and a superhyperfine sextet centered on each y component. These results were interpreted in terms of 0; adsorbed on the cyclopentadienyl ring of ferrocene and, assuming an ionic model, the yzzvalue of 2.0300 is broadly consistent with a 3 charge at the adsorption site. b. CohuIt Oxide. No E P R signal was observed by Dyrek (236) from COO prepared by decomposition of the carbonate, presumably because of a very fast relaxation of the Co2+ ions, and it was not possible to follow the oxidation of these ions on adsorption of oxygen. The results obtained by the iodometric method of Bielanski and Najbar (237) lead to the conclusion that oxygen is adsorbed as 0’- ions, similar to the FeO system discussed earlier. Using solid solutions of Coo-MgO with low COOcontents, Tabasaranskaya et al. (241) have observed the formation of an EPR signal 2.07 which was assigned to ions adsorbed on Mg” from 0; ions with y, sites. In a later more comprehensive study, Dyrek (242)found that for COO or its concentrated solid solutions (100-43.1 atom % Co), oxygen is adsorbed at room temperature as diamagnetic 0 2 -ions, whereas for moderately concentrated solid solutions (30.9-15.3 atom 7; Co), a poorly resolved E P R spectrum was assigned to superoxide ions 0; with approximate yzz values of 2.025-2.028 corresponding to Co3+ adsorption sites. On diluted solid solutions (10.4-3.0 atom % Co), 0; was adsorbed on Mg” sites with gzz = 2.07. Cordischi et al. (110)have extended the studies on the Coo-MgO system by impregnating magnesium hydroxide with low contents of COO (0.05-5 atom 7; Co) and showed unambiguously that oxygen is adsorbed on Co3+ at 77 K by the observation of a superhyperfine structure due to interaction of the unpaired electron of 0; with the nuclear spin of Co ( I = f). At higher temperatures of adsorption, the 0; ion is adsorbed on Mg” in agreement
+
-
56
M. CHE AND A. J . TENCH
with earlier results. Isotopic labeling experiments with " 0 indicate that in the latter case, the 0; lies parallel to the surface (243).In order to determine the role of the cobalt dispersion in the adsorption properties of the COOMgO solid solutions, Dyrek and Sojka (244a) have plotted the EPR signal intensity of 0; radicals adsorbed at room temperature as a function of the COO concentration. The curve passes through a maximum at 3.00 mole COOcorresponding to the maximum concentration of isolated Co2+ ions in tetrahedral coordination with trigonal distortion, which might imply that such ions are the adsorption sites. For higher COO concentrations, the number of Co2+ ions in clusters increases and this is thought to control the . the experiments form in which oxygen is adsorbed (0;or 0 2 - )However, were not designed to obtain information on the actual coordination of the surface Co2+ ions and how changes in the coordination would affect the adsorption of oxygen. Moreover, any contribution from the support ions to the adsorption properties of these solid solutions is not considered. Zecchina et al. (134e) conclude, from IR work, that surface Co2+ ions in a square pyramidal coordination can adsorb oxygen, but this is not substantiated by earlier results obtained by Hagan et al. (244b,c), who showed that surface Co2+ ions in Coo-MgO solid solutions were in tetrahedral coordination. c. Nickel Oxide. There is very little published work describing dioxygen species on NiO. Tsyganenko et al. (24e) have detected bands at 1500, 1140, and 1070 cm-' in the IR when oxygen is adsorbed at 77 K on NiO obtained from decomposition of the hydroxide in uacuo at about 550°C. Labeling experiments using various l 6 0 / l *O isotopic mixtures indicate that these bands correspond to dioxygen species. The band at 1500 cm- was assigned to neutral adsorbed oxygen, whereas the bands at 1140 and 1070 cm-' were attributed to O,, in reasonable agreement with the data in Appendix C. d. Manganese Oxide. Dyrek (236, 245) has investigated the adsorption of oxygen on MnO prepared in uacuo by decomposition of the carbonate in order to avoid any oxidation to Mn3+ or Mn4+ ions. A plot of the intensity of the single EPR line at g 2, due to Mn2+ ions, as a function of the amount of oxygen adsorbed shows a linear decrease suggesting oxidation of M n Z + into Mn3+.From results obtained independently by the iodometric analysis of Bielanski and Najbar (237), Dyrek concluded that oxygen was adsorbed in the form of diamagnetic 0'- ions. This author has also studied oxygen adsorption on MnO-MgO solid solutions (166)prepared by decomposition in uacuo of the parent coprecipitated carbonates. The solutions did not contain any Mn3+ or Mn4+ detectable by the iodometric method. Chemisorption of oxygen at room temperature on solutions containing 100-3.72 atom Mn was found to give rise to diamagnetic 0 2 -ions, whereas on more diluted solutions the oxygen gave an EPR signal with a g value of 2.007, thought to be from 0; or 0- ions. However, it is difficult to study
-
MOLECULAR OXYGEN SPECIES ON OXIDE SURFACES
57
these signals since at low concentration the Mn2+ ions give rise to a hyperfine structure, due to the interaction with the nuclear spin of Mn (I = ;), which overlaps with the oxygen signal. Since most 0 - ions are known to exhibit a characteristic EPR line shape (I),we would associate the line shape for the oxygen signal observed by Dyrek (166)with 0;rather than 0 - ,but this has not yet been confirmed by studies using " 0 . The results obtained by Dyrek (166) differ from those obtained earlier on the adsorption of oxygen on MnO-MgO solid solutions. Cordischi et al. (246), on MgO doped with 235 atom ppm of Mn, and Yamamura et al. (247),on MnO-MgO with low contents of Mn ( > 0, >> 0;. This is well illustrated by the reaction with ethylene where 0- ions react readily at -60°C 0, ions react at 25°C with a half-life of ca. 5 min, whereas only one-third of the 0; ions react after 2 hr at 175°C.The authors propose a number of surface intermediates (Table XIV) in the oxidation reactions based on analysis of the desorption products and IR studies. A number of generalized comments on the reactivity can be made for the MgO system. (a) In all cases, the principal initial reaction appears to be the abstraction of a hydrogen atom from the hydrocarbon by the oxygen, followed by subsequent surface reactions which may involve oxide ions of the surface.
1 I8
M. ('HE AND A. J . TENCH
(b) The intermediates for the oxidation of the alkanes always include alkoxide ions independent of the oxygen species involved. This probably reflects the stability of the alkoxide ion on the MgO surface. (c) Carboxylate ions are thought to be the intermediates in the reactions of C2 and C3 alkenes but the type of carboxylate ion formed with 0is different from that formed with 0; and 0; ; in the latter case, there is a scission of the C=C bonds following the initial step of hydrogen abstraction.
No similar comparative studies have been carried out using other supports and so the reaction behavior cannot be assumed to be general and probably, in part, it is controlled by the specific properties of the MgO surface. 3. Photo-oxidation Reactions
The 0; ion appears to play an important role in a number of photooxidation reactions (see Section VI,C); for example, the photo-oxidation of alkenes over Ti02. However, it seems likely that 0; is not, in many cases, active in the oxidation step but further conversion occurs to give a mononuclear species, not detected directly, which then oxidizes the adsorbed hydrocarbons. Photo-oxidation of lattice oxygen in the M=O systems (e.g., V 2 0 , supported on PVG) gives rise to an excited charge transfer state such as V4+-O-. This excited state can react as 0- either by addition to a reactant molecule or by an abstraction reaction (see Section V of Ref. I ) . In the presence of oxygen, 0; is formed which then reacts further with organic molecules.
4. Catalytic Reactions Selective oxidation by heterogeneous catalysis is of great industrial importance and accounts for no less than 21% of the major organic chemicals produced via reactions involving catalysis (428-431). The oxidation reactions include allylic oxidation to give aldehydes, nitriles, and acids; aromatic oxidation to give acids and anhydrides; epoxidation of olefins; methanol oxidation to give formaldehyde; and to a lesser extent, paraffin oxidation to give anhydrides (429-432). Because of this importance, there has been considerable effort to obtain a better understanding of the mechanisms of such reactions. However, there is only limited knowledge of the way in which oxygen is involved in the overall process. It is generally assumed that in, for example, allylic oxidation, an intermediate formed on the surface is oxidized by a specific type of lattice oxygen of the catalyst rather than an adsorbed oxygen species to form the reaction products such as acrolein (431). However, in both cases the adsorption of oxygen on the surface is of
MOLECULAR OXYGEN SPECIES ON OXIDE SURFACES
119
vital importance since the oxygen consumed in the reaction, whether from the lattice or from an adsorbed species, must be rapidly replenished. In heterogeneous catalysis, the oxide lattice can have a variety of different functions, but for this review we are concerned with its ability to provide oxygen in a suitably activated state to oxidize the reacting organic molecules to form the required products. From the evidence in the preceding sections on characterization it is probable that the adsorbed dioxygen can lie within a continuous range of species from electron deficient to electron rich, depending on the nature of the oxide and the reaction conditions. and On this basis, a division of oxygen species into electron rich (02-) electron deficient (e.g., 0;) is limited since it assumes that the species can be considered as entities separate from their environment. This approach is contrary to the picture that has emerged from the preceding discussion, where it became increasingly obvious that the nature of the species depends very much on its environment at the surface and that the formal description of a species as 0; does not necessarily accurately reflect its actual charge and bond order when adsorbed on the surface. In addition, it seems that many oxides can form a variety of different oxygen species on the surface depending on the reaction conditions. Bielanski and Haber (2) have divided the metal oxides into three main groups depending on their interaction with gaseous 0, : (a) p-type semiconducting oxides (NiO, MnO, etc.) which form electron-rich species (0-, 02-), (b) a group including n-type semiconductors (ZnO, TiO,, V,O,, etc.) and also dilute solutions of transition metal ions in diamagnetic matrices (e.g., COO in MgO) which form 0; and 0-,and (c) binary oxides where the lattice oxygen is present as 02-in well-defined oxyanions, e.g., Bi203.MOO,, and which do not form adsorbed oxygen species but only 02-ions. This would seem to be an oversimplification, since it seems likely that a range of oxygen species can be observed on all these oxides given the right conditions, although the thermal stability is likely to vary considerably. Oxygen can be involved in oxidation reactions in three distinct ways, more than one of which may be operative in any reaction mechanism. The first is the abstraction of a hydrogen or proton from an adsorbed organic molecule to give a radical or carbanion on the surface; the second is the attack on the organic species by a negatively charged oxygen ion whether lattice oxygen or an adsorbed oxygen; and the third is the replenishment of lattice oxygen which has been used in a direct oxidation reaction. The abstraction reaction appears to be very common and the preceding evidence shows that it occurs with Of;, 0-, O;, and 0; ions but the reaction with 0- is particularly fast. An exception to this, which at the same time provides strong evidence for the participation of a molecular oxygen species, is the selective oxidation of ethylene over silver catalysts to form
120
M. CHE AND A. J. TENCH
ethylene oxide (Section IV,D). A dioxygen ethylene complex is formed on the surface but it is not certain whether the precursor is O;, 0;-, or an intermediate between these two forms. This is the most direct evidence available for the insertion of oxygen into a C=C bond via a molecular oxygen species. Comparison with the reactivities measured by Iwamoto and Lunsford (393) on MgO would suggest that addition reactions to alkenes are consistent with 0; as the oxidizing species but hydrogen abstraction also occurs. For ethylene, only CO, was expected as a reaction product but the rate of reaction was slower than with propylene. This may indicate that the precursor of the ethylene complex on silver is a more electron-rich oxygen species such as O : - , where insertion to form a bridged structure is more likely because of the weaker bond strength. But there is no good evidence on the reactivity of this species. The interesting reactions are those involving a highly selective partial oxidation (428, 431), whereas unselective oxidation reactions are of interest only in very limited situations, for example, in the oxidation of car exhaust gases. The selective oxidation reactions, apart from the formation of ethylene oxide discussed earlier, are thought to occur via the Mars-van Krevelen (433) mechanism; the oxidizing agent is an 0,- ion from the lattice which, as it is incorporated into the hydrocarbon, donates its electrons to lattice cations. Such a mechanism is thought to be well established for the oxidation of olefins on Bi, O,/MoO, catalysts ( 4 3 4 , the ammonoxidation of propylene on Bi,O,/MoO, (435) and USb3OIo(436),and the oxidation of methanol on Fe,O,/MoO, (437a). For these reactions to occur the metal-oxygen bond energy of the active oxygen ions, under the reaction conditions, must be in a range where both removal and replacement can occur readily. Reoxidation by oxygen is assumed to lead to the formation of lattice 0,- ions via the simultaneous transfer of four electrons from the cations : 0,
+ 4e-
-+202-
However, in light of the earlier discussions, a stepwise transfer of electrons involving O,, Oi-, and 0- as intermediates might be expected; in this respect, surface potential measurements have proved useful (437b). For example, reoxidation studies on bismuth molybdate catalysts by Brazdil et al. (438)are consistent with a mechanism of oxidation where adsorption and dissociation of the dioxygen occurs before the rate-limiting incorporation step. It seems likely that these processes might be observed directly on this type of oxide if experiments were carried out under the correct conditions, but detailed investigations of these reactions have not been made. In these reactions which involve oxygen ions of the lattice, the actual nature of the intermediate species is not clear. The oxide structure appears to be associated with special defect arrays and there is evidence the lattice
MOLECULAR OXYGEN SPECIES ON OXIDE SURFACES
121
oxygen in different environments is responsible for the different steps in the process (439). There is also the possibility of a thermally activated charge transfer process leading to a reactive 0-species (see Section VII of Ref. 1 ) where the oxygen has a low coordination: o2-0 I1/
0-0
I/
0-M"' -0
O-M(n-l)+-
0
0
/
/
0
At lower temperatures the Mars-van Krevelen mechanism no longer applies. Sancier et al. (440) studied propylene oxidation in the presence of "0,over bismuth molybdate and found that the acrolein product contained '*O and not exclusively l60from the oxide lattice in contrast with results obtained by Keulks and co-workers (441, 442) at higher temperatures. This lower-temperature oxidation must involve adsorbed oxygen in some form but the nature is not clear. It is now accepted that not all these oxidation reactions do involve lattice oxygen (442,443). There are a number of other types of reactions where adsorbed oxygen species rather than lattice oxygen ions are thought to be the principal oxidizing agent. Tagawa et al. (444) have studied the oxidative dehydrogenation of ethylbenzene and concluded that gaseous oxygen forms 0-on the surface which abstracts a j?-hydrogen from the adsorbed complex. Akimoto and Echigoya (232,399,445)have given evidence that 0;ions are involved in the oxidation of butadiene to maleic anhydride over supported molybdena catalysts (Section VI). The authors suggest that the reactivity of the 0; ion and Mo=O are very similar, and that the Mo=O may behave like 0- during the oxidation reaction as indicated above. There is also evidence that 0; and 0-species take part in CO oxidation reactions on V 2 0 5 (365).Yoshida et al. (392)have reported the reaction of 0;on V 2 0 , with propylene and butene to form aldehydes while the lattice oxygen show little reactivity below 150°C. Fricke et al. (107) have concluded from a study of V,O5/P2O5 catalysts that both 0-and 0;lead to nonselective oxidation of butene in this system.
C . FUTURE DIRECTIONS Electron paramagnetic resonance has played a major role in the characterization of adsorbed oxygen species and the use of ''0,has enabled a major advance to be made in the understanding of the nature of the various oxygen species and how they can be bonded to the surface. The use of IR spectroscopy as a technique has tended to be neglected because of the
122
M. CHE AND A. J. TENCH
difficulty of unambiguous assignment of the bands to the various oxygen species. However, it is applicable to a variety of systems, particularly bulk transition oxides such as the iron group oxides, which EPR cannot easily probe. IR spectroscopy is also capable of providing information on nonparamagnetic species such as O:-. With an improved understanding of how the surface can perturb the adsorbed molecules/ions, IR, EELS, and possibly also resonance-enhanced Raman spectroscopy are likely to play an important role in the future, both in defining the nature of dioxygen ions of fractional charge and in the characterization of polynuclear species. Optical studies will grow in importance, e.g., for the 0;ion, and the related technique of photoluminescence spectroscopy is likely to be applied more widely. Both XPS and UPS have considerable potential but very careful experiments are necessary to improve the interpretation of the spectra. Solid-state NMR of adsorbed oxygen species labeled with "0 has not been reported up till now. If sufficient sensitivity can be obtained, this technique has potential in characterizing adsorbed species and their environment for nonparamagnetic oxides. Of the different oxygen species, the main interest has been in 0-, O;, and 0;.Relatively little attention has been paid to the characterization and reactivity of singlet oxygen, Oi-, lattice or adsorbed 02-species, and most importantly polynuclear species. The work on dioxygen species is likely to be related to the studies of oxygen carriers. Features of special interest in the future on the characterization side are likely to be the detailed geometry of the adsorption site, how the oxygen species is bonded to the surface, and its mobility. The majority of the work covered in the two reviews is concerned with studies at low temperature, whereas the catalytic oxidation reactions occur at elevated temperature. There are some indications that the different oxygen species may not be so clearly differentiated at higher temperature or that interconversion may occur readily. This is a particularly interesting area which deserves further exploration. In order to understand the mechanisms of catalytic oxidation, particularly those where selectivity is good, more work needs to be done with model systems. The stoichiometric reactivity experiments should be extended to other systems using IR spectroscopy as well as EPR to follow the oxygen species. Dynamic experiments at higher temperature on model systems are also required. In the future, emphasis needs to be placed on quantitative experiments where the kinetics are followed so that it is clear what reactions are being studied. Of particular interest in this area is the mechanism by which oxide ions of the lattice are replenished from the gaseous oxygen and how oxygen species are recreated by redox reactions following the oxidation reaction.
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MOLECULAR OXYGEN SPECIES ON OXIDE SURFACES
D. CONCLUSIONS In this article we have tried to make a realistic assessment of the present level of knowledge on the characterization and reactivity of various oxygen species on the surface. However, the form of a particular oxygen species in a specific oxidation reaction is yet to be conclusively established. Part of the reason for this is the need to break down any given reaction into its parts, which has often been necessary because of the limitations of the techniques employed, but speculation has all too frequently been accepted as fact. However, particularly the results from careful IR work point toward the need for a more flexible approach in understanding the nature of the species involved and are likely to lead to change in the paradigm presently adopted by many workers in the field. Above all, however, there is a need for a holistic approach to the reaction mechanism which combines both a study of the intermediates by a variety of techniques coupled with an overall analysis of the reaction pathway. It is difficult to combine a general semiempirical approach with specialized characterization but such a synthesis is most likely to lead us to a better understanding of the complex surface phenomena.
Appendix A.
Summary of g,, Values for 0; on Surfaces
~~
Systems
Bulk' NaO, KCI Surfacesd: Oxides MgO
CaO
SrO
gzz
2.175 2.436
2.0623-2.0895 2.0733-2.0779 2.077 2.0777 Range' 2.0623-2.0895 2.089-2.098 2.093 2.10 Range' 2.089-2.10 2.100 2.102 Range' 2.100-2.102
A (eV)*
Reference
0.16 0.06
446 42,69b
0.47-0.32 0.39-0.37 0.37 0.37 0.47-0.32 0.32-0.29 0.31 0.29 0.32-0.29 0.29 0.28 0.29-0.28
160 IS7 72, ISS 68 71 70 IS8 129 72
(Continued)
M. CHE AND A . J . TENCH
124
APPENDIX A (Confinued) Systems ZnO
CdO/A1,03 SnO,
Ce0,/AI,03 CeO,/SiO,
La203
sc,o, TiO,
TiO,/SiO, TiO,/PVG V,Os/SiO, V,O,/P20,/~iO, V,O,/MgO
OZZ
a
A (ev)”
2.042-2.051 0.71-0.57 2.0424-2.0519 0.70-0.56 0.68 2.0436 2.045 0.66 0.60 2.049 2.05 1 0.57 2.052 0.56 Rangee 2.042-2.052 0.71-0.56 2.039 0.76 2.024 1.29 2.026 1.18 2.0265 1.16 2.028 1.09 2.029 I .05 2.033 0.91 Range‘ 2.024-2.033 1.29-0.91 2.054-2.0589 0.54-0.49 2.0185 1.73 2.0312 0.97 2.0246 1.26 2.030 1.01 2.0266 1.15 2.028 I .09 Range‘ 2.0185-2.0312 1.73-0.97 2.035 0.86 2.060-2.12 0.49-0.24 2.055-2.121 0.53-0.24 2.063-2.093 0.46-0.31 2.047 0.63 2.019 I .68 2.020 1.58 2.021 1.so 2.0216 I .45 2.0225 I .39 2.023 1.35 2.0237 1.31 2.024 1.29 2.025-2.030 1.23- 1.01 2.020-2.026 1.58- I . I8 2.0223-2.0305 1.4-0.99 Range’ 2.019-2.0305 I .68-0.99 2.022 1.42 2.023 1.35 2.022 1.42 Range‘ 2.022-2.023 1.42- 1.35 2.026-2.090 1.18-0.32 2.07-2.09 0.41-0.32
Reference 167 80 174c 180 171, 172, I74a 155, 186 I 70 182 74 2?f 184 183, 186 I 70 185 189 191 73, 192 191 182 191 45 182 193 194 194 194 171, 186, 203 90,143 88 137 206 I 70 174c 75, 197 75 198 66
212h 106, 198 107 213, 214 198
MOLECULAR OXYGEN SPECIES ON OXIDE SURFACES
125
APPENDIX A (Continued) Systems
gzz
V2Os/AI@3 V,Os/ZrO, ZrO, Cr(CO),/SiO, CrOJSiO, CrOJSiO, Mo(CO),/SiO, MoOJSiO,
Range'
Range'
a-Al,O, Y-A1203
Range' Si02-AI,0, SiO, PVG Range'
A (eV)b
Reference
2.024 2.032 2.032 2.017 2.020 2.070 2.017 2.016 2.017 2.0173 2.0176 2.018 2.016-2.01 8 2.017 2.0155 2.0170 2.019 2.0155-2.019 2.039 2.035 2.039 2.070 2.025 2.0266 2.019 2.040 2.026-2.080 2.0300 2.025-2.028 2.062-2.098 2.070 2.124-2.138 2.0167 2.034 2.038 2.039 2.040 2.034-2.040 2.024 2.0250 2.0318 2.0310 2.0250-2.031 8
1.29 0.94 0.94 1.90 1.58 0.41 I .90 2.04 1.90 I .87 1.83 I .78 2.04-1.78 1.90 2.12 1.90 I .68 2.12- I .68 0.76 0.86 0.76 0.41 I .23 1.15 I .68 0.74 1.18-0.36 I .01 I .23- 1.09 0.47-0.29 0.41 0.23-0.21 1.94 0.88 0.78 0.76 0.74 0.88-0.74 1.29 1.23 0.95 0.98 1.23-0.95
2.022 2.032 2.038
1.42-0.94 830 0.78 121
213 I 98 171, 182 219 218 86,447 219,226 224 109 191 85, 108 81. 198 226 191 85. 108 213 226 213 108 108,213 219 61 233 233 233 120 242 110 24 I 110 111 248 111 108 112 254 252. 253 25 I 256
Zcolircs
HY
126
M. CHE A N D A. J. TENCH
Systcnis
ci:'i
A (eV)"
Rcference
AlHY H Mordenik Dehydroxylated I IY
103 I21 104 266
SCY TiY
2.038 0.78 2.040 0.74 2.038 0.78 2.0575 0.51 Alkali and alkalinc earth zeolites: See Tables IX and X 1.01 2.030 I.63 2.0195 I .6l 2.0197 0.74 2.040 0.66-0.56 2.045-2.052 0.9 1 2.033 2.034 0.88 0.67 2.044 1.29 2.024 0.49 2.060* 2.035- 2.0242 0.86- 1.28
103 279 280 280 281 285 26 7 103 105b
1.09- 1.01 I .09-0.50 0.74 0.77-0.19 0.96 1.16 0.82 0.20
56 289 288 290 299 299 299 302
TiA NiCaY Lax
La Y WHY WNaY CeX
105b 102u
Support14 metuls
2.028-2.030 2.028-2.058 2.040 2.0389-2.148 2.03 I6 2.0264 2.0366 2.141
Ag/PVG Ag/SiO, Ag/SiO, Ago Au/PVG RhjPVG Pt/PVG Pt/AI,O, Adducts'
[ CO"~(NH,),O;] ~~:C~Y [C~"'(NH,),~O;CO"'(NH,),]~~:COY [Co"'(CH,N H2),0;]' :COY [CO"~(P~NH,),O;]~+ :Coyu [Co"'(en),0;]2 :CoyB [Ru(C0)O,l4 :RuY Phthalocyaninato Co/AI 0 , y Coadsorbared MgO(C0,-0,) ThO,(CO-0,)A B TiO,(CO-0,)CrO,/SiO,(CO,-0,)MoO,/SiO,(CO,-O,)~ MoO,/SiO,(CO,-0,)HY(CO-0,)Orher systems' AlSb GaAs GaAs NaY +
+
,
2.084 2.072 2.075 2.079 2.084 2.056-2.083 2.098 2.040 2.019 2.040-2.088 2.046 2.046 2.0486 2.047 2.069 2.041 2.035 2.046 2.113
~
-
-
-
0.72 0.86 0.64 0.25
51 115 115 115
I16 83b 250 63 62 62 88 86 87 44 7 89
47 I13 I14 I276
127
MOLECULAR OXYGEN SPECIES ON OXIDE SURFACES APPENDIX A (Continued) Systems NaO, in krypton matrices KO, in krypton matrices RbO, in krypton matrices CsO, in krypton matrices SiOO' (bulk)' Polytetrafluoroethylene peroxy' Polypropylene peroxyf
Szr"
2.1106 2.1184 2. I227 2.1069 2.070 2.038 2.035
A (ev)" 0.26 0.24 0.23 0.27 -
-
Reference 44 44 44 44 91 58, 93 84
For each system, the order is given with increasing gzLvalues. Calculated using the simplified equation (6) (Section III,A,I,a); 1 has been taken equal to 0.014 eV (1276) so that comparison with earlier results can be made (3); A has been calculated for 0; ions only. ' Given for comparison. The systems are arranged using the same order as in Section IV. Some systems appear twice whenever the 0; ion can be stabilized either on the supported ion or on the support. g, values used to construct Fig. 3. The 0; ions on these systems do not fit the ionic model. Pr = propyl, en = ethylenediamine.
The choice of gzzvalues used to construct Fig. 3 has been restricted : 0 to oxides (bulk or supported oxides and zeolites) described by the ionic model 0 to gzz values which could be safely assigned to specific adsorption sites on the basis of superhyperfine interactions (see Table V and Section
111,A,3) 0 when superhyperfine interactions were absent, to gzz values which were confirmed independently by several laboratories 0 to gzz values relative to slightly reduced oxides. Stronger reduction usually results in a number of gzz values which are thought to be due to adsorption sites of various low oxidation states, different local coordinations, and/or different crystal planes (Section III,A, 1,a). This has been observed, for instance, in the case of T i 0 2 (20b, 170).
The gzz values were assigned to a given oxidation state of the adsorption site on the basis of spectroscopic and chemical evidence. For transition metal ions, the oxidation state was deduced from reactions of the type M'n+ O2-P Mn+O;, which were ascertained by a decrease in the EPR signal of M'"- l ) + ions and a parallel increase in that of 0; (Section IV). For nonreducible ions, the usual oxidation state has been taken. For the + 1 oxidation state observed only in alkali zeolites there is a large range of gzz values: 2.054-2.166 (Table X),which has been used in Fig. 3. It appears,
M. CHE AND A. J . TENCH
128
however, that the earlier value of 2.1 13 obtained in X-irradiated NaY zeolite (1276) and confirmed in Na-reduced NaY zeolite (272) corresponds to 0; adsorbed on Na' ions because of the presence of superhyperfine interaction due to the nuclear spin of Na (I = 3). Similar gzz values of 2.1 106, 2.1 184, 2.1227, and 2.1069 have been reported for the alkali superoxides NaO,, KO2, RbO,, and CsO,, respectively, trapped in krypton matrices (44).Thus the narrower range of gzz values 2.1 106-2.1227 for the + 1 oxidation state should be preferred. From Appendix A and Fig. 3, it is possible to deduce the oxidation state of the metal ion at the adsorption site and conclude whether the superoxide ion 0; is adsorbed on the supported oxide (or metal) or on the support by comparing gzz values relative to the supported system and to the support.
Appendix B. The Experimental '"0 Hyperf ine Parameters (in gauss) of Diatomic Oxygen Species (0; and ROO')
Systems
Bulk KCI K,S,O* SiOO'
A,,
Ayya
a
61.5 75.7 101.7, 43.2
-
A,,"
'isa
Reference
19.7 14.0 9,9.5
69b 448 91
0
15
-
-
~
9, 9.5
Sur/a:fbces
MgO y-irradiated MgO y-irradiated MgO (Pyridine + 0,) CaO CaO (Pyridine + 0,) SrO (Pyridine + 0,) ZnO SnO, CeO, CeO,/SiO, TiO, (anatase) TiO, (rutile I ) TiO, (rutile 11) TiO,
v,o,/sio, MoO,/AI,O, MoOJSiO, MoO,/SiO, Bi,O,, 3MoOJSi0, Bi,O,, MoOJSiO, IOBi,O,, MoO,/SiO,
71 11 76 17 76 76
~
~
-
-
-
-
-
-
68 334 72 71 70 72
15
80
80
0
80.5 75 75 77 76 72 80.3, 74.8 Notresolved 80.77 85,72 82,69
-
81 85,12 85.12
-
-
74 73 45 75 75 75
-
-
66
-
-
2120 85 82, 84, 85 81 231b 231b 231b
~
-
-
-
-
-
-~
-
-
~
-
~
~
-
129
MOLECULAR OXYGEN SPECIES ON OXIDE SURFACES APPENDIX B (Continued) Systems WO,/SiO, SiO, HY y-irradiated HY, HZ y-irradiated NaY NiCaX WHY WNaY CeX CeX Co ammonia adducts in Y zeolite Co amine adducts in Y zeolite RuY Pd mordenite MgO(CO2 + 0 2 ) ThO,(CO + 0,) A ThO,(CO + 0,) B TiO,(CO + 0,) TiO,(RCH,OO') TiO,( RR'CHOO') CrO,/SiO,(CO, + 0,) MoO,/SiO,(CO, + 0,) MoO,/SiO,(CO, + 0,) HY y-irradiated (CO + 0,) Organic peroxy Tetralin peroxy 1
I1 Triphenyl methyl peroxy 120 K 300 K Polytetrafluoroethylene peroxy (chain radical) 77 K 300 K (propagating radical) 71 K 300 K Polypropylene peroxy Enzyme Protein y-irradiated
FOO' CF,C000' (CH,),COO' ROO' C,H,(CH,),COO'
4,"
Ayya
aiaaa
Reference
74 77 82,63 84.5, 64.2 76 80 83 75 78 66 80.60 72 80,67 77 100, 50 95,65 75 104.42.5 95.35 94, 36 98,42 104,40 101,39.5 107, 37
61 255 81 83a 76 283 105b lO5b 102a 102a 51
13, 5
116 83b I52 63 62 62 88
-
90 90 86 87 447 89
87,59 88,60
1
91.61 40.29
}
449
107.46 89,40
}
58,93
107,46 98.60 72 68.46
26.5, 13
22.17, 14.5 23.3, 14.0 21.8, 16.4 23,18 21.8, 16.4
1
92
j8
84 100 101 450 45 1 946 452 94b
(Continued)
130
M. CHE AND A. J. TENCH
APPENDIX B (Continued) ~~~
Systems
Ax,
Inorganic peroxy' Bu'O(Ph,),AsOO' 93 K Bu'O(Me),AsOO' frozen solution Bu'O(OMe), POO' frozen Bu'O(Ph),POO' frozen 153 K Co oxygen carriers 71 K
A,,"
4,"
alsoa
Reference
91.61 86.5, 64.4 24.4, 18.6 76.2, 69.3 85.70 24.5, 15.5 88,60
300 K
4530
21.6
49 1026
~~
Only the absolute value of the hyperfine tensor is given; for the problem of the sign and the possible presence of a motion refer to original papers and discussion in Section III,A,2; for equivalent oxygen nuclei only one value is given. The systems are arranged using the same order as in Section IV. Bu' = tert-butyl; Ph = phenyl; Me = methyl.
Appendix C.
Characterization of Oxygen Species by Infrared Spectroscopy*
Although EPR has turned out to be the most important technique used so far in the characterization of adsorbed oxygen, there are a number of cases where it cannot be applied, for instance when paramagnetic oxygen radicals lead to linewidths broadened beyond detection or when the species are diamagnetic. In such cases IR has proved very useful, although the identification of the adsorbed species is not straightforward since the IR frequency can vary over a wide range. The various vibrational frequencies involving adsorbed oxygen are listed in Table XV together with those related to various model systems (gas phase, oxygen carriers, solid state and matrix isolated species), while Fig. 28 gives the frequency ranges of the oxygen species observed for the various sytems. The coordination of oxygen to transition metal ions which occurs mostly in the side-on fashion on surfaces (Section III,A,2 and Appendix B) can be described following the model of acetylene-metal complexes (467). Both nu and ng orbitals of molecular oxygen have proper symmetry to interact with the bonding set of s, p, and d orbitals on the metal. The bonding orbitals are shown in Fig. 29.
* See Ref. 4536.
MOLECULAR OXYGEN SPECIES ON OXIDE SURFACES
TABLE XV Selecird Vihrarional Frequencies ,for Species lnuolving O.xy(jrti Bods
Species Gas phase ('AJO, ("z, )O,
voo (cm
Reference
~
1483.5 1555 1586.I - 1596.6 1876
21 226 454
780-884
148
800-932
148
M O (superoxide-like) 0
1075-1122
148
0
1130-1 I95
148
738-794 800-900 900-1100 1137-1164 1825-1864
1496 1 and Refs. therein 1 and Refs. therein 21,132,455 456,457
(OdZ
0: Oxygen carriers O M
/ \ / M O (peroxide4 ke) 0 0
24f'
\ / M (peroxide-like) O M
/ k. /
9
I M (superoxide-like) Solid state
0: ~
M-0-M M=O 0; 0: Matrix isolated species 0 0
983
315
990-11 I5
127a, 133a-c, 458
\ / M
(peroxide-like) 0; (alkali and silver superoxides) Adsorbed species 0;M-0-M M=O 0; "Neutral" 0,
640-970 750-900 900-1 100 1015-1180 1460-1700
459-462 463-465 24e, 134d, 465 19a,b, 206, 24e, 134e, 463 20a,24d,e, 466
131
M. CHE AND A. J. TENCH
132 adsorbed oxyqen
‘
M-0-M
;-
1
MzO
L
1
0-
0
I
p e r t u r b e d 05 i
1
matrix i f o l a t e d l
n
I
n e u t r a l o2 1
02
species
1
I
M-0-M
solid
M= 0
’
i
oi
0;
oxyqen corriers M
(solution)
9-P
FIG.28. Infrared frequencies for species involving oxygen bonds (4536).
A D bond is formed by transfer of electron density from a filled dioxygen nu bonding orbital to s, p, and d orbitals of appropriate symmetry on the metal and two n bonds are formed by transfer from filled metal d orbitals into unfilled n g antibonding orbitals of dioxygen. This synergic bonding mechanism involves the drift of metal electrons (referred to as “n back-bonding”) into oxygen orbitals, thus making 0, as a whole negative, and at the same time the drift of electrons to the metal in the D bond, thus making O2 positive. The combination of both the D donor and n acceptor effects (relative to 0,) may lead to a large range of IR frequencies since the simple donation from dioxygen to metal will increase the voo stretching frequency as it does for CO and NO on Lewis acid centers, while the n back-bonding from metal to dioxygen will decrease the voo frequency (18b).
L FIG.29. The orbitals involved in the bonding of dioxygen to transition metal ions.
MOLECULAR OXYGEN SPECIES ON OXIDE SURFACES
133
The energy of the metal orbitals determines the extent of electron transfer between metal and oxygen (134c) and depends on the oxidation state of the transition metal ion, its coordination number, and the donor properties and strength of bonding of its ligands. For instance, the 71 back-bonding is expected to decrease when the oxidation state of the metal increases, e.g., when there is a decrease in the number of d electrons available (186).Similarly, the II back-bonding decreases on decreasing the coordination of the metal bonded mainly to 0 - donor ligands (468).It also decreases when the electron acceptor character of the ligand attached to the metal increases (1344. The adsorption of molecular oxygen on an oxide involves (in most cases) an oxidation of the metal with a concomitant reduction of the adsorbed oxygen. This arises from two factors: (1) Oxygen adsorption on oxides is observed mainly for oxides with metals in a reduced oxidation state, and (2) there are readily accessible reduced oxidation states for oxygen, i.e., 0;and 0; -. Thus, depending on the number of electrons involved in the adsorption, the nature of the oxide, and the relative energies of the metal valence s, p, and d orbitals and the dioxygen II, and n, orbitals, it is possible to envisage the transfer of none, one, two, or more electrons from the reduced oxide to the coordinated dioxygen moiety leading to the formation of O,, O i - , or to dissociative adsorption giving rise to M=O or M-0-M species. The IR data observed for oxygen adsorbed on oxides depend on a number of factors, such as the nature of the oxide (204. the pretreatment conditions (20b),and the temperature of oxygen adsorption ( 1 4 9 4 and this results in a wide range of frequencies. Although Drago has pointed out that the IR data should be used with caution (309),it is possible to give a reasonable assignment of the IR data for adsorbed oxygen by comparison with IR frequencies observed for the model systems given in Fig. 28. However, the similarity of the frequency values for example for M=O bonds and 0-0 bonds in certain dioxygen species shows that it is difficult to distinguish between mononuclear and molecular species. The nature of adsorbed oxygen can only be inferred from careful interpretation of experiments using 160/'80 isotopic mixtures. It is also important, whenever possible, to associate other techniques such as UV-visible or EPR in order to unambiguously identify the nature of adsorbed oxygen species. The data given in Table XV and Fig. 28 for adsorbed oxygen have been confirmed by experiments using l60/l8O isotopic mixtures in a few cases only (24d,e, 134d,e, 464) and should be used with caution. For instance, bands observed at 960-1200 cm-' were reported on SnO, and assigned to adsorbed 0;(469),but later experiments with "0-enriched oxygen did not confirm this assignment (464).Similarly, a band at 985 cm-' on Cr,03 was attributed to 0;- ( 1 4 9 ~ )Subsequent . work using isotopic labeling with
I34
M. CHE AND A. J. TENCH
160/’80 (134d)has not confirmed this assignment and it is more consistent with a species such as CFO. The vibrational data of adsorbed dioxygen have recently been reviewed by Busca ( f 8 h ) .
ACKNOWLEDGMENTS The authors acknowledge the facilities provided by AERE Harwell and UniversitC de Paris VI during the writing of this review; A. J. Tench acknowledges an appointment as Associate Professor at the Universitt de Paris VI and M. Che Vacation Associate appointments at Harwell. They are also very grateful to a number of people for helpful discussions, and in particular to Dr. C. B. Amphlett for encouragement and comment. Finally, they appreciate very much the help and support of their families during the writing of this review. 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.
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MOLECULAR OXYGEN SPECIES ON OXIDE SURFACES
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M. CHE AND A. J. TENCH
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398. Lyubimova, 0. l., Kotov, A. G., and Pshezhetskii, S. Ya., Kinet. Katal. 13, 1603 (1972). 399. Akimoto, M., and Echigoya, E., J. Chem. SOC.Faraday Trans. 173,193 (1977). 400. Parkes, D., J. Chem. SOC.Faraduy Trans. 168,613 (1972). 401. Takita, Y., and Lunsford, J. H., J. Phys. Chem. 83,683 (1979). 402. Aika, K., and Lunsford, J. H., J. Phys. Chem. 81, 1393 (1977). 403. Takita, Y., Iwamoto, M., and Lunsford, J. H., J. Phys. Chem. 84,1710 (1980). 404a. Ben Taarit, Y., Symons, M. C. R., and Tench, A. J., J. Chem. SOC.Faraday Trans. I73, 1149 (1977). 4046. Kuznicki, S. M., and Eyring, E. M., J . Am. Chem. SOC.100,6790 (1978). 405. Bickley, R. I., Catalysis (London) 5,308 (1982). 406. Formenti, M., and Teichner, S. J., Catalysis (London) 2, 87 (1978). 407. Breakspere, R. J., and Hassan, L. A. R., Aust. J. Chem. 30,971 (1977). 408. Munuera, G.,Gonzalez-Elipe, A. R., Soria, J., and Sanz, J., J. Chem. SOC.Faruday Trans. 176, 1535 (1980). 409. Haber, J., Kosinski, K., and Rusiecka, M., Discuss. Faruday SOC.58, 151 (1974). 410. Kubokawa, Y., Anpo, M., and Yun, C., Proc. Inr. Congr. Catul. 7th, 1980 B, 1170 (1981). 411. Tanaka, K., J. Phys. Chem. 78,555 (1974). 412. Tanaka, K., and Miyahara, K., J. Phys. Chem. 78,2303 (1974). 413. Courbon, H., Formenti, M., and Pichat, P., J. Phys. Chem. 81,550 (1977). 414. Herrmann, J. M., Disdier, J., and Pichat, P.,Proc. Int. Vuc. Congr., 7th. Int. Conf. Solid Surf, 3rd 2,951 (1977). 415. Herrmann, J. M., Disdier, J., and Pichat, P., J. Chem. SOC.Faruday Trans. 177, 2815 (1981). 416. Formenti, M., Juillet, F., and Teichner, S. J., BUN. SOC.Chim. pp. 1031, 1315 (1976). 417. Djeghri, N., and Teichner, S. J., J. Catal. 62, 99 (1980). 418. Yun, C., Anpo, M., Kodama, S., and Kubokawa, Y . , Chem. Commun.p. 609 (1980). 419. Volodin, A. M., Cherkashin, A. E., Zakharenko, V. S., React. Kinet. Catal. Lett. 11,277 (1979). 420. Wandelt, K.,Surf Sci. Rep. 2, 1 (1982). 421. Roberts, M. W., Sci. Prog. (Oxford) 68,65 (1982). 422. Au, C. T., Roberts, M. W., and Zhu, A. R., Surf. Sci. 115, LI 17 (1982). 423. Brundle, C. R., and Metcalf, L. P., J. Chem. SOC.Faruday Trans. I1 75, 1030 (1979). 424. Law, D. S., Lee, E. P. F., and Potts, A. W.,J. Chem. SOC.Furuduy Trans. 1178, 2101 ( I 982). 425a. lnoue, Y., and Yasumori, I., Bull. Chem. SOC.Jpn. 54, 1505 (1981). 4256. Tatibouet, J. M., and Germain, I. E., J. Chem. Res. ( S ) p. 268 (1981). (M) 3070 (1981). 425c. Tatibouet, J. M., Germain, J. E., and Volta, J. C., J. Catal. 82,240 (1983). 425d. Tatibouet, J. M., and Germain, J. E., J. Catal. 72,375 (1981). 425e. Tatibouet, J. M., and Germain, J. E., C.R. Acad. Sci. Ser. 11296,613 (1983). 425’ Volta, J. C., Desquesnes, W., Moraweck, B., and Tatibouet. J. M., Proc. Int. Congr. Catal., 7rh, 1980, E, 1398 (1981). 4258. Volta, J. C . , Forissier, M., Theobald, F., and Pham, T. P., Discuss. Furaday SOC.72,225 (1981). 425h. Phichitkul, C., Tatibouet, J. M.. and Germain, J. E., to be published. 425i. Murakami, Y., lnomata, Y.,Miyamoto, A., and Mori, K., Proc. Int. Congr. Catul., 7th. 1980, E, 1344 (1981). 425j. Miyamoto, A., Ui, T., and Murakami, Y., J. Carul. 80, 106 (1983). 425k. Volta, J. C.. and Morawek, B., Chem. Commun.p. 338 (1980). 4251. Marques, A. R., Davignon, L., and Djega-Mariadassou, G., J. Chem. Soc., Faraday Trans. 178,598 (1982). 426. Levine, J. D., and Mark, P., Phys. Rev. 144,751 (1966).
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p 325. McGraw-Hill, New York. 1979. Grasselli. R . K.. Cherii. Em/. Nrii..s 56, 49 (1978). Grasselli. R. K.. A . A . A . S . h r i u . M w i . . Jtrn. 4/h ( 1980). Grasselli. R. K.. and Burrington. J. D.. A h . Cu/u/.30, 133 (1981). Cullis. C. F.. and Hucknall. D. J.. Coral. (London) 5, 273 (1982). Mars. P.. and van Krevelen. D . W.. Chem. Enq. Sci. Suppl. 3,41 (1954). Sachtler. W. M. H.. and de Boer, N . H.. Proc. In/. Conqr. Ca/ul,.3rd. 1964. I , 252 (1965). Aykan. K.. J . Cu/irl. 12, 281 (1968). Grasselli. R. K.. and Suresh, I). D.. J . Card. 25. 273 (1972). Jiru. P., Wichterlova. B., and Tichy. J., Proc. Inr. Corryr. Coral.. 3 r d 1064, I , 199 (1965). Libre, J . M.. Barbaux. Y., Grzybowska, B.. and Bonnelle. J . P., to be published. Brazdil. J . F., Suresh, D. D.. and Grasselli. R. K.. J . Caral. 66, 347 (1980). Haber. J., and Witko, M., Act,. Chem. Res. 14. I (1981). Sancier. K. M.. Wentrcek. P. R.. and Wise, H.. J . C a r d 39, 141 (1975). Keulks. G . W.. J. Cord. 19, 232 (1970). Keulks. G . W.. and Krenzke. L. D.. Pro(,.In!. Congr. Carol., 6/11.1Y76. 2,806. 814( 1977). 443. Hoefs, E., Monnier. J. R., and Keulks. G . W., J. Cu/al. 57, 331 (1979). 444. Tagawa. T.. Hattori. T.. and Murakami, Y., J . Cnml. 75.66 (1982). 445. Akimoto. M.. and Echigoya, E.. J. Cum/. 35,278 (1974). 446. Bennett. J . E.. Ingram, D. J. E., and Schonland. D.. Proc. Phys. Soc. 69A, 556 (1956). 447. Lipatkina. N. I., Shubin. V. E., Shvets. V. A,, Chuvylkin, N . D.. and Kazansky. V. B.. Kina/. k’u/u/.23, 670 (1982). 448. Reuveni. A.. Luz, Z.. and Silver. B. L.. J . Mir~qnReson. 12, 109 ( 1973). 44Y. Melamud. E . . and Silver. B. L.. J. M a p R c w n . 14, I12 (1974). 450. Adrian. F. J .. J . Clitwr. P h j : ~46, . 1543 ( 1967). 451. Fessenden. R. W., J . Chon. Phjx. 48. 3725 (1968). 452. Fessenden. R. W.. and Schuler. R. H.. J. Chcrn. Phys. 44,434 (1966). 453a. Howard, J. A.. and Tail. J. C., Can.J. Chem. 56, 2163 (1978). 4536. Che, M., Kermarec, M.. Dyrek, K., and Tench. A. J.. Reo. Chim. Miner., in press. 454. Herzberg, G., “Molecular Spectra and Molecular Structure,’’ p. 560. Van NostrandReinhold, Princeton, N.J.. 1950. 455. Creighton, J. A,. and Lippincott, E. R.. J. Chem. Phys. 40, 1779 (1964). 456. Shamir. J.. Binenboym. J.. Claasen, H. H.. J. Am. Chem. Sor. 90, 6223 (1968). 457. Edwards. A. J., Falconer, W. E., Griffiths, J. E., Sunder, W. A,, and Vasile, M. J.. J. Chem. Soc., Dulron Trans. p. I129 (1974). 458. Andrews. L.. J. Phys. Chem. 73,3922 ( 1 969). 459. Metcalfe, A,, and Ude Shankar. S., J . Chem. Soc. Faraday Trans. 176,630 (1980). 460. Gland, J . L., Sexton, B. A., and Fisher, G. B., Surf: Sci. 95, 587 (1980). 461. Sexton, B. A,, and Madix. R. J., Chem. Phvs. Lett. 76,294 (1980). 462. Backx, C., de Groot. P. P. M., and Biloen, P.. Sur/: Sci. 104, 300 (1981). 463. Howe, R. F., Liddy, J. P.. and Metcalfe, A,, J. Chem. Soc. Faradav Trans. 168, 1595 (1972). 464. Harrison, P. G . , and Thornton. E. W., J. Cham. Soc. Faraday Trans. I74.2597 (1978). 465. Zecchina, A,, Coluccia, S., Cerruti, L., and Borello, E., J. Phys. Chem. 75, 2788 (1971). 466. Forster, H., and Schuldt, M., J. Chem. Phys. 66,5237 (1977). 467. Greaves, E. 0.. Lock, C. J. L., and Maitlis, P. M., Can. J . Chem. 46, 3879 (1968). 468. Kozuka, M., and Nakamoto, K., J . Am. Chem. Soc. 103,2162 (1981). 469. Gundrizer, T. A., and Davydov. A. A,. React. Kinet. Caral. Lett. 3, 63 (1975). 42Y. 430. 431. 432. 433. 434. 435. 436. 437a. 4376. 438. 43Y. 440. 441. 44-7.
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NOTEADDEDIN PRCOF The nature of the active species in the heterogeneous epoxidation of ethylene is still the subject of active debate. Using a UHV chamber linked to a high-pressure reactor cell, R. B. Grant and R. M. Lambert [Chem. Commun. p. 622 (1983)] have investigated this reaction on the (1 11) face of a silver single crystal. They conclude that chemisorbed atomic oxygen is the crucial surface species which selectively oxidizes ethylene to ethylene oxide, whereas adsorbed dioxygen plays no direct role in this reaction. This conclusion differs from that obtained in particular by Kilty e/ a/. (293) in the case of supported silver (see Sections IV,D, VI,B, VII, and VII1,B). The EPR spectra of the molecular 0;ion have now been obtained by reacting MO, + 0, (M = Na, K, Rb. Cs) in rare gas and nitrogen matrices [D. M. Lindsay, D. R. Herschbach, and A. L. Kwiram, J. Phys. Chem. 87,2113 (1983)]. Both the g and alkali hyperfine tensors suggest a dominantly ionic product M '0;. The EPR data are interpreted in terms of a model (0, - 0,)structure in which a relatively weak bond connects two equivalent 0, moieties. The EPR spectra do not allow one to distinguish between cis- and /rans-O;, but symmetry restrictions may preclude formation of the cis isomer. The g tensor of 0, differs substantially from those observed for the isoelectronic 25-electron radicals SO;, CIO,, and PO:- (see Section V,B). Finally, the origin of adsorbed oxygen on iron oxides has been further investigated by Borello and co-workers [C. Morterra, C. Mirra, and E. Borello, Chem. Commun. p. 767 (1983)l. It is found that an I R band observable at 1140 cm-' on a-FeOOH (goethite) which undergoes several reversible splittings and shifts on dehydration of the sample is the precursor of similar bands previously observed on a-Fe,O, (haematite) and assigned to adsorbed molecular oxygen (see Sections 11. 1V.B. and Appendix C).
ADVANCES IN CATALYSIS. VOLUME 32
Catalysis by Alloys in Hydrocarbon Reactions VLADIMIR PONEC Gorlaeus Laboratoria Rijksuniversiteit Leiden Leiden, The Netherlands
. . A. Electronic Structure of Alloys: Experimental Aspects . B. The Texture and Surface Composition of Alloys . . . C. Progress in the Theory of Alloys . . . . . . .
. . . . . D. Effects of Alloying o n the Chemisorption Bond Strength , . 111. Particle Size Effects. . . . . . . . . . . . . . I . Introduction.
11. Alloys
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A. Electronic Structure ol' Small Particles: Experimental Aspects and Theory . . . . . . . B. Effects of Particle Size on Chemisorption Behavior . . IV. Mechanism of Hydrocarbon-Hydrogen Reactions . . . . A. Introduction . . . . . . . . . . . . . B. Chemisorption Complexes of Hydrocarbons on Metals . C. Mechanism ol' Skeletal Reactions and Selectivity of Metals D. Particle Size EKects in Catalysis of the H C Reactions. . V. Hydrocarbon Reactions on Alloys . . . . . . . . A. Classification of the Reactions on Metals and General Description of Alloying EKects . . . . . B. Some Particular Alloy Systems . . . . . . . . C. Open Problems in Catalysis by Alloys . . . . . . D. A Speculative Model of Hydrocarbon Reactions on Metals and Alloys . . . . . . . . . . . VI. Conclusions . . . . . . . . . . . . . . . References . . . . . . . . . . . . . . .
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Introduction
Catalysis by metals and alloys plays an important role in industry as well as in laboratory-scale preparations. Catalyzed reactions are usually run at lower temperatures than the noncataiyzed ones and they are also more I49 Copyright 6 1983 by Academic Press. Inc. All rights 01reproduction in any form reserved. ISBN 0-12-007832-5
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selective.’ It is mainly the selectivity aspect that makes catalysis so important, but nevertheless, until recently, very few studies were devoted directly to selectivity. This was so because of the general belief that the problems of selectivity will be elucidated automatically one day, when the problems of activity have been solved. However, this appears to be too simplified an idea. In metal catalysis, for example, the activity of a given metal for a certain reaction is very often determined by the selectivity for the main reaction and for the side reaction of self-poisoningof the metal, and not by its activity in a given reaction only. Research on alloys as catalysts has recently contributed very much to the identification of the factors which determine the selectivity and, by that, the activity of metals. This progress, in particular in the field of hydrocarbon reactions, will be reviewed below. Research on alloy catalysts started in the 1950s with attempts to investigate the role in catalysis of the electronic structure of metals. This research was initiated by several papers of Dowden which, measured by their response in the literature, rank among the most important papers ever written on catalysis. However, it appeared later (for reviews, see 1-5) that two basic ideas, on which the so-called “electronic theory of catalysis” was built up, were not correct. These ideas were as follows : The rigid band theory (RBT) of solids according to which there is a considerable transfer of electrons among the alloy components ; an alloy surface is then a structureless plane with atoms indistinguishable for gas molecules. (2) The idea that to activate a molecule, an electron has to be transferred to or away from it, the main function of the catalyst being to mediate the transfer of electrons among reaction components. Due to the frequent use of inadequate techniques in the preparation and characterization of alloys, very controversial results were obtained. This, together with the failure of the above-noted ideas has led to a certain crisis in alloy research (see 1-5) and a loss of interest in this kind of investigations. About 10 years ago, several sources brought about a renaissance in alloy research : (1) A renewed industrial interest in catalysis by alloys or more generally
by bimetallic catalysts (see, e.g., 6). (2) Considerable progress achieved in the quantum theory of alloys and in the theory predicting the surface composition of alloys (see, e.g., 5 for a review).
’ Note: The ucrioity of a catalyst is usually defined as a rate per unit surface area or per site, measured at standard experimental conditions. The selectiuity is then a (normalized) ratio of rates or product concentrations in different reactions running simultaneously.
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(3) Fundamental investigations demonstrated that by alloying, dramatic changes can be achieved in the selectivity of metal catalysts (1-5, 7-9). It is not our intention to repeat all these results, including those already reviewed (1-3, and therefore only the most relevant and the most recent results will be discussed below. II. A.
Alloys
ELECTRONIC STRUCTURE OF ALLOYS : EXPERIMENTAL ASPECTS
From the various methods to be used to investigate the electronic structure of metals, probably the ultraviolet photoelectron spectroscopy (UPS) and X-ray photoelectron spectroscopy (XPS) methods brought forth the information most relevant for catalysis and surface science. These methods are best suited to monitor the changes in characteristics parameters of the d-bands by alloying, and since the most catalytically active metals are transition metals where d-orbitals are the frontier orbitals (Fermi level is cutting the d-band), the interest in these methods is not incidental. Since the pioneering paper by Seib and Spicer (10) convincingly demonstrated that the RBT did not hold, many other papers confirmed this conclusion and helped to create a new picture of the electronic structure of alloys. The main points may be summarized with the help of Fig. 1 as follows: (a) For endothermically formed alloys as well as for weakly exothermic cases, the position of the d-band does not change by alloying (Fig. 1). Due to the physics of the photoemission process, it is approximately the density of states N(E), in the d-band which is reflected by the distribution of the photoemit ted electrons I( E ) . (b) With these alloys the main change by alloying is the narrowing of the d-band (lower 6 ) . This indicates the decreasing overlap of the d-orbitals when the neighboring positions around a given metal atom are occupied by another component of the alloy. This effect-leading to an increase of the local density of states in a certain energy range-may influence the phenomena sensitive to the electron shift from the adsorbates to the metal and vice versa (21). This is, of course, a second-order effect as compared to the effects caused by the changes in position of the d-band (there, where they occur). Notice also that effects due to band narrowing are “collective” or “band” effects (in contrast to “local” or “ligand” effects) in the variations of the electronic structure of alloys. The above-listed conclusions are best demonstrated by the results of several
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EF.O
E
FIG. I . Photoemission from the valence bands of metals and alloys. Intensity of emission as a function of energy. E, is the Fermi energy; 6 is the bandwidth (schematically). Reprinted from Ref. 21.
papers (this is a selective, not a full list of references) which should be mentioned in this context (10, 12-16). It has also been signalized that in systems like Pt-Cu (17), where the position of the d-band does not change by alloying too much and where evidently no transfer of electrons from one component to the other takes place (by the way, there is no unequivocal evidence for such a transfer with any alloy of the group of alloys just discussed), alloying can cause certain rehybridization on one (or both) of the alloy components (Cu). Redistribution of electrons between the orbitals of predominantly s or d character is, of course, observed quite frequently (see Pd-Au, Pd-Ag) (12- 14).
A third conclusion is the following:
(c) Some intermetallic compounds reveal more pronounced changes due to alloying [Ni-A1 (18), Ni-Ga (18), Pd-Zr (19), etc.]: The position of d- (or p-) bands shifts by alloying and simulates changes, to be observed in the spectra of atoms while forming a chemical compound. If the data, conclusion (a) and (b) are considered, it is actually not that surprising that atoms of the solution alloys [discussed under (a) and (b)] preserve most of their individual chemisorption and catalytic properties also in alloys (for review, see 4, 5, 20, 21). It is, of course, more interesting what
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will happen with the intermetallic compounds. Unfortunately, it is too difficult to make these alloys in such a way that these alloys are homogeneous and any clustering of the Group VIII metal is excluded. But if the authors of Ref. 22 succeeded, indeed, in achieving that, their paper would then show that even formation of an intermetallic compound does not completely suppress, e.g., the ability of a Group VIII metal to form single and multiple metal-carbon bonds-a very essential feature of the formation of hydrocarbon chemisorption complexes. The same conclusions with the same remark of caution can also be drawn from the papers on the catalytic behavior of Pt-Sn (without carrier) alloys (23, 24).
B. THETEXTURE A N D SURFACE COMPOSITION OF ALLOYS Very small bimetallic particles on carrier are often X-ray diffraction amorphous and it is not easy to gain any information on their composition and structure. In spite of these difficulties, very important information on the texture of bimetallic particles has been obtained by extended X-ray absorption fine structure (EXAFS) (25-28). Results obtained by Sinfelt et al. (25,26) demonstrated that a very detailed picture can be obtained of, e.g., Ru-Cu catalysts by using this advanced technique, which may shortly appear to be the most important for the study of multicomponent catalysts. In the field of alloy surface composition, both theory and experimental determination achieved much progress in recent years. The present “state of the art” does not, unfortunately, allow one to predict quantitatively the surface composition from the bulk concentrations, but calculations on models allow one to estimate various effects and to make interesting conclusions and sometimes even semiquantitative predictions. The calculations are rather easy and have already been performed for models like (1) the ideal solution model where enrichment is always confined to the outmost layer (29),(2) the ideal or regular solution model with onelayer enrichment, taking into account the difference in atomic radii (strain energy) (30-32), (3) the regular solution model with enrichment spread over n (up to 4) layers (33), and (4) intermetallic compounds (37). For complicated systems semiempirical rules based on the phase diagrams (34) or data on the diffusion coefficient (35)might sometimes be quite useful. The equations binding bulk and surface concentration together have been derived by applying classical or statistical thermodynamics (33, 36, 37), kinetic considerations (5), or the Monte Carlo technique (38). The energy gain due to surface enrichment has usually been calculated by the “broken” model assuming additivity of bond enthalpies per bond pair. However,
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recently a quantum-mechanical estimate was also performed for this parameter and these calculations then allow one to go beyond the approximation of pairwise bonding (see, e.g., 39, 40). The early papers usually assumed that surface concentration can be estimated from the normalized ratio of the Auger peaks (41,42). However, soon it appeared that this had led to incorrect conclusions on, e.g., Ni-Cu (43) or Pd-Ag (44) alloys. In particular, the progress with pure Ni-Cu alloy systems suffered very much from the uncertainties caused by improper procedures. The main point to keep in mind in this respect is that the ratio of peak intensities for metals A and B in alloys related to those in the pure metals, as seen by Auger electron spectroscopy (AES), is given by (45, 47) the following equation :
or, when enrichment is confined to the first layer only, by
where Ni is the fraction of the total signal originating in the ith layer, Nrnetal is the number of atoms in one layer of a pure metal, and X is the A.B molar fraction. Fraction Ni can be calculated either by using a discontinuous Gallon model or a model of continuous attenuation of the signal (45, 47). When applying Eq. (l),independent information on the bulk depth profile of the enrichment is necessary ; or when another method (chemisorption, ion scattering) supplies the information on the surface concentration, Auger spectrometry can serve to make some estimate on this depth profile (43). Sometimes, a form of the depth profile is assumed and roughly checked by Auger spectral measurements at two different electron energies or at two different escape angles (46).A full discussion of these and related problems, as well as a review on the methods available for surface composition determination, is published elsewhere (47). Development in the field of measurements on metals without carrier can now be considered as satisfying and there is still some progress going on. However, the problem of a reliable determination of the surface composition of alloys on carriers is still too far from being solved. In particular, problems like the detection of small amounts of unalloyed active metals on carrier and the question of homogeneity in distribution of the active metals in an inactive matrix have not been solved yet, and just such problems are most likely responsible for some controversies in the results on alloys. More work has to be done in the future in this field.
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C. PROGRESS IN THE THEORY OF ALLOYS The most essential progress from the point of view of application of this theory in catalysis and chemisorption has actually been achieved by the very first papers (48-50), where the so-called coherent potential approximation (CPA) was developed and applied. By means of this, photoemission data were explained in a quite satisfying way and the catalytic research got full theoretical support for some of the ideas introduced in catalysis earlier on only semiempirical grounds ( 3 ) ; namely, individual components are distinguishable for molecules from the gas phase and the alloy atoms preserve very much of their metallic individuality also in alloys-something that was impossible according to the RBT and the early electronic theory of catalysis. It might be of some interest for the catalytic research that in the course of time the CPA theory was further modified and developed and also some alternative approaches were suggested. The most sophisticated versions of theory now also comprise effects like short- and long-range ordering, clustering, etc. This improves the agreement between theory and experimental data on the electronic structure (mainly UPS and XPS data), but does not change in any way the main conclusions mentioned above. The development of the alloy theory is best demonstrated by a selection of papers (51-57) or by a review (58). How was this development reflected by the theory of catalysis on alloys? An early and very important paper (9) discussed the selectivity and activity effects fully in terms of the old electronic theory of catalysis. Another paper (a), which appeared simultaneously with (9), turned attention to the fact that one must also consider effects other than only the changes in the electronic structure. The results on alloys should be rationalized on the basis of two aspects of alloying (8):
( I ) By alloying a metal A is dispersed (more or less, it depends on the type of alloys) in a metal B. If a certain reaction requires a big ensemble of contiguous atoms A in the surface of alloys, this reaction will be suppressed strongly by alloying. This may lead to selectivity changes, if other potential reactions in the system can occur on smaller ensembles or even individual atoms. This is true for systems when B is much less active than A. If both components are active, one has to consider also the possibility that a big ensemble required can be formed by a mixture of A and B. In some cases (Pt/Ir, Pd/Ni, Pt/Re, . . . ?) the mixed ensembles may even be suspected to be more active than the one-component ensembles. In the literature, this kind of effect is called an “ensemble size” effect (1-5). (2) In spite of the fact that the electron transfer among alloy components is much less frequent (or much less pronounced) than envisaged by the RBT,
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and that the individuality of components is preserved to a high extent also in alloys, small effects of alloying in the electronic structure of metals must always be considered. The question to be analyzed carefully is the following : How important are these (often just postulated) changes for (a) chemisorption and (b) a catalytic reaction?
To stress the localized character of chemisorption (a term surface pseudomolecules was introduced at that time), Sachtler introduced for the alloying effects discussed in paragraph (2) a term “ligand effect” (3). It was then a task for an experimentalist to establish how important-relatively-the effects (1) and (2) were. A general consensus now is that effect ( 1 ) is more essential than (2) in any case, but the discussion is still going on, on the reliability of some pieces of evidence which have been presented in the literature in favor of a role for effect (2).
D. EFFECTS OF ALLOYING ON THE CHEMISORPTION BONDSTRENGTH The simplest case to study is hydrogen. Earlier papers, where heats of adsorption were measured calorimetrically (59, 60), reported a decrease of the heat of adsorption when Ni was alloyed with Cu. The qualitatively same result was obtained later in a very detailed study by Prinsloo and Gravelle (61), although the decrease of the heat was less pronounced here. However, the thermal desorption studies revealed that on a pure Ni surface several states of Hadsare formed, each characterized by a peak maximum or a peak shoulder in the thermal programmed desorption (TPD) spectra, and when Ni is alloyed with Cu, the various states stay where they were on Ni and only their populations change. This has been found actually first for Pt/Au (62), but later also for Ni/Cu (63). The studies on monocrystal Ni/Cu planes (64-66) lead to the conclusion that the changes in the binding strength of individual states are either negligible or of a moderate size (10-15%; i.e., of a similar order of magnitude to variations in the heats of adsorption with the crystallographic planes of the small metal). It is a fact worth noting that the effect on the hydrogen heat of adsorption of a coadsorbed CO and the effect of replacing the Ni atom by a Cu atom are both the same. This is as if just dilution of the hydrogen layer caused a decrease in heats of adsorption. This is an effect one would expect if lateral attractive forces existed in the adsorbed layer and were suppressed by alloying or coadsorption. Another gas easy to measure is CO. This does not mean that better agreement has already been achieved here. One group of authors (67) did not find any systematic variation in the CO heat of adsorption with alloy composition of evaporated metal films; other authors (68) found a moderate
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variation (similar to that of hydrogen). A more pronounced variation was reported by Yu et al. (69), but this is in part due to the fact that Ni and Ni/Cu alloys were measured with different samples and in different apparatuses. However, the following remark must be made here. Those who worked with monocrystals, prepared varying compositions by first sputtering Cu away and then, by successive anneals, bringing it back into the surface. Since it is known (70) that the presence of defects increases the heat of adsorption of CO and brings about new states [found, indeed, in (69)], the data with sputtered and annealed monocrystals have to be discussed with some caution. Nevertheless, the conclusion can be made, very similar to that with hydrogen, that the variations in heats of adsorption with alloying are not very pronounced. This is also the conclusion of a most recent and very detailed study by Eley and Moore (71), who stress that the lowering of the heat of adsorption of CO with alloying (Pd/Au) is less pronounced than the observed decrease in the extent of adsorption. Before a conclusion is made that the above-mentioned findings evidence the variation by alloying in the chemisorption bond strength, the following must be considered. The heat of adsorption is an overall effect which comprises several contributions, including the heat of mutual interactions. Now, it has been shown recently that adsorption of CO and H, are accompanied by attractive and repulsive lateral interactions (72-74). With hydrogen, these two are in better balance; with CO the repulsive interactions (mainly, electrodynamic and electrostatic interaction of dipoles) clearly prevail. The cooperative action of both attractive and repulsive forces leads to the formation of ordered (i.e., observable by LEED) domains at rather low coverages. When a layer of adsorbed hydrogen is diluted by CO or by empty Cu sites, the measured heat may be lower if the attractive interactions are suppressed by it more than the repulsive ones. With CO, the opposite assumption would explain the observations. Sachtler and Somorjai (75) studied this question in great detail, with Pt/Au alloys. When an Au layer was epitaxially grown on Pt, the heats of various types of CO adsorption were independent of alloy composition. However, when Au was spread in and over the surface of Pt by annealing, the heats of CO adsorption were higher. Evidently, in the first case clusters of Pt atoms were sufficiently big to allow CO-CO interactions (leading to the clustering of CO) to occur freely; in the second case dilution of Pt in Au kept CO molecules a distance from each other, and at the same CO dosage heats of CO adsorption were higher. It is interesting to note at this point that also the selectivity effects were observable only with the “annealed” and not with the “epitaxial” alloys. It is questionable whether the heat measurements (calorimetric or by TPD) are sensitive enough to detect changes in the binding strength due to
158
VLADIMIR PONEC
alloying. Infrared (IR) spectra of adsorbed CO are usually mentioned as being better suited for the detection of small "ligand" effects of alloying (i.e., of small localized changes in the electronic structure due to the alloying). Indeed, numerous papers (results are summarized and analyzed in 11) exist showing that the frequency v(M/CO), where M stands for a Group VIII metal, decreases with increasing amount of the second alloy component (Group VIII or IB metal). This has been explained in the literature (76) and elsewhere as the consequence of an electron shift in the following sense: 0
111
YC/
CU~Pt+CU
However, since the data on v(M/CO) usually concern the situation where O(CO)+ 1 on M but atoms of Group IB metals either are unoccupied or bear a CO molecule vibrating with a frequency v(IB/CO) different from v(M/CO), the mutual interactions of CO vibrating dipoles are lower on alloys than on pure metals at standard experimental conditions. As a consequence, the v( M/CO) should be lower on alloys (dipole-dipole interactions cause a blue shift) than on pure Group VIII metals, when O,(CO) + 1. To decide between these two explanations, one has to perform experiments with 2CO/' 3 C 0 mixtures (77-79). Each "CO molecule functions as a free site since its dipole does not feel or cause the resonance-type interaction with the l 2 C 0 dipoles. If alloying causes an electronic structure or ligand effect, in the above-mentioned sense, the effect of alloying must be the same on pure l2CO as on mixed '2CO/'3C0 layers; the curves v vs '2CO/'3C0 composition must run parallel for a pure Group VIII metal and for an alloy. If the effect of alloying is purely a dilution effect in the dipole-dipole interactions, these two curves should converge into the same point, when '2CO/13C0 approaches the limit of zero. It has been found with Pt/Cu alloys (80) that the two curves indeed converge. No ligand effect could be detected in this way. It can be reasonably expected that the same conclusion holds also for other alloys which are formed with a smaller exothermic effect (Pd/Ag, Pd/Au, etc.) or which are formed endothermically (Pt/Au, Ni/Cu, etc.). Therefore, the authors (81) turned their attention to Pt/Pb alloys, where they found that, indeed, a part of the shift in v due to alloying might be caused by other than dilution effects. Preliminary experiments with Pt/Re and Pt/Sn alloys show that also with these alloys the contribution of other than dilution effects is very small (82). Those who prefer speculations on the ligand effects of alloying in hydrocarbon reactions may object that what is true for CO or H2 is not necessarily
'
CATALYSIS BY ALLOYS IN HYDROCARBON REACTIONS
159
true for hydrocarbon chemisorption complexes. This is in principle true, but with the data available today a statement seems to be justified that the effects of alloying should be primarily rationalized by the ensemble-size and ensemble-compositioneffects, since there are data supporting this idea, whereas for alloys studied to date, solid-state physics, heats of adsorption, and IR data offer very little support for speculation on an essential role of ligand and other electronic-structure (population of d-bands etc.) effects in hydrocarbon (and actually also other) reactions. Another objection is the following: “Is it actually reasonable that no effects due to electronic-structure changes (ligand effects) are found?” Is that not an artifact? However, there is a rational explanation for this (83). Baerends et al. calculated the binding strength of CO to a metal cluster of varying size and discovered that even second shell cluster atoms contribute considerably to the heat of adsorption. However, these atoms contribute to it mainly by their s-electrons. Now, when, for example, Pt is placed into a matrix of Cu (or Pd in Ag, etc.) it does not matter much whether the s-electrons of the next neighbors are supplied by Pt or by Cu atoms; these s-electrons are delocalized anyway. This leads, in our opinion, to the absence of any pronounced ligand effects in alloys of moderate exothermicity or those formed endothermically.
111.
Particle Size Effects
STRUCTURE OF SMALL PARTICLES : A. ELECTRONIC EXPERIMENTAL ASPECTS AND THEORY Most of the industrial metallic catalysts are metals on carrier. The main purpose of using a carrier is, of course, to achieve high dispersion of the metal component and to stabilize this form of metal against a spontaneous sintering. However, in important reactions (like reforming of hydrocarbons) a metal support is not inert and the overall reaction is actually an interplay of the two functions: that of the metal and that of the catalytically active “carrier.” Moreover, some other effects may also play a role: (1) Highly dispersed metals expose atoms to the gas phase in unusual geometric arrangements, which leads to particular low coordination (unsaturated) etc. (84-93). (2) Very small particles expose atoms of a lower coordination and therefore with a population of d-electrons different from that of atoms on flat
160
VLADIMIR PONEC
surfaces. Also the local density of states at the Fermi level is different on these atoms (94-97). (3) Small metal particles are usually anchored into the supporting surface by their own ions, which are built into the support surface; the smaller the particles are, the bigger is the expected effect of this anchoring (98). (4) Small metal particles reveal a not fully developed valence band (they have a system of discrete levels rather than a quasi-continuaus metallic-like band), which effect influences the binding energy as determined by XPS and might be, in principle, important also for chemisorption and catalysis (99, 100). ( 5 ) Small metal particles have a higher ionization potential and electron affinity, and both converge only slowly to the value of the work function (101, 102).
(6) Small metal particles have a compressed crystallographic structure. This effect is very well documented in the literature (103-107). (7) Small metal particles are frequently expected (however, the evidence is sometimes questionable) to experience an electron transfer with the carrier, which modifies the adsorption and catalytic properties of the metal particles [sometimes called the “Schwab” effect (108-116)]. In other cases, by special conditions under preparations of the catalysts, a so-called strong metal support interaction effect (SMSI) (117-12]) was evoked. In particular, with zeolites as carriers, there are pieces of experimental evidence reported (115, 116) in support of the existence of such transfer (for remarks on those conclusions, see 122, 123). The arguments of the pure theoretical predictions (94-97, 99-102) are very convincing. However, what are the relevant experimentally observed phenomena and their explanation? Important in this respect are the data obtained by XPS. Small particles (thin layers, or other atoms of lower coordination) reveal (a) a narrowing of the valence (d-) band; (b) a shift of the binding energy (BE) to higher values [Ekine, (measured) = hv - BE]; (c) disappearance of the effects like spin-orbital splitting of the band, etc. This is schematically shown in Fig. 2, which summarizes the data of various papers (123-135). These are all changes which the theory of metal bonding would predict for the transition of metal- small cluster+ isolated atoms (94-97, 99-102, 136). However, before we ascribe the higher BE of small particles to bonding effects, we have to consider the following. The observed BE value is influenced by relaxation and screening phenomena which effectively decrease the BE when going from a free atom to condensed matter (e.g., a metal). When these effects cannot operate on a full scale because the valence conduction band is not fully developed, the ob-
CATALYSIS BY ALLOYS IN HYDROCARBON REACTIONS
161
3
Group Vm- metal in an alloy
Group VIII- metal small particles
BE
€F
IE=Ol
FIG.2. Photoemission from a transition metal (e.g., Pd) in the state of a bulk crystal (top) or as a small cluster on a support (bottom). Comparison with photoemission of the same metal when dissolved in a matrix of a Group IB metal (middle).
served BE should be higher, as is indeed found experimentally. Whereas the bonding effects are expected to be influenced by the support used and should be different for different orbitals being ionized (i.e., for different initial states), the final state effects (relaxation and screening) should not show this sensitivity. The fact that the shifts in BE due to the particle size variations are mainly features independent of the support and are the same for various energy levels (see, e.g., 123-135) indicates strongly that the final state effects are most likely just those which are decisive for the observed BE shifts. A small cluster or an atom of, for example, Pd (see Fig. 2) in an alloy has the same position of the d-band centroid as the full developed band of pure Pd; this in contrast to the behavior of the small Pd particles. Such a difference can be rationalized if one assumes that due to the absence of a sufficient number of s-electrons in a small particle and due to their low mobility the screening is imperfect and thus the BEs are higher, whereas a hole on a Pd atom in a matrix of, say, silver can always be screened (extra-atomically) by s-electrons of the matrix. Photoemission experiments with flat surfaces revealed that atoms of lower coordination may have a different population of d-orbitals and a different local density of states (138-140). These effects have been also predicted and analyzed theoretically (94-97, 136, 137), and should be always considered. The only question is whether they manifest themselves in the chemisorption and catalytic behavior. In any case, the impression is that by making metal particles small in size, one can cause the electronic structure of a certain fraction of the metal atoms to vary more than by making a bulk “solution” alloy.
162
VLADIMIR PONEC
B. EFFECTS OF PARTICLE SIZEON CHEMISORPTION BEHAVIOR Important information on this problem has been obtained by Grunze (141). It appears that after CO chemisorption on Pd, the d-band photoemission (UV) is attenuated (differential spectra show a “negative” band) and two new bands appear due to the chemisorbed CO [(5a %)-band and (4a)-band]. A decreasing particle size causes an increase in the apparent BE of all three bands-the shift is almost the same for all three bands. This again indicates that the final photoemission effects could be responsible for the shifts observed. One may argue that UPS and XPS are not sensitive enough to detect subtle changes in bonding with varying particle size and that, e.g., IR spectra of adsorbed molecules might be a better tool. Therefore, the authors (142,143) compared the IR data for big particles, monocrystals, and sintered films with those obtained for very small particles (Pt, Ir, Cu). It appeared that the small particles behave differently, and since possible side effects due to the CO-CO interactions were excluded, the authors concluded that the reason for the different behavior must be sought in one of the effects listed on p. 159 [most likely (2) since the authors exclude explanation by effects (3) and (5)] or in the possibility that CO can approach small particles slightly closer than flat surfaces, which would enhance back donation effects on Pt or Ir, or direct donation effects on Cu. Although this effect is already more pronounced than anything ever observed by IR as a ligand effect of alloying it is still questionable whether this is important indeed for catalysis and whether other effects do not finally overshadow the subtle effects of the electronic structure variations with the particle size. We shall turn to this point later, in Section 1II.D.
+
IV.
Mechanism of Hydrocarbon-Hydrogen Reactions
A. INTRODUCTION One of the most important technological advances of the postwar period was when platinum/alumina (i.e., a “metallic”) catalyst was introduced in oil refining (144, 145). From the point of view of the mechanism of the reforming reactions, it was suggested in the very early stages of research on Pt reforming that under the industrial conditions (Pt-O.1- 1 .O wt.%/Al,O, (pure or modified); T = 470-530°C; pressure 10-30 atm) the catalyst is actually bifunctional (146-150) : Pt is mainly responsible for various dehydrogenation reactions, whereas the carrier (modified eventually by C1-
CATALYSIS BY ALLOYS IN HYDROCARBON REACTIONS
163
or F- ions) performs what is known as “catalysis by solid acids” (150, 151). This idea was confirmed in the meantime by many pieces of evidence and today the only modification of it is that the Pt or a Pt-alloy component is supposed to catalyze also some reactions other than dehydrogenation-the latter being nevertheless still considered to be the most important function of the metallic component. Authors in different countries (see, e.g., 152-154) studied Pt catalysts which were support-free (films) or which were prepared with such carriers that the acidic function of them could be neglected (active carbon, inert SiO,). They found that these Pt catalysts catalyzed almost all reactions which are known to occur with Pt/Al,O, catalysts under reforming. On the other hand, it became also clear that under the industrial conditions the surface of Pt is covered by sulfur, carbonaceous residues, coke, etc. to such an extent that most of the reactions are severely slowed, and only the simplest reaction of dehydrogenation (leading possibly to aromatization) can still proceed. These facts lead to the conclusion formulated above. Reactions of hydrocarbons on Pt, and to a lesser extent on other Group VIII metals as well, have already been the subject of three excellent reviews in this series (155-157), each review reflecting the views of the particular author(s). It is not this author’s intention to repeat information which is available elsewhere (155-157), but rather to focus on particular points; namely, those which help us to rationalize the data obtained with alloys, or vice versa, those which have been established by studies with alloys. Of course, the selection of data presented below, or the evaluation of discussions which have already taken place in the literature, is again unavoidably influenced by the author’s personal views. It is practical to discuss certain groups of reactions separately. Conveniently, the subdivision of reactions as presented in Table I may be used. A notation “3C-, 5C-” has been used in Table I to indicate the number of carbon atoms which form the essential part of the transition state complexes of the reactions mentioned. A more detailed definition of other terms will be given below. The three types of C-C bond fission concern the following reactions : (1) Terminal fission:
cccccc +ccccc + c +cccc + 2c +ccc + 3c (2) Internal fission :
cccccc ccc + ccc cccccc cc + cccc +
+
(3) Multiple fission : CCCCCC
+
6C
164
VLADIMIR PONEC
TABLE 1 Reactions of HCIH, Mixtures A. Hydrogenation, dehydrogenation, double bond isomerization, HC/D,“ exchange (all “nonskeletal” reactions)
9. Skeletal reactions
1. Nondestructive rearrangements
a. lsomerization via the 3C-intermediatesb b. Isomerization via the SC-intermediatesb c. Dehydrocyclization into a 5-ring d. Dehydrocyclization into a 6-ring e. Aromatization via a ring enlargement
2. Destructive reactions (hydrogenolysis or “hydrocracking”) a. Terminal fission b. Internal fission c. Multiple fission
HC denotes hydrocarbon. See the text for explanation of these terms
Information on the intermediates operating in the reactions of HC/H2 mixtures has been mainly obtained by the three following ways:
( I ) Exchange reactions in HC/D2 mixtures. In particular, the bonding metal-hydrocarbon fragment is conveniently studied by these reactions. The basis of these studies has been established by Kemball (see, e.g., 158, 159), Burwell (e.g., 160), Tamaru (161), Bond (162), and others. ( 2 ) Reactions of molecules with labeled carbon (13C, 14C). Gault and his co-workers pioneered (see, for review, 157) the use of 13C-labeled molecules in experiments by which the operation of adsorbed complexes with either three or five C atoms involved could be tested (see below). I4C has been repeatedly used in problems concerning the dehydrocyclization and aromatization reactions (157,163-166). The authors (163-166) in their research combined the use of labeled molecules with the third method, the study of “archetype” molecules. (3) The study of “archetype” molecules. This method has been proposed and widely used by Rooney, Burwell, Anderson, and others (see, for review, 155, 156, 160). In this method a molecule is used which can form an archetype of chemisorbed complex (“caged” molecules as derivatives of adamantane or ethane in its hydrogenolysis, neopentane in exchange with D2 or in reforming reactions, etc.) or which can form several complexes, but the contribution of these complexes to the overall mechanism is easily derived from the product spectrum [as is the case, for example, with neohexane (167, 168)].
CATALYSIS BY ALLOYS IN HYDROCARBON REACTIONS
165
It is a serious but frequently neglected problem that the analysis of the data obtained with the method (2) or (3) above is only straightforward when each molecule undergoes only a one-step reaction upon one adsorption sojourn on the catalyst surface. If several consecutive reactions (e.g., isomerization combined with hydrogenolysis or two isomerization steps in combination) follow each other before the molecules leave the surface, useful information is still gained (167, 168), but the discussion of data is more complicated. Metals like Pt or Pd do not seem to be a problem in this respect, as is the case with other metals at the lowest possible reaction temperatures. However, metals like Ir or Rh are apparently very active in performing several consecutive steps during one residence of the molecules on the surface, and at temperatures above 200°C it is difficult to avoid the multiple reactions (167).
B. CHEMISORPTION COMPLEXES OF HYDROCARBONS ON METALS Let us discuss first the binding of hydrocarbon molecules to the surface. 1. Metal- Carbon Single Bond
It seems beyond debate that when an exchange reaction of a hydrocarbon (HC) with D2 is observed and the initial product distributions are binomial (random distribution of D atoms), single a-metal-carbon bonds are being formed. Nevertheless, this conclusion was puzzling in the period when virtually no homogeneous alkyl-metal complexes were known and the stability of alkyl-metal complexes was doubted for “principal” reasons (see, e.g., 169). However, it appeared that these complexes can be rather stable when one blocks a very fast and easy elimination of one of the H atoms in the /?-position, which step decomposes the alkyl-metal bond into an olefin and a bound hydrogen atom (170,171). On the other hand, this means that the transition H
(where M denotes metal) must be considered in the schemes of catalytic reactions as a very easy running step, wherever the concentration of bound H and olefin is sufficiently high and this reaction is not blocked by other ligands on the same M atom.
2. Metal-Carbon Multiple Bonds The existence of such bonds was inferred first from the initial product distributions of the CH4/D, reactions. Kemball(158,159,172) reported that
166
VLADIMIR PONEC
with metals like Rh or Ni, even at lowest conversions, when no repeated desorption/readsorption process could be expected, distributions were observed with a very high content of the d4- and dJ-products. Kemball (172) suggested the following mechanism to explain it :
This mechanism employed the postulated multiple bonds. It might be that with some metals and at higher temperatures the dehydrogenation is deeper and the multiplicity of bonds is even higher (e.g., that HC-M is also formed). In spite of this uncertainty, the “multiple” exchange of CH4 became a very good diagnostic tool for the multiple metal-carbon bonds. It is an important question whether with higher hydrocarbons the ULY or uuu multiple bonds can exist as well. The fact that multiple exchange (i.e., more than one D atom enter the molecule during its one sojourn on the surface) of ethane and its homologs takes place at much lower temperatures than that of methane (158) indicates strongly that, wherever possible, the aP complexes (or its alternative, n-complexed olefins) are indeed formed. Moreover, they are formed more easily than the cia and related complexes. Nevertheless, the initial distributions obtained with mononuclear homogeneous complex (173-176) show very clearly that in this situation (an isolated center), the asymmetry of the C2H6/D2 exchange is high (clear maxima at d 3 and d6); in other words, carbene-like structures can be formed also with higher hydrocarbons than methane. Most likely, the correct conclusion is that the two types of multiple reactions-i.e., via the uu and the a/? complexes-run in parallel (177). The question arises whether also other double-site multiple bonds, like ay or ad, are equally possible. The answer seems to be a negative one (see, for review, 157-162): When the hydrocarbon chain is interrupted by a heteroatom (ethers etc.) or by a quarternary carbon (neopentane, neohexane, etc.), the exchange proceeds separately on both sides of this obstacle and does not go easily over to other parts of the molecules, so that one can conclude that the formation of cry (and analogous) complexes is always more difficult than the formation of the aP complexes. Even the formation of the ap complexes seems to be subject to certain limitations. Investigations with sterically well-defined and distinguishable hydrogen atoms (like, e.g., hydrogens of adamantanes) showed that only the mutually “eclipsed” hydrogen atoms can undergo the UP complex formation (160, 178, 179). Let us mention here another relevant fact : Formation of multiply bound complexes of the uu or UP type is not substantially altered (see below) when an active metal is diluted by alloying in a matrix of an inactive metal (e.g.,
CATALYSIS BY ALLOYS IN HYDROCARBON REACTIONS
167
Ni in Cu). This indicates that the formation of the “UP,’complexes is possibly a “one”-site process. This would point strongly in the direction of the alternative (exactly the same results should be obtained from reactions running through these two types of complexes) for the ap two-site complex, namely, a n-complexed, one-site-bound intermediate. M
/
CHz-CH \z
++
M
CHz+CHz
M
M
(This complex also involves the two a/? carbons in the metal-HC bonding.) However, it is also not possible to suggest the n complexes as the only form of the aP complexes. Burwell and Shrage (179) studied the exchange reactions of bicyclo-[3.3.1]-nonane, with which molecule one can reasonably expect some suppressing of the n-complex formation. Nevertheless, the authors found that the multiple exchange proceeded easily and was spread over the whole molecule, so that there was most likely a mechanism other than n-complex formation which allowed it. The question of the exact structure of complexes which are bound to the metal surface through two carbons is thus still open, and it is not impossible that both alternatives exist side by side. 3. Two Carbon Atom Complexes Formed upon OfeJn Chemisorption Reforming reactions comprise dehydro-/hydrogenations and olefins might also be intermediates of other reactions-such as the above-mentioned exchange reactions. The two forms of associatively adsorbed olefins have been already mentioned : n complexes and ap two-o-bonded complexes. The questions posed are as follows: (a) Do dissociative forms of olefins also exist; (b) are any of these forms reactive enough to be an intermediate of hydrogenation/dehydrogenation reactions? The answer to the first question is undoubtedly a positive one. The classical papers by Beeck et a f . , Rideal er al., and others have shown that ethylene disproportionates upon chemisorption into ethane and carbonaceous (adsorbed) residues (see 162). This disproportionation takes place at relatively low temperatures: at room temperature and lower (see 162 for review). Moreover, the intensity analysis of LEED data has shown that upon chemisorption of ethylene, ethylidyne structures are formed. Similar structures are also formed by dissociative adsorption of higher olefins (181,182). There is thus no doubt with regard the first question. The second question is, however, still being discussed (see, e.g., 180)Gault et al. studied the butenes/D, exchange and came to conclusions which were supported by the data of mass and microwave spectrometries (183-185)namely, the dissociative adsorption produces on some metals intermediates
168
VLADIMIR PONEC
which are comparable in activity with those of “associative” adsorption : a x-complexed olefin and an crP two-a-bonded olefin. It is a known fact that in many heterogeneous reactions propylene is more reactive than ethylene. The allylic hydrogen is labile and the dissociative (in C-H) adsorption in the allylic position is promoted by that. Possibly, the higher olefins can thus be adsorbed by two dissociative adsorptions: through the vinylic or the allylic position (186). The ease of dissociative adsorption of multiple exchange and of P-H elimination suggest that the transition
7
C
4
c
H C
M
M
I
II
should also be easy. This is a point to consider in suggestions on the mechanisms. 4.
Complexes Involving Three Carbon Atoms (3C, cry Complexes)
Exchange reactions of neopentane have already lead to the conclusion (155, 158, 162) that 3C complexes, bound to the surface by the ay carbon
atoms, may be formed on some metals (e.g., Rh or Pt). However, it was evident from those experiments that 3Ccry complexes are formed by metals much more reluctantly than the 2CaP or the crcr-bound complexes. It means that their formation can only be studied at (much) higher temperatures than those suited for the study of the HC/D2 exchange reactions. In this case one can advantageously use the skeletal reactions of neopentane themselves as evidence for the formation of 3C complexes. When neopentane is being isomerized or split into C1 and C3 fragments, 3C complexes are certainly
C
C
I c-c- c-c I ap
Ir.Ni Rh
I c-c-c I
+
c
C
C
aP FIG.3. 2C complexes from hydrogenolysis. Illustration of the experimental evidence available on their existence, the known (!) and unknown aspects (?) of 2C complex formation are also indicated.
CATALYSIS BY ALLOYS IN HYDROCARBON REACTIONS
169
being formed-there is even no other alternative (155). It appeared that under the temperatures of reforming-type skeletal reactions, neopentane reacts in parallel toward various products (155-157, 180, 187). Unfortunately, there is no way at the moment to establish how the 3C complexes are bound to the surface : as aay, aayy, or cry metallocyclobutane-like binding? We have already discussed the arguments which lead to the conclusion that under the conditions of the HC/D, exchange reactions, the 2CaP complexes are being formed. It is probable that also at higher temperatures this is the complex most easily formed. Nevertheless, it is important that we also have support for this statement in the form of other results. When a molecule like ethane undergoes hydrogenolysis into methane, at certain stages both carbons are bound to the surface, i.e., 2C complexes are formed. An alternative would be a radical-like fission with activation energy of 80-90 kcal/mol, i.e., almost two times higher than is found experimentally. Another molecule which demonstrates the same point is neohexane (167, 168 and references therein). When neopentane and methane are formed with this molecule, one can trust that 2CaP complexes had been formed. Again, it is not certain at which stage of dehydrogenation the splitting of a C-C bond takes place-sag, - act-@, or aaa-/?/?p?Therefore, the designation “2Ca/? complexes” must be understood (unless otherwise specified below) in a broad sense, i.e., as all complexes where two carbons are involved, irrespective of the number of bonds with the surface. This holds true not only for this review but also for most of the cited literature. The most relevant facts regarding the 2C and 3C complexes are once more summarized in Figs. 3 and 4. The still unsolved question: “How many bonds are formed upon chemisorption between the molecule and the surface?” is closely related to the problems of the detailed mechanisms. Various mechanisms have been suggested and for most of them good arguments have been made, but nevertheless, because of the above-mentioned uncertainty, they remain speculative. However, even speculative mechanisms may sometimes be helpful and therefore we shall turn to the problem of mechanisms in a separate section.
FIG.4. 3C complexes, the existence of which can be seen in experimental evidence from exchange [neopentane Rh (Pt)] and from hydrogenolysis and isomerization (neopentane, neohexane). As in Fig. 3, the known (!)and the speculative aspects (?) of the 3C complex formation are indicated.
I70 5.
VLADlMlR PONEC
Complexes Involving Five Carbon Atoms
Gault et al. noticed in their early papers (157)that the product pattern of methylcyclopentane (MCP) hydrogenolysis is sometimes surprisingly similar to that of hexane or methylpentane(s) isomerizations. They suggested that isomerization proceeded via a cyclic, methylcyclopentane-like intermediate. Later it appeared that the similarity was not always found, but an important idea was already born and, more importantly, was brilliantly confirmed by later papers from the laboratory of Gaults. The idea of the evidence is rather simple and can be elucidated by means of the following experiment. Let us consider, for example, a molecule of 2-methylpentane labeled in a branched position by 13C 2-methyl-I3C(2)pentane. If the consecutive reactions in the adsorbed state are with a given metal of low extent, and this is certainly true for Pt or Pd, then the appearance, among the product, of 3-methyl-13C(3)-pentane is very strong evidence of the operation of the 5C (cyclic) intermediates. Only via a ring closure at one place and an opening at another place of the molecule can a label move simultaneously with the branch. On the other hand, when the branch and labeled atom become separated by isomerization, this is evidence of the operation of the 3Cay complexes (see Fig. 5). Because of historic reasons, the mechanism employing the 3Ccry complexes is often called a “bond shift mechanism” and the mechanism with 5C complexes-“a cyclic mechanism.” However, both mechanisms involve cyclic intermediates at certain stages and for both mechanisms bonds are shifted. Therefore, notation specifying the number of carbon atoms involved seems to be preferable. As with 3C complexes, it is not clear how many bonds are formed with the metal and how many metal atoms are involved, when 5C complexes are formed. Some suggestions will be discussed below. The “state of the art” is summarized in Fig. 5 . The use of a labeled molecule is the only exact way to determine quantitatively the contribution of the 5C complexes to the overall isomerization. However, a rough estimate of it can, in favorable cases, also be made by comparing the isomerization of pentane (3C only) and hexane (both 3C and 5C complexes are possible) on the same catalyst and under the same conditions. For Pt and Pt/Cu alloys both methods have led to the same conclusions (188). It should be briefly mentioned that not only 2-methylpentane (2MP), but also other molecules can be used to establish the proportions of the 5C/3C mechanisms. With Pt, various molecules have lead to a similar result (157). However, with other metals the discrepancies are quite substantial (189). This can be rationalized either by assuming (189) that the 4Cu6 complexes
CATALYSIS BY ALLOYS IN HYDROCARBON REACTIONS
F‘*
-
c-c-c-c-c 2MP
171
cI *
c-c-c-c-c 3MP
3MP
/c\
c-c
*
c
c-c
/c\ ..
*
*
c
/\
*
*
C
FIG.5. 5C complexes. The most important piece of evidence for their existence is indicated in the reaction scheme. The known ( !) and speculative (?) aspects of the complex formation are indicated.
are also formed, or by admitting that more than one rearrangement of a molecule is possible during each sojourn on the surface of some very active metals. An inspection of the data on exchange reactions and on the reactions with “template” molecules (archetypes of certain intermediates) shows that the formation of various complexes takes place easily and the increasing difficulty in the formation can be indicated as follows: a
--ca/l --caa --c3Cay
-+
5Cac
From this comparison, the existence of the 4Ca6 complexes does not seem probable at low temperatures. However, it is not completely excluded either. It has been observed with Pt that molecules of 2,2,3,3-tetramethylbutane undergo hydrogenolysis mainly into two molecules of isobutane. This could be evidence of the formation of 4CaS complexes (190) on Pt. 6 . How Do Metals Difler in the Formation of DifSerent Complexes?
Investigations of the H C / D , exchange reactions have led to the following conclusions (158, 159, 162, 177, 191): Ni, Ru, Rh, and C o are the best catalysts for the formation of the aa complexes; Pt, Pd, and Ir are worse in this respect. On the other hand, the 2Caj complexes of exchange reactions are most easily formed on Pt, Pd, and Ir. Pt seems to be also the best metal to show ay binding (155, 156, 192). Results on reactions of neohexane and neopentane confirmed that Pt and, to a less extent, Pd are able to form 3Cay-type complexes rather than the
172
VLADlMlR PONEC
2CaP complexes of hydrogenolysis (this particular point will be discussed later). On the other hand, all other Group VIII metals showed a preferential formation of products which can be related to the 2CaP complexes. Only at higher temperatures or when a surface was (se1f)poisoned by carbonaceous, firmly adsorbed species did the 3Cay complexes show up (167, 168, 193-1 95).
By studying different metals and, with the same metal, catalysts with different particle sizes, various authors have shown that one has to assume at least two different mechanisms involving different 3Cay complexes (195198). Also, the work on alloys (see below) leads to such a conclusion. It is obvious that metals would differ in contributions by the respective mechanisms, but at the moment a generalization in this respect is not yet possible. Low temperatures and the simultaneous availability of a large number of contiguous sites (big ensembles)-the conditions usually met with “massive” carrier-free metals and with those metals which allow a fast removal of carbonaceous deposits from the surface-are the conditions which favor the 3Cay mechanisms (of isomerization, and possibly of hydrogenolysis as well). Relatively higher temperatures, smaller particle size, and alloying with an inactive metal all seem to promote the formation of 5C complexes. As an example, small Pt particles show almost exclusively 5C complex formation, whereas massive Pt, Ni, or Ni/Cu alloys show a prevailing 3Cay complex formation (157, 197-199). A definite theoretical explanation of this behavior is not available. It is important to realize that the preference of a metal for 3C as opposed to 2C complexes or for 5C as opposed to 3C complexes may be either intrinsic or induced by adsorption of less reactive carbonaceous fragments and carbon (for simplicity, we shall refer to both of these as “carbon”) on the metal (alloy) surface. Also, the choice of the reaction conditions (apparent contact time, poisoning or self-poisoning of the catalyst, etc.) influences the temperature range in which the catalysts can be tested, and since the selectivity in various complex formations is also temperature dependent, one must always analyze which aspects of the product distributions are intrinsic properties of a metal and which are induced by often unavoidable side reactions. C . MECHANISM OF SKELETAL REACTIONS AND
SELECTIVITY OF METALS
1 . Isomerizarion and 5-Ring Dehydrocyclization It can be now considered as well established that isomerization involves formation of (various) 3C and 5C complexes. However, all other details of
CATALYSIS BY ALLOYS IN HYDROCARBON REACTIONS
173
mechanisms discussed below (or in other papers) should be considered as no more than useful speculations. Results obtained with alloys (alloying causes variations in the distribution of the ensembles according to size) and with metals of varying particle size (due to the geometry of the curved surfaces and due to the deposition of “carbon,” the same effects are expected to operate here as in alloys) have lead to the conclusion that the various 3C and 5C complexes might differ in the size of ensembles which are required for the formation or the steadystate binding of the complexes. The “small” and “big” ensembles are, in the following, schematically represented by “one-site’’ and “two-site’’ ensembles. Figure 6 presents two suggestions from the literature for the possible pathways in the conversion of the 3Cay two-site complexes. Figures 7 and 8 present the suggestions for the 3Cay one-site complexes (155-157, 198). According to the mechanism in Fig. 6, isomerization is induced by dissociation of at least three C-H bonds; according to the other mechanisms, two or even one CH bond dissociation would be enough. Mechanisms like those in Fig. 7 were suggested (in various alternatives) when it became known that a 3C isomerization in a mixture with D, (instead of H,) produces only
FIG. 6. 3C complexes as intermediates of isomerization. A “two-site” (large ensemble) mechanism via a bond shift (left) or cyclopropane ring (right), as suggested by various authors (see text). Except for the number of C atoms involved, all other aspects of the mechanisms are speculative. The same remark holds for Figs. 7-9.
174
VLADIMIR PONEC
FIG.7. 3C complexes of isomerization-a the literature.
metallocarbenium intermediate as suggested in
d,-products. For the sake of simplicity, all reactions are shown with neopentane, but this does not mean that the suggestions are limited to this molecule. It has been already mentioned in passing that indications exist in the literature showing that the 3C isomerization can take place by formation of at least two different 3C complexes, having different activation energies of isomerization, different particle size effects, different responses to alloying, etc. (157, 195-198). The suggestions presented above offer a choice of different complexes for further speculations. However, a definitive description of isomerization mechanisms under different conditions (H, pressure, temperature, etc.) and with different catalysts (pure metals, alloys, etc.) is not yet possible. There is, of course, always something which supports one of the suggestions in Figs. 6-8. For example, the suggestion in Fig. 6 is supported by the fact that the results with alloys (see below) can only be rationalized when
FIG. 8. 3C complexes in isomerization. Various pathways, as suggested in the literature (see text), for the "one-site" (small ensemble) conversions of metallocyclobutane rings.
CATALYSIS BY ALLOYS IN HYDROCARBON REACTIONS
175
a “big ensemble” mechanism is assumed in addition to an “one site” mechanism. The suggestion in Fig. 8 is an analogy of the olefin metathesis (150), and the experiments with “caged” molecules support the suggestion in Fig. 7. However, it seems very probable that none of the mechanisms alone can explain all the data. Moreover, the stability of some intermediates (branched ally1 complex, cyclopropane highly strained ring) can be doubted. With one pathway of the mechanism in Fig. 8, there is also some additional trouble: With some metals or alloys the selectivity for isomerization might be very high, sometimes nearing 1007;. This implies that the reconstitution of the original alkane molecule should be 100% even under conditions (those of alkane skeletal reactions running) when the fragments would be thermodynamically more stable than the reconstituted molecule. Also, the free rotation of n-complexed olefin (the bond to the metal must not be too weak, otherwise olefins would desorb) raises some questions. In summary, probably all conversions in Figs. 6-8 should be considered when speculating on mechanisms, but caution is always called for. Isomerization via the 5C cyclic intermediates and the 5-ring closure can be discussed together. Figure 9 shows different one- and two-site intermediates which have been suggested in the literature for these two reactions (155-157, 198-201). Isomerization consists of a ring closure in these intermediates and a ring opening, both of which take place at different spots of a molecule. Upon dehydrocyclization, a desorption follows the ring closure. The Hungarian and Russian schools seem to prefer cyclization with five carbon atoms flatly lying on the surface (see, for review, 201). This is only possible when no more than one a-bond per carbon atom is formed toward the underlying metal atoms, since multiple bonds would probably lift the molecule away from the surface. However, the a-bonds are usually well localized (of course, one does not know for certain when metals are involved
*
*
*
*
*
*
FIG.9. Possible one- and two-site intermediates of reactions involving 5C complexes.
176
VLADIMIR PONEC
in this bonding) and unreactive, so that one would not expect a recombination of two 0- metal-C bonds into a new C-C bond of a cycloalkane. Another argument against such a recombination is the fact that C-C bond fission in alkanes and cycloalkanes is most likely preceded by C-H dissociation, which step automatically leads to the formation of one or more multiple metal-C bonds before or during C-C splitting. Thus, also with 5C complexes there remain open questions. 2. Hydroyenolysis Ethane can be hydrogenolyzed by all Group VIII metals, i.e., all these metals can form the 2C complexes. With all metals except Pt and Pd, the neohexane assay shows that the 2C hydrogenolysis is easier than the 3C splitting. However, neopentane is also hydrogenolyzed by all Group VIII metals, so that the difference in the ease of formation of the 3C and 2C complexes is not prohibitive for one or another mode of fission. Actually all 2C, 3C, and even 5C complexes can, at least in principle, be a starting point of C-C bond splittings. At the moment it is impossible to assess quantitatively the contribution of various complexes to the overall hydrogenolysis. A speculation in this respect will be presented at the end of this review. At that time we shall need the following information. At low temperatures, the splitting of hexane is mainly of the “internal” type for metals which are also good for isomerization, as is Pt or Ir (202204), but is of the terminal type for good hydrogenolytic catalysts such as Ni, Co, Rh, and Ru. Palladium stands between these two groups. When the temperature is increased, the selectivity for isomerization increases and that for hydrogenolysis decreases. The increase in temperature causes also a shift in the hydrogenolytic selectivities : from the internal type to the terminal splitting. The changes in isomerization and hydrogenolysis selectivities are almost mirror-like. This has been frequently observed (157, 198, 199), with various catalysts by F. Gault, who called the internal splitting “a frustrated isomerization.” Obviously, if such a close relation exists, the complexes leading to the internal hydrogenolytic splitting should be of the 3Ccry type. 3. Aromatization From the practical point of view, this is probably the most important reaction related to the metallic component of the reforming catalysts (despite the fact that a part of aromatization is acid catalyzed). There are certainly several pathways which can, at least in principle, lead to the aromatic products. Let us mention here the most relevant facts on aromatization of hexanes and higher hydrocarbons. Several authors (150, 151, 205-208) studied the formation of aromates
CATALYSIS BY ALLOYS IN HYDROCARBON REACTIONS
177
from pentanes which were substituted in such a way that the 5- but not the 6-rings could be closed in one step. Formation of aromates is then evidence that aromatization via a ring enlargement is in principle possible. However, Davis et a/. have shown, by also using labeled molecules (163-166), that wherever a direct 6-ring closure is possible, it is faster than the pathway via the 5-ring closure and ring enlargement. It has also been established that alloying, promotors, carriers, reaction conditions, etc. influence the aromatization and the 5-ring closure quite differently. This suggests that these two are, indeed, different pathways. The question now is: In which respect are they different? It is not excluded that the carbon 6-ring can be closed in a way analogous to that of 5-ring closure (see Figs. 5 and 9), i.e., via a metallocycloheptane or dimetallocyclooctane intermediate. However, the expected (283) lower stability of these rings does not make this idea very attractive. Another possibility is that aromatization is actually a consecutive dehydrogenation. Hexatriene, once it is formed, does not need any help from the catalyst to form cyclohexadiene (and benzene in the next step). Temperatures at which industrial aromatization (in the framework of reforming) takes place are high enough to make a sufficient concentration of olefins thermodynamically possible. This mechanism has been suggested and it is generally accepted for the first generation of reforming catalysts-namely, the oxides Cr,O,/MoO,/Al,O, etc., which operate at low pressures of hydrogen. However, the typical Pt/Al,O, catalysts (or Pt/Ir, Pt/Re, etc.) operate at 10-30 atm of total pressure, and under these conditions the presence of a sufficiently high concentration of olefins was doubted in the literature. Nevertheless, there are some data (201, 209, 211) which indicate that the mechanism of aromatization might also be on metals, as just described, a consecutive dehydrogenation. There is, moreover, one indirect indication for such a mechanism. As we shall see in Section V, from all possible reactions of reforming, the dehydrogenation/hydrogenation reactions seem to be able to proceed on the smallest ensembles of active sites; possibly they can be catalyzed even by a single atom. Therefore, these reactions are least affected by deposited “carbon,” by sulfur (always present in traces under industrial conditions), by alloying of Pt with inactive elements (Sn, Au), etc. Also, aromatization by the metallic component of the reforming reactions shows these features and this strengthens the belief that the metal-catalyzed aromatization, apparently always present under reforming conditions (212), is indeed a consecutive dehydrogenation and cyclization of trienes. Evidently, the “carbon” and sulfur deposits decrease the surface concentration of hydrogen, which effect leads to an increased dehydrocyclization and aromatization (see, e.g., 204). This might make the dehydrogenative pathway of aromatization feasible also at rather high total pressures.
178
VLADIMIR PONEC
On the other hand, hydrogen may have an accelerating effect (213) (positive order in P H I )because it keeps-by a continuous removal of carbon and sulfur-a small part (small ensembles) of the metal surface working. 4. Activity and Selectivity of Metals
The simplest situation is with ethane, which can be only hydrogenolyzed. A rather complete collection of data exists on this reaction due to Sinfelt and co-workers (214,215) and available reviews on this subject make profitable reading. The distinct features of ethane hydrogenolysis are the highly negative order in hydrogen pressure and the very low activity of Pt and Pd. The active metals are 0 s > Ru > Ni > Rh, Ir. Of these metals, Ru, Ni, and Rh are already known to form multiple carbon-metal bonds rather readily and this might be one of the factors favorably influencing their hydrogenolytic activity (159, 191). Cyclopropane hydrogenolysis to propane is a reaction which reminds one of the hydrogenation of olefins (216). This is due to the specific electronic structure of this molecule (217, 218). This hydrogenolysis is at low temperatures and is accompanied by hydrocracking into ethane and methane. At higher temperatures a multiple hydrocracking into three methane molecules may also take place. It is interesting to note that the propensity of metals to break the C-C bond is apparently closely related to the degree
60-
50-
LO302010-
0'
I
200
I
300
I
LOO
I
T(OC)
500
FIG. 10. Selectivities in hexane conversions versus temperature for benzene formation (Be), hydrogenolysis (Hy), methylcyclopentane formation (MCP), isomerization (ISOM), and dehydrocyclization (Dehy) (9 wt. % Pt on inert SO,).
CATALYSIS BY ALLOYS IN HYDROCARBON REACTIONS
179
to which the hydrocarbon molecule is dehydrogenated upon its chemisorption (216). The lower the ratio H/C of adsorbed species is, the higher the selectivity for hydrogenolysis. Hydrocarbons higher than C , can also undergo isomerization on Pt, running in parallel to hydrogenolysis. With butane or isobutane the selectivity for isomerization is rather low, also on Pt, but the higher hydrocarbons show more of isomerization reactions. With higher hydrocarbons some other metals (Ir, Pd) also show some isomerization selectivity. The following point should be noted. Guczi et al. (219) reviewed the data on the kinetics of hydrocarbon skeletal reactions and summarized the results as follows: (a) Log rate as a function of log pH2shows a maximum; at low hydrogen pressures the slope is positive, at higher pressures it becomes negative. (b) The maximum of the above-mentioned function shifts to higher pH2when the number of C atoms increases; skeletal reactions of heptane show already a positive order in pH,near atmospheric pressure. (c) This behavior is most likely related to the deposition of “carbon” on the metal surface. When this process is more extensive (molecules like heptane and bigger ones), the selectivity for isomerization is higher and that of hydrogenolysis lower than with smaller molecules. The same parallelism is found when different metals are compared with the same hydrocarbon molecule. Starting from C , molecules, dehydrocyclization (into cyclopentane and derivatives of cyclopentane) is also possible. From C6 on up, aromatization also occurs. These two reactions comprising a dehydrogenation step are only observable at temperatures which on most metals are higher than the region where hydrogenolysis (hydrocracking) is first observed.
3oi IS0
10-
MCP I
200
300
FIG. 11. Selectivities in hexane conversions versus temperature for hydrogenolysis (Hy), isomerization (ISO), benzene (Be), and methylcyclopentane (MCP) formations (1% I r on inert SiO,).
180
VLADIMLR PONEC
Thus, these reactions are only observable with metals which show a low intrinsic activity for hydrogenolysis (possibly Pt, Pd). Otherwise, the hydrogenolytic activity of a metal had to be suppressed by “carbon” deposits or by alloying with inactive metals (see Section V), by promoters and other modifiers (sulfur), etc. The last effects (e.g., “carbon” deposition) are also important with metals like Pt or Pd. The overall behavior of hexane/H, mixtures in contact with Pt or Ir on inert carriers is shown in Figs. 10 and 11. The data shown are derived from Refs. 202-204. These data have been collected by using different feed rates, and this is the reason why the Pt data do not form one smooth curve. A longer contact time leads, even at the lowest conversions, to a decrease in the C, dehydrocyclization and an increase in isomerization. This demonstrates the close relation (via the 5C intermediates) between these two reactions. Tables 11-IV demonstrate the behavior of various hydrocarbon molecules on different metals. The order of metals in their activities is also known for some reactions other than ethane hydrogenolysis. Maurel and Leclercq (220) found the following order for cyclopentane hydrogenolysis: Ru > Rh, Ir > 0 s > Ni > Pt > Pd > Cu, Fe. Various Co catalysts showed activity between that of 0 s and Pt (most likely the influence of an uncomplete reduction). Carter et a/. (221) found the following order for heptane hydrogenolysis : Ru > Rh, Ir >> Pt > Pd. The common features of these orders in activities are evident. The known data allow also a comparison of the selectivities in nonthe order usually found (decreasing Sndr) is destructive reactions (Sndr); Pt > Pd > Ir >> Co, Ru, Rh, 0 s . A comparison made by Davis and TABLE I I Product Distribution in n-Pentunc, Riwctions w i ~ hH , “
Molar percentage
Catalyst PtiSiO, (16 wt. ‘I,) Ni/SiO, (9 wt. y o )
(
pH2 Ppcnt C) (atm) (atm)
312 346 350 350 350
0.9 0.9
0.1 0.1
2.5 5.0 5.0
0.5 0.5 2.0
iso-
C, 5 6
C,
C,
17 15 20 18 85.9 6.6 4.7 77.0 8.5 8.3 52.0 0.5 3.0
C,
C,
c-C,
3 4 8.1 5.6 4.4
52 43 0 0 0
6 7 0 0 0
S,,” SC,d (‘lo)
(‘I fast isomerization and dehydrocyclization > slow isomerization and dehydrocyclization 2 hydrogenation/dehydrogenation, including a part of aromatization. While the two statements above are almost phenomenological, one inevitably enters the field of speculation upon attempting to find an answer for the following questions: Where does the reaction or binding to the surface take place-on the summits of the surface atoms or in the valley positions? Do metals and various molecules differ in this respect? Metals which with adsorbed CO prefer to form metal-carbon bonds on the summits are Pt and Ir (Cu?); metals which promote binding in the valley are Pd > Ni > Rh, Re. Metals promoting multiple metal-carbon bonds (with hydrocarbons) are Ni, Ru, Rh; Pt and Pd are much worse in this respect. Let us extrapolate and assume that what holds for C O also holds for hydrocarbon molecules, and that the characterization of the multiplebond formation propensity is valid also at higher temperatures than were established experimentally by exchange reactions. Then we can attempt to rationalize the available information on the formation and the role of various hydrocarbon complexes. We know already that metals differ in their ability to form 2C and 3C complexes. In general, under the conditions when solely the exchange reactions are running, formation of 2CaB complexes is always easier than formation of 3Cay complexes. However, the results obtained with alloys (active + inactive metals) show that we have to discern two different types of 2C complexes :
(i) One type of 2Ccomplex is not strongly affected by alloying (with alloys
CATALYSIS BY ALLOYS IN HYDROCARBON REACTIONS
203
like Ni/Cu, Pt/Au, Pd/Ag, etc.) (245) and is an intermediate of exchange reactions (ii) Another type of 2C complex, the formation of which is strongly suppressed by alloying, is the intermediate of hydrogenolysis Several suggestions have been made in the literature regarding these two groups : Group (i) R
R\
I
/R
CHI I
H/ C ’ f C L H
CHI I
Ni-Cu-NI-Cu-Cu-
N!
CH3
,R
c-C\
I
CH-R
/R
H I’ -NI-CU-I
I
-N~-~-NI-
H
Group (iil
H2C - CH2
/
NI
\
NI
HC - C H
/\
A
NI NI NI NI
C-
/I\
NI NI NI
C
/I\
NI NI NI
Most likely, the two groups differ in the sense suggested by Tetenyi (222), namely, in their degree of dehydrogenation and in the number of bonds formed between the individual carbon atoms and the metal surface atoms. It seems to be a reasonable assumption that not only the 2C complexes of group (i) but also those of group (ii) are formed more easily than the 3C complexes of exchange and skeletal reactions. (Note that the 3C complexes of skeletal reactions also require a dehydrogenation of the adsorbing molecule.) However, on some metals (see, e.g., the results on the mechanism of reactions on Pt and Pd) the products of 2C hydrogenolytic complexes do not appear in the gas phase before (i.e., at lower temperature) the 3C complexes. Thus, the assumption, when accepted, implies that when the 2C hydrogenolytic complexes [group (ii)] are formed more easily on all metals than are the 3C complexes, on some metals (Pt, Pd) they do not desorb readily and stick to the surface. Results on neohexane reactions (167) show that hydrogenolytic splitting can also occur with the 3C complexes (see, e.g., Pd). The question remaining is: Can the two types of hydrogenolytic splitting (2C and 3C) be related to two different ways of hydrogenolysis? There are indications that the answer should be positive. On Ni, for example, the terminal splitting of n-hexane is very rapid and occurs at such low temperature that the neohexane assay shows almost exclusively the formation of 2CaD complexes. In other words, the terminal splitting can likely be associated with the 2CaP splitting. On the other hand, Pt shows hydrogenolysis at higher temperatures than Ni does, and at low temperatures (i.e., relatively low; low for Pt) this splitting is of highly internal character. At slightly elevated temperatures this splitting
204
VLADlMlR PONEC
goes from this type to increasingly terminal splitting and to isomerization (3C). The internal fission of hexane is promoted by the same factors which promote formation of 3C complexes (see Fig. 15) from neohexane, namely “carbon” deposition and alloying like that of Ni with Cu and Pt with Au. Both factors are known (see above) to invalidate the valley position of metals which prefer to bind carbon atoms there, or one may assume that they strengthen the intrinsic preference of other metals for the summit positions. All these observations and tentative conclusions can be combined to form the following self-consistent picture : (1) The terminal splitting is a reaction via the 2C complexes which, when belonging to group (ii), are preferentially bound to the valleys. (2) The internal splitting is mainly a reaction initiated by the 3C complexes. This is a reaction which seems to prefer the summit position for the carbon atoms of the reacting hydrocarbon molecule. When conditions allow, this complex induces isomerization instead of hydrogenolysis. (3) The terminal 2C splitting complex formation always occurs more easily than the 3Cuy fission, but products of the former splittings stick to the surface of some metals too firmly and do not desorb under the temperature at which the 3C-induced hydrogenolysis or isomerization occurs. This seems to be the case for Pt and Pd and to some extent also for Ir. (4) “Carbon” deposition blocks the valleys rather than the summits, which relatively enhances formation of 3C complexes and relatively suppresses formation of 2C complex. Moreover, “carbon”-covered surface has a lower concentration of hydrogen and this relatively promotes isomerization and, with higher hydrocarbons-dehydrocyclization.
The ratio of rates of formation and removal (by H,) of firmly bound species (“carbon”) is different with different metals. Evidently, Pt and Pd keep more “carbon” on their surfaces than do the good methanation catalysts such as Ni, Ru, or Rh. The surface of, say, Pt is better blocked and thus protected against hydrogenolysis than are surfaces of other metals. The often-found particle size sensitivity of hydrocarbon reactions on Pt (less on other metals) might be related to this. The steady-state concentration of carbon is also a function of the hydrocarbon molecular structure: Higher hydrocarbons are more efficient in modifying the metal surfaces than smaller molecules are. Two extremes emerge from comparison of the Group VIII metals: Ni, Rh, Co, and Ru (the left corner of the Group VIII metal block of the periodic table) prefer terminal splitting, already show multiple splitting at rather low temperatures, are the best catalysts (with 0 s ) in hydrogenolysis of ethane (only 2C complexes possible), and catalyze well the reaction of carbon atoms to methane. Pt is the other extreme in all of these respects, with Pd and Ir
205
CATALYSIS BY ALLOYS IN HYDROCARBON REACTIONS
TABLE VII Complexes Operating upon Reforming of Hydrocarbons on Metals"
Function
Hydrogenolysis type
Hydrogenolysis
Terminal
2 or more
3Cuy
Isomerization or hydrogenolysis
lnternal
5c
Isomerization dehydrocyclization
Fast isomer: 2 or more Slow isomer: 1 Hydrogenolysis: 2 or more Slow dehydrocyclization: 1 Fast dehydrocyclization : 2 or more
Complex 2Cup
-
Required sites
Likely location At least IC in the valley ; most
likely- both "On top" position
"On top" position
T = 500-600 K.
somewhere in between. Based on these facts and on speculation, Table VII summarizes the features of various complexes and reactions in the skeletal reactions of hydrocarbons on the Group VIII metals. The metal not discussed yet is iron. It appeared to be a rather inactive metal. The possible reason for this is that iron is, under a running skeletal reaction or under conditions when more difficult dehydrogenation/hydrogenation can occur, covered by carbon to such an extent that one can rather speak of Fe carbides being the catalyst here. Most likely, the same holds for Group 111-VI transition metals. However, carbides (with an imperfect structure) of these metals are, in contrast to Fe, active in skeletal reactions.
VI.
Conclusions
Considerable progress has been made in accumulating information on the electronic structure of metals and alloys, on some aspects of the structure of hydrocarbon adsorption complexes, etc. Also, information on the relative importance of the electronic structure effects of alloying-as contrasted to the geometric, ensemble size effects-has grown appreciably. With solution alloys the effect of alloying on the electronic structure is surprisingly small, and also with intermetallic compounds these effects are not very pronounced. The effect of alloying on catalytic reactions depends
206
VLADIMIR PONEC
very much on the mechanism of the catalytic reaction and on the type of intermediates operating upon reaction. In parallel-running, simultaneous reactions, the alloying effect is very pronounced if in order for a reaction to occur, large ensembles of certain metal atoms are required. However, the fact that certain metal atoms are required is itself evidence of the importance for catalysis of the electronic structure of metal atoms. The selectivity in consecutive reactions (hydrogenation of multiple unsaturated molecules) sometimes depends on the (relative) heat of adsorption of the starting molecules and intermediate products (e.g., acetylene/ethylene), and since heats of adsorption are usually only marginally affected by alloying, alloying would not change this kind of selective behavior of metals. Various reactions or reaction systems (parallel or consecutive reactions) are influenced by alloying to a quite different degree. This should be kept in mind when attempting to find new or better alloy catalysts for a given reaction.
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CATALYSIS BY ALLOYS IN HYDROCARBON REACTIONS
207
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213. Sinfelt, J . H., in “Catalysis” ( J . R. Anderson and M. Boudart, eds.), Vol. 1, p. 258. SpringerVerlag, Berlin and New York, 1981. 214. Sinfelt. J . H., Adu. Catal. 23,91 (1973). 215. Sinfelt, J . H., Catal. Reu. Sci. Eng. 9, 147 (1974). 216. Merta, R., and Ponec, V.,Proc. h t . Congr. Catal., 41h, Moscow 2, 53 (1964). 217. Walsch, A. D., Trans. Faraday Soc. 45, 179 (1949). Randic, M., and Maksic, Z., Theor. Chim. Acta 3,59 (1965). 218. Bunker, R. J . , and Pyerimhoff, S. D., J. Phys. Chem. 73, 1299 (1969). 219. Guczi, L., Frennett, A., and Ponec, V., Acfa Chim. Acad. Sci. Hung. in press (1982). 220. Maurel, R., and Leclercq, G., Bull. SOC.Chim. France (4). 1234 (1971). 221. Carter, J . L., Cusumano, J. A., and Sinfelt, J. H., J. Catal. 20, 223 (1971). 222. Tetenyi, P., Acra Chim. Acad. Sci. Hung. 107,237 (1981). 223. Boudart, M., Proc. In/. Congr. Catal., 6th, London 1, 1 (1976). 224. Bond, G. C.. Proc. In!. Congr. Catal., 4th, Moscow 2,266 (1968). 225. Schlosser, E. G., Ber. Bunsenges. 73,358 (1969). 226. Burwell. R. L., Jr., Kung, H. H.. and Pellet, R. J . , Proc. In/. Congr. Caral., 6th. London p . I08 (1976). 227. Katzer, J . R.. J. Caral. 32, 166 (1974). 228. Ostermaier. J . J.. Katzer, J. R., and Manonque, W. H., J. Catal. 41, 277 (1976). 229. Somorjai. G. A.. “Chemistry in Two Dimensions: Surfaces.” Cornell Univ. Press, Ithaca, New York. 1981. 230. Somorjai, G. A,, and Blakeley, D. W., Nature (London) 258, 580 (1975). 231. Barbier, J . , and Marecot, P., Nouri. J . Chim. 5,393 (1981). 232. Carter, J . L.. Cusumano, J. A., and Sinfelt, J. H., J . Phys. Chem. 70, 2257 (1966). 233. Martin, G . A., J . Catal. 60,452 ( I 979). 234. Yates, D. J . C., and Sinfelt, J. H., J. Cola/. 8,348 (1967). 235. Vogelzang. M., and Ponec, V.. Adr. Card. Cheni. 11. Proc. Conf. Sdr Lcrke Ciry (1982). 236. Fuentes, S., and Figueras, F., J . Carol. 61,443 (1980). Fuentes, S., Figueras, F., and Comes, R., J . Caral. 68,419 (1981). Fuentes, F., Fuentes. S., and Leclercq. C.. in “Growth and Properties of Metal Clusters” (J. Bourdon, ed.), p. 525. Elsevier, Amsterdam, 1980. 237. Yao, H . C., Yao Yu, F.. and Otto, K., J. Caial. 56,21 (1979). 238. de Jongste, H . C., Kuijers, F. C.. and Ponec, V., Preparation of catalysts. Proc. In/. Synip. Sci. Btrses Prep. Heterog. Catal., Brussel p. 207 (1975). 23Y. Roberti, A.. Ponec. V., and Sachtler, W. M. H., J . Caral. 28, 381 (1973); (see fig. 6). 240. Franken, P. E. C., and Ponec, V., J . Caral. 42,398 (1976). 241. Elford, L..Muller, F., and Kubaschewskii, 0.. Ber. Bunsenges. Phys. Chem. 73,601 (1969). 242. Robbins, C . G . , Claus, H., and Beck, P. A,, Phys. Rev. Lert. 22, 1307 (1969). 243. Perrier, J . P.. Tissler, B., and Tournier. R., Phys. Rev. Lerr. 24, 313 (1970). 244. Vogt, E., Phys. Star. Sol. (b) 50, 653 (1972). 245. Ponec, V., J . Quanr. Chem. 12 (2), 1 (1977). 246. van Barneveld, W. A. A., and Ponec, V., Rec. Trou. Chim. 93,243 ( I 974). 247. Baiker, A., and Richarz, W., In!. Congr. Caral., 7th, Tokyo Communication DI (1980). 248. van Schaik, J . R. H., Dessing, R. P., and Ponec, V., J. Catal. 38, 273 (1975). 249. van Dijk, W. L., Groenewegen, J . A,, and Ponec, V., J. Catal. 45,277 (1976). 250. Duplyakin, V. K.. Yermakov. Yu. 1.- Belyi, A. S., Alfeev, V. S., and Kuzubor. B. N., Kine/. Karal. 19, 1605 (1978). 251. Masai, M., Moni, K., Muramoto, H.. Fuijiwara, T.. and Ohriaki. S., J . Card. 38, 128 (1975).
252. Moss, R. L.. Pope. P., and Gibbens, H. R..J . Catal. 55, 100 (1978).
CATALYSIS BY ALLOYS IN HYDROCARBON REACTIONS
21 3
253. Vlasveld. J. L., and Ponec, V., J . Caral. 44, 352 (1976). 254. Kane, A. F., and Clarke, J. K. A., J . Chem. Sor. Faruday Trans. 176,1640(1980);Bursian,
N. P., Proc. Conf: USSR I, 26 (1978). Biloen, P., Dautzenberg, R. M.. and Sachtler, W. M. H., J . Caiol. 50, 77 (1977). Hagen, D. I., and Somorjai, G . A., J . Catal. 41,466 (1976). Dessing. R. P., and Ponec, V., React. Kinei. Caral. Lett. 5,251 (1976). Gault, F. G.. Zahraa, 0.. Dartiques, J. M., Maine, G . . Peyrot, M., Weisang, F., and Engelhardt, P. A., Proc. Inr. Congr. Caral., 7th. T o k j ~ p. 199 (1980). 259. Karpinski, Z.. and Koscielski, T., J . Caial. 56,430 (1979). 260. Clarke, J. K. A.. Manninger, J.. and Baird, T., J . Cutal. 54, 230 (1978). 261. Plunkett, T.J., and Clarke, J. K . A,, J . Caial. 35, 330 (1974). 262. de Jongste. H. C., Kuijers, F. J., and Ponec, V., Proc. In/. Conyr. Carol., 6rh, London 2,915(1976). 263. Fischer. C. B., Surj: Sci. 62, 31 (1977). 264. Fischer. T.E., and Keleman, S. R., SurJ Sci. 69,485 (1977). 265. Rewick, R. T., and Wise, H., J . Phys. Cheni. 82,751 (1978). 266. Soma-Noto, Y.,and Sachtler, W. M. H., J . Caiul. 32,316 (1974). 267. Primet, M., Mathieu, M. V., and Sachtler, W . M. H., J . Caral. 44,324 (1976). 268. Delmon. J. A.. Primet, M., Martin, G . A,, and Imelik, B., Surf: Sci. 50, 95 (1975). 269. de Jongste, H. C., Ponec, V., and Gault, F. G.. J . Carul. 63,395 (1980). 270. O'Cinneide, A., and Gault, F. G.. J . Caral. 37,311 (1975). 271. Sinfelt, J. H., Synip. Bimei. Caral.. Meer. ACS, Las Vegas(l982);Carter. S. L.,McVicker, G . B.. Weissman, W., and Sinfelt. J. H., ACS Meer. Las Vegas Coll. 018, 183 (1982). 272. Leclercq, G . , Charcosset, H., Maurel, R., Bertizeau, C., Bolivar, C., Frety, R., Jaunay, D., Mendez, H., and Tournayan. L., Bull. Sor. Chin?. Belg. 88,577 (1979). 273. McVicker, G . B., A h . C a r d Clietn. I I , Symp. Sali k k e City (1982). 274. Haining, I. H. B.. Kernball, C., and Whan, D., J . Chem. Res. S p. 364 (1978);Haining, 1. H. B.. Kemball, C.. and Whan, D., J . Cheni. Res. S p. 170 (1977);Haining, I. H. B., Kemball, C., and Whan, D., J . Chem. Res. M. p. 2056 (1977). 275. Clarke. J. K. A.. and Taylor, J. F., J . Cheni. Soc. Faraduy Truns. I71, 2063 (1975). 276. Rasser, J. C., Beindorf, W. H., and Scholten, J. J. F., J . Carol. 59, 21 I (1979). 277. Ramaswamy, A. V., Ratnasamy, P., Sivasankar, S., and Leonard. A. L..Proc. I n / . Conyr. Caial., 6/h, London 2, 855 (1976). 278. Sinfelt. J. H., lecture presented during European F. G . Gault-lectureship, tour, 1980. 279. Menon. P. G . ,and Prasad. J., Proc. In!. Cnngr. Card., 6ih, Lonclon 2, 1061 (1976). 280. Biloen. P., Helle, J. N., Vebeek, H.. and Dautzenberg, R. M., J . Caral. 63, I12 (1980). 281. Ludlum, K. H., and Eischens, R. P., Am. Chem. Soc. 21, 375 (1976). 282. Yermakov, Yu.I.. Plenary lecture. Proc. I n / . Congr. Caial., 7rh, Tokyo p. 57 (1980). 283. Groenewegen, J. A., and Sachtler. W. M. H., J . Caiul. 33, 176 (1974). 284. Gomez, R., Corro, G., Diaz. G., Maubert, A,, and Figueres, F., Nouo. J . Chim. 4, 677 (1980). 285. Gomez. R.. DelAngel, G., and Corro, G.. Nouo. J . Cliim. 4, 219 (1980). 286. Guczi. L., Kemeny. E.. Matusek, K., and Mink. J., J . Chem. Soc. Farudq Trans. I76. 782 (1980). 287. Lam, Y.L., Criado, J., and Boudart, M., Nuurr. J . Chini. I, 461 (1967). 288. Inami. S.H.. and Wise, H., J . Ctiial. 26,92 (1972). 289. Eley. D.D., J . Rrs. Insr. Ctrtcrl.. Hokkoido Univ., Sapporo pp. 16-101 (1968). 290. Bond, G. C., and Allison, E. G., Carol. Reis. 7, 233 (1973). 291. McKee, D.W.. J . Phy.s. Clirm. 70, 525 (1966). 292. Hardy. W. A., and Linnett. .I.W., Trans. Farudav Soc. 66,447 (1970). 255. 256. 257. 258.
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293. Shamsuddin, M., and Kleppa, 0. J., J. Chem. Phys. 71,5154 (1979). 294. Wong. T. C.. Brown, L. F., Haller, G. L., and Kemball. C., J. Chem. SOC.Farada.v Trim. I 71,519 (1981). 295. Guczi, L., Proc. Int. Congr. Catal., 5th, Miami 1, 656 (1972). 296. Haller. G . L., Resasco, D. E.,and Rouco, A. J., Faraday Discuss. Chem. Soc. 72, 109 (1981). 297. Ponec, V.. Adv. SurL Membr. Sci. 13,2 (1979). 298. Visser, C., Zuidwijk, J. G. P., and Ponec, V., J . Catal. 35,407 (1974). 299. Don, J. A., and Scholten, J. J. F., Faraday Discuss. Chem. Soc. 72, 145 (1981). 300. Dominguez, J. M., Vazquez, A. S.,Renouprez, A. J., and Yacaman, M. J.. J . Card. 75, 101 (1982). 301. Anderson, J. R.,personal communication, University of Melbourne. 302. Dalmon, J. A., J . Catal. 60,325 (1979). 303. Luyten, L. J. M., van Eck, M., van Grondelle, J., and van HOOK,J. H. C., J. Phys. Chem. 82,2000 (1978). 304. Foger, K.. and Anderson, J. R., J . C a r d 61, 140 (1980). 305. Balaz, P., and Domansky, R., Petrochemia 18, 198 (1978). 306. Balaz, P., Sotak, I., and Domansky, R., Chem. Zuesti 32,444 (1978). 307. Nazymek, D., React. Kinet. Catal. Lett. 13, 155, 331 (1980). 308. Engels, S.,Langwitz, A., Schuster, L., and Wilde, M., Z. Chem. 20,305 (1980). 309. Burch, R., J . Catal. 71,348 (1981); Burch, R., and Garla, L. C., J . Catal. 71,360 (1981). 310. Johnson, M. F. L., and Keith, C. D., J . Phys. Chem. 67,200 (1963) and refs. therein. 311. Volter, J., Proc. All Union Con/: Catal. Novosibirsk 2, 153 (1978). 312. Driessen, J. M.,Poels, J. K., Hindermann, J. P., and Ponec, V.,J . Catal. 82.26 (1983). 313. Bursian, N. R., Kogan, S. B., and Davydova, Z . A., Kin. Katal. 8, 1283 (1967). 314. Bursian, N. R..Proc. All Union Conf. Catal. I, 26 (1978). 315. Puddu, S., and Ponec, V., Rec. Trav. Chim. 95,255 (1976).
ADVANCES IN CATALYSIS, VOLUME 32
Modified Raney Nickel (MRNi) Catalyst: Heterogeneous Enantio- Differentiating (Asymmetric) Catalyst YOSHIHARU IZUMI Institute for Protein Research Osaka University Osaka, Japan
1. What Is MRNi? . . . . . . . . . . . History of Discovery and Development of MRNi. . Profileof MRNiin Hydrogenation. . . . . . A. Hydrogenation Activity. . . . . . . . B. Kinetics . . . . . . . . . . . . IV. Profile of MRNi in Stereo-Differentiation . . . . A. Enantioface-Differentiating Ability . . . . B. Other Stereo-Differentiating Abilities . . . . V. Other Profiles . . . . . . . . . . . . VI. Surface Conditions . . . . . . . . . . A. Amount of Adsorbed Modifying Reagent . . B. Adsorption Mode of Modifying Reagent . . . VII. Mechanism of Enantio-Differentiation . . . . . VIII. Characterization of Catalyst by Modifying Technique . . . . . . . . . . IX. TA-NaBr-MRNi X. Other Investigations . . . . . . . . . . References . . . . . . . . . . . . . 11. 111.
1.
. . . . . .
. . . . . . . . . . . . . . . . .
.
. . . . .
. . . . .
. . . . .
. . . . . . . . . . . .
. . . . . . . . . . . . . . . . .
. 215 . 218 , , , , , , , , ,
224 224 225 229 229 245 248 249 249 . 250 . 254 ,262 .264 ,267 ,269
What Is MRNI?'
Metal catalysts can be endowed with various new properties by a simple chemical treatment. This catalyst with the new property is called a modified We shall use the following abbreviations throughout: AA for acetylacetone; DNi for nickel catalyst prepared by thermal decomposition of nickel formate; EDA for enantio(face)-differentiating ability (see footnote 2); GA for glycolic acid; HNi for nickel catalyst prepared from NiO by reduction; HNi-1 for HNi prepared from light-green NiO; HNi-2 for HNi prepared from dark-green NiO ; MAA for methyl acetoacetate; MHB for methyl 3-hydroxybutyrate; MRNi for modified Raney nickel catalyst (see p. 216); X-MRNi for Raney nickel catalyst modified with reagent X (see p. 216 and Scheme I ) ; RNiA for RNi pretreated with 2-hydroxy acid; OY for optical yield; and TA for tartaric acid (optically active). 215 Copyrighl 0 1983 by Academic Press. Inc All rights of reproducuon in any form reserved ISBN 0-12-007832-5
216
YOSHIHARU IZUMI sp2-prochiral center
I CH2COOCHj
CHzCOOCH3 I
HO-C -H I
CH3
f
1) generally contains more than N essentially different reactions or mechanisms, where the number varies from case to case and is not a function of N alone. SPACEOF ALL MECHANISMS A. THES-DIMENSIONAL AND THE Q-DIMENSIONAL SPACEOF ALLREACTIONS Let us begin here with a formalization of some of the ideas which have already been introduced. A chemical system contains species which will be denoted by a , , a,, . . ., aA and elementary reactions among these species which will be denoted by the S vectors in Eqs. ( l ) , rl = u l l a l r, = a Z l a l
+ a12a2+ . . . + ulAaA + a2,a2 + . . . + u2AaA
rs
+ us2a2 + . . + C(SAaA
=
aSlal
(1)
'
in which the a's are stoichiometric coefficients. Ordinarily, each elementary reaction will have one or two positive coefficients, one or two negative coefficients, and the remainder equal zero, but more nonzero coefficients are possible. Let us assume, however, that there is at least one positive coefficient and at least one negative coefficient. Any reaction in Eqs. (1) may be written as a conventional chemical equation by setting it equal to zero and transposing the negative terms to the other side of the equation. This notation has been discussed by Aris (14). Chemical equality, denoted by the symbol e, has been shown by Sellers (15) to be a group equivalence, thus satisfying ordinary rules of mathematical equality. Except when specific reservations are stated, every reaction is assumed to be reversible, that is, to be capable of any real rate of advancement, positive or negative. The elementary reactions in Eqs. (1) are not necessarily linearly independent, and, accordingly, let Q denote the maximum number of them in a linearly independent subset. This means that the set of all linear combinations of them defines a Q-dimensional vector space, called the reaction space. In matrix language Q is the rank of the S x A matrix ( 2 ) of stoichiometric coefficients which appear in the elementary reactions ( 1):
280
JOHN HAPPEL AND PETER H . SELLERS
Next, let us define the space of all chemical mechanisms in the chemical system under consideration. Let step si denote the molecular interaction which produces the reaction denoted by ri or R(si). Let a mechanism m be any linear combination of steps of the following form: m
=
alsl
+ a2sz +
*
+ asss
(3)
where each coefficient oi is a real number, describing the rate of occurrence of si. The set of all such mechanisms constitutes an S-dimensional vector space, called the mechanism space. The reaction r produced by m is found by applying the linear transformation R to Eq. (3), which gives the following: r
=
a,rl
+ a2r2+ . . . + asrs
(4)
Since we have explicit expressions (1) for each f i rthey can be substituted into expression (4),so as to express the reaction r as the following explicit linear combination :
thus, obtaining a general expression for any reaction r in the system. It is usual in algebra to express vectors by linear combinations, but it is conventional also, particularly in working examples, to use matrix notation, where only the scalar coefficients are written. Thus, m would be expressed by (al* . . as)and Eq. ( 5 ) by the following matrix equation :
B. THEP-DIMENSIONAL SUBSPACE OF ALLSTEADY-STATE MECHANISMS AND THE R-DIMENSIONAL SUBSPACE OF ALLOVERALL REACTIONS To define a system in a steady state it is necessary to distinguish two kinds of species, the intermediates a , , . ,,a, and the terminal species a,, ,. . . , a , + T , where I + T = A. In such a system a steady-state mechanism is one whose reaction only involves terminal species. The net rate of production of each intermediate in such a mechanism is zero, which is equivalent to saying that the first I coefficients are zero in the general expression (5) for a reaction. This gives us the characterization, introduced by Horiuti ( 4 , 7 ) ,for a steady-
.
,
ANALYSIS OF MECHANISMS FOR REACTION SYSTEMS
28 1
state mechanism as one whose coefficients p l , . . ., ps satisfy the I linear equations expressed by the following matrix equation: (PI . ’ ’ P S I
i“‘ us1
... @;I) ..*
= (O...O)
(7)
us1
If H denotes the rank of the S x I matrix in Eq. (7), then the dimension P of the space of all steady-state mechanisms equals S - H , and the dimension R of the space of all reactions which they produce equals Q - H . Let us describe the reactions in this R-dimensional space as the overall reactions, their essential property being that they involve terminal species only. Horiuti calls H the “number of independent intermediates.” Temkin (10) describes the equation P = S - H as Horiuti’s rule, and the equation R = Q - H as expressing the “number of basic overall equations.” To avoid confusion, let us confine the term basis and the concept of linear independence to sets of vectors, and let numbers such as H,P, Q, R, S be understood as dimensions of vector spaces. This makes it simple to determine their values and the relations among them, as will be done in Section 111. The dimension of a space equals the number of elements in a basis, which is defined as a set of elements such that every element in the space is equal to a unique linear combination of them. Therefore, P steady-state mechanisms can be chosen in terms of which all others can be uniquely expressed. This gives us a unique way to symbolize each steady-state mechanism and its overall reaction, but it does not provide a classification system for them which is valid from a chemical viewpoint, because the choice of a basis is arbitrary and is not dictated, in general, by any consideration of chemistry. A classification system for mechanisms is our next topic. C. DIRECT MECHANISMS AND SIMPLE REACTIONS
In a chemical system there is a unique collection of mechanisms, called the direct mechanisms of the system, which will be shown to be the fundamental constituents of any mechanism. Milner (8) called them “direct paths” and Sellers (9)-“cycle-free mechanisms.” Let m be any mechanism, and let r be the reaction which it produces; then m is defined as direct if it is minimal in the sense that, if one step is omitted, then there is no mechanism for r which can be formed from any linear combination of the remaining steps. Similarly, we can define r as a simple reaction if it is minimal in the sense that, if one of its reactants is omitted, then no reaction in the system involves only the remaining reactants. Let us use the word multiple as an antithesis to direct or simple. As we shall
282
JOHN HAPPEL AND PETER H. SELLERS
see in the examples of Sections V and VI, direct mechanisms for both simple and multiple reactions are of major interest. The set of all direct mechanisms in a system contains within it a basis for the vector space of all mechanisms. In general, there are more direct mechanisms than basis elements, which means that there can exist linear dependence relations among direct mechanisms but, even so, they differ chemically. That is, a direct mechanism with a given step omitted cannot be considered to result from a combination of two other mechanisms in which that step is assumed to occur. In the latter case the net velocity of zero for that step would result from a cancellation of equal and opposite net velocities rather than from the complete absence of the step. The set of all direct mechanisms (unlike a basis) is a uniquely defined attribute of a chemical system. In fact what we have called a direct mechanism is what is usually called a mechanism in chemical literature, even though the definition may be implicit. There are special cases where the direct mechanisms are linearly independent and constitute a basis. If all the direct mechanisms for a particular reaction r are disjoint, in the sense of containing no steps in common, then they are obviously linearly independent, or if there is only one direct mechanism for r, it is independent. This last case suggests a way of finding all the direct mechanisms in a chemical system. If we can find a subsystem which contains at most one mechanism m for any reaction r, then m is direct. In other words, m is a direct mechanism if S = Q in the chemical system, consisting just of the steps in m. Accordingly, to find all the direct mechanisms in a system where Q < S, we may consider each of the (i)subsystems, where
Hence in the entire system there are at most (i)direct mechanisms for r, but usually there are many less than this, not only because of the excluded subsystems, but also because different subsystems can contain the same direct mechanism for a particular r. This approach to finding direct mechanisms is implemented in Section IV, where an efficient procedure is given for making a complete list of all the direct steady-state mechanisms for a given reaction, simple or multiple. All mechanisms can be classified in terms of direct mechanisms. In this article we consider the problem of listing all direct mechanisms in a given system, but we do not undertake to list all combinations which consist of two or more direct mechanisms, advancing simultaneously at independent rates. Each direct mechanism contains a minimal number of elementary steps. Combinations of direct mechanisms in which additional steps must
ANALYSIS OF MECHANISMS FOR REACTION SYSTEMS
283
appear may also be termed mechanisms over the allowable range of such combinations without cycle formation. Such combinations are unique if they are composed of only two direct mechanisms. Combinations involving more than two direct mechanisms are not necessarily unique, in the sense that a given rate of reaction can no longer be represented by combinations of only those direct mechanisms. The combinations of increasing numbers of direct mechanisms will finally include all steps and thus constitute the most general mechanism consistent with the initial choice of elementary steps. The extent to which any given direct mechanisms may be combined without cycle formation can be determined by noting whether such combinations contain irreducible cycles. The latter are the cycles with a minimal number of steps which characterize a given system. They can be determined by a procedure that is analogous to that for finding direct mechanisms [Sellers (9a)I.For a multiple overall reaction, the relative degrees of advancement for each of the simple overall reactions chosen as a basis introduce additional restrictions on the allowable cycle free combinations) [Sellers (9b)I. Except for modeling isomerization systems involving multiple overall reactions, it is generally assumed that a single predominant direct mechanism is sufficient to characterize a given system. Usually the further simplifying assumption is made that there is a single rate-controlling step, other steps in a mechanism being taken at quasi-equilibrium. Another simplification is the assumption of unidirectional steps for reactions that are far from equilibrium.
111.
General Formulas for Mechanisms and Reactions
Equation (7) characterizes a steady-state mechanism algebraically, but it does not provide an explicit formula for any such mechanism. Therefore, using only the matrix (2) of elementary reaction coefficients and the knowledge of which columns in the matrix correspond to terminal species, let us derive a general formula for any steady-state mechanism. A. A
CHANGE OF
BASISFOR
THE
MECHANISM SPACE
The basis s l , . . ., ss for the mechanism space will be changed to one which contains H non-steady-state mechanisms and P steady-state mechanisms. The latter will be what Temkin (16) calls a “basis for all routes.” A route is what we are calling a steady-state mechanism, and a basis is a maximal linearly independent set of them. The basis is “stoichiometric” in Temkin’s
284
JOHN HAPPEL AND PETER H. SELLERS
terminology when it is selected as follows: R of the P basis elements are mechanisms whose reactions constitute a basis for the space of overall reactions, and the remaining C basis elements are what Temkin (10) calls “empty routes,” each of which has a reaction equal to zero. Let us use the mathematical term cycle to describe any element m of the mechanism space such that R(m) = 0, instead of “empty route.” The C elements mentioned are a basis for the space of all cycles. In algebra this space is known as the “kernel of R.” To construct the new basis, described above, we will change the basis s l , . . ., ss of the S-dimensional space of all mechanisms into a basis (9), which separates into 3 parts, ml,...,mH;
mH+,,...,mQ;
mQ+,,...,mS
(9)
where each mi is a linear combination of the original basis elements, such that (mQ+ , . . ., m,) is a basis for the subspace all cycles; (mH+ . . ., mQ) is a set of steady-state mechanisms no linear combination of which is a cycle, and (ml , . . ., mH) is a set of mechanisms no linear combination of which is a steady-state mechanism. To get basis (9), start with following matrix :
,
ss \us,
.’ ’
as1
%,I+,
’
.’
%,I
+TJ
which is the same as matrix ( 2 )except that we now require the first I columns to correspond to intermediates and the remaining columns to correspond to terminal species. Furthermore, the rows are labeled by the steps which correspond to them [i.e., aij is the stoichiometric coefficient of species a j in R(s,)]. Next, perform elementary row operations on matrix (lo), so as to put it in the diagonal form shown in Fig. 1. This will require interchanging some pairs of columns (but a column corresponding to an intermediate is never interchanged with one corresponding to a terminal species). Every time a row operation is performed it is also applied to the column of basis elements at the left of the matrix, with the result that the entries in this column become linear combinations of steps which will achieve form (9) when the diagonalization is complete. For instance, if the first row operation is to replace row j by row j minus row i, then the column entry sj is replaced by ( s j - si).There is no need to describe the diagonalization procedure in detail, except to say that it must be performed in such a way as to insure that the combinations m, , . . ., m, are linearly independent. This will be achieved if the elementary row operations are confined to changing a row by adding to
ANALYSIS OF MECHANISMS FOR REACTION SYSTEMS
285
FIG. I . This diagonalized matrix ( Pij) is formed from matrix ( a i j )by elementary row operations and column permutations. It has the same rank as (aij).
or subtracting from it one of the rows above. The only row permutation needed is moving a row of zeros to the bottom of the matrix. B.
BASICOVERALL REACTIONS, STEADY-STATE MECHANISMS, AND CYCLES
An explicit basis for the space of overall reactions is characterized by rows H + 1 through Q of the diagonalized matrix of Fig. 1. If the matrix is expressed by (Pij), then the desired basis is R(mH+I) = P H + l , l + l a l + l + ." + P H + l . A a A R ( m ~ + 2=) B H + ~ , I + ~ ~ I *+" I -I PH+2,AaA R(mQ)
=
DQ,l+laI+l
+
* * '
+
( 1 1)
PQAaA
A basis for the space of all steady-state mechanisms is (mH+1 , . . ., ms). Since mi is a linear combination of steps, this basis has the following form: mH+l
=
YH+1,lSl
+
" '
+ YH+l,SSS
286
JOHN HAPPEL A N D PETER H. SELLERS
The rows in (12) from m?+ through ms are a basis for the space of all cycles, and the coefficients in these rows form a C x S matrix (yij), which will be needed in Section IV for the construction of the unique set of all direct steady-state mechanisms.
C. ALGEBRAIC FORMULAS
Every element of a space is a unique linear combination of its basis elements. Therefore, a general expression for any steady-state mechanism m, including cycles, has the following form: (13) =pH+lmH+l + C1H+2mH+2 + " ' + h m S where the coefficients are any real numbers. A general expression for any overall reaction r is obtained by applying the function R to Eq. (13), which gives the following: = pH+
lR(mH+1)
+
' * *
+ PQR(mQ)
(14)
These expressions for m and r are made explicit by substituting into them the values for mi and R(mi)stated in bases (1 1) and (12).
D. MULTIPLE OVERALL REACTIONS
If R = 1 in a chemical system, it means that all steady-state mechanisms [i.e., all m which can be obtained by assigning particular numerical values to p l , . . ., ps in Eq. (13)] will have the same overall reaction r or a multiple of it, because then Eq. (14) reduces to r = pH+lR(mH+]). In this case the system is said to have a simple overall reaction, and, when we come to list all the direct mechanisms for r, there is no loss of generality if we take the multiple pH+ to be unity. If, however, R > 1, then the general formula (14) for an overall reaction r involves two or more independent parameters and is said to be a multiple overall reaction. In such a case each direct mechanism for r must also involve these parameters, unless we are prepared to choose particular overall reactions and determine a list of direct mechanisms for each. However, if we take a basic set of overall reactions and combine the direct mechansims of all of them, the process of combining them will lead to nondirect mechanisms. Accordingly, to generate all the direct mechanisms in a system, we must
287
ANALYSIS OF MECHANISMS FOR REACTION SYSTEMS
find them for the general overall reaction (14). This is contrary to the approach suggested by Lee and Sinanoglu (13).
IV.
A Procedure for Finding Every Direct Mechanism
A procedure, which was introduced by Happel and Sellers (I) for finding all direct mechanisms for a given reaction will be demonstrated here from the standpoint of how to apply it in practice. We demonstrate it by applying it to an arbitrary S-step chemical system, as defined in Section 11, to find all the direct mechanisms for the general overall reaction (14) derived in Section 111. A. THECYCLE-FREE SUBSYSTEM In a chemical system with S steps and a maximum of Q linearly independent elementary reactions every set of Q steps whose reactions are linearly independent constitutes a cycle+ee subsystem. It is apparent that, if R(sl),. . ., R(s,) are linearly independent, then no cycle can be formed with the steps s l , . . ., sQ(unless all the coefficients are zero). Accordingly, a laborious way ways in to find all the cycle-free subsystems would be to consider all the which Q rows can be chosen from the S x A matrix (2) and select those which are row independent. Since each row corresponds to a step, each selection of Q independent rows gives a cycle-free subsystem. The same result can be achieved by considering columns C at a time in the following smaller C x S matrix: . . . YQ+ 1.S ?Q+ 1 . 1 ?Q+ 1 . 2
(z)
?Q+2.1
YQ+2.2
Ys2
...
".
YQ + 2,s
Yss
called the cycle matrix, whose entries are defined by the basic cycles constructed in Section III,B and given explicitly in the last C rows of basis (12). Each column of this matrix corresponds to a step in the system, and each linearly independent set of C columns corresponds to a set of steps which, if removed from the system, would leave a Q-step cycle-free subsystem. This procedure for finding the Q-step cycle-free subsystems was introduced by Happel and Sellers ( I ) , and is equivalent mathematically to the procedure described in the last paragraph, but it takes advantage of the fact that the
288
JOHN HAPPEL AND PETER H. SELLERS
basic cycles for the system have already been determined by diagonalizing matrix (2).
B. THEDIRECTMECHANISMS Let us find every direct mechanism for a given overall reaction r. Assume r to be of the general form given in Eq. (14) and of multiplicity R ( R = Q - H), which means that an expression for it contains R parameters. Any mechanism for r is of the general form given in Eq. (13) and depends not only on the R parameters in its reaction, but on C additional parameters, where C is the dimension of the space of all cycles ( R C = S - H). Therefore, to determine a unique mechanism for r, we need to determine C parameters, and they can be chosen to make it a direct mechanism by the following reasoning: (i) Every direct mechanism belongs to a Q-step cycle-free subsystem and (ii) every Q-step cycle-free subsystem is obtained, as shown in Section IV,A, removing C appropriate steps from the system as a whole (Q = S - C ) . Taking (i) and (ii) together, we obtain a direct mechanism for r by rewriting the general mechanism (1 3) as a linear combination of steps, then setting C coefficients equal to zero, where they are the coefficients of C steps whose removal defines a cycle-free system. Let us clarify this procedure by applying it to four cases of increasing complexity: (1) C = 0; (2) C = 1, R = I ; (3) C>l,R=l;(4)C>l,R>l. Case I . If C = 0, there are no cycles in the system. For any reaction r in the system there is one mechanism, and it is direct. This applies to Example 1 in Section V and Example 6 in Section VI. Case 2. If C = 1 and R = 1, then the general mechanism, as given by To simplify the notation, Eq. (14), takes the form pH+lm,+l + pH+2mH+2. let this be written as pm + 4n, where m is a mechanism, n is a cycle, and the coefficients p and 4 are unrestricted. The overall reaction depends only on p and is a simple reaction. Ordinarily, with simple reactions the scalar coefficient is omitted, since r and pr are essentially the same reaction. Therefore, in the present case the general mechanism is taken to be of the form m + 4n. It will be recalled that m and n are fixed linear combinations of steps, which result from the diagonalization procedure described in Section III,A. Accordingly, m + 4 n can be expanded to a linear combination of steps, whose coefficients depend on 4 or else are constant. This becomes a direct mechanism if any single nonconstant coefficient is set equal to zero. In the first place this removes one step, which means that the mechanism belongs to a cycle-free subsystem and thus is direct. Second, setting a nonconstant coefficient equal to zero allows us to solve for 4, which means that all other coefficients are known, and the direct mechanism is completely
+
ANALYSIS OF MECHANISMS FOR REACTION SYSTEMS
289
determined. Furthermore, we get every possible direct mechanism if we go through the above procedure for every nonconstant coefficient. The above procedure applies to Example 2 of Section V, where it is carried out in full detail. Case3. If C > 1 and R = 1, then the general mechanism, as given by Eq. (1 4), takes the form p H+ mH+ + . . . + psms, in which mH+ is a mechanism and the other m’s are cycles. Rewrite it as pm + 41nl + . . . + &n,, where m is a mechanism, n, through n, are cycles and the coefficients are unrestricted, which means, as in Case 2, that we may set p equal to unity without loss of generality. To find each direct mechanism, expand m + 4 ,n, + . . . + &nc to a linear combination of steps, and set C coefficients equal to zero, where the choice of coefficients is made as follows: The cycles are of the following form:
,
n,
=
vllsl
+ . . . + vlsss
n,
=
vclsl
+ + vcsss
(16)
* * *
The C steps whose coefficients are set equal to zero and must correspond to C columns in (16) whose C x C matrix of coefficients is nonsingular. This choice will guarantee that the resulting mechanism is direct. Furthermore, each coefficient which we set equal to zero gives rise to a linear equation in the variables 41,. . ., 4,. We have C such equations, which are solvable because of the nonsingularity of the C x C matrix of coefficients. Accordingly, for each nonsingular C x C matrix of coefficients in (1 6), there is a complete solution for 41,.. ., &-,which means that the direct mechanism is completely determined. Furthermore, we get every possible direct mechanism, if we go through this procedure for every nonsingular C x C matrix of coefficients in (16). The above procedure applies to Examples 3, 4, and 5 of Section V, and it is carried out in full detail in Example 3. Case 4. If C > 1 and R > 1, then we have the general situation described in Section II1,B. Using Eqs. (12) and (13) from that section, we arrive at an explicit expression for the general mechanism as the sum of all of the following expressions : pH+1YH+1,ls1
+
~H+RYH+R.ISI
+ ... + P H + R Y H + R . S S S
PSYSlSl
” *
+ pH+lYH+l,SSS
+ . ’ . CCSYSSSS
The first R rows add up to a mechanism for the general overall reaction, and the remaining C rows are cycles. As in Case 3, our object is to choose coefficients for the cycles such that C of the columns in (1 7) add up to zero. The
290
JOHN HAPPEL A N D PETER H . SELLERS
C columns we choose must have the property that matrix ( y i j ) from these columns and from the last C rows of (17) form a nonsingular C x C matrix M. For simplicity of exposition let us suppose that the first C columns are the ones chosen, then M = (yij) where H + R + 1 5 i IS and 1 Ij IH . Our object now is to have the first C columns of (17) add up to zero. dSss,which is a Denote the sum of the first R rows of (17) by glsl * * mechanism for the overall reaction. Then the statement that the first C columns of (23) add up to zero is equivalent to the following matrix equation:
+
( O l * * ' O R )f (p(H+R+I"'pS)M
=
+
(18)
whose solution is (pH+R+I
" ' p S ) = (-Ol
* "
- CTR)M-
(19)
In other words, if the C coefficients P,,+~+ . . . p s are given the values determined by Eq. (19), then the total of the expressions in (17) will be a direct mechanism. Furthermore, if we go through this procedure for every selection of C columns in (1 7) such that the C x C matrix M is nonsingular, then we get every direct mechanism for the overall reaction (14). Altogether there are R + C undetermined coefficients pH+ 1 , . . ., ps in (17), the last C of which are determined for each direct mechanism. The remaining R parameters pH+ . ., P,,+~ are in the expression (14) for the overall reaction, which is of multiplicity R. Similarly each direct mechanism must be a function of these R parameters. The above procedure is carried out in the treatment of Example 9 in Section VI to obtain an initial direct mechanism. In all of the cases treated above the set of direct steady-state mechanisms which has been generated is exhaustive. However, it is possible for repetitions to occur among the mechanisms, but we can eliminate the possibility of repetitions in the following manner. Each time C values for pH + R + . . ., ps are determined, substitute them in the general mechanism, express it as the following linear combination of steps : ZilSl
+
..*
+
ZiSSS
(20)
and put its coefficients in the last row of Table I, which is a display of all direct mechanisms found, so far. To get the next line, choose a set of C columns from (17) which has not yet been considered. There are ( 5 ) ways of choosing columns which must be considered, but the computation (19) does not have to be carried out if either of the following conditions holds: (i) There is a row in Table I which already has zeros in the C columns we are considering. (ii) The matrix M determined by these columns is singular. These precautions will eliminate repetitions among the rows of Table I
ANALYSIS OF MECHANISMS FOR REACTION SYSTEMS
29 1
TABLE I
more simply than if they are thrown out after they are found. This saves unnecessary matrix inversions. Each of the (): column selections must be tested. Any systematic way of ordering these selections so that none are omitted is acceptable. This additional procedure to avoid repetitions is carried out in the last part of Example 9 in Section VI. V.
Systems with a Simple Overall Reaction
Most systems treated in the literature exhibit a simple overall reaction, which can be uniquely represented by a conventional chemical equation. In addition, the elementary reactions are usually selected so that all of them must be combined to form the overall reaction, which means that the system is cycle free and that there is no mathematical distinction between an elementary reaction and the step which produces it. Often the combination of steps giving the overall reaction is such that each intermediate is produced by exactly one step and consumed by exactly one step. The following example illustrates such a system. 1 . SULFURDIOXIDE OXIDATION (No CYCLES) EXAMPLE
The important commercial process of sulfur dioxide oxidation has been studied by a number of investigators. A set of steps that has been proposed for both platinum and vanadium oxide-based catalysis by Horiuti (7) for the overall reaction 2 S 0 2 + O2 + 2 S 0 3 is as follows: sl:
0,+ 2 1 s 2 0 1
s,:
so, + 1 ==S 0 , I
sj: S 0 , I s4:
+ 01=
S0,l
+1
so,/= so, + I
This is the form in which steps will be listed in all our examples-a
symbol
29 2
JOHN HAPPEL AND PETER H. SELLERS TABLE I1 OI
S0,I
2 0
0
SI s2
s, s4
I
-I
-I 0
0
so,/ 0 0 1 -1
0,
I
-2
so, 0
-1
-1 1
I
0
0 0 0 I
so,
-I
0 0 0
so,
0
TABLE I11 01
S0,I
SO&
I
0,
0 I 0
0 0 2
-2
-1
0
0
SI
+ 2s, + 2s,
2 0 0
SI
+ 2s, + 2s, + 2s,
0
SI s2
-I
-2
0 0 0
-1
-2
2
0
-1
-2 0
0
so,
-1
for the step, followed by a chemical equation for the elementary reaction produced by that step. In this example, but not in the subsequent ones, let us rewrite the elementary reactions in the following more formal vector notation: R(s,) = - 0 2 - 21 201
+ R(s2) = -SO2 - 1 + S 0 2 f R(s3) = - S 0 2 1 - 01 + S03f + 1 R(s4)= - S 0 3 1 + SO3 + 1
(22)
This displays the convention, tacitly assumed later, that the positive direction of a step corresponds to the advancement from left to right of the stated chemical equation. The matrix of stoichiometric coefficients for these reactions is shown in Table 11. The diagonalization of the matrix in Table I1 gives the matrix in Table 111, from which the steady-state mechanism is s1 + 2s, + 2s3 + 2s4. In Horiuti’s terminology the “stoichiometric numbers’’ are 1 for s1 and 2 for s 2 , s 3 ,and s4. Often, even in the case of simple reactions, it is possible to encounter systems with cycles. The following is an illustration of this situation.
EXAMPLE 2. THEHYDROGEN ELECTRODE REACTION ( 1 CYCLE) This system has received considerable attention. Milner (8) includes it in a study of several electrode reactions, and Horiuti (7) uses it as the basis of a general discussion.
293
ANALYSIS OF MECHANISMS FOR REACTION SYSTEMS
TABLE IV H
H,
Ht
*
The overall reaction of the system is given as 2H+ + 2eH,, and 3 steps are postulated whose elementary reactions are as follows : s,: H t + H + e - = H 2 s,:
+ e- + H +HeH,
H+
s3: H
Since H + and e- are always together, let us regard H+ + e - as a single component and write it simply as H + . The matrix of stoichiometric coefficients is given in Table IV, the diagonalization of which gives the matrix in Table V. From this we conclude that the general steady-state mechanism is as follows :
+ ~ 2 +) 4 ( ~ -1 sz
( ~ 1
-
(1
~ 3= )
+
+ (1
4)~1
-
4 1 ~ 2- 4
~
3
(24)
where the coefficient 4 is unrestricted. Following the method of Case 2 in Section IV,B, we find each direct mechanism by setting one of the coefficients equal to zero in the right-hand side of Eq. (24). This gives three possible values - 1, 1,0 for 4. Putting them in Eq. (24), we get all the possible direct mechanisms as follows : m,
=
2s,
+ s3
m,
=
2s,
-
m3 = s1
s3
(25)
+ s2
This result may be tabulated as shown in Table VI. In subsequent examples this sort of table will be the principal way of listing direct mechanisms. There are three elementary steps and one independent intermediate, so TABLE V
s2
H
H,
1
1 0
H+ -I
294
JOHN HAPPEL A N D PETER H . SELLERS T A B L E VI
m, m2
m,
0 2 I
2 0 1
I
-I 0
that there are 3 - I = 2 reaction routes according to Horiuti’s rule. Miyahara ( 1 7 ) and Horiuti (7) noted that any two of the three direct mechanisms could be combined algebraically to obtain the third. However, they are distinct from each other chemically since any two of them contain a step that is not involved at all in the third.
EXAMPLE 3. AMMONIA SYNTHESIS (2 CYCLES)
*
The reaction N, + 3H2 2NH, has been studied extensively from a mechanistic viewpoint. Horiuti (7) and Temkin (11) have proposed entirely different mechanisms for this reaction. Recognizing all steps in both mechanisms as possibilities, we find that there are in all 6 direct mechanisms, including the proposed ones, all of which produce the same overall reaction. In the following system for ammonia synthesis steps s l , s2,s 3 ,and s, were proposed by Temkin and s4, s5, s6, s8, and sg by Horiuti :
+ I* N,I N,I + H, * N,H,I N,H,I + I * 2NHI N, + 2/* 2NI NI + HI*NHI + I NHI + HI* NH,I + I NHI + H, * NH, + I H , + 21* 2HI NH21 + H I + NH, + 21
s , : N, s2:
s,: s4:
s5: s6: s,:
sg: sg:
The symbol I in system (26) refers to an active surface site on the catalyst. Every species with I in it is an intermediate and the rest are terminal species. For the purpose of our analysis we could omit I wherever it appears alone as a reactant. Notice that by including it as an intermediate, we get a case of a system where H < I , or in Horiuti’s terminology, the intermediates are not all independent. By diagonalizing the matrix of stoichiometric coefficients, we obtain the matrix given in Table VII. The seventh row shows that there is a simple
295
ANALYSIS OF MECHANISMS FOR REACTION SYSTEMS
TABLE VII N21 N,H21 NHI N I SI s2
+ s1
s3
s2
+ + s1
s4
+ sq - s3 - s2 - s, - 2s, - s4 + 2s3 + 2s, + 2s1 2% + s3 + s2 + s , sg + 2s, + s4 - s3 - s2 - S I + - s, + S 6 2s, 2s,
sg
sg
HI
NH21
I
N2
H2
-1 -I
-I
I 0
0 1
0
0
0
0
-
I
0
0
0
0
-
I
0 0 0 0
0 0 0 0
2 0 0 0
0 0 0 2 0 0 0 - 2 0 0 0 2
0
0
0
0
0
0 0
0
0 0
0 0
0 0
NH3
0
2 2 2 -2
-I -I
-I
-2
0
0
0
-I
-3
2
0 0
0
0 0
0 0
-I 0
0
0
1
0 ~
0
0
0 0
overall reaction, and the last two rows show that the space of all cycles is of dimension 2. The general mechanism is given' by the following equation: 2S7
+ + s2 + + s3
s1
&s8
+ 2 S 5 + s4 -
s3
4)(s1
+ $(s9 + s8 - s7 + s6) +
- s2 - s1)
+ s2 + s3) + &s4 + 2s5) + $ ( s 6 + ( 2 - $ b 7 f (6+ 9 b 8
= (I -
s9)
(27)
and $ are unrestricted. Notice that the combinations of steps (sg + sg) may be treated as single steps. The cycle matrix is given in Table VIII. There are two singular 2 x 2 submatrices in Table VIII, consisting of columns 1 and 2 and columns 3 and 4. The remaining eight 2 x 2 submatrices are nonsingular. This means, according to the Case 3 in Section IV,B, that each direct mechanism is obtainable by setting a pair of coefficients equal to zero in the second line of Eq. (27) and solving for 4 and $. This leads to six distinct solutions, as shown in Table IX, which considers every pair of coefficients. Cases 1 and 8, which correspond to singular 2 x 2 submatrices of Table VIII do not have solutions, and cases 7 and 9 are repetitions of previous solutions. Accordingly, we are left with six pairs of values for 4 and $, which may be substituted in Eq. (27) to get the direct mechanisms given in the following list: where
(sl
CI#
+ s2 + s3), (s4 + 2s,), and
+ 2s5) + 2s7 + s8 m2 = (s4 + 2s5) + 2(s6 + s9) + 3s8 (Horiuti) m3 = (s4 + 2s5) + s9) + 3s7 m4 = ( s , + s2 + s3) + 2s7 (Temkin) m5 = (sl + s2 + s3) + 2(s6 + s9) + 2.5, m6 = 3(s, + s2 + sj) - 2(s4 + 2 4 + 2(s6 + s9)
m, = (s4
- (sg
and tabulated in Table X.
0 0 0
296
JOHN HAPPEL AND PETER H. SELLERS
4 9
I 0
-I 0
0 I
0
1 1
-I
TABLE IX Case
Pairs of coefficients
1-4.4 I -4.G 1-4,2-9 1 -l#J,4+9 4, 4*2- 9 4.4 t G 9.2 - 9 9.4 + 9 2 - 9.4 t J,
I 2 3
*
4
5 6
I 8 9 10
Solutions none fj=I,+=O
4=1.+=2 +=I,$= -1 fj = 0, I j = 0 4=0,$=2
+=o,g=o 4 4
= =
none 0,9 = 0
-2.9
=
2
TABLE X SI
m1 m2 m3
m4
m, m6
+ s2 + sj 0 0 0 I I 3
s4
+ 2s,
S6
+ sg
s,
sg
0 2
2 0 3 2 0 0
1 3 0 0 2 0
I 1
I 0
0 -2
-I 0 2 2
m, and m4 are the mechanisms proposed by Horiuti and Temkin, respectively. m3 and m6 might be omitted on the grounds that some of their steps proceed in the wrong direction, but m, and m5 remain for consideration. Rate equations for simple reversible reactions are often developed from mechanistic models on the assumption that the kinetics of elementary steps can be described in terms of rate constants and surface concentrations of intermediates. An application of the Langmuir adsorption theory for such development was described in the classic text by Hougen and Watson (18), and was used for constructing rate equations for a number of heterogeneous catalytic reactions. In their treatment it was assumed that one step would be rate-controlling for a unique mechanism with the other steps at equilibrium.
297
ANALYSIS OF MECHANISMS FOR REACTION SYSTEMS
Such rate expressions are often termed Langmuir-Hinshelwood-HougenWatson (LHHW) equations and are widely used in chemical engineering [see Froment and Bischoff (19)].The usual procedure is to postulate plausible mechanisms without considering cycles, as in Example 1. In such cases it may be desirable to develop the complete list of possible direct mechanisms even if further considerations can rule out some as being unlikely. The following example illustrates a typical case. EXAMPLE 4. DEHYDROGENATION OF 1-BUTENE TO 1,3-BUTADIENE (3 CYCLES) Froment and Bischoff (19) report a study of the dehydrogenation of 1-butene to butadiene on a chromia-alumina catalyst. Neglecting isomerization of 1-butene, the following steps are postulated:
+ I * C,H,I + I*H21 H,I + I = 2HI H, + 2/* 2HI C,H,I + I * C,H,I + HI C,H,I + I* C,H,I + HI C,H,I * C4H6 + I C4H,I + I* C,H,I + H,I C,H,I + 21* C,H6/ + 2HI
s, : C,H, s,:
s3: s,: s5 :
s6: s, : s,:
s9:
H,
The overall reaction is
TABLE XI
1 0 0 0 0
'6 s5
s3 s2
s2
s, s, s*
+ s, + s5 - s3 - s2 + s, -
sg - s,
- S6 -
s9 - S6 -
s5 s5
+ s3
-
1
I 0 0 0
0
0
0 0 0
0 0 0
1
1 2
-
0 0
0
0 0 0 - 1 1 0 1 0 0 1
0 0 0 0
- 1 - 1 - 1 - 1 - 1 - 1
0 0 0 0
0 0 0
0 0 0 0
0 - 1
0 0 0
0 0 0
0 0
0
0 0 0 1 0
I
I
0 0 0
0 0 0
0 0 -
298
JOHN HAPPEL AND PETER H. SELLERS
The diagonalization of the matrix of stoichiometric coefficients is simplified in this case ifthe rows are not ordered as in steps (35).The result of diagonalizing is given in Table XI. Then, using the methods of Section IV,B, we find all the direct mechanisms of Table XI1 for the overall reaction
+ C,H,
C4H,
+ H,
Examination of the matrix in Table XI1 shows that 6 mechanisms are possible on the basis of the steps proposed by Bischoff and Froment. They identified m , , m2, m4, m5, and m6, and developed 15 rate equations corresponding to various choices of rate-controlling steps. After a set of experiments involving sequential testing and model discrimination, they retained mechanism m4 with step s1 as the rate-controlling step. According to the present procedure m, might be an additional mechanism to consider. A scheme which was considered by Hougen and Watson (18)and is slightly more complicated is the hydrogenation of isooctene codimer. This illustrates a reaction with a large number of cycles compared with the number of intermediates.
EXAMPLE 5. HYDROGENATION OF ISOOCTENES (4 CYCLES) A supported nickel catalyst was used to study the reaction in which mixed isooctenes, commercially known as codimer, are hydrogenated in the vapor phase to the corresponding isooctanes. Neglecting isomerization, the following steps are assumed to occur:
+ I*
s , : C,H,,
s,:
H,
+ I*
s3: C,H,,/
C,H,,I
H,I
+HJe
s4: C , H , , I ~ C , H , , s5:
H,
+I
+ 2 / * 2HI + 2 H I e C,H,,/ + 21
C,H,,I s,: C,H,, s,:
C,H,,I +I
s,:
C,H,,
sg:
H,
+ H,/* C,H,,I + 2 H I e C,H,,/ + /
+ C , H , , I ~C,H,,/
The overall reaction is It is not necessary to know in advance what the overall reaction is (Hougen and Watson assumed that it might also occur as an uncatalyzed reaction in the gas phase). For our purposes it is enough to know what the terminal
299
ANALYSIS OF MECHANISMS FOR REACTION SYSTEMS
TABLE XI1 s1
m1
+ s,
s,
I I I
m2 m 3 .
1
-1 -1 -1
m6
I
ss
s4
-I -1 -I 0 0 0
0 0 1 0
0 0 0
m4 m5
s3
-I -1
+
s*
sg
0
0
1
1
0 1 I
0 0 0
0 0
1 0
S6
0 0 0 0
species are, since the overall reaction among them is furnished by the diagonalization (see Table XIII) of the matrix of stoichiometric coefficients in steps 30. Using the methods of Section IV,B, we find all the direct mechanisms (see Table XIV) in this system for the overall reaction H2
+ CSH,, * CSH,,
The matrix given in Table XIV shows that there are nine direct mechanisms. Five of these, namely, m2, m5, m7, m,, and m,, were identified by Hougen and Watson. Seventeen different mechanisms with single ratecontrolling steps were modeled and tested for agreement with observed kinetic data. The model corresponding to m7 with s, as the rate-controlling step was chosen as the recommended rate equation. The mechanisms obtained by our procedure in addition to those developed by Hougen and Watson are m, , m3, m4, and m6. The large number of these mechanisms is related to the fact that four cycles appear in the diagonalized matrix. These models have unusual features that would probably not be noticed unless a procedure like this were employed. For example, in mechanism m, neither isooctane nor hydrogen adsorbs directly on the
TABLE XI11
-1
s3 S2
SS SS
+ sg + ss s1 + s, + -s2 + s5 + s2 - ss + s, S) + sg s4
s3
sg -
S)
S8
-s2
-
S6
s*
0 0 0 0 0 0 0 0
-1 1
1
0
0
0
0 0 0 0 0 0 0
0 0 0 0 0 0
1
-2
2
1 -1
I -2
-I
0 0 -1 0
0 0 0 0
-I
-1
1 0 0 0 0
0 -1 0
~
0
0
0 0 0 0
0 0 0 0
0 0 0 0
0
0 0 0
300
JOHN HAPPEL AND PETER H . SELLERS
TABLE XIV
m, m2 m , m 4 m5
m, m, mR m,
'I
'2
'3
0 0
0 0 0 0 I I 0 0 I
0 0 -1 -1 0 1 0 0 1
O 0 0 0 I I I
' 5
'6
1 1 1 1 1
0 1 0
-I
I
0
1 1 1
0 1 0
'4
'7
0 1
1 1 0
0 0 0
'8
'9
1 0 1
I
I
I
0 . 0
0
0
1
0 0 1 0
0
0 1
0 0
0 1 0 0
0
0
0
-I
0
catalyst, but instead reacts with adsorbed species. In this mechanism HI is not produced from the initial reactants and is recycled. If it is assumed that the reaction will be characterized by a single direct mechanism as well as a single rate-controlling step, one possible model is exhibited for every nonzero entry in the matrix of Table XIV. Tracer techniques with high-speed computers are useful in relaxing the requirement of a single rate-controlling step. Direct mechanisms can also be combined, if care is taken to avoid the possible occurrence of cycles. In Example 5 all nine direct mechanisms can be combined without cycle formation, though this is not always the case [Sellers (9a)I. VI.
Overall Reactions with a Multiplicity Greater Than One
Each system considered in this section has a space of overall reactions whose dimension exceeds one. In many industrial reactions involving organic substances a major product is formed, but a side reaction contributes to loss in selectivity or yield of the desired product. Such cases may be said to exhibit a multiple overall reaction, unless the ratio of desired product to by-product remains constant over a range of operating conditions, so that a simple chemical equation might be employed to express the stoichiometry. It is important to note that in these cases one cannot add up the separate direct mechanisms for all the simple reactions which add up to the overall reaction and expect, in general, to get direct mechanisms for the overall reaction, unless there are no common steps in the mechanisms of the simple reactions that form a basis for the system. The direct mechanisms for the reaction systems are, however, unique and any observed rates of appearance or disappearance of terminal species can
ANALYSIS OF MECHANISMS FOR REACTION SYSTEMS
30 1
be expressed in terms of these mechanisms. If one such mechanism predominates, it will be possible to express the rates of changes of terminal species in terms of step velocities for that mechanism. A typical process is the oxidation of ethylene over silver catalyst with a side reaction to produce undesirable carbon dioxide. Miyahara (17) was among the first to appreciate the problem of assigning unique mechanisms to this system and demonstrated that an appropriate set of steps could be chosen that corresponds to a single mechanism. In this mechanism, discussed by Happel and Sellers (Z), there are seven elementary steps and five independent intermediates, and following Horiuti’s rule there are 7 - 5 = 2 independent “routes.” This is also equal to the number of independent reactions and there are no cycles. Miyahara and Yokohama (20)considered this reaction further, employing a different mechanism with four steps and two intermediates and again arrived at a single unique mechanism for the system. Both these mechanisms followed the original view that dissociative adsorption of oxygen occurred and that the epoxide was formed as a result of interaction of ethylene with one adsorbed oxygen atom. However, later results [see Patterson (21)]suggest that diatomic oxygen is involved in the formation of the epoxide. The following two examples employ such mechanisms for the purpose of illustration, although a recent survey by Sachtler et al. (22)indicates that both views are still tenable.
EXAMPLE 6. ETHYLENE OXIDESYNTHESIS (No CYCLES) The following steps are postulated for the oxidation of ethylene:
sg:
+ 0,I + C,H, * C,H,O + OI C,H4 + 601 * 2C0, + 2H,O + 61
s4:
201*
s , : 0, I*O,I s,:
0,
(31)
+ 21
where all species involving the symbol I are intermediates and the rest are terminal species. From the diagonalization of the matrix of stoichiometric coefficients given in Table XV, the general mechanism (32) and the general overall reaction (33) can be found:
+
+ + ~ ( 2 +~ 21 ~ 2+
~ ( 6 ~ 61 ~ 2 sj) P(-7CzH4
-
~ 4 )
(32)
6 0 2 + 6 C 2 H 4 0+ 2 C 0 2 + 2 H 2 0 )
+ a(-2C2H4
-
0,
+ 2C2H40)
(33)
Since there are no cycles in the system, the general mechanism (32) is a direct mechanism for a multiple overall reaction (33), where p and a are unrestricted
302
JOHN HAPPEL AND PETER H. SELLERS
TABLE X V
6s, 2s,
+ 6s, + s3 + 2s, + s4
0,I
01
0 0
0 0
I
0 0
H,O
0,
6
2
-6
2
0
-I
CO,
C,H4
C,H40
2 0
-1 -2
values in both expressions. The form in which the multiple reaction (33) is written suggests that it is made up of two specific reactions operating independently, but this separation is not inherent in the chemical system. Any two basis elements for the reaction space could be used to construct an expression identical to (33). What is inherent in the system is the fact that there are four simple reactions in the space of all overall reactions. Their equations are
+ 0, * 2C,H40 + 4H,O * 2C,H40 + 5 0 , + 2H,O * C,H4 + 3 0 ,
2C,H4 4C0,
2C0,
2C0,
(34)
+ 2H,O + 5C,H4 * 6C,H40
having been determined by a procedure which is analogous to the procedure for finding direct mechanisms. Any two of them constitute a basis for the space of overall reactions and could have been used to express the general overall reaction. Notice that the first reaction, used in (331, is not simple. Let us collect terms and rewrite the general mechanism (32) and its overall reaction (33) in a more conventional way:
+ 2 0 ) ~ i+ (6p + 2 0 ) ~ 2+ + (35) ( 7 p + 20)C2H4 + (6p + 0)Oz * (6p + 20)C2H40 + 2pC02 + 2 p H 2 0 (6p
ps3
0 ~ 4
If the ratio p/rr remains constant, the last equation can be expressed with the ratio as a selectivity so that only a single free parameter remains, as recently suggested by Temkin (23). Examination of the matrix given in Table XV brings up an item of special interest. If the combination s4 of atomic oxygen were assumed not to occur, we would still be able to produce ethylene oxide by a combination of the first three steps. This scheme could place a lower limit on the selectivity at 6:7 or 85.7%, corresponding to a simple overall reaction rather than a multiple overall reaction. This serves to illustrate that we get fewer overall reactions than would be predicted by considering only the atom-by-species matrix, as a result of a more restricted choice of possible steps.
303
ANALYSIS OF MECHANISMS FOR REACTION SYSTEMS
EXAMPLE 7. ETHYLENE OXIDE SYNTHFSIS ( 1 CYCLE) Let us consider a more complicated case, similar to Temkin’s proposed system (5) for ethylene oxide formation, but without any prior assumption about the direction of any step in it. The overall reaction space remains the same as in Example 6, but there are additional intermediates. In particular, acetaldehyde (CH3CHO)is an intermediate which is not bound to the catalyst. Its role still requires clarification, as indicated in recent studies by Wachs and Chersick (24, but, whether or not Temkin’s scheme proves to be correct, it illustrates our method. The steps are as follows: sl:
0,+ I* 0,I
*0,I + 1 + C,H, * 01 + CH,CHO
s, : 201
s,:
0,I
s,:
+ I + C,H,O/ + C,H, * C,H,O + 01 + CH,CHO * 5 0 1 + 2C0, + 2H,O C2H,01 * 01 + C,H,
(36)
C,H,O
s s : O,/ s6: SO,/
s7:
Diagonalization of the matrix of stoichiometric coefficients gives Table XVI, from which we read off the general steady-state mechanism (37) and its overall reaction (38), where p , a, and C#J are unrestricted:
+ 3 4 s , + ( p + 30 + + d s 3 + + (2p + + 4(s4 + s,) p ( - 0, - 2C2H4 + 2C2H40)+ a( - 3 0 2 - C2H4+ 2 C 0 2 + 2 H 2 0 )
(p
$ 0 2
(37)
4)Sg
S6)
(38)
Following the method given in Section IV,B for determining those values of which make (37) into a direct mechanism, we arrive at the three direct mechanisms for the overall reaction (38), which are given by the matrix in Table XVII. Temkin (5)gave mechanism m3, but m, and m, would be equally valid choices, if all steps can occur in both directions. In interpreting his
4
TABLE XVI 0,I SI
-
s2
SI
+ s1 + s2
2s, sA
SI
+ s* + 2s,
3SI
+ 3s, + s, + S6
S,+S,+S,+S,
1 0
0 0
01 CH,CHO
C,H,O/
0
0
0 0
0 2 0
0 0 0 I
-2
1
C,H,O
-I
0 0 0
2 0 -1
-I
0
0
0
0
0 0
0 0
0 0
2 0
0
0
0
0
0
0
C,H,
0, C O , H,O
0 0 -2 0
-I
-2
-I -3
-1
0
I
-I 0
0
0 0 0 0
0
0 2
0 2
0
0
0 0 0
304
JOHN HAPPEL A N D PETER H . SELLERS
TABLE XVll SI
m, m2
p p p
m3
s2
+ 3a + 30 + 3a
’3
0
+ 3a + 3a
-p p
+ ‘h
SS
U
p-3a
U
0
U
2P
s4
+Sl
-p-3a
- 2P 0
results, Temkin concluded that step s7 should be included in a rate expression although it does not occur in mechanism m3. From the viewpoint developed in this article, we would conclude that if s, is to be employed, it must occur in the negative direction to accommodate the only possible other direct mechanisms m, and m 2 . Combinations of these three direct mechanisms which are cycle-free are discussed by Sellers (9b). There would be no essential change in the above results if we had diagonalized the matrix differently. If line 6 in the matrix of Table XVI were replaced by twice line 6 minus line 5, the matrix would still be in the appropriate diagonal form, but the general mechanism (37) would appear as (37’) and the overall reaction (38) would appear as (38’):
+ 5 0 ’ ) s , + ( p ’ + 50‘ + + 20’(s, + Sg) + 4”s4 + s7) + (2p’ - 20’ + 4 ’ ) S S p’( -02 - 2C2H4 + 2C,H40) + - 2C2H4O + 4 c o 2 + 4H2O)
(p’
4’)SZ
(37‘) (38’)
a’(-502
The matrix given in Table XVII would take the form given in Table XVIII. Each of the three mechanisms listed in Tables XVII and XVIII corresponds to elimination of the same step. The mechanisms in Table XVIII correspond to 0 = 20’ and p = p‘ - of, so we have simply changed the way of writing the two arbitrary advancement parameters without altering the mechanisms. One other item is worth noting in this example. Since s4 and s7 appear only in the cycle, to omit either one from the choice of possible steps would reduce the system to a unique direct mechanism with a multiplicity of two. But it would not be possible to eliminate further steps and still obtain a reaction among all the terminal species as we were able to do in Example 5. TABLE XVIII SI
m, m2 m3
p’ p’ p’
+ 5a’ + 5a’
+ 5a‘
s2
0 -p‘ p’
+ 7a’ +5d
s.3
+ Sh
s5
s4
+ s7
2a’
p, - 7a’
- p ( - 5ar
2a’
0 2p’ - 20’
2a’ - 2p‘ 0
2a’
ANALYSIS OF MECHANISMS FOR REACTION SYSTEMS
305
Another type of reaction that has received considerable study is that of hydrocarbon isomerizations [see Pines (25)l. These and similar multicomponent systems have been the subject of an elegant kinetic analysis by Wei and Prater (26).In this method pseudomonomolecular reactions following mass action laws are assumed. With these simplifications they treat the general mechanism which includes the occurrence of all possible steps without cycle formation. Data are obtained for concentration changes of reacting components versus time and individual rate constants are calculated by a novel method of integration of the differential equations that model the systems. The Wei and Prater method has been applied to n-butene isomerizations, as well as to several other systems [see Christoffel(27)l.The following example illustrates a different way of considering such systems. In this example models are first generated in which all possible elementary step velocities may not occur.
EXAMPLE8.
ISOMERIZATION OF
BUTENES (1
CYCLE)
Isomerizations (39) among the species 1-butene, trans-2-butene, and cis-2butene are postulated: s, : C4H8-1
+ I*C4H8-II
s*: C4H8-II= C4H8-2CI
s, : C4H8-lI* C4H8-2TI s4 : s5 : S6
:
C,H8-2CI= C4H8-2TI
(39)
* C4H8-2C + I C4H8-2TI * C4H8-2T + I C4H8-2CI
Let the isomers be denoted by 1,2C, and 2T. Then diagonalizing the matrix of stoichiometric coefficients gives matrix of Table XIX. From the matrix of Table XIX we obtain the general steady-state mechanism (40) with p and a unrestricted: P(S,
+ s5) +
+ S 6 ) + ( P + 4)sz + (a - 41% + 4 s 4
(40)
whose overall reaction is as follows : p( -(C4H8 - I)
+ (C4H8- 2C)) + a( -(C4H8 - 1) + (C4H8- 2T))
(41)
By the method of Section IV,B we arrive at three direct mechanisms tabulated in the matrix of Table XX. All six elementary steps assumed possible for any mechanism are shown in Fig. 2,
306
JOHN HAPPEL AND PETER H . SELLERS
TABLE XIX 11 S1
SI
SI
+ s2 + s3
I
2C
2T
I
-I -I -I
0
0
0 0
0 0
-I -I
I
0 1
0
0
0 0 I
0 0
0 0
0 0
I 0
0 I
0
0
0
0
0
S]
+ s2 + ss + s3 + s,
s2
-
+ s4
0
s3
2Tl
0
0 0
S]
2CI
-I
-1
-I 0
TABLE XX
Wei and Prater, in treating this system, assumed in effect that the surface species would be in equilibrium with butenes in the gas phase so that only reactions s2, s 3 , and s4 were employed, corresponding to six first-order rate constants. We have used the more general form shown in Fig. 2. The three direct mechanisms in the matrix of Table XX can be diagrammed as shown in Fig. 3. The arrows in Fig. 3 show the net direction the reaction velocity steps would take for the case where a mixture of reactants is employed such that 1-butene would produce both of the 2-butenes. The steps are not assumed to
C4Hg - 1
1lS1 C4Hg-2CP
54
C4Hg-2TP
551t C4Hg- 2C FIG.
C4Hg - 2 T
2. Possible elementary steps for isomerization of butenes. This diagram corresponds
to listing (39).
ANALYSIS OF MECHANISMS FOR REACTION SYSTEMS
307
R
FIG.3. Direct mechanisms for isomerization of butenes.
be unidirectional. The directions of these two independent reactions can be calculated from thermodynamic data. If experimental data cannot be satisfactorily modeled by one of the three direct mechanisms, a combination of m, and m, will accommodate any possible values for step s4, except when s4 does not occur in either direction, but that case would have already been ruled out since sq is not contained in the direct mechanism m3. If adsorption steps s l , s 5 , and s6 were considered to be at equilibrium, as assumed by Wei and Prater, the situation would be considerably simplified. Since the 2-butenes are connected by a direct reaction, it would be possible from thermodynamics to calculate the direction of the reaction connecting them. Thus, it would be necessary at the outset to model only two direct mechanisms, namely either m, or m2, and m3. If neither of these direct mechanisms were capable of modeling the data, a combination of the two would be sufficient to evaluate all six reaction velocity constants. Such modeling would, of course, be strengthened by supplementing the usual overall reaction rate experiments by tracer data. Recent studies of the kinetics and mechanism of n-butene isomerization over lanthanum oxide by Rosynek et al. (28) indicate that for this catalyst interconversion of the two 2-butene isomers (s4 in Example 8) is very slow and in that case the system could be described by mechanism m3. Studies by Goldwasser and Hall (29) indicate that as temperature is increased, there is appreciable direct conversion via s4 so that one or both of the other two direct mechanisms may be involved. These authors suggest that further studies with all three isomers, at several temperatures and with tracers, would be desirable. The monomolecular conversion of three components has also been considered in some detail by Kallo (30). Rate equations based on Langmuir adsorption were developed assuming a number of different mechanistic schemes including steps in which surface adsorption was not at equilibrium. Since the rate equations developed became complicated, practical application was devoted to cases in which only initial reaction rates were observed, so
308
JOHN HAPPEL AND PETER H. SELLERS
that the final steps involving product formation could be considered unidirectional. The scheme shown in Fig. 2 is one of a number of alternatives considered by Kallo (his scheme VII), and that corresponding to s4 = 0 in Fig. 3 was also identified by him, but not the alternatives with s2 = 0 and s3 = 0. A number of additional alternates could be developed following our procedure, consistent with an appropriate choice of possible elementary steps. Dehydrogenation, hydrogenation, and aromatization of hydrocarbons have also been widely studied [see Pines (25)]and applied industrially. Since olefinic compounds produced by such reactions can simultaneously react to form isomers, it is of interest to explore effects of such combined reactions on the mechanisms involved. Model studies of hydrogenation-dehydrogenation of the n-butenex-butene:hydrogensystem by Happel et al. (31)and of the isobutane:isobutene:hydrogensystem by Happel et al. (32) showed that they are governed by different kinetics. This seems to be due to the occurrence of isomerization reactions in the former case. Studies by Hnatow (33) indicated that in the case of n-butane dehydrogenation 1-butene is the primary product. However, the 2-butenes hydrogenate more rapidly than 1-butene. The following example illustrates how mechanisms can be developed to reflect these observations.
EXAMPLE 9.
BUTANE DEHYDROGENATION (3 CYCLES)
Following a variation of the well-known Horiuti-Polanyi mechanism, we consider the following steps as possible for the system n-butane-n-buteneshydrogen over chromia-alumina catalyst : C4Hl0
+ I*
C4Hl0I
+ I*H,I H21 + I* 2HI C4Hl0I + I* C4H9-II + HI C4HloI + I* C4H9-2/+ HI C4H9-11 + I*C4H,-II + HI C4H9-21 + I = C4H,-2CI + HI C4H9-21+ I * C4H8-2TI + HI H,
C4H8-II* C4H,-2CI C,H,-II*
C4H8-2TI
C,H,-ZTI*
C4H8-2T
C4H8-2CI
* C4H,-2C
C4H,-11* C4H8-2CI
C4H8-I
+I +I
+I
= C,H,-ZTI
TABLE XXI C,H8-2Tl
C4H,-2CI
C,H,-I1
C4H,-21 -1 -1
se
+ s5 + S8 + S I 1 + s5 + + sl, + s4 + Sb + SI3 - s5 + - S) + s9
S)
-
S8
- s,
'7
-
'8
+ '14
s1 - s, - s3 s1 - sz - s3 S6
S)
-- 1
1
0
+ SIO
0 0 0
0 0
0
0 0
0
0
0
0
0 0
0 0
0 0 0
0 0
0
0 0
0 0 0 0 0 s1 - s1 - s3
C4H9-II
0
0
1
H,I 0 0 0 0 0
C4HlOI HI 0 0 0 -1 -1
I
C4H8-2T
1 1
-1 -1
1
-I
1 1
-1 -1
-1
C4H,-2C
C4H,, 0 0 0 0 0
C4H,-I 0 0 0 0 0 0 0 0
H, 0 0
0 0 0 -1
0
1
0
0
0
0
0
I 0
0 0 2
0 0 0
0 0 0
0 0 0
0 0 0
0 0 0
1
0
-1
0
1
0 0
1
-1 -1
0 1
1 1
0 0 0
0 0 0
0 0 0
0 0 0
0 0 0
0 0
0 0 0
0 0 0
0
-1 -2
-1
0
0
0
0 0 0
0 0 0
0 -I
310
JOHN HAPPEL A N D PETER H . SELLERS
Diagonalization of the matrix of stoichiometric coefficients in (42) gives the matrix of Table XXI, from which we read off the steady-state mechanism (43) and its overall reaction (44), which has a multiplicity of three:
+ + s g + + a(s1 - s2 - + + + + T(s1 - s2 + + + + 4(s4 - + - + + s9 + + - f =(P + + T)(S1 - S2 psi1 + as12 + T S 1 3 ( P f a - 4 ) S g + (T 4 x S 4 + (a - 4 -t X +
p ( s , - s2 - s3
- s3
x(s7
s3
s11)
sg
sg
$(s7
slO)
s8 -
s13)
s6
s8
s5
s7
s12)
s5
s6
s7
sg)
s14)
S3)
Sg)
+(P - x
-
Il/ls8
$Is7
+ (4 - d S 9 + X s 1 0
(43)
f $s14
+ a(-C,Hio + (C4H8 - 2C) + + ~ ( - C 4 H l o+ (C4H8 - 1) + H2)
P ( - C ~ H , O+ (C4Hg - 2T) +
H2)
H2)
(44)
Now let us determine all the direct mechanisms for reaction (44) in detail to illustrate the method introduced in Case 4 of Section IV,B. Each direct mechanism for reaction (44) must depend on the parameters p, a, and T , because they appear in the reaction (44), but 4, X, and $ must be evaluated, which is done by means of formula (19) in Section IV,B. Start with the general mechanism (43) and extract from it the cycle the matrix of Table XXII and one mechanism ( 4 9 , written as a row vector: TABLE XXll
4 x
*
0
0 0 p+(T+T
0 0 0 /3
0 0 0
0 0 0
-1
T
p+a
0 0
1
-I
't
0
I
0 0
-I
1 - 1
a
p
1
-1 0
0
0
0
I
0
0
1
0
0
(45) The cycle matrix of Table XXll is a tabulation of mechanism (43) with y = 0, a = 0, and T = 0, and the row vector (51) consists of the coeEients in (43) with 4 = 0, x = 0, and Il/ = 0. Any three independent cycles could have been chosen to generate Table XXII and any mechanism for the overall reaction could have been chosen to establish the row vector (45).The choices we made are arbitrary and depend on the diagonalization procedure used to find the matrix of Table XXI, which is far from unique. The important point is that the list of direct mechanisms we are looking for is unique and independent of how the above choices are made. Starting from the left side of the matrix of Table XXII, choose the first three columns which constitute a nonsingular 3 x 3 matrix, which will be
31 1
ANALYSIS OF MECHANISMS FOR REACTION SYSTEMS
those headed by ss, s7, and s,. In these three columns of Table XXII and row vector ( 4 9 , we find matrix (46) and row vector (47):
I/ = (
p
+ 0,0,0)
(47)
Apply formula (19) to these to get an evaluation for (4,x, $) as follows:
\
0
-1
1/
+ 6,p + 0,-0) (48) Put these values, that is, 4 = p + 0,x = p + 0,and $ = -6, in the general =(
p
mechanism (43), which gives the first direct mechanism. It appears as the first row of the matrix in Table XXIII. The subsequent rows are determined in the same way with an additional procedure to avoid repetitions, which can be explained by first looking at the row m3 in Table XXIII, which is the direct mechanism corresponding to the columns s 5 , s,, s 1 4 .Since s I 4 is the last column, the next three columns to consider are s s , s8, s,. Before testing the 3 x 3 submatrix M of the matrix of Table XXII with those columns for nonsingularity, look for any mechanisms already listed in Table XXIII which have zeros in those columns. If any are found, dismiss that choice of columns. Mechanism m, has zeros in columns ss, s8, s9, wfiich means that this choice of columns would lead to the same mechanism as m , , unless M was singular, in which case we would also want to dismiss this choice. Thus, we avoid the test for nonsingularity in a case such as this. The next two choices would be s 5 , s 8 ,s l 0 and s S r sg, s14,which would be dismissed for the same reason. The next choice is ss, sg, s l 0 ,for which matrix M is singular. Finally, the next choice s 5 , s,, s14is accepted, and it determines the direct mechanism m4. A continuation of this procedure gives the complete set of direct mechanisms through m18. Mechanism m18 corresponds to the case where the isomerization steps s9, s l 0 , and s14are assumed to be very slow compared with steps of hydrogenation and dehydrogenation. In that case, if p and r~ were taken as negative and 5 as positive, we would model a situation in which n-butane was dehydrogenating to produce l-butene and at the same time the 2-butenes were being hydrogenated to produce n-butane. This qualitatively follows the observations of Hnatow (33). This mechanism is of interest because it calls attention to the limitation of a rule of thumb sometimes employed in catalytic research. It has at times
TABLE XXIII (si
-
'2
-
'3)
'11
'12
'13
'5.
(s4
+ '6)
s7
'8
s9
'10
0
P + U
m1 m 2
U
m4
0
m 5
0 -'I
m 8
U
0
m9 m10 m11
-p
m12 ml 3
m14 m15 m16 m17 m18
0
-U
-T
0 - T -T
- 'I
0
-T
0 U
0 0 0
-U
0
P + T
0
-T
0 0 0
P + U
0
m 7
P
P
P + U
m6
-U
0
P + O
m 3
' 1 4
P -'I
0 0 0 0 P
0
P
0 0 0 -U
0 P
0 0
T
ANALYSIS OF MECHANISMS FOR REACTION SYSTEMS
313
been suggested that a useful approach to improve catalyst development may be to study the formation of reactants from products instead of the desired forward reaction [see Bond (34)].This example shows how this idea may need to be modified for mechanisms involving isomer formation. Another point illustrated by Table XXIII is the need to carefully consider the effect of lumping isomers for convenience, when mechanistic models are generated. Thus, in Example 4, isomerization of 1-butene is neglected in selecting the elementary steps for butadiene production. In effect, it is assumed that all intermediates are indistinguishable whether 1-butene or a mixture of n-butenes reacts. If that scheme were used in the present case, we could consider a model in which only sl, s2,s3, s4, s6, and s I 3were retained, which would correspond to only a single direct mechanism. However, if instead we chose to retain all the elementary steps as possibilities except s l , and s12,we would obtain five direct mechanisms for a system producing only 1-butene (in which p = (T = 0). Reactions involving the catalytic hydrogenation of carbon monoxide to produce hydrocarbons and oxygenated products are important for chemical and fuel production from coal. The following example, in which the methanation of synthesis gas is simulated, illustrates a typical system.
EXAMPLE 10. METHANATION OF SYNTHESIS GAS(3
CYCLES)
A 16-step system for methanation over a nickel catalyst is selected that includes many features discussed in the literature :
+ I*CI + 01 + HI* CHI + I CHI + HI* CH,I + I CH,I + HI * CH,I + I CH,I + HI* CH, + 21 OH/ + HI* H,O + 21 co, + I* COJ co + I*COI COI CI
+ 21* 2HI C0,I + HI* CHOOI + I CHOOI + HI* CHOI + OH1 01 + HI* OH1 + I COI + OI* C0,I + I CHOOI + I * OH1 + COI CHOI + HI* CHI + OH1 COI + HI*CHO/+ I H,
(49)
314
JOHN HAPPEL A N D PETER H. SELLERS
It is by no means exhaustive but does exhibit the main possibilities for reaction mechanisms. The idea of a CHO,,, intermediate has been advocated by Pichler and Schultz (35),although other intermediates such as CHOH,,, could have been equally well postulated to take into account the hypothesis that C O dissociation might be assisted by hydrogen. The water gas shift reaction is considered to occur via the CHOO,,, intermediate according to studies of Oki et al. (36).The hydrogenation of carbidic species by atomic rather than molecular hydrogen follows the findings of Happel et a/. (37). By diagonalizing the matrix of stoichiometric coefficients in (49), we get the matrix of Table XXIV, from which we can read off the overall reaction
+ 0(-2CO
-
2Hz
+ CH4 + C 0 2 )
( 50)
which has a multiplicity of two. Then we can apply the methods of Section IV,B to determine all the direct mechanisms for r (shown in Table XXV) from which are omitted the following steps with constant coefficients:
+ 0)(s3 + s4 + +
+ ( p + 20)sS + (3p + 20)s9 (51) The maximum number of steps is obtained ( H + R = 13) for m,, m I 2 , (p
sg)
ps6
- cs7
m I 3 , m14, and m15.If the CHOI is thought to be unlikely [see Yates and Cavanagh (38)],we are left with only m a , m,,, and m, for consideration, so it can be seen that establishment of the presence or absence of such an intermediate is important. Mechanism ma is probably the simplest since it does not require the CHOOl intermediate that appears in the water gas shift reaction. However, if it were assumed that over certain catalysts the reaction of O/ with COI ( ~ 1 3 )could not occur, then mechanism m,, should be considered as a possibility. In this example we only listed the following two overall simple chemical reactions to represent the system: r,:
3H,
r,:
2H,
+ CO * CH, + H,O + 2CO * CH, + CO,
Altogether there are five simple overall reactions in the system, the others being as follows :
+ H,O * CO, + H, (water gas shift) CO, + 4H, * CH, + 2H,O (Sabatier) 4CO + 2H,O + CH, + 3C0, (Kolbel-Engelhardt)
r 3 : CO
r4: rs :
(53)
Any 2 of these 5 equations could be employed to represent the overall reaction with a multiplicity of 2 and the same 15 direct mechanisms would be obtained. In this example we expressed the general overall reaction as
I
l
l
oooooooo--m I
0 0 0 0 0 0 0 - 0 0 0
I
I
1
0 0 0 0 0 0 - 0 0 N N
m w I I I
I
- N
0 -
--
0 0 0
0 0 0
0 0 0
0 0 0
0 0 0
0 0 0 0 - 0 0 0 0 0 0
- 0
0 0 0
I
3 0
I 0 0 0
I 3 0
0 0 0
I
3 0
0 0 0
I
3 0
0 0 0
I
I
0 - - 0 0 0 0 0 0 0 0
I
0 0 - - 0 0 0 0 0 0 0
I
0 0 0 - - 0 0 0 0 0 0
I
- 0 0 0 0 0 0 - 0 0 0
I
oooooooooom
3 0
0 0 0
- - - - N N - - N N m
O O O O O O O O O N O
D O
0 0 0
I
3 0
0 0 0
I
3 0
0 0 0
I
0 - - - - - 0 0 N 0 0
3 0
0 0 0
I
0 0 0 0 0 0 - 0 0 0 0
3 0
0 0 0
I
0 0 0 0 0 - 0 0 0 0 0
3 0
I
- - 0 0 0 0 0 0 0 0 0
m
I
m
n
0 0 0
N - 0 1-
-
3 0
- 0 0 0 0 0 0 0 0 0 0
m
N
c
m
N
mm
' D P
m v I
t +
v)N
m m
::
'1
m
N
m-
m
m
m
d v)
n m N
m
II
m
I
p
s
b c
+
o+
m-
m-
m N
N I m-
lo+ v)-
I
+ " Z
0 1
m-
- + I f + my + + d, + + + 7 A: + - I y + 4
mo" vm)w N
mm
++,
++
t l + -
316
JOHN HAPPEL AND PETER H. SELLERS
TABLE XXV 6 1
+ s2) 0 0 0 0 0 0
m1 m2
%I
0 0
p+a
-a -a -a P + U
0
-a p + a P+a
0 0
a P + O
P
0 0
P + U
p + a
P+
--d
P fP
s12
P
P
tP
m13
tP +a
mi4
/J + O
-0
-a
ml5
p+2a
-a
-a
'1.3
0 0 0 -p
- 2a
0 P
0 0 P+a
0 0 P+a p+2a
'14
s1s
0
P + U
a
-a -a
0
0 0 0
a
m12
810
p
a
- p - a
0 0 0
-a
+ 20 a a a
P + U
0
-U
P+a fP +
0 0
0 - 2a 0 0 0 -P P
-- P
=
0 0 0 0
p+a p+a P + O
p + a P + a
P
0 P
0 0 0 PI2 0 -a
'16
P + O
0 p + a p 0 2a
+
0 P
0 0 0 0 -P
0 a
0
pr, + or,. If we had separately calculated direct mechanisms for r l and r2 instead of considering the combined reaction system, i.e., by leaving out the terminal species CO, for r1 and H,O for r,, we would have obtained 7 direct mechanisms for rl and 10 direct mechanisms for r,. Therefore, if we took m, as 1 of the 7 direct mechanisms for rl and m, as 1 of the 10 direct mechanisms for I-,, then there would be 70 combinations of the form pm, am2. However, if we proceeded to examine the result obtained in this manner, we would find that only 15 of them are direct mechanisms and the other 55 are combinations of 2 of the 15 without any canceled steps. Similarly, the numbers of direct mechanisms for r3, r4, and r5 are 7, 10, and 15, respectively. Obtaining the direct mechanisms for the multiple reaction at the outset is clearly a more effective procedure than combining results for separate individual simple reactions. In fact, without the complete list of possible steps for all reactions, it would not be possible to know whether the complete list of direct mechanisms for each had been derived. The reverse reaction, steam cracking of methane, involves the same elementary steps as the methanation reaction. The kinetics for that reaction have been developed for a single direct mechanism by Snagovskii and Ostrovskii (39). A recent comprehensive review of a set of elementary reactions that can be chosen to represent the Fischer-Tropsch synthesis has been presented by Rofer-De Poorter (40). Such sets of elementary reactions form the starting point for our treatment as discussed in Section 11.
+
ANALYSIS OF MECHANISMS FOR REACTION SYSTEMS
VII.
317
Discussion
This survey has been concerned with the enumeration of all possible mechanisms for a complex chemical reaction system based on the assumption of given elementary reaction steps and species. The procedure presented for such identification has been directly applied to a number of examples in the field of heterogeneous catalysis. Application to other areas is clearly indicated. These would include complex homogeneous reaction systems, many of which are characterized by the presence of intermediates acting as catalysts or free radicals. Enzyme catalysis should also be amenable to this approach. The subject of reaction mechanism also has a bearing on other fundamental problems of physical chemistry. In the following two sections the relationship of the material presented here to thermodynamics and chemical kinetics is considered. A. THERMODYNAMICS In carrying out the procedure for determining mechanisms that is presented here, one obtains a set of independent chemical reactions among the terminal species in addition to the set of reaction mechanisms. This set of reactions furnishes a fundamental basis for determination of the components to be employed in Gibbs’ phase rule, which forms the foundation of thermodynamic equilibrium theory. This is possible because the specification of possible elementary steps to be employed in a system presents a unique a priori resolution of the number of components in the Gibbs sense. The number of comporienfs as defined for application in the phase rule is equal to the number of terminal species minus the number of independent chemical reactions and minus the number of any restrictive conditions, such as material balance or charge neutrality. The number of independent reactions includes only those that actually occur under the conditions in question. For example, Gibbs (41), in citing the case of equilibrium in a vessel containing water and free hydrogen and oxygen, states that we should be obliged to recognize three components in the gaseous part since the reaction H, + +02 H,O is assumed not to occur (i.e.,3 - 0). If, however, a suitable catalyst were present, or if the temperature were high enough for reaction (2),the system would reduce to 3 - 1 = 2 components. If we imposed the condition that all the H, and O2 came from dissociation of H 2 0 , we would have a system of only one component. Thus, Gibbs clearly recognized that it was the reactions that would actually occur that should be employed in thermodynamic calculations. A method for calculating the number of independent reactions discussed
318
JOHN HAPPEL A N D PETER H. SELLERS
by Gibbs has been called Gibbs’ rule of stoichiometry by Christiansen (3). Aris and Mah (42) have discussed this rule at length and stated that it gives an upper bound to the number of independent reactions. This method is based on determination of the number of independent reactions by means of an atom by species matrix, as also developed by Amundson (43).Such a method will give the maximum number of independent reactions, assuming no isomerizations, rather than those required by the phase rule itself. Aris and Mah developed an ingenious scheme whereby experimental observations of rates of change of species consumed and produced can be used to determine an observational matrix which will show how many independent reactions are actually accounting for the observed composition changes. They applied this procedure to an example given by Beek (44) in which ethylene was converted to ethylene oxide as well as to CO, and HzO, so that an atom-species matrix would correspond to at most two independent overall reactions, similar to Examples 6 and 7 in Section VI of this article. It was shown that the data indeed were consistent with two independent overall reactions, but this procedure may be difficult to apply practically given the usual uncertainties in kinetic data. Bjornbom (45)has discussed the relation between reaction mechanism and the stoichiometric behavior of chemical reactions. He pointed out that the set of linearly independent reactions which is obtained from the atom-by species matrix may be larger than the number of independent reactions required to describe an actual physical system because the latter must be related to its reaction mechanism. He gave several examples of complex homogeneous reaction systems including a trial-and-error procedure for hydroperoxide decomposition, based on data obtained by Hiatt and coworkers (46).The system considered was the metal catalyzed decomposition of secondary butyl hydroperoxide in n-pentene solution to give secondary ethyl butyl alcohol, methyl butyl ketone, water and oxygen. The atom-byspecies matrix analysis showed that the maximum number of independent reactions was two, but by forming matrices of the intermediates he showed that the number of independent reactions was actually equal to one. Then by trying linear combinations of the elementary reaction such that the intermediates cancel, he found the single reaction consistent with Hiatt’s experiments: 3C4H,00H
* 2C4H,0H + CH,COC,H, + 0, + H,O
(54)
The procedure described here is consistent with his approach but is simpler to use in more complicated cases. Often thermodynamic calculations for complex systems are made assuming that all chemical changes can take place that are allowed within the framework of the atomic material balances. This approximation may be appropriate at high temperatures but is often not true for catalytic systems.
ANALYSIS OF MECHANISMS FOR REACTION SYSTEMS
319
Smith (47)and Bjornbom (48)have discussed the introduction of restrictions into equilibrium calculations in addition to the elemental abundance constraints. Often only a restricted set of terminal species is chosen, but it would seem logical in choosing such additional restrictions to revert to Gibbs’ original idea of specifying possible reactions as well as possible species. The choice of elementary reactions, rather than of overall reactions with complicated stoichiometry, would be simplified by modern developments in theoretical chemistry and surface science. B. KINETICS Our object has been to enumerate all sets of steps corresponding to possible direct mechanisms. Insight into how to choose the elementary steps themselves can often be obtained from physicochemical principles and experimental surface examination as well as from rate data. This information will also throw light on the most likely mechanisms from among those generated. The magnitudes of concentrations of intermediates and of step velocities appearing in these mechanisms are the parameters in kinetic models that form the next step for further discrimination. A detailed treatment of model building for this purpose is beyond the scope of this article. The subject is briefly discussed here in the context of the methods presented. At the outset it may also be advantageous to consider problems of structural identifiability and distinguishability. A model is not identifiable if in principle it is impossible to determine the desired parameters on the basis of proposed data to be obtained even if there is no experimental uncertainty or inadequacy in computer programming. Even when it is theoretically possible to determine the parameters for a given model, it may also be possible to compute those for a competing model from the same data and the models will then be indistinguishable. Resolution of these problems is often not simple. The subject is discussed in an advanced monograph by Walter (4Y) as well as in papers by Park and Himmelblau (49a)and Walter et al. (4%). Most standard chemical engineering tests on kinetics [see those of Carberry (50), Smith (52), Froment and Bischoff (19), and Hill (52)],omitting such considerations, proceed directly to comprehensive treatment of the subject of parameter estimation in heterogeneous catalysis in terms of rate equations based on LHHW models for simple overall reactions, as discussed earlier. The data used consist of overall reaction velocities obtained under varying conditions of temperature, pressure, and concentrations of reacting species. There seems to be no presentation of a systematic method for initial consideration of the possible mechanisms to be modeled. Details of the methodology for discrimination and parameter estimation among models chosen have been discussed by Bart (53)from a mathematical standpoint.
3 20
JOHN HAPPEL AND PETER H. SELLERS
Information on the steps in a reaction mechanism can be extended significantly by isotopic tracer measurements, especially by transient tracing [see Happel et al. (54,55)].Studies by Temkin and Horiuti previously referenced here have been confined to steady-state isotopic transfer techniques. Modeling with transient isotope data is often more useful since it enables direct determination of concentrations of intermediates as well as elementary step velocities. When kinetic rate equations alone are used for modeling, determination of these parameters is more indirect. The use of tracers in this manner has also been considered by Le Cardinal et al. (56), with special reference to homogeneous systems, and discussed by Happel (57) and Le Cardinal (58).Such an approach parallels the viewpoint of Aris and Mah (42) in which they distinguished between the kinematics and kinetics of overall reactions. Rates of change of species are considered without reference to their correlation in terms of rate equations related to particular physical conditions. To summarize, the type of information that has been presented in this review should be useful in furnishing a logical first step in comprehensive understanding of complex chemical reaction systems. Consideration of a chemical system in terms of unique direct reaction mechanisms required to produce observable rates of change of terminal species has distinct advantages, especially when multiple overall reactions are involved. The required necessary assumptions regarding possible elementary reaction steps are becoming increasingly accessible through modern tools for surface spectroscopy and fundamental theories of chemical kinetics of elementary reaction steps. A number of examples worked out in detail illustrate that the procedure can be rather readily followed even if one does not wish to go into mathematical details. In many cases insights are obtained that are not immediately obvious from more superficial considerations.
VIII. A
(4 C=S-Q
H I I M m m, N n,
P=S- H
List of Symbols
Number of species in chemical system. ith species in chemical system. Dimension of cycle space. Rank of matrix of stoichiometric coefficients of intermediates only. Number of intermediates in chemical system. Active site on a catalyst. C x C submatrix of matrix (oij). Mechanism. ith element in a basis for mechanism space. Dimension of unspecified vector space. ith cycle. Dimension of steady-state mechanism space.
ANALYSIS OF MECHANISMS FOR REACTION SYSTEMS
Q R = Q - H
R R(m) ri = R(si) S S Si
T = A - I V aij Bij
Yij
Pi v..
P U bi
T T..
i’ 4i x
*
32 1
Dimension of reaction space. Dimension of overall reaction space. Linear transformation of mechanism space to reaction space. Reaction produced by mechanism m. ith elementary reaction in chemical system. Dimension of mechanism space; also the number of steps (types of molecular interaction) in chemical system. Step in a mechanism. ith step (type of molecular interaction) in chemical system. Number of terminal spacies in chemical system. Row vector for a steady-state mechanism. Stoichiometric coefficient of species a j in elementary reaction T i . Stoichiometric coefficient of species a j in overall reaction R(mi). Net rate of advancement of step sj in steady-state mechanism mi. Coefficient of mi in a mechanism m. Net rate of advancement of step sj in cycle ni. Rate of advancement of a mechanism. Rate of advancement of a mechanism. Coefficient of step si in a mechanism m. Rate of advancement of a mechanism. Net rate of advancement of step sj in the ith direct mechanism. Rate of advancement of a cycle. Coefficient of cycle ni in a cycle. Rate of advancement of a cycle. Rate of advancement of a cycle. REFERENCEB
1. Happel, J., and Sellers, P. H., Ind. Eng. Ckem. Fundam. 21,67 (1982). 2. Moore, W. J., “Physical Chemistry,” 4th ed., p. 332. Prentice-Hall, Englewood Cliffs, New Jersey, 1972. 3. Christiansen, J. A,, Ado. Catal. 5, 31 1 (1953). 4. Horiuti, J., and Nakamura, T., Adu. Caral. 17, I (1967). 5. Temkin, M . I . , Adu. Catal. 28, 173 ( I 979). 6. Boudart, M . , “Kinetics of Chemical Processes,” Chap. 3. Prentice-Hall, Englewood Cliffs, New Jersey, 1968. 7. Horiuti, J., Ann. N.Y. Acad. Sci. 213, 5 (1973). 8. Milner, P. C., J. Electrochem. SOC.111,228 (1964); Milner, P. C., J. Elec/roclrem. Sor. 111, 1437 (1964). 9. Sellers, P. H., Arch. Ration. Mech. Anal. 44, 23 (1971); Sellers, P. H., Arch. Ration. Mecli. Anal. 44, 376 (1972). 9a. Sellers, P. H.. SIAM J. Appl. Math. (1983) in press. 9h. Sellers, P. H., in “Proceedings of the Symposium on Chemical Applications of Topology and Graph Theory” (R. B. King, ed.). Elsevier, Amsterdam, 1983. 10. Temkin, M. I., In!. Chem. Eng. 11,709 (1971). 11. Temkin, M. I., Ann. N.Y. Acad. Sci. 213, 79 (1973). 12. Sinanoglu, O., J. Am. Cliem. SOC.97,2309 (1975). 13. Lee, L..and Sinanoglu, O., Z. Phys. Chem. 125, 129 (1981). 14. Aris, R., “Elementary Chemical Reactor Analysis,’’ p. 8. Prentice Hall, Englewood Cliffs, New Jersey, 1969. 15. Sellers, P. H., SIAM J. Appl. Math. 15, 637 (1967).
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JOHN HAPPEL AND PETER H. SELLERS
16. Temkin, M. I., I n / . Chrm. Eng. 16, 264 (1976). 17. Miyahara, K . , J. Res. Inst. Catal. Hokkaido Uniu. 17, 219 (1969). 18. Hougen, 0. A,, and Watson, K . M., “Chemical Process Principles,” Vol. 3. Wiley, New
York, 1947. 19. Froment, G . F.. and Bischoff, K. B., “Chemical Reactor Analysis and Design.” Wiley,
New York, 1979. 20. Miyahara, K.. and Yokohama, S., J . Res. Inst. Catal. Hokkaido Uniu. 19, 127 (1972). 21. Patterson, W. R., in “Catalysis and Chemical Processes” (R.Pearce and W. R. Patterson, 22. 23. 24. 25.
26. 27. 28. 29. 30. 31. 32. 33. 34. 35. 36. 37. 38. 39. 40. 41. 42. 43.
44. 45. 46. 47. 48. 49. 49a. 496.
50.
eds.), p. 253. Blackie, London, 1981. Sachtler, W. M. H., Backx,C.,and Van Santen, R. A,, Catul. Rev. Sci. Eny. 23,127(1981). Temkin, M. I . , Kinet. C a d . Engl. Trans. 22,409 (1981). Wachs, I . E., and Chersick, C. C., J. Card. 72, 160 (1981). Pines, H., “The Chemistry of Catalytic Hydrocarbon Conversions.” Academic Press, New York, 1981. Wei, J., and Prater, C. D., Adu. Catal. 13, 206 (1962). Christoffel. E. G . . Catal. Reu. Sci. Eng. 24, 159 (1982). Rosynek, M. P., Fox, J. S., and Jensen, J. L., J . Catal. 71.64 (1981). Goldwasser, J., and Hall, W. K., J . Catal. 71, 53 (1981). Kallo, D., J. Catal. 66, 1 (1 980). Happel, J . , Hnatow, M. A., and Mezaki, R., Adu. Chem. Ser. (97). 92 (1970). Happel, J . , Kamholz, K., Walsh, D., and Strangio, V., I&EC Fundam. 12, 263 (1973). Hnatow, M. A., Ph.D. thesis, New York University, 1970. Bond, G . C., “Heterogeneous Catalysis-Principles and Applications,” p. I I I . Clarendon, Oxford, 1974. Pichler, H., and Schultz, H., Chrm. l n g . Tech. 42, I162 (1970). Oki, S., Happel, J., Hnatow, M. A., and Kaneko, Y., Proc. In!. Congr. Caral., 5th. p. 173 (1972). Happel, J., Fthenakis, V., Suzuki, I . , Yoshida, T., and Ozawa, S . , Proc. Inr. Congr. C u td ,, 7th, Tokyo p. 542 (1981). Yates, J. T., Jr., and Cavanagh. R.R., J . Caral. 74,97 (1982). Snagovskii. Y. S., and Ostrovskii, G . M., “Modeling the Kinetics of Heterogeneous Catalytic Processes” (in Russian). p. 232. Khimia, Moscow, 1976. Rofer-De Poorter, C. K., Chem. Rev. 81, 447 (1981). Gibbs, J. W., “The Scientific Papers of J. Willard Gibbs,” Vol. 1, Thermodynamics, pp, 63. 70, 138. Dover, New York, 1961. Aris. R.,and Mah, R. H. S . , Ind. Eng. Fundam. 2,90 (1963). Amundson, N.R.. “Mathematical Methods in Chemical Engineering,” p. SO. PrenticeHall, Englewood Cliffs, New Jersey, 1966. Beek, J . , A h ) . Chem. Eng. 3,204 (1962). Bjornbom, P. H., AIChEJ. 20,285 (1977). Hiatt, R., Irwin, K. C., and Gould, C. W., J . Org. Chem. 33, 1430 (1968). Smith, W. R.. Ind. Eng. Chem. Fundam. 19, l(1980). Bjornbom, P.. Ind. Eny. Chem. Fundam. 20, 161 (1981). Walter, E.. “Identifiability of State Space Models.” Lecture notes in “Biomathematics.” Vol. 46. Springer-Verlag, Berlin and New York, 1982. Park, S. W., and Himmelblau, D, M., Chem. Eng. J . 25, 163 (1982). Walter, E., Le Courtier, Y., and Happel, J., 1EEE Trans. Autom. Conrrol, in press (1983). Carberry, J. J., “Chemical and Catalytic Reaction Engineering.” McGraw-Hill, New York. 1976.
ANALYSIS OF MECHANISMS FOR REACTION SYSTEMS
323
51. Smith, J. M., “Chemical Engineering Kinetics,” 3rd ed. McGraw-Hill, New York, 1981. 52. Hill, C. G., Jr., “An Introduction to Chemical Engineering Kinetics and Reactor Design.” Wiley, New York, 1977. 53. Bard, G., “Nonlinear Parameter Estimation.” Academic Press, New York, 1974. 54. Happel, J., Suzuki, I., Kokayeff, P., and Fthenakis, V.,J. C u d . 65, 59 (1980). 55. Happel, J., Cheh, H. Y., Otarod, M., Ozawa, S., Severdia, A. J., Yoshida, T., and Fthenakis, V.,J. Curd. 75, 314 (1982). 56. Le Cardinal, G., Walter, E., Bertrand, P., Zoulalian, A,, and Gelas, M., Chem. Eng. Sci. 32, 733 (1977). 57. Happel, J., Chem. Eng. Sci. 33, 1567 (1978). 58. Le Cardinal, G., Chem. Eng. Sci. 33, 1568 (1978).
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ADVANCES I N CATALYSIS. VOLUME 32
Homogeneous Catalytic Hydrogenation of Carbon Monoxide : Ethylene Glycol and Ethanol from Synthesis Gas B . D . DOMBEK Union Carbide Corporation South Charleston. West Virginia 1. Introduction . . . . . . . . . I1. Cobalt Catalysts . . . . . . . . A . Background . . . . . . . . B. Catalytic Activity and Selectivity . . C . Solvents . . . . . . . . . D . Catalyst Stability . . . . . . . E . Mechanism . . . . . . . . 111. Rhodium Catalysts . . . . . . . A . Background . . . . . . . . B. Catalytic Activity and Selectivity . . C . Solvents and Promoters . . . . . D . Catalyst Stability . . . . . . . E . Mechanism . . . . . . . . 1V . Unpromoted and Carboxylic Acid-Promoted Ruthenium Catalysts . . . . . . . A . Background . . . . . . . . B. Catalytic Activity and Selectivity . . C . Solvents . . . . . . . . . D. Catalyst Stability . . . . . . E . Mechanism . . . . . . . . V . Lewis Base-Promoted Ruthenium Catalysts A . Background . . . . . . . . B . Catalytic Activity and Selectivity . . C . Solventsand Promoters . . . . . D. Catalyst Stability . . . . . . . E . Mechanism . . . . . . . . VI . Other Catalysts . . . . . . . . . VII . Conclusions . . . . . . . . . References . . . . . . . . . .
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325 Copyright Q 1983 by Academic Press. Inc . All rights of reproduction in any form reserved . ISBN 0-12-007832-5
326
B. D. DOMBEK
1.
Introduction
Hydrogenation of carbon monoxide by heterogeneous catalysts has been extensively researched and reviewed ( I - - ] @ ,but homogeneous catalysts for this reaction have received less attention (7-11). Hydrocarbon synthesis by heterogeneous catalysts, often broadly referred to as the Fischer-Tropsch synthesis, has been studied in depth as a route to liquid fuels from synthesis gas (H2/CO), which is obtainable from fossil and renewable raw materials. Products of this catalytic process are usually paraffinic and olefinic hydrocarbons, although some heterogeneous catalysts can produce oxygenates such as alcohols, aldehydes, acids, esters, and ketones. Oxygenates are very desirable as chemical intermediates and products, apart from their possible use as fuels or fuel additives. For chemical purposes, however, high selectivity with respect to both chemical functionality and molecular weight distribution must be achieved, a result rarely found in heterogeneously catalyzed synthesis gas conversion. An exception is methanol synthesis, which is conducted commercially with high selectivities over heterogeneous catalysts (9, 12-14). Certain two-carbon oxygenates can also be formed with good selectivities by supported rhodium catalysts (15- 19). Polyhydric alcohols such as ethylene glycol and glycerine, however, are very rarely found as products of heterogeneously catalyzed reactions. Indeed, Storch noted in 1948 that “the synthesis of polyalcohols . * . [from] hydrogen and carbon monoxide remains as one of the prizes of further research” ( 1 ) . Although a 1951 patent (20) claimed the production of hydroxylic compounds including ethylene glycol and glycerol by the use of solid Mn-Cr catalysts under at least 2000 atm of H2/C0, this work does not appear to have been further developed or replicated since. The only other reports known to describe the production of diols by heterogeneous catalysis involve the use of supported palladium catalysts (14, 21). Traces of ethylene glycol and other diols were obtained along with the major products methanol and methyl formate. Thus, heterogeneous catalysts, with these exceptions, are not known to yield polyalcohols from HJCO, whereas homogeneous catalysts, with few exceptions (22-26), produce nearly exclusively monoand polyalcohols. Indeed, when nonoxygenated hydrocarbons are observed as products in supposedly homogeneous systems, the possibility of a heterogeneous component must be considered. In view of this background, the present review is limited to reactions employing homogeneous catalysts for the direct conversion of synthesis gas to oxygenates, and is especially directed toward the synthesis of ethylene glycol. Ethanol is sometimes observed as a major project in such catalytic systems, and these reactions are also considered. Other two-carbon oxy-
ETHYLENE GLYCOL AND ETHANOL FROM H2 AND CO
321
genates, acetaldehyde and acetic acid or its esters, are sometimes observed in these same catalytic systems as minor products; their formation will not be addressed except incidentally. The thermodynamics of CO hydrogenation has been discussed in several publications (3, 4, 8, 27). Reactions forming hydrocarbons from H, and CO are thermodynamically favorable over a wide range of temperatures and pressures; the equilibrium constants decrease with increasing temperature. Methane formation has the most negative free energy change per carbon atom. Alcohols are produced with smaller negative free energy changes than corresponding hydrocarbons; this effect is particularly pronounced for the formation of methanol and polyalcohols, reactions in which water is not produced as a by-product. Lower operating temperatures and higher pressures are beneficial for formation of these oxygenated products. It is evident that for reactions converting CO and H, to methanol and polyhydric alcohols, which are not the most thermodynamically favored products, kinetic control is important in determining product selectivity. Direct conversion of synthesis gas to chemicals is not presently commercialized in any homogeneously catalyzed process. The major impetus behind the consideration of synthesis gas-based processes is the anticipated widening in the price differential between coal and petroleum or natural gas. Synthesis gas from coal is expected to be a very competitive feedstock for the production of chemicals having relatively simple molecular structures, including ethylene glycol (28, 29). The actual timing in the replacement of petroleum-based technology by coal-based technology will be dependent on numerous factors, but most directly on hydrocarbon feedstock prices and capital costs for the coal-based route being considered. Although chemicals can be made from H2/C0 in ways which involve several processing steps, direct routes are ideally the most economical. Selectivities are potentially higher, whereas capital costs and energy consumption may be lower because of the fewer reaction steps required. Technological improvements may be required before a homogeneous synthesis gas conversion process is commercialized, but the impact of such improvements could be very great. For these reasons, the literature reviewed here is approached from a chemical and mechanistic point of view, in terms of what it may signify for future research, rather than entirely from a practical applications standpoint. For the sake of consistency and convenience, rates to products quoted throughout this review are given in terms of turnover frequency, defined here as the number of moles of product formed per gram-atom of metal per hour. This provides units of reciprocal hours, rather than the more commonly accepted reciprocal seconds. For the reactions described here, however, use of the former units affords numbers of a magnitude convenient for ready comprehension and comparison.
328
B. D. DOMBEK
II. Cobalt Catalysts
A.
BACKGROUND
Reactions of carbon monoxide and hydrogen appear to be almost universally catalyzed by cobalt complexes. Indeed, HCo(CO), has been termed “the quintessential catalyst” by Orchin (30),because of its participation in hydroformylation, olefin isomerization, hydrogenation, and carbonylation reactions. It is therefore not surprising that the first known references to the conversion of H, and CO to ethylene glycol involve the use of cobalt catalysts. Several patents issued to DuPont in the early 1950s which claim the use of cobaltous fluoride (31) and other cobalt complexes (32,33) as catalyst precursors for an apparently homogeneous reaction under high H ,/CO pressures. Solvents employed included acetic acid, water, toluene, benzene, heptane, and cyclohexane; pressures and temperatures reported in examples ranged from 1400 to 3000 atm and from 180 to 290”C, respectively. Products listed were ethylene glycol diacetate and glycerol triacetate, with smaller amounts of methyl acetate, ethyl acetate, and n- and i-propyl acetates when acetic acid was employed as the solvent. For reactions conducted in other solvents, the only products specifically reported were ethylene glycol and its mono- and diformate esters, although lower-boiling fractions are mentioned. During this same time period, ethylene glycol was reported as a product in cobalt-catalyzed methanol homologation reactions; methanol, ethanol, and glycol ethers were also found in similar reactions carried out to homologate n-propanol and n-butanol(34). Reaction conditions were 800-1000 atm of HJCO at 225°C. These reports appear to have attracted little attention at the time of their publication, perhaps because the high pressures and low productivities made any potential application or thorough study of such reactions seem difficult or impossible. However, a renewed interest in the use of H,/CO as a feedstock stimulated further investigation (or rediscovery) of this chemistry by several groups. Indeed, in the absence of the more recent investigations, the reliability and interpretations of the earlier work might be open to question.
B. CATALYTIC ACTIVITY AND SELECTIVITY An initial report by Rathke and Feder of Argonne National Laboratory on cobalt-catalyzed hydrogenation of CO (35)has been followed by extensive study of this chemistry by the same workers (36-38). Reaction conditions adopted for these investigations (pressures below 375 atm and temperatures below 230°C) are much less severe than those employed in the earlier work
ETHYLENE GLYCOL AND ETHANOL FROM H2 AND CO
329
noted above, so a systematic study of the catalytic chemistry was more easily undertaken. Products observed in these studies (conversions up to 1 M total products) are alcohols from methanol to pentanol and their formate esters, acetaldehyde and propionaldehyde, ethylene glycol and its formate esters, are carbon dioxide and water. A first-order dependence of total activity on the concentration of HCo(CO), was noted (35, 38). Under conditions of higher pressure (2000 atm), Keim and co-workers (39,40,40u)have observed ethylene glycol, glycerol, and other polyols, consistent with the earlier findings of Gresham (32). Under some conditions, substantial amounts of methyl acetate were also observed (Table I). Individual rates are not given, but the amount of H2/C0 consumed by the reaction in toluene solvent was about five times that in N-methylpyrrolidone, CH3NC(0)CH2CH2CH2 (NMP). The distribution of products in these reactions can change very substantially with reaction time, as illustrated in Fig. 1. The rate of ethylene glycol formation remains quite constant, but rates to other products change markedly as the reaction proceeds, indicating that secondary reactions are taking place. Some of the plausible secondary reactions in this system have been listed by Feder et ul. (37): I
I
(a) alcohol homologation, the conversion of an alcohol to the next higher alcohol through metal alkyl, metal acyl, and aldehyde intermediates (41); TABLE I Products from CO Hydrogenation by Cobalt CataIysisa.b Solvent Product (wt. %)
Toluene
Methanol Methyl formate Methyl acetate Ethanol n-Propanol Ethylene glycol Ethylene glycol monoformate I ,2-Propyleneglycol 1.3-Propyleneglycol Glycerol
31.4 34.1 0. I 0.4 0. I 25.1 2.1 0.6 0.6 2.1
Data from Ref. 39. Reaction conditions: 2000 atm, HJCO Co, 1 hr. N-Methylpyrrolidone.
=
NMP' 14.4 6.4 28.9 5.9 ~
-
I , 230°C 0.2 M
330
B. D. DOMBEK
(b) transesterification, the conversion of various alcohols to their formate esters by reaction with methyl formate; (c) methanolysis, the interception of an intermediate metal acyl by methanol to produce methyl esters such as methyl acetate; (d) alkane formation by hydrogenolysis of a metal alkyl; (e) aldolization, reactions of intermediate aldehydes ; and (f) C 0 2 formation, by hydrolysis of formate esters to give formic acid which decomposes. It appears that methanol, methyl formate, and ethylene glycol are primary reaction products. Upward curvature in the amounts of higher alcohols (Fig. 1) indicates that their formation rates are dependent on the concentration of the next lower alcohol. Indeed, methanol homologation under these conditions was shown to be first-order in methanol concentration (38).Thus, during CO hydrogenation the methanol concentration can reach a steadystate level where its rate of production is balanced by the rate at which it is converted to higher products. Fahey has reported results which show ethanol as the major reaction product (42, 4 3 , with a molar selectivity on the order of 50%. These results are obtained at 31 3 atm and 200% conditions under which the overall rate of CO hydrogenation is slow and secondary reactions such as homologation may be more significant. Similar reactions under higher pressures were found to give a lower proportion of ethanol (42). However, decomposition of the CH3(0CH2CH,),0CH, (tetraglyme) sol-
0
10
20
30
40
YIIO~. M set
FIG.I . Product distribution as a function of reaction time in cobalt-catalyzed CO hydrogenation. (Reprinted from Ref. 38. by courtesy of Marcel Dekker. Inc.) Reaction conditions: 26.5 atm H , , 340 atm CO, 182 C, 1.4-dioxane solvent. Y = fo[HCo(CO),]dr; cobalt concentration changes throughout reaction because of sampling. Average HCo(CO), concentration is 0.05 I M .
ETHYLENE GLYCOL AND ETHANOL FROM H2 AND CO
33 I
vent to the observed products could be significant in these experiments (vide infra). Formate esters of the various alcohols formed are observed as major products in these cobalt-catalyzed reactions, and the mole ratio of formates to alcohols remains constant throughout a reaction. This observation would be consistent with the occurrence of a rapid carbonylation equilibrium process, ROH
+ CO * R0,CH
(1)
a reaction known to be catalyzed by bases (44). Further work by Rathke and Feder, however, indicates that addition of alkyl formate to a catalytic reaction does not lead to a rapid formate/alcohol equilibrium (38). Instead, it appears that methyl formate is produced as a primary product, and transesterification with other alcohols present in the system, as in reaction (2), leads to the various formates observed: CH,O,CH
+ ROH 5 CH,OH + RO,CH
(2)
The equilibrium constant for R = CH2CH, was observed to be about 0.8 at 200°C. It also seems probable that some of these formate esters are produced when intermediate aldehydes, such as acetaldehyde, are hydrogenated : C H , C H O B CH,CH,OH
+ CH,CH,O,CH
(3)
These results are presented here to emphasize the fact that selectivity and rates to various products can be subject to great variation as a result of secondary reactions. Any attempt to determine the fundamental responses of a catalytic system to changes in reaction variables must recognize the potential complications of such secondary reactions. Rathke and Feder have carried out calculations to determine the amounts of primary products actually produced by the cobalt system, assuming that these products are methanol, methyl formate, and ethylene glycol (38). The amounts of these primary products were estimated by the following relationships : [CH,OH], = [CH,OH] -
+ [higher alcohols] + [aldehydes]
[HOCH2CH202CH] - 2[C2H,(O2CH),]
[CH302CH], = [CH,O,CH]
(4)
+ [higher formate esters]
+ [HOCH2CH202CH] + 2[C2H4(02CH)J
(5)
+
[HOCH2CH20H], = [HOCH2CH20H] [HOCH2CH202CH]
+ [C2H4(02CH)2]
(6)
According to Eq. (4), the amount of primary product methanol is related to the amount of methanol observed plus the number of moles of alcohols
B. D. DOMBEK
332
and aldehydes produced from methanol via homologation; concentrations of glycol formates are subtracted because they are assumed to be equivalent to the amount of methanol produced from methyl formate via transesterification: CH,O,CH
+ HOCH,CH,OH
-+
CH,OH
+ HOCH,CH,O,CH
(7)
(Higher alcohol formate esters do not appear here because an equivalent amount of methanol is consumed in the homologation process which produced them.) In reactions using trifluoroethanol as solvent, the observed
0.10
0.00
z! 2
0 t
a
0.06
2
w
u
z 0
u
0.04
0.02
0
0.04
0.00
0.12
0.16
TP. M FIG.2. Calculated amounts of primary reaction products formed relative to total products in cobalt-catalyzed CO hydrogenation. (Reprinted from Ref. 38, by courtesy of Marcel Dekker. Inc.) Reaction is the same as that of Fig. I .
333
ETHYLENE GLYCOL AND ETHANOL FROM H 2 AND CO
amount of trifluoroethyl formate (a secondary product formed by transesterification with methyl formate) is also subtracted. Similar reasoning leads to Eqs. ( 5 ) and (6) for the other primary products. The success of this approach is illustrated by Fig. 2, which displays results for the reaction of Fig. 1. This plot depicts the linear increase in the amounts of primary products (including their derivatives) as a function of total product concentration, which is also linear in Y , the abscissa of Fig. 1. Experimental data treated in this manner thus indicate the selectivity to primary products formed by a reaction, regardless of the extent of secondary reactions, and meaningful interpretation of the effects of reaction variables becomes possible. Under some conditions the secondary reactions can be suppressed, and the concentrations of primary products then increase linearly as the reaction proceeds. For example, Fig. 3 gives the results of a reaction in which the homologation process has been inhibited by added tri-n-butylphosphine. A similar inhibition of the homologation reaction was found in experiments carried out in 2,2,2-trifluoroethanoI solvent (36). This alcohol is itself resistant to homologation under these conditions, but does undergo transesterification with methyl formate. Data listed in Table I1 are those obtained after derivation of the actual rates of primary product formation, using Eqs. (4)-(6). These data are
? i
0.08
50: 2 t
0
4
12
8 YIIO?
16
20
M soc
FIG.3. Product distribution in cobalt-catalyzed CO hydrogenation with added tri-n-butylphosphine. (Reprinted from Ref. 38, by courtesy of Marcel Dekker, Inc.) Reaction conditions: 149 atm H,, 149 atm CO, 182’C, 1,4-dioxane solvent, 3(n-C4H,),P/Co, average HCo(CO), concentration is 0.03 M .
TABLE I1 Rates to Primary Products in Cobalt-Catalyzed CO Hydrogenation ai 182°C"
Rateb (hr-' x 10') Expt. 1 2 3 4 5 6 7 8 9 10 I1 12
Notes C C
C C C
C
c, d
c, e
f f f
f,9
[HCo(CO),l Av. ( M ) 0.12 0.10 0.078 0.089 0.051 0.1 1 0.03 0.084 0.081 0.034 0.052 0.0025
PH(atm)
96.5 I92 30.2 100
26.5 102(D2) 149
Pc0 ( a m )
115 121 113 123 340 111 149
115
115
154 97.1
179 170 108 108
161
161
Data from Ref. 38, by courtesy of Marcel Dekker, Inc.
* Turnover frequency, moles of product per mole of catalyst per hour. 1,4-Dioxane solvent. = 3. ' 8.5 M H,O. 2,2,2-Tri!luoroethano1 solvent. 214°C.
'(n-C,H,),P/Co
107k (atm-' sec-')
2.0 1.7 1.8 2.0 1.9 2.7 0.69 2.4 2.7 2.0 2.0 46
CH,OH 3.48 5.15
1.47 3.46 1.18 4.26 2.66 5.17 6.75 3.11 7.89 239
CH,O,CH
I .74 2.93 0.45 I .37 0.51 1.98 0.25 0 3.90 1.68 1.39 26.6
HOCH,CH,OH 1.75 3.86 0.05 2.38 0.14 3.47 0.81 4.17 4.35 1.54 2.32 0
ETHYLENE GLYCOL AND ETHANOL FROM H2 AND CO
335
consistent with the rate equation d[P]/dt = k[H c O ( c o ) , ] P ~
(8) where [PI represents the molar concentration of total products, and PH is the partial pressure of hydrogen in atmospheres. The temperature dependence of the reaction rate constant has been determined (37) to fit the equation log,& (atm-' sec-') = A - 41,000/2.303RT
(9)
where A is 12.40 in 1,Cdioxane and 13.02 in trifluoroethanol. Thus, the free energy of activation is approximately 41 kcal/mol in both solvents. The indicated first-order rate dependence on [HCo(CO),] is observed in 1,Cdioxane solvent, as shown by the constant value of k over a range of HCo(CO), concentrations. A slightly higher concentration dependence may be observed in 2,2,2-trifluoroethanol (e.g., Expt. 9). It is proposed that such behavior, if significant, could be the result of greater ionic dissociation of HCo(CO), at the lower concentrations in this more polar solvent. A catalyst concentration effect on product selectivity was reported by Keim et al. (39), but rate effects are not reported and possible secondary reactions are not taken into account. Also listed in Table I1 are turnover frequencies to the primary products, in units of moles of product per mole of HCo(CO), per hour, which allow a comparison of the relative activity of the catalyst to the primary products under various reaction conditions. It is evident that the activity of this system is quite low under the conditions recorded. The maximum turnover frequency to the glycol product is below 0.05 hr-', and the highest rate to methanol is slightly above 2 hr-'. When the primary products of CO hydrogenation are estimated according to Eqs. (4)-(6), Rathke and Feder have suggested (37,38)that the following relationships hold :
The ratio of formates to alcohols [relation (lo)] was observed to remain constant throughout a catalytic reaction, as had earlier been reported in the hydrogenation of an aldehyde by HCo(CO), (45).This observation provides further evidence that methyl formate is actually a primary product of the reaction. The formate/alcohol ratio was found to be independent of H, partial pressure, but to increase substantially with CO pressure (nearly
336
B. D. DOMBEK
second-order in P , in 2,2,2-trifluoroethanoI solvent, and somewhat less dependent on CO in 1,Cdioxane). This finding is taken as evidence that C H 3 0 H and HOCHzCHzOH are formed from a single intermediate and CH30,CH is formed from a different intermediate. An increase in temperature was also noted to cause a slight decrease in the formate/alcohol ratio. Increasing temperature appears to have an adverse effect on the ethylene glycol/methanol ratio (39), although little information is available. This ratio is found to increase rapidly with increasing pressure (37, 39). Fahey has reported similar results concerning ethylene glycol/methanol selectivity (43).Experiments over a wide range of pressures (1 26- 1973 atm) are reported to show a first-order dependence of the total product formation rate on overall pressure, in agreement with the results of Rathke and Feder which were observed in a much smaller pressure range. (A first-order rate of gas uptake was also reported, but this result appears fortuitous. As noted above, the product distribution can change markedly with time because of secondary reactions, and the number of moles of gas consumed need not be simply related to the number of moles of organic products formed.) The mole ratio of ethylene glycol-derived products to methanol-derived products was found to change substantially with pressure, as shown in Table 111. (Alkyl formates are apparently regarded as methanol-derived products in this analysis.) Because of the overall first-order dependence of reaction rate on pressure (specifically, on hydrogen partial pressure) in combination with the rather complex selectivity relationships among primary products, it is regarded as quite probable that all of the primary products (and the separate interTABLE 111 Product Selectivity and Rate as a Funcrion of' Pressure for Cobalt-Catalyzed C O Hydrogenation"*h Pressure (atm)
Ethylene glycol/methanol'
Rate (hr-')d
313' 1361 1973
0.10 0.19
6.8 22.3 42.1
0.60
"Data from Ref. 43. (Copyright 1981 American Chemical Society. Adapted with permission.) Reaction conditions: 100 ml tetraglyme solvent, 2 mmol Co, HJCO = I , 230°C. Mole ratio of ethylene glycol-derived products t o methanolderived products. Rate of total product formation. 200°C.
ETHYLENE GLYCOL AND ETHANOL FROM H2 AND CO
337
mediates leading to them) derive from a single common intermediate whose rate of formation is first-order in (hydrogen) pressure. Implications of this assumption will be addressed further in Section II,E. A reaction related to CO hydrogenation to alcohols has recently been reported (46,46a,466,46c);the products are apparently siloxanes. Reactions were carried out in the presence of high concentrations of hydrosilanes, as follows : CO
+ H,
Catalyst R,SiH
.XOCH,CH,OX
+ CH,OX
(12) where R = n - C6HI3,C,H,; X is not specified, but presumably is mainly R,Si. Catalysts reported for this reaction include Co2(CO), as well as complexes of rhodium and ruthenium, and reported conditions range from 272 to 544 atm of H 2 / C 0 at 270°C. Turnover frequencies are reported to be as high as 7 hr-’, with up to 45% molar selectivity (among reported products) to the glycol derivative. This represents “a remarkable acceleration by hydrosilanes of the transition-metal-catalyzed reduction of carbon monoxide to compounds containing methoxy and 1,2-ethanedioxy groups.” The acceleration noted appears to result from carrying out the reactions at much higher temperatures than those commonly used for comparable CO hydrogenation experiments. (The maximum temperature is usually dictated by the stability of the catalyst.) The hydrosilane thus seems to provide a remarkable degree of stabilization to the cobalt catalyst. The reaction rate, however, is reported to decline with time, perhaps suggesting catalyst instability since silane conversions appear to be low. The most significant result of this work is the observation of relatively high selectivity to the two-carbon product at the high temperature employed. It is suggested that this effect may be a result of the hydrosilane functioning as a trap for a reversibly formed intermediate which may be a precursor of the two-carbon product.
C. SOLVENTS The solvent is a sine qua non of a homogeneous catalyst system. Solvent properties are indeed very important in determining the activity, selectivity, and stability of a catalyst. Solvent stability is also essential, if the catalytic system as a whole is to be stable. As described above, several solvents have been employed in studies of cobalt-catalyzed CO reduction. Keim et al. (39) noted a substantial difference in activity and selectivity between catalyst solutions in toluene and N-methylpyrrolidone (Table I). Most of the information in this area again comes from the work of Feder and Rathke (36). Listed in Table IV are their results showing changes in the activity of the cobalt catalyst corresponding to changes in solvent polarity. The rates
338
B. D. DOMBEK
TABLE IV Rates of'Cobalt-Catalyzed CO Hydrogenation in Dij'erent Solvents at 200°C"
Solvent
0)
Pressure (atmy
Rate (hr-' x 10')
Heptane Benzene 1 ,4-Dioxaned 2,2,2-TrifluoroethanoI
I .9 2.3 2.2 26.7
290 290 296 306
zl.8 > 7.9 > 13.3 46.4
" Data from Ref. 36. Room temperature dielectric constant. HJCO = 1. I96 'C.
are seen to increase by a factor of about 20 from the least polar solvent, heptane, to the most polar, trifluoroethanol. In view of the relatively large span of dielectric constants covered by these solvents, an ionic mechanism is concluded to be unlikely. Instead, the more polar solvents are proposed to better stabilize a polar transition state or intermediate. These solvents are all relatively nonbasic compared to N-methylpyrrolidone (NMP), a highly polar solvent [t = 32 (47),pK2" = -0.9 (48)] which Keim er cil. found to give relatively poor activity compared to toluene (39). The low activity in this solvent is perhaps caused by appreciable deprotonation of the relatively strongly acidic HCo(CO), [pK, < 2 (49)], to give the inactive Co(C0)i anion. Changes in solvent were also found to have significant effects on product selectivity, mainly relating to secondary product formation. As mentioned above, the use of trifluoroethanol solvent was found to suppress alcohol homologation. and only primary products were observed (36). This inhibition of the homologation process was attributed to strong hydrogen bonding of the relatively acidic fluoroalcohol to the product alcohol oxygen atoms. Such an interaction would make the alcohol less available for protonation by HCo(CO), , a step believed to be involved in the homologation sequence (41). Addition of a Lewis base, tri-n-butylphosphine, was also found to inhibit alcohol homologation (38).Such behavior could be due to both the competition of the more strongly basic phosphine with alcohol for the HCo(CO), proton, and to the conversion of HCo(CO), to a more weakly acidic phosphine-substituted complex. High-pressure infrared studies have indicated that the complex HCo(CO),P(C,H,), is the predominant equilibrium species when HCo(CO), and tri-n-butylphosphine are heated at 190°C under 80 atm of HJCO (50). The addition of 16% water to dioxane solvent was found to increase the
ETHYLENE GLYCOL AND ETHANOL FROM H2 AND CO
339
rate of reaction slightly (Table 11, Expt. 8), which could be the result of increasing the solvent dielectric constant. The addition of water also has significant effects on product selectivity. Formate esters are apparently hydrolyzed to alcohol and formic acid, which decomposes to CO, and H, . Selectivity to ethylene glycol was also substantially increased by the addition of water. Since hydrolysis of the dioxane solvent in the presence of acidic HCo(CO), could lead to ethylene glycol, similar experiments were carried out with methanesulfonic acid instead of the cobalt hydride. Slow hydrolysis of dioxane was indeed observed, but the major product was diethylene glycol. Examination of cobalt-catalyzed experiments did not reveal this product, so it was concluded that little if any of the observed ethylene glycol was the result of solvent decomposition (36). Similar rates to ethylene glycol have been observed under comparable conditions in a solvent, trifluoroethanol, which cannot produce ethylene glycol upon decomposition. Other studies of cobalt-catalyzed CO reduction must be viewed critically from the standpoint that solvent decomposition could contribute significantly to the observed products. Cobalt carbonyl in (CH,OCH,CH,),O (diglyme), CH30CH2CH20CH3(glyme), and CH,OCH,CH,OH solvents has been reported to catalyze rapid CO hydrogenation at H,/CO pressures below 300 atm and temperatures ~ 2 0 0 ° C(51, 52). High selectivities to ethanol and other two-carbon oxygenates are claimed, and novel mechanisms are presented to account for this activity and selectivity. Although it was noted that solvent decomposition could be partially responsible for the observed results, recent work by others suggests that only a minor fraction of the products originates from CO hydrogenation (53).This study confirmed that high apparent selectivities to ethanol could be realized when diglyme, tetraglyme, or 2-methoxyethanol were used as solvents (Table V). No ethanol was detected, however, after similar experiments in sulfolane, tetrahydrofuran, or dioxane solvents. (Relatively short reaction times are reported, which could limit secondary reactions.) Use of 2-ethoxyethyl ether as solvent gave substantial yields of n-propanol, suggesting that the terminal alkyl group in the ether solvents is being cleaved and homologated. Because of the complexity of the solvent mixture, the source of alcohol products in the experiments was investigated by use of 13C0 in the H,/CO reaction mixture. An experiment in diglyme under the conditions of Table V gave ethanol product of which only 5% was derived totally from CO hydrogenation. [The instability of related solvents in cobalt-catalyzed methanol homologation has also been reported (M).] These results also raise doubts about the source of ethanol in the similar experiments of Fahey described above (e.g., Table 111), which were performed in tetraglyme solvent. Indeed, it is noted that “at very high pressures, the tetraglyme solvent reacted and yielded both higher and lower molecular weight materials” (42). It is not
B. D. DOMBEK
340
TABLE V Product Distribution in Cobalt-Catalyzed C O Hydrogenation in Various Solvents",b Selectivity (mol %) Solvent (MeOCH,CH,),O (EtOCH2CH,),0
CH,
C,H,
CH,OH
CH,CH,OH
21.0 30.0
-
2.6
67.8 28.0
5.2 39.9
68.0 72.1
22.7
-
-
-
__
26.0 54.5
-
5.8
52.6
-
I
OCH,CH,CH,CH~ I ,4-Dioxane Sulfolane' MeOCH,CH,OH" MeOCH,CH,OMe a
9.8
~
CH,CH,CH,OH
~
CH,CHO 2.2
Total products (mmol)
-
40 16
-
-
10
-
-
4
-
11.4
20
-
-
-
Data from Ref. 53. Reaction conditions: 0.1 M Co, 200 atm, HJCO = I , 220°C. 4 hr. I
I I
' S0,CH,CH,CH,CH2, trace conversion noted. Ethylene glycol also detected.
clear how solvent decomposition was established in these experiments or excluded in others.
D. CATALYST STABILITY Relatively little reference has been made to cobalt catalyst stability in the studies cited above. Fahey has reported that reactions in tetraglyme at 200°Cunder 136 atm of 1 : 1 HJCO gave precipitates of cobalt metal ( 4 2 ) . Some experiments which produced elemental cobalt were found to give "nonliquid products," presumably methane and other light hydrocarbons. Experiments under higher pressures showed no evidence of cobalt precipitation. Heterogeneous cobalt catalysts are known to be effective for the Fischer-Tropsch process; supported cobalt metal is a catalyst for the synthesis of methane and other hydrocarbons at 240°C under I atm of HJCO (55). The relatively high fraction of methane (and the presence of ethane) reported in Table V may thus be the result of partial cobalt precipitation, although no comment is made regarding catalyst instability. Investigations of cobalt stability as a function of catalyst concentration, temperature, and CO partial pressure have been carried out in connection with cobalt-catalyzed hydroformylation (56-58). The stability of Co,(CO), in heptane is shown by Fig. 4, which relates to the equilibrium
q,, + 8CO,,I
* CO,(~O),,,,,,,
(13)
ETHYLENE GLYCOL AND ETHANOL FROM H2 AND CO
34 1
r' YYJ
f 2m YI
-K 100 x m
:.
E M LO
20 lemperoture
-
0
FIG.4. Stability of cobalt carbonyl catalyst [Co,(CO), and HCo(CO),] as a function of CO partial pressure and reaction temperature (57,58). (Reproduced with permission of Ernest Benn Ltd. and Springer-Verlag.)
The partial pressure of CO necessary to maintain Co2(CO), in solution rises rapidly with temperature. The decomposition of C o 2 ( C 0 ) , may, however, be kinetically slow in the absence of compounds which could catalyze this conversion (56), and the decomposition process is reported to be autocatalytic (59). Thus, a catalytic reaction was possible in an unstable temperature-pressure region for some time before cobalt metal precipitation became noticeable. Although operation with a metastable catalyst may be possible in short batch experiments, it would be undesirable in a continuous reaction where stability over extended periods is essential. It should be recognized that the stability of cobalt complexes under carbon monoxide can be enhanced by the addition of ligands, as is the case for phosphine-modified cobalt hydroformylation catalysts (57, 58). The stability will also probably depend on properties of the solvent employed. Nevertheless, the plot shown in Fig. 4 appears to be quite useful for assessing long-term cobalt stability under HJCO in the absence of strongly coordinating solvents or ligands. Inspection of the reaction conditions adopted for the experiments reported in Tables I-IV suggests that, in general, they are within the region of cobalt stability. An obvious exception is Expt. 12 of Table 11. Indeed, significant quantities of precipitated cobalt were reported for this experiment (36). Experiment 8 also gave precipitated cobalt, which suggests that the presence of water may lower cobalt stability. Both of these reactions produced, in addition to the normal products, significant amounts of methane and a distribution of straight-chain alkanes. The reaction conditions of Table V also appear to lie in the unstable region where cobalt precipitation might be expected. Indeed, heterogeneous catalysis by precipitated cobalt metal
342
B. D. DOMBEK
may be the cause of the relatively high fraction of methane product observed in these experiments.
E. MECHANISM Much of the information bearing on a catalytic mechanism comes from kinetic observations made on the catalytic process. Any proposed mechanism must then be consistent with the cyclic conversion of the starting catalyst species to the transition state of the rate-determining step. The identity of the stable catalyst species under reaction conditions must therefor be known in order to interpret kinetic results. Although a variety of cobalt complexes may be used as catalyst precursors in the reactions described above, most of these complexes are found to be converted largely to HCo(CO), under reaction conditions. King has found that Co,(CO), , (C,H,),C,Co,(CO), ,CH,CCo,(CO), , and (C,H,P),Co,(CO),, are transformed into HCo(CO), in n-tetradecane solution under 200 atm of 1 : I H,/CO at 110-190°C (60). The same precursors were found to provide active catalytic systems when studied in preparative reactions (dioxane solvent, 173-200"C, 200 atm of I : 1 H,/CO). Two other cobalt complexes, (CH,),SnCo(CO), and [CH3N(PF2),Co2(C0),, were found not to be converted to HCo(CO), ,and did not catalyze CO hydrogenation. Other studies by high-pressure infrared spectroscopy under conditions similar to those used for catalytic CO hydrogenation have been reported by Whyman (50). At 150°C and 290 atm of 1 : 1 H,/CO in heptane solution the equilibrium 2HCo(CO),
* H, + CO,(CO),
(14)
was found to lie well toward the hydride complex, although small absorptions due to the carbonyl dimer could be detected. Rathke and Feder have employed Co2(CO),as the catalyst precursor in their studies. Samples withdrawn from reactions under pressure were analyzed for both total cobalt and for HCo(CO), (35);conversion to HCo(CO), was observed to the extent of 50-90%, varying according to (14) with temperature and hydrogen pressure. Experiments with different levels of catalyst showed that the overall rate of CO reduction was first-order in the HCo(CO), concentration, as determined by titration of reaction samples. Thus, there is substantial evidence that the catalyst in this system (or more precisely, the species present in the transition state of the rate-determining catalytic step) is a mononuclear cobalt complex. The observed kinetic dependences [Eq.
ETHYLENE GLYCOL A N D ETHANOL FROM H2 A N D CO
343
(8)] are useful in consideration of how HCo(CO), is transformed on the pathway to the transition state. Fachinetti el nl. have proposed that a trinuclear hydroxymethylidyne cobalt cluster could possibly be an intermediate in CO hydrogenation (6163). Evidence for the equilibrium
has been presented, including isolation of the amine adduct of the cluster from a reaction between HCo(CO), , Co,(CO), , and triethylamine (62). This trinuclear cluster is very thermally labile, decomposing to HCo(CO), and Co,(CO),. Model studies were carried out with a related derivative, Co,(CO),CO-X (X = CH3), which was found to react with H2/C0 (1 15 atm, 120°C) in the presence of Co,(CO),, yielding the organic products X-OCH,, X-OCH2CH20H, and X-OCH2CH202CH (61).This is presented as evidence that the trinuclear hydroxymethylidyne cluster could be a precursor of methanol and two-carbon alcohols in CO hydrogenation. It is difficult, however, to reconcile the observed kinetic dependences of the cobalt-catalyzed process with a mechanism which involves conversion of HCo(CO), to the trinuclear cluster. If the rate-determining step is proposed to occur after cluster formation, a dependence on HCo(CO), concentration of greater than first-order would be expected, in disagreement with observed results. On the other hand, plausible rate-determining steps before or during cluster formation also do not lead to the observed dependences. For example, the forward step in (14) has been shown to have an inverse kinetic dependence on CO partial pressure and a second-order dependence on HCo(CO), concentration (64, 65). Involvement of (14) as a preequilibrium step in cluster formation would contribute an inverse hydrogen partial pressure dependence as well as a higher cobalt concentration dependence. For these reasons, it appears probable that under conditions employed for the research described here, cobalt cluster formation is not involved in the major catalytic pathway. The HCo(CO), complex is therefore inferred to be involved in initial hydrogen transfer to carbon monoxide. This step was initially proposed to comprise rate-determining hydrogen atom transfer from HCo(CO), to free CO, affording a formyl radical, HCO; subsequent reaction with further HCo(CO), would lead to the observed products (35). However, kinetic observations (the zero-order dependence on CO partial pressure) were later made which are inconsistent with such a process (36). All experimental observations are consistent with the first step in the CO hydrogenation process being a rapid, reversible equilibrium in which the
344
B. D. DOMBEK
hydrogen atom in HCo(CO), migrates to a carbonyl ligand :
In the subsequent, rate-determining step
hydrogen is added to the coordinatively unsaturated formyl complex, and the formyl ligand is converted to coordinated formaldehyde. The dihydride species may be a transient intermediate or may be representative of the transition state in the rate-determining step. A mechanism involving rapid, equilibrium formation of H,CO followed by rate-determining addition to HCo(CO), would also be consistent with the observed kinetics. This pathway is regarded as less probable, however, because added formaldehyde is not rapidly converted back to H2 and CO as would be expected if such an equilibrium were important ( 3 6 ) . The conversion of a metal hydride into a metal formyl by hydride migration, as in (16), has been regarded for some time as a very difficult or improbable reaction. Certainly, there is evidence to suggest that for a number of metal carbonyl complexes alkyl migration to a carbonyl ligand, as follows: R
I
L,,M-CO
0
II
+ L’ * L,L’M-CR
is more easily accomplished and observed than analogous hydride migration (66-69). Nevertheless, there are recent ceports which show that hydride-toformyl conversion is an observable process in certain stoichiometric systems ( 7 0 7 2 ) .This step is therefore plausible in catalytic systems, although it is probably an endothermic process with a very small equilibrium constant. Participation of the hydride-formyl equilibrium in (16) is also plausible in light of an apparent inverse kinetic deuterium isotope effect for the catalytic process. Use of deuterium gas instead of hydrogen (cf. Expts. 6 and 4 in Table 11) causes an increased rate, with k,/k, = 0.73 (37).The existence of an isotope effect implies that hydrogen atom transfer occurs before or during the rate-determining step, and an inverse kinetic isotope effect may be possible in the case of a highly endothermic, product-like transition state (73). On the other hand, Bell has concluded that inverse kinetic isotope
ETHYLENE GLYCOL A N D ETHANOL FROM H2 A N D CO
345
effects are not observed for single-stage hydrogen transfer reactions (74). The observation of such an inverse rate effect can instead imply that the observed process includes an endothermic equilibrium which possesses an inverse equilibrium isotope effect. An equilibrium will exhibit such inverse behavior ( K H / K D< 1) if the hydrogen vibrational frequencies are higher on the product side than on the reactant side. This may be seen to be the case for (16), where (Co-H) z 1830 cm-' and (C-H) z 2900 cm-'. If any normal equilibrium or kinetic isotope effects in other steps before or during the rate-determining step are not large enough to negate this inverse effect, an apparent inverse kinetic isotope effect will result. This appears to be a plausible situation for the combination of (16) and (17). Fahey presents the products of (1 7) as uncomplexed formaldehyde and HCo(CO), rather than a bound-formaldehyde species (43). Free formaldehyde is a thermodynamically unfavorable product from H2 and CO (a),and significant stabilization may be expected as the result of coordination in a metal complex. However, thermodynamic calculations are presented which indicate that small equilibrium concentrations of formaldehyde could be present under the conditions of these cobalt-catalyzed reactions (43). Although small amounts of uncoordinated formaldehyde are indeed expected as a result of the following endothermic ( 3 6 , 3 7 ) equilibrium: H I 0 (CO)&o- II C H
/ \
H
I 0 * (CO),CO + II C H
H
/ \
H
the significance of this equilibrium is minor if it is the complexed formaldehyde which is consumed in subsequent steps. The concept of a (bound) formaldehyde intermediate in CO hydrogenation is supported by the work of Feder and Rathke (36) and Fahey (43). Experiments under H,/CO pressure at 182-220°C showed that paraformaldehyde and trioxane (which depolymerize to formaldehyde at reaction temperatures) are converted by the cobalt catalyst to the same products as those formed from H,/CO alone. The rate of product formation is faster than in comparable H,/CO-only experiments, and product distributions are different, apparently because secondary reactions are now less competitive. However, Rathke and Feder note that the formate/alcohol ratio is similar to that found in H2/CO-only reactions (36). Roth and Orchin have reported that monomeric formaldehyde reacts with HCo(CO), under 1 atm of CO at 0°C to form glycolaldehyde, an ethylene glycol precursor (75). The postulated steps in this process are shown in (19)-(21), in which complexes not observed but
346
B. D. DOMBEK
presumed as intermediates are shown in brackets : HCo(CO),
+ HCHO
4
[(CO),Co-CHZOH]
(19)
0 [(CO),Co-CH,OH]
+ CO
II
4
0
(20)
0
I1
[(C0)4C~-CCH,0H] + HCo(CO),
[(CO),Co-CCH,OH]
II
4
HCCHzOH
+ Co,(CO),
(21)
Other work has shown that cobalt complexes under HJCO pressure and at higher temperatures can catalytically convert formaldehyde to glycolaldehyde (76) or ethylene glycol (77-79); methanol is observed as a product as well. It has therefore been well demonstrated that formaldehyde can be converted by cobalt catalysts to the same products observed from CO reduct ion. A very simplified possible scheme for subsequent reactions of a formaldehyde intermediate, modeled largely after one presented by Rathke and Feder (38),is the following: 0
0
Hydrogenolyses leading to products or their organic precursors are shown as reactions with [HI, which may be either H2 or HCo(CO), . Recent studies by Pino and co-workers (80) suggest that in the related cobalt-catalyzed hydroformylation process, it is H, which is largely responsible for this hydrogenolysis. It has also been shown that H2 can undergo analogous oxidative addition to (CO),CoH, forming H,Co(CO), at low temperatures (81). The essentially irreversible hydrogenolysis reactions are presumed to occur for the coordinatively unsaturated (tricarbonyl) intermediates shown.
ETHYLENE GLYCOL AND ETHANOL FROM H2 AND CO
347
The "CO insertion" reactions shown are likely to involve the coordinatively saturated (tetracarbonyl) complexes, and are expected to be reversible. The coordinatively unsaturated tricarbonyl and saturated tetracarbonyl complexes are presumed to be in equilibrium, the position lying predominantly toward the coordinatively saturated species. Glycolaldehyde is postulated as the ethylene glycol presursor, as will be further described below. This scheme appears to be consistent with the selectivity relationships (10) and ( 1 I ) reported by Feder er a/. (37). Changes in H, or CO partial pressure are not expected to significantly affect the first branching. Increases in CO partial pressure are anticipated to favor CO insertion reactions (22) and (24) by increasing the concentration of coordinatively saturated, tetracarbonyl, reactants. This will lead to the observed increases in selectivity to glycolaldehyde (ethylene glycol) and methyl formate. The effects of increased H, partial pressure are less easily rationalized. Increased hydrogen pressure could perhaps increase the fraction of glycolaldehyde relative to methanol, consistent with ( 1 I), by a concerted H, addition-alkyl migration process involving the tetracarbonyl hydroxymethyl complex. Coordinating solvents have been shown to be involved in alkyl migration processes (82), and it appears possible that H2 could behave similarly. Rathke and Feder make the assumption that the pathway to methanol of (25) is minor relative to (23), based on the observation of a relatively constant formate/alcohol ratio at different hydrogen partial pressures, (10). However, this does not seem to be a necessary conclusion. Cobalt-catalyzed hydroformylation of olefins proceeds, under similar reaction conditions (59), through a cobalt alkyl complex analogous to the hydroxymethyl complex shown in (22) and (23). In the hydroformy lation process, little hydrocarbon, a hydrogenolysis product analogous to the methanol produced by (23), is normally produced. Indeed, acyl complexes of the formula 0
II
(CO),CoCR
are observed as a major component of the hydroformylation system under catalytic conditions (50). This indicates that alkyl migration (22) competes very effectively with hydrogenolysis (23) in hydroformylation, and a similar result should probably be expected in CO hydrogenation. All of the observed effects appear to be consistent with a mechanism as shown in (22)-(25) in which substantial amounts of methanol are formed via (25), if the first branching is indeed reversible and if the CO insertion process in (24) is slower than that in (22)-a reasonable assumption (67,83).However, in the absence of additional data relevant to (22)-(25), further rationalizations of product selectivities do not appear to be warranted. The fate of a glycolaldehyde intermediate in these reactions may be somewhat analogous to that of formaldehyde, as outlined by Fahey (43)
348
B. D. DOMBEK
Interaction of glycolaldehyde with the catalyst can thus lead to glyceraldehyde (a glycerol precursor), ethylene glycol, or ethylene glycol monoformate. Pathway (26) represents an aldehyde hydroformylation, or chain-growth route, and provides a plausible pathway to the glycerol which is observed as a product at higher pressures (53).Higher polyols may also be expected as products by continuation of this chain-growth process through a sequence of aldehyde intermediates. However, the selectivity of branching for higher aldehydes will presumably differ from that in the formaldehyde case because of steric and electronic effects. As a result of these influences, the first branching in (26)-(29) should lead more selectively to the metal-oxygen bonded intermediate than was the case for formaldehyde. This will cause a rapid diminution in chain growth beyond the two-carbon polyol stage. An increased proportion of glycerol and possibly other polyols at higher pressures (32) can be rationalized, since increases in both CO and H, pressure are expected to enhance the proportion of CO insertion (26) relative to hydrogenolysis (27). Semiempirical molecular orbital calculations based on modified extended Hiickel theory have been used to model the geometries and energies of the transition-state species of (16), (17), and (22)-(25) (37). Within the uncertainty limits, these calculations are supportive of the proposed mechanism. A formaldehyde complex of the formula (CO),CoH(CH,O) was found to have the cobalt atom closer to the formaldehyde carbon atom (2.34 A) than the oxygen atom (2.85 A). This agrees well with data reported for a formaldehyde complex of osmium, [(C,H,),P],(CO),Os(CH,O) (84).A slightly less
ETHYLENE GLYCOL A N D ETHANOL FROM H2 A N D CO
349
stable structure was found with the oxygen atom closer to the cobalt center. This, it is suggested, may indicate that hydrogen atom transfer in the next reaction step can proceed either to the carbon or the oxygen atom. A hydroxycarbene isomer, (CO),CoH(=CHOH), was found to be significantly less stable. The possible intermediates (CO),Co(CH,OH) and (CO),Co(OCH,), formed by hydrogen atom transfer in (CO),CoH(CH,O), were also briefly examined ; it was concluded that highly accurate calculations would be required to gain information about factors possibly responsible for the branching to thsese two species. Further calculations have been done which suggest that a hydroxycarbyne complex, (CO),CoECOH, may be more stable than the (CO),CoCHO formyl complex. These results have led Nicholas to propose that this species is a possible intermediate in the CO reduction process (84a).Comparison of the relative stabilities of the above two complexes is not altogether valid, however, since the hydroxycarbyne species is coordinatively saturated whereas the formyl is coordinatively unsaturated. In order for the former complex to react with H z , as proposed (84a), loss of a carbonyl ligand to give a higher-energy species would presumably be required (which could also lead to kinetic dependences at variance with those observed). Reaction with hydrogen and CO might then be expected to involve transient formation of the hydroxycarbene intermediate (CO),CoH(=CHOH); as noted above, this is calculated to be a relatively unstable complex in comparison with the formaldehyde species derived from a formyl intermediate. The intermediacy of a hydroxycarbyne species therefore does not appear to be energetically feasible for CO reduction processes. Selectivity observations described here for cobalt catalysts and in Section IV for ruthenium catalysts are also not explained by invoking a hydroxycarbyne intermediate.
111.
Rhodium Catalysts
A. BACKGROUND Knowledge of patents claiming cobalt catalysts for the conversion of HJCO mixtures to ethylene glycol (31-33) appears to have led to initial investigation of rhodium catalysts for this reaction at Union Carbide (27, 85-87). Early experiments by Pruett and Walker at pressures of about 3000 atm indicated that the activity of rhodium was notably greater than that found for cobalt. Several other potential catalyst precursors, including compounds of Sn, Ru, Pd, Pt, Cu, Cr, Mn, Ir, and Pb, were screened for activity under pressures of about 1500 atm and found not to produce detectable
350
B. D. DOMBEK
amounts of polyhydric alcohols. Solvents initially examined with the rhodium catalyst included tetrahydrofuran, water and alcohols, NMP, dioxane, o-dimethoxybenzene, tetraglyme, and toluene. Experiments with added methanol or in methanol solvent indicated that this alcohol is not an intermediate in glycol formation (27). Various organic “ligands” were found to promote the glycol-forming reaction; among those reported were pyrocatechol, bipyridyl, piperazine, and various substituted pyridines, especially 2-hydroxypyridine (85). With the realization that anionic rhodium complexes are present in catalyst solutions, cationic counterions such as alkali and alkaline-earth metals were found beneficial in the process (86). Experiments first reported in the patent literature were carried out largely at pressures x3000 atm, although some at pressures as low as 400 atm were also described. After these initial findings, a considerable amount of progress has been made in improving the activity and stability of this catalytic system, and increasing our understanding of it.
B. CATALYTIC ACTIVITY AND SELECTIVITY Products formed by the rhodium catalyst are generally the same as those produced by the cobalt system. Table VI (Expt. A) presents a summary of products reported by Fahey (43) from a rhodium-catalyzed reaction at 1973 atm. Ethylene glycol and methanol are the two major products, comprising more than 60% of the total product mixture. The overall rate of product formation in this example (moles of products per gram-atom of metal per hour) is 473 hr-’, and the rate to ethylene glycol is 222 hr-’. Some of the compounds found to be secondary products in the cobalt system are also observed in this reaction ; ethanol and higher linear alcohols (and their formate esters) are presumed likely to be secondary products in this reaction also. 1,3-Dioxolane and 2-(hydroxymethyl)-l,3-dioxolaneare the ethylene glycol acetals of formaldehyde and glycolaldehyde, respectively. Their observation is evidence that these aldehydes may be intermediates in the CO reduction process (43). Examples from the patent literature, as shown in Expts. B and C of Table VI, are generally found to contain possible secondary products in much lower proportions. These examples illustrate selectivities to the two major primary products, methanol and ethylene glycol, of greater than 90%. The higher proportion of probable secondary products in Expt. A would appear to be a result of carrying out this reaction to very high product concentrations ; the volume of liquid products formed during the experiment is approximately equal to the initial volume of solvent. Results of experi-
ETHYLENE GLYCOL AND ETHANOL FROM H2 AND CO
35 1
TABLE VI Products from Rhodium-Catalyzed CO Hydrogenation
Amt. (mmol) Product
Expt,: A0.b
Methanol Methyl formate Ethanol Ethyl formate 1 -Propano1 I-Propyl formate 1-Butanol Ethylene glycol Ethylene glycol monoformate Propylene glycol Glycerol 1,3-Dioxolane 2 4 Hydroxymethy1)- 1.3-dioxolane
284 52 191 20 121 29 < 31 1000 118
118 120 2 20
B
d
Ce.I
281 25 6.3
52 2.5
329 14
69 0.9
22
0.5
1.0
Data from Ref. 43. (Copyright 1981 American Chemical Society. Adapted with permission.) Reaction conditions: 100 ml tetraglyme solvent, I mmol Rh, 5 mmol 2-hydroxypyridine, 1973 atm, HJCO = I , 230°C. 4.5 hr. Data from Ref. 86. Reaction conditions: 76 ml tetraglyme solvent, 3 mmol Rh, 12 mmol2-hydroxypyridine, I333 atm, H J C O = 1.5, 220°C 4 hr. ‘ Data from Ref. 88. Reaction conditions: 75 ml tetraglyme solvent, 3 mmol Rh, I0 mmol 2-hydroxypyridine, 0.5 mmol cesium formate, 544 atm, HJCO = I , 220°C. 4 hr.
ments run to more moderate conversions, such as Expts. B and C of Table VI, are therefore believed to be representative of largely primary reactions. Although secondary reactions which increase with product concentrations in these rhodium-catalyzed reactions may not complicate rate interpretations as severely as in the cobalt systems, other complexities can arise as product amounts increase. Results presented in a patent (89) indicate that increased concentrations of ethylene glycol in the reaction medium have an adverse effect on the ethylene glycol/methanol ratio produced. High concentrations of ethylene glycol cause a decrease in the rate of its formation and may enhance the rate of methanol formation. In another patent, apparently referring to this effect, it is reported that ethylene glycol is “destroyed” during the catalytic process (90).Other studies in which I4C-labeled ethylene
352
B. D. DOMBEK
glycol was added to experiments showed that the label was recovered essentially quantitatively as ethylene glycol, indicating that there was no destruction or transformation of this product under the particular conditions employed (91). The buildup of glycerol is also reported to have an adverse effect on the rate of ethylene glycol formation (90), which suggests that the inhibitory effect noted for ethylene glycol is common to polyalcohols. Addition of methanol at similar levels does not inhibit the glycol-forming process (91). Ethylene glycol concentrations of less than 5 M (about 30 wt. %) are claimed to be desirable in order to minimize its inhibitory effects (89).Since most of the experiments reported here do not involve such high product concentrations, major changes in selectivity during these reactions are not expected. Nevertheless, the existence of this effect must be recognized. The reaction rates in this system are presumably first-order in catalyst concentration, as implied by the scaling of product formation rates proportionately to rhodium concentration (90, 92, 93). Responses to several other reaction variables may be found in both the open and patent literature. Fahey has reported studies of catalyst activity at several pressures in tetraglyme solvent with 2-hydroxypyridine promoter at 230°C (43). He finds that the rate to total products is proportional to the pressure taken to the 3.3 power. A large pressure dependence is also evident in the results shown in Table VII. Analysis of these results indicates that the rate of ethylene glycol formation is greater than third-order in pressure (exponents of 3.23 . 9 , and that for methanol formation somewhat less (exponents of 2.3-2.8). The pressure dependence of the total product formation rate is close to third-order. A possible complicating factor in the above comparisons is the increased loss of soluble rhodium species in the lower-pressure experiments, as seen in Table VTI. Experiments similar to those of Fahey have also been
eflecr nJ' Pressure on
TABLE VII Rhodium-Catalyzed CO Hydrogenurion",h
Solvent
Pressure (atm)
Ethylene glycol rate (hr-')
Methanol rate (hr-')
Rhodium recovery rAy
Sulfolane Sulfolane SITd
544 1020 544 I020
9.5 72.5 10.25 92.5
13.5 71.5 15.5 67.5
81 I 07 67 91
SITd
Data from Ref. 94. Reaction conditions: 75 ml solvent, 3 mmol Rh, 7 mmol ethylenedimorpholine, 260°C. ' Amount of rhodium soluble or suspended in solution after reaction, based on amount charged; determined by atomic absorption spectroscopy. 1 : 1 volume ratio of sulfolane and tetraglyme.
ETHYLENE GLYCOL AND ETHANOL FROM H, AND CO
353
noted to give considerably less than quantitative recovery of soluble rhodium at the lower pressure (94, 95). Such catalyst instability effects could lead to abnormally high apparent pressure dependences. Another study which provides information on the pressure dependence is shown in Fig. 5. This series of experiments, which gave rhodium recoveries within the range of 78-930/, (96),exhibits somewhat lower pressure dependences. The total rate of product formation is proportional to the pressure taken to the 2.3 power, and the reactions which produce methanol and ethylene glycol have very similar pressure dependences in this study. The ratio of formates to alcohols, however, was seen to increase with pressure. Rate dependences on CO and H, partial pressures have not been determined, but the experiment in Fig. 5 at a different HJCO ratio (0.67 instead of 1) suggests that the dependence on hydrogen pressure is similar to, or perhaps slightly greater than, the dependence on CO pressure. Results obtained by other workers (97, 98) lead to the same conclusion.
Pressure, atm
FIG.5. Plot of log(rates) vs. log(pressure) for rhodium-catalyzed CO hydrogenation. Reaction conditions: 75 ml sulfolane, 3 mmol Rh, 1.25 mmol pyridine, HJCO = I , 240 C, 4 hr (96). Total rate includes rates to methanol, methyl formate, ethanol, ethylene glycol monomethanol;).( ethylene glycol. Open figures are formate, and propylene glycol: (A)total;).( for an experiment with HJCO = 0.67.
354
B. D. DOMBEK
The effect of increasing the reaction temperature under otherwise constant conditions may be seen in Fig. 6 . Changes in temperature have very little effect on the ethylene glycol/methanol ratio. The routes to both products have Arrhenius activation energies of approximately I8 kcal/mol, as estimated from Fig. 6 . A possible deviation from this dependence at the highest temperature in this series (25OOC) may suggest partial catalyst instability, but rhodium recoveries were not reported. The effect of increased temperature on catalyst stability (based on analyses of rhodium in solution at the end ofcatalytic experiments) is shown in Table VIII ; rhodium recovery is significantly lower for the higher-temperature experiments. Before proceeding to a more detailed description of the effects of various solvents and promoters on catalyst activity and stability, it should be noted that the responses described above are possibly, or even probably, influenced by solvents and promoters. The responses shown, however, appear to be generally characteristic of these rhodium-containing systems. It is apparent that the rate of product formation is significantly accelerated by increases in reaction temperature. Higher temperatures, however, can bring about catalyst instability unless the pressure is simultaneously increased. Higher Temp., "C
10.0
I
2
8.0
6.0
0)
2
40
2.0
i.01 . 1.90
I 2.00
T - ' , K-' x
2.10
lo3
FIG. 6. Effect of temperature on rhodium-catalyzed CO hydrogenation: (u)methanol; ( 0 )ethylene glycol. Reaction conditions: 75 ml y-butyrolactone solvent, 3 mmol Rh, 10.5 mmol 2-hydroxypyridine, 0.5 mmol cesium 2-pyridinolate, 544 atm, H,/CO = I , 4 hr (88).
ETHYLENE GLYCOL AND ETHANOL FROM H2 AND CO
355
TABLE V l l l Acriuirv und Rhodium Recowry in Tetraglyme Soluent at 544 arm and 220-240 C" ~~
Temp. ( C) 220' 240h 220h
240' 220' 240'
Ethylene glycol rate ( h r - ' )
Methanol rate ( h r - ' )
Rh recovered
Salt Cs-2HP Cs-2HF PPN-OAcd PPN-OAcd BTB' BTB'
5.51 6.59 4.97 7.12 5.13 6.90
4.43 9.64 3.39 7.55 3.98 6.66
81 64 81 61 83 46
r 5 , generally provide an optimum rate to glycol at a relatively low amine/rhodium ratio, on the order of 0.1-0.5. For amines with a pK < 5 , a higher amine/rhodium ratio is required to achieve the highest ethylene TABLE X Amounts of Amines Required for Optimum Promoter Eflect” Amine
pK (H,O, 25°C)
1,8-Bis(dimethylamino)naphthalene Sparteine Dibutylamine Triethylamine N-Methylpiperidine Piperazine Ammonia I ,4-Diazabicyclo[2.2.2]octane 2,4,6-Trimethylpyridine N-Methylmorpholine Pyridine 1,IO-Phenanthroline Aniline
12.3 12.0 11.3 10.7 10.4 9.7 9.3 8.8 7.4 7.4 5.2 4.8 4.6
Amine/rhodiumb -0.1
0.2-0.3 0.3-0.5 0.3 0.3-0.5 0.3-0.5 0.3 0.2-0.3 0.2-0.3 0.3-0.4 -0.1 1.6 2.3
’ Data from Ref. 109. Minimum ratio of moles of amine to moles of Rh which provides the maximum rate to glycol.
358
13. I). DOMBEK
glycol productivity; bascs progressively weaker are required in greater amounts. Shown in Fig. 7 are the results of adding increasing amounts of two “strong bases,” 1,4-diaza[2.2.2]bicyclooctane (dabco) and N-methylmorpholinc (NMM), to rhodium catalyst solutions. The optimum ethylene glycol productivity is achieved in both cases at amine/rhodium ratios of less than 0.5. Differences in behavior are observed as further amounts of these amines are added. Higher concentrations of dabco cause diminished productivity, while increased amounts of N M M have little effect (109). Rates to methanol are affected little by either amine. Since the basicities of these amines are similar, other factors appear to be important in determining the degree of ethylene glycol rate decrease with excess amine. Kaplan has proposed that ion pairing between rhodium complex anions and the positively charged counterions has an adverse effect on catalytic activity for ethylene glycol formation (96, 109, 110).The following scheme: Aminc
+
[Amine H f &-I&!= Amine
[amine H t B-1
(30)
[amine H]++ &--catalysis
(31)
+ (alcohols. Rh species, etc.) % [amine HIf
r
P
(32)
Dabco
2 N- Methylrnorpholine
I
0.5
I
I
1.0 1.5 2.0 Arnine/Rh, Mole Ratio
1 2.5
FIG.7. Effect of amine/Rh ratio on product rates: (m) methanol; ( 0 )ethylene glycol. Reaction conditions: 75 ml sulfolane, 3 mmol Rh, 544 atm, HJCO = 1, 240’C,4 hr (109). Upper graph is for I ,4-diazabicyclo[2.2.2]octane(dabco); lower graph is for N-methylmorpholine.
ETHYLENE GLYCOL AND ETHANOL FROM H2 AND CO
359
is based on this assumption, and is consistent with the observation that amines act as promoters but can inhibit the reaction of interest at higher levels. In the scheme, &is a generalized rhodium complex, and &-is the catalyst or its immediate precursor. The amine and hydrogen act to reduce the rhodium species in (30), thus forming an anionic rhodium species as the ammonium salt. Less weakly basic amines may be required in larger amounts to force equilibrium (30) to the right, but strong. bases can apparently accomplish this step quantitatively. (Note the nearly constant amine/ rhodium ratio in Table X for bases with a pK > 5.) Increases in the concentration of added amine will also provide higher concentrations of [amine HI', by deprotonation of other rhodium species and hydroxylic products in solution as in (32); the more strongly basic the amine, the greater the fraction of it which will be converted to its conjugate acid. This increased concentration of [amine H]+ can be expected to shift equilibrium (31), an ion pairing process, toward an inactive (or less active) anion-cation complex. Thus, two factors determine the inhibitory ability of an amine: its basicity (K32)and the ion pairing ability of its conjugate acid ( K 3 , )(109). A weakly basic amine will produce little [amine H I + , and an excess of amine will therefore not greatly affect equilibrium (31). On the other hand, a strongly basic amine may be an excellent promoter even if it produces a high concentration of [amine HI', so long as this cation is poor at forming an ion pair in equilibrium (31). This would seem to be the optimum situation, since even that [amine HI+ formed by the (promoting) process of (30) will not inhibit the catalytic reaction. In general, it appears that amines which are found to cause little inhibition at higher levels [such as NMM, 2-hydroxypyridine, and 1,8-bis(dimethyIamino)naphthalene] have structures which allow, in the protonated form, delocalization of the positive charge over more than one atom. The freedom to use an amine at levels higher than optimum without causing a large glycol rate inhibition presents several benefits. The criticality of precisely controlling the amine/rhodium ratio is reduced, and higher levels of amine have been noted to improve the catalyst stability (109). Changes in reaction temperature and solvent dielectric constant are expected to affect the equilibria (30)-(32), and such effects are indeed observed (108).The addition of salts as promoters can also alter the optimum amounts of amine promoters to be used. 2. Salt Promoters Counterions for the anionic rhodium complexes present in catalyst solutions may also be provided by the addition of salts. A salt may be used as the sole promoter, but it appears that under many conditions a combination of salt and amine provides the best results. Table XI indicates that
360
B. D. DOMBEK
TABLE XI Efccr of Salts as Promorers in Rhodiutn-Carulyzcd CO Hydrogenation".' ~
Cation None Li+ Na+ Kt Rbt cs Mg' Sr2 Bat' PPN +
+
f
r.d
~~
~
Ethylene glycol rate ( h r - ' )
Methanol rate ( h r - I )
1.34 2.22 3.45 2.89 3.01 5.24 0.94 I .01 0.87 4.91
5.49 4.47 3.49 2.32 9.35 4.1 I 3.46 3.07 I .82 3.39
" Data from Ref. 88.
' Reaction conditions: 75 ml tetraglyme solvent, 3 mmol Rh, 10 mmol 2-hydroxypyridine, 0.45-0.50 mmol of acetate anion with cation specified, 544 atm. HJCO = I , 2 2 0 C , 4 hr. ' Data from Ref. 102. Bis(triphenylphosphine)iminium.
TABLE XI1 Effecr ?/'Cesium Salr.~on Rhodium-Catalyzed CO Hydrogenationanh
Anion
Ethylene glycol rate ( h r - l )
Methanol rate (hr- ')
Fluoride Chloride Bromide Iodide Formate Acetate 2-Pyridinolate Sulfate
5.51 4.77 3.49 2.42 5.71 5.24 5.85 I .94
3.85 4.97 4.40 4.53 4.32 4.1 I 4.43 4.04
Data from Ref. 88. Reaction conditions: 75 ml tetraglyme solvent, 3 mmol Rh, 10 mmol 2-hydroxypyridine, 0.45-0.50 mmol ofcesium cation with anions specified, 544 atm. HJCO = 1, 220°C. 4 hr. a
ETHYLENE GLYCOL A N D ETHANOL FROM H2 A N D C O
36 1
addition of various acetate salts to a catalyst solution already containing an amine promoter can cause substantial improvements in rate and selectivity to ethylene glycol. Under the conditions of these experiments (in tetraglyme solvent), cesium and bis(tripheny1phosphine)iminium ([(C,H,),P],N+ ; PPN') salts provide the best rates to the glycol product. In Table XI1 are shown the results of a series of experiments with a variety of cesium salts; formate and pyridinolate anions are found to be the most effective under these conditions. Halides were found not to be as effective in these experiments, although the use of higher levels of iodides is reported in a patent as being useful ( I ION). The amount of cesium salt added to these reactions was found to have a dramatic effect on the activity and selectivity of the system, as illustrated by Fig. 8. Maximum rate and selectivity to ethylene glycol are observed at a Rh/Cs' ratio of 6. At higher levels of salt the overall activity increases very substantially, but the amount of the glycol product diminishes rapidly. A similar series of experiments using the bulky PPN' cation was found to give somewhat different results (102), as shown in Fig. 9. The ethylene glycol rate maximum is observed at a rhodium/cation ratio of 4, and changes in glycol/methanol selectivity are not as pronounced as was found for the cesium cation. The bulky PPN' cation therefore has the advantage of providing a system in which the ethylene glycol rate and selectivity are not
Cs*/Rh. Mole Ratio
FIG. 8. Plot of rates as a function of C s + / R h ratio: ( 0 )methanol;).( ethylene glycol Reaction conditions: 75 ml tetraglyme solvent, 3 mmol Rh, 10 mmol 2-hydroxypyridine, 544 atm, H,/CO = I . 220 C, cesium formate promoter as indicated. 4 hr (88). Methanol and ethylene glycol rates at C s + / R h = 0 are 5.21 and 1.34 hr- I , respectively.
362
B. D. DOMBEK
as sensitive to increased cation/rhodium ratios. Similar results have been observed for quaternary phosphonium (If I) and ammonium cations (1f2), and the 3,3-bisdimethylamino-N,N,N"'-tetramethylacrylamidinium cation ( 9 9 ) .An advantage of being able to operate with higher salt concentrations is the possible increase in catalyst stability in the presence of higher promoter concentrations ( I f2, 113). Ion pairing interactions appear to be an important factor in rate inhibition for the glycol product as indicated above for amine promoters, and such effects apply in the case of salt promoters as well. The PPN' ion, for example, is large and has delocalized charge; it is therefore expected to interact only weakly with anions in solution. Alkali metal cations, however, may interact with rhodium complex anions as well as with the solvent and other anions in competitive processes. Cations with varying degrees of solvation will exhibit differences in ion pairing ability. A number of possible equilibria must then be considered if the effects of salt promoters are to be rationalized. Consistent with experimental observations, anionic rhodium complexes essentially free of ion pairing interactions are proposed (I 13) to provide the best catalyst. The best salt promoter would then be one in which the cation ion pairs the least with the rhodium catalyst. The anion of the salt is believed to act as a promoter by forming or transforming the anionic rhodium complexes involved in catalysis. Possible equilibria representing processes involved in inhibition by the cation are the following:
+ Rh--catalysis [ M + X - ]& M + + X-
[M+Rh-]%
M+
(33) (34)
x - + ROHK".XH + RO-
(35)
+ RO-
(36)
[MtRO-]&Mt
where Rh- is the active rhodium species formed by interaction of the salt promoter, MX, with the rhodium precursor. Any effects which reduce the interaction between M and jth- are expected to lead to increased rates of product (ethylene glycol) formation. It follows that the use of cation complexing agents o r high-dielectric-constant solvents, as described below, may enhance catalytic activity by increasing the amount of free &-. On the other hand, increased amounts of free M + will diminish the rate of catalysis by a mass law effect on equilibrium (33), decreasing the amount of free Rh-. The anion X- of the salt promoter can also have an effect on the amount of free M + (inhibitor) through the equilibria (34)-(36). A more +
ETHYLENE GLYCOI.. AND ETHANOL FROM H, A N D CO
-
0.083
0.167
0.250 0.333
363
0.5000.667
PPN + / R h , Mole Ratio
FIG.9. Plot of rates a s a function of P P N + / R h ratio: (m) methanol; ( 0 )ethylene glycol. PPN acetate was used as promoter ( / 0 2 ) .Reaction conditions are the same as those of Fig. 8.
basic anion will decrease the concentration of free M f both through a smaller equilibrium constant K , , and a larger equilibrium constant K , , . Because an excess of salt promoter has, under some conditions, been found to improve catalyst stability, it is regarded as desirable to employ as much salt as possible without substantially reducing the rate to ethylene glycol (113). By use of cesium carboxylates with a range of basicities, it was shown that, consistent with (33)-(36), the salts with the most basic anions gave the least rate inhibition when added at increasing levels (113). The ion pairing ability of cations can also be reduced by addition of complexing agents. For example, the use of alkali metal salts in the presence of “cryptands” ( 114, 1140) such as 4,7,13,16,2 I ,24-hexaoxa-1,1 O-diazabicyclopentatriacontane (222-crypt) :
is claimed to give increased catalyst stability and higher rates to ethylene glycol (103). The alkali metal cation can be enveloped in the interior of this and related molecules, thus preventing it from direct interaction with anions in solution. The size and geometry of the cryptand may be varied to provide an optimum fit for different metal cations. Similar results have been reported for “spherand” compounds, which also complex alkali or alkaline-earth metal cations (1146).
364
B. D. DOMBEK
3. Solvents The solvents used in these rhodium-catalyzed reactions may also act as complexing agents for counterions of the anionic rhodium complexes. For example, tetraglyme is known to coordinate alkali metal cations. Such solvation decreases the possibility of the cation interacting with the anionic rhodium catalyst and lowering its activity or solubility. The crown ethers, such as [ 181-crown-6
cone> to ,4
(38)
u
and [15]-crown-5, comprise a class of compounds which also can complex cations very effectively (114); these cyclic polyethers have been found to be excellent solvents for rhodium-catalyzed CO hydrogenation (92). The cyclic structures are multidentate ligands for the metal cations, and the effectiveness of the complexation is a function of the degree to which the ether can conform to the size of the metal cati0.n. An effective combination for the catalytic process has been found to be [18]-crown-6 with cesium salts as promoters (92). Although the cesium cation is too large to fit within this polyether, effective complexation is apparently achieved by the formation of a 2 : 1 crown ether-cesium “sandwich” complex, as indicated by crystal structure determinations (115, 116). Since the crown ethers are very effective complexing agents, the amount of free M + in solution, as in (33)-(36), is expected to be small; the crown ether competes very well with &-and X - for M’. Indeed, it is found that the addition of excess salt causes a much lower degree of rate inhibition in [18]-crown-6 as compared to some other solvents. For example, Fig. 10 illustrates the differences between [ 18]-crown-6 and tetraglyme as the level of salt promoter is increased. The capability of using an excess of salt reduces the criticality of precisely controlling the salt concentration and is beneficial for the stability of the catalyst (92). Another method of reducing ion pairing is to use a solvent having a high dielectric constant, such as sulfolane: n
o*
s) %O
(39)
This material has a dielectric constant of 43.3 at 30°C; it has very low proton basicity (pK,,,+ = - 12.9) and is a weak Lewis base (117). Indeed, sulfolane is an excellent solvent for the rhodium catalytic system, giving good rates
ETHYLENE GLYCOL AND ETHANOL FROM H2 AND CO
I
0.167
0.217
365
I 0.267
CS+/ R h , Mole Ratio
FIG. 10. Effect of cesium concentration on ethylene glycol rates in 18-crown-6 ( 0 )and tetraglyme (B)(92). Reaction conditions: 75 ml solvent, 3 mmol Rh, cesium benzoate, 544 atm, HJCO = I , 220'C, 4 hr.
and high rhodium recoveries (96). The potential instability of this material at high temperatures and resultant adverse effects on catalyst activity have been noted (94). Another patent shows that copper and its salts are useful in negating adverse effects of sulfur compounds which may be present as a result of the decomposition of sulfur-containing solvents (118). Another class of solvents having high dielectric constants is the lactones, such as y-butyrolactone :
and 6-valerolactone (95). The dielectric constant of butyrolactone is 39 at 20°C (47), and this solvent appears to give good rates and improved catalyst recoveries. These lactones, however, will polymerize to some extent during the reaction ( 9 3 , and may also react with hydroxylic products in a saponification process (93). Substituted butyrolactones are more stable toward these ring-opening reactions, and 2,2-dimethyl-y-butyrolactone, I
I
C(CH,),CH,CH,OC(O), has been shown to be superior to the unsubstituted analog (93). N-Methylpyrrolidone was noted by Keim et ul. (39) to be a good solvent for CO hydrogenation by rhodium catalysts in the absence of added promoters. The basicity of the compound [pKBH+ = -0.9 (48)]probably allows it to serve the same function as weakly basic amine promoters. The high
366
B. D. DOMBEK
dielectric constant of this solvent [ 3 2 at 25°C (47)] also indicates that it should be effective at separating ion pairs. Compounds related to this are the cyclic ureas, such as 1,3-dimethy1-2-imidazolidinone
which has been shown to be a very effective solvent ( 1 19). Some of the above solvents, such as tetragiyme and crown ethers, are effective because of their ability to complex cations, whereas others, such as sulfolane and butyrolactone, are useful by virtue of their high dielectric constants. Mixtures of these two types of solvents can lead to improved results, better than those obtainable in the single solvents. For example, Fig. 1 I shows that mixtures of sulfolane (high dielectric constant) and tetraglyme (complexing) solvents give improved rates and selectivities to the glycol product. The catalyst stability is substantially better in sulfolane than in tetraglyme under the conditions of these experiments, but it may be seen that a large fraction of the sulfolane can be replaced by tetraglyme before an adverse effect on stability is observed. However, as catalytic conditions become more severe (higher temperature or lower pressure), a higher sulfolane/tetraglyme ratio must be used to maintain the stability of the catalyst (94). The use of mixed complexing and high-dielectric-constant solvents is also illustrated for cyclic ureas (119), crown ethers (Y2), and phosphine oxides (YO), in the appropriate combinations.
0
20
40
60
80
100
Sulfolone, Volume Percent
FIG. I I . Effects on rate and catalyst stability of using sulfolane-tetraglyme mixtures as solvent: ( 0 )methanol; (m) ethylene glycol; ( A ) rhodium recovery. Reaction conditions: 75 mi solvent, 3 mmol Rh, 0.65 mmol cesium benzoate, 544 atm, HJCO = I . 240°C 4 hr (Y4).
ETHYLENE GLYCOL AND ETHANOL FROM H2 AND CO
361
Organic phosphine oxides are reported to be useful solvents for several reasons (90). Phosphine oxides are strong Lewis bases and can complex the positively charged counterions, thus increasing the amount of non-ionpaired rhodium complexes in solution. These compounds also possess relatively high dielectric constants, which decreases the attractive forces between anions and cations in solution. Additionally, the strong hydrogen bond acceptor capability of phosphine oxides allows them to complex with ethylene glycol and glycerol, which are bidentate hydrogen bond donors. Increasing concentrations of these polyalcohols, as noted above, can cause a decreased rate of ethylene glycol production and may also lower the catalyst stability. Apparently, phosphine oxides can counteract these harmful effects by the hydrogen-bonding interaction.
D. CATALYST STABILITY Many of the reactions described above are seen to give less than quantitative recovery of the rhodium catalyst component. The amount of rhodium remaining in a catalyst solution was determined by atomic absorption spectroscopy, and is reported as the percent of the rhodium charged which remains soluble or suspended in the reaction mixture at the end of the reaction (95). After some experiments a wash procedure was employed to dissolve rhodium complexes possibly left in the reactor; heating a charge of pure solvent in the reactor under H 2 / C 0 pressure sometimes dissolved substantial amounts of rhodium species (94-96, 104, 108, 109). High recoveries of rhodium are essential in a practical process because of the scarcity and high price of this metal (120, 121). The form of the unrecovered rhodium in experiments shown above is not certain. Although metallic rhodium has been found to be taken into solution to provide an active homogeneous catalyst under some conditions (85, 91), it appears unlikely that metallic rhodium is being dissolved in the reactor washes described above. In reaction media similar to those employed for catalytic reactions (tetraglyme or [ 181-crown-6 solvents, amine and salt promoters), rhodium clusters of decreased solubility may be produced from more soluble precursors. These clusters include [Rh,4(C0)2,]4- (122, 123), [Rhi 5(cO)27l3- (123), [Rh22(CO)35Hx+nI'5-')- (115, 124), clusters which are possibly [Rh,3(CO)24H,]'5-"'- (x = 0, 1) (122), and [Rh,,(C0)3,]4(116). These are relatively large, highly charged metal complexes whose growth may be reversed by the application of carbon monoxide pressure (115, 125, 126). It therefore appears possible that at least some of the precipitated rhodium noted in catalytic experiments could be in the form of similar highly reduced clusters, perhaps even higher in nuclearity. Under some conditions, reactions of ionic clusters under hydrogen ( 1 atm, 25°C)
368
B. D. DOMBEK
have been reported to give a precipitate of rhodium metal (126). Effects of the counterion on the growth and decomposition of anionic rhodium clusters have been reported (126a). The enhanced stability of the rhodium catalyst system under carbon monoxide is the subject of several patents. Separation of the alcohol products from a catalyst solution, by distillation for example, might involve the use of conditions under which the catalyst becomes insoluble. Contacting the catalyst-product mixture with carbon monoxide while distilling the product is claimed to minimize catalyst instability (127). Examples show the use of CO gas to strip products from heated catalyst solutions. Comparisons with the use of N, gas demonstrate that CO does indeed stabilize the catalyst. Another patent describes the use of a continuous reactor with catalyst solution recycle in which such a CO stripping column is used to remove products from the reaction mixture (128). The rhodium catalyst was found to be continuously lost from solution at a low rate during extended operation of the continuous unit. The rhodium level, however, could be increased nearly to the initial value by lowering the H2/C0 ratio in the reactor vessel, i.e., increasing the CO partial pressure while decreasing the H2 partial pressure (128). This apparently resolubilizes rhodium species which have been lost from solution in the reactor vessel. Another patent shows that periodically lowering the temperature of the reactor also has the effect of raising the rhodium concentration, apparently by causing the solubilization of precipitated rhodium species (129). Other patents illustrate the use of a solvent extraction process to separate the alcohol products from the catalyst (130, 131). When a catalyst solution containing alcohol products is mixed with water and a water-immiscible solvent, the alcohol products are extracted into the aqueous phase and the rhodium species enter the water-immiscible solvent. The effectiveness of the extraction and the stability of the rhodium catalyst can be greatly increased by carrying out the process under CO pressure (131). The general behavior of rhodium catalysts with respect to stability thus appears to be similar to that seen for cobalt catalysts; an inverse relationship between carbon monoxide partial pressure and reaction temperature is apparent. Stability decreases rapidly with increasing temperature, and raising the pressure tends to improve catalyst stability. It is not certain whether the adverse effects of increasing the H2/C0 ratio are merely the result of a decreased CO partial pressure, or whether increased hydrogen partial pressure induces catalyst instability. Catalyst stability in this system is substantially influenced by the characteristics of solvents and promoters. Indeed, the properties of solvents and promoters which improve the catalytic activity for ethylene glycol production (increased dielectric constant, greater cation complexing ability, or
ETHYLENE GLYCOL AND ETHANOL FROM H 2 A N D C O
369
lower ion-pairing ability) also appear, in general, to improve catalyst stability.
E. MECHANISM The characteristics of the rhodium catalytic system described above suggest that this is a very complex system. No simple concentration dependences are evident, and subtle ion-pairing effects can have a large influence on activity and selectivity. Studies of the rhodium chemistry also indicate a high degree of complexity in this system. Although many rhodium complexes may be used as catalyst precursors (86),the most commonly used precursor is Rh(CO),(acac) (acac = acetylacetonate). It is known that reduction of this and similar mononuclear rhodium species by bases under carbon monoxide affords a rhodium cluster anion, [Rh 2(CO)3,]2- (106, 132-135), whose structure is shown in Fig. 12. This complex is useful as a catalyst precursor, and various mixed-metal clusters of the same general structure containing cobalt, rhodium, and iridium in the cluster framework have been prepared (136-138). Although these compounds are reported to be useful as catalyst precursors, no catalytic results are given. The [Rh,,(CO),,12- cluster is in equilibrium under CO with another anionic cluster, initially identified as [Rh, 2(CO),,]2- (106). Later studies, which involved the low-temperature isolation of this very labile complex, showed it to be [Rh,(CO),,]- (107),whose solution structure (139) is shown in Fig. 13. Studies by infrared (126, 140) and NMR (139) spectroscopy have shown that [Rh, 2(CO)3,]2- is essentially completely converted to [Rh,(CO),,]- under relatively low CO pressures (5 atm at 25°C) as follows:
a
3[Rh, r ( C O ) 3 0 1 Z ~ 6[Rh5(C0),s1-
b F c . 12.
+ RhJCO),,
d
Molecular structure of [Rh,,(CO),o]Z- (135).
(42)
B. D. DOMBEK
370
0
FIG. 13. Solution structure of [Rh,(CO),,]- (139). x represents a bridging carbonyl ligand.
Catalytic reaction solutions prepared from Rh(CO),(acac) in the presence of amines and/or carboxylate salts show the presence of [Rh,(CO),,]- and a mononuclear species, [Rh(CO),] - , when observed by high-pressure infrared spectroscopy (86, 140). The spectral features of these mixtures remain unchanged as the temperature is increased up to 180°C (at 500 atm of HJCO) of 210°C (at 1000 atm). At temperatures above these values significant broadening and shifting of the absorption bands occur, and the species present cannot be identified with certainty (140). (It is interesting to note in this regard that most of the reported catalytic experiments have been carried out at temperatures above these values.) The changes reported at higher temperatures are reversible, and bands assignable to [Rh5(CO)l,I- and [Rh(CO),]- reappear upon cooling the solutions below 180-210°C under pressure. Although the identity of rhodium complexes present at temperatures and pressures generally employed for catalytic experiments cannot be determined with certainty by infrared spectroscopy, rhodium clusters of the formula [Rh, 3(CO),,H,]'5-"'- (x = 2, 3) are possibly present under such conditions. These are known complexes (141), stable under an atmosphere of carbon monoxideat room temperature, and obtainable from [Rh,,(C0),,]2by reaction with H, ( W C , 1 atm). The structure of [Rhl,(CO),,H,]2- is shown in Fig. 14. It has also been found that [Rh,(CO),,]- can be converted to these clusters by the following reaction with HJCO, the temperature required being determined by the H,/CO pressure (140): [Rh,(CO),,]-
-+
[Rh,3(C0)24H,]'5-X)-(x
=
2 , 3)
+ higher clusters
(43)
ETHYLtNE GLYCOL A N D ETHANOL FROM H 2 AND CO
37 I
d FIG.14. Molecular structure of [Rh,,(CO),,
1’-
(13
The temperature was 80, 190, and 240 C, respectively, at HJCO pressures of 1, 600, and 1000 atm. [The higher clusters presumably include [Rh,,(CO)z5]4- and [RH I 5(CO)27]3-(142).] Such clusters therefore seem likely to be present in reaction solutions during catalysis. The shifting of equilibria to higher clusters with an increasing number of metal-metal bonds upon raising the temperature appears to be a general phenomenon (126,143); it has been observed for other rhodium clusters at lower temperatures and pressures (135). The 13-Rh-atom clusters of (43) are interconvertible by protonation and deprotonation; amines such as N M M serve to deprotonate the trihydride. These clusters can be fragmented by carbon monoxide as follows :
These interconversions illustrate the general finding that CO pressure causes cluster fragmentation, whereas replacing CO by H, allows cluster growth to occur. An important feature of the fragmentation process shown is the removal ofmononuclear [Rh(CO),]- or HRh(CO), units from larger clusters by reaction with CO. Evidence for further fragmentation of [Rh,(CO),,]to [Rh(CO),]- and Rh,(CO), is reported based on high-pressure infrared spectroscopy (126), but such a transformation is not observed in highpressure NMR experiments (139, 144).
372
B. D. DOMBEK
One function of amine and other basic promoters may be to facilitate cluster transformation by allowing facile protonation/deprotonation processes to occur. The [Rh,3(C0)24H3]2-cluster is a sufficiently strong acid to be deprotonated by NMM, as seen in (44).The HRh(CO), species has been found to be fully deprotonated by N M M and N,N-dimethylaniline (145). Protonation and deprotonation in reaction systems containing hydrogen are linked to oxidation and reduction processes. For example, addition of hydrogen to a metal cluster followed by deprotonation leads to a net reduction of the metal species; conversely, protonation and elimination of H 2 causes oxidation of the metal species (123, 125). The basicity and amount of promoter employed can therefore determine the metal oxidation state of the system. The preferred oxidation state of rhodium in the catalytic system for ethylene glycol formation appears to be between the extremes of 0 and - 1. Catalytic systems containing a high proportion of [Rh(CO),](oxidation state - 1) produce largely methanol and little ethylene glycol (102), whereas systems containing neutral rhodium complexes such as Rh,(CO)], (oxidation state zero) have low activity without a basic.promoter (reducing agent). Another possible function of certain promoters may be to facilitate the transfer of “Rh(C0);” fragments between clusters (123, 126), as follows: [Rhl,(C0)z,]3- t 2L
* [Rh14(CO)z5]4-+ [Rh(CO),L,]+
(45)
The ligand which removes the “Rh(C0);” fragment may be a halide ion (146), an amine, or a solvent molecule (123). Reactions according to (45) were observed for N M M , bis(N,N-dimethyl)-ethylenediamine,and 1,lOphenanthroline. In reactions of the first two amines, infrared absorptions were detected which could possibly arise from the [Rh(CO),(amine),]+ species; such complexes have been studied previously (147). The [Rh,, .(CO)2,]3- cluster is observed to react also with certain solvents according to (45). Although this cluster is stable at ambient conditions in acetone and tetraglyme (relatively low-polarity solvents), it reacts in sulfolane and [ 181-crown-6 to form [Rh,,(C0)2,]4- (123). It is suggested that a decrease in ion pairing in the latter two solvents may have facilitated release of the “Rh(C0):” fragment. Carboxylate promoters may also be able to coordinate to an “Rh(C0):” fragment, and therefore facilitate a process such as that shown in (45). Reactions of [Rh, s(CO)2,]3- with H2 in the presence of cesium carboxylates are reported (123) to be consistent with the formation of small equilibrium amounts of Rh(CO),(O,CR) by reaction (45). Carboxylates could therefore be involved in cluster growth or transformation during catalysis. The possibility that a cluster framework could be an important feature in determining the activity of a catalyst has led to investigations of less labile
ETHYLENE GLYCOL AND ETHANOL FROM H2 AND CO
373
rhodium clusters as catalyst precursors. A number of rhodium clusters containing encapsulated main-group elements have been prepared, including [Rh, 7S2(CO)32]3- (148-15O), [Rh,C(CO), 532- (151), [Rh,P(CO), (152- 154), [Rh ,Sb(CO), J 3 - ( 1 5 9 , [Rh oP(CO),,]3- (156), and [Rh ,As (C0),,l3 - (157). The encapsulated atoms stabilize the cluster framework, and several of these complexes have been observed to remain intact under high pressures of H,/CO in the presence of promoters, conditions which normally lead to CO hydrogenation (4). Table XI11 shows that the activity of such systems is substantially lower than that of a normal rhodium catalytic system, so the possibility that small amounts of the clusters are fragmented (perhaps reversibly) to form catalytically active species cannot be excluded. The role of metal clusters in the rhodium-catalyzed hydrogenation of CO remains uncertain. It is evident that various cluster species are present during catalytic operation, but it is also clear that labile fragmentation and rearrangement processes are possible. Indeed, these processes are facilitated by the species observed to promote catalytic activity. Criteria set forth by Laine for identifying cluster-catalyzed reactions (158) are not definitive for this process. The pressure dependence of the reaction has been suggested to be attributable to a shifting of equilibria between clusters in solution at varying H,/CO pressures (43); however, the identity or characteristics of the active species are not apparent. The general outline of steps leading to the primary oxygenated products presented above for cobalt catalysts (a chain growth process which proceeds through aldehyde intermediates) may also apply to the rhodium system. Certainly, the same array of products is observed in both systems, although secondary reactions are evidently less predominant in most of the rhodium TABLE Xi11 Catalytic Activity of Systems Based on Stabilized Clusters"~h
Complex
Ethylene glycol rate (hr- I )
Methanol rate (hr- I )
Data from Ref. 140. (Adapted with permission. Copyright 1980 American Chemical Society.) Reaction conditions: 75 ml sulfolane, 3 mmol Rh, 5 mmol NMM, 1000 atm, H,/CO = 1,260"C. 0.375 mmol cesium benzoate added.
374
B. D. DOMBEK
reactions. The pathways leading to 1,2-propyleneglycol (101)and n-propanol (100) in certain promoter-modified rhodium reactions are not certain. Although the products are possibly formed entirely by secondary reactions, definitive experiments are not reported. The possible intermediacy of formaldehyde in CO hydrogenation has been addressed above with regard to the cobalt catalytic system. Fahey has observed a small amount of 1,3-dioxolane (the ethylene glycol acetal of formaldehyde) as a product of the rhodium system (43). Thus, there is evidence that formaldehyde or a complexed form of this molecule could be an intermediate in the CO reduction process by this system. Rhodium catalysts are indeed found to be useful for the hydroformylation of formaldehyde to glycolaldehyde (159-261); methanol is a by-product in these reactions. An experiment in which I4CH2Owas added to a rhodium-catalyzed CO reduction system showed that the label was incorporated into all of the expected products, including ethylene glycol, methanol (and their formate esters), ethanol, and the ethylene glycol acetal of glycolaldehyde (91). The label was not found in CO or C 0 2 .These results support the general mechanism described above in which (coordinated) formaldehyde is a precursor of methanol and glycolaldehyde, which is itself a precursor of ethylene glycol and higher polyalcohols. An interesting result of this experiment is that the ethylene glycol/methanol mole ratio from I4CH2O (3.5) is substantially higher than that for the overall reaction (1 3.This may indicate an alternative route to glycol via hydrodimerization of formaldehyde, which would result in a higher concentration of label in the glycol product. This finding may also imply that there are slight differences in selectivity for the hydroformylation of free (labeled) formaldehyde and coordinated (unlabeled) CHzO-an entirely plausible possibility. The absence of I4CO at the end of this experiment may indicate that formation of formaldehyde by this system is an essentially irreversible process. This cannot be definitely concluded, however, in light of the probable reactivity differences between coordinated and uncoordinated formaldehyde. Several differences between the cobalt- and rhodium-catalyzed processes are noteworthy with regard to mechanism. Although there is a strong dependence in the cobalt system of the ethylene glycol/methanol ratio on temperature, CO partial pressure, and Hzpartial pressure, these dependences are much lower for the rhodium catalyst. Details of the product-forming steps are therefore perhaps quite different in the two systems. It is postulated for the cobalt system that the same catalyst produces all of the primary products, but there seems to be no indication of such behavior for the rhodium system. Indeed, the multiplicity of rhodium species possibly present during catalysis and the complex dependence on promoters make it
ETHYLENE GLYCOL AND ETHANOL FROM H2 AND CO
315
appear possible that several species may be catalytically active, each having its own product-forming selectivity. Any conclusions about the mechanism of the rhodium-catalyzed process, however, must await more detailed catalytic and chemical studies. IV.
Unpromoted and Carboxylic Acid-Promoted Ruthenium Catalysts
A. BACKGROUND Homogeneous ruthenium catalysts have been reported to convert H 2 / C 0 to methane (22) and a distribution of linear alkanes (10, 162). The argument was presented that a metal cluster, such as Ru,(CO),~,was an essential catalyst component in order to form the alkane product, and it was claimed that the mononuclear Ru(CO), was inactive (10). Further studies have shown, however, that methane and alkanes are formed by rutheniumcontaining catalytic systems only when metallic ruthenium (formed by decomposition of the homogeneous catalyst) is present (163-167). Strictly homogeneous solutions of ruthenium complexes are found not to produce alkanes, but instead usually yield methanol as the major product. The use of certain additives or solvents has been found to improve the activity or selectivity to the methanol product (164-168). It has further been found that carboxylic acids can cause this system to form ethylene glycol as its carboxylate ester in addition to the methyl ester (166, 167, 169-172). The similarity in many details of the chemistry in the above homogeneous systems suggests that they are modifications of the same basic ruthenium catalyst system. The most notable common feature of these reactions, to be discussed in this section, is the presence of predominantly Ru(CO), as the stable ruthenium species during catalysis. Addition of ionic promoters, particularly halide salts, to ruthenium-containing solutions has been found to provide catalytic systems with very different characteristics. Such systems contain ruthenium complexes other than Ru(CO), during catalysis, and will be described in Section V.
B. CATALYTIC ACTIVITY AND SELECTIVITY Catalyst solutions generated by the reaction of Ru(acac), or Ru,(CO),, with H 2 / C 0have been reported by Bradley to produce methanol and methyl formate as the major products (264, 165). Methyl formate is produced at a constant rate, suggesting that it is a primary product and not derived from
376
B . D. DOMBEK
initially formed methanol. Reactions at lower pressure were found to give much smaller relative yields of methyl formate (166), consistent with the effects of pressure on the formate/alcohol ratio observed for the cobalt and rhodium catalytic systems. Reactions in a variety of solvents give a range of activities, as seen in Table XIV. Tetrahydrofuran (THF) and sulfolane provide lower activities than those observed in certain other solvents, such as ethanol and ethyl acetate. Carboxylic acids provide methanol (ester) rates comparable to those found in the latter solvents, but are notable in that ethylene glycol esters are also observed. Ethylene glycol is not normally observed as a product in the absence of carboxylic acids. Small amounts of ethylene glycol have been reported as products after ruthenium-catalyzed reactions in NMP and toluene solvents at 2000 atm (39). However, observations of minor amounts of this product must be viewed with caution unless great care is taken in the experimental procedure. For example, it was earlier reported that a catalyst derived from Ru,(CO),, TABLE XIV Ruthenium-Catalyzed CO Hydrogenation ~
Expt.
Solvent
I 2 3 4 5 6 7 8 9
TH F Ethanol Ethanol THF Sulfolane Ethyl acetate Acetic acid Acetic acid Propionic acid THF THF THF
10
II 12
~~
Pressure Temp. -OCH2CHLONotes (atm)" ("C) rate ( h r - l ) b C C C C
d e e
.f g
h i
1300 340 340 340 340 340 340 340 340 1200 1200 1200
268 260 230 230 230 230 230 260 230 275 275 290
-
0.29 0.34 0.22
HJCO = I unless specified otherwise. HJCO = 1.5 (164). 'Methyl formate levels very low (166). Includes methyl acetate formed by transesteritication (166). 'Methanol and ethylene glycol detected as acetate esters (166). Methanol and ethylene glycol detected as propionate esters (166). From Ref. 165. Includes added triphenylphosphine, P/Ru = 3 (165). H,/CO = I . 5 . Methanol rate includes methyl formate (164). j Not detected.
~
CH30rate (hr-I) 27.5 23.2 11.8 4.12 3.19 7.04
Methyl formate rate (hr-l) 3.06 n.d.1 n.d. n.d. n.d. n.d.
11.1
29.6 13.0 35 51 I I6
-
18
3.7
ETHYLENE GLYCOL AND ETHANOL FROM Hz AND CO
377
and 2-hydroxypyridine in tetraglyme solvent under pressures of at least 1700 atm gave substantial amounts of ethylene glycol (97, 98). Later studies by the same researchers (173, 174) led them to conclude that the ethylene glycol in these experiments was actually produced by rhodium species leached from the walls of the vessel during reaction, and originating from earlier rhodium-containing experiments. A patent also claims the production of ethylene glycol by a ruthenium catalyst in the presence of 2-hydroxypyridine in n-propanol at 1700 atm (175); this result appears suspect in light of the conclusions cited above. Traces of ethylene glycol have been detected in catalytic solutions derived from RU,(CO),~in T H F solvent, after reaction at pressures of 1000-1500 atm (176); a blank run containing no RU,(CO),~immediately preceding these experiments produced no detectable glycol. The major products of these ruthenium-catalyzed experiments were found to be methanol and methyl formate. Other products have also been reported from ruthenium-catalyzed CO reduction experiments. Ethylene glycol and its ethers are reported as products from ruthenium-containing solutions in the presence of polyhydric phenols and, optionally, mineral acids (177). The fact that the only solvent employed in the examples cited is the dimethyl ether of tetraethylene glycol (tetraglyme) suggests that solvent decomposition could lead to these products, particularly in the presence of the acidic “promoters.” The decomposition of glyme solvents in catalytic solutions containing acidic HCo(CO), has already been discussed (53). Methanol, methyl formate, dimethyl ether, and acetone are reported as products from ruthenium-containing solutions in 2-methoxyethanol(144). Once again, solvent decomposition is implicated, and the source of these products remains uncertain. Disregarding ambiguous results perhaps caused by metal precipitation, catalyst contamination by rhodium and possibly other metals, and catalyzed solvent decomposition, it appears probable that homogeneous ruthenium catalysts in the absence of ionic promoters can produce essentially only the one-carbon products methanol and methyl formate. [Under much higher pressures (3000-3600 atm) it is reported that higher linear alcohols may also be obtained (174).] Ethylene glycol is at most a trace product, except in the presence of carboxylic acids which cause the formation of ethylene glycol esters. [Siloxane derivatives of ethylene glycol are also obtainable by carrying out ruthenium-catalyzed reactions in the presence of reactive hydrosilanes (46, 46b, 46c).] The amount of ethylene glycol product formed in acetic acid solvent is usually minor relative to the methanol product. Table XIV, for example, shows examples in which the C , / C , product ratio is within the range of about 35-90. Esters of the three-carbon polyalcohol, glycerol, have also been
378
B. D. DOMBEK
detected in these product mixtures. The CJC, product ratio is even smaller, generally within the range of 200-300 (179). Other minor products in these reaction solutions include esters which result from the hydrogenation of carboxylic acid to alcohol (166). A first-order dependence of the CO reduction rate on metal concentration has been observed in these systems (164-167). Dependences on hydrogen and carbon monoxide partial pressures are less simple. A study of these effects on the rate of ethylene glycol ester formation in acetic acid solvent (164) showed that the dependence on hydrogen partial pressure was constant, with an order of about 1.3 (Fig. 15). The dependence on CO partial pressure is more complex, exhibiting a high dependence at low partial pressures (100 atm at 23OoC),but showing zero-order dependence at higher CO partial pressures (200 atm). (The effect of these partial pressures on the rate of methyl ester formation was parallel ; only minor changes in selectivity were noted over the range of pressures investigated.) Bradley has reported that under yet higher partial pressures of CO (ca. 500 atm), increases in CO pressure cause decreased rates of CO hydrogenation (164). Increases in CO partial pressure have also been reported to enhance the
0.30
r
170 atrn
co
170 atrn H2
-
0.05
75
100 150 200 250 Pressure of CO or Hz, atrn
FIG. 15. Dependences of ethylene glycol diacetate yield (formation rate) on CO and H, partial pressures. (Adapted from Ref. 166 with permission. Copyright 1980 American Chemical Society.) Partial pressure of reagent not being varied is 170 atm. Reaction conditions: 50 ml acetic acid, 2.35 mmol Ru, 2 3 0 C , 2 hr.
ETHYLENE GLYCOL A N D ETHANOL FROM H, A N D CO
379
formate/alcohol ratio. It was further found that addition of triphenylphosphine had a negligible effect on the overall rate of product formation, but caused a substantial decrease in the formate/alcohol ratio [cf. Expts. 10 and 1 1 in Table XIV (1641. The temperature dependence of the CO reduction process has also been studied. Over the range 250-290°C under 1200 atm of H,/CO, an Arrhenius activation energy of 32 kcal/mol was reported (164). The activity of this ruthenium system is comparable to, or somewhat greater than, that of cobalt catalysts under the same conditions of temperature and pressure. Rhodium catalysts provide substantially higher activity than either of these systems. As will be seen later, however, addition of ionic promoters can greatly increase the activity of ruthenium-based catalysts. C. SOLVENTS As described above and shown in Table XIV, the identity of the solvent may have significant effects on the rate of CO reduction. Alcohols, esters, and carboxylic acids appear to provide the highest rates, whereas T H F and sulfolane are somewhat less effective. Heptane solvent has been reported to afford poor rates of CO reduction by this system (163). Differences in rates among these solvents appear small enough to be attributable to an effect such as the enhanced stabilization of a polar transition state by the more polar solvents. The presence of certain additives, such as boric acid and aluminum alkoxides, has also been found to increase the rate of CO reduction, perhaps for similar reasons (168). Carboxylic acids appear to have a role in this system more complex than that of other solvents, since they are incorporated into the products and alter the selectivity by promoting formation of the two-carbon glycol product. Noncarboxylic acids have not been found to possess this ability to induce glycol formation; Br~nstedacids with a range of acidities have been investigated but found not to be effective (166). For example, pentachlorophenol, which has a pK, similar to that of acetic acid, does not promote the formation of significant amounts of a glycol product when used as solvent (167). Addition of other acids, such as H,PO,, to carboxylic acid solvents is not observed to enhance the rate or selectivity of two-carbon product formation. However, a variety of carboxylic acids promote formation of the glycol ester products (166, 167, 171). Thus, carboxylic acids are quite specific promoters for glycol formation, and acidity alone is not the source of this promoter effect. The influence of carboxylic acid concentration on the rate of two-carbon product formation has been investigated. Dilution with other solvents causes
380
B. D. DOMBEK
.-
0.30
0.20 0.15
-
0.10
-
rn U
3
2 0.08-
0 3
P
5 0.06 -
.0
- 0.04
F 0.030
(3
t
5 0.02-
5
L 4
6
8 10
15
20
Acetic Acid, M
FIG.16. Log-log of ethylene glycol diacetate yield vs acetic acid concentration when diluted with varying amounts of methyl acetate and water. (Adapted from Ref. 166 with permission. Copyright 1980 American Chemical Society.) Reaction conditions: 50-75 ml solvent, 2.35 mmol Ru, 340 atm, HJCO = 1, 230 C, 2 hr.
a rapid decline in the rate of glycol formation (166, 169). The rate of glycol production is approximately proportional to the second power of carboxylic acid concentration, as shown in Fig. 16. In contrast, the rate of methanol or methyl ester formation is changed little upon altering the acid concentration. Studies of ruthenium-catalyzed reactions in carboxylic acid solvents have been reported by Knifton (171, 172), but most of these experiments contain added salt promoters which greatly modify the catalytic behavior. These experiments will be considered in Section V, along with other Lewis basepromoted ruthenium systems.
D. CATALYST STABILITY The stability of soluble ruthenium carbonyl species toward decomposition to metal is a function of both carbon monoxide partial pressure and reaction temperature, similar to the situation described earlier for cobalt complexes and shown in Fig. 4. However, a quantitative study of these variables on ruthenium stability has not yet been reported.
ETHYLENE GLYCOL AND ETHANOL FROM H2 AND CO
38 1
Solutions of ruthenium carbonyl complexes in acetic acid solvent under 340 atm of 1 : 1 H,/CO are stable at temperatures up to about 265°C (166). Reactions at higher temperatures can lead to the precipitation of ruthenium metal and the formation of hydrocarbon products. Bradley has found that soluble ruthenium carbonyl complexes are unstable toward metallization at 271°C under 272 atm of 3 : 2 H,/CO [I09 atm CO partial pressure (165)]. Solutions under these conditions form both methanol and alkanes, products of homogeneous and heterogeneous catalysis, respectively. Reactions followed with time exhibited an increasing rate of alkane formation corresponding to the decreasing concentration of soluble ruthenium and methanol formation rate. Nevertheless, solutions at temperatures as high as 290°C appear to be stable under 1300 atm of 3 :2 H,/CO. Careful studies by Doyle et ul. (163) have also shown that soluble ruthenium species are inactive for hydrocarbon formation. A soluble system could be maintained in heptane solvent at 250°C under 100 atm of 1 : 1 H,/CO for many hours by taking precautions to avoid the possible introduction of impurities into the system and by slowly raising the temperature. No hydrocarbon formation was observed in this reaction. Only upon heating to about 260°C was the disappearance of soluble ruthenium complexes noted, along with the formation of linear alkanes. These results may suggest that metastable homogeneous ruthenium solutions can be formed, as has been reported for cobalt complexes (56); precipitation of the metal may be an autocatalytic process.
E. MECHANISM Information from several sources is relevant to the identity of ruthenium species present in these catalytic solutions. Reactions of ruthenium complexes under 200 atm of 1 :2 H,/CO at 180°C (180) or 80 atm of CO at 150°C (181) have been reported to produce mainly Ru(CO), . The mononuclear species is formed as an equilibrium product from Ru,(CO),, under CO pressure, the position of the equilibrium +RU,(CO),~+ CO + Ru(CO),
(46)
depending on the temperature and CO pressure. Studies of this equilibrium have been carried out at temperatures of 75-125°C and CO pressures of 10-60 atm (182). At 100°C under 60 atm of CO, an equilibrium solution approximately 3 x M in ruthenium contains Ru(CO), to the extent of about 99%. Under 10 atm of CO, slightly more than 50% of the ruthenium is in the form of Ru(CO),. Higher temperatures favor the equilibrium formation of the Ru,(CO),, cluster.
High-pressure infrared studies of ruthenium carbonyl solutions under HJCO at temperatures employed for CO reduction have also been reported. I n Ir-tetradecane solution at 180 C under 1 : 1 H 2 / C 0 , mainly Ru(CO), is detected (60).In acetic acid solvent at 200 C, only Ru(CO), is detected under 400 atrn of 1 : 1 H,/CO; at HJCO pressures of 200 atm, Rii3(CO)12is also observed (166). Reaction solutions have also been studied by sampling under reaction conditions, rapidly cooling the samples to low temperatures, and analyzing them by infrared spectroscopy; after reaction at 265 atm of 1 : 1 HJCO at 180 C , only Ru(CO), could be detected (164). At higher temperatures and lower pressures (100 atm of 1 : 1 HJCO and 250°C). evidence was seen for the clusters Ru,(CO),, and H,Ru,(CO),, as well as Ru(CO), ( 163). The presence of mainly the mononuclear Ru(CO), species under catalytic conditions, in combination with the observation of first-order rate dependences on ruthenium concentration, indicates that a mononuclear catalyst is involved in this CO reduction process. Reactions before or during the rate-determining step which involved metal cluster formation, or other processes requiring more than one metal species, would cause rate dependences of higher order in metal concentration. The observation of equilibrium (46) can also explain the changing rate dependences on CO partial pressure (Fig. 15). Under high CO partial pressures (e.g., > 200 atm under the conditions of Fig. 15) zero-order or negative rate dependences are observed; under these conditions equilibrium (46) is shifted essentially completely to the mononuclear species. Under lower partial pressures of CO (e.g., 100 atm under conditions of Fig. 15) cluster formation is observed to occur, thus lowering the concentration of effective (mononuclear) catalyst present. Therefore, the initial rate enhancement observed upon increasing a low CO partial pressure is the result of increased cluster fragmentation and the generation of active catalyst species. At the same time, however, CO may be an inhibitor of the actual hydrogenation process. Once the cluster fragmentation equilibrium has been shifted far toward the mononuclear species, only the inhibition effect is observed. The effect on CO reduction rates of H, partial pressure has been found to be somewhat greater than first order (ca. 1.3) (166). The existence of a nonintegral dependence on hydrogen pressure suggests the participation of an equilibrium involving hydrogen addition prior to the rate-determining step. It is known that Ru(CO), reacts with H z under pressure to form H,Ru(CO), (181), and this reaction is a plausible equilibrium process under catalytic conditions. A scheme consistent with the observed behavior of the system can be constructed if a second molecule of H z reacts with a catalytic intermediate before or during the rate-determining step, as follows: RU(CO),+
H,Ru(CO),&
product
(47)
ETHYLENE GLYCOL AND ETHANOL FROM H2 AND CO
383
Assuming steady-state behavior, the rate law for this process is as follows:
d [ product] - k k , [Ru(C0),]Pi dt k-IP,, + k2PH This equation predicts a rate dependence on H, partial pressure of between first- and second-order and a CO dependence of between zero- and negative first-order, as well as first-order dependence on Ru(CO), concentration. Based on this scheme, it may be suggested that H,Ru(CO), is the active catalyst species in this system. The steps by which this metal hydride forms the observed organic products are perhaps similar to those already discussed for cobalt catalysts. Steps which may be involved are intramolecular hydride migration to produce a formyl ligand : (CO),RuH,
* (CO),HRu-CHO
(49)
followed by rate-determining H, addition and production of a formaldehyde intermediate :
H
7\ H
(As discussed previously, thermodynamics indicate that free formaldehyde
will not be a major product of this reaction, although a small equilibrium concentration may be formed.) 0
0
II II ( C O ) ~ H R U - C C H Z O H-HCCHzOH
cY \
cv
(51)
(CO), HRu-CHZOH
0 (c0)3H1zRu
-Id. H' \ H
(52)
CH3OH
0 II
k
-
Y
0 11
(CO~HRU-COCHJ-HCOCH~
(53)
CH30H
(54)
(CO). HRu OCHS
384
B . D. DOMBEK
The glycolaldehyde shown in (51) results from a “CO insertion” reaction followed by reductive elimination, and is presumed to be a precursor of ethylene glycol. Since ethylene glycol is, however, at most a trace product of this catalytic system, step (51) appears to be essentially inoperative. Methyl formate, a major primary product of this system under some conditions, is also presumed to be formed by a CO insertion process, (53). Methanol may be formed by a reductive elimination (hydrogenolysis) of either a hydroxymethyl ligand, (52), or of a methoxy ligand, (54). The scheme in (51)-(54) is useful in considering possible reasons that ethylene glycol is not produced by this catalytic system. One possibility is that the alkyl migration, or “CO insertion” process of (51) is particularly unfavorable for the hydroxymethyl ligand. However, Roth and Orchin have demonstrated a reaction which apparently involves alkyl migration of a cobalt hydroxymethyl complex, (20), at quite low temperatures (75). Also, ethers and esters of the hydroxymethyl ligand (which are expected to be very similar electronically to the parent hydroxymethyl) have been shown to undergo CO insertion in manganese complexes under mild conditions (68). These studies suggest that differences in alkyl migration behavior between hydroxymethyl complexes and other simple alkyl complexes are small and of a quantitative rather than qualitative nature. Another possible reason that ethylene glycol is not produced by this system could be that the hydroxymethyl complex of (51) and (52) may undergo preferential reductive elimination to methanol, (52), rather than CO insertion, (51). However, CO insertion appears to take place in the formation of methyl formate, (53), where a similar insertion-reductive elimination branch appears to be involved. Insertion of CO should be much more favorable for the hydroxymethyl complex than for the methoxy complex (67, 83). Further, ruthenium carbonyl complexes are known to hydroformylate olefins under conditions similar to those used in these CO hydrogenation reactions (183, 184). Based on the studies of equilibrium (46) previously described, a mononuclear catalyst and ruthenium hydride alkyl intermediate analogous to the hydroxymethyl complex of (5 1) seem probable. In such reactions, hydroformylation is achieved by CO insertion, and olefin hydrogenation is the result of competitive reductive elimination. The results reported for these reactions show that olefin hydroformylation predominates over hydrogenation, indicating that the CO insertion process of (51) should be quite competitive with the reductive elimination reaction of (52). The evidence then suggests that the reason ethylene glycol is not formed by this system is that its hydroxymethyl precursor is not efficiently produced in the first step of (51). It follows that most of the methanol produced by this catalytic system must be formed by pathway (54), through a methoxide
ETHYLENE GLYCOL AND ETHANOL FROM Hz AND CO
385
intermediate. The very low selectivity of this system for the glycol product then appears to be determined by the preferred conversion of coordinated formaldehyde into a methoxy ligand. [Added paraformaldehyde has also been observed to be converted quite effectively to methanol by this system (179), presumably via (54).] Factors influencing the direction of formaldehyde insertion into a metalhydrogen bond, and thus the product selectivity in a scheme such as (51)(54), are expected to include the acidity of the hydride ligand. A highly acidic hydrogen atom may be more selectively transferred to the formaldehyde oxygen atom, producing the hydroxymethyl ligand as in (55) :
This appears to be the case for cobalt catalysts, which can hydroformylate formaldehyde to glycolaldehyde with high selectivity, apparently through the hydroxymethyl intermediate (75). The HCo(CO), hydride is known to be a strong acid, having a pK, < 2 (49).In contrast, reactions of HMn(CO), , a much weaker acid [pKa = 7 (49)],with formaldehyde under H 2 / C 0 have been found to produce instead the hydrogenated product methanol (1 79). The hydroxymethyl intermediate appears not to be formed, since under the identical conditions its ether and ester derivatives are converted in high yield (via CO insertion and hydrogenation) to ethylene glycol-containing products according to the following reaction (68, 179) : (co),M~--cH,oRJ!+
HOCH,CH,OR
(57)
It may be concluded that formaldehyde inserts into the less acidic manganese-hydrogen bond to form a methoxide ligand as shown in (56). As already mentioned, ruthenium catalysts also convert added formaldehyde into methanol rather than hydroformylated products. The H,Ru(CO), complex is not highly acidic (185) and is expected to have a pKa somewhat greater than that of H,Fe(CO), [pK, = 6.8 (186)].In this respect, H,Ru(CO), is anticipated to resemble HMn(CO), more than HCo(CO), . Reactions of added formaldehyde may differ somewhat from those of a possible coordinated formaldehyde intermediate generated by CO hydrogenation. It may be unnecessary for an added formaldehyde molecule to be coordinated to the metal before reacting with its hydride ligand. Such an addition could take place by an ionic or even a radical process. However, the trends in selectivity appear to be consistent in those systems which both
386
B. D. DOMBEK
reduce carbon monoxide to organic products and convert added formaldehyde; the cobalt and rhodium catalytic systems produced ethylene glycol (or glycolaldehyde) both from H,/CO and HCHO, and the ruthenium system produces only methanol from both. Slight quantitative variances in product selectivity from the two reactants have already been noted for the rhodium system (91), suggesting that there may indeed be small differences in the conversion mechanisms of bound and added formaldehyde. It is interesting to consider the function of carboxylic acids in promoting ethylene glycol ester formation by this system. Knifton has presented a thermodynamic argument, suggesting that the more favorable free energy of ethylene glycol ester formation, relative to that of free ethylene glycol, may be responsible for the formation of this product (171). However, the reactants and products in this system are not at equilibrium, and thermodynamics past the transition state of the product-determining step seem unlikely to be applicable to product selectivity. (As described in the next section, other ruthenium catalysts can produce ethylene glycol under similar conditions of temperature and pressure without the need for carboxylic acids.) Spectroscopic and chemical studies indicate that the ruthenium species present during catalysis in carboxylic acid solvents is Ru(CO), , as in other solvents (166).The chemical behavior, including the overall rate of CO reduction and the responses to reaction variables, is very similar to that observed in other solvents. The effect of carboxylic acids on product selectivity therefore appears to occur at a reaction stage after the rate-determining step of CO reduction. Remembering that the promoter function of a carboxylic acid is not a result only of its acidity, the scheme shown in (51)-(54) may be examined for intermediates which could interact with a carboxylic acid. One such intermediate is the hydroxymethyl complex of (51). If this intermediate were reversibly formed in small concentrations, it could be converted to the carboxylate ester by reaction with a carboxylic acid, as follows :
H
/-\
H
in a process similar to that of simple alcohol esterification. Hydroxymethyl complexes are normally quite unstable, apparently decomposing via a P-hydride shift from the oxygen atom (187),i.e., the reverse of hydroxymethyl formation in (58). However, carboxylate esters of the hydroxymethyl ligand are stable and not readily converted back to formaldehyde (68).Such an acyloxymethyl ligand would then be capable of undergoing CO insertion and hydrogenation to glycolaldehyde (ester) or ethylene glycol (ester) products, analogous to the reaction of (51). The carboxylic acid may also react
ETHYLENE GLYCOL AND ETHANOL FROM H2 AND CO
387
directly with the coordinated formaldehyde species, acylating it to form the ester without proceeding through a hydroxymethyl intermediate. Both of these reactions forming the acyloxymethyl ligand could be expected to exhibit the observed second-order dependence on acid concentration (166, 167). The direct acylation of a coordinated formaldehyde species has recently been observed (188). Reaction of the vanadium formaldehyde complex Cp2V(CH20)(Cp = qs-CsHj) with benzoyl chloride yields the 0
1I
Cp2(CI)V-CH20CC,H,
analog of the ruthenium ester shown in ( 5 8 ) . Formaldehyde added to ruthenium catalytic solutions in acetic acid is found to be converted, as expected, to ethylene glycol and methyl esters (179). Somewhat different selectivity is observed, however, from that found in standard H 2 / C 0 runs. Under conditions which give a C,/C2 ratio of 38 from CO hydrogenation, a reaction with added formaldehyde gave a C,/C2 ratio of 21 (179). As described above, slight differences in the reaction mechanisms involving bound and added formaldehyde could be responsible. Additionally, formaldehyde is known to react with carboxylic acids to form acetals and hemiacetals (189). These species could also react with ruthenium hydrides to form acyloxymethyl complexes without ever proceeding through hydroxymethyl or coordinated formaldehyde intermediates. The observation of glycerol triacetate as a trace product of CO hydrogenation by this ruthenium system in acetic acid solvent ( I 79) suggests that glycolaldehyde (ester) can undergo further chain growth by the process outlined in (26) for the cobalt system. As with formaldehyde, however, a carboxylic acid is apparently necessary to promote formation of the metalcarbon bonded intermediate which can produce the longer-chain product. Except for the modification of (58) in reactions (51) and (52), the scheme of (49)-(54) appears to apply to catalytic systems containing Ru(CO), in carboxylic acids and a variety of other polar and nonpolar solvents. As described in the next section, introduction of ionic promoters brings about significant changes in the catalyst chemistry. V.
Lewis Base-Promoted Ruthenium Catalysts
A.
BACKGROUND
The addition of certain ionic promoters to ruthenium catalytic solutions has been found to dramatically affect the rate and selectivity of CO hydrogenation. Whereas ruthenium solutions do not otherwise produce ethylene glycol as a significant product (except as its derivatives in in reactive solvents),
388
B. D. DOMBEK
ionic promoters can cause the formation of large amounts of this product in a variety of solvents. The species employed as a promoter need not be added in an ionic form, but it appears that the capability of forming ionic species under reaction conditions is essential. Halides are particularly effective as ionic promoters. Although the ruthenium species observed during CO reduction in the absence of promoters is Ru(CO),, its concentration can be reduced to unobservable levels by promoters which cause the formation of ionic ruthenium complexes. Because this system differs from unpromoted ruthenium catalysts in as many respects-rates, selectivities, catalytic species observed, and mechanism-it is addressed separately in this section. B. CATALYTIC ACTIVITY AND SELECTIVITY
Ruthenium complexes in nonreactive solvents such as sulfolane and NMP in the presence of halide promoters are found to possess high activity for the nium catalysts in many respects-rates, selectivities, catalytic species observed, and mechanism-it is addressed separately in this section. TABLE XV Hydrogenation o f C 0 by Halide-Promoted Ruthenium Catalysts
Expt.
Solvent
Notes”
1 2 3 4 5 6 7 8
Sulfolane NMP 18-Crown-6 Sulfolane 18-Crown-6 NMP Acetic acid Bu,PBr
b c
d e
f g
h i
Pressure (atm)
Temp. (“C)
Ethylene glycol rate (hr-’)
408 544 544 850 850 1020 430 430
200 250 250 1no 200 240 220 220
0.65 4.68 4.10 2.05 8.60 48.4 0.74 I .63
Methanol rate (hr-I)
Ethanol rate (hr-’)
2.86 36.8 141 1.71 30.4 384 6.46 19.8
0.29 4.7 1 I 8.4 0.35 1.51 24.5 0.70 5.63
HJCO = I, ruthenium source is Ru,(CO),,, unless otherwise specified. 75 ml solvent, 30 mmol Ru, 180 mmol KI, 2 hr (190). 75 ml solvent, 6 mmol Ru, 18 mmol KI. 0.77 hr (190). 75 ml solvent, 3 mmol Ru, 60 mmol KI, 0.63 hr, H,/CO = I .5 (190). 75 ml solvent, 30 mmol Ru, 180 mmol KI, 1.68 hr (190). 75 ml solvent, I5 mmol Ru. 60 mmol KI, 0.47 hr (190). 75 mol solvent, 6 mmol Ru, 120 mmol KI,0.13 hr (190). 50 g solvent, 3.75 mmol RuCI,. H,O, 37.5 mmol heptyl(triphenyl)phosphonium acetate, 18 hr (171). Products are acetate esters. 15 g tetrabutylphosphonium bromide “solvent,” 4 mmol RuO,. H,O, 2 hr (199).
ETHYLENE GLYCOL AND ETHANOL FROM H2 AND CO
389
glycerol, the ethylene glycol acetals of acetaldehyde, glycolaldehyde and formaldehyde, and small amounts of methane. Free acetaldehyde is also sometimes observed. Methanol and ethylene glycol are the major primary products; ethanol is largely a secondary product derived from methanol via homologation, through acetaldehyde (191). Reactions which are allowed to proceed for extended periods are thus found to produce higher relative yields of ethanol and other related secondary products. Increased product levels, particularly ethylene glycol, can reduce the rate of glycol production (192). Under such conditions, higher relative yields of glycol derivatives such as acetals and ethers may be observed. For these reasons, experiments designed to study chemical responses of this catalytic system were carried out using relatively short reaction times to avoid large contributions of secondary reactions and product interactions with the catalyst (191). Iodide-promoted reactions in phosphine oxide solvents have been observed under some conditions to produce ethanol from H2/C0 with good rates and high selectivities (193-195) (Table XVI, Expts. 1-3). Experimental evidence suggests that the ethanol is a secondary product, although its selectivity is high even after very short reaction times (193). An acid component is believed to be involved in alcohol homologation by this system, which will be described in more detail below. Related work has been reported in amide solvents with halide or hydrohalic acid promoters (196). Ethanol and acetaldehyde as well as methanol are observed. Enhanced yields of acetaldehyde appear to be obtainable by operating such a system at reduced temperatures, although overall rates of CO reduction suffer. Reactions of ruthenium catalyst precursors in carboxylic acid solvents with various salt promoters have also been described (170-172, 197) (Table XV, Expt. 7). For example, in acetic acid solvent containing acetate salts of quaternary phosphonium or cesium cations, ruthenium catalysts are reported to produce methyl acetate and smaller quantities of ethyl acetate and glycol acetates (170-172). Most of these reactions also include halide ions; the ruthenium catalyst precursor is almost invariably RuCl, . H 2 0 . The carboxylic acid is not a necessary component in these salt-promoted reactions : as shown above, nonreactive solvents containing salt promoters also allow production of ethylene glycol with similar or better rates and selectivities. The addition of a rhodium cocatalyst to salt-promoted ruthenium catalyst solutions in carboxylic acid solvents has been reported to increase the selectivity to the ethylene glycol product (198). Very similar reactions using a ruthenium catalyst, carboxylic acid solvent, and a slightly different promoter system have been reported (197) to give increased amounts of ethyl ester product (Table XVI, Expts. 4 and 5). Most examples show the use of RuO, * H,O as the catalyst precursor in a carboxylic
390
13. D. DOMBBK
acid containing a quaternary phosphonium bromide salt, under 430 atm of H,/CO at 220°C. The long reaction times reported (1 8 hr) and the observation of higher alcohol ester products suggest that secondary, alcohol homologation, processes are involved in the formation of the ethyl ester. Amounts of methane (calculated in approximate fashion based on reported typical levels in the vented gas, the free reactor volume, and the reaction temperature and pressure) found in these reactions appear to be quite high, TABLE XVI Eilianol Production Jrom HJCO by Ruthenium Catcilys/s"
Expt.
Solvent Pr PO Pr PO Pr,PO Propionic acid Propionic acid Bu,PBr Bu,PBr Bu,PBr
"
HJCO
=
Pressure Temp. (am) ( C)
408 408' 850 430 430 272 430 272
210 240 230 220 220 220 220 220
Promoter or cocatalyst
Ethanol rate (hr- I )
CH,(C,H,),PBr C,H I ,(C,H5),PBr TiO,(acac) Zr(acac), Co(acac),
6.6 17.4 14 2.4" 2.6' 2.6 2.6 2.3
12
4
I2
Ethanol C efficiency
("6)"
Notes
59 51 59 31 34' 31 31' 20
d e
./' 11
.I
k 1?1
n
I unless specified otherwise.
* Carbon efficiency to ethanol, defined as moles of CO converted
to ethanol divided by total moles of CO converted to organic products. ' H,/CO = 2. Other products and rates (hr- ')are methanol (4.0),n-propanol(O.25). n-butanol(O.1). ethylene glycol ( O . l ) , and methane (3.6)(194). ' Methanol ( l O . l ) , n-propanol(O.38), ethylene glycol (0.07).and methane (20)(194). Methanol ( 5 . 5 ) , n-propanol (0.87), n-butanol (0.30), ethylene glycol ( I .8), and methane (6.9) (193). Product is ethyl propionate. Methyl propionate (2.58),n-propyl propionate (0.43),n-butyl propionate (0.04).and methane (6.84)(197). ' Assuming reported "typical" methane levels are formed in this experiment ; ethanol carbon efficiency among liquid products is 63%. Methyl propionate (1.3), n-propyl propionate (0.5).n-butyl propionate (0.04). and ethylene glycol dipropionate (0.04)(197). Methanol ( I .47),n-propanol (0.21).n-butanol (0.17).methyl acetate (0.14).ethyl acetate plus methyl propionate (0.26).propyl acetate plus ethyl propionate (0.13), and methane (6.67)(204). ' Assuming reported "typical" methane levels are formed in this experiment; ethanol carbon efficiency among liquid products is 52%. Methanol ( I .07,n-propanol (0.30).n-butanol (0.22),methyl acetate (0.08), ethyl acetate plus methyl propionate (0.36).and n-propyl acetate plus ethyl propionate (0.08) (204). " Methanol ( I .17).n-propanol(O.83), n-butanol (0.10), methyl acetate (0.29).ethyl acetate (0.83). n-propyl acetate (0.35).and methane (8.22)(203).
ETHYLENE GLYCOL AND ETHANOL FROM H 2 AND CO
39 1
and selectivities to the ethyl ester are