Advances in Catalysis, Volume 28

ADVANCES I N CATALYSIS VOLUME 28 Advisory Board G. K. BORESKOV Novosibirsk, U.S.S.R. M. BOUDART M. CALVIN Berkelzy...

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ADVANCES I N CATALYSIS

VOLUME 28

Advisory Board G. K. BORESKOV Novosibirsk, U.S.S.R.

M. BOUDART

M. CALVIN Berkelzy, California

Stanford, Calijornia

P. H. EMMETT

A. OZAKI

G.- M. SCHWAB

Portlund, Oregon

Tok90, Japan

Munich, Germany

G. A. SOMORJAI Berkeley, Calfornaa

R. Uco Milan, Italy

ADVANCES IN CATALYSIS VOLUME 28

Edited by

D. D. ELEY The University Nottingham, England

HERMAN PINES Northwestern University Evanston, Illinois

PAULB. WEEZ Mobil Research and Development Corporation Princeton, New Jersey

1979 ACADEMIC PRESS A Subsidiary of Harcourt Brace Jovanovich, Publishers

New York

London Toronto Sydney San Francisco

COPYRIGHT @ 1979, BY ACADEMIC PRESS, INC. ALL RIGHTS RESERVED. NO PART O F THIS PUBLICATION MAY BE REPRODUCED OR TRANSMITTED IN ANY FORM OR BY ANY MEANS, ELECTRONIC OR MECHANICAL, INCLUDING PHOTOCOPY, RECORDING, OR ANY INFORMATION STORAGE AND RETRIEVAL SYSTEM, WITHOUT PERMISSION IN WRITING FROM THE PUBLISHER.

ACADEMIC PRESS,INC.

111 Fifth Avenue, New York, New

York 10003

United Kingdom Edition published b y ACADEMIC PRESS, INC. (LONDON) LTD. 24/28 Oval Road, London N W 1 7 D X

LIBRARY OF CONGRESS CATALOG CARDNUMBER: 49-7755 ISBN 0-12-007828-7 PRINTED IN THE UNITED STATES OF AMERICA 79 80 81 82

9 8 7 6 5 4 3 2 1

CONTRIBUTOS............................................................. ................................................................... PREFACE

vii ix

Elementary Steps in the Catalytic Oxidation of Carbon Monoxide on Platinum Metals T. ENGEL AND G . ERTL I. 11. 111. IV.

Introduction

.

Adsorption of Oxygen

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

\I.

VI. VII. VIII. IX.

Oxidation of CO on Iridium Oxidation of CO on Rutheniu Factors Influencing the Catalytic Activity . . . . . . . . . . . References.. . . . ...........................................

24 39 51 59 63 65 65

73

The Binding and Activation of Carbon Monoxide, Carbon Dioxide, and Nitric Oxide and Their Homogeneously Catalyzed Reactions RICHARD EISENBERG AND DANE. HENDRIKSEN 1.

11. 111. 1V.

Introduction . . . . . ................. Carbon Monoxide Carbon Dioxide ................................................. Nitric Oxide ....... ................. References ..................................................

79 119 144 164

The Kinetics of Some Industrial Heterogeneous Catalytic Reactions M. I. TEMKIN I. 11. 111. I\?. 17.

Introduction .................................................... The Measurement of Reaction R Mass Transfer in Heterogeneous Ideal Adsorbed Layers ........................................... The Routes of Complex Reactions V

173 184

vi

CONTENTS

Vl . 1711 .

VIII.

IX . X.

XI . XI1 . XI11 . XIV .

xv .

XVI . XVII . XVIII . XIX .

xx.

XXl .

The Steady-State Conditions ...................................... Some General Relations for Steady-State Reactions . . . . . . . . . . . . . . . . . Reversible Many-Stage Reactions .................................. Nonuniform Surfaces ............................................ Adsorption Equilibrium and the Kinetics of Reaching it on Nonuniform Surfaces ............................................ The Kinetics of Reactions on Nonuniform Surfaces . . . . . . . . . . . . . Oxidation of Ethylene into Ethylene Oxide ........................ Hydroxylamine Synthesis ........................................ Reaction of Methane with Steam .................................. Ammonia Synthesis-Simple Kinetics ............................. Ammonia Synthesis-Complicated Kinetics . . . . . . . . .... Carbon Monoxide Conversion .................................... Isotopic Exchange between Steam and Hydrogen . . . . . . . . . . . . . . . . . . Phosgene Synthesis .............................................. Reactions of Carbon with Carbon Dioxide and Steam ............... Ammonia Oxidation ........... .. ... References ......................................................

192 197 203 207 213 223 230 239 244 250 257 263 267 270 273 279 287

Metal-Catalyzed Dehydrocyclizationof Alkylaromatics SIGMUND M . CSICSERY

I. I1 .

111. IV . V.

v1. Vll . VIII .

............ Introduction ............... The Dehydrocyclization of C , enzenes ....................... The Dehydrocyclization of Clo and Higher Alkylbenzenes ........... Double Cyclization of CBand Higher Paraffins ..................... The Dehydrocyclization of Alkylbenzenes Over Chromia-Alumina Catalysts .................................. The Dehydrocyclization of Alkylnaphthalenes ....................... Dehydrocyclization of Diphenylalkanes ...... ............... Conclusions ..................................................... References ...................... ................... ....

293295 296 312

I

314 315 318 319 319

Metalloenzyme Catalysis JOSEPH

I. I1. Ill . I V.

J . VILLAFRANCA AND FRANK M . RAUSHEL

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

...... 324

Therrnolysin .................................................... Yeast Hexokinase ............................................... Glutamine Synthetase ..................................... References ......................................................

AUTHOR INDEX .............................................................

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

'326 336 349 366 371 387 397

Contributors Numbers an parentheses andzcate thp pages on whach the authors' contrabutaons b e p

SIGMUND M . CSICSERY, Chevron Research Company, Richmond, California 94802 (293) RICHARD EISENBERG, Department of Chemistry, University of Rochester, Rochester, New York 14627 (79) T. ENGEL,* Institutf u r Physikalische Chemie, Universitat Miinchen, 8 Munchen 2, West Germany ( 1 ) G. ERTL,Institut f u r Physikalische Chemie, Universitat Miinchen, 8 Munchen 2, West Germany ( 1 ) DANE . HENDRIKSEN, Corporate Research Laboratories, Exxon Research and Engineering Company, Linden, New Jersey 07036 (79) FRANK M . RAUSHEL,Department of Chemistry, Pennsylvania State University, University Park, Pennsylvania I6802 (323) M . I . TEMKIN, Karpov Institute of Physical Chemistry, Ul. Obukha 10,107120, Moscow B-120, U S S R (173) JOSEPH J . VILLAFRANCA, Department of Chemistry, Pennsylvania State University, University Park, Pennsylvania I6802 (323)

*Present address: IBM Research Laboratory, Zurich CH-8803 Riischlikon, Switzerland. vii

This Page Intentionally Left Blank

The five chapters in this volume cover the whole spectrum of modern catalysis research. Two deal with industrial catalysis. Professor Temkin’s chapter reviews his thorough kinetic studies of several reactions which, apart from their great theoretical interest, have found applications in the USSR in process optimization and reactor design. Here we have the applied mathematics of catalysis. Dr. Csicsery’s article examines the value that studies of molecular mechanisms have had in the development of dehydrocyclization, a subject extensively studied in the USSR, United Kingdom, and United States. The article by Dr. Engel and Professor Ertl is complementary to that of Professor Eisenberg and Dr. Hendriksen, the first dealing with the heterogeneous catalysis of carbon monoxide and the second with its homogeneous catalysis, together with that of carbon dioxide and nitric oxide. Carbon monoxide over the years has proved to be one of the most self-revealing of molecules to surface spectroscopic techniques. Couple this with its importance for exhaust gas studies and one can easily appreciate why it has attracted so much attention, from the infrared studies of Eischens in the 1950s to the present. Finally, the article by Professor Villafranca and Dr. Raushel presents a fascinating account of metalloenzyme catalysis. It is one of the paradoxes of modern science that, once the original x-ray crystallographic problem was solved by Perutz, Kendrew, and their colleagues, the very complexity of enzyme molecules has enabled very rapid progress toward the definition of the activated complexes involved. However, we may expect heterogeneous catalysis to catch up soon, as advances in knowledge due to surface spectroscopy and homogeneous catalysis analogies play their part. Engel and Ertl’s article already provides evidence for this progress. The tendency for enzyme catalysis to run neck and neck or even lead heterogeneous catalysis goes back one hundred years. Indeed, in 1890 Cornelius O’Sullivan and Frederick W. Tompson had already produced a fairly clear idea of intermediate complex formation between “invertase” and sucrose, as a result of their careful kinetic investigations ix

X

PREFACE

u. Chem. SOC.Trans. 57,834 (1890)-the

fact that these two scientists worked at Bass’s Brewery at Burton-on-Trent, United Kingdom, may also endear their memory to workers in the catalysis field]. The labile character of intermediate complexes is really what catalysis is all about. D. D. ELEY



ADVANCES IN CATALYSIS. VOLUME 28

Elementary Steps in the Catalytic Oxidation of Carbon Monoxide on Platinum Metals T. ENGEL* and G . ERTL Institui fur Physikalische Chemie Universiia i Miinchen Miinchen. West Germany

I . Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . I1. Adsorption of Carbon Monoxide . . . . . . . . . . . . . . . . . .

I11.

IV .

V.

VI . VII . VIII . IX .

2 2 A. Mechanism of Bond Formation . . . . . . . . . . . . . . . . . 3 B. Surface Structure . . . . . . . . . . . . . . . . . . . . . . . 6 C. Energetics of Adsorption . . . . . . . . . . . . . . . . . . . . 14 D . Kinetics of Adsorption and Desorption . . . . . . . . . . . . . . 19 Adsorption of Oxygen . . . . . . . . . . . . . . . . . . . . . . . 24 A . Characterization of the Surface Species . . . . . . . . . . . . . . 24 B. The Nature of the Chemisorption Bond . . . . . . . . . . . . . . 28 C . Surface Structure . . . . . . . . . . . . . . . . . . . . . . . 30 D . Surface Mobility . . . . . . . . . . . . . . . . . . . . . . . 34 E. Kinetics of Adsorption . . . . . . . . . . . . . . . . . . . . . 34 F . .Kinetics of Desorption and Adsorption Energies . . . . . . . . . . 35 Oxidation of CO on Palladium . . . . . . . . . . . . . . . . . . . 39 A . Surface Interaction between Oxygen and Carbon Monoxide . . . . . 40 B . Reaction Mechanism . . . . . . . . . . . . . . . . . . . . . . 42 C. Reaction Kinetics . . . . . . . . . . . . . . . . . . . . . . . . 46 Oxidation of CO on Platinum . . . . . . . . . . . . . . . . . . . 51 Oxidation of CO on Iridium . . . . . . . . . . . . . . . . . . . . 59 Oxidation of CO on Ruthenium . . . . . . . . . . . . . . . . . . . 63 Oxidation of CO on Rhodium . . . . . . . . . . . . . . . . . . . 65 Factors Influencing the Catalytic Activity . . . . . . . . . . . . . . 65 A . Influence of the Surface Structure . . . . . . . . . . . . . . . . 65 B. The Role of Surface Impurities . . . . . . . . . . . . . . . . . 69 C . ThePressureGap . . . . . . . . . . . . . . . . . . . . . . . 71 References . . . . . . . . . . . . . . . . . . . . . . . . . . . . 73

* Present address: IBM Research Laboratory.

Zurich. CH-8803 Ruschlikon. Switzerland.

1 Copyright 0 1979 by Academic Press. Inc. All rights of reproduction in any form resewed . ISBN 0-1 2-007828-7

2

T. ENGEL AND G . ERTL

1.

Introduction

Catalytic oxidation of carbon monoxide over catalysts from the platinum group metals has been investigated in numerous studies dating back to the classic work of Langmuir (1). Apart from its enormous practical importance, this reaction is considered to proceed through a relatively simple mechanism since only diatomic molecules are involved and product formation occurs presumably only over a very few steps. Nevertheless, only about 10 years ago firm knowledge about the microscopic elementary processes of this reaction was very limited. In the interim, a whole series of powerful surface-sensitive spectroscopic techniques was developed that yielded direct information on the chemical, electronic, energetic, and dynamic properties of clean and adsorbate covered surfaces that are kept under well-defined conditions. The oxidation of CO over a Pd(ll0) single-crystal surface was the first example of a catalytic reaction that was studied under these aspects (2) and, in the interim, many investigators with similar model systems have increased our knowledge of this reaction enormously. It is thus believed to be, at present, the best understood catalytic reaction, although a series of open questions still needs clarification by future work. This article is concerned with a discussion of the adsorptive properties of the reactants, their mutual interaction, and finally the mechanism and kinetics of product formation, whereby the primary emphasis will be given to investigations with welldefined single-crystal surfaces. The conditions applied with such model studies are very far away from those found with “real” catalysis. This often discussed “gap” concerns the surface structure and cleanliness as well as the applied gas pressures. These aspects will be briefly discussed in the last section, where it will be shown that in the present case the situation is rather favorable, so that firm conclusions for real catalysis can also be drawn from the results obtained for model systems by methods of “surface science.” II. Adsorption of Carbon Monoxide

Carbon monoxide is the molecule whose adsorptive properties have certainly been investigated in most detail. A review of the work prior to 1970 was published by Ford (3)in this series, but in the interim so much additional information has been obtained that a complete compilation of the data would be far beyond the scope of this article. Instead, the main emphasis will be given to a discussion of essential features that fortunately have much in common with different platinum group metals as well as with various crystal planes. Adsorption normally takes place in molecular form, although sometimes

CATALYTIC OXIDATION OF CO ON PLATINUM METALS

3

some tendency for dissociation is observed. However, since there is no evidence for the participation of dissociatively adsorbed CO in the catalytic C 0 2 formation, this spurious effect will not be considered further. Pd, Pt, and Ir are the most frequently studied metals. Less information is available for Ru and Rh, and so far no results have been published for 0 s .

A. MECHANISM OF BONDFORMATION It is commonly accepted that chemisorption of CO on transition metals takes place in a way that is quite similar to bond formation in metal carbonyls (4). First experimental evidence for this assumption was obtained from a comparison of the C - 0 stretching frequencies (5) and was later confirmed by data on the bond strength (6) as well as by valence and core level ionization potentials obtained by photoelectron spectroscopy (7). Recent investigations have in fact shown that polynuclear carbonyl compounds with more than about 3-4 metal atoms exhibit electronic properties that are practically identical to those of corresponding CO chemisorption systems (8, 9), thus supporting the idea that the bond is relatively strongly localized to a small number of metal atoms forming the chemisorption site. Accordingly, binding occurs through a concerted electron transfer from the highest filled molecular orbital of CO (50, which is predominantly a lone electron pair at the carbon end) to the metal and back-donation of metallic d electrons to the lowest unfilled CO orbital (the antibonding 27r* level), as illustrated by Fig. 1. Accordingly, the energetic separation between the COderived 5~ and In levels should become smaller or eventually the 50 orbital should be pulled down below the In: level. This model for the chemisorption bond was first proposed by Blyholder (lo),and later confirmed by different theoretical models (11-29). Although most of these calculations were performed for nickel, there is experimental as well as theoretical (22,23)evidence that the platinum metals exhibit similar behavior. The most detailed experimental information on the electronic properties of adsorbates is obtained by ultraviolet photoelectron spectroscopy (UPS) (7) and a great amount of spectral data has been published for adsorption of CO on Pd (30-33), Pt

FIG.1. Schematic representation of the donor-acceptor model for CO adsorption.

4


E,=O

2

C 6 6 Energy bdw E,

10

12 [el

FIG.2. UPS data for the adsorption of CO on Pd(l10). (a) clean surface, (b) with CO adsorption, and (c) difference spectra, b - a (31).

(34-38), Ir (39-46), Ru (47), and Rh (48)surfaces. A characteristic spectrum is reproduced in Fig. 2. All systems exhibit a relatively broad maximum extending from the Fermi level (EF)to about 6-8 eV below EFwhich arises from emission from the metallic d states. Additional maxima arise at approximately -8 and - 11 eV, which originate from the CO-derived (50 + 171) and 40 levels. In the case of Ir, the former is clearly composed from two contributions, indicating that in this case the energetic separation between the 50 and In levels is presumably somewhat larger than with the other metals. Angular resolved spectra for CO/Ni( 100) (49) as well as for CO/Pd( 1 1 1) (33) were interpreted in a way that the 50 level has even a slightly larger ionization energy than the 171 level that is mainly unperturbed by chemisorption. This conclusion would be in agreement with cluster calculations (15,23) for these systems, but further results are needed for a general solution of this question. In addition, it must be pointed out that an evaluation of the “bonding shift” (i.e., the lowering of the 5 0 orbital due to formation of the chemisorption bond) from the observed photoelectron spectra is rather questionable owing

5


to other different effects that may influence the experimentally accessible ionization energies (21). The UPS spectra exhibit a strong damping of the emission from the d states after adsorption of CO, an effect that is mainly caused by an antiresonance phenomenon of the emitted electrons rather than by a true depletion of d states (22). For this reason the chemisorption level derived from the “back-donation” effect, which is predicted to be just below E F ,can never be clearly identified from photoelectron spectroscopic results. On the other hand, the energy distribution of electrons emitted from surfaces under the action of a high electric field (FEED spectroscopy) strongly suppresses contributions from d states in favor of sp-like states and the so-called enhancement factor R(E).which is equal to the ratio of the energy distribution from the adsorbate covered to that of the clean surface, is simply proportional to the local density of (occupied) states on the adsorbate pa(&),provided that the clean surface exhibits a free-electron-like behavior (50). Recent FEED measurements (51) with the system CO/Ir( 100) revealed the formation of a state at 0.6 eV below EF that exhibits at saturation coverage an additional shoulder at about - 1.25 eV as reproduced in Fig. 3. This effect is ascribed to the formation of the discussed back-donation state that predominantly exhibits the character of metallic wave functions, but is responsible for partial transfer of electronic charge to the adsorbate. Semiempirical calculations for the system COiPd (which is similar to the CO/Ir system) predicted the formation of such a state at 0.7 resp. 0.2 eV below EF (16, 22). The shoulder at - 1.25 eV probably arises from the occupation of different geometric sites at high surface concentrations, since adsorbed

n I

n‘2p

-.IS

\

4 \ I

- 40

\

I \

\

1 \

1 \

-*S

FIG. 3. (a) Enhancement factor R(E)as a function of the energy relative to the Fermi level for field emission from a CO/Ir(lOO). (b) Schematic representation of the CO levels that give rise to the enhanced emission just below the Fermi level (51).

6


CO shows a general tendency for the formation of close-packed arrangements near saturation as will be outlined in the next section. Indirect experimental evidence for the partial occupation of the antibonding 2n* orbital of CO upon adsorption is given from measurements of the variation of the work function, A q , as well as from data on the C-0 vibration frequency. With Pd (52), Ru (5.9, and Ir (54-56) surfaces, the work function was found to increase upon adsorption of CO. indicating a net transfer of electronic charge from the metal to the adsorbate. With different Pd single crystal faces, for example, dipole moments of the adsorbate complexes ranging from 0.29 to 0.35 D have been determined (52) that are considerably larger than the dipole moment of free CO [O. 11 D (531. The only exception known thus far is presented by the system CO/Pt( 11 1) (58, 59). With increasing coverage the work function at first decreases, passes through a minimum at 0 = 1/3 and reaches nearly the initial value at higher coverages. This effect resulted from the occupation of two different types of sites on the surface with dipole moments p , = 0.2 D and pz z 0 D (58). It might be caused by a smaller degree of electronic charge transfer from the metal to the adsorbate, although the adsorption energy and the C-0 vibration frequencies are quite similar to the data obtained with the other metals under discussion. A further consequence of the back-donation effect consists of a weakening of the C-0 bond strength that manifests itself in a lowering of the C-0 stretching frequency as compared with the value of free CO ( = 2143 cm- '). Vibration spectra from CO adsorbed on single crystal surfaces have thus far been obtained with Pd (60), Pt (59, 61-64), and Ru (65). In all cases bands below 2100 cm-' were observed. Some data for Pd are included in Table I and a more detailed discussion will be presented in the following two sections.

B. SURFACE STRUCTURE I . General

Numerous studies with low-energy electron diffraction (LEED) revealed that most of the clean surfaces of the platinum group metals exhibit an atomic arrangement that is identical to that expected from an undistorted termination of the bulk. Variations of the vertical lattice spacings between the topmost atomic layers are very small, if present at all (66). Exceptions are, however, found with the (100) and (1 10) planes of Ir and Pt. The clean and thermodynamically stable structures of the Pt( 100) (67--69) and Ir( 100) (70, 71) surfaces were found to reconstruct and to exhibit 5 x 1-LEED patterns. A plausible explanation (72) is that in these cases the topmost atomic layer forms a hexagonal arrangement, similar to that within the (1 1 1)

7


TABLE I Chemisorption of CO on Pd Single Crystal Planes''.h

Density of surface atoms x (cm-2) Initial isosteric heat of adsorption (kcal/moi) Maximum increase of the work function A q (eV) Dipole moment of adsorbed CO (Debye) Maximum density of adsorbed molecules x (cm-') C-0 stretch frequency at low coverage (cm- ') Adsorbate structures with increasing coverage

1.33

1.53

34

36.5

0.94

40

0.80

35.5

0.53

35

0.98

0.75

I .26

0.33

0.29

0.35

0.33

1.0

0.94

0.89

I .o

1823

8x

,/3R30 (0 = 0.33)

1 compression

1 c4 x 2 (0 = 0.5)

1

-

I895

c 4 x 2 R45 (0 = 0.5)

1.27

I .06

1878

c2 x 2 (0 = 0.5)

1 x I (0 = I )

1

1

1

compression (0 = 0.7)

4 x 2 (0 = 0.75)

1 x 2

1 2 x 1 (0 = 1.0

one-dimensional disordered 2 x 1

1 one-dimensional disordered 3 x 1

compression (0 = 0.6)

1 hexagonal (0 = 0.63)

1 hexagonal (0 = 0.66) "

H. Conrad, H. Ertl, J. Koch. and E. E. Latta. Surface Sci. 43,462 (1974).

* A. M. Bradshaw and F. M. Hoffmann. Surfnce Sci. 72, 513 (1978).

plane. More sophisticated models have recently been proposed by Miiller et al. ( 7 3 , but final clarification can only be obtained from a complete dynamic analysis of the LEED intensities. Similarly, the clean Pt(ll0) (74) and Ir( 110) (54, 75, 76) surfaces are reconstructed into 1 x 2 structures. In

the latter case a dynamical LEED analysis revealed evidence that every

8

T. ENGEL AND G. ERTL

second atomic row on the surface is missing, thus producing this periodicity (77). It is interesting to note that the surface reconstruction may be. removed by properly chosen reaction cycles so that (metastable) clean surfaces with the ideal 1 x 1 structure are prepared that may exhibit characteristic differences as compared to the corresponding reconstructed phases (56, 69, 78). It is generally accepted that the CO molecule is chemisorbed through the carbon atom, the molecular axis being perpendicular to the surface plane. This conclusion is supported by the following evidence : i. The analogy with transition metal carbonyls where thus far almost no evidence exists for bonds of the type M-CO-M or M-OC (isocarbonyl). ii. Model calculations that always reveal that the M-CO bond is energetically most favorable (12, 79). iii. Ion scattering experiments from CO/Ni( 1 11) revealed that the noble gas ions “see” almost exclusively only oxygen atoms which therefore form the outmost atomic layer (80). iv. Angular resolved photoelectron spectroscopy. The differential photoionization cross section for electrons ejected from the 40 orbital of an oriented CO molecule has been calculated by Davenport (81). Angular resolved photoelectron spectra at various photon energies were recorded for CO/Ni(l00) by Allyn et al. (82) that could show conclusively that in this case the perpendicular configuration of the adsorbate on the surface exists and confirmed conclusions from previous measurements with CO adsorbed on Ni(ll1) (36), Pt(lI1) (35),Ru(OO1) (83),and Ir(100) (84),although in the latter case the experimental evidence is not yet completely clear. v. LEED: Quite recently a LEED intensity analysis was performed with the system CO/Pd(100) that confirmed the perpendicular orientation of the CO molecule that is attached to the surface through the C atom (85). It must be mentioned that a LEED analysis for CO/Ni(100) led to the conclusion that in this case the molecular axis is tilted (86),which is, however, in disagreement with the UPS results obtained with the same system (82).

Although thus far only one complete structural analysis from CO overlayers on the platinum metals has been performed, in many cases the LEED patterns reveal direct information on the mutual surface configurations of the adsorbates. In a few cases even the geometry of the adsorption sites is known with a great degree of certainty by including vibrational spectroscopic data into the discussion. The general features are quite similar for all systems. The adsorbed CO molecules form simple overlayer structures up to medium coverages and are presumably located in identical, highly symmetrical sites. At higher coverages compression of the overlayer unit cells takes place and close-packed arrangements are formed. The maximum density of adsorbed particles is always approximately 1 x 10’’ cm-’ and is determined by the


9

effective diameter of the CO molecule ( - 3 A). The same crystallographic planes of the different metals exhibit usually quite similar overlayer structures and differences arise mainly at high coverages owing to the slightly varying lattice constants. The LEED patterns observed with various Pd planes are included in Table I and the following discussion will be mainly concentrated on this metal. 2. (111)Plane On Pd( 11 1) a

8x

@ R 30" structure is fully developed at 6 = 1/3 (60,

87,88). The C-0 stretch frequency of 1823 cm-' observed at low coverages

indicates that the molecules are most probably located in threefold coordinated sites (60)as indicated by the structure model of Fig. 4a. Upon further increasing the coverage the overlayer unit cell is continuously compressed, as indicated by the arrows in Fig. 4a, until at 8 = 0.5 a c4 x 2 structure is reached. This compression is connected with an enormous shift of the IR band that suggests that the CO molecules now move into bridge sites as drawn in Fig. 4b (60).The compression continues until 6 = 0.6. Additional

(C)

(d)

4

FIG.4. Structure models for the adsorption of CO on Pd(ll1). (a) x J j R 30" structure (threefold sites), (b) c 4 x 2 structure (bridge sites), (c) 0 = 0.63, and (d) 0 = 0.66 (88).

10


uptake of CO causes the formation of two structures with hexagonal arrangements of the adsorbed particles with 8 = 0.63 (Fig. &) and 0 = 0.66 (Fig. 4d) (88). Beyond 0 = 0.5 the mutual repulsion between adsorbed molecules becomes already so strong that additional adsorption can only be observed far below room temperature. The highest coverage is associated with an adsorbed particle density of 1.01 x 10’’ cm-2 and a CO-CO distance of 3.38 A. The structure models reveal that above 8 = 1/3 the CO molecules are relatively easily displaced from their highly symmetric sites and form partly incoherent structures, a point that will be further discussed in the following section. The IR spectra indicate that even at relatively low coverages sites with different geometry may be occupied, indicating some degree of disorder in the overlayer (60). Identical x $ R 30” structures at 8 = 1/3 were found with all the hexagonal planes of the other metals: Pt(ll1) (58), Rh(l11) (48, 89), Ir(ll1) (40, 90, 91), and Ru(0001) (53, 92). Above 8 = 1/3 the situation is slightly different, but adsorption leads always to the formation of close-packed layers. On Pt(ll1) compression into a c4 x 2 structure occurs and finally an incoherent hexagonal layer is formed (58). On Rh(l1 I) a 2 x 2 structure with hexagonal symmetry is formed (89) (Fig. 5). On Ir(ll1) transformation into a 2 $ x 2 $ R 30” structure and its further compression is observed (40, 90, 9 4 , and on Ru(OO1) a 2 x 2 structure similar to Rh(ll1) is formed (53). The LEED patterns observed with CO/Pt(l 1 1) exhibit a rather high degree of disorder at room temperature, which indicates the existence of nearly equivalent sites. From vibrational spectroscopic data it was concluded that, in contrast with the situation shown for Pd(ll1) in Fig. 4a, at 0 = 1/3 “on top” positions rather than threefold coordinate sites are preferentially occupied, whereas at 0 = 1/2 a mixture of on top and “bridge” bound molecules exist (62,63). Since, however, with this system the work function change exhibits an anomalous behavior (as out-

8

- - w

FIG.5. Structural model for the 2 x 2 pattern observed for CO adsorption on Rh(l11) (89).


11

lined above) that suggests a smaller degree of electronic back-donation than with other metals, it appears to be questionable if in this case the normal assignment of C-0 stretch frequency with the type of adsorption site is applicable.

3. (100) Plane Adsorption of CO on Pd(100) leads to the formation of a c4 x 2 R 45" structure at 8 = 0.5 which is unique in so far as all the adsorbed molecules may be placed into equivalent positions only with the occupation of bridge sites (93-95). In fact, IR spectroscopic investigations revealed only the existence of a single species at this coverage with ? = 1895 cm-' at very low 8 that continuously increases to 1949 cm-' at 8 = 0.5. The latter effect is correlated with the variation of the adsorption energy with coverage as will be discussed in the following section. The structure model reproduced in Fig. 6, as proposed originally by Park and Madden ( 9 3 , can therefore be regarded as being well established even without a LEED intensity analysis. It must be pointed out that two equivalent domain orientations that are rotated by 90" with respect to each other exist on the surface. The validity of this structure model was recently completely confirmed by a LEED intensity analysis (85). The derived atomic distances are included in Fig. 6 and are quite comparable with the structural data of transitions metal carbonyls. Beyond 8 = 0.5 the unit cell of the adsorbate layer is again continuously compressed along the direction indicated by the arrows in Fig. 6 until saturation at 8 = 0.8 with a mutual CO-CO distance of about 3.1 A is reached (94, 95). The situation with the other metals is somewhat different. On Rh(100) at 8 = 0.5 a c2 x 2 structure is formed where the adsorbed molecules have been proposed to be located in fourfold sites as illustrated by Fig. 7a (89). Saturation of the adlayer is characterized by a "split" 2 x 1 structure that is repro-

FIG. 6. Structural model for the adsorption of CO on Pd(100) showing the arrangement and bond distance involved (52,85).

12


ww FIG.7. Structural models for the adsorption of CO on Rh(100) showing (a) the c2 x 2 and (b) the split 2 x 1 structures(89).

duced in Fig. 7b (89). This exhibits a hexagonal arrangement with a CO diameter of about 3.2 A. With Ir and Pt the situation is more complex owing to the existence of two structurally different modifications of the clean surfaces. With Ir(100) 5 x 1 as well as with Ir( 100) 1 x 1 formation of c2 x 2overlayer structures was observed (96). Such a structure was also reported for adsorption of CO on a Pt( 100) 1 x 1 surface (69). However, in addition, a 2 x 3 f i R 45" structure at 8 = 0.66 and a c4 x 2 structure with 8 = 0.85 corresponding to a density of adsorbed particles of 0.98 x 1015 cm-' were also reported (97). On the other hand, with Pt(100) 5 x 1 Kneringer and Netzer (98) observed a fairly rapid transformation into the unreconstructed 1 x 1 pattern upon CO admission at room temperature which was also followed by the formation of a c4 x 2 structure and, under special conditions, of a 2 x 2 structure.

4. (110) Ptane On Pd(ll0) a c2 x 2 structure is formed at 0 = 0.5. The CO molecules are presumably located in fourfold sites as shown by Fig. 8a. Further adsorption leads to the formation of a 4 x 2 structure at 0 = 0.75 where the adsorbed particles form incoherent chains along the troughs of this plane (Fig. 8b). Saturation is characterized by a 2 x 1 structure in which some characteristic spots are missing (52), which was interpreted (99) as arising from the existence of glide lines within the overlayer arrangement. The result-


t lo011

13

t loo11

FIG. 8. Structural models for the adsorption of CO on Pd(ll0) showing (a) the c2 x 2 structure (0 = 0.5) and (b) the 4 x 2 structure (0 = 0.75) (52).

ing structure model corresponds to 0 = 1. It is interesting to notice that in this case the CO molecules do not occupy the sites of highest symmetry. It is tempting to speculate if the molecular axis is still perpendicular to the surface plane or, rather, tilted toward the troughs formed by the metal atoms of the second layer. A structure of the latter type was also observed with Ir( 110) ( 5 4 7 6 ) and Pt( 110) surfaces (74).There is some indication that with Ir( 110) some oxygen adsorption on the surface is needed for ordering of this structure (76). At lower coverages 2 x 2 and 4 x 2 structures were observed with this plane (76), the former of which is not present on Pd( 1 10) and corresponds to 8 = 0.25. 5. Other Planes

Only very few investigations have been performed with surfaces not belonging to the most densely packed planes. On Pd(210) the formation of 1 x 1- and 1 x 2-overlayer structures were observed for which the proposed structure models are reproduced in Fig. 9 (52, 60). For the 1 x 1 structure a C - 0 stretch frequency of 1878 cm-' was observed which is, interestingly,

(a 1 (b) FIG.9. Structure models for the adsorption of CO on Pd(210) showing (a) the 1

ture and (b) the 1 x 2 structure (60).

x 1 struc-

14


[OIOI

FIG.10. Structural model for the adsorption of CO on Ru(lO1) (100).

rather similar to the value found for Pd( 100). The initial adsorption energies are also similar (52).From this it was concluded that also in the present case the CO molecules occupy bridge sites which, however, in this case are formed by atoms from the first and second layer, which would necessarily imply that the molecular axis is tilted. On Pd(3 1 1) one-dimensional disordered structures with periodicities of 2a, and 3a,, respectively, in the [Ol 11 direction were observed (52). An interesting study was performed by Reed et al. (ZOO) with a Ru(lO1) surface for which an ordered overlayer was detected whose proposed structure model is reproduced in Fig. 10.

c. ENERGETICS OF ADSORPTION Table I1 lists a series of values for the adsorption energy of CO on the three most densely packed planes of the platinum metals at low coverages; i.e., at surface concentrations that are believed to be small enough to rule out noticeable effects on these energies owing to mutual interactions between the adsorbates or the occupation of different adsorption sites. These values are therefore believed to represent the strength of the proper metal-CO bond with the adsorbate in its energetically most favorable position on the surface. These data were obtained in two different ways : i. Through equilibrium measurements by applying the Clausius-Clapeyron equation to the analysis of adsorption isotherms. This procedure yields the isosteric heat of adsorption and yields reliable values to within f 1 kcal/mol since it is free from any additional assumptions. ii. From an analysis of thermal desorption spectra (TDS). This procedure is more convenient, but needs certain assumptions on the reaction order of desorption (which can safely be taken as one) and on the preexponential. The resulting quantity is the activation energy for desorption that should be equal to the adsorption energy since the kinetics of CO adsorption is nonactivated. The weak point of the analysis is the unknown preexponential, which is frequently assumed to be 1013 sec-'. It could, however, be shown


15

TABLE I1 Adsorption Energies for C O at Low Coverages Plane

(111)

Ru Rh Pd Ir Pt

29".' 31' 34f 34' 33'

(100)

(110)

-

26b.d

29' 36.5g 35' 32'

40h 37' 26"'

-

~~

* (001) plane. (101) plane.

T. E. Madey and D . Menzel, Jpn. J . Appl. Phys., Suppl. 2, Pt. 2, 229 (1974) (isost.). P. D . Reed, C. M. Comrie, and R. M. Lambert, Surface Sci. 59, 33 (1976) (TDS). D. G . Castner, B. A. Sexton, and G . A. Somorjai, Surface Sci. 71, 519 (1978) (TDS) G. Ertl and J. Koch, 2. Naturforsch., Teil A 25, 1906 (1970) (isost.). J. C. Tracy and P. W. Palmberg, J. Chem. Phys. 51,4852 (1969) (isost.). H. Conrad, G. Ertl, J . Koch, and E. E. Latta, Surface Sci. 43,462 (1974) (isost.). ' J . Kiippers and A. Plagge, J . Vac. Sci. Techno/. 13,259 (1976). ' J . Kuppers and H . Michel, in preparation (TDS). K. Christmann and G. Ertl, 2. Naturforsch., Teil A 28, I144 (1973) (isost.). ' G . Ertl, M. Neumann, and K. M. Streit, Surface Sci. 64, 393 (1977) (isost.). R. W. McCabe and L. D . Schmidt, Surface Sci. 60,85 (1976) (TDS).

'

(91) that the resulting data are frequently too low by several kcal/mol if compared with the corresponding isoteric heats of adsorption, indicating that the preexponential might in fact be up to several orders of magnitude larger. A direct determination (ZOZ), which will be described in Section II,D, shows that for the system CO/Pd(l11) this is indeed the case. Moreover, as shown recently by King (102), simple analysis of thermal desorption spectra in terms of the Redhead model may lead to erroneous results if the adsorption kinetics obey a precursor state model [which is the case for CO adsorption (see below)]. Therefore, the values in Table I1 that were obtained from TDS data should be considered with some care.

In any case the data are relatively similar and range between 26 and 40 kcal/mol, which is quite analogous to the M-CO bond strengths in transition metal carbonyls (6). For a given metal the adsorption energies vary by less than 15% between the different crystallographic planes. Inspection of Table I shows that for Pd this also holds for the more open (210) and (31 1) planes. This result is also supported by theoretical treatments (16, 22) and appears to be a quite general phenomenon for all adsorption systems (6). This effect yields some indication that the "structural" factor in the kinetics of surface reactions may be of relatively small importance. This statement is, in fact, true for the oxidation of CO, but may by no means be generalized.

16


An interesting point in this connection is the role of “active” sites. This concept was originally introduced by Taylor (103) and was particularly concerned with structural imperfections on low-index planes such as steps, corners, etc., that will certainly be present on the surface of a real catalyst in a great variety. Convenient models for such defects are offered by single crystal surfaces with periodic arrays of atomic steps that were extensively studied by Somorjai (101). Several studies on the energetics of CO adsorption have also been performed with stepped surfaces in comparison with the corresponding low-index planes. CO was found to be bound more strongly by about 6 kcal/mol at the steps of a Pt( 1 1 1) surface ( 1 0 3 , whereas practically no effect on the adsorption energy by the presence of monatomic steps was found with Pd( 11 1) (52) and with Ir( 11 1) (106, 107). The influence of steps on the adsorption energy is generally of the same order of magnitude as the variation between different low-index planes (6). Consequences for the kinetics of the catalytic COz formation will be discussed below. The adsorption energy varies with coverage owing to the operation of lateral interactions and to the occupation of different adsorption sites. It is frequently rather difficult to separate these effects from each other which is, however, quite possible in a few cases of CO adsorption. Figure 11 shows the variation of the isosteric heat of CO adsorption on Pd( 1 1 1) with coverage. Ead is constant up to 0 = 1/3, indicating the occupation of equivalent sites with negligible repulsion between the adsorbates (87).(There is presumably even some lateral attraction leading to island formation.)

IVl 0 a2 a3 ah 0.5 e FIG.11. Adsorption energy as a function of the CO coverage and work function change

on Pd(ll1) (87).


17

At 0 = 113 Eadmakes a sudden drop by 2 kcal/mol, which is correlated with the onset of the compression of the overlayer unit cell; i.e., the adsorbed particles are displaced from the energetically most favorable sites (see Fig. 4). Near 6 = 0.5 Ead decreases steeply, which is ascribed to the onset of strong repulsive interactions owing to orbital overlap. The variation of the C-0 stretch frequency with 8 for the same system is reproduced in Fig. 12 and follows nicely the trend of the adsorption energy (60). This quantity remains almost constant up to 8 = 1/3, but then exhibits a strong shift suggesting a movement of the adsorbed molecules into other sites. There is, however, substantial evidence for the operation of another effect at higher coverages; namely an in-phase vibration of the whole coupled system of admolecules owing to intermolecular interactions (60, 61, 64, 107, 108). Thus IR spectra are believed to indicate the localized adsorption site only at low surface concentrations, whereas at high coverages this influence is virtually removed owing to strong intermolecular interactions. A somewhat different behavior is found with the system CO/Pd(l00). Figure 13 shows the variation of Eadwith coverage (94).As can be seen, this

1620-+

e

e

e me

1840-

1860-

1680-

1900-

a

1920-

+

W1

9

6

0 a1

- 1 a2 a3 a4 Covorago 8

05

FIG. 12. CO stretch frequency as a function of coverage on Pd(ll1) (60).

18


~

01

02

0.3

04

05

06

07

e

Coverage (monolayerr)

FIG.13. Adsorption energy as a function of the CO coverage on Pd(100) (94).

quantity starts to decrease continuously even at rather low surface concentration, which can be ascribed to the operation of long-range indirect repulsive interactions between the admolecules. As has been shown theoretically, such indirect interactions via the metal electrons may well be different in character and magnitude on different crystal planes (108, 109). At 0 = 0.5 Eadfalls steeply, again owing to the onset of compression of the overlayer unit cell and displacement of the CO molecules from their favorable bridge sites (see Fig. 6). At still higher coverages Eaddecreases strongly, owing to the onset of direct orbital overlap, which could fairly well be described by the interaction potential between free CO molecules (94). The variation of the adsorption energy with coverage is again nicely reflected in the change of the C-0 stretch frequency (60)(Fig. 14), which in this case has by now increased at low coverage as expected. These two examples demonstrate that energetic differences between dif-


--

1

1900 - OO 0 0

-5 1920 n

E

z

1940

-

octagonal ring Icompression octagonal ring Icompression

-

"c (4x 2 ) 45' O

0,

>

e

1

0 0

L Q)

19

"

I

0

1960 -

0 0

-

0 0

00

1980 '

1 FIG. 14.

ferent geometric locations may be rather small, which is also supported by corresponding theoretical treatments (16, 22). Figure 15 shows the "energy profile," i.e., the variation of the adsorption energy for CO across the substrate unit cell, for Pd(100) (22) which exhibits nearly equivalent bond strengths for bridge and fourfold sites and predicts a rather high mobility of the adsorbed CO that is in agreement with the experimental experience. Experimental evidence for only small energetic differences between different sites was found for CO/Pt(l 11). Here the adsorption energy between bridge and linear sites (which might eventually also be threefold sites) differs by less than 1 kcal/mol(58, 110). D. KINETICS OF ADSORPTION AND DESORPTION

The rate of adsorption onto a solid surface is given by r,

=

dn,/dt = p/(2~mkT)'~~.s(9)

(1)

where n, is the density of adsorbed particles per cm2,p the gas pressure, m the adsorbate mass, k the Boltzmann constant, T the absolute temperature,

20


FIG.15. Representation of the variation of the adsorption energy with position of the CO molecule on a Pd(100) surface (22).

and s(9) the probability for adsorption (= sticking coefficient) which is a function of the fractional coverage 9 = 0/9,,,. The term p/(27~mkT)''~ denotes the gas flux to the surface and thus the kinetics of adsorption is described by the sticking coefficient s(9). This definition is based on the assumption that a particle, once adsorbed, stays on the surface indefinitely, which, of course, may be approximately fulfilled only if the surface temperature is low enough so that the rate of desorption is negligible. Accordingly, the sticking coefficient may be derived from the variation of the coverage with time, or by measuring the fraction of particles back-scattered from the surface. The latter procedure (e.g., with molecular beam experiments) is experimentally more difficult, but yields much more accurate data if s is close to unity. This is, in fact, the case for the initial sticking coefficient so for CO on the platinum metal surfaces, i.e., for the adsorption probability extrapolated to zero coverage. In a recent molecular beam experiment for the system CO/Pd(l 11) (101) so = 0.96 was determined at 300 K. For the same system so was found to be independent of temperature between 300 and 650 K, which indicates that CO adsorption is a nonactivated process. In addition, the sticking coefficient was found to be independent of the angle of incidence of the gas flux. Similar values for so were obtained for other surfaces (111). With a series of differently oriented Pt single crystal surfaces McCabe and Schmidt (112) obtained a somewhat larger variation of so between 0.24 and 0.95. The angular dependence of the sticking coefficient that has been reported (113) is an interesting effect that needs further clarification.


21

Interesting information on the mechanism of adsorption dynamics is contained in the functional dependence of the sticking coefficient on coverage. For example, the simple Langmuir model assumes that a particle is adsorbed only if it strikes a free adsorption site from the gas phase which yields s(9) = s,(l - $), and s(9) = s,(l - 9)2 if two adjacent free adsorption sites are needed for the initial substrate-adsorbate bond formation. A dependence of the latter type was surprisingly found for CO adsorption on Ir(ll1) ( 9 4 , although no plausible explanation thus far can be given. In all the other cases investigated thus far, s(9) falls off much less at low coverages or remains practically constant. This can be (at least qualitatively) attributed to a precursor-state model (114-116). The basic idea of this model is that a molecule that is initially located (presumably in a physisorbed state) above a particle which is already chemisorbed may diffuse in this second layer until it encounters a vacant adsite onto which it may then chemisorb. In the firstorder precursor model (114) the resulting variation of the sticking coefficient with coverage is given by 4 9 ) = s o u - 9141 - C9)ll

(2)

whereby the parameter c is determined by

ct;; +fb -fb’)/U + A ) .

(3) As illustrated by Fig. 1 6 , i is the probability for adsorption of the precursor molecule on an empty surface site, & is the probability of desorption of the precursor from a bare site, and&’ the probability of desorption from an occupied site. For c = 0 the Langmuir model is retrieved, whereas for c > 0 the physically adsorbed molecule stays on the surface a relatively long time, during which it can encounter a site for chemisorption. A series of data was satisfactorily analyzed in terms of this model (111),yielding a good descripc =

\

’/////////// FIG. 16. Schematic representation of adsorption and desorption including a precursor state. is the probability for adsorption in the precursor state. f,’,fd’ and f,, fd are the probabilities for diffusion and desorption from the precursor and chemisorbed states, respectively. f, and x’ are the probabilities for the transition from the precursor state to the chemisorbed state and for the inverse process, respectively (101).

22

T. ENGEL AND G . ERTL 0

0.1

0.5

0.3

0.5-

0

1

1

1

2

4

6

8

“r I

IOX

1

,o14 molecules

[

cmz

FIG.17. The relative sticking coefficient as a function of the CO coverage on Pt(ll1) (58).

tion of the experimental data with c = 0.35 for Ru(001), c = 0.5 for Pd(l1 l), c = 0.85 for Pt(l1 l), and c = 0.88 for Pt( 1 lo). Photoemission experiments by Norton et al. (110) suggest some direct experimental evidence of the existence of a precursor state on Pt(ll1) that can probably tumble on the surface. A more detailed inspection of the data reveals, however, that the simple first-order precursor-state model does not yield quantitative agreement with the experimental results. This was, for example, shown for the system CO/Pt( 111) [for which excellent agreement between the data reproduced in Fig. 17 of three independent studies was found (58, 61, 110)] as well as for CO/Ir(110) (76). In the latter case the best fit could be achieved with a second-order precursor model in which the sticking coefficient depends on the lifetime of the precursor state as well as on the fraction of doublets of neighboring empty sites (116). A satisfactory physical explanation for such an effect is still missing. An attempt to distinguish between direct reflection of a molecule from the surface and desorption from a short-living precursor state can be made by inspection of the angular distribution of a molecular beam scattered from the surface, since in the case of elastic reflection the particles should exhibit some “memory” of their initial momentum leading to a specular contribution in the angular distribution. As can be seen from Fig. 18, for the system CO/Pd( I 1 1) this is not the case, but instead an almost perfect cosine distribution is observed (101). This indicates complete energy and momentum exchange with the surface, even at a surface temperature of 1000 K. However, detailed studies with a modulated CO beam on the phase lag between impinging and scattered particles could only be explained on the basis of the assumption of at least three different species on the sur-


23

FIG.18. Angular distribution for CO scattering from Pd(l11). At 300 K, Oco x 0.5 whereas at 1020 K,,,O

0,', the reaction is first order in pco and zero order in 0, . A different type of nonstationary state experiment that has been reported by Bonze1 and co-workers (193, 194) is illustrated in Fig. 44. A Pt( 110) surface was initially exposed to a mixture of CO and 0 2 .At the substrate temperature and pco and p o 2 , B,, > 65, and the sticking coefficient for 0, is drastically reduced. Ifp, is suddenly reduced, it is seen that after an induction period which is strongly temperature dependent the rate of CO, production passes through a maximum before decreasing sharply. A similar experiment in which the surface was initially covered with oxygen after pco was

CATALYTIC OXIDATION OF CO ON PLATINUM METALS 1

-

15

-

-

Lo-

-

0.5

483K

-

I

FIG.43.

I

I

20

Pt

55

1

1

1

reaction order plot for CO, formation on a platinum surface (132).

Po2=I.S*IO-'Amp

t(sec)

FIG.44. Rate of CO, formation on Pt(ll0) under nonstationary conditions. The surface was initially exposed to a mixture of CO and 0,. At t = 0, the CO pressure was abruptly reduced and the subsequent rate of CO, formation was registered. The time corresponding to the maximum rate shifts to lower values with increasing surface temperature (193).

suddenly increased is shown in Fig. 45. The first arrow in Fig. 45 indicates the time when poz was reduced. After a time At, pco is suddenly increased and it can be seen that the rate of CO, production rises rapidly after the admission of CO. Furthermore, the maximum rate of CO, production is directly proportional to pco .

56


.-u) C

. I -

+

p

9

0

10

20

30

40

t (sec)

FIG.45. Rate of CO, formation on Pt(ll0) under nonstationary conditions. Oxygen adsorption was terminated at the time correspondingto the first arrow. After a time At, CO was admitted into the chamber (193).

Although molecular beam relaxation investigations such as those that have been carried out on the Pd( 1 11) surface (176) are not yet available for Pt surfaces, a number of molecular beam studies that yield information additional to that described above have been reported (189, 190, 195-197). Pacia et al. (189) have carried out integral angle scattering experiments in which the variation in the angular scattering distribution of CO and 0, with coverage as well as that for the CO, formed are averaged out and the total scattering of each gas is registered as a pressure rise in the vacuum chamber. For mixed beams of CO and 0,, it is possible to measure the absolute reaction probability for an impinging 0, molecule. Under conditions corresponding to the maximum reaction rate in Fig. 40, the reaction probability for 0, is equal to the zero coverage 0, sticking probability at the corresponding substrate temperature. By impinging CO on the surface with the beam and supplying 0, through an isotropic pressure in the gas phase, the reaction probability for CO can be determined. The results in Fig. 46 show that the reaction probability is independent of substrate temperature up to 900 K at 0.8, after which it falls rapidly with increasing temperature. Angular resolved measurements of the CO, formed in the reaction have yielded rather surprising results. Palmer and Smith (196) have reported that the CO, formed on a Pt( 1 1 1) surface is emitted in a peaked distribution that can be described by a cos60! function where LY is the emission direction measured from the surface normal. Becker et al. (190) have reported similar results for a polycrystalline Pt surface. They obtain angular distributions of


57

I 04 0.2

0.00

*

T(K)

300 500 700 900 1100 1300 FIG.46. Reaction probability for a CO beam impinging on a platinum surface that is immersed in an isotropic 0, pressure as a function of the substrate temperature (189).

the form cosdcx with d between 2 and 3. In their experiment it was possible to determine the velocity of the emitted CO, molecules and it was found that the molecules emitted normal to the surface had a substantially higher velocity than those emitted at 45 degrees. The origin of such peaked distributions [that differ from the cosine distribution that was measured for CO, formed on Pd( 111) (176)] and the dependence of the particle velocity on the emission angle is as yet not completely understood (198). The experiments discussed above were all carried out with total pressures below Torr. However, Hori and Schmidt (187) have also reported nonstationary state experiments for total pressures of approximately 1 Torr in which the temperature of a Pt wire immersed in a CO-0, mixture was suddenly increased to a new value within a second. The rate of CO, production relaxed to a steady-state value characteristic of the higher temperature with three different characteristic relaxation times that are temperature dependent and vary between 3 and 100 seconds between 600 and 1500 K. The extremely long relaxation time compared with the inverse gas phase collision rate rule out an explanation based on changes within the chemisorption layer since this would require unreasonably small sticking coefficients or reaction probabilities of less than The authors attribute the relaxation times to characteristic changes of surface multilayers composed of Pt, CO, and 0. The effects are due to phases that are only formed at high pressures and, therefore, cannot be compared to the other experiments described here. The central question in this discussion is again whether the reaction takes place between two adsorbed species (Langmuir-Hinshelwood mechanism) or between a gas phase CO molecule and an oxygen adatom (Eley-Rideal mechanism). All of the investigation cited indicated that under steady-state

58


-=

conditions with pco * p o l , T T,,, and for total pressures below Torr the reaction proceeds via an LH mechanism. The evidence is seen most clearly in the inhibiting effect of CO adsorption on the rate at high 8,, (193, 194). However, for T > T,,, the CO inhibition effect is no longer observed and the reaction rate falls off slowly with increasing temperature, whereas Oco decreases by several orders of magnitude in the same temperature range. For these reasons, most authors have proposed an ER mechanism for T > T,, . An ER mechanism has also been proposed to explain that the rate of CO, production for CO impinging on an oxygencovered surface is largely independent of the substrate temperature (189, 193, 194). However, more recent studies have shown that the reaction rate for T > T,, is not first order in B0 (see Fig. 42). Therefore, the kinetics cannot be explained by a simple ER mechanism. Furthermore, the observation that above T,,, the reaction is first order i n p , is not conclusive evidence for the ER mechanism since in all the investigations for which the total Torr, l3,, will be linearly proportional to pco for pressure is below T > T,,, . In order to explain these contradictions, the zero-order dependence on 8, has been attributed (132) to a modified ER mechanism in which the CO molecule is trapped in a mobile precursor state. In this state, the CO can diffuse over the surface while sampling a number of sites and, therefore, the reaction order in 8, will be zero as long as the diffusion radius in the precursor state is large in comparison with the mean separation of the oxygen adatoms. However, a distinction between this process and the diffusion of a chemisorbed CO molecule can only be made on the basis of modulated beam relaxation studies. Similarly, the nonstationary state experiments of Bonze1 and co-workers (149, 193, 194 in which no measurable induction period is observed for CO impinging on an oxygen precovered surface is not conclusive evidence for an ER mechanism since the same behavior occurs on Pd( 111) surfaces and has been shown to be consistent with an LH mechanism (176) when the smallest detectable relaxation time is not limited by the experimental configuration to several seconds. The determination of the activation energy of the reaction is difficult since the coverages 8, and 0, under reaction conditions must be known. The values that have been determined for the LH reaction vary from 8 kcal/mol (194 to 23 kcal/mol (186, 189), whereas under the conditions attributed to an ER mechanism, no appreciable effect of temperature on the rate was seen (132, 149, 189. 193, 194). The activation energy of 8 kcal (194) was determined under conditions in which 8, and Oc0 varied appreciably and with the assumption that the reaction is first order in both coverages. In view of the results shown in Fig. 42, this assumption may not be justified. Pacia et af. (189) have obtained their activation energy between 800 and 1 150 K ( T > T,,,) for which Oco is very small. Their value agrees well with


59

that determined on Pd(ll1) in the same coverage range (176). Winterbottom (186) has determined the activation energy in an intermediate range of Oc0. Investigations on Pd( I 1 1) (176) have shown that the activation energy for the catalytic oxidation of CO depends on the surface configuration and is appreciably lower for high 8, than for low 8,. This and the questionable assumption that the reaction rate is first order in both 8, and do over a wide coverage range may explain the differences in the reported activation energies. In summary, it may be safely concluded that the sequence of reaction steps as established for palladium [i.e., (7)-(9)] also holds for platinum and that the reaction will be governed by essentially similar rate laws although the experimental evidence is not yet as complete and conclusive. Gradual variations between these two metals can be ascribed to differences in the sticking coefficients and adsorption energies of the reactants. In addition, it appears that incorporated oxygen (which is less reactive than the chemisorbed species) may play a more important role with Pt, as indicated, for example, by the results of Hori and Schmidt (187). VI.

Oxidation of CO on Iridium

The catalytic oxidation of CO on iridium has not been as extensively studied as with palladium and platinum. However, as for these metals, both steady-state (40, 54, 124, 199-202) and nonsteady-state investigation (124, 200-203) have been carried out on both polycrystalline and single crystal surfaces. As the results are for the most part very similar to those obtained on palladium and platinum surfaces, only those results that shed additional light on the kinetics and mechanism basic to the reaction will be emphasized here. Figure 48 shows a series of curves for the rate of CO, production on Ir(ll0) as a function of the substrate temperature for two different total pressures and for various ratios of pco to po2 (124). As was observed on Pd and Pt, the reaction rate increases sharply between 400 and 500 K, and reaches a maximum between 500 and 600 K,after which it decreases with increasing substrate temperature. Thus far the results are quite similar to those described for Pd and Pt. However, Fig. 47 also shows that the shape of the curve for Ir( 1 10) is different for increasing and decreasing temperatures. This hysteresis is independent of the heating rate below 4 K/sec (142). It has been suggested (124) that this behavior is due to two stable reaction configurations that are separated by an activation barrier. Since this behavior has not previously been reported on other platinum metals, it may be due to the conversion of chemisorbed to incorporated oxygen on Ir( 110). Figure 48 shows the variation of the concentration of adsorbed CO and oxygen with

60


TEMPERATURE ( K ) FIG.47. Rate of CO, production on Ir(ll0) showing the hysteresis obtained in a temperature cycle. Q is the ratio of the CO and 0, pressures (124).

temperature under reaction conditions. The determinations that were made with XPS (200) show that the maximum rate is reached at a temperature for which the CO coverage is already rather small and the oxygen concentration is at saturation. This effect completely parallels the findings with Pd and Pt; i.e., adsorbed CO acts again as an inhibitor for oxygen chemisorption. As has been discussed above for palladium, the kinetic parameters for the CO oxidation reaction depend strongly on the surface configuration which is a function of the CO and oxygen coverages. Moreover, the sticking coefficients for CO and 0, are strongly coverage dependent and those for 0, decrease with increasing CO coverage. These properties of the reaction system make it difficult to postulate a simple kinetic model. However, Kiippers and Plagge (203) have simulated the Langmuir-Hinshelwood reaction with the same reaction steps as formulated in Section IV,B by (7)-(9). At the temperatures of interest, the thermal desorption of 0, again can be neglected. Using the parameters po2 = pcs3= 1 x lo-’ Torr, density of adsorption sites 1.6 x 1015 cmP2, k, = 10 exp[-E~~(0,,)/RT], k, = 6.1 x lo6 sec- exp( - 10.7/RT) and the experimentally determined quan6,) and Edes(OCO),where s(O) and E(0) are the sticking cotities s,,(Oo,


I

I

I

1

61

I

TEMPERATURE (K)

FIG.48. Rate of CO, formation as well as 0, and 8, perature for CO, formation on Ir(l11) (200).

as a function of the substrate tem-

efficient and activation energy for desorption as a function of the relevant coverages, the steady-state reaction rate could be simulated numerically. The results that are shown in Fig. 49 are in good agreement with the experimental results and show that the reaction kinetics can be described quantitatively with a Langmuir-Hinshelwood mechanism. A somewhat higher activation energy of between 13 and 16 kcal/mol, depending on the coverages, has been reported by Taylor et a/. (124) for the oxidation reaction on Ir(ll0). These investigators have also compared the reaction rates on clean and oxidized Ir(l10) surfaces and have found that the activation energy for the reaction is essentially unchanged by the oxide layer. This is of importance in comparing the results of kinetic investigations obtained at low pressures in UHV systems with those obtained at high pressures since oxygen chemisorption at elevated temperatures and high pressures often leads to the formation of an oxide layer, as discussed above. The occurrence of a hysteresis in the Y( T ) curves (as reproduced in Fig. 47) is, however, difficult to explain in view of these findings.

62


10

08 06 04

02

LOO

co 0

06

-

0-

500

600

TlKl

coverage 0-

0oO o >,

04

02

\O ‘0.

LOO

006

I

o-o--o-o-

500

600

TlKl

500

600

T IK1

0 coverage

004

002 LOO

FIG.49. Computer simulation of the L-H rate of CO, production on Ir(ll1) (see text) (203).

Fewer nonsteady-state measurements have been carried out on iridium than on platinum and palladium. Figure 50 shows the results of a 0,-CO coadsorption experiment on Ir( 11 1) (203). Initially 0, was adsorbed, followed by CO adsorption, after which the crystal was heated with a linear temperature rise. It is seen that the peak temperature for CO, desorption is shifted to lower values with increasing CO coverage. This may be due to a second-order desorption effect (203) or a reduced activation energy for the reaction owing to interactions in the adlayer, as was found on Pd(l11) (176). 2


63

FIG.50. Nonstationary CO, production from an adlayer on Ir(l11) containing both oxygen and CO (203).

In conclusion, a qualitatively similar behavior to that of Pd and Pt is found with iridium. Whether the activation energy for the LH reaction under comparable conditions is indeed appreciably lower with Ir than with the other two metals needs, however, some further clarification.

VII

.

Oxidation of CO on Ruthenium

The oxidation of CO has been studied on the Ru(001) (148) and Ru(l0l) (159) surfaces. Some significant differences have been observed with respect to the behavior observed on Pd, Pt, and Ir surfaces. Figures 51 and 52 show the steady-state rate of CO, production as a function of temperature on both surfaces under the same reaction conditions (pco= po2 = 4 x Torr). Madey el af. (148) found that the reaction rate on Ru(OO1) is unappreciable below 400 K and slowly approaches a maximum that is reached only at 950 K, after which the rate again decreases. These data are shown in Fig. 51. Figure 52 shows the results obtained by Reed el al. (159) on Ru(lO1). On this surface the maximum rate is reached at 650 K, which is quite similar to the value reported for Pd, Pt, and Ir. However, a further increase in the temperature to 1100 K reduces the reaction rate by only 30% (159).In analogy to the analysis of the palladium results, this suggests


Temperature ( K

I

FIG. 51. Steady-state rate of CO, production on Ru(001) as a function of the substrate temperature (148).

FIG. 52. Steady-state rate of CO, production on Ru(l0l) as a function of the substrate temperature (159).

that the activation energy for the reaction and that for CO desorption are nearly equal. The structural specificity observed on these two ruthenium surfaces is greater than that which has been observed on the other platinum group metals. Madey et al. (148) concluded that the desorption of CO is not responsible for the form of the reaction rate curve on Ru(OO1) for temperatures up to the maximum since CO is desorbed at temperatures considerably below 950 K. Reed et a / . (159)concluded that CO desorption is also not responsible for the form of the reaction rate curve on Ru(lO1) since the


65

apparent activation energy obtained from a plot of the logarithm of the reaction rate vs 1/T is lower than the activation energy for the desorption of CO. This conclusion is, however, misleading since without correcting for the temperature dependence of the CO and oxygen coverages and the coverage dependence of the sticking coefficients, in particular that for oxygen on the CO coverage, no meaningful activation energy can be extracted from the reaction rate vs. temperature curve. In any case there is clear evidence that the kinetics of CO, formation on Ru is somewhat different from the behavior of the other platinum metals which is possibly caused by variations of adsorption and activation energies.

VIII.

Oxidation of CO on Rhodium

Thus far no results for the catalytic oxidation of CO on single crystal planes of Rh have been reported, but since the adsorption properties of the reactants are quite similar to those for palladium, close similarities of the reaction mechanism and kinetics on these two metals can be expected. Quite recent nonstationary measurements by White and co-workers (204206) with polycrystalline Rh wires support this suggestion. Numerical simulation of the transient CO, formation on the basis of the reaction sequence described by (7)-(9) yielded agreement with the experimental results under the assumption of a single LH mechanism with an activation energy of 25 kcal/mol above 530 K (corresponding to low CO coverages) and of 14.3 kcal/mol below this temperature (higher CO surface concentrations) that agree almost perfectly with the values derived for Pd(ll1) by the molecular beam technique as outlined in Section IV,C. An essential difference between both metals, however, appears to exist since with Rh it was concluded that oxygen may inhibit the adsorption of CO which was never observed with Pd. Further investigations using surface spectroscopic techniques are certainly needed in order to resolve this question. IX.

Factors Influencing the Catalytic Activity

A.

INFLUENCE OF THE SURFACE STRUCTURE

The activity of a catalyst for a particular reaction may be strongly dependent on the surface structure. Reactions of this type are called “structure sensitive” or “demanding,” whereas with “structure insensitive” or “facile” reactions this effect is of minor importance (206). With real catalysts this distinction is usually obtained by varying the mean particle size (and thereby

66


the relative proportions of the different exposed crystal planes and lattice imperfections), whereas with single crystal studies which form the main scope of the present contribution a more direct investigation may be performed by varying the crystallographic orientation of the surface. In addition, surfaces with periodic arrays of monoatomic steps can be prepared which form convenient models for “active sites” (104). Figure 53 shows relative rates of CO, formation under steady-state conditions that were recorded with various single-crystal surfaces of Pd as well as with a polycrystalline Pd wire (173). It must be noted that with these experiments no determination of the effective surface areas was performed so that no absolute turnover numbers per cmz are obtained. Instead, the reaction rates were normalized to their respective maximum values. As can be seen from Fig. 53, all data points are close to a common line which indicates that, in fact, with this reaction the activity is influenced very little by the surface structure. As has been outlined in Section 11, the adsorption of CO exhibits essentially quite similar behavior on single-crystal planes with varying orientation. Since the adsorption-desorption equilibrium of CO forms an important step in the overall kinetics of steady-state CO, formation, this effect forms at least a qualitative basis on which the structural insensitivity may be made plausible. A more direct investigation of the influence of the surface structure on the catalytic activity was performed by Hopster et al. (151) by using a

FIG.53. Steady-staterate of C02 production on various palladium surfaces as a function of the substrate temperature. The maximum rate has been normalized to the same value for all samples (156).

67


platinum single crystal whose surface was curved in such a way that not only the (1 1 1) plane, but also vicinals with varying step density of two different crystallographic directions were present. The steady-state reaction rate as well as 0, and the sticking coefficient for oxygen could be measured as a function of p o 2 , pco, and the substrate temperature. The area upon which the reaction was studied was determined by the electron beam diameter which was less than 0.5 mm. Figure 54 shows the variation in 0, with the step density. It is seen that 0, increases with increasing step density, which is partly due to an enhanced sticking probability. Figure 55 shows the reaction probability for an impinging CO molecule as a function of 0, for the Pt(ll1) and the Pt[14( 11 1) x (1 1 I)] surface. The reaction probability is larger on the flat surface for low 8, which may be attributed to a higher binding energy of the oxygen adatom at step sites that reduce the reaction probability on the stepped surface. This effect almost compensates the higher sticking coefficient for oxygen. For larger B,, the reaction probabilities become equal. As was observed by Matshushimaet al. (132),the rate of CO, production is zero order in 0, above a critical oxygen coverage 6; that for the Pt(ll1) surface is BOc = 0.04. This value compares well with the value of 0,' = 0.08 that was observed on the Pd(l11) surface (176). Figure 56 shows the CO, yield determined for the Pt( 1 1 1) and the Pt[ 14( 1I 1) Y (1 1 l)] surfaces as a function of the partial pressure ratio pOI/pCO.Both curves deviate only very slightly from each other. Although the yield varies with step density, the effect for 7% +

QI

Yq

3

0.10

-

3x120

OP) \ d f l

0 V

B8

0 2x20

ODs-

-\\ 0 1

I

0-d 1

0.1

I

kp' 1

0

-n

t.SSOK I

I

0.1

n(rooti

ln-1(min

fn-'(ml~(ll?tl

1

0.2

,

68


OXYGEN COVERAGE 9

FIG.55. Reaction probability for CO impinging on flat and stepped Pt(ll1) surfaces as a function of the oxygen coverage (151).

PARTIAL PRESSURE RATIO b2/b FIG.56. Catalyticyield as a function of the partial pressure ratiop,,/p,, for flat and stepped Pt(ll1) surfaces (151).

step sites is less than lo%, indicating that the reaction is not structure sensitive to a significant degree. This justifies the comparison made in this review between polycrystalline and single-crystal surfaces. This result is not very surprising for T > T,, since on a flat surface the reaction yield per impinging CO molecule is already on the order of 0.5, so that at most an increase of a


69

factor of 2 in the reaction probability is possible through geometrical optimization of the surface structure. These results also demonstrate that for the systems discussed here there is no evidence that the reaction occurs preferentially at surface imperfections. It thus becomes also clear why the overall kinetics with polycrystalline samples and “real” catalysts is quite similar to the behavior of flat singlecrystal surfaces as long as the degree of surface cleanliness and the pressure conditions are comparable. In this regard it is interesting to notice that in a study under “real” catalytic conditions with alumina-supported platinum catalysts McCarthy et al. (207) found no dependence of the catalytic activity with particle size (i.e., structure insensitivity as observed with the low-pressure single crystal studies) between 480 and 550 K at high CO concentrations, whereas in excess oxygen the reaction became structure sensitive. In the latter case partial oxidation presumably occurred which causes reconstruction of the surface and the by also a variation of the effective surface area.

IMPURITIES B. THEROLEOF SURFACE

A further serious factor determining the apparent gap between model studies with well-defined clean surfaces and the conditions encountered with real catalysts is the fact that the latter are never prepared in a way that the degree of surface contamination is well controlled. Elements such as sulfur or carbon might always be present in considerable concentrations on the surface so that the chemical nature of the catalytically active surface could possibly be quite different. Bonze1 and Ku performed a detailed study on the influence of sulfur on the adsorption of CO (208) as well as on the catalytic CO, formation (209) on a Pt( 1 1 1) surface. It was found that preadsorbed sulfur affected both the adsorption energy of CO as well as the total amount of CO adsorbed. A Pt surface saturated with sulfur (0, x 0.75) was no longer able to adsorb any CO. Consequently, the rate of CO, formation also decreased continuously with increasing sulfur content of the surface and became practically zero for 0, x 0.28. These results demonstrate the role of sulfur as a rather effective catalyst poison. On the other hand, it was found (209) that interaction between oxygen and sulfur led to the formation and desorption of SOz in a LangmuirHinshelwood type reaction whereby sulfur is removed from the surface. Figure 57 shows the decrease of the relative S concentration on the surface with time under the influence of an oxygen atmosphere at 200°C. For low S concentrations this reaction proceeds quite rapidly, whereas at higher 8, an increasingly longer induction period occurred. In this latter range it was

70


L

I

1

I

I

I

I

I

0

5

10

15

20

25

30

35

4

[

t min]

FIG.57. Surface reaction between sulfur and oxygen. The relativeconcentration of adsorbed sulfur is shown as a function of time for constant substrate temperature and oxygen coverage. R refers to the initial sulfur concentration (209).

observed that SO, formation proceeded inhomogeneously along the island boundaries from adsorbed sulfur. Similar processes presumably occur if carbon atoms that may also react with oxygen to form CO are present on the surface. These results suggest that under catalytic steady-state conditions in a CO/O, mixture the main impurities may be reacted off to volatile compounds so that the reaction proceeds, in fact, on an essentially clean surface even without the need for an ultrahigh vacuum and stringent cleaning procedures of the surface. This suggestion was confirmed by an experiment in which a Pd wire was exposed to a CO/O, reaction mixture without any previous surface cleaning (173). The stationary rate of CO, formation as a function of temperature (as shown in Fig. 58) is at first rather low, but increases steeply above a certain activation temperature. In any further runs the catalyst exhibited its normal activity, as is known from atomically clean surfaces. Auger spectroscopy demonstrated that at the beginning the surface was heavily contaminated by S and C , whereas these elements were completely absent after the sample had reached its full activity. Thus the well-known “break-in” effect of the catalyst (210) can in this case be simply explained by the removal of the inhibitors by surface reaction with one of the reactants, and the abovestated assumption is confirmed.


71

T 300 Loo 500 FIG.58. Rate of C 0 2 formation as a function of temperature on an initially contaminated Pd wire surface showing the effect of activation (156).

C. THEPRESSURE GAP The results discussed in this article were mostly obtained with ultrahigh vacuum systems at total pressures not exceeding Torr, whereas real catalysis is performed in the atmospheric pressure regime. This general “pressure gap” raises the serious question to which extent experiments of the type described using the spectroscopic techniques of “surface science” are relevant at all for real-life catalysis. A general answer to this problem can certainly not yet be offered. However, a rather favorable situation is found in the present case, as long as the discussion is confined to temperatures below T,, at which the reaction rate reaches is maximum r,,, (cf., for example, Fig. 35). This situation has been discussed in detail in Section IV for palladium and holds as well for the other platinum metals since the shape of the r(T) curve is always quite similar. It has been shown that the kinetics may then approximately be described by

(16) The interesting point is that only the ratio of the partial pressures of the reactants enters this equation, but not the total pressure. The physical reason is obvious. Both reactants are competing for free adsorption sites on the surface and since CO inhibits the dissociative adsorption of oxygen the partial pressure of the latter compound enters into the denominator of the rate expression. This result indicates that the actual surface composition and, in addition, the reaction mechanism will not be affected by the total pressure. The only effect of increasing the total pressure (at a fixed po,:pco ratio) will be that rmax will be shifted towards higher temperatures since reduction of r = k‘.Po21Pco.

72


the CO coverage requires continuously higher temperatures. It is thus predicted that r,,, will presumably never be reached at reasonable temperatures ( < SOOT) if the total pressure lies within the Torr range, so that (16) yields a reasonable description of the kinetics over the entire temperature range. This conclusion was indeed confirmed in a series of previous investigations with Pd catalysts under real pressure conditions (163, 165, 166) where, in fact, rate laws according to (16) were established. Baddour et al. (166) determined the reaction rate and the relative surface concentration of adsorbed CO (by means of infrared spectroscopy) simultaneously under reaction conditions with a silica supported Pd catalyst. With a mixture of 0.2 Torr 0, and 0.04 Torr CO, they found the onset of the reaction as well as of desorption CO to occur above 370 K. This observation is in complete accordance with the findings obtained with Pd single crystals at much lower total pressures. The only deviation from the above rate law was reported by Daglish and Eley (164) who found a rate law of the type r of p0,/p& with a Pd wire. If this equation would, in fact, be generally valid, an increase of the total pressure at fixed stoichiometry of the gas mixture would cause a continuous decrease of the rate that would eventually become vanishingly small at atmospheric pressure which would not be in agreement with the general practical experience. The situation may become more complicated at higher temperatures where the eventual formation of oxides (or less active surface oxygen species) may come into play. Their role in the catalytic CO, formation must still be explored in more detail. The experiments by Hori and Schmidt (187) with Pt as described in Section V give some indication of the additional effects that may occur at higher pressures. Neglecting this complication, however, we may safely conclude that the kinetics and mechanism of the catalytic oxidation of CO on the platinum metals with real catalysts and pressure conditions are essentially governed by the same microscopic elementary processes as established experimentally with the model systems comprising single-crystal surfaces and rigorous vacuum conditions. Similar conclusions were, for example, also reached by Hanson and Boudart (211) by comparing their kinetic data on the oxidation of hydrogen on Pt/SiO, catalysts under normal pressure conditions with the results of corresponding low-pressure studies (212, 213). These conclusions on the pressure gap that may be regarded as being firmly established for the catalytic oxidation of CO on metals of the platinum group, illustrate the power that surface science techniques in conjunction with single-crystal studies under ultrahigh vacuum have when applied to practical catalytic reactions. ACKNOWLEDGMENT We gratefully acknowledge the receipt of manuscripts from W. H. Weinberg and J. M. White prior to publication.


13

REFERENCES I. 2. 3. 4. 5. 6.

Langmuir, I., Trans. Furuduy Soc. 17,672 (1922). Ertl, G., and Rau, P., Surt Sci. 15,443 (1969). Ford, R. R., Ada. Curd. 21, 51 (1970). Blyholder, G., J . Phys. Chem. 68,2772 (1964). Eischens, R. P., and Pliskin, W. A., Adv. Carul. 10, 1 (1958). See, for example, Ertl, G., in “The Nature of the Surface Chemical Bond” (T. N. Rhodin and G. Ertl, eds.), Chapter 5. North-Holland Publ., Amsterdam, 1978. 7 . See, for example, Gadzuk, J. W., and Rhodin, T. N., in “The Nature of the Surface Chemical Bond” (T. N. Rhodin and G. Ertl, eds.), Chapter 3. North-Holland Publ., Amsterdam, 1978. 8. Conrad, H., Ertl, G., Knozinger, H., Kiippers, J., and Latta, E. E., Chem. Phys. Leu. 42, 115 (1976). 9. Plummer. E. W., Salanek, W. R., and Miller, J. S., Phys. Rev. A 18, 1673 (1978). 10. Blyholder, G., J. Phys. Chem. 68,2772 (1964). 11. Grimley, T. B., in “Molecular Processes on Solid Surfaces” (E. Drauglis, T. R. Gretz, R. J. Jaffee, eds.), p. 299. McGraw-Hill, New York, 1969. 12. Robertson, J. C., and Wilmsen, C. W., J . Vuc. Sci. Technol. 9,901 (1972). 13. Blyholder. G., J . Vuc. Sci. Techno/. 11, 845 (1974). 14. Blyholder, G., J. Phys. Chem. 79, 756 (1975). IS. Batra, I. P., and Bagus, P. S., Solid Srure Commun. 16, 1097 (1975). 16. Doyen, G., and Ertl, G., Surf: Sci.43, 197 (1974). 17. Cederbaum, L. S., Domcke, W., von Niessen, W., and Brenig, W., Z. Phys. 821,381 (1975). 18. Bagus, P. S., and Hermann, K., Solid Stare Commun. 20, 5 (1976). 19. Kasowski, R. V., Phys. Rev. Lett. 37,219 (1976). 20. Waber, J. T., Adachi, H., Averill, F. W., and Ellis, D. E., Jpn. J. Appl. Phys., Suppl. 2, Pt. 2, 695 ( 1 974). 21. Hermann, K., and Bagus. P. S., Phys. Rew. B 16,4195 (1977). 22. Doyen, G., and Ertl, G., Surf: Sci. 69, 157 (1977). 23. Chen, B. H., Foyt, D. C., and White, J. M., Surf: Sci. 67, 218 (1977). 24. Yu,H. L., Phps. Rea. B 15,3609(1977). 25. Politzer, P., and Kasten, S. D., Surj. Sci. 36, 186 (1973). 26. Politzer, P., and Kasten, S. D., J. Phys. Chem. 80, 385 (1976). 27. Anderson, A. B., and Hoffmann, R., J . Chem. Phys. 61,4545 (1974). 28. Anderson, A. B., Surf: Sci. 62, 119 (1977). 29. Itoh, H., Jpn. J . Appl. Phys. 16, 2125 (1977). 30. Conrad, H.. Ertl, G., Kiippers. J., and Latta, E. E., Furuduy Discuss. Chem. Soc. 58, 116 (1974). 31. Kiippers, J., Conrad, H., Ertl, G., and Latta, E. E., Jpn. J. Appl. Phys., Suppl. 2, Pt. 2, 225 ( 1974). 32. Gustafsson, T., Plummer, E. W., Eastman, D. E., and Freeouf, J. L., Solid Srure Commun. 17,391 (1977) 33. Lloyd, D. R., Quinn, C. M., and Richardson, N. V., Solid Stare Commun. 20,49 (1976). 34. Norton, P. R., and Richards, P. J., Surf. Sci. 49, 567 (1975). 35. Apai, G., Wehner. P. S., Williams, R. S., Stohr, J., and Shirley, D. A,, Phys. Rev. Lett. 37, 1497 ( I 976). 36. Smith, R. J., Anderson, J., and Lapeyre, G. J., Phys. Rea. Left. 37, 1081 (1976). 37. Collins, D. M., and Spicer, W. E., Surf: Sci. 69, 114 (1977). 38. Bonzel, H. P., and Fischer, T. E., Surf. Sci. 51,213 (1975). 39. Brodkn, G., Rhodin, T. N., Brucker, C., Benbow, R., and Hurych, Z . , Surf. Sci. 59, 59 (1976).

74


40. Kiippers, J., and Plagge, A., J . Vuc. Sci. Technol. 13, 259 (1976). 41. Broden, G., and Rhodin, T. N., SolidState Commun. 18, 105 (1976). 42. Zhdan, P. A., Boreskov, G. K.,Boronin, A. I., Egelhoff, W. F., and Weinberg, W. H., Chem. Phys. Lett. 44,528 (1976). 43. Zhdan, P. A., Boreskov, G. K., Boronin, A. I., Schepelin, A. P., Egelhoff, W. F., and Weinberg, W. H., Surf. Sci. 71, 267 (1978). 44. Rhodin, T. N., Seabury, C., Traum, M.. and Hurych, Z., in “Vacuum Ultraviolet Radiation Physics” (M. C. Castex. M. Poney, and N. Poney, eds.). Vol. 11, p. 259. Montpellier, Mendon, 1977. 45. Rhodin, T. N., and Brucker, C. F., Solid State Commun. 23,275 (1977). 46. Rhodin, T. N., Kanski, J., and Brucker. C. F., Solid State Commun. 23, 723 (1977). 47. Fuggle, J . C., Madey, T. E., Steinkilberg, M., and Menzel, D., Surf. Sci. 52, 521 (1975). 48. Neumann, M., personal communication. 49. Allyn, C. L., Gustafsson, T., and Plummer, E. W.. Chem. Phys. Lett. 47, 127 (1977). 50. Plummer, E. W., in “Interactions on Metal Surfaces” (R. Gomer, ed.), p. 143. SpringerVerlag, Berlin and New York, 1975. 51. Klapper, K.,Kempin, H. F., and Ertl, G., Phys. Rev. Lett. 41, 333 (1978). 52. Conrad, H., Ertl, G.. Koch. J., and Latta, E. E., Surf. Sci. 43,462 (1974). 53. Madey, T. E., and Menzel, D., Jpn. J . Appl. Phys.. Suppl. 2, Pt.2, 229 (1974). 54. Christmann, K., and Ertl, G.. Z. Nuturforsch., Teil A 28, 1144 (1973). 55. Kempin, H. F.. Klapper, K., and Ertl, G., Nouu. J. Chim. 1, 295 (1977). 56. Kiippers, J., and Michel, H., to be published. 57. “CRC Handbook of Physics and Chemistry,” 48th ed., p. E-65. Chem. Rubber Publ. Co., Cleveland. Ohio. 58. Ertl, G., Neumann, M., and Streit, K. M.. Surf. Sci. 64, 393 (1977). 59. Horn, K., and Pritchard, J., J. Phys. (Paris) 38, 164 (1977). 60. Bradshaw, A. M., and Hoffmann, F. M., Surf. Sci. 72,513 (1978). 61. Shigeishi, R. A., and King, D. A,, Surf. Sci. 58,484 (1976). 62. Froitzheim, H., Hopster, H., Ibach, H., and Lehwald, S., Appl. Phjs. 13, 147 (1977). 63. Krebs, H., and Liith, H., Appl. Phys. 14, 337 (1977). 64. Crossley, A., and King, D. A,, Surf. Sci. 68, 528 (1977). 65. Thomas, G. E., and Weinberg, W. H., J. Chem. Phys. 70,954, 1437 (1979). 66. Kesmodel, L. L., Stair, P. C., and Somorjai, G. A,, S u f . Sci. 64, 342 (1977). 67. Hagstrom, S., Lyon, H. B., and Somorjai, G. A., Phys. Rev. Lett. 15, 491 (1965). 68. Fedak, D. G., and Gjostein, N. A , , Surf. Sci. 8, 77 (1967). 69. Helms, C. R., Bonzel, H. P., and Kelemen, S., J. Chem. Phys. 65, 1773 (1976). 70. Grant, J. T., Surf. Sci. 18,288 (1969). 71. Ignatiev, A., Jones, A. V., and Rhodin, T. N., Surf. Sci. 30,573 (1972). 72. Palmberg, P. W., and Rhodin, T. N., Phys. Rev. 161,586 (1967). 73. Heilmann, P., Lang, E., Heinz, K., and Miiller, K., Surf. Sci. 83,487 (1979). 74. Comrie, C. M., and Lambert, R. M., J . Chem. Soc., Furaduy Trans. I 72, 1659 (1976). 75. Nieuwenhuys, 9. E., and Somorjai, G. A., Surf. Sci. 72,8 (1978). 76. Taylor, J. L., Ibbotson, D. E., and Weinberg, W. H., to be published. 77. Chan, C. M., Van Hove, M. A., Weinberg, A. H., and Williams, E. D., SolidSrure Commun. 30,47 (1979).

78. Grant, J. T., Surf. Sci. 18,228 (1969). 79. Veillard, H., Nouv. J . Chim. 2, 215 (1978). 80. Engler, W.. Heiland, W., and Taglauer, E., Verh. Dtsch. Phys. Ges. [6] 13, 592 ( I 978). 81. Davenport, J. W., Phys. Reu. Leu. 36, 945 (1976). 82. Allyn, C. L., Gustafsson, T., and Plummer, E. W., Chem. Phys. Lett. 47, 127 (1977).


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83. Fuggle, J. C.. Steinkilberg. M., and Menzel, D., Chem. Phys. 11, 307 (1975). 84. Rhodin, T. N., Seabury, C., Traum, M., and Hurych, Z., in “Vacuum Ultraviolet Radiation Physics” (M. C. Castex, M. Poney, and N. Poney, eds.), Vol. 11, p. 259. Montpellier, Mendon, 1977. 85. Behm, R. J., Christmann, K., Ertl, G., Van Hove, M. A., Thiel, P. A., and Weinberg, W. H., Surf. Sci. (in press). 86. Anderson, S., and Pendry, J. B., Surf. Sci. 71,75 (1978). 87. Ertl, G., and Koch, J., Z. Naturforsch., Teii A 25, 1906 (1970). 88. Conrad, H., Ertl, G., and Kuppers, J., Surf. Sci. 76,323 (1978). 89. Castner, D. G., Sexton, B. A., and Somorjai, G. A., Surf. Sci. 71, 519 (1978). 90. Hagen, D. I., Nieuwenhuys, B. E., Rovida, G., and Somorjai, G. A., Surf. Sci. 57, 632 ( 1 976). 91. Comrie, C. M., and Weinberg, W. H., J. Chem. Phys. 64,250 (1976). 92. Grant. J. T., and Haas, T. W., Surf Sci. 21,76 (1970). 93. Park, R. L., and Madden, H. H., Surf. Sci. 11, 188 (1968). 94. Tracy, J . C., and Palmberg, P. W., J. Chem. Phys. 51,4852 (1969). 95. Ertl. G., and Koch, J., 2. Phys. Chem. [N. F.] 69, 323 (1970). 96. Grant, J. T., Surf. Sci. 18,228 (1969). 97. Broden, G., Pirug, G., and Bonzel, H. P., Surf. Sci. 72,45 (1978). 98. Kneringer, G., and Netzer, F. P., Surf. Sci.49, 125 (1975). 99. Lambert, R. M., Surf. Sci. 49, 325 (1975). 100. Reed, P. D., Comrie, C. M., and Larnbert, R. M., Surf. Sci.59, 33 (1976). 101. Engel, T., J . Chem. Phys. 69,373 (1978). 102. King, D. A., Surf. Sci.64,43 ( I 977). 103. Taylor, H. S., J . Phys. Chem. 30, 145 (1926). 104. Somorjai, G. A., Adu. Catal. 26, 1 (1977). I05. Collins, D. M., and Spicer, W. E., Surf. Sei. 69,85 (1977). 106. Nieuwenhuys, B. E., Hagen, D. I., Rovida, G., and Somorjai, G. A., SurJ Sci. 59, 155 ( 1 976). 107. Hagen, D. I., Nieuwenhuys, B. E., Rovida, G., and Somorjai, G . A,, Surf. Sci. 57, 632 ( I 976). 108. Einstein, T. E., and Schrieffer, J. R., Phys. Rev. E 7,3629 (1973). 109. Johansson, P., and Hjelmberg, H., Ned. Tijdschr. Vacuumtech. 16, 119 (1978). 110. Norton, P. R., Goodale, J. W., and Selkirk, E. B., to be published. 111. Weinberg, W. H., Comrie, C. M., and Lambert, R. M., J. Catal. 41, 489 (1976). 112. McCabe. R. W., and Schmidt, L. D., Surf Sci. 66, 101 (1077). 113. Steinbriichel, C., and Schmidt, L. D., Phys. Rev. E 10,4209 (1974). 114. Kisliuk, P. J., J. Phys. Chem. Solids3, 95 (1967); 5, 78 (1958). 115. Kohrt, C., and Corner, R., J. Chem. Phys. 52,3283 (1970). 116. King, D. A,, and Wells, M. G., Proc. R . Sac. London, Ser. A 339,245 (1974). 117. Redhead, P. A,, Vacuum 12,203 (1962). 118. Schmidt, L. D., Caral. Rev. 9, 1 I5 (1974). 119. King, D. A,, Sur$ Sci. 47,384(1975). 120. Menzel, D., in “Interactions on Metal Surfaces’’ (R. Gomer, ed.), p. 101. Springer-Verlag, Berlin and New York, 1975. 121. Chan, C. M., Aris, R., and Weinberg, W. H., AppL Sui$ Sei. I, 360(1978). 122. Pfnur, H., Feulner, P., Engelhardt, H. A,, and Menzel, D., Ned. Tijdschr. Vacuumtech. 16, 232 (1978). 123. Norton, P. R., Surf. Sci. 47,98 (1975). 124. Taylor, J. L., Ibbotson, D. E., and Weinberg, W. H., Surf. Sci. 79, 349 (1979).

76


125. Wilf, M., and Dawson, P. T., Surf. Sci.65,399 (1975). 126. Fuggle, J. C., Madey, T. E., Steinkilberg, M., and Menzel, D., Phys. Letr A 51, 163 (1965). 127. Egelhoff, W. F., Jr., Linnett, J. W., and Perry, D. L., Furuduy Discuss. Chem. SOC.58, 37

(1974). 128. Madey, T. E., Yates, J. T., Jr., and Erikson, N. E., Chem. Phys. Lett. 19, 487 (1973); Surf. Sci. 43, 257 (1974). 129. Conrad, H., Kuppers, J., Nitschke. F., and Plagge, A,, Surf. Sci.69, 668 (1977). 130. Conrad, H., Ertl, G., Kiippers, J., and Latta, E. E., Surf. Sci. 65, 235 and 245 (1977). 131. Tucker, C. W., Jr., J. Appi. Phys. 37,4147 (1966); 38,2996 (1967). 132. Matshushima, T., Almy, D. B., and White, J. M., Surf. Sci. 67,89 (1977). 133. Ivanov, V. P., B0reskov.G. K., Savachenko, V. I., Egelhoff, W. F., Jr., and Weinberg,

W. H., Surf. Sci. 61,207 (1976). Ducros, R., and Merrill, R. P., Surf Sci.55, 227 (1976). McCabe, R. W., and Schmidt, L. D., Surf: Sci.65, 189 (1977). Kuppers, J., and Ertl, G., Surf Sci. 77, 647 (1978). Collins, D. M.. and Spicer, W. E., Surf: Sci. 69, 114 (1977). Eastman, D. E., and Cashion, J. K..Phys. Rec. Leu. 27, 1520 (1971). Batra, I . P.,and Robaux, 0.. J. Vuc. Sci. Technol. 12, 242 (1975). Batra, 1. P., and Robaux, O., Surf: Svi. 49, 653 (1975). Rosch, N., and Menzel, D., Chem. Phys. 13,243 (1976). 142. Chen, B. H., Foyt, D. C., and White, J. M.. Surf. Sci.67,218 (1977). 143. Doyen, G . , and Ertl, G., J . Chem. Phys. 68, 5417 (1978). 144. Jacobi, K., Scheffler, M., Kambe, K., and Forstrnann, F., Solid State Commun. 22, 17 ( I 977). 145. Chan, C. M., Luke, K. L., Van Hove, M. A., Weinberg, W. H., and Withrow, S. P., Surf. Sci. 78,386 (1978). 146. Thiel, P. A,, Yates, J. T., and Weinberg, W. H.. Surf. Sci.82,22 (1979). 147. Grant. J. T., Surf. Sei. 25,451 (1971). 148. Madey, T. E., Engelhardt, H. A., and Menzel, D., Surf Sci.48,304 (1975). 149. Bonzel, H. P., and Ku, R., Surf: Sci.40,85 (1973). 150. Weber, B., Fusy, J., and Cassuto, A,, J . Chim. Phys. 71, 1551 (1974). 151. Hopster, H., Ibach, H., and Comsa, G.. J. Cutul. 46,32 (1977). 152. Pirug. G., Broden, G., and Bonzel, H. P., Proc. Int. Vuc. Congr.. 7th. 1977 p. 907 (1977). 153. Zhdan, P. A., Boreskov, G. K., Boronin, A. I., Egelhoff, W. F., h., and Weinberg, W. H., Surf. Sci.61, 25 (1976). 1.54. Morgan, A. E.. and Somorjai, G. A., Surf. Sci. 12,405 (1968). 155. Ertl. G., and Koch, J., Z. Phys. Chem. [N.S.] 69,323 (1970). 156. Koch, J., Thesis, Technische Universitat, Hannover (1972). 157. Demuth, J. E., J . Colioid Inrerfuce Sci. 58, 184 (1977). 158. Orent, T. W., and Hansen, R. S.. Surf Sci. 67,325 (1977). 1.59. Reed, P. D., Comrie, C. M., and Lambert, R. M., Surf: Sci. 64,603 (1977). 160. Blakely, D. W., and Somorjai, G. A,, Surf: Sci. 65,419 (1977). 161. Klein, R., and Shih, A,, Surf. Sci. 69,403 (1977). 162. Engel. T., von dem Hagen, T., and Bauer, E.. Surf. Sci.62, 361 (1977). 163. Schwab. G. M., and Gossner, K., Z. Phys. Chem. [N. S.] 16,39 (1958). 164. Daglish, A. G., and Eley, D. D., Acres Congr. Int. Cutul., 2nd. 1960 Vol. 2, p. 1615 (1961). 165. Tajbl, D. G., Simons, J. B., and Carberry, J. J., Ind. Eng. Fundurn. 5, 171 (1966). 166. Baddour, R. F., Mndell, M., and Heusser, U. K., J. Phys. Chem. 72,3621 (1968). 167. Stephens, S. J., J. Phys. Chem. 63, 188 (1959). 134. 135. 136. 137. 138. 139. 140. 141.


77

168. Alexander, E. G . , and Russell, W. W., J. Phys. Chem. 68, 1614 (1964). 169. Park, R. L., in “Fundamentals of Gas-Surface Interactions” (H. Saltsburg, J. N. Smith, and M. Rogers, eds.), p. 295. Academic Press, New York, 1967. 170. Kawasaki, K., Sugita, T., and Ebisawa, S., J . Chem. Phys. 44,2313 (1966). 171. Kawasaki, K., Sugita, T., and Ebisawa, S., Surf Sci. 8,485 (1967). 172. Ertl, G., and Koch, J., in “Adsorption-Desorption Phenomena” (F. Ricca, ed.), p. 345. Academic Press, New York, 1972. 173. Ertl, G.,and Koch, J., Caial., Proc. Inr. Congr., 5th. 1972 p. 969 (1 973). 174. Ertl, G . ,and Neumann, M., Z. Phys. Chem. [N. F.] 90, 127 (1974). 175. Engel, T., and Ertl, G., Chem. Phys. Leit. 54,95 (1978). 176. Engel, T., and Ertl, G., J. Chem. Phys. 69, 1267 (1978). 177. Close, J . S . , and White, J. M., J. Caial. 36, 185 (1975). 178. Matsushima, T., and White, J. M., J. Caial. 39, 265 (1975). 179. Matsushima, T., Almy. D. B., Foyt, D. C., Close, J. S., and White, J. M., J. Catal. 39,277 (1975). 180. Matsushima, T., and White, J. M., J. Carat. 40, 334 (1975). 181. Matsushima, T., Mussett, C. J., and White, J. M., J. Caial. 41, 397 (1976). 182. Ertl, G., in “Molecular Processes on Solid Surfaces” (E. Drauglis, T. R. Gretz, and R. J. Jaffee, eds.), p. 147. McGraw-Hill, New York, 1969. 183. Heyne, H., and Tompkins, F. C., Proc. R. Soc. London, Ser. A 292,460 (1966). 184. Sklyarov, A. V., Tretyakov, 1. I., Shub, B. A,, and Roginskii, S. Z., Dokl. Akad. Nauk SSSR 189,302 ( 1 969). 185. McCarthy, E., Zahradinck, J., Kuczynski, G. C., and Carberry, J. J., J. Catal. 39,29(1975). 186. Winterbottom, W. L., Surf Sci. 36,205 (1973). 187. Hori, G . K., and Schmidt, L., J . Caial. 38, 335 (1975). 188. Nishiyama, Y., and Wise, H., J. Caial. 32, 50 (1974). 189. Pacia, N., Cassuto, A., Pentenero, A., and Weber, B., J. Cafal.41,455 (1976). 190. Becker, C. A,, Cowin, J. P., Wharton, L., and Auerbach, D., J. Chem. Phys. 67,37 (1977). 191. Bonzel, H. P., and Ku, R., J . Vac. Sci. Technol. 9,663 (1972). 192. Bonzel, H. P., and Ku, R., J . Chem. Phys. 59, 205 (1973). 193. Bonzel, H. P., and Ku, R., Surf Sci. 33, 91 (1972). 194. Bonzel, H. P., and Burton, J. J., Surf: Sci.52,223 (1975). 195. Palmer, R. L., and Smith, J. N., Jr., Caial. Rev. 12,279 (1975). 196. Palmer, R . L., and Smith, J. N., Jr., J . Chem. Phys. 60,1453 (1974). 197. Palmer, R. L., J. Vac. Sci. Technol. 12, 1403 (1975). 198. Comsa, G., Proc. Vac. Congr., 7ih, 1977 p. 1317 (1977). 199. Ageev, N. V., and Ionov, N. I., Kine!. Katal. 14,687 (1973). 200. Zhdan, P. A,, Boreskov, G . K., Egelhoff, W. F., and Weinberg, W. H., Surf Sci. 61, 377 ( 1976). 201. Ivanov, V. P., Boreskov, G . K., Savachenko, V. I., Egelhoff, W. F., and Weinberg, W. H., J. Cafal.48,269 (1977). 202. Weinberg, 0. H., Egelhoff, W. F., Ivanov, V. P., Tatanrov, V., and Boreskov, G. K., Proc. h i . Vac. Congr., 7rh, 1977 p. 1 I51 (1977). 203. Kiippers, J., and Plagge, A., Z. Naturforsch. 3C,81 (1979). 204. Campbell, C. T., and White, J. M., J . Caial. 54,289 (1978). 205. Campbell, C. T., Shi, S. K., and White, J. M., Appl. Surf. Sci. 2,382 (1979). 206. Campbell, C. T., Shi, S. K., and White, J. M., in preparation. 207. McCarthy, E., Zahradinck, J., Kuczynski, G. C., and Carberry, J. J., J. Caial. 39,29(1975). 208. Bonzel, H. P., and Ku, R., J. Chem. Phys. 58,4617 (1973).

78


209. Bonzel, H. P., and Ku, R., J . Chem. Phys. 59, 1641 (1973). 210. Baddour, R. F., Modell, M., and Goldsmith, R. L., J . Phys. Chem. 74, 1787 (1970). 211. Hanson. F. V . , and Boudart, M.,J . Curd. 53, 56 (1978). 212. Pacia, N., and Dumesic, J. A., J . Curd. 41, 155 (1976). 213. Boudart, M., Collins, D. M., Hanson, F.V.,and Spicer, W. E., J . Vuc. Sci. Technol. 14, 441 (1977).

ADVANCES IN CATALYSIS, VOLUME 28

The Binding and Activation of Carbon Monoxide, Carbon Dioxide, and Nitric Oxide and Their Homogeneously Catalyzed Reactions RICHARD EISENBERG Department of Chemistry University of Rochester Rochester, New York AND

DAN E. HENDRIKSEN Corporate Research Laboratories Exxon Research and Engineering Company Linden, New Jersey

I. Introduction. . . . . . . . . . . . . . . . . . . . . . . Homogeneous Catalysis-Basic Concepts . . . . . . . . . 11. Carbon Monoxide . . . . . . . . . . . . . . . . . . . . A. Features of CO Coordination . . . . . . . . . . . . . . B. Models for CO Activation . . . . . . . . . . . . . . . C. Catalyzed Reactions of CO . . . . . . . . . . . . . . . 111. Carbon Dioxide . . . . . . . . . . . . . . . . . . . . . A. Carbon Dioxide as a Ligand in Transition Metal Complexes. B. Insertion Reactions of CO, . . . . . . . . . . . . . . . C. C02 Reduction and/or Incorporation . . . . . . . . . . IV. Nitric Oxide . . . . . . . . . . . . . . . . . . . . . . . A. The Coordination of Nitric Oxide . . . . . . . . . . . . B. Reactivity Patterns of Coordinated Nitric Oxide . . . . . . C. Catalyzed Reduction of Nitric Oxide . . . . . . . . . . . References . . . . . . . . . . . . . . . . . . . . . .

I.

. . . . .

_ . _ . . . . . . . . . . . .

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

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

19

_ . 81 . 84 . 84 . 88 . . . . . . . . .

95 119 121 128 141 144 146

149 157 164

Introduction

The ability of metal complexes in solution to catalyze reactions efficiently has served to make homogeneous catalysis a vigorously studied area over the 19 Copynght 0 1979 by Academic Press, Inc All nghts of reproduction in any form reserved ISBN 0-1 2-007828-7

80

RICHARD EISENBERG AND DAN E. HENDRIKSEN

past two decades (1-4). The efficiency of the homogeneously catalyzed reaction lies in several factors including the mild conditions under which the reaction may proceed, the selectivity of the reaction, and possibly its stereospecificity. Several important successes have been achieved in translating homogeneously catalyzed reactions into viable commercial processes, and mechanistic studies of homogeneous catalysis have led to a detailed understanding of catalyst requirements and individual reaction steps (5-8). With regard to the former, some of the successes include the Wacker process for olefin oxidation based on a palladium catalyst (9-10), the hydroformylation process using a rhodium(1) phosphine complex ( I I ) , and the carbonylation of methanol to acetic acid employing rhodium carbonyl iodide (12, 13). In the present review, we focus on the binding and activation of carbon monoxide, carbon dioxide, and nitric oxide by transition metal complexes, and examine selected reactions of these molecules which are, or may be, catalyzed by metal complexes in solution. These three simple oxides exhibit significantlydifferent coordination chemistries and diverse reactivity patterns when bonded to transition metal ions. The activation of each thus poses different problems in the field of homogeneous catalysis. By considering these three simple inorganic oxides together in this review, we shall be able to draw comparisons noting similarities and differences, in addition to examining each independently. From this exercise, we hope that a paradigm for the activation of small molecules will emerge. Reasons for interest in the catalyzed reactions of NO, CO, and C 0 2 are many and varied. Nitric oxide, for example, is an odd electron, heteronuclear diatomic which is the parent member of the environmentally hazardous oxides of nitrogen. Its decomposition and reduction reactions, which occur only in the presence of catalysts, provide a stimulus to research in nitrosyl chemistry. From a different perspective, the catalyzed reactions of CO and C 0 2 have attracted attention because of the need to develop hydrocarbon sources that are alternatives to petroleum. Carbon dioxide is one of the most abundant sources of carbon available, but its utilization will require a cheap and plentiful source of hydrogen for reduction, and the development of catalysts that will permit reduction to take place under mild conditions. The use of carbon monoxide in the development of alternative hydrocarbon sources is better defined at this time, being directly linked to coal utilization. The conversion of coal to substitute natural gas (SNG), hydrocarbons, and organic chemicals is based on the hydrogen reduction of CO via methanation and the Fischer-Tropsch synthesis. Notable successes using heterogeneous catalysts have been achieved in this area, but most mechanistic proposals remain unproven, and overall efficiencies can still be improved. Recent developments in organometallic chemistry and homogeneous

BINDING AND ACTIVATION OF CO, CO,, AND NO

81

catalysis suggest that significant progress in these catalyzed reactions is underway, and it seems appropriate at this time to bring these studies together, examine them in perspective and possibly discern what future directions these investigations will take. HOMOGENEOUS CATALYSIS-BASIC CONCEPTS In general, homogeneously catalyzed reactions can be viewed as occurring in a sequence of steps, each of which involves a change in one or more of the following properties of the catalyst metal complex : coordination number (i.e,, the number of ligands covalently bonded to the metal ion), metal ion oxidation state (based on the assumption that each metal-ligand bond is of the coordinate covalent type L+M), and valence electron count (the number of electrons of the particular metal ion plus the electrons involved in metal-ligand bonding). Of these properties, the metal ion oxidation state seems most susceptible to confusion, especially with ligands such as nitrosyl, 71-bonded allyl, and bridging hydride. We will avoid further clouding the oxidation state problem here, and suggest that the examples later in the text may help in clarifying it. Additionally, other references should be consulted (14, 15).

With regard to the valence electron count, this number determines whether the transition metal ion is using its full complement of valence shell orbitalsi.e., the five nd’s, the (n + l)s, and the three (n 1)p’s. If the valence electron count is eighteen, all of the orbitals are fully utilized in bond formation and electron pair storage, ihe effective atomic number (EAN) rule is fulfilled and the metal ion is said to be saturated. If it is seventeen, the metal ion is covalently unsaturated, and if it is sixteen or less, the metal ion possesses at least one vacant coordination site and is said to be coordinatively unsaturated. The importance of the valence electron count in homogeneously catalyzed reactions has been discussed by Tolman (7). The sequence of steps in homogeneously catalyzed reactions can be classified into three categories: the binding and activation of substrates, substrate transformation or coupling, and product elimination. Regeneration of the catalytically active species, if not implicitly carried out in the product elimination step, must also occur. The binding and activation of substrates involves the attachment of the reactant molecules to the catalyst complex. Except for those cases in which the metal complex donates an electron pair to the substrate, as for example in metal hydride formation from H’, the catalyst center must possess some degree of unsaturation. For even electron transformations (i.e., those not involving radicals) the unsaturation is of the coordinative type. Substrate activation is simply the process of making a chemical entity

+

82


more reactive than it would otherwise be in the absence of the catalyst, or imparting to this entity a new or different reactivity pattern. Activation is achieved by changing the electron density distribution in the substrate, and from an orbital standpoint this usually means transferring electron density from substrate bonding orbitals to the catalyst center and back-donating electron density from the complexed metal ion into orbitals that are antibonding with regard to substrate bonds. The catalyst-complex substrate interaction is thus intrinsically synergistic in which substrate metal donation and metal substrate backbonding reinforce each other cooperatively. For an X-Y substrate, activation in this manner can ultimately lead to cleavage of the X-Y bond in a process called oxidative addition (5, 6, 16). The metal complex center has undergone an increase in both coordination number and oxidation state since in this formalism the electron pairs in metal-ligand bonds are associated with the ligands (14, 15). While substrate activation by oxidative addition occurs in this way for H2 (16), the term oxidative addition really represents a stoichiometric transformation, and does not necessarily imply a specific mechanism. In fact, studies over the past decade have shown that the interaction of X-Y with a metal center to give X-M-Y proceeds by any of a variety of mechanisms determined by the substrate and the metal complex (16-18). However, once the X-M-Y species is formed, the X-Y substrate can be viewed as activated. In the case where the XY bond is a multiple one, binding and activation of the substrate leads to a reduction in the XY bond order (19). This is because the substrate orbitals involved in accepting electron density from the metal center are of the n* type. For donation of electron density, the substrate employs either a nonbonding lone-pair orbital that gives end-on coordination (1) or a nb-type orbital that yields side-on or q2 attachment. Since the oxides CO, CO, , and NO all have multiple bonds, the notions outlined above will be important in considering their activation. --+

--+

M +X=Y (1)

M

41 (2)

Once activated, the substrates are transformed via a number of different possible steps including ligand migration, insertion, elimination or extrusion, and external attack on bound substrate. Of these, the last is most easily envisioned-a reagent not coordinated to the metal center of the catalyst attacks the bound substrate whose coordination has rendered it chemically reactive. Insertion reactions (20-22), which are among the most widely found of the substrate transforming steps, are represented by the general conversion M-X + Z + M-Z-X. The inserting moiety Z may be either a free species

BINDING AND ACTIVATION OF CO, COz, AND NO

83

in solution or a prior coordinated ligand. In the latter case the insertion is referred to as rnigralory. Two especially important classes of migratory insertions are the formation of metal acyls from carbonyl insertion into metalalkyl bonds (20, 20a, 22) and the formation of metal alkyls from olefin insertion into metal-hydride bonds (21). Both of these reactions have been studied using kinetic and stereochemical methods of analysis, and a reasonable understanding of them is emerging. Extrusion or elimination reactions represent the microscopic reverse of insertions. In decarbonylation reactions the alkyl group migrates from the acyl carbon atom to an adjacent but unoccupied coordination site of the metal center giving rise to an alkyl carbonyl species (23). The reverse of olefin insertion into M-H is fi-elimination, which represents the primary pathway for decomposition of metal alkyls (24, 25). This reaction can be suppressed by using alkyl ligands containing no 8-hydrogen atoms or by preventing the metal-alkyl group from achieving the planar M-C-C-H arrangement requisite for elimination through the use of metallocycle ligands (24). The product elimination step proceeds with cleavage of the catalystsubstrate bonds. This may occur by dissociation, solvolysis, or a coupling of substrate moieties to form the product. The last of these involves covalent bond formation within the product, and corresponds to the microscopic reverse of oxidative addition. Upon reductive elimination both the coordination number and formal oxidation state of the metal complex decrease. In most homogeneous catalytic processes, the product elimination step, while essential, is usually not rate determining. The larger kinetic barriers are more frequently encountered in substrate activation and/or transformation. The fact that each of the reaction steps outlined above has an “inverse” that is energetically accessible leads readily to the formulation of catalytic cycles. This is not simply a restatement of microscopic reversibility. The distinction is that the activation energies for all of these reactions are not prohibitively high so that catalysis can occur. While catalysis is a kinetic phenomenon, one must come to grips with certain thermodynamic considerations in analyzing any potential catalysis. For a multistep catalytic cycle, the sum of the free energy changes for the individual steps must equal the free energy change of the catalyzed reaction. In order to achieve catalysis under mild conditions, each step should be slightly favorable from a thermodynamic standpoint, and certainly none should be more than slightly unfavorable. This means that, except for the most exothermic reactions, no step should become too favorable as this would necessitate a positive AG for a different step. Even in a highly exothermic reaction, an extremely favorable step may lead to a deep potential well on the surface of the catalyzed reaction, resulting in a large and in-

84


hibiting barrier to catalysis. Thus thermodynamic considerations of individual steps are important in an analysis of real or proposed catalysis. The basic notions of homogeneous catalysis have been amplified and discussed more extensively elsewhere, as have examinations of individual reaction steps. In briefly reviewing these concepts, we are constructing a framework in which to consider the activation of CO, C 0 2 , and NO by metal complexes in solution and homogeneously catalyzed conversions of these simple oxides. II. Carbon Monoxide

The binding of CO to transition metals and the formation of metal carbonyl compounds have been recognized since Mond’s famous discovery of Ni(CO), in 1891 (26). Since that time the chemistry of metal carbonyls has developed into one of the most vigorously studied areas of inorganic and organometallic chemistry. Many new and interesting compounds have been synthesized in these studies, novel and unusual structures have been determined, and a rich reaction chemistry has been recorded. The growing research effort in homogeneous catalysis has been closely intertwined with these studies because of the catalytic properties that certain metal carbonyl compounds exhibit. Each year thousands of publications appear on metal carbonyl chemistry. Many excellent treatises and reviews have appeared (27-32). An important current stimulus to studies of metal carbonyl chemistry revolves around the need to develop new and efficient procedures for the facile transformation of CO into hydrocarbons and organic chemicals, and to understand the ways in which existing catalyzed reactions occur. Such conversions, which are discussed more fully under the section on CO reduction, are of key importance in utilizing coal more effectively in the nation’s energy program.

A. FEATURES OF CO COORDINATION From the standpoint of the present review, we will focus on compounds in which “unusual” and potentially reactive features of CO coordination are revealed, and on certain classes of reactions in which the conversion of CO to other chemical entities is achieved. For the purposes of determining what a reactive mode of CO coordination is and in what way the CO ligand can be activated, it is first instructive to briefly review the primary modes of carbon monoxide coordination. These are illustrated as Lewis structures

BINDING AND ACTIVATION OF CO, C O l , AND NO

85

(3)-(5) and are called the terminal, bridging, and triply bridging modes, respectively.

The terminal mode of coordination (3) is by far the most frequently encountered. It involves the synergic interactions of a-donation and n-back bonding between metal and ligand. From these interactions, the stretching frequency of a coordinated carbonyl is reduced from that of free CO (2155 cm-') to the range 2100-1850 cm-l. In fact, the sensitivity of v, to the backbonding interaction allows it to be used as a probe in determining the extent of this interaction. For example, in the isoelectronic and isostructurd series Mn(CO),+, Cr(CO), and V(CO),-, the carbonyl stretching frequencies are 2096, 2000, and 1859 cm-', respectively (28). This trend, which represents a weakening of the C - 0 bond, is explained by the view that as the positive charge on the central metal atom decreases (Mn > Cr >V), the energy of filled metal d orbitals rises relative to the energetically higher n*CO acceptor orbitals, producing an increase in d, --t n* backbonding (Mn < Cr V). This same explanation predicts an inverse correlation for the metal-carbon stretching frequencies, and this is precisely what is observed with values for vM+ of 416, 441, and 460 cm-', respectively. The greater backbonding for V over Cr over Mn corresponds to a relative increase in the importance of resonance structure (3b) in a bonding description of the MCO unit. This line of argument assumes for carbonyls that force constants can be derived directly from observed CO stretches and that there exists a direct force constant-bond order correlation. Although not rigorously correct, the validity and value of this approach has been shown by many investigators, especially Cotton and his co-workers (33). While the lowering of the carbonyl stretching frequency reflects a reduction in the C-0 force constant and bond order, the resultant carbonyl ligand can hardly be construed as activated. In general the terminal M-CO unit is unreactive to most reagents. One of the reasons for this lies in the fact that the bond order in a terminally bonded CO is not greatly reduced, i.e., the CO triple bond has not been reduced beyond a bond order of 2, with the exact value depending on the particular compound. A different reason for the lack of reactivity of terminal carbonyl ligands is based on the notion that an unreactive charge distribution results from the interaction that leads to the bond order reduction. This view requires

-=

86


recognizing how carbonyl molecular orbitals that interact with metal valence shell orbitals are polarized. Although free CO has a very small dipole moment, the two component atoms have vastly different electronegativities resulting in bonding molecular orbitals polarized toward oxygen, and antibonding m.o.’s polarized toward carbon (34). Upon CO coordination to a metal, donation of an electron pair from a carbon based lone pair orbital is compensated for by back donation into n* orbitals of the carbonyl ligand (35).The latter interaction, shown in orbital terms in (6), not only achieves a delocalized electronic structure over the three atom MCO moiety, but also maintains a relatively neutral charge distribution by placing greater electron density on carbon than on oxygen. The buildup of a positive charge on the o-donating carbon is thus avoided by effective backbonding.

The bridging mode of CO coordination (4) is sometimes called the “ketonic” mode because of its similarity with organic carbonyl compounds. In this manner of coordination, the carbonyl forms a covalent bond with each of the bridged metals, the CO bond order is reduced to two-or with additional backbonding, less-and vco falls into the range 1850-1700 cmHowever, some important differences between (4) and an organic carbonyl group exist that give rise to different relative reactivities. One of these differences derives from the fact that the metal atoms are less electronegative than carbon, leading to a polarization of the M-C covalent D bonds in (4) toward the carbonyl carbon atom. Another factor is that the occurrence of the bridging mode of CO coordination relates to the size of the bridged metal atoms (29).Initially it was thought that (4) occurred only in compounds of the smaller first transition series elements but recent structure studies, such as of Rh4(CO)12and Ir,(CO),(PPh,), , have shown that CO may also bridge second and third transition series elements in select instances (28). In all of these cases, the bridge exists between metal atoms that are bonded directly to one another. There is thus bonding electron density along the metal-metal axis at the point closest to the bridging carbonyl. Often the carbonyl bridges occur in pairs or complementary sets over two or more metal atoms. Examples include Fe,(CO),, , (rf’-CSH,),Fe2(CO), (bridged isomers) and Rh4(CO),2. In the first two, the bridging carbonyls

BINDING AND ACTIVATION OF CO, C 0 2 , AND NO

87

are in pairs (7) while in Rh4(C0)12 three bridging CO's span the edges of a triangular array of Rh atoms, (8) (28).

0 (7)

(8)

Unlike organic carbonyl compounds, the MCM bond angle in bridging carbonyl compounds is significantly less than the value expected for an sp2 hybridized C, being, for example, 83" in Co2(CO), (36) and 77" in Fe2(C0)9 (37). Finally, the four atom system M(C0)M appears to have complex multicenter interactions that have been discussed qualitatively by Braterman (38). All of the above features serve to make briding carbonyl groups less reactive than organic carbonyls to nucleophilic attack at carbon. The triply bridging mode of CO coordination (5) is relatively uncommon although it has been observed in an increasing number of structures. A simple valence bond structure is inadequate to describe the bonding in the triply bridging carbonyl, and one is required to use a molecular approach delocalized over the carbonyl and metal atom framework. As with regularly bridging carbonyls, the MCM angles in (5) are acute. Examples of triply bridging COs are found in [($-C,H5)Fe(CO)l4, (q5-C5H5)3Ni3(C0)2 and Rh,(C0)16 (28). While the simple valence bond structure is inadequate for (5) (38), it is clear that significant reduction of the CO bond order has occurred. Typical values for the carbonyl stretches of p3-CO's are around 1600 cm-'. One aspect of metal carbonyl chemistry that should be mentioned in surveying the more commonly found modes of CO coordination is the stereochemical nonrigidity of carbonyl clusters. This aspect has received considerable attention over the past decade, especially as 13C nmr instrumentation has become more readily available. In many carbonyl clusters, terminal and bridging carbonyls as established by x-ray structural studies are equilibrated on the nmr time scale ( 3 7 , 3 9 4 1 ) . The manner of equilibration takes place in a concerted way in order that each metal center maintains a constant electron count. For example, bridge *terminal interconversion, (I), proceeds via complementary unsymmetrical CO bridges.

88


Cotton has discussed this phenomenon extensively (37). For complementary sets of more than two carbonyl bridges, the equilibration is also probably concerted. All of the carbonyls in Rh,(CO), appear magnetically equivalent at room temperature, with bridge terminal interconversion occurring via the formation of an unbridged species whose structure is analogous to Ir4(C0)12 (39). For the equilibration processes to take place on the nmr time scale, the energetic barrier for interconversion of CO bonding modes cannot be great (usually c 22 kcal/mol). While the fluxional properties of metal carbonyls will not be of concern to us here, they illustrate the facility for interconverting CO bonding modes, and suggest that if CO activation can be achieved beginning with a particular mode of coordination, carbonyls bound in other ways to metal atoms may also be “activatable.”

*

,

B. MODELS FOR CO ACTIVATION Having examined the more commonly found modes of CO coordination, we now consider the specific question of carbon monoxide activation. There appear to be two strategies that can be followed. In the first, maximal reduction of the CO triple bond is achieved upon interaction with two or more metal centers (42), and in the second a reactive charge distribution is effected, possibly with the aid of an external Lewis acid. The former of the two strategies most closely relates to dissociative CO adsorption on metal surfaces and heterogeneous catalysts, while the latter consists of enhancing the reactivity of the carbonyl ligand to nucleophilic attack at carbon, regardless of whether reduction or oxidation of the carbonyl ligand ultimately occurs. Based on values of vco, it would appear that triply bridging carbonyls among the more common modes of CO coordination most closely approach the notion of activation by maximal CO bond order reduction. However the reactivity of the p3-carbonyl ligand has not been demonstrated in a systematic way, and the lack of many triply bridging carbonyl species cannot be taken as a definite indication of enhanced reactivity. We now turn our attention to other possible modes of CO coordination in which vco has been significantly reduced upon binding. One intriguing report in which the carbonyl stretch is unusually low (1645 cm- ’) is that of ) ~ = bis(dipheny1phosphino)methane) (43). The strucMn, (CO) ( d ~ m (dpm ture of this complex, shown as (9),reveals that the bridging carbonyl is o-bonded to one of the Mn atoms and n-coordinated to the other.


P 0

\

I

89

P

1P

P

From an electron-counting standpoint, the bridging carbonyl in (9) can be viewed as a four-electron donor, two to each Mn, with concomitant backbonding from both metal atoms into carbonyl n* orbitals. This contrasts sharply with regular carbonyl bridges which formally donate one electron to each metal atom and have relatively little backbonding. The 4-electron bridging carbonyl in (9) may thus be an activated CO, and its formation can be envisioned by removing a carbonyl ligand from a saturated binuclear unit that is held together by bridging ligands. It is interesting to note that the compound Fe,(CO)s(dfpa)2 (dfpa = bis(Muorophosphino)methylamine or (PF2)2NMe)(44)has a structure virtually identical to that of the Mn system above, except that the bridging carbonyl appears normal and does not have to adopt the potentially activating o,n-mode of coordination because the Fe complex possesses two additional electrons from the metal atoms. Recently the existence of a 4-electron bridging CO has been proposed for the Pt complex Pt,Cl,(dpm),(CO) that shows vm of 1638 cm-' (45). The formation of this carbonyl complex, which is reversible, takes place by addition of CO to the metal-metal bonded species Pt,Cl,(dpm), and is thought to result in two coordinatively saturated Pt centers as shown in (10).

90

RICHARD EISENBERG AND D A N E. HENDRIKSEN

However, a structural study will have to verify this proposal. In a separate study on closely related Pd systems, Balch and co-workers have synthesized and characterized analogous bridging CO and CNR complexes with the former having vco of 1704 cm-' (46).A structural investigation of the latter, however, reveals a symmetric bridge and the absence of a direct metal-metal bond. Balch's proposal for the carbonyl bridged system is thus of a symmetric CO bridge connecting two nonbonded metal atoms, the first time this mode of coordination has been observed without the support of a metal-metal bond (46). The structure of Pd,Cl2(dpm),(CO), which has just been reported (46a,b)confirms this proposal. Within the last year further developments on CO bridged binuclear syst e m have appeared. Mague, Cowie, and Sanger (46c) have reported the structure of a rhodium dpm complex in which a bridging CO occupies the proposed active site of a binuclear A-frame type structure (46d).The A-frame arrangement is shown as (lOa), and the CO-bridged structure is (lob). While v m of the bridging carbonyl in (lob) is 1863 cm-', suggesting little activation, a related Ir system exhibits significantly different properties with the carbonyl stretching frequency of the p-CO ligand at 1680 cm-' (46e). The structure, however, is similar to (lob).

( 10 a)

x = sz;

c1-

Perhaps the most interesting class of compounds from the viewpoint of models for CO activation are those systems in which coordinated carbonyls have interacted with Lewis acids such as trialkylalanes (47-54). Research on these compounds, especially by Shriver, Burlitch, and their co-workers, has shown that the alanes bind to oxygen lone pairs of carbonyl ligands with concomitant lowering of vco by over 100 cm-'. For example, in [(~5-C5HS)Fe(CO)(p-COA1Et,)]z which has structure (ll), the bridging carbonyls show a single stretch at 1682 cm-' compared with a value of ca. 1800 cm-' in the parent carbonyl compound (50).

BINDING AND ACTIVATION OF CO, CO,

, AND NO

91

AlEt,

/

Et&' ((1)

Shriver has also observed that coordination of a carbonyl by a Lewis acid serves as a driving force to convert a terminal carbonyl into bridging coordination (49,51).In the presence of AlR, , the unbridged isomer of [CpRu(CO),], is converted entirely into a bridged carbonyl species with alanes bonded to the oxygen atoms of the bridging carbonyls. This is not surprising since bridging carbonyls are expected to have more negative oxygen atoms and thus act as better Lewis bases than terminal carbonyls. In select instances, however, terminal carbonyl ligands have been observed to bind AlR, Lewis acids with highly significant reductions in vco. For example, Burlitch (52)has found that (q5-C5HS)W(CO),- forms an adduct with AlPh, that exhibits vco at 1600 cm-' owing to the M-CO-Al moiety, and Kotz (53) has reported a similar lowering with AlMe, and a neutral Mo carbonyl complex. The synthesis and structure of a closely related Al-OC complex has also been reported by Burlitch (54). The complex is A1[(q5-C5H5)W(CO),],(THF), with octahedral coordination about A1 comprised of one oxygen-bonded CO ligand from each organometallic moiety and 3 THF ligands. The structure is as shown in (12). THF,

(C,H,)(CO),W-C--~&

THfiO-C-W(CO)2(C5H5) A1 \THF \

i:

\

(CsHdW(C0)z (12)

-

While the W-(2-0 angles remain essentially linear ( 176"),the C-0-A1 angles range between 140 and 162". The C-0 distances in the bridging carbonyls are nearly 0.1 longer than in the terminal CO's, while the W-C distances for these ligands are 1.85 A versus 1.95 for the terminal carbonyls. The significance of these results is clear. The electron acceptor property of the Lewis acid coordinated carbonyl is greatly enhanced. This is not only

92


evidenced by the structural parameters and by significant lowering of vco of the Lewis acid coordinated carbonyl, but also by the fact that the stretching frequencies of other terminal carbonyls in these compounds increase indicating that electron density has been removed from the metal center. The possible significance of Lewis acid binding of CO will be seen below in a recent study of CO reduction using metal clusters. The problem of developing a reactive charge distribution on a coordinated carbonyl is dependent on the type of attack desired and the reactivity of the reagent. Here we focus solely on even electron modes of reaction. In addition to the weak oxygen basicity, the carbonyl ligand exhibits a susceptibility to nucleophilic attack at C that varies widely as a function of the metal, the complex charge, the ligand array, and the medium. This site of attack is extremely important since all redox reactions of coordinated CO necessarily involve a change in the carbon atom’s oxidation state. With strongly basic and nucleophilic reagents, little activation of the carbonyl ligand is required, but attack at C occurs readily. Two classic examples are shown as (2) and 13) (55956). Fe(CO),

+ 20H--+ HFe(CO),- + HC0,-

(2)

0-

Cr(C0)6

+ LiPh+

(CO),Cr-C

//

Me30+ __t

\Ph (CO),Cr=C

/OMe

(3)

b h

In the first, Fe(C0)5 is converted into the carbonyl hydride anion by reductive decarboxylationfollowing initial OH- attack on a carbonyl carbon (55). The carbonyl group is thus oxidized to carbonate in basic medium. The second reaction, (3) (56), is one of E. 0. Fischer’s celebrated carbene forming reactions (57). Phenyl lithium reacts with Cr(CO), leading to an anionic acyl complex. A subsequent alkylation step using Me,Of yields the methoxyphenylcarbene complex Cr(CO),(CPh(OMe)). Recent attention on reactions of this type with unactivated carbonyls has focused on the possibility of using an active hydride as the attacking nucleophile. The resultant formyl species is then viewed as an intermediate on the path to CO reduction. For example, Casey and Neumann (58) have shown that trialkoxyborohydride reacts with various metal carbonyls according to (4)to give identifiable, and in one case, isolable, formyl complexes. HB(OR)3-

+ L,M(CO)

-+ B(ORL

+ L,M-C

p H ‘

-

(4)

BINDING AND ACTIVATION OF CO,

COz, AND

NO

93

Prior to this report the only known formyl complex had been Fe(CO),(CH0)- synthesized by Collman and Winter from Fe(C0);- and acetic formic anhydride (59). Within the last year, additional reports of formyl complexes have appeared (59u-594. Most notable among these is the work by Casey, Gladysz, and their co-workers on the neutral formyl complex (C,H,)Re(CO)(NO)(CHO) made by the reaction of R,BH- with the cationic dicarbonyl system (C,H,)Re(CO),(NO)+ (59b,c).This neutral formyl species is, like its anionic predecessors (58), coordinatively saturated. A significant feature of the chemistry of the limited number of formyl complexes reported to date is their ability to act as hydride donors, thus reversing the formation process ( 5 9 4 . However, Casey has just reported (594 the bimolecular reaction of the neutral formyl species to form a dimeric metallo ester type complex (12a) which is important in modeling CO reduction chemistry, and is discussed below.

(120)

The fact that there is such a paucity of metal formyl complexes is both interesting and significant because of the proposed intermediacy of coordinated formyl in CO reduction, and the sharply contrasting abundance of metal acyl complexes. Since many of the acyl complexes are known to form by migratory insertion of CO in an alkyl carbonyl complex (20,20a, 22), the lack of formyl complexes from hydride carbonyls relates to the thermodynamic difference in the equilibrium ( 5 ) when Y is alkyl and when it is hydride. M-CO

I

Y

-L M-C

P

(5)

Y '

It is significantthat the only stable formyl complexes isolated to date (58-60) are coordinatively saturated, which eliminates the possible conversion to a carbonyl hydride without the prior loss of a ligand. The unfavorable thermodynamics of (5) for formyl formation are a necessary consideration in developing schemes for CO reduction by this method. Reaction of coordinated CO with less active nucleophiles can take place when the carbonyl ligand is sufficiently activated. The manner in which this activation occurs is by a reduction in the backbonding interaction (60) that may be achieved when the CO-bound metal ion is in a higher oxidation state.

94


It is well known, for example, that CO in the presence of water is an effective reducing agent for more oxidized metal ions (27, 61). The mechanism for this reduction has long been established, and the case of rhodium, as studied by James and co-workers (62), is typical. These investigators find that coordination of CO to rhodium in the + 3 oxidation state is followed by nucleophilic attack of water and reductive decarboxylation to give Rh(I), COz, and protons as shown in (6). Rh"'

+ CO -+

Rh"'-CO

Rh"'-C

B

J.

'OH

Rh'

+ CO, + H+

(6)

The binding of CO to Rh(II1) in this example or to a more oxidized metal ion in general produces a reactive charge distribution since a-donation from carbon is not adequately compensated for by backbonding into the C-polarized .n* orbitals of carbonyl. Thus a residual positive charge develops at the carbonyl carbon, enhancing its electrophilicity. The stretching frequencies of Rh(II1) and Ir(II1) coordinated carbonyls, for example, are in the range of 2150-2040 cm- with greatest propensity to nucleophilic attack found amongst the systems having the highest vco values. Additional activation of coordinated CO is possible in strongly acidic media in which protonation of the terminal oxygen lone pair can further enhance the charge polarization of the carbonyl ligand and the susceptibility of the C atom to undergo nucleophilic attack. While stable species showing this type of binding have not been isolated, their possible kinetic role cannot be ruled out. In a recent report, protonation of a bridging carbonyl oxygen has been described by Shriver and co-workers ( 6 2 ~ )The . compound is proposed to have structure (12b), and represents an extension of Shriver's previous studies on Lewis acid adducts of bridging carbonyls (49-51). The compound is prepared under rigorously anhydrous conditions and the protonation of the 0 atom is ascribed to the enhanced basicity of 0 in the bridged bonding mode (4).

BINDING AND ACTIVATION OF CO, COz, AND NO

95

The oxygen end of the carbonyl ligand is a hard, albeit weak, base (48) and is thus compatible with binding H ' . The major problem with this type of carbonyl activation is that because the M-CO bonding has been weakened via a reduced synergic interaction, carbonyl dissociation may become competitive. An example of the way in which activation in the acidic regime can be used is provided by the as yet hypothetical case of hydroxycarbene formation. This species is generally viewed as a key intermediate in CO reduction, whether by homogeneous or heterogeneous catalysis. Previously we noted that the formyl ligand is produced via hydride attack on bound CO because of the great reactivity of the hydride employed. We propose that in strong acid medium with protonation of the carbonyl, a nonnucleophilic hydride may be used for the corresponding transformation, and that if the hydride is prior coordinated to the metal carbonyl, a facile, direct conversion to a hydroxycarbene species may be achieved. This proposal as embodied in (7) may have interesting ramifications in the catalyzed reduction of CO. The formation of the 0-protonated carbonyl has been suggested in a study by Ellis (63) on the reaction of V(CO),(diars)' - (diars = 1,Zbis(dimethylarsino)benzene) with HX, which is described under catalyzed oxidations of co.

M-co I

H

H+

M=c

/o-H+ '€I

T h s section has attempted to delineate the possible ways in which CO activation can be achieved by discrete metal complexes in order that homogeneous catalysis of CO reactions can be better understood. Principal means of activation are by significant bond order reduction and/or development of reactive charge distributions on the coordinated carbonyl. Oxidation or reduction of the CO ligand will transpire at carbon, and the primary mode of attack at that site will be by nucleophiles. C. CATALYZED REACTIONS OF CO In this section the focus of attention will be on homogeneously catalyzed reactions of CO that involve either reduction or oxidation. The products of these reactions will be hydrocarbons and alcohols on the one hand, and C 0 2

96


and carbonates on the other. In surveying these reactions, we will not discuss as a separate entity carbonyl insertion reactions, which have been well studied mechanistically, and which play a key role in many catalytic processes such as hydroformylation ( I I ) , methanol carbonylation (12) and acrylic acid synthesis (22). In these reactions, the carbonyl insertion into a metalcarbon bond is more aptly described as an alkyl group migration from the metal to the carbonyl carbon. In decarbonylation reactions, on the other hand, the microscopic reverse of this migration serves as the critical step. A number of excellent reviews covering carbonyl insertion and “de-insertion” reactions have appeared and these are cited (20, 20a, 22, 64, 65). In the catalyzed reductions of CO we concentrate mainly on H, as the reductant, while in the catalyzed oxidations we focus primarily on the role of water and hydroxide. 1. Reduction of CO by Hydrogen

The reduction of CO by hydrogen represents one of the most interesting and challenging problems to the practitioner of homogeneous catalysis today. The current impetus for study of this problem comes mainly from the need to develop efficient utilization of our coal and other carbon-rich, hydrogen-deficient resources for fuels and chemicals. While the solution to this multifaceted problem clearly involves technology and economics as well as chemistry, the orientation here is on the desired chemical transformations and how they may be achieved in the laboratory using homogeneous catalytic systems. The conversion of coal into hydrocarbons and organic chemicals begins by reacting coal with steam and oxygen at elevated temperatures to produce a mixture of gases composed mainly of CO and hydrogen. The reduction of CO by H, can then lead to a variety of different products, as represented by equations (8)-(12) (66-71).

+

(2n

+

+

(8) (9) (10) (1 1)

1-E

(12)

3Hz CO --t CH4 HzO 2H2 CO + CHjOH 1)HZ nCO ---t C.H,.+Z + nH2O 2nHz nCO --f C,H,. + nHzO or C.Hz,+IOH + (n - 1)HzO 2CO + CHI 3Hz

+ +

+

+

OH

Prior to carrying out the reduction, the required ratio of H, : CO is attained using the water gas shift reaction, (1 3), which is discussed under CO oxidation.


97

The hydrogenation of CO using heterogeneous catalysts has been known since the early part of this century. In 1902 Sabatier and Senderens reported the synthesis of methane by (8) using a nickel catalyst (71, 72). After 75 years, Ni still remains the element of choice in catalyzing the methanation reaction. As a heterogeneous catalyst, it is generally supported on materials such as alumina, or used in alloys and as Raney Ni. The formation of methanol, (9), was carried out in 1923 by BASF using a ZnO-Cr,O, catalyst (70).This conversion, which is highly specific, is run under relatively vigorous conditions (250-400°C and 100-600 atm). In the early 1920s, Fischer and Tropsch reported the synthesis of higher hydrocarbons from CO and Hz using alkalized iron turnings as the catalyst (70, 73).This work inaugurated an extensive research effort by Fischer and his co-workers on the conversion of coal to hydrocarbon fuels that eventually led to its utilization by Germany during World War I1 (74). After a period of relative inactivity, interest in the Fischer-Tropsch (F-T) synthesis has renewed over the past decade as petroleum reserves have diminished and prices have risen sharply. At present, however, there is only one major commercial facility for Fischer-Tropsch synthesis, which is at Sasolburg, South Africa, and a number of pilot plant projects and demonstration units that have been developed in recent years for the purpose of improving the efficiency of the F-T process (75). All of these projects utilize heterogeneous catalyst systems, mostly based on iron. The mechanisms of the hydrogenation of CO have been studied and debated since the 1920s when the reactions were first reported. An initial proposal by Fischer and Tropsch (73) involved the formation of a surface carbide, but this scheme was subsequently rejected for a number of reasons, including the fact that ruthenium, which does not form a carbide, catalyzes F-T synthesis. Today the two most widely postulated models for F-T chemistry are those put forth by Storch et al. (74)and by Pichler (76). Both involve oxygen-containing surface intermediates suggested as CHzO. In the former, a surface-bound carbonyl is hydrogenated to give a bound hydroxycarbene species, and chain growth occurs via a condensation of two hydroxycarbene moieties followed by hydrogenation as illustrated in (14) and (15). 0

C

I

M+2H-

H , C ,OH

It

M

(14)

98


The hydrogens in these reactions are surface bound. Termination of the chain is proposed to occur via one of several possible paths given as (16) and (17). R--C&..,OH L

RCHzCHO R-CH,

‘C

,OH

II

M

-

-

acids, esters, alcohols

(16)

RCH,CH,OH + hydrocarbons

H,

,OH

;i

M

+ R’=CH,

S R C H ,

(17)

Pichler’s mechanism differs primarily in the manner in which chain growth takes place. Pichler proposes that repetition of the sequence of carbonyl insertion into a M-C bond followed by reduction of the resultant acyl leads to formation of the higher hydrocarbons as in (1 8). R

R

I co

M’

R

I

tco

oc,M/

C==o

R

I

0

& c,M/

CHOH Hz

-H,d

0 CLM/

-+etc.

(18)

Although chain growth is not a feature relevant to methanation, the initiation and termination steps of the Anderson model for F-T synthesis are believed by at least some workers in the field to be applicable to the mechanism of the highly specific methanation reaction (71). The formation of methane is proposed to follow from the surface bound hydroxycarbene species by (1 9).

The hydroxymethyl intermediate could either form methanol or upon further reaction with hydrogen, eliminate H,O and go on to form methane. In presenting the proposed mechanisms of CO hydrogenation reactions, we have omitted detailed examination of any experimental evidence, which can be found elsewhere (70, 71,74, 76), and have attempted to summarize only select conclusions. It is important to remember that these proposals are still only that, with definitive proof lacking, and that alternative postulations remain viable. Surface bound intermediates have not been conclusively established as to either composition or reactivity. With the established success of heterogeneous catalysts in the hydrogenation of CO via methanol synthesis, methanation, and F-T synthesis, it is justifiable to question the interest in investigating these reactions under


, AND

NO

99

homogeneous catalytic conditions. A number of answers exist, some of which are provided by considering the deficiencies of present heterogeneous catalyst systems. For example, while Ni is preeminent as the metal for catalyzing methanation, it is easily poisoned by sulfur. Moreover, since reaction (8) is so exothermic (AG" = - 33.9 kcal/mol at 300"K), dissipation of the heat released under conditions of catalytic methanation (typically 250-400°C) becomes problematic, and may render the catalyst inactive (71). For methanol synthesis, (9), the primary problem with existing catalyst systems lies in the vigorous conditions required, and in particular, the higher pressures (70). This appears true even for a newer Cu-ZnO catalyst that operates in the range of 50-100 atm. Also, it should be noted that AG" for reaction (9) becomes positive at 425"K, resulting in a very unfavorable equihbrium at the temperatures required for methanol synthesis (68).The primary problem with the F-T synthesis is the relative lack of selectivity in the products obtained (68, 70). A wide range of products is obtained including paraffins, olefins, and oxygenated organic compounds with a broad distribution of molecular weights. While reaction conditions and catalyst design influence these distributions, the ability to form selectively only a single product or a small range of products has not yet been achieved. Thus, while many scientists believe that the major problems with CO hydrogenation are technological, none will deny that homogeneous catalysis may provide important breakthroughs in the chemistry of CO reduction, its understanding and its possible utilization. In general, homogeneous catalyst systems appear more resistant to poisoning and offer the possibility of greater efficiency of operation via more specific product distribution. Homogeneously catalyzed reactions are often performed under significantly milder temperatures, and can render the heat dissipation problem more tractable through yigorous stirring. In addition, homogeneous catalyst systems have great flexibility of design (they represent the ultimate in dispersion) and at the same time, have consistency of character in preparation. The latter allows for the transferability of data between studies for comparison and interpretation. Vannice has described this inability to compare results meaningfully as a problem with earlier studies of heterogeneously catalyzed CO hydrogenation (77). Finally and perhaps most importantly, homogeneous catalyst systems make detailed mechanistic studies of these reactions more feasible by chemical and physical methods. Solutions can be examined spectroscopically, kinetic measurements can be performed, product distributions can be readily analyzed, labeling studies can be carried out, and model systems can be devised, synthesized and characterized. Once initimate details of the chemical transformations are known, a reexamination and modification of the catalyst can be done in a systematic way.

-

100


In this regard, it is noteworthy that while surface bound hydroxycarbenes are postulated species, discrete complexes containing hydroxy- and alkoxycarbenes have been known since E. 0. Fischer’s studies beginning in 1964 (56, 57). These complexes are possibly analogous to proposed surface intermediates, and their chemistry may model some of the heterogeneously catalyzed transformations. Coupling of alkoxy carbenes, for example, gives dialkoxy olefins as observed in (20).

However, condensation of hydroxycarbene ligands, as in (15), has not yet been observed, and hydroxycarbene ligands in general show a propensity to eliminate with a hydrogen shift as aldehydes (57). With all of the reasons for studying CO reduction by H2 using homogeneous catalysts, it is somewhat surprising that so little has been done to date. However the problem is rapidly gaining greater attention, and devefopments over the past 4 years have been very encouraging. We now briefly summarize those hydrogenations of CO that are homogeneously promoted or catalyzed that have been reported to date with very recent developments included at the end of this section. In 1976 Bercaw and co-workers reported the stoichiometric reduction of CO to methanol using derivatives of bis(pentamethylcyclopentadieny1)zirconium (78). Bercaw had shown previously (79) that permethylated cyclopentadienyl ligands greatly enhance the stability of these complexes, permitting their isolation and/or in situ identification. In their study, these workers observed and partially characterized the novel hydride carbonyl species (13) which reacts with the dihydride complex Zr(C5Me5),H2 at room temperature to form the methoxy species (14).

L-/

(131

\

(14)

Hydrolysis of (14) readily yields MeOH (78). The carbonyl hydride complex (13) forms by the addition of CO to Zr(CSMes)2H2and possibly by the reaction of Zr(C5Me5)2(CO)2with H2, although only the former route allows its observation. In the absence of the dihydride complex Zr(C,Me5),H2, (13) reacts with itself on warming from - 80°C to 25°C yielding the unusual enediolate dimer (15) with the structure postulated as shown (78).


, AND

NO

101

(15)

While the mechanism for the formation of the methoxy complex (14) is not established, it is significant that the dihydride Zr(C,Me,),H, is needed for the reduction of the CO coordinated in (13). A reasonable proposal for this reaction can be formulated if it is assumed that since complex (13) is formally do, Zr-CO backbonding will not be of major importance, and that hydride complexes of the group 4 elements possess substantial hydridic character. The first assumption may lead to a more favorable equilibrium constant for carbonyl hydride +formyl interconversion as in (5), while the second suggests H- attack in this sequence; presumably on a coordinated formyl. If the latter results in Zr-H addition across C=O, then reductive elimination of a C-H bond leads to the observed product. This is shown in (21).

H,

J 43

(21)

CH3-0/'b

Alternatively, the formyl species may prove analogous to the acyl ligand in Zr(C,H,),(R)(RCO) in which the acyl group is 7c-coordinated to the metal serving as a 3-electron donor (SO). Attack by hydride may then lead to methoxy formation with the oxygen bonded to the original Zr atom of carbonyl complex (13). The dimerization reaction leading to (15) may provide an important clue into the mechanism of this stoichiometric reduction of CO to methanol by Zr species, and further studies are in progress. Recently, Schwartz and Shoer have described the use of (C,H,),ZrCI, as a catalyst in the reduction of CO by diisobutylaluminum hydride (DIBAH) (81). The products of this reaction, after aqueous acid workup, were methanol, ethanol, l-propanol, and l-butanol in decreasing amounts. A labeling

102


experiment using l 3CO verified that the alcohols produced were indeed formed by CO reduction and homologation. Isobutylene was also evolved indicating that the DIBAH reducing agent can serve as a source of more than one hydride unit in the CO reduction. The reaction sequence is believed to commence via the formation of (16), (22).

Dissociation of DIBAH from (16) then creates a vacant site for CO coordination. Hydride migration from Zr or external attack by hydride from DIBAH leads to the formation of a formyl that is then reduced by the aluminum hydride species (82). The steps for chain propagation and termination are outlined below. Since the exact nature of the Zr and Al species is unknown in these steps, they are represented by (Zr) and (Al), respectively.

+ (Zr)-CH,-R

(Al)-O--(Al)

homologation

chain propagation chain termination

R

Zr-H

I + (AI)-CH-O-(Al)

co (R

=

H, alkyl)

A total of three equivalents of DIBAH and two equivalents of CO are consumed in each run with eventual precipitation of (C,H,),ZrCI,. The sequence can be repeated with no loss of activity. This system may provide valuable insight into the notion of chain growth in CO reduction chemistry despite the fact that H, is not employed directly as the reductant. In another recent study of CO reduction using organometallic compounds of the early transition elements, Caulton et al. report that CH4 is evolved when toluene solutions of (C5HJ2Ti(CO), are heated at 150°C under H, or H, + CO atmospheres (83). When deuterium is used, the observed product is CD4, although some deuteration of the cyclopentadienyl rings is noted. Caulton does not report the percent yield of methane, but does state that the reaction isnot catalytic. One of the most novel aspects of this investigation is the isolation of a blue Ti, cluster that is inert to further reaction

,


103

with Hz + CO. The complex has the formula Ti60~(C5H5)6, and possesses structure (17) (83). FP

cp\Tfi?p I' ,Ti-Ti, CP

\,

?;i

,y CP

Cp (each = qface 5 -of C5H, the octahedron has a y3-oxygen)

CP (17)

Its formation may well arise from the reaction of water with (C,H,)?Ticontaining species, the water being produced in the methanation reaction. The study underscores the necessity of having a catalyst compatible with the products of the catalyzed reaction. In the case of CO reduction by H, , this implies at least modest aqueous stability. A reaction that may also give insight into the reduction of CO by hydridic species, although again Hz is not the ultimate source of the reductant, is found in the report by Treichel and Shubkin in 1967 (84).These investigators found that (C,H,)M(CO),(PPh,)+ (M=Mo,W) and NaBH, react to form the methyl species (C,H,)M(CO),(PPh,>(CH,) in reasonable yields. The sequence of steps in the reduction of the CO ligand is presumed to begin with hydride attack to give a formyl intermediate that, because of the saturated nature of the complex, is relatively stable to a-elimination to form a metal carbonyl hydride. Reduction of the formyl group then proceeds through a hydroxymethyl species and ultimately to the methyl product. Further support for this sequence comes from another report of CO reduction by BH4- by Graham and co-workers (85). These investigators find that borohydride reduces a carbonyl to a methyl in (C,H,)Re(CO),(NO)+ when the reaction is carried out in THF, but when done in benzene-water an unstable hydroxymethyl species, (C,H,)Re(CO)(NO)(CH,OH), is claimed (86). The reduction steps embodied in these reactions with BH4- are important in understanding metal-promoted hydrogenations of CO. A most intriguing and successful Fischer-Tropsch synthesis employing a homogeneous catalyst system has recently been reported by Muetterties and Demitras (87). The catalyst system is composed of certain group 8 carbonyl cluster compounds in molten NaCI.ZAlC1,. While most of the reactions in the study used Ir4(CO)12 as the carbonyl cluster, Rh,(CO),, , Rh6(CO)16, and Ru,(CO),, were also found to be active. The most salient features of this homogeneously catalyzed F-T synthesis may be summarized as follows: (a) of all metal complexes tried, including several mononuclear

104


ones, only certain cluster compounds were active; (b) the molten NaCl. 2ACl, medium is highly acidic and this appears to be essential for catalytic activity; and (c) the primary product from this reaction, which is carried out at 180°Cand 1.5 atm under a 3 : 1 H2:CO ratio, appears to be ethane (87). The first point may be intimately connected with the notion that CO activation is achieved in this system by maximal reduction of the carbonoxygen triple bond via interaction with more than one metal center, while the second is suggestive of the type of Lewis acid binding to carbonyl oxygens as seen by Shriver, Burlitch, and others (47-54). It is noteworthy that molten NaAlCl, , which lacks acidic character, is a relatively ineffective medium under the reaction conditions. That the primary reaction product in this study is ethane is highly significant. The major reaction products, C,H, and CH,, differ in their ratios depending on the reaction time with the highest values (10: 1 to 4: 1) occurring relatively early in the sequence. Thus methane is viewed in this study as a secondary product with aluminum chloride acting as an effective catalyst for carbonium ion based ethane fragmentation (87). The relatively select product distribution in this study underscores one of the potential advantages of homogeneous catalysis and provides impetus to further studies in this direction. The homogeneously catalyzed hydrogenation of CO, which may prove to have the most immediate and important commercial value, is the Rh complex catalyzed synthesis of ethylene gIycol, (23), developed by Pruett and Walker (82). 2CO

+ 3H2

Rh complex LS0-260"C+ > 3000 psi

HOCH$H,OH

This catalysis has been studied over an extensive pressure range from 3000 to 50,000 psi because of the large, unfavorable entropy term that renders AG for the reaction positive above 110°C. The selectivity of the reaction is quite good with a product distribution of HOCH2CH20H,CH,OH, and higher polyols in the ratio of 7 :2 : 1 (88). The active catalyst in this Union Carbide process appears to be a polynuclear rhodium carbonyl anion that is generated in situ from a variety of Rh complexes serving as catalyst precursors. Among species that have been isolated are [Rh,2(C0)34]2-,[Rh,3(CO),4H2]3-, and [Rh,3(CO)2,H,]Z-, the latter two being a conjugate acid-base pair. The catalyst is not poisoned by the addition of H2S, which in fact leads to the formation of a more stable cluster that has been isolated and structurally characterized as [Rh,,(CO),2S2]3-(88). Although the mechanism of ethylene glycol formation remains as yet undetermined, it may proceed by coupling of hydroxymeth-

BINDING AND ACTIVATION OF CO, C O , , AND NO

105

ylene species, and the cluster may be required for activation of carbon monoxide. In summary, this section has outlined what we believe to be one of the most interesting areas in carbon' monoxide chemistry. The hydrogenation of CO has great practical importance, and there exists much reason to pursue this reduction chemistry using homogeneous catalysts. To date, bona ja'e homogeneous catalyst systems have been found for ethylene glycol synthesis, and hydrocarbon formation mainly as ethane. The former is based on Rh cluster species while the latter uses different noble metal carbonyl clusters. Stoichiometric formation of CH,OH has been achieved using cyclopentadienyl zirconium derivatives and a catalytic reduction of CO by diisobutyl aluminum hydride using related zirconium species has appeared. A noncatalytic amount of methane was found in the reaction of (C5H,)Ti(CO), with hydrogen, and BH,- was observed to reduce a coordinated carbonyl to coordinated methyl in at least two cases. The years 1978 and 1979 have witnessed continuing activity on the catalytic reduction of CO and models for it. Both Casey and Gladysz have established that the neutral formyl complex (C5H5)Re(CO)(NO)(CHO) which they synthesized (59b,c) is the first intermediate in the borohydride reduction of coordinated CO to methyl as reported by Graham and co-workers (85). When the neutral formyl complex is reacted with BH, . THF, the species (C,H,)Re(CO)(NO)(CH,) results. A similar reduction does not occur when H, is used as the reductant, however (59b). While the previous report by Nesmeyanov et al. (86) of a hydroxymethyl species in the BH,- reduction process is now viewed as incorrect (59b,e), Casey has recently described (59e) unequivocal characterization of this species, and has shown how the formyl complex (C,H,)Re(CO)(NO)(CHO) can lead to its formation as shown in (23a). (C,H,)Re(CO)(NO)(CHO) +

(C,H,)Re(CO)(NO)(CH,OH)

+ (C,H,)Re(CO)(NO)(COOMe)

Both Gladysz and Casey report the synthesis of a new diformyl complex in their studies also (59b,c). Bercaw and co-workers have published a full paper (88a) on the Zr(C,Me,),H, reduction of CO to the methoxy compound (14), and suggest that the q2-formyl oxycarbene bonding mode (17a) analogous to Floriani's .n-acyl ligand (80) may be important in stabilizing the first intermediate in the reduction process. f-f

106


(174

The hydride donor ability of ZrH,(C,Me,), was demonstrated convincingly in a separate study (88b) when it reacted with (C,H,),W(CO) to yield the zirconoxycarbene compound of structure (17b).

A U7b)

In an attempt to develop similar hydride transfer reagents using less active hydrides, Labinger and co-workers (884 reacted a Nb hydride with Fe(CO), to give ( I ~ c ) ,but no subsequent reduction chemistry was observed. (C5Hs)2NbeFe(CO)4 (17e)

Marks and co-workers (88d,e)have found that permethylcyclopentadienyl complexes of the actinides have a chemistry closely paralleling that of the Zr compounds reported by Bercaw et al. (78, 79, 88a, 88b). Specifically, these investigators have found that M(C,Me,),(Me), (M = U, Th) react with CO to give butenediolate dimers similar to (15), but with two bridging ligands. The activation of CO is thought to arise from the facility of the migratory insertion reaction which then yields an q2 stabilized acyl or oxycarbene of structure type (17a), but with methyl replacing hydrogen (88e). A second postulated intermediate in the catalyzed reduction of CO by H, is coordinated formaldehyde which is a tautomer of hydroxycarbene. Early in 1979 Roper (88f) reported the formation and structural characterization of a stable formaldehyde complex, thus providing support for the proposed intermediacy of this species in homogeneously catalyzed CO reduction schemes. The CH,O complex has the specific structure (17d), and eliminates H, on heating to reform the starting Os(CO),L, complex.

(174

Further model reaction work has been reported recently by Masters et al. (889). Using AlH, in THF as both a stronger Lewis acid and more powerful hydride donor than BH, , these investigators found that reaction with


107

Ru,(CO),, yields up to 30% conversion of CO to methane, ethane, and propane, while reaction with M(CO), (M = Cr, Mo, W) leads more selectively to ethene formation. These observations, which relate to previously cited studies by Muetterties (87) and Schwartz (84, are interpreted to mean that dual binding of CO is essential for reaction and that a carbenoid species is a probable intermediate, notably M=CH,. In a potentially significant development in CO reduction chemistry, Rathke and Feder (88h) have found that HCo(CO), serves as a catalyst for conversion of CO/H, mixtures to MeOH and HCOOMe under forcing conditions (300‘atm, 200°C), albeit at very slow rates. An activation energy of 40.7 kcal/mol was obtained, and a scheme based on formyl radical chemistry was proposed. While the latter seems premature, the fact that reduction products were observed in the absence of cluster compounds brings into question the necessity of such a structural arrangement. A study by Henrici-Olive and Olive (884 using W(CO), and AlCl, in benzene under CO/H, reveals the formation of alkyl benzenes with chain lengths corresponding to a Schulz-Flory distribution similar to that observed in “classical” heterogeneous Fischer-Tropsch catalysis. This is the first example of such chain growth in a homogeneous system. There is a slight predominance of ethylbenzene which may indicate its formation by more than one path. Finally, Basset and co-workers ( 8 8 ~report ) that impregnation of alumina with metal carbonyl clusters leads to CO reduction through initial H, formation from CO + OH- (or adsorbed water) followed by catalyzed reaction on the surface. Above 25OoC, the selectivity of the reaction toward methane formation increases greatly, but so does decomposition of the surface bound clusters. In all cases about half of the carbon monoxide was converted to CO, as would be expected for the production of H, reducing equivalents. These recent developments indicate the current level of activity in the homogeneous catalysis of CO reduction, and it is expected that these efforts will continue at least at this level in the near-term future. 2. Oxidation of Carbon Monoxide Here we consider carbon monoxide as a chemical reducing agent, with the oxidation product being CO, in all cases. Carbon monoxide may act as a reductant in three different ways: (1) as a direct oxygen atom acceptor; (2) as a two-electron reductant with water as the source of oxygen; and ( 3 ) as an indirect reductant in which the reducing power of CO is used to make H, which then carries out the desired reduction. The three ways are shown schematically as (24)-(26) (in (24) [O] represents an oxide source). Equation

108

RICHARD EISENBERG A N D D A N E. HENDRIKSEN

(26a) is the water-gas shift reaction which is of fundamental importance in schemes for CO reduction.

+

--

CO [0] C 0 2 + [2e-] CO HzO C02 + 2H+ CO + HzO +COZ + Hz H2 2H+ 2 e -

+

+ 2e-

+

(24) (25)

(264 Wb)

It is currently used in ammonia synthesis as a means of removing CO (89), and its employment as a hydrogen source is expected to increase in the future (67).We shall discuss catalysis of the water gas shift reaction more extensively later in this section. The first class of reactions, direct transfer of an oxygen atom from the oxidant to carbon monoxide, has not been commonly observed as a reaction catalyzed by metal complexes in solution. One example, derived from a preparative procedure developed to substitute carbonyls with other ligands (90),is the reaction of trimethylamine oxide, Me3N0, with carbonyl clusters such as Os3(CO),,in the presence of excess CO. The net reaction is shown as (27). Me,NO

+ CO

-

Me3N + COz

(27)

The mechanism of this oxygen transfer involves attack of the amine oxide on a coordinated carbonyl to yield the indicated products, with the resulting vacant coordination site again being filled by a carbonyl. A second example is reaction (27a) which has recently been examined in detail by Feltham and Kriege (90a). The reaction is presumed to occur via an intramolecular oxide ion transfer through structure (17e).

Roundhill has recently reported the catalytic oxidation of CO to CO, by O2 using noble metal complexes, but no mechanistic details are given (92). Another direct transfer of oxygen to CO may take place in certain homogeneous catalytic systems for the reduction of nitric oxide by carbon monoxide according to (28) when no water is present to serve as the oxygen transfer agent (92,93). The catalyzed reductions of NO by CO are examined in the nitric oxide section of the review.

BINDING AND ACTIVATION OF CO,

2N0

CO,,

AND NO

+ CO --t N2O + COZ

109 (28)

The second class of reactions, with CO and H 2 0 serving as a source of two electrons, is not a catalytic reaction as such but is examined closely here since this is an integral step in the homogeneous catalysis of the water gas shift. The reaction may take two forms, with the reduction product being either a reduced metal center, (29), or a metal hydride, (30). M"+

M"+

+ CO + H 2 0 +Mn-2 + CO, + 2H+ + CO + H 2 0 M-H"-' + C 0 2 + H+ --f

(29) (30)

The reduction of metal ions in higher oxidation states by CO and H,O has been known for many years. Work on the reduction of Hg2+, Ag+, Ni2+,Cu2+,and Pd2+ has been summarized recently (4). The reduction of these metal ions does not proceed via a stable intermediate carbonyl. Since a metal carbonyl must be an intermediate in this reaction, however, the coordinated carbonyl must be very susceptible to attack by water, reacting as soon as it is formed. The ability of a metal in a higher oxidation state to activate a coordinated carbonyl to attack by as weak a nucleophile as water was noted previously in the description of the work by James et al., on the reduction of rhodium(II1) by carbon monoxide and water (62).Here a stable rhodium(II1) carbonyl, Rh(CO)Cl:-, can be observed as the initial product of reaction of RhC1,. 3H20 with CO. The Rh(II1) is then efficiently reduced to the rhodium(1) anion [RhCl,(CO),]-, even in nonaqueous solvents such as dimethylacetamide, where the only water available for reaction is the water of hydration of the starting rhodium chloride. The key intermediate in the reduction of metal ions by carbon monoxide and water is the hydroxycarbonyl (18). Initially (18) was proposed to form by a migratory insertion of CO into a M-OH bond, but more recent studies have favored a direct attack of water or hydroxide on a coordinated carbonyl (4,62).This latter view is in accord with the expected reactivity of coordinated CO toward nucleophiles. Intermediate (18) may then decarboxylate to give CO, and either a reduced metal ion or a metal hydride, as in (29) and (30), respectively.

(18)

CO activation via coordination to metal ions in higher oxidation states is especially well illustrated by the work of Halpern and co-workers (94)on the catalyzed oxidation of CO to CO, by ferricyanide according to (31).

110


In this study a cobalt(1) carbonyl complex serves as the catalyst. The complex [Co(CO)(CN),(PEt,),]- is oxidized stepwise by two equivalents of Fe(CN);leading to formation of the Co(II1) carbonyl species (19) which is readily attacked by water or hydroxide in the pH range 6-12.5. This attack produces the hydroxycarbonyl intermediate (20) which is identified as the main species in solution.

\OH (20)

(19)

Reductive decarboxylation of (20) yields CO, , H + , and a Co(1) species at a measurable rate (94). In the presence of CO, the starting cobalt complex is regenerated, and a catalytic system for the oxidation of CO by ferricyanide is established. It is significant that in this system the metal-carbonyl bond is formed when the cobalt is in a reduced state. It is the subsequent oxidation of the cobalt by electron transfer that activated the carbonyl to attack by water or hydroxide. That this activation results in a weaker metal-carbonyl bond is evident since the Co(II1)-carbonyl may be hydrolyzed in acidic solution with loss of the carbon monoxide ligand (94). The formation of a metal hydride by reduction with CO and H,O is closely related to the preceding examples of reduction at the metal center. This relationship is clear when a metal hydride is viewed as the conjugate acid of the reduced metal center, (33). M'N'

+ H+ + [M(N+Z)-H]+

N and N

+ 2 are formal

metal oxidation states

(33)

There is a mechanistic relationship too, as both reduction of a metal center and formation of a metal hydride may proceed through a hydroxycarbonyl intermediate (18) as in (34).

+ H+ Metal hydride formation from the hydroxycarbonyl can be a concerted process, involving /?-elimination of metal and hydrogen, leading directly to the products. A requirement of this pathway however, is that the metal hydroxycarbonyl species, (18), be coordinatively unsaturated. A complex lacking this prerequisite cannot form a metal hydride directly.


, AND NO

111

A well-studied reaction yielding a metal hydride is that of trans-[PtCl(CO)(PEt3),]’ with water to yield CO, and trans-[PtC1(H)(PEt3),], (35) (95). truns-[PtCl(CO)(PEt,),]+

+ HzO+ truns-[PtC1(H)(PEt3),1 + C 0 2 + H+

(35)

The kinetics of this reaction have been studied in detail and a hydroxycarbonyl is specifically proposed as an intermediate consistent with the kinetic data. Decomposition of this intermediate hydroxycarbonyl may proceed by p-elimination of the platinum hydride product since the hydroxycarbonyl is a 16-electron coordinatively unsaturated complex. Another well-known example of metal hydride formation from CO and H 2 0 is the reaction of iron carbonyl in aqueous alkali (55) (36). Fe(CO),

+ 20H-

+ HFe(CO),-

+ HC03-

(36)

If this iron carbonyl hydride is formed by a p-elimination from the intermediate (CO),Fe(COOH)-, a carbonyl must be lost in a prior step since this hydroxycarbonyl is coordinatively saturated. An alternative mechanism would be reductive decarboxylation followed by protonation of the very basic Fe(C0): - intermediate. From the preceding discussion and examples it is evident that a metal hydroxycarbonyl intermediate is central to all reductions by carbon monoxide and water. Consequently it is surprising that only two reports of stable, isolated hydroxycarbonyls have appeared, in contrast to the numerous isolated complexes of analogous alkoxycarbonyls, M-C’

/o ‘OR

and carbamoyls,

An iridium(II1) hydroxycarbonyl, [IrC1,(C0)(PMe2Ph),(COOH)], has been synthesized from the reaction of [IrC1,(CO),(PMe2Ph),]+ with wet ether (96). The complex is well characterized, is rather thermally stable, and is coordinatively saturated, which no doubt contributes significantly to its stability. This neutral Ir(II1) species is expected to be rather inert, and the lack of labile ligands prevents 8-elimination to produce metal hydride and CO, . A pair of coordinatively unsaturated platinum hydroxycarbonyls has been reported, but little experimental evidenceis given (97). These complexes, which have the formula PtR(COOH)(dppe) (R = methyl, cyclohexenyl; dppe = Ph2PCH,CH2PPh,), are formed by the reaction of carbon monoxide with the corresponding hydroxy compounds, PtR(OH)(dppe), an apparent example of CO insertion into the metal hydroxide bond. The hydroxy-

112


carbonyl complexes are reported to be stable in the solid state, in chloroform, and in dichloromethane. Their reluctance to undergo /3-elimination is striking, especially in contrast to the solution chemistry of trans-PtCl(CO0H)(PEt3)2 (95). Two additional metal hydroxycarbonyl complexes have been isolated and characterized quite recently. As part of the development of the extensive chemistry of the (C,H,)Re(CO),(NO)+ cation, Casey et al. have obtained the compound (C,H,)Re(CO)(NO)(CO,H) by the addition of one equivalent of NaOH to the rhenium cation complex in ether-water ( 9 7 4 . Treatment of the hydroxycarbonyl with CF,CO,H in acetone regenerates (C,H,)Re(CO),(NO)+, while treatment with Et,N in acetone converts the hydroxycarbonyl to the hydride, (C,H,)Re(CO)(NO)H. The second recently isolated metal hydroxycarbonyl is (C,H,)Fe(CO)(PPh,)(CO,H), obtained by treatment of the chloride salt of (C,H,)Fe(CO),(PPh,)+ with one equivalent of KOH in a benzene-water mixture (97b). This iron hydroxycarbonyl reacts with excess KOH to yield the potassium salt of the acid, and also reacts with excess HC1 to re-form the chloride salt of the starting dicarbonyl cation. This amphoteric behavior may be attributed to the equilibria in (36a).

+ OH- +MCOOH +MCOO- + Hf

(36a) These proposed equilibria are supported by changes in the IR spectra of the hydroxycarbonyl in solvents of different polarity. In benzene the spectrum corresponds to the covalent species MCOOH, while in the highly polar solvent formamide, the spectrum corresponds to MCO+, with hydroxide as the counterion. The decomposition of this iron hydroxycarbonyl proceeds rapidly upon warming in benzene. In contrast, the potassium salt does not readily decarboxylate in formamide. This indicates that in this system the decarboxylation involves transfer of the hydrogen to the metal. The third general reaction type, with carbon monoxide acting as a reducing agent through the intermediate production of hydrogen, proceeds via the water-gas shift (26a). While currently employed for CO removal as well as H, production, the water-gas shift will play an increasingly important role as a source of H, in the future as coal and other hydrogen-deficient carbon sources are used as starting materials for synthetic hydrocarbon fuels and chemicals (67). The heterogeneous catalysts presently in commercial use fall into two general classes (89). Those based on Fe,O, operate at ca. 350°C and are termed high temperature shift catalysts. Those based on finely divided copper metal with zinc oxide are active in the lower range 200250°C and are hence termed low-temperature shift catalysts. The reaction reaches a thermodynamic equilibrium that is not greatly affected by pressure, but is quite sensitive to temperature, favoring H, production at lower MCO+

BINDING AND ACTIVATION OF CO, C02, AND NO

113

temperatures. Hence the copper-based low-temperature shift catalysts can achieve a greater conversion, but suffer the drawback of susceptibility to sulfur poisoning, a problem the high-temperature shift iron catalysts do not share. There are thus good reasons for exploring metal complexes as water-gas shift catalysts in the expectation of finding new catalysts that are resistant to sulfur poisoning, and that have higher activity at lower temperatures than are presently found with the metal and metal oxide heterogeneous catalysts. Two homogeneous metal complex water-gas shift catalyst systems have recently appeared (98, 99). The more active of these comes from our Rochester laboratory (99,99a).It is composed of rhodium carbonyl iodide under CO in an acetic acid solution of hydriodic acid and water. The catalyst system is active at less than 95°C and less than 1 atm CO pressure. Catalysis of the water-gas shift reaction has been unequivocally established by monitoring the CO reactant and the H, and CO, products by gas chromatography The amount of CO consumed matches closely with the amounts of H2 and CO, product evolved throughout the reaction (99). Mass spectrometry confirms the identity of the CO, and H, products. The reaction conditions have not yet been optimized, but efficiencies of 9 cycles/day have been recorded at 90°C under 0.5 atm of CO. Appropriate control experiments have been carried out, and have established the necessity of both strong acid and iodide. In addition, a reaction carried out with labeled 13C0 yielded the same amount of label in the CO, product, ruling out any possible contribution of acetic acid decomposition to CO, production (99). Recent studies (ZOO) at the University of Rochester laboratory have been directed toward the interrelated goals of optimizing the reaction rate and determining the details of the reaction mechanism, including a determination of the important reactive metal complexes in this system. There appear to be three major rhodium carbonyl iodide species. Univalent rhodium is present as [RhI,(CO)J , while the two rhodium(II1) complexes have been identified as the monocarbonyl, [Rh15(CO)]2-, and a dicarbonyl, trans[RhI,(CO),]- . These rhodium(II1) carbonyls have been characterized by the positions of their single carbonyl stretches in the infrared spectrum, as reported previously by Forster (ZOZ, 102). The two are in equilibrium under a CO atmosphere, and apparently under 1 atm of CO at room temperature the equilibrium (36b) is shifted to favor trans-[RhI,(CO),]-. Rh15(CO)*-

+ CO + RhI,(CO),- + I-

(36b) This knowledge that the dominant rhodium species in solution are mononuclear can be combined with the results of a brief kinetic study (100)that yields an order in the total rhodium concentration of very nearly one. The conclusion is that the active species in this catalyst system are mononuclear,

114


and metal cluster catalysis is not necessary in this system in order to activate the reactant CO. A brief investigationof the effect of CO pressure on the rate of the catalyzed reaction has been carried out. In the range 0.5-1.0 atm at 8 0 T the rate appears to be approximately first order in CO pressure. This dependence may be assumed to be associated with the CO, producing steps in the catalytic cycle, rather than with the H, producing reactions, and evidence supporting this assumption has been obtained by substituting 1,- for protons as the oxidizing agent. A similar dependence on the pressure of CO is found for this reaction, (37), which corresponds to the catalytic reduction of I, to HI by CO (103). I,

+ CO + HZO + 2H+ + 21- + CO2

(37)

Our present working hypothesis for the mechanism of this water-gas shift catalyst system remains essentially the same as it was initially presented (99). The catalytic cycle consists of two separate reactions, with the rhodium alternating between the Rh(1) and Rh(II1) oxidation states. Rh(II1)

+ CO + HZO 4 Rh(1) + C02 + 2H+ Rh(1) + 2H+ --.t Rh(II1) + Hz CO + HzO -+ Hz + COZ

The production of CO, involves activation of CO by coordination to a Rh(II1) center, followed by attack of water to form a transient hydroxycarbonyl species. This species decarboxylates to yield a Rh(1) intermediate that rapidly forms RhI,(CO),-. If the dominant Rh(II1) species under 0.5-1 .O atm of CO is trans-RhI,(CO),-, the approximate first-order dependence on CO pressure indicates equilibrium or transient formation of a tricarbonyl, possibly RhI,(CO), , in which the CO ligands are sufficiently activated to facilitate ready attack by water. The reaction is very sensitive to acid concentration and the ratio of [H+]/[H,O]; the rate is maximized at approximate values of 2 M and 0.1, respectively (100). In the context of CO activation, the strongly acidic medium may lead to the formation of a kinetically significant species in which the carbonyl ligand is protonated and thus further activated to nucleophilic attack by water. The resultant species must lose protons to form the hydroxycarbonyl intermediate (18) which in turn may have to lose H + to decarboxylate. In extremely strong acid, these deprotonation steps may be inhibited, thus offering a possible explanation for the reduced rate of catalysis when [H'] is greater than ca. 3 M and [H+]/[H,O] is greater than 0.2. The other half of the catalytic cycle is the reduction of protons to yield H,, with [RhI,(CO),]- postulated as the intermediate reducing agent. Of

BINDING AND ACTIVATION OF CO, C O ,

, AND

NO

115

the Rh(1) halo carbonyl anions, only the iodo complex reacts with its respective hydrohalogenic acid to form a Rh(II1) carbonyl halide and H, (101). Certainly protonation of the Rh(1) species is the initial reduction step (the proton being reduced to hydride) and strong acid should facilitate this process. The need for iodide in this system may be to activate the Rh(1) center to protonation via the formation of [Rh13(CO),]2- which as a dianion would react more readily with H+ to form a Rh(II1) hydride. (In the absence of excess iodide, catalysis does not occur and the solutions remain light yellow, characteristic of Rh(1) halo carbonyl anions.) The production of H, may then proceed from the Rh(II1) hydride via a hydride/iodide interchange, transferring H-to H+ and I- to Rh(II1). In recent work the experimental studies described above have been extended and a mechanistic scheme has been developed which is consistent with all the experimental evidence (103~).The most significant new information pertains to the temperature dependence of the catalyzed reaction rate. Two different temperature regimes have been found. In the range 80-100°C the apparent activation energy is 9.3 kcal/mol, there is a first-order kinetic dependence on carbon monoxide pressure, inverse dependence on acid and iodide concentrations, and rhodium(II1) species are the predominant metal species in solution. In contrast, below 65°C the apparent activation energy is 25.8 kcal/mol, there is no kinetic dependence on carbon monoxide pressure, a second-order dependence on iodide concentration, and a positive dependence on acid concentration. The rhodium(1) species RhI,(CO),- is the predominant metal species in solution. These data are interpreted in tsrms of a change in the rate-limiting step in the catalytic cycle in the two different temperature regimes. In the high-temperature limit when rhodium(II1) species are predominant, the limiting step is rhodium(II1) reduction or C 0 2 production. In the low-temperature limit when rhodium(1) species are present, the rate-limiting step is rhodium(1) oxidation or hydrogen production. These new observations have significantly extended our understanding of the mechanism of this catalytic reaction, and a reasonably detailed mechanistic scheme has been presented which is consistent with all the kinetic data and with the studies on the complexes present in the catlytically active solutions ( 1 0 3 ~ ) . A group at Monsanto has also studied the catalysis of the water-gas shift reaction by rhodium carbonyl iodide (10%). The main difference between their work and our own is the choice of reaction conditions. Their study was conducted at 185°C under _200-400 psig carbon monoxide. Despite this drastic difference in reaction conditions, the studies are surprisingly consistent. In particular, the Monsanto group also finds evidence for two ratelimiting reactions. They did not find this by temperature variation, but instead, consistent with our own work, find that at low acid and iodide

116


levels oxidation of Rh(1) or hydrogen formation is rate determining, while at higher concentrations of these species Rh(II1) reduction or CO, formation becomes rate limiting. Further agreement is found in the rates observed. If the rate of -1 turnover/hr obtained in our study at 90°C and P, of 400 Torr is extrapolated to 185°C and 400 psig of CO assuming an activation energy of 9.3 kcal/mol and a first-order CO dependence, one obtains a rate of -700 turnovers/hr. This value is of the same magnitude as the -400 turnovers/hr reported by the Monsanto group. The principle difference between the Monsanto work and our own occurs in the observations of CO dependence. They find that increasing P, inhibits the reaction in the regime where Rh(1) oxidation is rate limiting, but find that the reaction rate is independent of P, when Rh(II1) reduction is limiting; we find that the catalysis is independent of Pcoand first order in P, in those regimes, respectively. In view of the much higher temperature and pressures of the Monsanto study, one possible explanation of this difference is that different steps in the catalytic pathway could change in relative importance, thus altering the kinetic dependences on Pco . In summary, the phenomenon of catalysis of the water-gas shift reaction by this homogeneous metal-complex system has been firmly established, and the considerable progress that has been made toward elucidating the mechanism of the reaction has been outlined. The Monsanto work on this system is interesting because it was done under industrially significant conditions, thus verifying the utility and stability of this catalyst system at these higher temperatures and pressures. Whether this catalyst system can demonstrate significant advantages over the presently used heterogeneous catalysts remains to be determined. Ford and co-workers have also recently developed a homogeneous catalyst system for the water-gas shift reaction (98).Their system consists of ruthenium carbonyl, Ru,(CO),, , in an ethoxyethanol solvent containing KOH and H,O under a CO atmosphere. Experiments have been conducted from 100-120°C. The identity of the H, and CO, products has been confirmed, and catalysis by both metal complex and base has been verified since the total amount of H, and CO, produced exceeds the initial amounts of both ruthenium carbonyl and KOH. The authors point out that catalysis by base in this system depends on the instability of KHC03 in ethoxyethanol solution under the reaction conditions (98).Normally the hydroxide is consumed stoichiometrically to produce carbonate, and this represents a major reason why a water-gas shift catalyst system has not been developed previously under basic conditions, As has been noted above, coordinated carbonyl does not have to be greatly activated in order for it to undergo attack by the strongly nucleophilic hydroxide ion. Because of the instability of KHCO,

BINDING AND ACTIVATION OF

CO, C O , ,

AND NO

117

in ethoxyethanol the ruthenium carbonyl catalyst system bypasses the problem of stoichiometric base consumption, while coupling the oxidation of CO to the production of hydrogen. Many aspects of the mechanism remain unestablished at present although a generalized reaction sequence based on metal carbonyl anions has been put forth (104). The detection and isolation of H3Ru4(C0);, as a major species present in the catalyst solution lends support to this generalized scheme. Recent work by Ford et af. demonstrates that a variety of metal carbonyl clusters are active catalysts for the water-gas shift under the same reaction conditions used with the ruthenium cluster (1%). In particular, the mixed metal compound H,FeRu3(CO),, forms a catalyst system much more active than would be expected from the activities of the iron or ruthenium systems alone. The source of the synergetic behavior of the iron/ruthenium mixtures is under investigation. The ruthenium and ruthenium/iron systems are also active when piperidine is used as the base, and in solutions made acidic with H,S04 as well. Whether there are strong mechanistic similarities between the acidic and basic systems remains to be determined. The use of nitrogen-containing bases, and in particular ammonia, in homogeneous catalyst systems for the water-gas shift reaction is described in a patent by Fenton (105). The reaction media in his systems are moderately to strongly basic with pH’s in the range 7.5-12.0. Many noble metal compounds are given as catalyst precursors. The main thrust of the patent, however, is an observation of reaction catalysis through an analysis of products. In the runs given in the patent, a pressure of 56 atm of CO and a temperature of 200°C were employed. A recent development at Rochester (106) has been the discovery of a new homogeneous catalyst for the water-gas shift reaction that appears to be at least twice as active as the rhodium carbonyl iodide system described above. The new system is based on a platinum chloride-tin chloride complex in acetic acid-HCI solvent. Preliminary results suggest that the Sn(II)/ Sn(1V) couple plays an integral role in the catalysis. With all of the reactions proceeding through a platinum complex species, it appears that Sn(I1) is reducing protons to H, and that Sn(1V) is oxidizing CO to CO,. Two platinum-tin complexes, cis-[PtCl(CO)(SnCl,),]- (107) and PtCI,(CO)(SnCl,)] - , have been isolated from the solution under catalytic conditions. When a sample of the first complex was heated in the catalyst solution under nitrogen, hydrogen evolution was observed with little CO, being evolved. When the second complex was dissolved in the catalyst solution and heated under CO, CO, evolution corresponding to -0.6 mol of CO, per Pt complex was observed within 10 minutes with only a trace of H, observed. From these observations it appears that [PtCI(CO)(SnC13),]- is most likely responsible

118

RICHARD EISENBERG AND D A N E. HENDRIKSEN

in the catalysis for H2 formation, and that [PtC12(CO)(SnC1,)]- is most likely responsible for CO, formation. A proposed mechanism described as “coupled cycles” has been presented (106). The elements of yet another water-gas shift catalyst system, this one photochemically assisted, have been recently outlined, though the authors do not make any reference to this aspect of their work (63).The focus of the research is on the preparation of a vanadium carbonyl trihydride, [H3V(CO),(diars)], from the reaction of (Et,N)[V(CO),(diars)], with excess aqueous or anhydrous HX (X = C1, Br, I) in THF solution. The initial product is HV(CO),(diars) which then reacts with excess HCl to yield the vanadium trihydride and phosgene, or CO, in the presence of water. The postulated mechanism of this reaction is most interesting, especially in light of the preceeding proposals for CO activation through an induced charge separation. Since both strong acid and halide are demonstrated as necessary, the mechanism is proposed to involve protonation of the complex and halide attack, presumably on a coordinated carbonyl. The suggested mechanism (63) is given as (38). HV(CO),(diars)

+ HX ---L H,V(CO),(diars)+X-

H,V(CO),(diars)1C(OH)X,} X = C1, Br, I

-P

HX THV(CO),(diars){C(OH)X} T H,V(CO),(diars) + COXz

(38)

Protonation of a carbonyl oxygen rather than the metal may be encouraged in this case by the high coordination number of vanadium. This would then promote halide attack on the carbonyl carbon to yield an intermediate hydroxyhalocarbene, which reacts further to yield the indicated products. This system represents a potential photoassisted water-gas shift catalyst system since H,V(CO),(diars) upon photolysis with a mercury vapor lamp yields H2, and in the presence of CO regenerates the starting complex HV(CO),(diars). The feasibility of coupling these two reactions in the same reaction solution remains to be demonstrated. Recently three more laboratories have presented systems for the homogeneous catalysis of the water-gas shift reaction. Pettit et al. report a variety of metal carbonyl clusters that are active catalysts in aqueous trimethylA. D. King, Jr., and co-workers find that the monoamine-THF (107~). nuclear metal carbonyls of iron, chromium, molybdenum, and tungsten can act as catalysts for the water gas shift in aqueous organic solvents in the presence of base (107b). In one run with tungsten carbonyl, the reaction was catalytic with respect to the added KOH as well. Otsuka et al. find that the platinum(0) complex Pt[P(i-Pr),], will catalyze the water-gas shift reaction in aqueous acetone or THF in the absence of added acid or base ( 1 0 7 ~ ) . This system apparently activates the water molecule towards this reaction as well as the carbon monoxide. Oxidative addition of water to a species


, AND

NO

119

PtL, is inferred since a catalytic intermediate, trans-{PtH(CO)[P(i-Pr),l,)OH could be isolated as the BPh,- salt. A scheme is presented involving CO, production via the hydroxycarbonyl PtH(CO,W)L, , and hydrogen production via the known dihydride tran~-PtH,[P(i-Pr)~],. Some of the catalyst systems for the water-gas shift that have been described above have also been found to be useful in a variety of other reactions that can use CO and H,O as reactants in place of molecular hydrogen. The key catalytic intermediate in these reactions is a metal hydride formed from CO and H,O. Examples of such reactions include the hydroformylation of olefins using metal carbonyl cluster complexes (107a, 1 0 7 4 , the reduction of aldehydes both in conjunction with the hydroformylation ( 1 0 7 4 and separately using a rhodium-triethylamine catalyst system (107e), the reduction of aromatic nitro compounds to amines using both metal clusters and iron carbonyl (107f),and the catalytic exchange of deuterium for hydrogen at saturated carbon atoms in tertiary amines (1079). This section has dealt with the oxidation of CO to CO, , especially as it enters into the water-gas shift reaction (26a). A reasonable view of the homogeneous catalysis of this reaction, whether in basic or acidic media, is emerging in which CO formation proceeds from nucleophilic attack of water or OH- on an activated carbonyl followed by either reductive decarboxylation or hetero-atom /?-eliminationyielding, respectively, a reduced metal or a metal hydride species.

111.

Carbon Dioxide

The chemistry of carbon dioxide with transition metal complexes is a field of research that has only recently received wide attention. One goal of research in this area is the development of efficient catalytic processes in which carbon dioxide is reduced by molecular hydrogen and/or incorporated into an organic molecule. Some examples of such desirable transformations are (39)-(42). CO2

+ H2 + HCOOH

CO, + HZ + ROH ---t {HCOOR or RC02H} + H 2 0 C02 + H2 R2NH + HCONR2 HZO C02 + Hz + R,C=CH2 + RZCHCH2COOH

+

+

(39) (40) (41)

(42)

Each of these reactions requires the presence of a catalyst, and it is through the development of transition metal-CO, chemistry that complexes capable of catalyzing these processes will be discovered. Carbon dioxide does have an extensive chemistry with main group organometallic compounds (108). Reactions of CO, with such species as Grignard

120


reagents and alkyl lithium compounds are well known and documented. While reaction of CO, with such main-group organometallic species can readily lead to the products indicated above, the reactions are invariably stoichiometric with respect to the organometallic compounds, which in turn places definite economic constraints on their use in the large-scale production of oxygenated organic chemicals. As a source of carbon, CO, is both abundant and readily available. However, at the present time a fundamental obstacle exists to the use of CO, as a carbon source in the production of more reduced organic compounds, even if the necessary chemistry is developed and proven to be efficient on an individual reaction basis. In all of the preceding examples of CO, reduction and incorporation, hydrogen is a necessary reactant. For the more generalized reductions of CO, to hydrocarbons and carbohydrates, (43) and (44),respectively, even greater mole ratios of H, : CO, are required. m COz

+ (3m + l)Hz + C,H,,+, + 2m HzO n COZ + 2nHz --+ C.HZ.0. + n H 2 0

(43) (44)

While some of the hydrogen consumed in (43)and (44)is incorporated into the organic products, much of it is also consumed in the production of water. Thus, water in these reactions is serving as a “sink” for the excess oxygen in CO, over that in the desired product and the formation of water is providing a thermodynamic driving force for the reaction. When the focus is not on CO, as a reactant, it too is an often used “oxygen sink,” a thermodynamically very stable repository for oxygen. An example of the utility of water and CO, as such stable oxygen sinks is in the generalized reaction for carbon monoxide reduction with a 1: 1 H, :CO ratio (45). 3CO + 3Hz + 2fCHZ+ + COz + HzO (45) [This equation actually represents a combination of (11) and (13) given above.] Hence if CO, is to be used as a raw material for the production of compounds containing C-H bonds, then H,O must serve as the ultimate destination of the unwanted oxygen. This in turn requires hydrogen. Consequently, the ready utilization of CO, as a source of carbon requires simultaneously a cheap and abundant source of hydrogen. Such a source is not presently available. The present route to the large-scale production of hydrogen (89)involves steam reforming of natural gas (methane) at high temperature, (46), CH4

+ CO

(46)

+ HzO --.t Hz + COZ

(47)

+ HZO

--.f

3Hz

and the water-gas shift reaction (47). CO


121

Since hydrogen production by currently economic methods requires concurrent generation of CO, as the oxygen sink, the dilemma in the present utilization of CO, as a starting material for more reduced organic compounds such as petrochemicals seems complete. Stated simply, an oxygen sink is required in the formation of C-H bonds from CO, , but at present that sink is ultimately CO, itself. The dilemma is not hopeless, however, if one realizes that an alternative means of storing “unwanted” oxygen in these reduction reactions is as metastable molecular oxygen. But such a procedure will require an input of energy. This, of course, is precisely what happens in photosynthesis whereby carbohydrates and 0, are produced from CO, and H,O using sunlight, (48) (109- 1I I ) . nCO, + n H,O

-%c.H,.o.

+ no,

(48)

A discussion of photosynthesis is beyond the scope of this article, but it should be remembered that (48) represents the most fundamental and important reaction in the fixation of CO, . The purpose of this section is to review the known chemistry of CO, with transition metals, including (a) the coordination chemistry of CO,; (b) insertion reactions of CO, into metal-hydrogen, metal-carbon, metalnitrogen, and metal-oxygen bonds; and (c) the catalytic reduction of CO, and its incorporation into organic products. There have been recent reviews on the reactions of CO, by Volpin and Kolomnikov (108, 112) and Kolomnikov and Grigoryan (112a), as well as other less accessible reviews on CO, chemistry including that with transition metals (113). While there is some overlap with the 1975 Volpin and Kolomnikov review and the 1978 Kolomnikov and Grigoryan review, we include a good deal of later material, and our purpose and perspective differ from their reviews. A. CARBON DIOXIDE AS A LIGAND IN TRANSITION METALCOMPLEXES

The notion of CO, activation by transition metal complexes is intimately related to CO, coordination chemistry, since it is through coordination that the electronic structure of the CO, molecule can be sufficiently perturbed to result in altered reactivity. Different modes of CO, coordination will manifest themselves in different types and degrees of altered reactivity. Whether these changes in the CO, electronic structure and consequent reactivity correspond to true activation must be examined on an individual reaction basis. In addition, there exists the larger question of whether coordination of CO, to a metal center is really a necessary step in promoting reaction of CO, with different substrates. Two obvious alternatives are: (1) occurrence of direct substrate-CO, reaction in solution with the product

122


then being stabilized by coordination to the metal center, and (2) direct CO, attack upon a coordinated substrate, in which case it is essentially the substrate that is being activated rather than the CO, . Exploring the chemistry of C0,-coordination complexes should provide some answers as to whether these complexes efficiently activate the carbon dioxide molecule toward reaction with a variety of substrates. Carbon dioxide coordination chemistry is a very recent development in organometallic chemistry. The first claims for coordination complexes of CO, did not appear until 1969, and only recently have any number of wellcharacterized complexes been reported. Consequently, no sweeping generalizations can be made regarding the nature of CO, coordination complexes, but a few tentative guidelines can be inferred. The carbon dioxide molecule exhibits several functionalities through which it may interact with transition metal complexes and/or substrates. The dominant characteristic of CO, is the Lewis acidity of the central carbon atom, and the principle mode of reaction of CO, in its main group chemistry is as an electrophile at the carbon center. Consequently, metal complex formation may be anticipated with basic, electron-rich, low-valent metal centers. An analogous interaction is found in the reaction of the Lewis acid BF, with the low-valent metal complex IrCl(CO)(PPh,), (114). These species form a 1 : 1 adduct in solution; evidence for an Ir-BF, donor-acceptor bond includes a change in the carbonyl stretching frequency from 1968 to 2067 cm-'. In addition to interacting with the Lewis acid center of the CO, molecule, these same low-valent metal complexes may also interact with the carbonoxygen n-bonds in CO, , in much the same way as olefins interact with electron-rich complexes. Finally, the oxygen atoms in CO, may be expected to show weak electron donating ability, possibly coordinating to a very electron-poor metal, although this mode of coordination of CO, is not presently known. The known carbon dioxide coordination complexes seem to fall into three categories. In the first, carbon dioxide functions as a discrete ligand; the only interaction binding the CO, is with the metal itself. There is evidence for two such forms of unsupported coordination: through the carbonoxygen double bond, (21), and through the carbon atom alone, (22).

(21)

(22)

The second category might be called supported CO, coordination, where an additional labile interaction stabilizes the coordinated CO, beyond that


123

achieved by the metal-CO, interaction alone. As in unsupported coordination, the CO, molecule is still readily displaced from the metal center. A general form of such support is an interaction of one CO, oxygen atom with an acidic center in the coordination sphere, (23)(A = Lewis acid center).

The third form of CO, coordination involves a generally irreversible reaction with other ligands or external reactants, leading to “CO, coordination” as carbonate, acetate, etc. Much of this coordination mode will be dealt with in the sections on CO, insertion reactions. In this section we will not attempt to present a comprehensive literature survey. In considering which reports and claims of CO, complexes to present, we have limited the discussion to those that seem to offer the most insight into both CO, complex formation and activation, Only two structural reports of CO, coordination complexes have appeared (225,116).The first illustrates the unsupported mode of coordination through n-bonding, and the second, a bis-CO, adduct, may be viewed as an example of supported coordination. The first structural report was of (carbon dioxide)bis(tricyclohexylphosphine)nickel, [Ni(CO,)(PCy,),] .0.75 C6H, Me ( 2 2 5 ) . This was prepared by reaction of the zerovalent nickel complexes Ni(PCy,), or [{Ni(PCy,),},(N,)] with CO, in toluene. The x-ray crystal structure shows that the coordinated CO, is strongly bent and q2 coordinated through carbon and one oxygen, with the P,Ni(CO,) arrangement essentially planar, (24).

While the uncoordinated carbon-oxygen bond (1.17(2)A) is unchanged from that in free CO, (1.16 A), the coordinated C-0 bond (1.22(2)A) appears somewhat longer, though still significantly shorter than the carbonoxygen single bond length of about 1.45 A. In a subsequent paper (127) the authors conclude that the bonding in [Ni(CO,)(PCy,),] is very similar to that in [R(CS,)(PPh,),] as described by Mason and Rae (228).The bondingis dominated by the synergic interactions of n-electron donation to the metal and backbonding to the CO, n* orbitals. The CO, ligand in this complex is not bound strongly, being released quantitatively when a stream of argon is bubbled through a solution of the complex in toluene at room temperature. The CO, complex can be re-formed by bubbling CO, through the solution. This work has recently been extended to complexes with other phosphine

124


ligands (117). The complexes [Ni(CO,)L,], L = PEt, and PBu,", were obtained by reacting toluene solutions of NiL4 with 1 atm of CO, and cooling to - 30 and - 70", respectively, to obtain crystals. By a comparison of their ir spectra, the Ni-CO, interaction appears stronger in these complexes, and this is attributed to the smaller steric bulk of these phosphine ligands compared to PCy,. Infrared spectral evidence is also presented to support formation of [Ni(CO,)L,] complexes in solution. These are tentatively described as distorted planar, with CO, bound only through carbon. The other carbon dioxide complex characterized by x-ray crystallography contains two linked CO, molecules in the coordination sphere (116). This complex, [IrC1(C20,)(PMe3),], was prepared by the interaction of CO, with chloro(cyclooctene)[tris(t~methylphosphine)]iridium(I), [IrCl(C,H,J(PMe,),], in benzene solution. The structure, (25), shows essentially octahedral coordination about the iridium center, with one metal-carbon bond and a five-membered chelate ring formed with the second CO, molecule.

'P L1

Presumably the complex forms by electrophiiic attack of the CO, carbon on the electron-rich metal center, followed by a similar electrophilic attack of the second CO, on the more basic oxygen of the coordinated CO, , forming an oxygen-carbon bond. The metallocycle ring closing then completes the complex formation. Support for this mechanism comes from infrared spectra implicating a mono-CO, adduct that is observed when the starting metal complex reacts with less than two equivalents of CO, . Subsequent work from the same laboratory indicates that coordination of only one carbon dioxide molecule can also be achieved in similarly electron-rich and sterically unhindered systems (119). The complex [Ir(dmpe),Cl] (dmpe = Me,PCH,CH,PMe,), when suspended in benzene or as a solid powder, takes up one molecule of carbon dioxide, which is easily displaced by other ligands or upon heating. Similarly, Ir(diars),Cl (diars = o-C6H4(As(CH3),),) takes up CO, to form white 1r(diars),C1.COz. If the chelating ligands remain co-planar as in the starting complexes, they presumably block the coordination sites c k to the bound CO,, preventing addition of the second CO, as in [IrCl(C,O,)(PMe,),]. A bonding configuration


CO, C O , ,

AND NO

125

as in (26) is suggested, as opposed to the side-on bonding described above in "i(C02)(PCY3)21* Two additional reports of structurally characterized carbon dioxide coordination complexes have recently appeared. The first is a striking example of supported CO, coordination (119~).In [Co(pr-salen)KCO,THF], , the carbon dioxide molecule is bent and is coordinated to the electron-rich cobalt(1) at carbon, to one acidic K + through one oxygen, and to two potassiums at the other oxygen. This bifunctional coordination is possible since the potassium ion is held close to the cobalt in the precursor [Co(pr-salen)K] by coordination to the oxygens in the pr-salen ligand. This structural report follows earlier work in which CO, was found to bind reversibly in a 1 :1 stoichiometry with Co(salen)Na, salen = N,N'-ethylenebis(salicy1ideneiminato) (119b). The second new structural report of CO, coordination is of [HOs,(CO),,,&2-c~,)os&o)1,]-, prepared by reaction of Os,(CO),, with [Os3(CO),,H]- (119~).The bent briding CO, in this complex is bound through carbon to the Os, cluster and through both oxygens to the Os, cluster. A later study shows that oxygen is necessary for this reaction, and also reports three more related p,-CO, complexes (1194. Flynn and Vaska have reported CO, adducts of trans-[M(OH)(CO)(PPh,),] (M = Rh, Ir) (120) that apparently are analogous to the bis-CO, adduct described above in that the coordinated CO, is stabilized by an interaction between one of the oxygens and an acidic center, the carbon atom of a second carbon dioxide molecule in the previous case and the proton of a coordinated hydroxide in the present case. The adducts were formed by the reaction of the dry powdered complexes with 1 atm of carbon dioxide. The rhodium complex formed the adduct much more quickly and completely, and on continuous pumping decarboxylated much more slowly than the analogous iridium compound. Significantly, the voH stretch in the starting complex disappeared on complex formation with the CO, . Also, coordination of CO, to the analogous complexes containing F, C1, or ClO, in place of OH could not be achieved. The change in vm on complex formation is surprisingly small, only 16 and 27 cm-', respectively, for the rhodium and iridium complexes, indicating that the complexes are best considered as containing univalent metal centers and that the metal-CO, interaction

126


is rather weak. A possible structure for the adduct in which the addendum is stabilized by a C0,-Lewis interaction is (27), with hydrogen bonding a crucial stabilizing factor.

(27)

Alternatively, a true CO, adduct may not be formed, and instead an isomer of the M(1) bicarbonate species M(CO)(HCO,)(PPh,), may be generated (see below). A recent report (121) on the reactions of tetrakis(trimethy1phosphine) iron, Fe(PMe,),, with carbon dioxide reveals a rich and varied chemistry, illustrating many of the reaction modes of CO, with low-valent transition metal complexes. Two primary reactions of CO, with Fe(PMe,)4 are noted, as a consequence of the two isomers in equilibrium (49). Fe(PMe,),

=+

(49)

L,(H)Fe(PMe,CH,)

Reaction of the metallated species proceeds in THF at room temperature to yield a C0,-inserted product and a species that has taken up a second CO, by insertion into the metal-hydride bond, (50).

/pa L,(H)Fe h e ,

+ coa

-

H \

::

L,(H)Fe/ 0, o , \ $€IP Me/P\ Me

p

+Co,

-L

~

0

o \ j e / o ~ o(50) L’

I

\

MeAP\Me L

The iron(0) tetrakis(ph0sphine) on the other hand reacts in pentane at 0°C to form a CO, adduct, L,Fe(CO,), which has been isolated. This may react further with CO, to yield a disproportionation product, (51).

All of the products of these reactions were characterized and identified by their infrared, proton NMR, and proton decoupled phosphorus NMR spectra. Some of the difficulties and ambiguities presently plaguing carbon dioxide coordination chemistry are illustrated in the following summaries of re-


CO, C O , ,

127

AND NO

search reports. In particular, they point out the problems of assigning a CO, coordination mode without a structural study. Chatt et al. (122) have found that [Mo(N,),(PMe,Ph),] reacts rapidly with carbon dioxide in toluene at room temperature to give a product analyzing as [Mo(CO,),(PMe,Ph),]. A structure analysis was attempted but was not successful due to poor quality crystals and decomposition of the adduct in the x-ray beam. On standing in THF solution this adduct spontaneously forms a new compound, which was characterized by x-ray crystallography. This proved to be a dimeric complex with two bridging carbonates, [(PMe,Ph)3(CO)Mo(C03)2Mo(CO)(PMe2Ph)3], each carbonate bidentate at one molybdenum atom and monodentate at the other. The authors speculate on a reduction and disproportionation of coordinated C 0 2 to form the carbonyl and carbonate ligands: 2 C 0 2 2e- -+CO;CO. In contrast, neither Mo(N,),(dppe), nor W(N,),(PMe,Ph), reacts with CO, . The reactions of dialkyl palladium complexes with carbon dioxide are reported to yield two types of complexes (123). When trans-[Pd(CH,),(PEt,),] or cis-[Pd(CH3),(PMePh2),] are reacted with CO, in hexane or toluene solution, respectively, for ca. 15 hours, a white precipitate accumulates accompanied by the evolution of methane. The authors formulate these complexes as dimers of [Pd(CH3)(C0,)L2] and also consider a structure involving phosphine ligand metallation. However, these species have recently been reformulated as monodentate bicarbonate Pd(I1) complexes, trans-[PdMe(HOCO,)L,], and the structure of trans-[PdMe(HOCO,)(PEt,),] has been determined by x-ray crystallography (1234. Furthermore, water has been demonstrated to be necessary for the reaction, and in the presence of alcohols analogous monoalkyl carbonato complexes, trans[PdMe(ROCO,)(PEt,),], are obtained. In contrast, when the diethyl complexes trans-[Pd(C2H,),Lz] react with C 0 2at - 20 to - 40"for 15 hours (1,= PEt, , hexane solution; L = PMePh, , toluene solution), ethane is evolved and the products are formulated as zero-valent [Pd(C,H,)(CO,)L,] (123). These complexes are much less stable, decomposing at room temperature and 50°, respectively, with approximately stoichiometric evolution of CO, and ethylene. The bonding mode of CO, in these complexes does not appear to be firmly established, but CO, may well function as a discrete ligand, as in the known Ni(0) complexes

+

+

(115,117).

Results from the same laboratory have appeared on carbon dioxide complexes with copper(1) compounds (124). Phosphine-containing copper(1) carboxylate complexes, formed by the insertion of CO, into a copper-alkyl bond, take up additional CO, . A complex, formulated 'as [(RCOO)Cu(CO,)(PPh3),1 has been isolated and the CO, shown to be labile, i.e., the CO, was lost on attempted recrystallization. The authors speculate that the car-

128


boxylate groups may be monodentate in these complexes, but do not consider the possibility of an additional stabilizing interaction between the monodentate carboxylate and the bound CO, as in (28). .

\

0-d

’R

(28)

A number of CO, adducts with low-valent rhodium complexes have been C6H6 (125), reported, the most recent being [(PPh3)3Rh2(C02),(CO)z]. formed by the reaction of carbon dioxide with a benzene solution of triphenylphosphine and [{(PPh,)Rh(CO),}, .C,H,]. The complex has been characterized by analysis and its infrared spectrum. A previously reported compound formulated with only one molecule of CO,, [(PPh3)3Rh,(C0,)(CO),], exhibits a different infrared spectrum (126), and its formulation as an adduct of Rh(0) now seems unlikely, since an apparently identical complex is obtained on treatment of [Rh*(CO)(PPh,),(OCO,H)] with air (127). A third complex, formulated as [(PPh3)5Rh,(C0,)Cl,], is obtained by reaction of CO, with [Rh(PPh,),Cl] (128). This class of adducts appears to be still too ill-characterized to speculate profitably on the bonding modes of the coordinated carbon dioxide. It should be pointed out that a fourth purported CO, adduct of rhodium, [Rh,H,(CO,)(PPh,),], (129), has recently been characterized structurally by x-ray crystallography and reformulated as [Rh,(C03)(PPh3), C6H,] (130), with the bridging carbonate bidentate at one rhodium and monodentate at the other. It is evident from the preceding discussion that the field of CO, coordination chemistry is still a new area, but enough work has been reported to allow the several different modes of CO, coordination to be identified. These include (1) binding to the metal center as a Lewis acid through carbon alone, (2) mcomplexation involving one of the C--V units, and (3) supported coordination in which an additional weak interaction stabilizes the bound CO,. In all cases except for the “irreversible” reactions of CO, leading to bound carbonate, carboxylate, etc., the metal-CO, interaction is relatively weak and the CO, is easily liberated.

B. INSERTION REACTIONS OF co, Carbon dioxide insertion reactions are potential intermediate steps in catalytic cycles leading to reduction of CO, or its incorporation into organic molecules. Analogies with carbon monoxide chemistry may be drawn, e.g., insertion reactions of carbon monoxide (20) play a key role in both the

129


catalytic carbonylation of methanol to form acetic acid (12, 13) and in the hydroformylation reaction (11). Relatively little mechanistic work has been reported on the insertion reactions of CO, . The mechanism seems to be established only for the insertion of CO, into the dialkylamides of the early transition metals (131). We will speculate on probable mechanisms for the various types of insertion reactions that follow. Future work will undoubtedly shed more light on .these processes, leading to a better understanding of the reaction, and enabling a more rational design of catalyst complexes in order to incorporate the insertion process into an efficient catalytic cycle. Since the insertion of carbon dioxide apparently involves a sequence of reactions, it is instructive to consider the available pathways that can lead to insertion. A simple picture of the available mechanisms for CO, insertion is outlined schematically in (52). There are three components: the metal center, the ligand, and the inserting molecule, CO, . Any two of these components may be bound together initially. (a) M-C02 (b) 0,C-X

(c) M-X

X

1

* M-(CO,)-X

M

co* M-X M-X-CO,

M+O,CX-

In (52a), attack of free ligand on coordinated CO, will complete the process, and in (52b) attack of the “ligand-bound” CO, on the metal center stabilizes the X-CO, adduct, completing the reaction. Each of these processes requires some free ligand in solution, either simply dissociated or perhaps protonated as well. If the metal-ligand bond remains intact as in (52c), CO, may initially bind to the (presumably unsaturated) metal center, followed by a ligand migration to the bound CO, , or alternatively, it may attack directly at the bound ligand, resulting in ion pair formation with the metal complex and subsequent bonding to the metal through CO,. The pathway involving ligand migration is analogous to CO insertions (20) and that involving external electrophilic attack is analogous to SO, insertions into metal-alkyl bonds in saturated complexes ( 2 0 4 There are two reasonable structural results of CO, insertion since the metal-ligand pair may add across the carbon-oxygen double bond in two different ways yielding (29) or (30), depending in part on the polarity of the M-X bond.

130

RICHARD EISENBERG AND DAN E. HENDRIKSEN 0

M-U-C-X (2%

1. Insertion into M-H

0

II

and M-C

II

M-C+X (30)

Bonds

The insertion of carbon dioxide into a transition metal-hydrogen bond may be seen as the first and crucial step in the reduction or fixation of CO, . This insertion could proceed in either of two ways: to produce a formate complex, either mono- or bi-dentate [(31) or (32),respectively], or to form a metallocarboxylic acid, (33).

Of the presently known reactions, production of the formate complex predominates. Catalysis of the hydrogen reduction of CO, , which apparently involves insertion into a metal-hydrogen bond, is considered later. Here we consider the insertion reaction itself. The review by Volpin and Kolomnikov (108) surveys the insertion of CO, into metal-hydrogen bonds. Since their review was completed, several additional papers have been published on one of the most interesting and perhaps representative systems. The reaction of [RuH,(PPh,),] with C O , was first reported by Komiya and Yamamoto (129). This complex reacts slowly in toluene solution with CO, to yield the formate complex [(HCOO)RuH(PPh& toluene]. The same workers now report that the analogous complexes with the phosphines PPh,H, PPh,Me, or PPhMe, do not react (132). Evidence was also presented indicating that one triphenylphosphine must dissociate before the reaction with CO, can proceed. Kolomnikov el al. (133) have also examined this system and firid the same product starting with both [Ru(N,)H,(PPh,),] and [RuH,(PPh,),]. They have further confirmed that the product is the formate complex by an x-ray structure determination (108, 133). The most interesting aspect of this system is that the solution chemistry of the product formate complex indicates a formulation as a carbon dioxide-hydride complex, [Ru(C02)H,(PPh3),]. For instance, reaction of the formate with ligands such as PPh,, N, , H2, or CO results in evolution of CO, and addition of the new ligand to the ruthenium center (108, 133), (53a)-(53d). PPh,

d HZRU(PPh3)d

co

H,Ru(CO)(PPh,),

(534


131

The appropriate NMR experiment to determine whether the carbon dioxidehydride complex is formed in equilibrium amounts in solution does not appear to have been done, but recent work on a similar system seems to support this hypothesis. Direct formation of the formate complex [(HCOO)Ru(PPhMe,),]+ was achieved, and when this complex is dissolved in CD,Cl, at 30°C the NMR spectrum shows a broad hydride resonance centered at z = 17.4, indicating the presence of a metal-hydride complex (134), with the CO, possibly coordinated to the ruthenium. Formate complex formation also occurs when solid [H,Fe(PPh,Et),] or [H,Fe(N,)(PPh,Et)3] is reacted with CO, in sunlight, yielding (HCOO),Fe(PPh,Et), (135). No reaction was observed in the dark or at 0°C in the light, indicating the reaction is most likely initiated by a photochemical or thermal dissociation of H2 or N, , respectively. In these reactions it appears that coordination of CO, to the metal center is the first step in the mechanism. Though all the reacting complexes are initially coordinatively saturated, they contain the easily dissociable ligands N,, (H),, and/or PPh3 which can be lost either thermally or photochemically, leaving an electron-rich unsaturated metal center. After initial C 0 2 coordination, the reaction sequence might then proceed by migration of the hydride to the carbon atom through a four-center transition state, followed by formation of the bidentate formate, (54). The analogy of ethylene insertion into a metal-hydrogen bond may be appropriate here. 0

40 H-M’]

-M--O-C’

0

-M

\

\H

/ .%A :C-H \ :Y

(54)

0

The only claim for the production of a metallocarboxylic acid from the insertion of CO, into a metal-hydrogen bond in the opposite sense is based on the reaction of CO, with [HCo(N,)(PPh,),] (108, 136). The metallocarboxylic acid is said to be implicated since treatment of the product in benzene solution with Me1 followed by methanolic BF3 yielded a considerable amount of methyl acetate as well as methyl formate derived from the cobalt formate complex. Metallocarboxylic acid species formed by attack of H,O or OH- on a coordinated carbonyl are considered in the section on CO oxidation. The insertion of carbon dioxide into a transition metal-carbon bond is of importance since it forms a new carbon-carbon bond. The general reaction, ( 5 9 , transforms a coordinated alkyl or aryl group into a coordinated carboxylate. 0

M-R

+ C02

--f

/ ‘..A : C-R \ :Y

M

0

(55)

132


Reactions of CO, with main group organometallic compounds such as organolithium compounds and Grignard reagents (108) are common, but these compounds differ from transition metal alkyls in that they are much less covalent, reacting more as R-. Among the transition metal organometallics, work has been reported with copper(1) alkyls, group 8 complexes, and with early transition metals such as titanium and zirconium. These last reactions, with do transition metal complexes, have been reviewed (108) and may in many respects bear a close resemblance to known nontransition metal reactions. Consequently, these reactions will not be examined, with attention instead focused on those organotransition metal complexes with partially filled or filled d-electron valence shells. Some very interesting work on the C 0 2 insertion reaction has been reported by English and Herskovitz (137). When Ir(depe),Cl (depe = EtzPCH2CH2PEt2)in acetonitrile is placed under CO, , the iridium carboxylate hydride is produced, (56). Ir(depe),CI

+ CO, + CH,CN + [HIr(depe),(02CCH2CN)] + C1+

(56)

This product was characterized by its NMR spectrum and also by reaction with HCl followed by BF3/methanol to yield methylcyanoacetate ester. The reaction occurs readily, and in the absence of detectable amounts of the oxidative addition product of acetonitrile with the iridium complex, [HIr(depe),CH,CN]+. In contrast, neither Rh(depe),Cl nor Rh(dmpe),CI (dmpe = Me2PCH2CH2PMe,) react with CO, in acetonitrile, though Rh(dmpe),Cl does react with CO, in nitromethane to form the analogous nitro-acetate hydride complex, (57). Rh(dmpe),Cl

+ C 0 2 + CH3N02+[HRh(dmpe)2(02CCH2N02)]f + Cl-

(57)

While Ir(depe),Cl does not react detectably with acetonitrile, Ir(PMe3)4Cl does to form the oxidative addition product, (58), and if CO, is present over the reacting solution, its insertion is effected, yielding the product [HIr(PMe3)4(0,CCH,CN)I+ (137). Ir(PMe,),CI

+ CH,CN -+

[HIr(PMe,),CH2CN]+

+ CI-

(58)

Significantly,however, if the CO, is introduced into the solution after formation of the oxidative addition complex with acetonitrile, CO, insertion does not proceed. This is positive evidence that the mechanism of insertion in this particular reaction does not involve direct insertion into the metalcarbon bond. These observations allow a possible mechanism to be outlined in (59)-(6 1). In these systems there is a competition between two bases and two acids, the metal complex and the anion of the solvent (NCCH2- or O,NCH,-) on the one hand, and C 0 2 and H+ on the other. A crucial step in the mech-

133


anism involves the abstraction of a proton from the solvent by the metal complex, resulting in a resonance-stabilized carbanion. The anion may then attack dissolved COz to form -OZCCHzCN or -O2CCHzNO2, which will be further stabilized by coordination to the protonated M(I1I) center.

+

Ir'(PMe,),+ CH,CN CH2CNCO,

+

-+

-+

[HIr"' (PMe3),lZ+ -O,CCH,CN

+ CH,CN-

[HIr(PMe,),JZ+ + -O,CCHzCN -+ [HIr(PMe,)4(0,CCHzCN)]

+

(59)

(60) (61)

Alternatively, the CO, may coordinate to the metal to a kinetically significant extent either before or after proton abstraction, resulting in attack of the solvent anion on a coordinated CO, . In either case it is significant that it is not the COz that is being activated by the metal complex to promote the insertion reaction, but rather the carbon-hydrogen bond in the solvent. Once the metal-carbon bond is formed in [HIr(PMe,),CH,CN]+, this combination of a high-valent metal and an alkyl group with a strongly electron withdrawing group on the alpha carbon is not amenable to either C0,-metal coordination or electrophilic attack of CO, on the metal-bonded carbon, and consequently does not insert CO, . Some closely related chemistry has more recently been reported ( 1 3 7 ~ ) . The complex HFeCH,CN(dmpe), dissolved in CH,CN rapidly takes up CO, to form a cyanoacetate complex. Treatment with Br, or I, yields free cyanoacetic acid, confirmed by treatment with CH,OH/BF, to yield NCCH,COOMe. A case of COz incorporation by apparent insertion into a metal-carbon bond, though not directly detected, is found in the thermolysis of the previously described COz coordination complex, Ir(dmpe),Cl. COz (129).Upon heating at 120°C the coordinated CO, carboxylates a methyl group in the phosphine ligand, formally an insertion into a metallated ligand, (62).

This proceeds in spite of the fact that no metallation of the phosphine ligand could be detected by NMR in the absence of CO, . This may be a common fate of coordinated CO, in complexes where such a reaction may occur. A second example of such metallation and carboxylation is found in the reactions of Fe(PMe,), (221). Carbon dioxide insertion into a rhodium-carbon bond has been found in the reaction of Rh(Ph)(PPh,), with CO, under 20 atm of pressure at room

134


temperature, the product being tris(tripheny1phosphine)rhodium benzoate, Rh(O,CPh)(PPh,), (108, 138). The identity of the product has been confirmed by comparison with the authentic complex formed from benzoic acid and Rh(OH)(PPh,),, and also by an x-ray crystal structure (108, 138). The structure shows the complex to be distorted square planar with a monodentate benzoate ligand. All known CO, insertions into a metal-carbon bond result in carboncarbon bond formation, except in one instance. The insertion of CO, with formation of a metallocarboxylate ester is claimed in the reaction of CO, with a cobalt complex (108, 139). Two species were isolated from the reaction of CO, with a mixture of acrolein and the complex CoH(N,)(PPh,), , which was assumed to form Co(CO)(C,H,)(PPh3f2 before the CO, was introduced. The two reaction products were characterized by their ir spectra and chemical reactions, and formulated as Co(O,CEt),(PPh,), and the metallocarboxylate ester Co(CO,),(COOEt)(PPh,), n = 0.5- 1.O. Metallocarboxylate esters are well-known products from the attack of alkoxides on metal carbonyls. Carbon dioxide insertion into copper(1)-carbon bonds has received particular attention from two groups of workers. An interesting reversible insertion of CO, has been observed by Saegusa et al. (140) in DMF solution with the complex PhC-CCu(PBu,"), , yielding an equilibrium concentration of PhC~CCO,Cu(PBu,"), at room temperature under 1 atm of CO, . The equilibrium is demonstrated by approaching the same relative concentrations from either direction. Alternate heating and cooling of the reaction solution promotes CO, expulsion and insertion, respectively. Other phosphorus ligands such as PPh, and P(OMe), were not as effective at promoting CO, insertion, and phosphine-free DMF solutions of PhCrCC0,Cu were found to decarboxylate irreversibly. Directly analogous chemistry has more recently been reported with the complex NCCH,Cu(PBu,"),, with x = 3 most likely ( 1 4 0 ~ )This . complex inserts CO, in DMF solution to form NCCH,CO,Cu(PBu,"), . Yamamoto et al. (124,141) have recently demonstrated that alkylcopper(1) complexes with phosphine ligands insert carbon dioxide into the carboncopper bond to yield carboxylates. The complexes MeCu(PPh,), -iEt,0, and RCu(PPh,),, R = Et, n-Pr, and i-Bu, in THF solution react with a stream of CO, initially at -40°C with gradual warming to 0°C. In each case the complex (RCOO)Cu(PPh,), is formed. Complexes with other phosphine ligands, PPh,Me, PBu,", PEt,, and P(C,H1 all react analogously to form the carboxylate complexes. In addition, coordinatively saturated methylcopper complexes with three tertiary phosphine ligands such as MeCu(PPh,), .toluene and MeCu(PPh,Me), react slowly with CO,

BINDING AND ACTIVATION OF CO, C 0 2 , AND NO

135

to give acetato complexes, (MeCOO)Cu(PPh,), and (MeCOO)Cu(PPh,Me), (124).

The mechanism of formation of these carboxylate Complexes is not obvious. The authors consider the CO, insertion reaction as proceeding through an intermediate C0,-coordinated alkylcopper complex, but were not able to detect or isolate such species even at low temperatures (124). An alternative proposal involves direct electrophilic attack of CO, on the alpha carbon of the alkyl group, analogous to the mechanism of insertion of SO2 into metal-carbon bonds in coordinatively saturated complexes ( 2 0 4 . This electrophilic attack by CO, on the carbon bound to the metal would be followed by intermediate formation of an ion pair, and ultimately metal-oxygen bond formation to yield the carboxylate complex. Such a mechanism would be consistent with the observed reaction of CO, with the coordinatively saturated tris(phosphine) alkylcopper complexes. A mechanism proceeding through coordination of CO, to the metal would require prior dissociation of one phosphine ligand followed by alkyl migration to the coordinated C 0 2 , also a possible but unproven pathway. Insertion of carbon dioxide into a nickel-carbon bond has recently been demonstrated. The complex Ni(C,H,),(bpy), (bpy = 2,2'-bipyridine), reacts with CO, in benzene at 40-50°C as shown in (62a) (1414. The intermediate nickel(I1) alkyl carboxylate can then react with additional CO, to form either a dicarboxylate or diethyl ketone, as in (62b) and (62c). Ni(C,H,)Abpy)

(33s)

(33a)

+ CO,

-

Ni(C,H,)(OCOC,H,)(bpy)

+ CO, --+Ni(OCOC,H,),(bpy) + CO, Ni(bpy)(CO,) + (C,H,),CO

(33s)

(624 (62b) (624

The intermediate Ni(C,H,)(OCOC,H,)(bpy) (33a) has been prepared separately from Ni(C,H,),(bpy) with one equivalent of propionic acid, and isolated. Carbon dioxide does not insert into all metal-carbon bonds, as demonstrated by the lack of reaction with [HIr(PMe,),CH,CN]+ (137). Similarly, benzylchromium complexes failed to insert CO, (142), and in the reaction of Me(Me,N),WZW(NMe,),Me with CO, , the molecule inserts into the metal-nitrogen bonds, but not into the metal-carbon bonds (143). More recently a benzylchromium complex has been found to insert CO,, though in this case the benzyl ligand is also bound by an o-NMe, inserts one CO, into a linkage (1434. Thus, Cr(CH,C,H,-o-NMe,), Cr-C bond to yield a product formulated with a bidendate carboxylate. In summary, activation of CO, by metal coordination apparently is required for insertion into a metal-hydrogen bond, and may well also be

I36


required to promote insertion of CO, into a metal-alkyl bond. In contrast, when the organic ligand can dissociate as a resonance-stabilized anion, CO, insertion into the M--C bond can proceed instead through direct attack of this free anion on free CO, in solution, followed by coordination of the product carboxylate to the metal.

2. Insertions into M-N

and M-0

Bonh

Very little work on the insertion of CO, into transition metal-nitrogen bonds was published prior to the recent series of papers by Chisholm and Extine (144-148). The only previous unrelated work seems to be the observation that CO, inserts in Cr(NPr,'), (149). Chisholm and Extine have examined the formation, structures, exchange reactions, and mechanisms of formation and exchange of the early transition metal dimethylamides M(NMe,), (where M = Ti, Zr, V and n = 4; M = Nb, Ta and n = 5), W,(NMe,)&, and W(NMe2), . These compounds all react rapidly in solution with carbon dioxide to form, respectively, the N,N-dimethylcarbamato compounds M(OzCNMe2),, W,(O,CNMe,),, and W(NMe,h(O,CNMe,h, the last of which is unique in its incomplete insertions of CO, . The structures of this tungsten complex (144) and of the niobium product (146) of reaction with CO, have been determined by x-ray crystallography. Other dialkylamide complexes of these metals were found to react similarly. Information regarding the mechanism of CO, insertion in these systems has been gained from extensive studies of the exchange reactions (63) and (64) (148).

+ 13C0, MO2l3CNMeZ+ M0,WN(CD3), * MOZWNMe, + M0213CN(CD3), M0213CNMe2+ 12C0, T .M0212CNMe,

(63) (64) These exchange processes were found to proceed by a catalyzed reaction sequence involving adventitious free amine, (65) and (66). The second exchange process, (64), requires fortuitous CO, as well.

+ HNMe, + H0212CNMe, + H0212CNMez-+ MOZ1TNMe,+ H0213CNMe,

'TO2 M0213CNMe,

(65)

(66)

The insertion reaction itself doubtless proceeds by a very similar amine catalyzed process composed of (65) and (67). MNMez

+ HO,CNMe,

+ MO,CNMeZ

+ HNMe,

(67)

Compelling evidence for this mechanism was obtained when a freshly prepared batch of the tungsten complex W,(NEt,)4Me, failed to react with CO, over a period of 24 hours, but upon addition of a trace of free amine, reaction commenced immediately and was complete within minutes (148).


137

The possibility of insertion of CO, into metal-amide bonds via a mechanism not unlike SO2 insertion into metal-alkyl bonds remains a viable possibility in CO, chemistry although it was not realized in these studies by Chisholm and co-workers. In this mechanism, CO, acting as an electrophile, attacks an amide nitrogen, with the intermediate then rearranging to the metal carbamate, (68).

The M-NMe, moiety in these complexes, however, is planar, indicating that the nitrogen lone pair is involved in n-bonding with the metal. It is worth noting that in neither the mechanism discerned by Chisholm nor in the mechanism speculated upon here is activation of C0 2 by the metal complex an integral part of the insertion process. The insertion of carbon dioxide into a transition metal-oxygen bond, e.g., a metal alkoxide, results in an organic carbonate ester, coordinated in either a monodentate or bidentate manner. Only a limited number of such reactions have been observed, and little mechanistic information is available. The reactions may proceed by interaction of CO, with ROH or RO- in solution followed by metal coordination, in a manner similar to the CO, reactions with the early transition metal dialkylamides. Alternatively, direct attack of CO, on the alkoxide oxygen might occur, or a CO, adduct may form as an intermediate. One of the first such reactions to be reported did not involve a simple metal alkoxide compound, but rather alkoxide groups adsorbed on a manganese surface (150). This reaction was monitored by infrared spectroscopy. Initially a manganese surface was evaporated under vacuum onto an oil-covered infrared cell window. Introducing a small pressure of alcohol in the cell, either methanol or ethanol, resulted in bands characteristic of surface metal alkoxides. Addition of C 0 2 to the cell, followed by pumping away the excess after reaction, resulting in bands attributed to the surface organocarbonate, most likely coordinated in a bidentate manner, (69). Heat treatment with pumping leads to regeneration of the metal alkoxide species. $4

ROH

+ Mn -+M n t O R

heat,COl vacuum ~

Mn L $ C - 0 - R

(69)

Treatment of the metal surface with CO, alone forms surface CO, species that do not react with alcohol to form the organic carbonate. Consequently, the mechanism must not involve formation of CO, groups on the metal

138


surface. Direct addition of surface metal alkoxide across the CO,' carbonoxygen double bond is favored, though migration of alkoxide from metal to adsorbed CO, is not ruled out. Other reports confirm similar reaction of CO, with metal alkoxide compounds. Benzene solutions of the alkoxide compounds Ti(OBu"), , Zr(OBu"), , Zr(OEt), , Nb(OEt), , and Fe(OEt), were found to absorb carbon dioxide in various amounts (151). Only Zr(OBu"), reacted with an amount approaching one equivalent per metal complex. The solution ir spectra indicate formation of coordinated organic carbonates. From this evidence, equilibrium formation of the species M(OR),- l(O,COR) under 1 atm of CO, is inferred. In contrast, a suspension of Cu(OCH3), in pyridine was found to react with exactly two equivalents of CO, to yield a blue-green solution of Cu(O,COMe), (152). This reaction could be reversed by heating the solution under a stream of nitrogen. Pyridine must stabilize this complex by coordination to the metal, since removal of this solvent by drying resulted in the loss of CO, . A variety of other solvents failed to promote the reaction. The dimeric compound [Cu(acac)(OMe)], was also found to take up CO, in pyridine solution with formation of [Cu(acac)(O,COMe)(py)] (152), which again could not be isolated since CO, was lost on removal of the pyridine solvent. No reaction of CO, was observed with Co(OMe), or Ni(OMe), (152). The same workers have been able to isolate another similarly reversible copper(1) carbon dioxide carrier, Cu(0,CO-t-Bu)(CN-t-Bu), (153). This complex forms in benzene solution at 10°C according to (70).

-

Cu(0-t-Bu)(CN-t-Bu)

+ C 0 2 + 2CN-t-Bu + Cu(0,CO-t-Bu)(CN-t-Bu),

(70)

The reaction is reversed by refluxing the solution under nitrogen. The coordinated organic carbonate product is assumed monodentate in this complex in accord with the 18-electron rule. The role of the isocyanide ligands as a-donors is emphasized, as evidenced by the ir spectra. The similarly a-donating ligands PEt3 and PPh3 are also found effective in promoting CO, uptake by the copper alkoxide. A recent report also demonstrates reversible insertion of CO, into two of the alkoxy groups in M0,(OCH2CMe3)6 to yield Mo,(OCH,CMe,),(0,COCH,CMe3), (154). This contrasts with the reaction of the related dialkylamide complex W,(NMe,), , which reacts completely with CO, to yield W,(O,CNMe,), . More recently this work has been extended and it has been demonstrated that the Mo,(OR), compounds where R = Me3Si, Me3C, Me,CH, and Me3CCH, all react readily and reversibly both in solution and in the solid state with CO, to give the insertion products Mo,(OR),(O,COR), ( 1 5 4 ~ ) .


, AND

139

NO

The mechanism of CO, insertion in solution is proposed to be analogous to that observed for the dialkylamide complexes discussed earlier, and is shown as (70a) and (70b).

=+ROCOOH MOOR + ROCOOH +Mo0,COR + ROH ROH + C 0 2

(70a) (70b)

This mechanism involves catalysis by traces of alcohol in solution and is supported by a lack of reaction in one case when rigorously pure materials were used. It is further supported by the observation that the reaction proceeds faster in the presence of added alcohol. Alkoxide ligand interchange has been demonstrated to occur during the course of this CO, insertion in solution. This mechanism, however, cannot be operative in the solid state. The heterogeneous reaction between the solids and CO, gas occurs much more slowly than in solution, and there is no alkoxide ligand exchange, indicating a different kinetic pathway. The mechanism proposed in the solid state reaction is a direct electrophilic attack of CO, on the Mo-OR bond, possibly proceeding through a Mo-CO, intermediate. The reactions of [RuH(PPhMe2)JPF6 were previously mentioned in the context of formate complexes and their equilibrium formation of C0,hydride complexes (134). This compound also has an interesting chemistry in its formation of alkyl carbonate complexes (134). When [RuH(PPhMe,),]PF6 is dissolved in methanol or ethanol under COz, the alkyl carbonate complex forms, (71), [RuH(PPhMe,),]+

+ CO, + ROH

--f

[ R U ( O ~ C O R ) ( P P ~ M ~ , )+~ ]H3 +

+ PPhMe, (71)

which is shown not to be the corresponding carboxylate complex by comparison of the two infrared spectra. In spite of the fact that the formate complex also yields the alkyl carbonate on recrystallization from alcohol, (721, [Ru(O,CH)(PPhMe,),]+

+ ROH -+ [Ru(O2COR)(PPhMe,),] + Hz +

(72)

the authors consider that the alkylcarbonate is not formed through alcohol attack on the formate, but rather by interaction of CO, with an intermediate alkoxide complex, (73). Ru-H

+ ROH

co

--f

0

RuOR ---% RU + H2

: C-OR

\ ,,Y

(73)

0

The intermediate formation of a carbon dioxide-hydride complex from the formate might facilitate such a mechanism, (74).

140

-

RICHARD EISENBERG AND DAN E. HENDRIKSEN Ru(O2CH) ---L HRu(CO2)

ROH

Ru(CO,)(OR) + Ru(O2COR)

(74)

+ H2

In exploring the scope of this reaction, the authors found similar chemistry with the complex [H,Ru(PPh,),] (134). The alkyl carbonate could be formed directly by reaction with CO, in alcohol, or the formate could be produced as an intermediate by reaction with CO, in benzene, and then converted to the alkylcarbonate by recrystallization from alcohol. This analogous behavior may be taken as possible evidence for the equilibrium formation of a C0,-hydride complex in solution from the formate, [RuH(O,CH)(PPh3)31. Strong evidence for an insertion mechanism involving an intermediate CO, adduct is found in the reactions of [M(CO)(OH)(PPh,),], M=Rh, Ir. The authors initially found that CO, reacts with these hydroxide complexes in ethanol to form the bicarbonato complexes (123, postulated as containing a monodentate ligand. The CO, adducts, formed from the metal hydroxide complexes in the solid state, react irreversibly, quantitatively, and rapidly when dissolved in ethanol to form the bicarbonato complexes (120). This suggests that formation of these bicarbonato complexes in ethanol proceeds through the CO, adduct. The insertion of CO, into a metal alkoxide bond is unique in that it can occur into either the M - O bond or the 0-R bond leading in both cases to the same product, (75). 0

M-*OR

+ C02

II

---f

M-OC-*OR

0 or M-*O-COR

II

(75)

There is actuqlly good evidence for the latter possibility in the formation of certain transition-metal bicarbonato complexes from metal hydroxides. Detailed kinetic and mechanistic data have been obtained for the reaction of CO, in water solution with [M(NH3)50H]2f,M = Co, Rh, Ir (155, 156). It was found that of all the species connected by acid equilibria, CO, uptake was effected by only one reaction, (76).

+ COz =+ M(NH3)50C02HZ+

M(NH3)PHZ+

(76) The decarboxylation reactions were similarly studied (155, 156). In these complexes, prior coordination of CO, to the metal is blocked by the amine ligands, and is further unlikely because of the high oxidation state of the metal. All the evidence indicates very strongly that reaction occurs by direct attack of CO, on the coordinated hydroxide oxygen, and that the bicarbonate ligand is formed without rupture of the metal-oxygen bond. However, the hydroxide group is probably unique in facilitating this mechanism through the lability and ready transfer of the proton, as opposed to transfer of an alkyl group.


CO, C O , ,

AND NO

141

In general then, activation of CO, by coordination to the metal is apparently not required for insertion of CO, into a metal-amide or metal-alkoxide bond. Instead, reaction between free ligand and CO, in solution may take place, or the CO, may attack directly the coordinated ligand, followed by a rearrangement to yield the product. C. CO, REDUCTION AND/OR INCORPORATION In the preceding discussion various simple reactions of CO, with transition metal complexes have been described, namely complex formation with CO, and insertion of CO, into metal-ligand bonds. In addition to the insight and knowledge gained directly from this work, it serves to lay the groundwork for development of catalytic processes leading to CO, reduction and/or incorporation into organic molecules. Some such processes have been reported. Two reports have appeared on the catalyzed reaction of CO, with epoxides to form alkylene carbonates. One of the processes uses phosphine complexes of zerovalent nickel as the catalyst (159, and appears closely related to the more recent isolation of (PCy3),Ni(C02) (115). Ethylene oxide reacts in benzene under 500 psi pressure of CO, in a stainless steel autoclave at 100°C to form ethylene carbonate with 95% selectivity, (77), using as the catalysts NiL, , L = PCy, or PPh, .

0

/ \

CH2-CHz

+ C02 +

(77)

The free phosphines alone show some catalytic activity under these conditions. Other epoxides are also shown to react, and some intermediates, (34)-(36), are proposed.

/"k

L.Ni \C/

il

0

(34)

(35)

(36)

A second report of organic carbonate production from epoxide and CO, utilizes copper(1) cyanoacetate, Cu(O,CCH,CN), as a carrier of activated CO, (158). Reaction of propylene oxide with Cu(O,CCH,CN) at 130°Cfor 10 hours yields propylene carbonate in 83% yield, based on the

142


cyanoacetate. It was also found that either Cu(CH,CN) or Cu(O,CCH,CN) acts as a catalyst for the reaction when it is carried out under CO, . Several reports have described the formation of alkyl formates or formamides from CO, , hydrogen, and an alcohol or amine. The earlier catalytic preparations have been reviewed (108).A recent paper describes the production of alkyl formates catalyzed by several transition metal complexes and tertiary amines under 25 atm each of CO, and H, at 140°C, (78) (159). ROH

+ CO, -t H2 + HCOOR + H,O;

R = Me, Et, Pr

(78)

Typical catalysts are Pd(diphos), and RuH,(PPh3), , while Et3N is a representative co-catalyst amine. The same workers have more recently used a similar catalyst system to produce formic acid (formate ion) from hydrogen and carbon dioxide (160). The catalyst system is again Pd(diphos), and Et,N in benzene solution under 25 atm each of CO, and H,. The reaction is run at room temperature or higher. A small amount of water dramatically accelerates the rate of the reaction, and a mechanism is proposed to account for this effect. Control experiments are stated to rule out the initial reduction of CO, to CO by H, , followed by reaction with water to yield the formic acid. It should be noted that the production of formic acid from CO, and H, proceeds with a net increase in free energy under ambient conditions, but that an increased pressure of H, and CO, will shift the equilibrium favorably. When alkyl formates or formamides are produced from CO,, H,, and the amine or alcohol, the stability of the water formed in the reaction provides a powerful driving force, and the thermodynamics of these reaction are favorable under ambient conditions. More recently another report on the catalytic conversion of CO,, H, , and alcohols into formate esters has appeared (160~).This work uses as catalysts the anionic iron carbonyls HFe(CO),- and HFe,(CO), , and reports modest conversions to alkyl formates under conditions of elevated pressure and temperature. The synthesis of 1,2-propanediol formates from carbon dioxide, hydrogen, and propylene oxide has been achieved using RhCl(PPh,), as the catalyst, (79) (161). CH3 COZ

+ H2 + H2C'0'

I

CH

-

CH3 I HzI-IH H + H HC

II

0

CH3

CH,

- H 6CH 0

f

I

H - H

I-I:l

I1

(79)

ICH

0

The two mono-formates are formed in comparable amounts, suggesting little regioselectivity in the oxirane ring opening reaction. Propylene carbonate, 1,2-propanediol, and formic acid were also found as reaction prod-


143

ucts. This reaction was typically carried out at 100°C under 40 atm each of H, and CO, for 40 hours. A proposed mechanism involves initial production of a metal formate, which then enters the reaction cycle, shown schematically as (80).

0

A = HO

Two parallel routes for the elimination of glycol formate are suggested, involving either reaction with H, or with cocatalyst water. The detection of formic acid in the reaction products suggests another mechanism, with initial production of formic acid from H, and CO,, followed by reaction with the oxirane. This mechanism is not favored however since the yields of glycol formates varied substantially when various substituted oxiranes were reacted. This would not have been expected in a mechanism with formic acid as an intermediate. A third mechanism, not considered by the authors, could proceed through initial production of propylene carbonate, followed by reduction to the mono- or di-formate. Two recent reports of noncatalytic metal complex-promoted reduction of CO, merit discussion here. Work by Floriani et al. details reactions of CO, with early transition metal complexes in toluene solution leading to deoxygenation or disproportionation and reduction (161~).The three reactions are shown as (80a)-(80c). [CpzTiC112 + COz

90". 18h atm

4 c ~ , T i ( C 0 )+~ 4*C02

52",2d

+ 3co2

74", 4d

3Cp,Zr(CO),

. [Cp,TiCl],O + CO . ~ ( C P ~ T ~ ) ~ ( * C +O ~2*CO ) I Z + 8CO (Cp,ZrO)3

+ 9CO

(80a) (80b) (80c)

The disproportionation of the CO, in (80b) to CO and C0:- has been verified by isotope labeling as shown, and the structures of the metal-containing products in (80b) and (80c) have also been determined. A report of the disproportionation and reduction of CO, by reaction with

144


a variety of metal carbonyl anions has also appeared (161b). In particular, the reaction with CpFe(CO),- was studied, which proceeds as in (80d). 2Na+[CpFe(CO),]-

+ 2C02

THF

[CpFe(CO),I1 + NazC03 + CO

(god) A proposed mechanistic sequence involves stepwise coordination of two molecules of CO, to CpFe(CO),-, followed by formation of Na,CO, and transient formation of CpFe(CO),+ which quickly reacts with more of the anion to form the dimer. Support for this mechanism comes from the observation that reaction with labeled CO, results in extensive and rapid incorporation of labeled CO in the product dimer. Finally, the hydrogen reduction of CO, to CO and water, (81), has been achieved, with the product CO bound to the metal promoter (162). RhCI(PPh&

+ H,(20 atm) + COZ(20atm)

RhCl(CO)(PPh,),

+ H,O

(81)

The solvent used was hexamethylphosphoramide. As in the production of formic acid (160), this reaction is markedly accelerated by the addition of water, and a proposed mechanism includes the formation of formic acid and its subsequent decomposition to CO and H,O, with the CO being bound by the rhodium, (82). H2 + C02 --t HCOOH + HzO

+ CO (as metal carbonyl)

(82)

Note that this overall reaction is essentially the reverse of the water-gas shift reaction, and would be thermodynamically slightly unfavorable if the product CO was not strongly bound in the metal complex. These few known examples of catalytic reactions in which CO, is either reduced by hydrogen or incorporated into an organic molecule provide only a modest indication of the potential utility of CO, as a feedstock for organic chemicals production. The mechanisms of these reactions in most instances have not been unequivocally established, especially with regard to the metal-CO, interaction and the notion of CO, activation via coordination. Future work will extend our knowledge of the nature and scope of these catalytic reactions, and will establish a firmer base for an understanding of their mechanisms, including the necessity of activating the stable CO, molecule. IV.

Nitric Oxide

One of the relatively few simple odd electron species, nitric oxide is an intriguing heteronuclear diatomic and the parent member of the oxides of nitrogen. Like carbon monoxide, nitric oxide has a long and distinguished coordination chemistry, but unlike CO, it forms very few binary metal


CO, COz , AND

NO

145

complexes of formula M(NO),. Instead, NO is found in transition metal complexes in conjunction with ligands ranging from “harder” or “type A ’ ligands such as ammonia and chloride to “softer” a-acceptor type ligands such as phosphines and carbon monoxide (163, 164). Nitrosyls are also well known in organometallic systems in which alkyls, olefins and a-complexed carbocycles occur (165). A number of excellent reviews on nitrosyl complexes, their structures, bonding and reaction chemistry have appeared (163-1714.

The main stimuli to the investigation of coordinated NO reactivity come from the widespread occurrence of oxides of nitrogen as an environmental hazard produced in many combustion processes (172, 173), and from the fact that NO is only a proton and an electron removed from CO with orbitals that are changed slightly, but significantly, in terms of energy and polarization (174). Each year over lo6 tons of oxides of nitrogen, NO,, are produced in fossil fuel combustion processes with approximately 60% of this amount generated from stationary sources such as electric power plants and industrial boilers. At flame temperatures, the oxides of nitrogen are produced mainly by @3a) fN2

+

402

--+ NO

(834

as NO which then undergoes aerial oxidation to the highly toxic species NO,. One line of investigation in the environmental control of NO, is the development of catalysts for the facile conversion of NO into less harmful chemical entities. Despite its thermodynamic instability (AGf0 = 20.7 kcal/ mol) and its radical nature, NO is kinetically inert with respect to decomposition and reduction, and requires the presence of a catalyst for many of its reactions. Since the 1920s, studies of the catalyzed decomposition of NO using numerous metals and metal oxides have been performed (175-178). Catalytically active oxides include Rh,O, , Co,O,, CuO, and La,O, , both with and without the aid of supports. In addition, researchers in the field have examined the reduction of nitric oxide using gaseous species such as CO, NH, and H, (179-185). It is important to realize that in the overwhelming majority of these studies (175-182) the catalyst systems have been heterogeneous. Despite the considerable efforts in this area, a detailed mechanistic picture of the stoichiometrically simple, but kinetically complex reactions of NO has not fully emerged. Recently some successes using homogeneous catalyst systems have appeared (183-189). In addition, stoichiometric conversions of coordinated NO have been observed (190-194) which also provide a clue to the mechanism of catalyzed conversions and the development of new catalysts.

146


A. THECOORDINATION OF NITRIC OXIDE In considering the activation of nitric oxide by coordination, it is advisable to review briefly the modes in which NO binds to transition elements. It is here that an interesting contrast with CO emerges. Whereas carbon monoxide coordinates as a terminal ligand in only one way [structure (3)], nitric oxide shows a variability in its manner of terminal coordination. This is evidenced most prominently in the structural arrangement of the MNO unit as linear or bent (167-169). The linear mode, as shown in resonance structures (37a) and (37b) parallels directly carbonyl coordination [cf. (311 and exhibits a significant n-bonding interaction between the metal and the nitrosyl. In the linear mode of coordination, the nitrosyl acts as a 3e- donor according to the electroneutral formalism, or as NO' by the oxidation state formalism. Whatever the formalism employed, the metal nitrosyl n-interaction is a dominant feature of the bonding, and the value of vN0 provides direct evidence of the relative importance of (37b).

j +

:0:

I1

N

I1

M (b)

:M-

(4 (37)

As with carbonyl coordination, the degree of backbonding in the linear mode of coordination is influenced by the complex charge, metal oxidation state and nature of the ancillary ligands. An example of how both the oxidation state and the relative d orbital energies influence vN0 is shown dramatically in the pentacyanonitrosyl systems M(NO)(CN);- ( z = 2, M = Fe, Mn, vN0 = 1939, 1885 cm-'; z = 3, M = Mn, Cr, vN0 = 1725, 1645 cm-') (195).While the M-NO n-bonding in all of these pentacyanonitrosyls is strong, dominating the electronic structure, the transfer of electron density into nitrosyl n* orbitals relates to the difference in energy between metal d, and ng0 orbitals, and hence the composition of the filled metal nitrosyl n-bonding molecular orbital (195). When relatively little electron density is transferred to the nitrosyl by this interaction, as evidenced by vN0 greater than ca. 1880 cm- the nitrosyl exhibits true NO' character and is activated to attack by nucleophiles (171).Below that value, sufficient electron density has been back-donated by the metal to remove a reactive charge distribution. The bent mode of nitrosyl coordination, as represented by Lewis structure (38),provides the nitrosyl with structural uniqueness.

',

.. . H?.

"

I

M


147

In this mode of binding, the nitrosyl serves as a 1 e- donor, coordinating formally as NO-. Although proposed many years ago (196), the bent mode of coordination was first established in 1967 by Ibers and Hodgson for [Ir(NO)Cl(CO)(PPh,),]+ (193, and since then has been determined for many compounds by x-ray structural techniques (167-169). In (38),the donor N atom possesses a nonbonding lone pair of electrons. Thus the charge affinity of the N atom has been inverted. Whereas the linear nitrosyl undergoes nucleophilic attack when vNo is sufficiently high, the bent nitrosyl is set for electrophilic attack at the donor N atom. Indeed this has been observed in a number of cases discussed below. The structural studies of nitrosyls have shown to date that the bent nitrosyl ligand invariably occurs at the apical position of a square based pyramid or a distorted octahedron in which the metal ion configuration assuming NO- coordination is d6. Despite numerous electronic structural descriptions (168, 169, 198-200), it is not totally clear why fully bent nitrosyls with M - N - 0 bond angles of 120" have not been found in other geometries such as the square plane and the trigonal bipyramid. The NO' and NO- modes of coordination differ by two electrons in terms of formal charge. Interconversion of these two bonding modes becomes feasible when the bound metal ion possesses two complementary oxidation states. It has been proposed that this interconversion, which corresponds to an intramolecular redox reaction, represents a unique and facile way to achieve coordinative unsaturation at the metal center with the nitrosyl acting as an electron pair reservoir (201). Interconversion of linear and bent nitrosyls has been reported in the unusual complex Ru(NO),CI(PPh,),' that possesses one linear and one bent nitrosyl, structure (39) (202).

-

By carrying out the synthesis of this complex with one of the nitrosyls labeled, Collman (203)observed four vNo bands corresponding to the labeled and unlabeled nitrosyl in each position, thus indicating that an interconversion of the two bonding modes was occurring. In their extensive studies of metal nitrosyl chemistry, Enemark and Feltham showed that the mode of NO bonding can be altered by the simple addition of another ligand (168, 204, 205). An example of this phenomenon is illustrated by reaction (83b), and has been described by these investigators as stereochemical control of valence (204).

148


As. ..

(As

As = diars)

In this reaction the nitrosyl ligand bends, undergoing a 2 e- reduction, rather than form a 20 e- complex. This reaction can also serve to activate the nitrosyl ligand. Whereas the linear nitrosyl may be unreactive if vNo is sufficientlylow, the bent nitrosyl is reactive to electrophiles. The intermediate mode of coordination for terminally bonded NO occurs in certain paramagnetic systems and dinitrosyl complexes. In these systems, the M-N-0 bond angle is in the approximate range 150-165" and the values of vNo suggest a strong backbonding interaction. For example, in Ir(NO),(PPh,),+ the M-N-0 angle is 164(1)O and vNo are 1760 and 1715 cm-I (206), while in the S = 4 complex Fe(NO)(diars),(NCS)+ the ,v are 1 5 9 ( 1 ) O and 1620 cm-' (207, respective values of M-N-Oand 208). Bonding descriptions have been put forth to explain the partial bending of the MNO unit in low spin S = 3 complexes of this type (168,198-200): It appears that in these systems the nitrosyl ligand exhibits neither nucleophilic nor electrophilic character and that the primary mode of reaction of these odd electron complexeslies in the realm of one electron redox chemistry as in (84) and (85) [Fe(NO)(diars),X]+ -+ [Fe(NO)(diars),X]*+ + e[Fe(NO)(rnnt),lZ- -+ [Fe(NO)(mnt),]e-

+

(84) (85)

where diars = o-phenylenebis(dimethy1arsine) and mnt = maleonitriledithiolate (207, 209). These reactions result in a change in nitrosyl bonding mode with little other structural modification. There has not been an extensive investigation of the behavior of these complexes as radical species. The most extensive group of dinitrosyl complexes are four-coordinate, pseudo-d" systems which possess distorted tetrahedral geometries (167, 168). Of these, only a few have slightly bent nitrosyls including Ir(NO),(PPhJ2+ ( 2 0 , Rh(NO),(PPh,),+ (210), and the tetranuclear system Co,(NO),(NO,),(N,O,) (211). The chemistry of these systems with respect to the conversion of the nitrosyl ligand into other chemical entities has been addressed briefly and is discussed under the catalyzed reduction of NO. The tetranuclear Co complex is very intriguing since it forms from Co(N0)(CO), + NO in a sealed tube. The observation of NO,- bridges suggests that some of the NO was reduced in the formation of the complex but no analysis of the product gases was reported (211).

BINDING AND ACTIVATION OF CO, CO2 , AND NO

149

The bridging mode of nitrosyl coordination, (40), has been reported in a relatively limited number of complexes. :0:

II

N,

M’

M ‘ (40)

Both symmetrical and unsymmetrical bridges (i.e., bridges having equal and unequal M-N distances, respectively) have been observed. An example of the former is represented by R U , ( N O ) ~ ( C O(212) ) ~ ~ while that of the latter is Mn,(NO),(NO,)(q 5-C5H5)2 (213). The bridging mode of NO coordination does not appear at this time to be an important structure in activating NO except possibly to Lewis acid attachment at oxygen. Reactions of coordinated NO that involve a change in the oxidation state of nitrogen will probably not proceed via this coordination. In this section the bonding modes of nitric oxide have been described. We next focus on the reactivity patterns that the modes of nitrosyl coordination impart to the bound nitric oxide molecule, and how they have been observed in nitrosyl reaction chemistry. B. REACTIVITY PATTERNS OF COORDINATED NITRICOXIDE

The linear and bent modes of NO coordination give the nitrosyl ligand true amphoteric character in the kinetic sense. The bent or NO- ligand is activated to electrophilic attack at the nitrosyl nitrogen, while the linear or NO’ moiety is subject to nucleophilic attack if insufficient electron density has been back-donated to remove the positive charge at N. In surveying briefly the reactivity patterns of coordinated NO through some of its stoichiometric conversions, the ultimate product of the nitrosyl ligand becomes secondary. Both oxidation and reduction of the coordination nitrosyl can take place by reactions commencing with either nucleophilic or electrophilic attack. The stoichiometric conversions simply illustrate the facility with which coordinated NO is activated to each type of attack. 1. Electrophilic Attack In the bent mode of bonding, the nitrosyl N has approximately sp2 hybridization with a lone pair of electrons. Electrophiles that have been reported to attack at this position include O2 (192,214,215), H+ (190, 205), and NO (193). The first of these reacts with nitrosyl to give nitro and nitrato species, depending on the particular complex system in question. In a study of the

150


oxygenation of cobalt nitrosyls having quadridentate SchiiT base ligands, Basolo and Clarkson (192) observed solely NO,- formation that was believed to take place by the sequence (86)-(89) (L4= acacen, benacen, etc.; B = pyridine, PMe,Ph, etc.). CoLd(N0) BCoLd(N0)

+B

+

0 2

(86)

BCoLd(N0) /O -+ BCoLd(N

1

(87)

\O-O /O BCOL~(N

)

+ BCoL4(NO)

BCoL,(N

/O

\O-O

\O-O

BCoL,(N

---f

/O

O \ /N)CoL4B

(88)

O\

(89)

N)CoL4B -+ 2BCoL4(N02) \O-O/

Kinetic measurements showed a first-order dependence in both complex and oxygen concentrations. The rate determining step was therefore (87), corresponding to an electrophilic attack at NO- (192). Base is required to bend the nitrosyl, and in its absence, the reaction is greatly retarded. The oxygenation of the related Co system BCo(DH),(NO) where DH = dimethylglyoximate,B = pyridine derivatives, PPh, and N-methylimidazole, has been studied by Trogler and Marzilli (214). The main difference from the study by Basolo is the observation of a mixture of products. An 0-bonded nitrato compound is obtained as the major product in this reaction with the expected nitro species giving less than 50% of the product mixture. The authors were unable to explain the occurrence of the product mixture, or elucidate the mechanism of 0-bonded NO,- formation. It does seem clear, however, that electrophilic attack by 0, at NO- is the key step in the reaction. In the absence of added bases, the reactivity of the nitrosyl should be diminished with respect to electrophilic attack, and it is found that only nitro products are obtained (214). Exclusive oxidation of the bent nitrosyl ligand to NO,- has been observed by Kubota and Phillips (215) for the complexes Ir(NO)ClX(CO)(PPh,), (X = C1, Br, I, NCS, N,). These six-coordinate octahedral complexes have bent nitrosyl ligands with vNo in the range 1560- 1540 cm- which are among the lowest reported for terminally bonded NO. Again oxygenation proceeds via electrophilic attack on the nitrosyl ligand. It is tempting to speculate that nitrato formation then occurs by rearrangement of the peroxynitrite intermediate, (90), but this proposal is as yet without proof.

'


151

When the oxygenation is carried out on Ir(NO)Cl(CO)(PPh,),+ in the presence of bases such as pyridine, nitro and nitrato mixtures are once again obtained (215). Electrophilic attack by H+ on bent nitrosyl ligands leads to reduction (290, 205, 216). Complexes containing HNO, NHOH, and NH,OH have all been prepared in this manner, raising the possibility, as yet unrealized, of a catalytic reduction of NO to hydroxylamine. The simplest and best characterized of these reactions is (91) reported by Enemark et al. (205). [Co(NO)(diars),Br]

+

+ H+

--f

[Co(HNO)(diar~),Br]~+

(91)

The geometry of the starting complex is established from an x-ray study to be octahedral with a bent NO- ligand. The related 5-coordinate complex [C~(NO)(diars),]~+ in which the nitrosyl is linearly coordinated does not protonate. The reversible formation of hydroxylamine was reported in 1970 by Roper et al. (190). In this reaction, (92), the bending of the nitrosyl ligand is presumed to occur in situ followed by successive protonations. NHzOH

The electrons for the reduction come from the metal that is initially in the -I oxidation state and which ends up as an Ir(II1) ion. A completely analogous reaction commencing with the rhodium species Rh(NO)L, has also been reported by Collman (216). The electrophilic attack of nitric oxide on a bent nitrosyl is now realized to be the path by which hyponitrite-bridged Co species are formed. Reaction (93) was known since the time of Werner ( 2 1 3 , but the black and red isomers of [CO(NO)(NH,),]~+obtained from this reaction defied definitive characterization for many years. It has now been established that the black isomer is a mononuclear, octahedral complex of Co(1II) and NO- (218) while the red isomer is a hyponitrite bridged system containing two Co(I11) ions (219). Co"

+ NO + NH,

---t

[Co(N0)(NH,)J2+ black and red isomers

(93)

152

RICHARD EISENBERG A N D D A N E. HENDRIKSEN

In reaction (93) the black isomer is initially formed ; electrophilic attack of NO followed by coordination of the N,O,; ligand to a second Co(I1) center may then be proposed to yield the red isomer. This sequence is shown as (94).

. / o 2+

N

An additional example of electrophilic attack on NO- by free NO appears to occur in (95) which corresponds to a metal promoted disproportionation of NO into N,O and coordinated NOz- (193). CoCl,

+ 2en + 3 N 0 + Co(en),(NO,)Cl, + N,O

(95)

en = NH2CH2CH2NH2

In (95), the initial formation of the bent nitrosyl species [Co(NO)(en),Cl]+ (220) is followed by electrophilic attack of free NO and then by reaction with a second NO molecule leading to products. Since the nitro complex produced in this reaction has the same stereochemistry as the reactant Co**'-NO- species, it has been proposed that the Co-N(nitrosy1) bond remains intact throughout the reaction sequence, thus requiring free NO to attack via its oxygen end (193).This proposal requires further proof. All of the reactions outlined above show by example that bent nitrosyls are activated to electrophilic attack. 2. Nucleophilic Attack The propensity of the linearly bonded nitrosyl group to undergo nucleophilic attack is inversely related to the effectiveness of back-donation in metal nitrosyl n-bonding. Below a vNo value of ca. 1860 cm-', the nitrosyl is unreactive with nucleophiles. Above that value, and especially above 1900 cm-', the nitrosyl ligand reacts with a variety of nucleophilic reagents at the N atom (171). This type of attack'is the most widely known and best studied in metal nitrosyl chemistry. Reactions (96)-( 101) provide examples of nucleophilic attack on linearly

BINDING AND ACTIVATION OF CO, COz , AND NO

153

coordinated nitrosyl. The reaction of the nitroprusside ion with OH- leading to the formation of a nitro complex, (96),

+

Fe(NO)(CN)52- 20H(vNo 1939 cm-I) Ru(NO)(NH&~+ NH,OH (1925 cm-') Ru(NO)(NH3)2+ + NzH4

+

-

Fe(N02)(CN),4-

+ HzO

RU(N~O)(NH,),~+ + H30+

R U ( N ~ O ) ( N H ~ )+ , ~Ru(NdNH,?,+ +

RU(NZ)(NH~)~~+ IrC13(NO)L2+ ROH ---+ IrCl3(RONO)LZ+ H+ (1945 cm-l; L = PPh3) RuCl(A-A)2(NO)Z' 20H- aRuCl(A-A)Z(N02) + H20 (vNo = 1927, 1886 cm-l; A-A = bipy, diars)

+

+

(96) (97) (98a) (98b) (99) ( 100)

[sol (=solvent) RuCl(A-A),(Sol) + N, N,O

+

+

is probably the earliest characterized of these transformations (163-221). As illustrated in (96)-(lOl), both oxidation (191,222) and reduction (223225) can occur by this manner of attack. Nitrogen nucleophiles lead to reduction of the coordinated nitrosyl yielding either free or complexed nitrous oxide, dinitrogen, azide or hydrazine derivative (223-225). The review by Bottomley (171) discusses the electrophilic behavior of nitrosyl in reactions of the type shown above. A novel illustration of the electrophilic behavior of coordinated NO+ is provided in a recent report by Feltham (226)in which the specifically labeled nitro nitrosyl complex cis-Fe(' 5N0)('4N02)(S,CNMez)z was found to undergo an intramolecular oxide transfer according to (102).

The reaction suggests the possible employment of the nitrosyl+ nitro conversion for oxide ion acceptance and donation, and this in turn could play a role in the future development of nitrosyl complexes as oxygenation agents. One objective in elucidating the chemistry of coordinated NO is to estab-

154


lish the ease with which nitrogen-carbon bonds can be formed in these systems. To date, C-N bond formation arising from nucleophilic attack on linearly bonded nitrosyl has been observed in only a few selected instances. In (103), the electron-rich n-system of the N,N-dialkylaniline attacks the nitrosyl, binding it at the position para to the - NR, functionality (227). [R~Cl(bipy),(NO)]~+ + C6HsNR, ---f [R~Cl(bipy)~(N(O)c~HyR2)1' + H+

( 103)

Viewed in another way, the arene is nitrosated by metal coordinated NO+. A second example of C-N bond formation occurs in the reaction of Fe(NO)(CN):- with ketones in base giving rise to red intermediates which decompose to give oximes (228). The reaction appears to proceed via the sequence (104)-(106) in which an enolate ion of the organic carbonyl nucleophilically attacks the nitrosyl ligand giving (41) which then hydrolyzes to Fe(H,O)(CN):- and the oxime. Further studies to develop nitrosyls as nitrosating agents are in progress, but the transformations reported to date are stoichiometric and, at this point, limited in scope. 0

0-

II

+ OH-

CH3-C-R

I T- C H T C - R

+ HzO

( 104)

1 O V C H

I ,CN + Fe(NO)(CN)sz- + NC, NC,Fe,

-R

4-

+ H+

(105)

L CN

N (41) (41)

+ H,O+ + Fe(H,0)(CN)s3-+ RCC-NOH I1

(106)

0

3. An Important Diference The use of metal nitrosyl complexes for the purpose of forming carbonnitrogen bonds represents an appealing entry into the synthesis of organonitrogen compounds. One aspect of nitrosyl chemistry that poses an interesting and significant problem in this regard is the apparent reluctance of nitrosyls to undergo migratory insertion via (107). This apparent inability stands in sharp contrast with the prevailing situation in metal carbonyl chemistry in which alkyl carbonyls readily convert to acyl species.

2.

M-NGO:

B

( 107)

+ M-N

R'


155

The symmetry of transformation (107) is identical to that of the alkyl carbonyl-to-acyl conversion, ( 9 , and therefore the difference in the propensities of these insertions to occur must relate to differences in the orbital energetics of the two processes. Since the metal-nitrosyl n-interaction is a dominant feature of the electronic structure in NO' complexes, and since it will be weakened by the formation of a nitroso species, one can expect (107) to be correspondingly less favorable than acyl formation in (5). If reaction (107) does occur, however, one can readily envisage tautomerism of the nitroso ligand to an oxime. The reactive nature of the oxime may then create difficulties in the identification and isolation of organonitrogen products, and indeed, this may have obscured general observation of transformation (107) in the past. In view of the dual bonding modes of NO, migratory insertion (107) is not the only transformation of this type that can be envisioned. In (108), the nitrosyl is coordinated as NO-, and the migration of the R group to form a coordinated nitroso or oxime compound can be viewed, at least heuristically, as a 1,2-shift of Rf . 0

M-N'

d

+M-N

No

(1 08)

R'

This represents an inversion of the charge affinity of the migrating alkyl group from what it is in (107), and raises many interesting speculations for future investigation. However, at the present time the viability of any migratory insertions in alkyl nitrosyls must first be established. One intriguing reaction that may bear on this problem is reported by McCleverty (229).The Ru(NO),L, (L = PPh,) complex in (109) is known to be a coordinately saturated, pseudo-tetrahedral complex With linearly coordinated nitrosyls (230), Ru(NO),L,

+ PhCHzX + CO -+ RU(CO)~X,L~ + PhCH=NOH + PhCN

(109)

but in the presence of CO may form a transient adduct in which one of the nitrosyls is forced to bend, (1 10).

The bent nitrosyl may then act as a nucleophile on the benzyl halide, leading to PhCH=NOH and PhCN, the latter arising from the former by dehydration.

156


A parallel situation appears to obtain for the mixed allyl nitrosyl complex Ru(NO)(C,H,)L, prepared by Schoonover and Eisenberg (231). This complex which is coordinatively saturated (NO’ and q3-allyl), forms a CO adduct which is assigned a bent nitrosyl structure (231). Further reaction under CO leads to the formation of Ru(CO),L, with the possible elimination of acrolein oxime. The coupling of the allyl and nitrosyl ligands can be viewed in this case as nucleophilic attack of NO- on an y3-allyl species. Unlike in reaction (1 lo), both of the moieties to be coupled lie within the same coordination sphere. The significance of these results is that it lends viability to the notion embodied in (109) in which a migratory insertion of nitrosyl occurs as NO-. The coupling of allyl and nitrosyl moieties is relevant to modeling the heterogeneously catalyzed synthesis of acrylonitrile from propylene NO, equation (1 1 1) (232).

+

4A

+ 6N0

---t

+

+

4 f i C p ~ 6H20 N2

(1 11)

In this catalysis a surface allyl species is formed which combines with NO to form 3-nitrosopropene. Tautomerism and dehydration then lead to the acrylonitrile product. The coupling of the allyl and nitrosyl ligands in Ru(NO)(CO)(C,H,)L, (231), via (1 12) represents a key step in the proposed reaction sequence and suggests the necessity of a bent nitrosyl, at least in a discrete complex case.

A similar view has very recently been presented by Clement, Klabunde, and Parshall (233) who examined the reactions of NO with various q3-allyl complexes. In particular these investigators found that the bridged binuclear complex [Ni(q3-C,H,)Br], reacts with NO to form a Ni complex of acrolein oxime. Prior coordination of the nitric oxide and bending of the nitrosyl ligand are thought to occur before coupling to give the acrolein oxime precursor (233). It now seems clear that a systematic exploration of the relative propensities of (107) and (108) for the formation of C-N bonds is needed. Recently Schoonover et al. (23.31) have reported the mixed allyl nitrosyl complex Ir(NO)(C,H,)(PPh,)2t which exhibits solution behavior showing a facile transformation between the linear and bent bonding modes (37) and (38), and which decomposes rapidly under CO to yield acrolein oxime, presumably by the intramolecular coupling shown in (1 12). A recent presentation of McCleverty’s work (17Za) offers a differing view for the role of


157

CO in (109) since the reaction also proceeds in the absence of CO, but the formation of the organic products is thought to begin with nucleophilic attack on benzyl bromide by a coordinated NO group. McCleverty's recent review ( 1 7 1 ~also ) gives an extensive discussion of the reaction chemistry of coordinated NO. Two recent papers by Schug and Guengerich (233b,c) present additional examples of the electrophilic behavior of nitrosyl in [RU(NO)(NH,),]~+. ) the conversion of NO+ to an organic nitrile via a One of these ( 2 3 3 ~shows sequence similar to (104)-(106). Additional studies of the reactions of coordinated NO can be found in other reviews, but the basic reactivity patterns outlined in this section can serve to develop a comprehension for most of the nitrosyl chemistry which has been reported to date.

c. CATALYZED REDUCTION OF NITRICOXIDE The reactions of nitric oxide involve either oxidation or reduction (or both simultaneously in disproportionation and decomposition). Except for oxidation, these reactions of nitric oxide require catalysts for them to proceed at significant rates. An important stimulus to studying these catalyzed reactions lies in the environmental hazards posed by oxides of nitrogen, of which NO is the parent member, as discussed at the beginning of this section. Particular attention in the area of metal complex catalyzed reactions has focused in the last five years on the reduction of nitric oxide by carbon monoxide, (113). 2N0

+ CO

--f

N2O

f

CO2

/113)

This reaction represents a conversion of two noxious combustion effluents into relatively innocuous and inert products, and at the same time, possesses great intrinsic interest because of the extensive bond reorganizations that take place including the transfer of an oxygen atom. Several different homogeneous catalyst systems for the reduction of NO by CO have been described to date (183-185, 187, 189), and in all cases, the reduction follows (1 13) with nitrous oxide as the reduced N-containing product, this despite the fact that reduction to N, is more favored thermodynamically. The reason for adherence to the stoichiometry of (1 13), rather than (1 14), for example, may relate to the fact that N,O, once formed, is a very poor ligand. NO

+ CO

+ C02

( 1 14) In addition, N,O is a relatively inert product, being reduced to N, only with difficulty under mild thermal conditions. For the different homogeneous --t

jN,

158


catalyst systems reported for (1 13), there appear to be different reaction mechanisms operative. We now outline the available evidence and basic features for each of the catalyst systems. The best understood and most thoroughly investigated catalyst system is that developed in our laboratory at Rochester (183, 187, 188, 234). The catalyst system consists of a rhodium complex generated in situ in an ethanol solution containing water and hydrochloric acid. Although initial studies establishing the catalysis used RhCl, . xH,O, the observation of an induction period and the isolation of a Rh(1) complex after catalysis led to the determination that [RhCl,(CO)J was the best precursor to the active rhodium species in this catalysis (187).Under a mixture of NO and CO gases, a yellow solution of [RhCl,(CO),]- , whether introduced as an authentic salt or generated in situ, is transformed within an hour at room temperature into an olive-green solution of the active rhodium species. The nature of the active catalyst has been probed spectroscopically (188). The visible/UV spectrum of this species was monitored over several hours under a CO atmosphere, during which time a peak at 332.5 nm characteristic of [RhCl,(CO),]- grew in, and an isosbestic point at 360 nm was observed. The isosbestic point in this series of spectra indicates that only one rhodium species is present to significant amounts in the green solutions of the active catalyst. Through the use of a flow technique, the infrared spectrum of the active catalyst was measured, and found to contain vco at 2095 cm-' and vNo at 1715 and 1680 cm-'. These bands indicate that the single rhodium species in solution is a dinitrosyl carbonyl complex. The presence of two nitrosyls suggests a diamagnetic formulation for the active catalyst, which is supported by the absence of any esr signal from the catalyst solution at low temperature. From the high energy position of the carbonyl stretch, we are able to assign the active catalyst species as a complex of Rh(III), and from the low values of vNo, we can surmise that both nitrosyls coordinate in a bent manner as NO- (188). The active catalyst is thus formulated as (42) in which S is a solvent molecule which may or may not be coordinated in this Rh(1II) species. 0 0-

c c1\ I

II

/N. C,/Rh\N.

L

0 (42)

The active catalyst is maintained only under NO and CO gas mixtures. Under CO alone the carbonyl catalyst precursor [RhCl,(CO),]- is reformed, and under NO alone a nitrosyl complex is formed which also functions as a


159

precursor to the active catalyst. This nitrosyl complex, originally formulated (187) as [RhCl,(NO),]-, now appears to be of a more complex formulation (188).

The mechanism of the catalysis of the CO reduction of NO via (113) using this rhodium system has been investigated by kinetic studies (187) and by an isotope labeling experiment (234). Scheme I presents the mechanism as currently viewed. From initial rate studies it was established that the reaction is first order in total rhodium concentration, indicating that N-N bond formation occurs intramolecularly. The rate also appears to be first order in the partial pressure of CO and zeroth order in the partial pressure of NO over the range 200-400 Torr. We conclude from these results that the rate determining step is attack of CO on the catalytically active species (42) (187). Both H+ and H,O are found to be cocatalysts in this system. The reaction does not proceed in strictly anhydrous solutions, and the rate is markedly enhanced when HC1 is added to an aqueous ethanolic solution of the active catalyst. Our proposal to account for these effects postulates that water acts as the oxygen transfer agent in this catalysis, being consumed in the production of CO, and regenerated in the formation of N,O. Conversely in this proposal, protons are consumed in N,O formation and liberated in the production of CO,. Thus the coupling of the two NO- ligands in this scheme is viewed as occurring with protonation leading to the formation of a hyponitrous acid (H,N,O,) or hyponitrite ligand species which then decomposes to N,O and H 2 0 . The role of water as the oxygen transfer agent in this catalysis was probed directly by an isotope labeling experiment in which 180-labeled water was employed as the co-catalyst (234). If the oxygen in the product CO, did indeed come from water as postulated, and not directly from NO, then the product CO, in this experiment should show 80-incorporation. In order to minimize scrambling between the product CO, and the water in solution, that proceeds through the intermediate formation of carbonic acid, a flow reactor was employed in which product gases were swept from the catalyst solution, and trapped and isolated downstream. With this experimental set-up, the most rigorous control run showed only 2%exchange of the oxygen atoms in CO, with labeled water in solution (234). Mass spectrometric analysis of the product CO, confirmed substantial '80-incorporation, which was viewed as occurring via (115), and thus provided concrete evidence for the role of water in this catalysis.

Rh"'--CO

+ H2B-+

Rh"'-C/

/ o + H+ -+ Rh' + COB + H+ ' B H

(1 15)

160


[Equation (115) is the same as (6) studied by James and co-workers (62) in the CO reduction of RhCl,.] The labeling experiment also revealed information on the stability of the hydroxycarbonyl intermediate in (115). If this species, Rh-COOH, was formed in an equilibrium concentration, then proton transfer and the reverse reaction would lead to incorporation of labeled oxygen in the carbonyl ligand and therefore to the observation of doubly labeled CO, . However, comparison of the abundances of the three isotopic carbon dioxide molecules found (masses 44, 46 and 48) with distributions calculated assuming (i) equilibrium formation of the hydroxycarbonyl and (ii) immediate decomposition of the intermediate clearly showed that the hydroxycarbonyl intermediate reacts to form CO, immediately after it is formed, with no indication of a substantial equilibrium or incorporation of l 8 O in the carbonyl ligand. Having established the role of water in this catalytic cycle, and by inference the role of protons also, it is instructive to view the overall reaction as the sum of the two redox reactions : CO 2N0

+ H 2 0 + C02 + 2H+ + 2e+ 2Ht + 2e- -+ N20 + H20

2N0

+ CO + NZO + COZ

and if one introduces the Rh(I)/Rh(III) couple, these may be rewritten as Rh"' -CO Rh'

+ 2N0

2N0

--

+ H20 -Ht- Rh"' (COOH) Rh"' (NO-),

+ CO

catalyst complex

__+

+ CO, + Ht + NzO + HzO

Rh'

Rh"'

N,O

+ COz

While water acts as the oxygen transfer agent in this catalysis, the electrons are transferred via the rhodium center. A different homogeneous catalyst system for the carbon monoxide reductition of NO is based on transition metal dinitrosyl complexes and has been investigated by both Johnson and co-workers (184,186,235),and Ibers and co-workers (185, 236). In 1973 Johnson and Bhaduri (184) reported the separate reactions (1 16) and (1 17) which appeared to presage catalysis of NO reduction via (1 13). Shortly thereafter, Haymore and Ibers (185) showed that Ir(NO),L, + and related dinitrosyl Ir(NO),(PPh3)2+ + 4CO -+Ir(CO),(PPh3),+

+

Ir(C0)3(PPh3)2+ 2 N 0

+ N,O + CO,

+

Ir(N0)2(PPh3)2+ 3CO

(1 16) (1 17)


161

Po (42 1

[Rhc1,(co)l-

+ CO, + H,O + H+ N,O

+

Scheme I. Catalytic cycle for the reduction of N O by CO based on [RhCl,(CO),]- in aqueous acidic ethanol.

complexes did indeed catalyze (1 13) under mixtures of NO and CO although the rates of catalysis were slow. These investigators also proposed that an important intermediate in this catalysis is a saturated complex in which the two nitrosyl ligands of the starting complex have coupled to form a cisdinitrogen dioxide moiety as in structure (43).

(43)

This structure should be contrasted with an alternative formulation of this species as a complex containing one linear and one bent nitrosyl which is suggested in (1 10) for the Ru analog. The dinitrogen dioxide ligand differs

162


from hyponitrite in terms of its charge and electron density distribution, and the implication in invoking it in this catalysis is that the coupling of the nitrosyls precedes charge transfer from the metal center to the two NO ligands. Johnson el al. (186)have presented a series of isotope labeling experiments using I5NO in support of the M(N202)proposal, but their results can be interpreted as indicating simply that coordinated NO exchanges rapidly in this system when CO is present. More recently this same group of investigators (235) has determined the structure of a complex long thought to be a model of the metal-coordinated dinitrogen dioxide intermediate. However, the structure analysis shows “Pt(N,O,)(PPh,),” to be in reality the cishyponitrite complex in which the N202- ligand is chelated to a Pt(I1) center through the oxygen atoms, (44). Complex (44)reacts under CO to form N 2 0 and C 0 2 . Bhaduri and Johnson discuss the mechanism of the Ir(NO),(PPh,),’ catalyzed reduction of NO by CO, (1 16) and (1 17), in a recent full paper (235a).

(4)

In what is the most extensive kinetic study of catalysis by the dinitrosyl complexes to date, Kaduk and Ibers (236) have determined a rate law for the system [Rh(NO),L,]+ in DMF, and have examined catalyst solutions using 31P nmr spectroscopy. These investigators find that [Rh(NO),L,] serves only as the catalyst precursor, and is not a major species present under catalytic conditions. After the addition of NO, at least three different species form which are proposed as isomers of Rh(NO),L2+ (236). The observation of free triphenylphosphine may also suggest the existence of dimeric species, but it is clear that further characterization of these catalyst solutions is necessary. The rate law determined by Kaduk and Ibers (236) is: rate = k [Rh]’I2It is also found that water accelerates the rate of the reaction while addition of PPh, suppresses the rate. The reaction rate is independent of the pressure of CO and is unaffected by the addition of acid as HPF6. The authors consider the suppression of the rate by PPh, to be due to an equilibrium formation of the inactive [Rh(NO),(PPh,),]+, and propose the dependence on water to be only a solvent effect. The observed dependences on the total rhodium concentration and the pressure of NO may at first glance seem curious, and indeed the authors state (236) that the exact functional forms of these dependences are not unambiguously established since the kinetic studies were run over a limited range of conditions. A possible explanation for the observed fractional dependences can be developed, +

BINDING AND ACTIVATION OF CO, CO2, AND NO

163

however, if one assumes the existence of a dimer under NO that dissociates to a catalytically active monomer Rh* with concomitant loss of NO, (1 18). Rh,(NO)

K + 2Rh* + NO

(118)

From (1 18) one obtains [Rh*] =

K1/2[Rh2(NO)]1/2 [NO]'f2

If the rate is proportional to [Rh*],then the observed dependences are obtained ifthe concentration of the precursor dimer is proportional to the total rhodium concentration, an assumption that may be consistent with the 31P nmr observations. In any case, catalysis of NO reduction by dinitrosyl bis-phosphine complexes is less well understood than the previously described catalyst system based on [RhCl,(CO),]- in aqueous acidic ethanol. No concrete evidence to support or refute the postulated M(N202) intermediate exists, and the manner in which CO,is formed is undefined. The study by Kaduk and Ibers suggests that it will be very difficult to identify both the active metal species and the precise mechanism of catalysis. A third homogeneous catalyst system for the reduction of nitric oxide by carbon monoxide has recently been developed. Kubota et al. (189) find that an aqueous acidic solution of PdCl, and CuCl, or CuCl under NO and CO produces various yields of N,O and CO, depending on the relative concentrations of the two metals. Typical conditions are 0.01 M PdCl,, 0.20 M CuCl,, and 2 M HCl under 1 atm total of NO and CO. Formation of CO, is achieved via (1 19), and is fast relative to N 2 0 production. CO

+ Pd2+ + HzO ---t Pd(0) + CO, + 2H+

(1 19) Several processes appear to account for the reduction of NO to N,O. This reaction proceeds slowly in the presence of either Pd(0) or Cu(1) alone, but is much faster when both metals are present, leading the authors (189) to propose a Pd(I1)-Cu(1) chloride-bridged species as the most active reductant. Significantly, N 2 0 and CO, are produced in 1 : 1 proportions when CuCl is used in place of CuCl,, and the rate of the reaction increases as the initial concentration of CuCl is increased from to 1.0 M . The three catalyst systems discussed in this section for the reduction of NO by CO underscore the mechanistic complexity of a reaction which is stoichiometrically simple. Extensive bond reorganization is required in the reduction via (1 13), and each of the catalyst systems appears to proceed by a different mechanism. While two of the three systems possess a common feature in terms of CO, formation, each appears to be different with respect to N 2 0 production. The systematic development of new homogeneous

164


catalyst Systems for NO reduction in the future will undoubtedly be coupled to our growing knowledge of nitric oxide coordination chemistry and the influence of coordination on nitric oxide reactivity. ACKNOWLEDGMENTS We wish to acknowledge the many useful and informative discussions with colleagues which have been held in the course of preparing this review and to thank Profs. J. E. Bercaw, J. A. Ibers, J. A. Labinger and M. Kubota for permitting us to see manuscripts prior to publication. We also wish to acknowledge the experimental efforts of Dr. Carol Meyer and Mr. Chien-Hong Cheng for their work at Rochester on nitric oxide reduction and the water gas shift reaction respectively, and the financial support of the National Science Foundation. R. E. gratefully acknowledges the John S . Guggenheim Foundation for a fellowship. Finally, we wish to give special thanks to Ms. Lorraine Bianco without whose help this manuscript would not have been brought successfully to publication.

REFERENCES

I. General or representative references are given when a broad subject area is cited or discussed. 2. “Homogeneous Catalysis,” Adv. Chem. Ser. No. 70 Chapters 1, 2, 4-7, and 10. ACS, Washiggton, D.C., 1968; D. Forster and J. F. Roth, eds., “Homogeneous Catalysis. 11,” Adv. Chem. Ser. No. 132, and references therein. ACS, Washington, D.C., 1974. 3. B. R. James, “Homogeneous Hydrogenation.” Wiley, New York, 1973, and references therein. 4. M. M. Taqui Khan and A. E. Martell, “Homogeneous Catalysis by Metal Complexes.” Academic Press, New York, 1974. 5. J. Halpern, Acc. Chem. Res. 3,386 (1970); Annu. Rev. Phys. Chem. 16, 103 (1965). 6. J. P. Collman, Acc. Chem. Res. 1, 136 (1968); J. P. Collman and W. R. Roper, Ado. Organomet. Chem. 7,53 (1968). 7. C. A. Tolman, Chem. SOC.Rev. 1, 337 (1972). 8. A. J. Deeming, “M.T.P. International Review of Science, Inorganic Chemistry,” (M. L. Tobe, Ed.), Series 11, Vol. 9, p. 271. Buttersworth, London, 1974. 9. P. M. Henry, Adv. Organomet. Chem. 13, 363 (1975). 10. P. M. Henry, 1.Am. Chem. Soc., 94,4437 (1972), and references therein. 11. G. Yagupsky, C. K. Brown, and G. Wilkinson, J. Chem. SOC.A p. 1392 (1970); p. 2753 (1970), and references therein. 12. J. F. Roth, J. H. Craddock, A. Hershman, and F. E. Paulik, Chem. Tech. p. 600 (1971). 13. D. Forster, J. Am. Chem. SOC.98, 846 (1976). 14. B. Chiswell and D. W. James, “Fundamental Aspects of Inorganic Chemistry.” Wiley, New York, 1969. 15, T. Moeller, “Inorganic Chemistry.” Wiley, New York, 1952. 16. R. G. Pearson, “Symmetry Rules for Chemical Reactions.” Wiley (Interscience), New York, 1976. 17. P. B. Chockand J. Halpern, J. Am. Chem. SOC.88,3511 (1966). 18. J. S. Bradley, D. E. Connor, D. Dolphin, J. A. Labinger, and J. A. Osborn, J. Am. Chem. SOC.94,4043 (1972). 19. L. Vaska, Acc. Chem. Res. 4,335 (1971).

BINDING AND ACTIVATION OF CO, CO2, AND NO

165

A. Wojcicki, Adv. Organomet. Chem. 11,87 (1973). A. Wojcicki, Adv. Organomet. Chem. 12, 32 (1974). R. Cramer, Acc. Chem. Rex 1, 186 (1968). F. Calderazzo, Angew. Chem., Int. Ed. Engl. 16,299 (1977). J. K. Stille et al., J. Am. Chem. SOC.96, 1508, 1514, and 1518 (1974); C. P. Casey, C. A. Bunnell, and J. C. Calabrese, ibid. 98, 1166 (1976). 24. J. X. McDermott, J. F. White, and G. M. Whitesides, J. Am. Chem. SOC.95, 4451 (1973); 98,6521 (1976). 25. J. Schwartz and J. B. Cannon, J. Am. Chem. SOC. 96,2276 (1974), and references therein. 26. L. Mond, C. Langer, and F. Quinke, J . Chem. SOC.57,749 (1890). 27. F. Calderazzo, R. Ercoli, and G. Natta, in “Organic Syntheses via Metal Carbonyls” (I. Wender and P. Pino, eds.), p. 1. Wiley (Interscience), New York, 1968. 28. E. W. Abel and F. G. A. Stone, Q. Rev., Chem. SOC.23,325 (1969); 24,498 (1970). 29. F. A. Cotton and G. Wilkinson, “Advanced Inorganic Chemistry,” 3rd ed. Wiley (Interscience), New York, 1972. 30. B. H. Robinson, “M.T.P. International Review of Science, Inorganic Chemistry,” (M. .I. Mays, Ed.), Series 11, Vol. 6, p. 1. Buttersworth, London, 1975. 31. J. C. Kotz and K. F. Purcell, “Inorganic Chemistry.” Saunders, Philadelphia, Pennsylvania, 1977. 32. See Advances in Organometallic Chemistry for more specialized reviews. Often a tabulation of the organometallic chemistry literature for the preceding year is given [M. I. Bruce, Adv. Organomet. Chem. 11,448 (1973); 12,380 (1974)]. 33. F. A. Cotton and C. S. Kraihanzel, J. Am. Chem. SOC.84,4432 (1962); Inorg. Chem. 2, 533 (1963); F. A. Cotton, ibid. 3,702 (1964). 34. A. Streitwieser, Jr. and P. H. Owens, “Orbital and Electron Density Diagrams.” Macmillan, New York, 1973; L. C. Snyder and H. Basch, “Molecular Wave Functions and Properties.” Wiley (Interscience), New York, 1972. 35. N. A. Beach and H. B. Gray,,J. Am. Chem. SOC. 90,5713 (1968); R. F. Fenske and R. L. DeKock, Inorg. Chem. 9, 1053 (1970). 36. G. G. Sumner, H. P.Klug, and L. E. Alexander, Acta Crystallogr. 17,732 (1964). 37. F. A. Cotton, Prog. Inorg. Chem. 21, 1 (1976). 38. P. S. Braterman, Struct. Bonding (Berlin) 10, 7 (1972). 39. F. A. Cotton, L. Kruczynski, and A. J. White, J. Am. Chem. SOC. 94,6191 (1972); J. Evans, B. F. G. Johnson, J. Lewis, J. R. Norton, and F. A. Cotton, Chem. Commun. p. 807 (1973). 40. F. A. Cotton and D. L. Hunter, Inorg. Chim. Acta 11, L9 (1974); D. A. Gansow, A. R. Burke, and W. D. Vernon, J. Am. Chem. SOC.94,2550 (1972). 41. J. G. Bullit, F. A. Cotton, and T. J. Marks, Inorg. Chem. 11,671 (1972); R. D. Adams and F. A. Cotton, Inorg. Chim. Acta 7, 153 (1973). 42. E. L. Muetterties, Bull. SOC.Chim. Belg. 85, 451 (1976). 43. R. Colton, C. J. Commons, and B. F. Hoskins, Chem. Commun. p. 363 (1975). 44. M. G. Newton, R. B. King, M. Chang, and J. Gimeno, J. Am. Chem. SOC.99,2802(1977). 45. M. P. Brown, R. J. Puddephat, M. Rashidi, Lj. ManojloviC-Muir, K. W. Muir,T. Solomun, and K. R. Seddon, Inorg. Chim. Acta 23, L33 (1977). 46. M. M. Olmstead, H. Hope, L. S. Benner, and A. L. Balch, J. Am. Chem. SOC.99, 5502 (1977). 46a. R. Colton, M. J. McCormick, and C. D. Pannan, Chem. Commun. p. 823 (1977). 466. R.Colton, M. J. McCormick, and C. D. Pannan, Aust. J. Chem. 31, 1425 (1978). 46c. M. Cowie, J. T. Mague, and A. R. Sanger, J. Am. Chem. SOC.100,3628 (1978). 46d. C. P. Kubiak and R. Eisenberg, J. Am. Chem. SOC.99,6129 (1977). 46e. C. P. Kubiak and R. Eisenberg, unpublished results.

20. 2Ua. 21. 22. 23.

166


47. D. F. Shriver, J. Orgunomet. Chem. 94,259 (1975); Chem. Br. 8,419 (1972). 48. A. E. Crease and P. Legzdins, J. Chem. Educ. 52,499 (1975). 49. J. S. Kristoff and D. F. Shriver, Inorg. Chem. 13,499 (1974), and references therein. 50. N. E. Kim, N. J. Nelson, and D. F. Shriver, Inorg. Chim. Acru 7,393 (1973); N. J. Nelson, N. E. Kime, and D. F. Shriver, J. Am. Chem. SOC.91,5173 (1969). 51. Sr. Agnes Alich, N. J. Nelson, and D. F. Shriver, Chem. Commun. p, 254 (1971). 52. J. M. Burlitch and R. B. Petersen, J. Orgunomet. Chem. 24, C65 (1970). 53. J. C. Kotz and C. D. Turnipseed, Chem. Commun. p. 41 (1970). 54. R. B. Petersen, J. J. Stezowski, C. Wan, J. M. Burlitch, and R. E. Hughes, J. Am. Chem. SOC. 93,3532 (1971). 55. W. Hieber and F. Leutert, Z . Anorg. Allg. Chem. 204, 145 (1932); J. W. Reppe, Justus Liebig’s Ann. Chem. 582, 121 (1953). 56. E. 0.Fischer and A. Maasbol, Chem. Ber. 100,2445 (1967); Angew. Chem. 76,645 (1964); Angew. Chem., Int. Ed. Engl. 3,580 (1964). 57. E. 0. Fischer, Rev. Pure Appl. Chem. 24,407 (1970); Adv. Organomet. Chem. 1 4 , l (1976). 58. C. P. Casey and S. M. Neumann, J. Am. Chem. SOC.98,5395 (1976). 59. J. P. Collman and S. R. Winter, J. Am. Chem. SOC.95,4089 (1973). 59u. T. J. Collins and W. R. Roper, Chem. Commun. p. 1044 (1976). 59b. C. P. Casey, M. A. Andrews, and J. E. Ring, J. Am. Chem. SOC.101,142 (1979). 59c. W. Tam, W. K. Wong, and J. A. Gladysz, J. Am. Chem. SOC.101, 1589 (1979). 59d. J. A. Gladysz and J. H. Merrifield, Inorg. Chim. Acfu 30, L317 (1978). 59e. C. P. Casey, M. A. Andrews, and D. R. McAlister, J. Am. Chem. SOC. 101,3371 (1979). 60. M. Y. Darensbourg, H. L. Conder, D. J. Darensbourg, and C. Hasday, J. Am. Chem. SOC. 95, 5919 (1973), and references therein. 61. W. Reppe, Justus Liebigs Ann. Chem. 582,116 (1953). 62. B. R. James, G. L. Rempel, and F. T. T. Ng, J . Chem. SOC.A p. 2454 (1969). 62u. H. A. Hodali, D. F. Shriver, and C. A. Ammlung, J. Am. Chem. SOC.100,5239 (1978). 63. J. E. Ellis, R. A. Faltynek, and S. G. Hentges, J. Am. Chem. SOC.99,626 (1977). 64. R. F. Heck, Adu. Cutul. 26, 323 (1977). 65. G . P. Chiusoli, Acc. Chem. Res. 6,422 (1973). 66. “The National Energy Plan.” US.Govt. Printing Office, Washington, D.C., 1977. 67. I. Wender, Curul. Rev.-Sci. Eng. 14,97 (1976). 68. M. A. Vannice, Curd Rev.-Sci. Eng. 14, 153 (1976). 69. “Report of the Workshop on Fundamental Research in Homogeneous Catalysis as Related to U.S. Energy Problems.” Stanford University, Stanford, California, 1974. 70. Comprehensive discussions of CO/H2 reactions, methanation and methanol synthesis are given in “Encyclopedia of Chemical Technology” (R. E. Kirk and D. F. Othmer, eds.), 2nd ed. Wiley (Interscience), New York, 1963-1970. For CO/H2 reactions, see Vol. 4, p. 446; for methanation, see Vol. 13, p. 364; for methanol synthesis, see Vol. 13, p. 370. 71. G . A. Mills and F. W. Steffgen, Caful. Rev. 8, 159 (1973). 72. P. Sabatier and J. B. Senderens, C. R. Hebd. Seances Acud. Sci. 134,514 (1902). 73. F. Fischer and H. Tropsch, Chem. Ber. 56,2428 (1923); 59,830,832, and 923 (1926). 74. H. H. Storch, H. Golumbic, and R. B. Anderson, “The Fischer-Tropsch and Related Syntheses.” Wiley, New York, 1951. 75. Pilot plant and demonstration plant scale projects include the SYNTHANE process, the Hygas process, and the Bigas process among others. See W. P. Haynes and A. J. Forney, “High BTU Gas from Coal: Status and Prospects,” Tech. Rep. PERC/IC-76-1. U.S.E.R. D.A. Office of Public Affairs, 1976. 76. H. Pichler and H. Schulz, Chem.-Ing.-Tech.42, 1162 (1970). 77. M. A. Vannice, J. Cutul. 37,449 and 462 (1975).


167

78. J. M. Manriquez, D. R. McAlister, R. D. Sanner, and J. E. Bercaw, J. Am. Chem. Soc. 98,6733 (1976). 79. J. E. Bercaw, J. Am. Chem. SOC.96,5087 (1974). 80. G. Fachinetti, C. Floriani, F. Morchetti, and S. Merlino, Chem. Commun. p. 522 (1976). 81. L. I. Shoer and J. Schwartz, J. Am. Chem. SOC.99,5831 (1977). 82. R. L. Pruett and W. E. Walker, U. S. Patent 3,833,634 (1974); Ger. Offen. Patent 2,262,318 (1973); W. E. Walker and R. L. Pruett, Ger. Offen. Patent 2,426,411 (1975). 83. J. C. Huffman, J. G. Stone, W. C. Krusell, and K. G. Caulton, J. Am. Chem. SOC.99, 5829 (1977). 84. P. M. Treichel and R. L. Shubkin, Inorg. Chem. 6,1328 (1967). 85. R. P. Stewart, N. Okamoto, and W. A. G. Graham, J. Organomet. Chem. 42, C32 (1972). 86. A. N. Nesmeyanov, K. N. Anisimov, N. E. Kolobova, and L. L. Krasnoslobodskaya, Izv. Akad. Nauk. SSSR, Ser. Khim. p. 860 (1970). 87. G. C. Demitras and E. L. Muetterties, J. Am. Chem. SOC.99,2796 (1977). 88. R. L. Pruett, presented at Workshop in Organometallic Chemistry, University of Rochester, Rochester, New York, 1977; Ann. N . Y. Acad. Sci., 295, 239 (1977). 88a. J. M. Manriquez, D. R. McAlister, R. D. Sanner, and J. E. Bercaw, J. Am. Chem. SOC.100, 2716 (1978). 886. P. T. Wolczanski, R. S. Threlkel, and J. E. Bercaw, J. Am. Chem. SOC.101,218 (1979). 88c. J. A. Labinger, K. S. Wong, and W. R. Scheidt, J. Am. Chem. SOC.100,3254 (1978). 88d. J. M. Manriquez, P. J. Fagan, and T. J. Marks, J. Am. Chem. SOC.100,3939 (1978). 88e. J. M. Manriquez, P. J. Fagan, T. J. Marks, C. S. Day, and V. W. Day, J. Am. Chem. SOC. 100,7112 (1978). 88J K. L. Brown, G. R. Clark, C. E. L. Headford, K. Marsden, and W. R. Roper, J. Am. Chem. SOC.101,503 (1979). 889. C. Masters, C. van der Woude, and J. A. van Doom, J. Am. Chem. SOC.101,1633 (1979). 88h. J. W. Rathke and H. M. Feder, J. Am. Chem. SOC. 100,3623 (1978). 88i. G. Henrici-Olive and S. Olive, Angew. Chem. Int. Ed. Engl. 18, 77 (1979). 88j. A. K. Smith, A. Theolier, J. M. Basset, R. Ugo, D. Commereuc, and Y. Chauvin, J. Am. Chem. SOC.100,2590 (1978). 89. “Catalyst Handbook,” Chapters 5 and 6. Springer-Verlag, Berlin and New York, 1970. 90. M. Tachikawa, G. F. Stuntz, and J. R. Shapley, unpublished results, as quoted in R. J. Lawson and J. R. Shapley, J. Am. Chem. SOC.98,7433 (1976). For other Me,NO reactions with metal carbonyls, see H. Alper and J. T. Edward, Can. J. Chem. 48, 1543 (1970); J. Elzinga and H. Hogeveen, Chem. Commun. p, 705 (1977). 90a. R. D. Feltham and J. C. Kriege, J. Am. Chem. SOC. 101,5064 (1979). 91. G. D. Mercer, W. B. Beaulieu, and D. M. Roundhill, J. Am. Chem. SOC.99,6551 (1977). 92. B. L. Haymore and J. A. Ibers, J. Am. Chem. SOC.96,3325 (1974). 93. B. F. G. Johnson and S. Bhaduri, Chem. Commun. p. 650 (1973). 94. J. E. Bercaw, L.-Y. Goh, and J. Halpern, J. Am. Chem. SOC.94,6534 (1972). 95. H. C. Clark and W. J. Jacobs, Inorg. Chem. 9, 1229 (1970). 96. A. J. Deeming and B. L. Shaw, J. Chem. SOC.A p. 443 (1969). 97. T. G. Appleton and M. A. Bennett, J. Organomet. Chem. 55, C88 (1973). 97a. C. P. Casey, M. A. Andrews, and J. E. Rinz, J. Am. Chem. SOC.101, 741 (1979). 976. N. Grice, S. C. Kao, and R. Pettit, J. Am. Chem. SOC.101, 1627 (1979). 98. R. M. Laine, R. G. Rinker, and P. C. Ford, J. Am. Chem. SOC.99,252 (1977). 99. C.-H. Cheng, D. E. Hendriksen, and R. Eisenberg, J. Am. Chem. SOC.99,2791 (1977). 99a. R. Eisenberg and C.-H. Cheng, U.S.Patent 4,107,076 (1978). 100. D. E. Hendriksen and R. Eisenberg, unpublished results. 101. D. Forster, Inorg. Chem. 8,2556 (1969).

168


102. J. J. Daly, F. Sanz, and D. Forster, J. Am. Chem. SOC.97,2551 (1975).

103. B. P. Curtis, Jr., L. W. Fanin, F. E. Paulik, and J. L. Price, Ger. Offen. Patent 2,211,231 (1973); Chem. Abstr. 78, 138415f (1973). 103a. E. C. Baker, D. E. Hendriksen, and R. Eisenberg, J. Am. Chem. SOC.,102 (1980). 103b. T. C . Singleton, L. J. Park, J. L. Price, and D. Forster, ACS Pefroleum Division Preprints 24,329 (1979). 104. R. M. Laine, R. G. Rinker, and P. C. Ford, 173rd Meet., Am. Chem. SOC.,1977 Paper INOR-99 (1977); P. C. Ford, personal communication. 104u. P. C. Ford, R. G. Rinker, C. Ungermann, R. M. Laine, V. Landis, and S. A. Moya, J. Am. Chem. SOC.100,4595 (1978). 105. D. M. Fenton, Union Oil Co., U. S. Patent 3,539,298 (1970). 106. C.-H. Cheng and R. Eisenberg, J. Am. Chem. SOC.100,5968 (1978). 107. J. V. Kingston and G. R. Scollary, J. Chem. SOC.A p. 3765 (1971). 107a. H. C. Kang, C. H. Mauldin, T. Cole, W. Slegeir, K. Cam, and R. Pettit, J. Am. Chem. SOC. 99,8323 (1977). 1076. R. B. King, C. C. Frazier, R. M. Hanes, and A. D. King, Jr., J. Am. Chem. SOC.100,2925 (1978). 1 0 7 ~ T. . Yoshida, Y. Ueda, and S. Otsuka, J. Am. Chem. SOC.100,3941 (1978). 107d. R. M. Laine, J. Am. Chem. SOC.100,6451 (1978). 107e. Y. Watanabe, K. Takatsuki, and Y. Takegami, Tefrahedron Lett. 3369 (1978). 107J K. Cann, T. Cole, W. Slegeir, and R. Pettit, J. Am. Chem. SOC.100, 3969 (1978). 1079. R. M. Laine, D. W. Thomas, L. W. Cary, and S. E. Buttrill, J. Am. Chem. SOC.100,6527 (1978). 108. M. E. Volpin and I. S. Kolomnikov, Organomet. React. 5,313-386 (1975), and references therein. 109. D. W. Krogmann, “The Biochemistry of Green Plants.” Prentice-Hall, Englewood Cliffs, New Jersey, 1973. 110. J. A. Bassham and M. Calvin, “The Path of Carbon in Photosynthesis.” Prentice-Hall, Englewood Cliffs, New Jersey, 1957. I l l . E. Rabinowitch and Govindjee, “Photosynthesis.” Wiley, New York, 1969. 112. M. E. Volpin and I. S. Kolomnikov, Pure Appl. Chem. 33,567 (1973). 1120. I. S. Kolomnikov and M. Kh. Grigoryan, Russian Chemical Reviews 47, 334 (1978); translated from Uspekhi Khimii 47,603 (1978). 113. T. Ito and A. Yamamoto, J. SOC.Org. Synth. Chem., Tokyo 34,308 (1976); Chem. Absrr. 85, 136305 (1976); M. Ichikawa, Kyoritsu Chem. Libr. 11, 32 (1976); Chem. Abstr. 85, 10513(1976). 114. R. N. Scott, D. F. Shriver, and D. D. Lehman, Znorg. Chim. Acfu 4,73 (1970). 115. M. Aresta, C. F. Nobile, V. G. Albano, E. Forni, and M. Manassero, Chem. Commun. p. 636 (1975). 116. T. Herskovitz and L. J. Guggenberger, J. Am. Chem. SOC.98, 1615 (1976). 117. M. Aresta and C. F. Nobile, J. Chem. SOC.,Dalton Trans. p. 708 (1977). 118. R. Mason and A. I. M. Rae, J. Chem. SOC.A p. 1767 (1970). 119. T. Herskovitz, J. Am. Chem. SOC.99,2391 (1977). 119a. G. Fachinetti, C . Floriani, and P. F. Zanazzi, J. Am. Chem. SOC.100, 7405 (1978). 1196. C. Floriani and G. Fachinetti, Chem. Commun. 615 (1974). 119c. C. R. Eady, J. J. Guy, B. F. G. Johnson, J. Lewis, M. C. Malatesta, and G. M. Sheldrick, Chem. Commun. 602 (1976). 119d. G. R. John, B. F. G. Johnson, I. Lewis, and K. C. Wong, J. Organomefal.Chem. 169, C23 (1979).


169

120. B. R. Flynn and L. Vaska, Chem. Commun. p. 703 (1974). 121. H. H. Karsch, Chem. Ber. 110,2213 (1977). 122. J. Chatt, M. Kubota, G. J. Leigh, F. C. March, R. Mason, and D. J. Yarrow, Chem. Commun. p. 1033 (1974). 123. T. Ito, H. Tsuchiya, and A. Yamamoto, Chem. Lett. p. 851 (1976). I23a. R. J. Crutchley, J. Powell, R. Faggiani, and C. J. L. Lock, Inorg. Chim. Acta 24, L15 (1977). 124. A. Miyashita and A. Yamamoto, J. Organornet. Chem. 113, 187 (1976). 125. I. S. Kolomnikov, T. S. Belopotapova, T. V. Lysyak, and M. E. Vol’pin, J. Organornet. Chem. 67, C25 (1974). 126. Y. Iwashita and A. Hayata, J. Am. Chem. SOC.91,2525 (1969). 127. B. R. Flynn and L. Vaska, J. Am. Chem. SOC.95,5081 (1973). 128. M. E. Vol’pin, I. S. Kolomnikov, and T. S. Lobeeva, Izv. Akad. Nauk SSSR, Ser. Khim. p. 2084 (1969). 129. S . Komiya and A. Yamamoto, J. Organomet. Chem. 46, C58 (1972). 130. S. Krogsrud, S. Komiya, T. Ito, J. A. Ibers, and A. Yamamoto, h o r g . Chem. 15, 2798 (1976). 131. M. H . Chisholm and M. W. Extine, J. Am. Chem. SOC.99,792 (1977). 132. S. Komiya and A. Yamamoto, Bull. Chem. Soc. Jpn. 49,784 (1976). 133. I. S. Kolomnikov, T. S. Belopotapova, and M. E. Vol’pin, J. Gen. Chem. USSR (Engl. Transl.) 45, 1958 (1975). 134. T. V. Ashworth and E. Singleton, Chem. Commun. p. 204 (1976). 135. V. D. Bianco, S. Doronzo, and M. Rossi, J. Organomet. Chem. 35, 337 (1972). 136. I. S. Kolomnikov, G. Stepovska, S. Tyrlik, and M. E. Vol’pin, Zh. Obshch. Khim. 42, 1652 (1972); J. Gen. Chem. USSR (Engl. Transl.) 42, 1645 (1972). 137. A. D. English and T. Herskovitz, J. Am. Chem. SOC.99, 1648 (1977). 137a. S . D. Ittel, C. A. Tolman, A. D. English, and J. P. Jesson, J. Am. Chem. SOC.100, 7577 (1978). 138. I. S. Kolomnikov, A. 0. Gusev, T. S. Belopotapova, M. Kh. Grigoryan, T. V. Lysyak, Yu. T. Struchkov, and M. E. Vol‘pin, J. Organomef. Chem. 69, C10 (1974). 139. I. S . Kolomnikov, G. Stepovska, S. Tyrlik, and M. E. Vol’pin, J. Gen. Chem. USSR (Engl. Trans[.)44, 1710 (1974). 140. T. Tsuda, Y. Chujo, and T. Saegusa, Chem. Commun. p. 963 (1975). I40a. T. Tsuda, Y. Chujo, T. Saegusa, J . Am. Chem. SOC.100,630 (1978). 141. A. Miyashita and A. Yamamoto, J. Organomef. Chem. 49, C57 (1973). 141a. T. Yamamoto and A. Yamamoto, Chem. Lett. 615 (1978). 142. F. A. L. Anet and E. Leblanc, J. Am. Chem. SOC.79,2649 (1957). 143. M. H. Chisholm, F. A. Cotton, M. Extine, and B. R. Stults, J. Am. Chem. SOC.98, 4683 (1976). 143a. L. E. Manzer, J. Organometal. Chem. 135, C6 (1977). 144. M. H. Chisholm and M. Extine, J. Am. Chem. SOC.96,6214 (1974). 145. M. H. Chisholm and M. Extine, Chem. Commun. p. 438 (1975). 146. M. H. Chisholm and M. Extine, J. Am. Chem. SOC.97,1623 (1975). 147. M. H. Chisholm and M. W. Extine, J. Am. Chem. SOC.99,782 (1977). 148. M. H. Chisholm and M. W. Extine, J. Am. Chem. SOC.99,792 (1977). 149. C. W. Newing, Ph.D. Thesis, University of London (1971); in Chisholm and Extine (147, footnote 6). 150. A. J. Goodsel and G. Blyholder, J. Am. Chem. SOC.94,6725 (1972). 151. M. Hidai, T. Hikita, and Y. Uchida, Chem. Lett. p. 521 (1972). 152. T. Tsuda and T. Saegusa, Inorg. Chem. 11,2561 (1972).

I70


153. T. Tsuda, S. Sanada, K. Ueda, and T. Saegusa, Inorg. Chem. 15,2329 (1976). 154. M. H. Chisholm, W. W. Reichert, F. A. Cotton, and C. A. Murillo, J. Am. Chem. SOC.99,

1652 (1977). 1 5 4 ~ M. . H. Chisholm, F. A. Cotton, M. W. Extine, and W. W. Reichert, J. Am. Chem. SOC.100,

1727 (1978). 155. E. Chaffee, T. P. Dasgupta, and G. M. Harris, J. Am. Chem. SOC.95,4169 (1973). 156. D. A. Palmer and G. M. Harris, Inorg. Chem. 13,965 (1974). 157. R. J. DePasquale, Chem. Commun. p. 157 (1973). 158. T. Tsuda, Y. Chujo, and T. Saegusa, Chem. Commun. p. 415 (1976). 159. Y. Inoue, Y. Sasaki, and H. Hashimoto, Chem. Commun. p. 718 (1975). 160. Y. Inoue, H. Izumida, Y. Sasaki, and H. Hashimoto, Chem. Lett. p. 863 (1976). 1 6 0 ~ G. . 0. Evans and C. J. Newell, Inorg. Chim. Actu 31, L387 (1978). 161. H. Koinuma, H. Kato, and H. Hirai, Chem. Lett. p. 517 (1977). 1 6 1 ~ G. . Fachinetti, C. Floriani, A. Chiesi-Villa, and C. Guastini, J . Am. Chem. SOC.101, 1767

(1979). 161b. G. 0. Evans, W. F. Walter, D. R. Mills, and C. A. Streit, J. Orgunornetul. Chem. 144,C34

(1978). 162. H. Koinuma, Y. Yoshida, and H. Hirai, Chem. Letl. p. 1223 (1975). 163. B. F. G. Johnson and J. A. McCleverty, Prog. Inorg. Chem. 7,277 (1966). 164. N. G. Connelly, Inorg. Chim. Actu, Rev. 6,48 (1972). 165. W. P. Griffith, Adv. Orgunomef. Chem. 7,211 (1968). 166. J. A, Masek, Inorg. Chim. Actu, Rev. 3,99 (1969). 167. B. A. Frenz and J. A. Ibers, MTP Int. Rev. Sci.Phys. Chem., Ser. 1 11, 33 (1972). 168. J. H. Enemark and R. D. Feltham, Coord. Chem. Rev. 10,339 (1974). 169. R. Eisenberg and C. D. Meyer, Acc. Chem. Res. 8,26 (1975). 170. K. G. Caulton, Coord. Chem. Rev. 14,317 (1975). 171. F. Bottomley, Acc. Chem. Res. 11, 158 (1978). 1 7 1 ~ J. . A. McCleverty, Chem. Rev. 79,53 (1979). 172. W. Bartok, A. R. Crawford, and A. Skapp, Chem. Eng. Prog. 67,64 (1971).

173. “Air Pollution,” ACS Repr. Collect. Am. Chem. SOC.,Washington, D.C., 1973. 174. P. E. Cade, W. M. Huo, and J. B. Greenshields, At. Data Nucl. Data Tables 15, 1 (1975). 175. T. E. Green and C. N. Hinshelwood, J. Chem. SOC.p. 1709 (1926). 176. C. S. Howard and F. Daniels, J. Phys. Chem. 62,215 (1958). 177. M. Shelef, K. Otto, and H. Gandhi, Atmos. Environ. 3, 107 (1969). 178. E. R. S. Winter, J . Cutal. 22, 158 (1971). 179. K. Otto and M. Shelef, J. Phys. Chem. 76,37 (1972); M. Shelef and K. Otto, J . Cutul. 10, 408 (1968). 180. R. L. Klimisch and G. J. Barnes, Environ. Sci. Technol. 6, 543 (1972). 181. R. J. H. Voorhoeve, J. P. Remeika, and D. W. Johnson, Jr., Science 180,62 (1973). 182. M. Shelef and H. S. Gandhi, Ind. Eng. Chem., Prod. Res. Deu. 11,2 (1972). 183. J. Reed and R. Eisenberg, Science 184,568 (1974). 184. B. F. G. Johnson and S . Bhaduri, Chem. Commun. p. 650 (1973). 185. B. L. Haymore and J. A. Ibers, J. Am. Chem. SOC.96,3325 (1974). 186. S. Bhaduri, B. F. G. Johnson, C. J. Savory, J. A. Segal, and R. H. Walter, Chem. Commun. p. 809 (1974). 187. C. D. Meyer and R. Eisenberg, J. Am. Chem. SOC.98, 1364 (1976). 188. D. E. Hendriksen, C. D. Meyer, and R. Eisenberg, Inorg. Chem. 16,970 (1977). 189. M. Kubota, K. J. Evans, C. A. Koerntgen, and J. C. Marsters, Jr., J. Am. Chem. SOC.100, 342 (1978). 190. K. R. Grundy, K. R. Laing, and W. R. Roper, Chem. Commun. p. 1501 (1970).


191. 192. 193. 194.

171

J. B. Godwin and T. J. Meyer, Znorg. Chem. 10,471 and 2150 (1971). S. G. Clarkson and F. Basolo, Inorg. Chem. 12, 1528 (1973).

D. Gwost and K. G . Caulton, Inorg. Chem. 13,414 (1974). F. Bottomley, E. M. R. Kiremire, and S. G. Clarkson, J. Chem. Soc., Dalton Trans. p. 1909 (1975). 195. P. T. Manoharan and H. B. Gray, Inorg. Chem. 5,823 (1966); J. Am. Chem. SOC.87,3340 (1965). 196. N. V. Sidgwick and R. W. Bailey, Proc. R. SOC.London, Ser. A 144, 521 (1934); N. V. Sidgwick, ‘‘Chemical Elements and Their Compounds,” Vol. I, p. 685. Oxford Univ. Press (Clarendon), London and New York, 1950. 197. D. J . Hodgson and J. A. Ibers, Znorg. Chem. 7,2345 (1968); D. J. Hodgson, N. C. Payne, J. A. McGinnety, R. G. Pearson, and J. A. Ibers, J. Am. Chem. SOC.90,4486 (1968). 198. C. G. Pierpont and R. Eisenberg, J. Am. Chem. SOC.93,4905 (1971). 199. D. M. P. Mingos, Inorg. Chem. 12, 1209 (1973). 200. R. Hoffmann, M. M. L. Chan, M. Elian, A. R. Rossi, and D. M. P. Mingos, Inorg. Chem. 13, 2666 (1974). 201. J. P. Collman, N. W. Hoffman, and D. E. Morris, .I Am. . Chem. SOC.91, 5659 (1969). 202. C. G. Pierpont and R. Eisenberg, Inorg. Chem. 11, 1088 (1972); C. G. Pierpont, D. G. VanDerveer, W. Durland, and R. Eisenberg, J. Am. Chem. SOC.92,4760 (1970). 203. J. P. Collman, P. Farnham, and G. Dolcetti, J . Am. Chem. SOC.93, 1788 (1971). 204. J. H. Enemark and R. D. Feltham, Proc. Nail. Acad. Sci. U.S.A. 69,3534 (1972). 205. J. H. Enemark, R. D. Feltham, J. Riker-Nappier, and K. F. Bizot, Znorg. Chem. 14,624 (1975). 206. D. M. P. Mingos and J. A. Ibers, Inorg. Chem. 9, 1105 (1970). 207. T. E. Nappier, R. D. Feltham, J. H. Enemark, A. Kruse, and M. Cooke, Inorg. Chem. 14, 806 (1975). 208. J. H. Enemark, R. D. Feltham, B. T. H i e , P. L. Johnson, and K. B. Swedo, J . Am. Chem. SOC.99,3285 (1977). 209. J. A. McCleverty, N. M. Atherton, J. Locke, E. J. Wharton, and C. J. Winscom, J. Am. Chem. SOC.89,6082 (1967). 210. J. A. Kaduk and J. A. Ibers, Inorg. Chem. 14, 3070 (1975). 211. R. Bau, I. H. Saberwhal, and A. Burg, J. Am. Chem. Soc. 93,4926 (1971). 212. J. R. Norton, J. P. Collman, G . Dolcetti, and W. T. Robinson, Inorg. Chem. 11,382 (1972). 213. J. L. Calderon, F. A. Cotton, B. G. DeBoer, and N. Martinez, Chem. Commun. p. 1476 (1971). 214. W. C. Trogler and L. G. Marzilli, Znorg. Chem. 13, 1008 (1974). 215. M. Kubota and D. A. Phillips, J. Am. Chem. SOC.97,5637 (1975). 216. G. Dolcetti, N. W. Hoffman, and J. P. Collman, Inorg. Chim. Acta 6, 531 (1972). 217. J. Sand and 0. Gender, Ber. Dtsch. Chem. Ges. 36,2083 (1903); A. Werner and P. Karrer, Helu. Chim. Acta 1, 54 (1918). 218. C. S . Pratt, B. A. Coyle, and J. A. Ibers, J. Chem. SOC.A p. 2146 (1971). 219. B. F. Hoskins, F. D. Whillans, D. H. Dale, and D. C. Hodgkin, Chem. Commun. p. 69 (1969). 220. D. A. Snyder and D. L. Weaver, Inorg. Chem. 9,2760 (1970). 221. J. H. Swinehart, Coord. Chem. Reo. 2,385 (1967). This reference is a comprehensive review of the chemistry of Fe(NO)(CN):-. 222. C. A. Reed and W. R. Roper, J. Chem. SOC.,Dalton Trans. p. 1243 (1972). 223. F. Bottomley and J. R. Crawford, Chem. Commun. p. 200 (1971); J. Am. Chem. SOC.94, 9092 (1972). 224. P. G. Douglas and R. D. Feltham, J. Am. Chem. SOC.94, 5254 (1972).

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225. S . A. Ademyi, F. J. Miller, and T. J. Meyer, Inorg. Chem. 11,994 (1972); F. J. Miller and T. J. Meyer, J . Am. Chem. SOC.93, 1294 (1971). 226. 0. A. Ileperuma and R. D. Feltham, J . Am. Chem. SOC.98,6039 (1976). 227. W. L. Bowden, W. F. Little, and T. J. Meyer, J. Am. Chem. SOC.96,5605 (1974); 98,444 (1976). 228. J. H. Swinehart and W. G. Schmidt, Inorg. Chem. 6, 232 (1967); S. K. Wolfe and J. H. Swinehart, ibid. 7, 1855 (1968). 229. J. A. McCleverty, C. W. Ninnes, and I. Wolochowicz, Chem. Commun. p. 1061 (1976). 230. A. P. Gaughan, Jr., B. J. Corden, R. Eisenberg, and J. A. Ibers, Inorg. Chem. 13,786 (1974). 231. M. W. Schoonover and R. Eisenberg, J. Am. Chem. SOC.99,8371 (1977). 232. J. M. Kruse, E. I. duPont de Neumours, U.S. Patent 3,342,847 (1967). 233. R. A. Clement, U. Klabunde, and G. W. Parshall, J . Mol. Catal. 4,87 (1978). We thank Dr. Parshall for a preprint of this manuscript. 233a. M. W. Schoonover, E. C. Baker, and R. Eisenberg. J. Am. Chem. SOC.101, 1880 (1979). 233b. C. P. Guengerich and K. Schug, Inorg. Chem., 17 1378 (1978). 233c. K. Schug and C. P. Guengerich, J . Am. Chem. SOC.101,235 (1979). 234. D. E . Hendriksen and R. Eisenberg, J . Am. Chem. SOC.98,4662 (1976). 235. S. Bhaduri, B. F. G. Johnson, A. Pickard, P. R. Raithby, G. M. Sheldrick, and C. I. Zuccaro, Chem. Commun. p. 354 (1977). 235a. S. Bhaduri and B. F. G. Johnson, Transition Met. Chem. 3, 156 (1978). 236. J. A. Kaduk and J. A. Ibers, to be submitted. We thank Prof. Ibers for a preprint of this work.

.


The Kinetics of Some Industrial Heterogeneous Catalytic Reactions M . I . TEMKIN Karpov Institute of Physical Chemistry Moscow. USSR

I. I1. 111. IV . V. VI . VII . VIII . IX. X.

Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . The Measurement of Reaction Rate . . . . . . . . . . . . . . . . Mass Transfer in Heterogeneous Catalysis . . . . . . . . . . . . . Ideal Adsorbed Layers . . . . . . . . . . . . . . . . . . . . . . The Routes of Complex Reaction . . . . . . . . . . . . . . . . . The Steady-State Conditions . . . . . . . . . . . . . . . . . . . Some General Relations for Steady-State Reactions . . . . . . . . . Reversible Many-Stage Reactions . . . . . . . . . . . . . . . . . Nonuniform Surfaces . . . . . . . . . . . . . . . . . . . . . . Adsorption Equilibrium and the Kinetics of Reaching it on Nonuniform Surfaces . . . . . . . . . . . . . . . . . . . . . . XI . The Kinetics of Reactions on Nonuniform Surfaces . . . . . . . . . XI1. Oxidation of Ethylene into Ethylene Oxide . . . . . . . . . . . . . XI11. Hydroxylamine Synthesis . . . . . . . . . . . . . . . . . . . . . XIV . Reaction of Methane with Steam . . . . . . . . . . . . . . . . . XV . Ammonia Synthesis-Simple Kinetics . . . . . . . . . . . . . . . XVI . Ammonia Synthesis-Complicated Kinetics . . . . . . . . . . . . . XVII . Carbon Monoxide Conversion . . . . . . . . . . . . . . . . . . XVIII . Isotopic Exchange between Steam and Hydrogen . . . . . . . . . . XIX . Phosgene Synthesis . . . . . . . . . . . . . . . . . . . . . . . X X . Reactions of Carbon with Carbon Dioxide andsteam . . . . . . . . XXI . Ammonia Oxidation . . . . . . . . . . . . . . . . . . . . . . . References . . . . . . . . . . . . . . . . . . . . . . . . . . .

I

.

173 174 178 184 188

192 197 203 207 213 223 230 239 244 250 257 263 267 270 273 279 287

Introduction

The phenomenon of catalysis is a subject of chemical kinetics. as follows from Ostwald’s definition of catalysis. Therefore. the accumulation of data on the kinetics of concrete catalytic reactions favors the progress of the theory of catalysis. It is expedient to choose for kinetic studies the reactions of industrial importance since the kinetics of these reactions are not only a source of in173

Copyright 0 1979 by Academic Press. Inc . All rights of reproduction in any form reserved. ISBN 0-12-007828-7

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formation for the theory, but also find immediate practical applications in process optimization and reactor design. The calculations of this kind are performed now with electronic computers on a large scale. This chapter presents an account of the main results of studies of kinetics of some industrial heterogeneous catalytic reactions. The studies have been carried out by the author with his co-workers at the Karpov Institute of Physical Chemistry (Moscow, USSR). The presentation is not chronological ; the reactions are arranged based on the character of interpretation of their kinetics. For compactness and to avoid repetition, some general aspects of experimental methods and theoretical principles developed in the course of our investigations are formulated prior to the discussion of individual reactions. A more detailed analysis of many of the subjects dealt with in this part of the chapter can be found in the contributions by Kiperman (1) and Snagovskii and Ostrovskii (2). The limits of this article did not allow inclusion of the results of the comparison of kinetic equations with experimental data. They are given in the referenced original papers. II. The Measurement of Reaction Rate

The rate of a reaction is its intensity expressed quantitatively. Let a reaction be described by a chemical equation of the form aA + a’A = bB

+ b‘B

(1)

where A and A are reactants, B and B are products, and a, a’, b, b’ are stoichiometric coefficients. If a molecules of substance A and a’ molecules of substance A‘ have reacted and formed b molecules of substance B and b‘ molecules of substance B’, we shall say that there has been one run of the reaction along (1). The total number of reaction runs divided by the Avogadro number ( N = 6.02 x mol-’) is called “the extent of reaction” (3); this variable, which was introduced into chemical thermodynamics by de Donder, is usually denoted as 5. The extent of the reaction, as the amount of matter, is measured in moles. Some authors state that the reaction rate is dt/dt where t stands for time. But dt;/dt is proportional to the size of the reactor and, hence, is an extensive property like t, and not an intensive property, as should be the reaction rate, according to the definition of the term. The derivative dt/dt is to be called the reactor productivity, but not the reaction rate. In general, when reaction rate, Y, varies with time and is not the same in

INDUSTRIAL HETEROGENEOUS CATALYTIC REACTIONS

175

different parts of the reaction space, it is defined as follows. If the reaction is homogeneous r = a2gat a0

(2)

where u is the volume in which the reaction occurs. For a heterogeneous reaction

r = a2t/at as (3) where s is the area of the surface on which the reaction occurs. Thus the rate of homogeneous reaction is the density of the reactor productivity, and that of heterogeneous reaction is the surface density of the reactor productivity. It is more convenient for theoretical considerations to define the reaction rate on the basis of the number of its runs without dividing the number by N . Engineering calculations often need the rate of heterogeneous catalytic reaction to be referred not to the surface, but to the mass of the catalyst or the volume of the bed of the catalyst grains. The relation between different expressions of a heterogeneous catalytic reaction rate is determined by the values of specific surface area, o, of the catalyst and the bulk density, P b , of the catalyst bed. The total surface of the catalyst, s = om, where m is the mass of the catalyst; m = PbU, where u is the volume of the bed. The notion “reaction rate” will be discussed more completely later (see Section VI), including the cases when the reaction simultaneously follows several pathways with the formation of different products. In most of the investigations described below the reaction rates were measured by the circulation flow method proposed in 1950 (4). This method offers a possibility of realizing a steady-state heterogeneous catalystic reaction without any concentration and temperature gradients; i.e., it belongs to the group of methods which were called “nongradient” (5). The scheme of the circulation flow system is shown in Fig. 1. The apparatus for measurements at atmospheric pressure is made of glass. The mixture of gaseous reactants is fed into the system at a constant rate through pipe 1. The converted gas mixture flows out through pipe 2 ; it is analyzed, e.g., chromatographically. Pump 3 provides for the intense circulation of the mixture. The hollow glass piston of the pump is packed with iron and is actuated by an electric magnet; the valves of the pump are arranged in such a way that both strokes of the piston are utilized. Reactor 4 (containing catalyst grains) is placed in an oven. Before contacting the catalyst, the gas mixture is heated in a coil also immersed in the oven. Circulation rate (i.e., the velocity of the passing of the gas mixture through reactor 4) is many times higher than the flow rate (the velocity of the gas mixture in pipes 1 and 2).

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M. I. TEMKIN

- II

FIG.1. Circulation flow system.

Some time after the beginning of the operation, the system comes to a steady state and the composition of the outgoing gas mixture in pipe 2 ceases to change. In essence, it is sufficient to pass through the system a volume of the gas mixture 5 times that of the system. After that the reaction continues indefinitely under constant conditions. Because of the large circulation rate, the degree of conversion at each single passage of the gas mixture through the reactor is small; therefore, the composition and temperature of the mixture at the reactor inlet and outlet are almost the same. Conversely, the composition of the mixture at the cycle inlet and outlet can differ considerably; the overall degree of conversion possibly being high since the gas mixture repeatedly passes through the reactor and is renewed to only a small extent by the feedstock. It has already been mentioned that the experiments of this sort realize a practically nongradient course of reaction. Specifically, large-scale nongradient conditions exist in the reactor, viz., the composition of the gas mixture and its temperature in the space between catalyst grains are virtually identical in the whole volume of the reactor. But this does not preclude considerable temperature and concentration gradients directed from the surface to the center of the catalyst grain. These gradients can be eliminated,


177

however, if sufficient small grains are used. On the other hand, experiments with grains of the size which is used in industry can provide data on the effect of diffusionin the pores of the catalyst and heat transfer in the grains on the course of reaction under the industrial conditions. The measurements with small grains characterize a reaction in the kinetic region. The composition of the gas mixture flowing out of the cycle, which is determined by analysis, is just the composition of the reaction medium. The reaction rate is easily calculated from the same results of analysis; the extent of the reaction, t, during some time, t , is equal to the increment of the amount of any product or to the decrement of the amount of any reactant when the gas mixture has passed through the system divided by the respective stoichiometric coefficient. Since the reaction is steady state, 5 is proportional to t ; and since the reaction proceeds under nongradient conditions, 5 is proportional to the surface of the catalyst, s. As a result, the general Eq. (3) reduces to r = lyts.

(4)

If the reaction is conducted in a closed system and an identical course of the reaction on the whole surface of the catalyst is achieved, e.g., by circulation, then r = (l/s)(d(/dt).

(5)

If the reaction is steady state and carried out in a flow reactor, then r = (l/t)(d 0; if, on the other hand, B , is a reaction product, bi < 0 and, usually, do/&, < 0.Thus, all nonvanishing products bi(8w/dci), are usually negative (although exceptions do exist). The term of (10) containing u4 is, as a rule, positive. The greater q differs from 1, the slower the series (10) converges; when the difference is large, the series becomes divergent. Limiting the expansion as given above, i.e., terminating it by the term with u4, one may obtain q with a sufficient accuracy at the level not less than ca. 0.85. In order to judge whether the kinetic region is achieved, a formula adequate for q values still nearer to 1 is sufficient; therefore, the series can be terminated by the term with u2. Let us introduce a symbol Y = - a2 1 (bi/Di*)(do/dci),. i

The minus sign in the definition makes Y usually positive. For q close to unity q x 1-Ar.

(12)

Therefore, the effect of diffusion may be neglected if the absolute value of the dimensionless criterion, Y, is sufficiently small. Thus, if /TI6 0.3, then, according to (12), 0.98 < q 6 1.02.


181

In a number of cases the kinetics is described by an equation of the form w =k

n cYi i

where ni is the order of the reaction with respect to substance Bi,k is the rate constant. Then, according to ( l l ) ,

r = -azw, ci (binjoi*(ci),). For first-order reactions, w = kc, n = 1, b = - 1 ; therefore, T Thiele (12) has obtained an exact expression for this case:

(14) =

a2w,/D*c.

where cp = a(k/D*)”’.

(16)

Therefore, cp = f l . There was an opinion that a2w,/D*c [or ~ ( w l , l D * c ) ’ / ~is] a generally applicable criterion of the attainment of the kinetic region. However, this is not always true. In the case of a reversible first-order reaction A P B.

(17)

w = k + c A - k - c , where k , and k - are rate constants for the forward and reverse reactions, bA = -1 and bB = 1. It follows from (11) that T = a2(k+/DA*+ k-/D,*). Since the problem is linear, an exact solution can be obtained; if we have

cp = a(k+/D,*

+ k-/DB*)“2,

(18)

then Eq. (15) is again valid (13). According to (15) and (18), the efficiency factor, q, is constant regardless of the degree of approach to the equilibrium. It follows that can be much smaller than unity no matter how small the reaction rate on the periphery of the grains, 0,. Thus, in the case of a reversible reaction, one cannot conclude that the reaction occurs in the kinetic region only on the basis that its rate is small; for a correct judgment, the value of the criterion Y should be found as defined by (1 1). One should take into account the specific features of gas diffusion in porous solids when measuring effective diffusion coefficients in the pores of catalysts. The measurements are usually carried out with a flat membrane of the porous material. The membrane is washed on one side by one gas and on the other side by another gas, the pressure on both sides being kept

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M. I. TEMKIN

identical. Often hydrogen and nitrogen are used. The amount of one gas passing through the membrane per unit time is determined from its content in the stream of the second gas after the membrane and the velocity of the stream. Having known the thickness and the area of the membrane, the effective diffusion coefficient of the first gas is calculated. Starting from this, further calculations of effective diffusion coefficients of participating gases for the required values of temperature and pressure are performed, calculation depending on the interpretation of diffusion; namely whether it is of the Knudsen, bulk, or intermediate type. Bulk diffusion coefficients in binary gas mixture are almost independent of the ratio of components of the mixture. Therefore, it was supposed that if diffusion in the measurements described above is of the bulk type, i.e., the free path of molecules is much lesser than the diameter of pores, then the first gas diffuses into the second gas at the same rate as the second gas diffuses into the first. However, as early as in 1833 Graham described in a number of articles (14) the results of his experiments on diffusion of various gases in porous bodies, mainly gypsum. He concluded that gases diffuse into each other not at equal velocities, but at velocities that are inversely proportional to the square roots of their densities or, in modern terms, to the square roots of molecular weights. In the concluding paragraph Graham wrote : The law at which we have arrived (which is merely a description of the appearances, and involves, I believe, nothing hypothetic), is certainly not provided for in the corpuscular philosophy of the day, and is altogether so extraordinary, that I may be excused for not speculating further upon its cause, till its various bearings, and certain collateral subjects, be fully investigated.

Later, after the works of Knudsen on gas flow at low pressures, the result of Graham was explained supposing that in his experiments the free path of molecules exceeded the diameter of pores [see, e.g., the book by Herzfeld (IS)].Only in 1955 Hoogschagen proved experimentally that even with pores considerably larger than the free path of the molecules, the velocities of the counterflows of gases are inversely proportional to the square roots of molecular weights (16). Subsequently, this result was confirmed by a number of authors (17, 18); none of them, however, like Hoogschagen, mention Graham’s work. In justice, this relationship must be called “Graham’s law,” as named by Maxwell (19). The explanation of Graham’s law given by Hoogschagen is not complete, as subsequent authors (17,18) stated. However, the attempts of these authors to give a more complete explanation for the law are not convincing. It is known that at conventional measurements of diffusion coefficients in binary gas mixtures using wide capillaries, equal velocities of counterdiffusion of the components are observed. From the considerations developed


183

in the referenced articles, the reason for the difference between narrow capillaries where Graham’s law is followed and wide capillaries where this law is not obeyed cannot be seen. Let steady-state diffusion of components A and B of a gas mixture occur in opposite directions at a capillary diameter that exceeds many times the free path of the molecules, the pressure along the capillary being constant. The concentration of A decreases from left to right and that of B from right to left. Let us assume that the gas mixture at the beginning does not move. A small section of the wall of the capillary is impinged upon by the molecules hitting the wall from the distance of the order of the average free path. Since the concentration of the molecules of the component A on the left is larger than on the right, more such molecules will be coming from the left. As a result, the section will be subjected not only to the normal force of partial pressure of A, but also to a tangential force directed in parallel to the axis of the capillary. The tangential force owing to component B has the opposite direction. However, if the molecular weights of A and B are different, these tangential forces do not equilibrate. According to the law of action and counteraction, the gas mixture in this case will be affected by a force equal in value to the resultant of these tangential forces, but oppositely directed. Such forces will arise at all elementary sections of the capillary surface. Under their action the gas mixture will be accelerated until, as the result of addition of the velocity of the gas mixture as a whole and velocities of molecules relative to the gas mixture, the velocities of the molecules relative to the capillary will change so that the total momentum transfer to the wall would vanish. Then the force applied to the gas mixture and, consequently, the acceleration will also vanish and the gas mixture will move at a constant velocity. This gas flow, the necessity of which has not been recognized by the previous authors, may be called concentrational effusion, in analogy to a long known phenomenon of thermal effusion. The flow occurs in a form completely different from the usual flow of a gas in a capillary under the effect of pressure difference, viz., the velocity of the gas mixture in the case of concentrational effusion is identical at all distances from the axis of the capillary, except for the layer adjacent to the wall, the thickness of which is of the order of the free path, in which the velocity decreases up to zero. The requirement that the tangential momentum transfer to the wall per unit time be equal zero immediately leads, as it has been already demonstrated by Hoogschagen, to Graham’s law. Such is the explanation of Graham’s law from the viewpoint of “corpuscular philosophy”; from it follows that this law should be valid for a capillary of any width on the condition that the pressure along the capillary is constant. Why then does Graham’s law not manifest itself in wide capillaries? The reason is that the requirement of the constancy of pressure becomes stricter, the wider the capillary. It is known that, according to Poiseuille’s

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law, the volume flowing through the capillary per unit time is proportional to d4 where d is the diameter of the capillary. A small and therefore imperceptible difference in pressures in wide capillaries creates a Poiseuille’s flow compensating the concentrational effusion. It can be calculated, for instance, that in the case of counterdiffusion of H2 and N, at 1 atm and 20°C, this compensation at d = 0.01 mm requires the pressure difference, AP = 3.4 Torr and at d = 2 mm it is sufficient to have AP = Torr. Let a vessel containing hydrogen be connected through a valve and a capillary of 2 mm diameter to a vessel containing nitrogen and let, as long as the valve is closed, the pressure in both vessels be exactly 1 atm. When the valve is opened, concentrational effusion directed towards the vessel with nitrogen starts ; at the moment the pressure in the vessel with nitrogen becomes higher than in the vessel with hydrogen by Torr, effusion will be compensated by Poiseuille’s counterflow and the gases will counterdiffuse at equal velocities. The discussion of diffusion in pores of a catalyst requires taking into account the specific features of diffusion in narrow capillaries.

IV.

Ideal Adsorbed Layers

The regularities of reactions on the catalyst surfaces are of a very complicated nature and their description is only possible on the basis of schematic and simplified physical models. A model of this kind should, on the one hand, reflect the main features of the phenomenon and, on the other hand, result in comprehensible mathematical expressions. The model of an ideal adsorbed layer or, in terms of the author of the model, Langmuir, “simple adsorption” (20)is the simplest and historically the first of the models retaining their importance until now. It is assumed that the surface of a solid consists of a definite number of areas of atomic size, “sites”; each of these sites is capable to bind one particle, atom or molecule, in an adsorption process. The adsorbed layer is ideal if (1) all surface sites are identical and (2) the interaction between adsorbed particles may be neglected. An ideal adsorbed layer possesses the properties of a perfect (ideal concentrated) solution formed by adsorbed particles of one or several species and free sites. Therefore, mass action law for the rates of surface reactions and corresponding equilibria is formulated quite similar to the law for volume reactions in ideal systems with the only difference being that the equations may also contain, along with surface concentrations of substances, surface concentrations of free sites. Mass action law cannot, however, be applied to a reaction rate if two (or more) adsorbed particles participate in an elementary act when the surface


185

mobility of the particles is low and the exchange of particles between the surface and the volume is not sufficiently rapid. Then, as a result of the reaction, the probability of encountering the particles of reactants on neighboring sites is lower than the probability at an entirely chaotic distribution of the particles on the surface. This reservation is not important for the examples we shall discuss below; mass action law holds true for them. In contrast to spatial distribution, the equilibrium energy distribution of adsorbed particles cannot be violated to any substantial degree by reaction since energy is rapidly transferred between adsorbed particles and solids. Therefore, the activated complex method may be applied to rates of surface reactions. For this we consider the activated complex (transition state) of a surface reaction as a likeness of adsorbed particle (21). But, assuming that each adsorbed particle occupies only one site, it is necessary, even in the simplest kinetic model, to consider that activated complexes are able to occupy not only one, but also several surface sites (21). For example, the usual picture of a reaction between two particles adsorbed on neighboring sites involves, in fact, the notion that the activated complex occupies both sites. When the activated complex occupies several sites, this does not create any difficulty for the theory since the surface concentration of activated complexes is an infinitesimal quantity, and so the possibility of overlapping the required sites is excluded. If particles enter the surface activated complexes directly from the volume, e.g., when the reaction occurs as a result of impact of molecules from the gas phase upon the adsorbed molecules, the expression for the reaction rate will contain, together with surface concentrations, the values of volume concentrations. These impact mechanisms were long ago proposed by Langmuir (22) for the reactions of CO and H, with 0, on the surface of platinum; the reaction occurs at the impact of a CO or H, molecule against an adsorbed 0 atom. Such reactions seem to be numerous (23). Along with this, the above-mentioned adsorption mechanisms that involve the reaction between two adsorbed particles are possible. Elementary acts of surface reactions in which more than two particles participate are hardly probable. It is convenient to consider the processes of adsorption and desorption, being the constituents of heterogeneous catalytic reactions, as particular cases of surface reactions. Thus, the reversible process of adsorption and desorption of substance A can be expressed as follows : AfZ#ZA

(19)

where Z is the free site of the surface and ZA is the site occupied by the molecule A. Applying mass action law, we obtain

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M. I. TEMKIN

for the rate of adsorption and

r - = k_[ZA] (21) for the rate of desorption. Here PA is the partial pressure of substance A in the gas phase, and [ZA] are surface concentrations of free and occupied sites, k, and k- are rate constants. We shall in the following denote by [Z] and [ZA] the fractions of the free and occupied surface, respectively; this only results in a change in the numerical values of k, and k- compared t:, the case when surface concentrations are expressed as number particles on the unit surface area. With this notation

[a

[Z]

+ [ZA] = 1.

(22)

Equations (20) and (21) are Langmuir expressions for adsorption and desorption rates. If the equilibrium is reached with respect to the process (19), then, in accordance with mass action law, [zAI/PA[zl

(23)

=a

where a is the equilibrium constant of adsorption (19) which is usually called “adsorption coefficient.” From (22) and (23) it follows that

[ZA] = aP,/(1

+ UP,),

(24)

i.e., the Langmuir adsorption isotherm; [ZA] is usually denoted as 8, or simply 8 if only one substance is adsorbed. The value a, being an adsorption equilibrium constant, is related to the standard Gibbs’ energy of adsorption, AGOo,by the equation

AGOo= -RTlna.

(25)

Here the state with [ZA] = [ZIis taken as a standard state of the adsorbed layer; thus, in the case when only one gas is adsorbed, the layer is in the standard state at the coverage 1/2. It can be easily seen that l/a is the equilibrium pressure at [ZA] = [Z], i.e., at the standard state of the adsorbed substance. This value may be called “desorption pressure”; we shall denote it as b. It is analogous to vapor pressure or dissociation pressure in monovariant systems (24). Indeed, in the case of equilibrium of liquid with its vapor, the surface from which evaporation occurs is equal to the surface for condensation; the same equality is realized at the adsorption equilibrium if the fraction of the occupied surface is equal to that of the free surface. This analogy explains the applicability of the Nernst approximate formula to desorption pressure (24): log,, b

=

-(AHO/4.58T)

+ 1.75 log,,

T + i*

(26)


187

where AHa is molar adsorption heat (isosteric), z* is a conventional chemical constant. The latter, as was suggested by Eucken, can be taken as 3 for all molecules except H, and 1.5 for atoms and H, . With (26) the order of magnitude of b can be easily estimated if AHa is known, or, vice versa, AHa estimated from b. An elementary reaction may be, in the general case, represented as follows: aA + a‘A‘

+ iZI + i’Z1’ + zZ+products.

Here A and A are the substances entering the reaction from the gas phase, I and I’ are the substances entering the reaction from the adsorbed state, a, a’, i, i’ are stoichiometric coefficients,z is the number of free sites required, in addition to the sites occupied by particles rand I’,for the formation of the activated complex. In reality, the elementary reaction involves no more than three particles including free sites, Z; therefore, a + a‘ i ‘i + z I 3 and some of the values in the left-hand part of the inequality equal zero. If the gas phase and the adsorbed layer are ideal, then the rate, r, or the reaction will be given by the equation

+ +

where k, is the rate constant, CA and CAt,are the concentrations of the ., by the substances A and A in the gas phase. We shall express CA and C number of molecules in the unit volume. Reaction rate, r, will be defined as the number of reaction acts per unit area per unit time. According to the theory of absolute rates of surface reactions (22),

k, = K(k,T/h)L(gf’/fAafdt~~~~,)e-(~a’kB~). (28) Here K is the transmission coefficient, for the adiabatic reactions (in the quantum mechanical sense) K z 1, k, is the Boltzmann constant, h is the Planck constant, L is the number of sites for adsorption on the surface of a unit area, g is the number of positions of the activated complex when one of the sites occupied by the complex is fixed (when the activated complex occupies two sites, g = 2, 4, or 6 depending on the symmetry of the twodimensional lattice of the crystal face),fA, etc. are partition functions (sums over states) of particles A, etc. calculated with the energies of microstates measured from zero energies of given gaseous or adsorbed particles, i.e., including only the energy of thermal motion; for gas particlesf values refer to unit volume; f # is the partition function of the activated complex not including the factor corresponding to the “reaction path” coordinate, E, is the energy barrier of the reaction (activation energy at T = 0) for one reaction act. It is not difficult to obtain from the formula (27) containing concentrations the corresponding formula with partial pressures of the components of the

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M. I. TEMKIN

gas mixture, by means of the equations :

CA = PA/kBT, CAP = PAr/k,T. (29) The constant k, is substituted then by the rate constant k ; according to (271, k

= k,-/(kBT)"+"'.

(30)

The expression for k obtained from (28) and (30) can be put in the following form :

k = tc(kBT/h)Lge-AG*'kBT

(31)

where AG' is the Gibbs activation energy of the reaction (a standard one, but it is not denoted since for the process of activation other AG values are never considered). Referring to one reaction act, c+

For an estimation of the order of magnitude ofthe preexponential factor in the Arrhenius equation for a surface reaction between simple molecules, the termsf' and&, may be taken as equal to 1 since they include only vibrational modes that are degenerate if hv k J ( v being the vibration frequency) (21). The values offA for molecules in the gas phase are calculated in the usual manner.

+

V. The Routes of Complex Reactions We define an elementary reaction as a multitude of reaction acts identical in the nature of participating particles and in the direction of change. A set of different elementary reactions occurring jointly and related each to other by having some of the participating species in common will be called a complex reaction. Catalytic reactions, homogeneous and heterogeneous (as well as chain reactions, etc.), are complex. The result of a complex reaction as given by chemical analysis is described by one or several stoichiometric equations. For instance, the reaction of ethylene with oxygen on the surface of silver results in the formation of ethylene oxide, carbon dioxide, and water vapor; it is described by the equations

+ $0,= C,H40 + 30, = 2C0, + 2 H 2 0

C,H4 C,H4


189

Such equations will be called “overall chemical equations,” reactants and reaction products will be termed “reaction participants.” Along with reaction participants, the chemical equations of elementary reactions comprising the complex reaction include other species that do not appear in the overall equations. They are called “intermediates.” For the purpose of the kinetic analysis of a complex reaction, its elementary reactions are grouped into stages (or steps). A stage, first of all, is a pair of mutually reverse elementary reactions or one elementary reaction if it is irreversible. Such stages will be called “simple stage.” Sometimes several simple stages may be united into one complex stage; this is permissible when the rates of the constituent elementary reactions are extremely high compared to the rate of a complex reaction as a whole. Chemical equations of the stages contain, like the chemical equations of elementary reactions, not only reaction participants but also intermediates. Overall reaction equations are linear combinations of chemical equations of stages, i.e., they are obtained by the addition of chemical equations of stages multiplied by certain numbers (positive, negative, or zero). The numbers must be chosen in such a way that the overall equations contain no intermediates. According to Horiuti, such numbers are called stoichiometric. To illustrate this notion by a simple example, let us assume that the mechanism of oxidation of SOz on the surface of a solid catalyst, e.g., platinum, is described by the following scheme : 1.

2.

2 2 + o2P 2 Z O zo + sozP Z + so, 0,

1

2

(33)

+ 2SOZ = 2s0,

In this scheme, as always in the following, Z denotes a site on the surface of the catalyst, the arrows denote elementary reactions, and in the overall equation the equality sign is employed. Stoichiometric numbers by which the stage equations should be multiplied to obtain the overall equation are given on the right side. In this instance intermediates are Z and ZO. Another scheme equivalent to the previous one may be used : 1.

2.

22

+ o2 P 2 Z O

zo + so* # Z + so,

4 I

(34)

$0, + SOz = SO,

Thus fractional stoichiometric numbers are also allowable (although it is always possible to write overall equations so that all stoichiometric numbers are integers).

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M. I. TEMKIN

The overall equation of scheme (34) is obtained from the overall equation of scheme (33) when it is multiplied by 1/2. The possibility of multiplying overall equations by arbitrary numbers is shown by the use of the equality sign. Such an operation is not permissible for the equations of simple stages; e.g., the equation Z++02#ZO

is senseless: it would mean that 1/2 of the 0,molecule participates in an elementary act. On the contrary, the equation

:o, + so,

=

so,

means only that from n 0, molecules and 2n SO, mdlecules 2 n SO3 molecules are formed, where n is some number; the equation describes the stoichiometry of the process, but not its mechanism. In this example the exclusion of the intermediates could be effected, essentially in one way only, since the column of stoichiometric numbers (‘i2) is obtained from the column (i)by multiplication by an arbitrary factor; this holds for any other suitable column of stoichiometric numbers. However, this is not always the case-even when the reaction is described by one overall equation. Thus, a possible mechanism of oxidation of hydrogen on platinum is (25) Mi)

+ +

1. 02 2Z+2ZO 2. H, 2ZFt2ZH 3. ZO + ZH+ZOH Z 4. ZOH + ZH + H 2 0 + 2 2 5. ZO + H, + Z H 2 0

+

+

1 0 0 0 2

”2’

1 2 2 2 0

(35)

N’);M2);0,+ 2H2 = 2Hz0

This scheme means that H,O can be formed as a result of an impact of a H, molecule upon an adsorbed 0 atom, stage 5, or in the reaction between particles adsorbed on the adjacent sites of the surface, stage 4. A set of stoichiometric numbers of the stages producing an overall reaction equation is called, after Horiuti, a “reaction route.” Thus, in scheme (35) there are two reaction routes, N “ ) and N(’); route N ” ) cannot be obtained from N(l) through multiplication by a number. Therefore, routes N ( ’ )and N(’) are essentially different, although their respective overall equations are identical. Let us denote the stoichiometric number of stage s of route p as vp’. If the stoichiometric numbers of all stages of route M3)are obtained from the , to the equation stoichiometric numbers of routes N(’)and N ( 2 ) according vs( 3 )

= c I v s( l ) + Czv;2’

(36)


191

where C, and C2 are arbitrary constant numbers, we shall say that route

N 3 )is a linear combination of routes N(’)and N(’) and express this relationship by a symbolic equation ”3’

=

C,”1’

+ C,”?

(37)

Such a definition can, evidently, be extended to any number of routes. It is clear that if N ( l ) ,N(’), N 3 ). . . are’routes of a given reaction, then any linear combination of these routes will also be a route of the reaction (i.e., will produce the cancellation of intermediates). Obviously, any number of such combinations can be formed. Speaking in terms of linear algebra, the reaction routes form a vector space. If, in a set of reaction routes, none can be represented as a linear combination of others, then the routes of this set are linearly independent. A set of linearly independent reaction routes such that any route of the reaction is a linear combination of these routes of the set will be called the basis of routes. It follows from the theorems of linear algebra that although the basis of routes can be chosen in different ways, the number of basis routes for a given reaction mechanism is determined uniquely, being the dimension of the space of the routes. Any set of routes is a basis if the routes of the set are linearly independent and if their number is equal to the dimension of the space of routes. It may occur that when stoichiometric numbers of stages are chosen to form a route, the cancellation of one intermediate at the summing up of the stage equations inevitably leads to the cancellation of another. Thus, in the mechanism (33) this is the case with the intermediates Z and ZO. This results from the equality [Z]

+ [ZOI = 1.

(38)

We shall call such relationships “balance equations.” If there were two types of sites on the surface, we would have two balance equations. It can be shown algebraically that the number of basic routes, P,is determined by

P=S+ W-J (39) where S is the number of stages, W is the number of balance equations, and J is the number of intermediates. In the example (35) S = 5, W = 1, and J = 4 (intermediates Z, ZO, ZH and ZOH); therefore, P = 2. Since N”) and MZ) are linearly independent and P = 2, N1) and N2)form a basis. The relationship (39) was given in a somewhat different form by Horiuti (26). A linear independence of routes does not imply linear independence of the respective overall equations. For instance, as mentioned, the basic routes of the reaction (35) result in the same overall equation. A more complicated

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M. I. TEMKIN

example is offered by the reaction of oxidation of ethylene on silver; although its stoichiometry is described by two equations given above, three basic routes correspond to the reaction mechanism discussed later in Section XII. A set of linearly independent overall equations such that any other acceptable overall equation of the reaction in question is a linear combination of equations of this set forms a basis of the overall equations. Let the number of basic overall equations be Q;it is evident that Q IP. Let us denote the number of substances participating in the reaction as M and the number of independent components, in the sense this notion is used in the Gibbs’ phase rule, as C ; then Q=M-C.

(40) In the reaction (35) M = 3 (0,,H,, and H,O), and 0, and H, can be taken as independent components; thus C = 2. Therefore, Q = 1. In the cases when P > Q the basis of routes can be always chosen in such a way that the P - Q routes will have overall equations 0 = 0. Such routes do not result in any chemical transformation, they will be called “empty.” Thus, instead of the basis of routes in the scheme (35), a basis may be used consisting of route N ( ’ ) (or W 2 ) )and an empty route [ M 2 )- N “ ) ] . We should not be embarrassed by the negative sign of the stoichiometric number of stage 5 in the route [N”) - N‘”] indicating that the stage follows the reverse direction. Although stage 5 in the scheme (35) is shown to be irreversible, it implies only that the rate of the reaction in the reverse direction is so small that it can be ignored; in general, all elementary reactions are reversible and the reaction along the empty route under discussion is possible. Moreover, this is of no importance since the operation of the substitution of one basis of routes for another is of a purely algebraic nature. The basis of routes with Q nonempty and P - Q empty will be called “stoichiometric basis of routes” because, being a basis of routes, it determines simultaneously a basis of stoichiometric overall equations. VI.

The Steady-State Conditions

In deriving the kinetic equations of heterogeneous catalytic reactions, the surface concentrations are assumed to be steady state (or stationary), as has been done by Langmuir in the previously mentioned study of the reactions of CO and H, with 0, on platinum (22). The treatment of surface reactions as including adsorption equilibria widely used by Hinshelwood and other authors is a particular case of this more general approach of Langmuir. Catalytic reactions in industrial reactors are, as a rule, steady state.


193

Laboratory studies of the reactions at steady-state conditions have the advantage of the much simpler mathematical analysis of the results compared to nonsteady processes since the problem of deriving kinetic equations corresponding to a given reaction mechanism is reduced to the solution of a set of algebraic equations instead of differential equations in the general case of a nonsteady reaction. Since mass action law for elementary reactions in ideal adsorbed layers (including also adsorption and desorption processes) coincides in its form with mass action law for elementary reactions in volume ideal systems, general results of the theory of steady-state reactions are equally applicable to volume and to surface reactions. They are very useful when the reaction mechanism is complicated. The concepts of the theory of steady-state reactions are better explained with examples, but here they will be formulated mainly in a general form; the examples will be found below in the discussion of concrete reactions. A reaction is at steady state if the concentrations of all species in each element of the reaction space (i.e., volume in the case of a homogeneous reaction or surface in the case of a heterogeneous reactions) do not change in time. Evidently, this is possible only in an open system such as a tubular reactor or a circulation flow system. It has been stated in Section V that the species participating in elementary reactions that comprise a complex reaction are grouped into two categories : intermediates and reaction participants. In a steady-state reaction, the concentration (volume or surface) of an intermediate does not change in time because the sum of rates of formation of this species in elementary reactions is equal to the sum of rates of its consumption in other elementary reactions. The concentration of a reaction participant does not change in time because, along with the formation and consumption in elementary reactions, this species enters an element of the reaction space or leaves it; in the first case this is a reactant, in the second case this is a product. For an elementary reaction, the concept “run” introduced in Section I1 is equivalent to a single act of the reaction. The number of runs of a simple stage is defined as the difference of the number of runs of the forward and the reverse elementary reactions comprising the stage. We define the run along a route as a set of runs of stages in numbers equal to the stoichiometric numbers of these stages for the given route. For a steady-state reaction the formation of a molecule of an intermediate in an act of an elementary reaction must be compensated by consumption of this molecule in some other elementary reaction. If a new molecule of the same or another intermediate is formed in this reaction, it must also be consumed. Sooner or later a moment comes when a complete compensation of the formation and consumption of the intermediate molecules occurs in

194

M. I. TEMKIN

a certain group of reaction acts. This implies that a run along a certain route of a given complex reaction is completed, the route by no means being an obligatory basic one. Now each route of the reaction can be represented as a linear combination of its basic routes; therefore, a run along any route can be expressed as a corresponding linear combination of runs along the basic routes. In accordance with this, let us substitute for the runs along nonbasic routes (i.e., the runs that are the actual constituents of a reaction) the equivalent runs along basic routes. This means, for example, that if there was a run along a route [C,N(’) + C2iV2)+ C,M3)] where N”), M2), and M3) are basic routes, we shall say that there was C , runs along the route N”), C2 runs along the route M2), and C , runs along the route M3). Let us consider a sufficiently long period of time to neglect the runs along the routes that have begun before this period and not terminated before it, or the routes that have begun during this period and not terminated before its end. It is always possible, since the reaction is at steady state and, hence, the period of time is not limited. Then, neglecting uncompleted routes, we shall have all eventuated reaction acts distributed singularly and without remains between basic routes, forming runs along these routes. The ratio of the number of runs along the basic route obtained in this manner to the period of time and the size of the reaction space (volume or surface) is the rate of the reaction along this basic route. The notion “rate along a basic route” has a rigorous sense only for steadystate reactions; in the general case the runs along stages cannot be singular and without remains divided between basic routes. The rate of a steady-state reaction is determined by the rates along the basic routes, as a vector in the usual three-dimensional space is determined by its components along the coordinate axis. Let us denote the rate of the sth stage as rlsl;if this is a simple stage, then rlsl= r,

- rPs

(41)

where r, and r - s are the rates of the forward and the reverse elementary reactions comprising the stage. The rate along the basic route W’)is denoted as d p ) ; according to its definition there are v$‘)#~)runs of the sth stage per unit time per unit reaction space, which was the share of the NP) route after the distribution of the runs of stages between basic routes. The total number of runs of the sth stage per unit time per unit reaction space is obtained by summation over all P basic routes. Thus, P

This equation was advanced originally by Horiuti and Nakamura (26) in a


195

formal manner, the rates along basic routes being defined as the coefficients in linear relations connecting vip) and rlslvalues. In the most general case the rates of all elementary reactions are mutually comparable, then all stages are simple and not at equilibrium. Equation (42) is applicable to each stage. Expressing rls,in this equation through the concentrations of species (volume of surface concentrations) in accordance with (41) and mass action law, we obtain S equations according to the number of stages that relate rates along basic routes with concentrations. To these S equations, W balance equations are added. J concentrations of intermediates and P rates along basic routes are the unknowns. According to (25)

J + P = S + W,

(43)

so the number of equations is equal to the number of unknowns. If, for some stage r, x r - , B rlslthe stage is called “fast,” it may also be called a quasi-equilibrium stage. In addition, strictly equilibrium stages are also possible; namely, if stoichiometric numbers of a stage for all basic routes are zero, then r, = r - , (ie., rlsl= 0). Herewith r, and r - s may be small. Reversible adsorption of a catalytic poison may serve as an example of such a stage. Quasi-equilibrium and strictly equilibrium stages will be united under the general name “equilibrium stages.” If a stage is an equilibrium one, then instead of the steady-state condition (42) the equilibrium condition according to the mass action law is employed; this does not change the number of equations and unknowns. Fast nonequilibrium stages do not give equations, but then the number of unknowns decreases correspondingly since some concentrations of intermediates are not considered. Thus, (42) always solves completely the problem of the rate of a steady-state reaction, at least in principle. The problem is reduced to the solution of a set of algebraic equations. We shall call (42) “stage steadystate” conditions. In fact, we need to apply this equation only to nonequilibrium, and hence, simple stages; therefore we may substitute for rlsl the difference r, - r - s without loss of generality and use the equation in the following form : P

When a reaction occurs in an ideal system (i.e., in ideal gas mixture, ideal solution, or ideal adsorbed layer), then r, and r - , in (44)are determined by simple mass action law. We shall call “linear” the stages whose rate, rlsl= r, - r - , , depends linearly on the concentrations of intermediates (including free sites of the surface); the stages whose rate depends nonlinearly on the concentrations of intermediates (i.e., includes squares of concentrations of

196

M. I. TEMKIN

intermediates, their products, etc.) will be called “nonlinear.” A stage is linear if each of the two elementary reactions comprising a reversible stage (or the only elementary reaction of an irreversible stage) include only one particle of an intermediate. A reaction mechanism, all stages of which are linear, will be called linear. For such a mechanism, (44)produces a linear set that always has one solution. This solution can be obtained algebraically in an explicit form. If the reaction mechanism is nonlinear (i.e., if it includes nonlinear stages along with linear ones), the existence of several solutions of a system of (44)(i.e., of several steady states of the reaction) is possible in some cases (28). Sometimes the steady-state course of a reaction is not reached at all and sustained oscillations of the rate (29) or a continuous acceleration of the reaction of an exponential type (30) occur. In the industrial catalytic processes discussed below, these possibilities are not realized even in the cases when the mechanisms are nonlinear; therefore, it is not expedient to discuss these possibilities here in more detail. The original more formal proof (31) of (44)consisted in the demonstration of its equivalence to the steadyLstate condition in the Bodenstein’s form; i.e., the condition that the rates of formation of intermediates are equal to zero. Let xis be the stoichiometric coefficient of an intermediate X j in the chemical equation of a stage s. It is understood that xjs > 0 if X j is being formed, xjs< 0 if X j is being consumed in stage s, and xjs = 0 if Xjdoes not participate in this stage. By definition, stoichiometric numbers for the route N(P)satisfy the equation

s:

vpxjs= 0.

s= 1

(45)

Let r ( X j ) be the rate of formation of an intermediate Xj; i.e., the number of molecules X j formed per unit time per unit reaction space. Then F

Substituting rISlfrom (42)into (46),we obtain

c c P

S

V(Xj) =

xjs

s=l

p=l

v:p)r(p)

197


From this and (45) it follows that r ( X j ) = 0. This conclusion holds for all intermediates. Thus, if (42) is fulfilled for all stages, the reaction is at steady state. The conditions of steady state in the form of Bodenstein are obtained having r ( X j ) = 0 in (46). In contradistinction to (42), these conditions may be called the intermediates steady-state conditions. They define only a part of the unknowns, viz., the concentration of intermediates, Xj, but there is usually no difficulty in the determination of the reaction rate from these. In a number of cases the application of states steady-state conditions for obtaining kinetic equations is more convenient. Besides that, these conditions are a source of certain general relations to be discussed later. If the number of basic overall equations, Q, is smaller than that of basic routes, P,it is expedient to use a stoichiometric basis of routes for a description of the reaction. In this case we need to know only Q rates along nonempty basic routes; the remaining P - Q rates are not required (but this by no means implies that they equal zero).

VII.

Some General Relations for Steady-State Reactions

The choice of a basis of routes for a given reaction is ambiguous and, therefore, to a certain degree arbitrary. Let us consider the following problem: having known the rates r“), d2), . r(‘) along the basic routes N1), N2),. . N p ) ,find the rates r“)’, r(’)‘ . . . r“)’ that answer to another basis of routes, N’)‘,N2)‘. . MP)’(31, 32). Let the passage from the first basis (not primed) to the second (primed) be defined by equations of the form “i”‘ = c1 1 v(’) s + c12 v s( 2 ) + . . . C,,v:p’ e

+

v p ‘ = cp 1 v(sl) + cPZ P) s + . . . + cppv:p’ where C l l , etc. are coefficients. Then the sought for rates will be found by solving, with respect to r(’)’,P)’,. . . r(‘)’, the set of equations = r(2) =

r‘p’

=

. . . + Cplr(P)’ + . . . + cP2+P)’ C1 2r(1)’ + cZ2+2)’ . . .

c11r(l)’ + c

2 1r ( 2 ) ’

c1Pr(l)’+ cZ P r(2)’

+

+

...

(48)

+ Cppr(P)‘

It is composed in such a way that the coefficients of each row of the set (47)

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M. I. TEMKIN

are situated in a corresponding column of the set (48). Indeed, according to the equalities (48) D

,..

+ (CPlv:l) + CP2v:') +

. . + Cppvip))r(p)'

Hence, and from (47) P

P

According to (42), the s u m in the left-hand side of (49) equals rlsl;therefore, the right-hand side of the equation also equals rlsi.This means that the set of rates r(')',r(2)',. . r"" satisfies the stages steady-state conditions (i.e., it is the sought for solution). It has been already noted that the rate of a steady-state reaction can be regarded as a vector in the P-dimensional space specified by its components, which are the rates along the basic routes. In terms of linear algebra, the above result means that when the basis of routes is transformed the reaction rate vector along these routes is transformed contravariantly. The equations determining the kinetics of a steady-state reaction comprise, together with unknown rates along the basic routes, unknown concentrations of intermediates. Only the rates as functions of the concentrations of reaction participants are usually required; therefore, the unknown concentrations are to be excluded from the equations. In many cases this is made easier by application of an equation that is obtained as follows (33). We form an identity including the rates of in stages with numbers, sl, s2, . . . s,,,, chosen arbitrary from the total number of stages, S: +

(rsl - r-sl)rszrs3* . . rs,

+ r-&s2

+ r-s,r-sz(rs3 - r - s 3 )

*

- r-s2)f-s3 - - rs,

. . rsm +

*

*

+ r-s,r-szr.-s3.

= r s l r s z r s 3* ~ ram

-

(rs, - r - s m ) rslr-szr-s3. . . r-s,. (50) *


199

It is easy to be satisfied that this identity is correct by opening the parentheses. Expressing the rates of the stages rsl - r - s l - r(s1lY rs2 - r - s 2 = rls*l9 etc. in accordance with the steady-state condition (44)and regrouping the terms one obtains

...

= r s1y s2 rs 1

rs, - r-s1r-s2r-s3 . . * r-s,.

(51)

The order of stages in (51) is arbitrary, their number 1 Im I S. At m = 1 (51) becomes (44). The direction opposite to that accepted in the record of the reaction mechanism may be assigned to a stage, but then the signs of the stoichiometric numbers of this stage should be changed (- instead of and + instead of -). It can be easily seen that the rule on writing the lefthand side of (51) is as follows: the items of the factor at the rate along the basic route r’p) are obtained from the product r,,r,, * . . r,, by a successive substitution of v::), v$), etc., for the first, second, etc., factor, the stage number subscripts acquiring a minus sign in the factors preceding that replaced. Equation ( 5 1) will be called the steady-state reaction equation. We shall denote the rate of a single-route reaction as r and the stoichiometric numbers of its stages as v,. The equation of single-route steady-state reactions follows directly from (51):

+

In a number of reaction mechanisms groups of stages can be separated, such that the stages participate not in all, but only in a part of basic routes. This means that the stages of such groups have nonzero stoichiometric numbers only in these routes and are related to the remaining basic routes only via intermediates. Such group of stages will be called a “block”.’ If a reaction mechanism consists of blocks, it is expedient to apply steady-

’ This term is used in the theory of graphs (34).

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M. I. TEMKIN

state reaction equation to each block separately including in it only the stages of this block. Since stoichiometric numbers of stages not comprising a given block equal zero, the equation will contain only the basic routes of the block and have such a form as if there were no other blocks. In particular, if a block includes only one route, then the single-route steady-state reaction equation can be applied. A general method of application of (51) for the elimination of unknown concentrations with the help of a graph of a reaction mechanism is described elsewhere (27). Here we limit ourselves to the application of (52) to a simple two-stage mechanism : 1. Z + A , # Z I + B , 2. ZI + A , P Z + B2

A,

+ A,

=

1 1

(53)

B, + B2

In scheme (53) I denotes an intermediate particle adsorbed on site Z,A,, A,, B,, and B, are the molecules of gaseous reaction participants. The rates of the elementary reactions are rt = kl[Z]PAl; r - l = k-l[ZI]PB, T2 =

k,[zI]p~,;

Y-2

(54)

= k_2[z]P~,.

In addition we have a balance equation:

[Z]+ [ZI]= 1. Let us take s1

=

1, s2 = 2; then from (52) we obtain (since v1 = v 2

(55) =

1)

Canceling [ZI], we find

[Z]= r

lCZpA2 -k k-lPB, klPAlk2PA2

Taking s1 = 2, s2

=

- k-1PB1k-2PB2

(57)

1, in accordance with (52)

Now we can cancel [Z]; this gives

[ZI]= r

klPA~ klPAlk2PA2

+ k-2PB2 - k-lPB,k-2PB2'

(59)


20 1

Adding (57) and (59), we obtain

whence, according to ( 5 9 ,

This deduction illustrates the method of application of (52) that also can be used in more complicated cases. The intermediate steady-state conditions lead to the same result (35). If both stages of scheme (53) are irreversible, then, having k - = k - 2 = 0 in (61), we obtain r=

k l P A l k2PA2 klPAl

+ k2PA2'

This equation can be written as follows:

llr = ( l / k l p A ~ ) + (l/kZpAz)*

(63)

It can also be easily obtained directly from the stages steady-state conditions. Indeed, applying these conditions to each stage, we have

Therefore,

[z] = r/klPAlr [a]= r/k2PA2. As a result of (55), we obtain (63) by summing these equations. In general, (51) and (52) are useful for the deduction of kinetic equations of such reactions, all stages of which (or at least a considerable part of those stages) are reversible. If all stages are irreversible, it is convenient to use stage steady-state condition (44). It is easy to see how the obtained kinetic equations are to be changed when, in a reaction analogous to that described by the scheme (53), the number of gaseous participants differs from that in the scheme. If, e.g., two species A, and A,' participate in the forward direction of the stage 2, there will be the product P A I P A 2 t , instead of PA2,in the kinetic equation. If species A2 is absent, 1 must be substituted for PA2,and so on. The steady-state rate of a many-stage reaction is not attained immediately;

202

M. I. TEMKIN

we shall consider the establishment of a steady state using a simple twostage catalytic mechanism as an example (36). Let the concentrations of A,, A2, B1, and B, be constant and let the coverage of surface [ZI] at a certain initial moment be not equal to the coverage that would correspond to a steady-state reaction at the given concentrations. A simple calculation (36) shows that the reaction rate, r, will change in time t , according to the law r = r,

+ (ro - r,)e-r’c

(66)

where r , is the steady-state reaction rate corresponding to (61), which is attained asymptotically (formally, at t = co). ro is the value of r at t = +O (i.e., at t = E where E is a positive value approaching zero) and z is the relaxation time of the reaction rate

Here L is the number of catalytic centers (i.e., sites for adsorption of particles I) on unit surface. The ratio r/L is called “turnover number” of a catalytic site. The reverse value u = L/r is of the dimension of time. It may be called “turnover time”; it is the average time required for a catalytic center to complete the cycle of transformations and to return to the initial state. The designation u corresponds to the German ‘word die Umschlagszeit [i.e., the time of turnover (of stock)]. On the basis of (61) and (67), it can be demonstrated that (68)

t 0 and T < 0 as it has been demonstrated by Zel’dovich (43) and the logarithmic isotherm corresponds to distribution (94) (40). Besides that, if we accept Assumption 1, then distribution (94) will give the equation of adsorption kinetics by Zel’dovich and Roginskii and distribution (93) will result in the equation of adsorption kinetics by Bangham. Finally, if Assumption 1 is correct, then distribution (93), including distribution (94) as its particular case, follows from the kinetics of fractional order reactions (44). The assumptions of the special model of a nonuniform surface were formulated above in terms of changes of the standard Gibbs energy, AGO, and the Gibbs activation energy, AG’. It can be assumed that standard entropy, So, of each kind of adsorbed particles and activated complexes on different sites of a nonuniform surface does not differ substantially; in this case there can be given practically equivalent formulation of the model

* Later the result of Polanyi was again obtained by Semenov.


21 1

involving (instead of AGO and AG’) the variation of internal energy change A U (which in an adsorbed layer virtually coincide with enthalpy change AH) and activation energy E. Then Assumption 1 (i.e., rule of transfer) takes the form of a linear relationship between E and AU, etc. To avoid the necessity of making assumptions about entropy, we prefer the formulation in terms of the Gibbs energy. Let the gas pressure that corresponds to adsorption equilibrium be denoted asp. The value o f p for the standard state of a site is denoted as b and is called “desorption pressure” of the site (Section IV). Each site of a nonuniform surface is characterized by a certain b value or by an adsorption coefficient, a = I/b (for a given temperature). At adsorption equilibrium, the probability that a site is occupied 0

= p/(p

+ b) = ap/(l + up).

(96)

It is often more convenient to consider not b or a but ( = In b = -In a ; it was proposed to call this value the “desorbtion exponent.” According to (25)

4; = AG,O/RT.

(97)

Let us number the sites of a surface of unit area in the sequence of increasing b or ( (i.e., in the sequence of decreasing adsorbtion power). We shall denote the number of a site as I ; its ratio to the total number of sites on a section of unit area,

(98) will be called “relative number” of the site. The value of s varies from 0 to 1. Experimentally observed quantities pertaining to the whole surface, such as the amount of adsorbed substance, heat of adsorption, reaction rate, are sums of contributions of surface sites or, since the number of sites is extremely great, the respective integrals. As 5 increases monotonously with s, each of them can be taken as variable for integration; both methods of calculation are used. If 4; is chosen as an independent variable, a differential function of distribution of surface sites with respect to desorption exponents, (p((), is introduced such that p(()d( is the number of sites with AG:/RT values within the limits of ( and ( + d5. Let us denote the largest and the smallest values of 4; as (,, and C1, respectively, and introduce the notation s = l/L,

y

= T/O.

(99)

Then by virtue of the second assumption the function ~ ( 4 ; )is expressed as follows :

212

M. I. TEMKIN

at

to < 5 < t1

at 5 > l

p(t) = AeYC

445) = 0

1

where A is a constant. The value of A is determined by the normalization condition

where L is the whole number of sites on the unit surface. According to (100) and (102) the condition (103) may be written thusly:

Substituting expression (101) with y # 0, we obtain A = yL/(eYt' - eyto).

(105)

At y = 0, i.e., at an evenly nonuniform surface, p({) = A, and (104) gives '4 = U(5, -

50).

(106)

The same result is obtained from (105) evaluating the indeterminate form in the usual manner. Some statements referring to nonuniform surfaces are valid irrespective of the form of the distribution function of surface sites with respect to their AG,' values. Using the relative number of site, s, as an independent variable, we obtain the following expression (40) for the equilibrium surface coverage, 8:

where a is a function of s (monotonously decreasing). Let the equilibrium gas pressure, p, be so small that a,p

.

(122)

This value is called the nonuniformity exponent. At high p values, when a$ 9 1 and a l p 9 1, unity may be neglected as compared to aop and a$ and then from (121) 8 = l/fln(ao/a,), and from (122) 8 = 1 (i.e., the surface is saturated). In contrast, at low p , when a$ Q 1 and a# 4 1, we expand In( 1 + a d ) and In( 1 + a l p ) in series and, restricting to the first terms, obtain 0 = (a, - a l ) / f ) p (i.e., the Henry law). If a surface is strongly nonuniform (ie., a, B a,), then a range of p values exists where a,p p 1 and simultaneously a l p 1. This means that for the most strongly adsorbing sites the probability of being occupied is almost 1 and for the most weakly adsorbing sites it is almost 0. This will be called the region of medium coverages. We obtain for this region from (121)

+

8 = ( 1I f ) Waop).

(123)

According to this equation, 0 depends linearly on In p ; therefore, we call (1 23) “logarithmic adsorption isotherm.” This isotherm, as it was mentioned in Section IX, describes experimental data well in many cases.


215

Isosteric molar heat of adsorption, q, is determined by the equation

(a InplaT),

=

q/RT2.

(124)

Sincefdepends on T according to (1 17), it follows from (123) that q depends on 8 linearly: q = 40 - RTf8

(125)

where q,

=

RT2[dIn (l/ao)/dT].

( 126)

Thus qo is the heat of adsorption on a uniform surface with adsorption coefficient a,. According to (1 25), nonuniformity exponent may be regarded as the ratio to RT of the slope of the straight line in the plot of q as function of 8. For strongly nonuniform surfacesf is much greater than 1. If a homonuclear diatomic molecule (e.g., H, , 0, , N, , etc.) is adsorbed with dissociation into atoms, the adsorption isotherms (121) and (123) give 8 as the functions of partial pressure of atoms in gas phase, pat,that would correspond to dissociation equilibrium at the partial pressure of diatomic molecules, p. These values are linked by P 2 P = K9 where K is the dissociation constant. Thus, (Kp)'', must be substituted in the equation of isotherm for p. Instead of (123) we obtain

9 = (1/2f) WOKP) (i.e.,the linear dependence of 8 on Inp is retained for dissociation adsorption). To derive adsorption isotherm at the exponential nonuniformity as described by (100)-(102) at 1 > y > 0, it is convenient to use the desorption exponent, t = -In a, as integration variable. The probability of a site being occupied is B = ap/(l up) = 1/(1 + l/ap), and at adsorption equilibrium

+

For the distribution function considered

Let us introduce the symbol

216

M. I. TEMKIN

Inasmuch as u = (1 - a)/a, the meaning of u is the ratio of the probabilities for a site to be free or occupied. According to (129),

d< = du/u

(130)

and since =

(128) takes the form

Y Y

P U ,

1

(131)

l h p

e = ( I / L ) ~ Y1 loop [uY-'/(1 + u)]du.

(132)

This integral generally cannot be expressed through elementary functions. Because of it, we consider medium coverages only when alp 4 1 and a,p % 1. In this case 0 and 03 may be substituted for the limits of the integral without an appreciable error. We assume that 0 < y < 1. It is known that at 0 < h < 1 Jom

+ u)]du = n/sin hn.

[uh- '/( 1

(133)

As etl = l/al, eco= l/a, and a, % a,, we may neglect eyro compared to (105). Then AIL = yaIYand, hence

eYc1in

t? = (yn/sin yrt)(alp)Y.

(134)

e = cpy

(135)

Therefore

where

C = (yn/sin yn)aIY.

( 136)

Equation (135) is the well-known Freundlich adsorption isotherm. In a number of instances this isotherm accurately describes experimental data. The interpretation of the Freundlich adsorption isotherm as resulting from exponential nonuniformity of surface is due to Zel'dovich (43). A kind of nonuniformity of surface with negative y is also conceivable. If -1 y < 0, then a similar derivation gives for the region of medium coverages the following adsorption isotherm (44) :

-=

e = 1 - (C/rlYI)

(137)

where Iy( is the absolute value of y. To derive equations of the rates of adsorption and desorption we shall assume that the rate of migration of adsorbed molecules along the surface


217

is sufficiently high for the equilibrium within the adsorbed layer to exist even when the layer is not in equilibrium with the gas phase. Then the probability for each surface site to be occupied is determined by a certain value p that is one and the same for the whole surface, p being the pressure in the ideal gas phase that would be in equilibrium with the surface at the given coverage. This value will be called fugacity of adsorbed substance. When there is no equilibrium with gas phase, p differs from the actual gas pressure, which we shall denote as P ; if p < P, the amount of adsorbed particles increases with time; if p > P, this amount decreases. We shall consider the rate of change of number of adsorbed particles on unit surface, r ; it may be positive or negative, as it is the difference between the rates of adsorption (i.e., “condensation”),r + ,and desorption (ie., “evaporation”), r - : r

=

r+ - r - .

(138)

On a uniform surface Y+ =

K + p ( 1 - 6) = (K+P)/(l

+

Up)

(139)

where K + is the adsorption rate constant. Thus, the contribution of one + up), site of a nonuniform surface to the rate of adsorption is (l/L)(~+p)/(l and the rate of adsorption on a nonuniform surface is

Since we assume the linear relationship between AG+* and AG,’, (91), we obtain for a nonuniform surface, on the basis of (25) and (31), K + = K+Oe-a(C-Co)

(141)

where K + ’ is the value of K at ( = to(i.e., on the most strongly adsorbing sites). Applying (101) for q(l),we obtain from (140) and (141) r+

(142)

=

Equation (142) contains a new symbol, m : m=u-y.

(143)

We now turn to the integration variable u determined by (129). With the help of (105) we find that y+

=-.

(144)

218

M. I. TEMKIN

where n=1-m.

(145)

For the region of medium coverages we approximate the integral substituting integration limits 0 and co for uo and u l . Making use of (133) under the condition that 0 < n < 1 (and hence, 0 < m < 1) and bearing in mind that t1 - to= f and sin nn = sin mn, we obtain

Y K+OP n r + =-*-. sin mn eyf - 1 (sop)"' Equation (146) is also valid at y = 0 if, for the factor y/(eyJ- l), its limit at y = 0 (i.e., l/f) is substituted; additionally, at y = 0, m = a. Therefore, for evenly nonuniform surface we have

Thus, the special model of a nonuniform surface that we consider gives for r , the expression r + = k,P/p"

f 148)

where k , and m are constants; m satisfies the inequalities 0 < m < 1. Subscript ''&" is used at k because, as will be seen later, the expression for desorption rate, r - , contains the same constant, k , . It is evident that at low coverages the adsorption rate is proportional to P and is independent of p ; this result can be obtained from (148) with m = 0. For large coverages the probability that a site is not occupied is proportional to l/p; therefore, the adsorption rate is proportional to P / p ; this is obtained from (148) with m = 1. Thus, (148) can be regarded as valid for 0 < m < 1. Equations (146) and (147) express the dependence of the rate of adsorption on the fugacity of the adsorbed layer, p . In order to obtain r + as a function of 8, p should be expressed in terms of 0 by means of a corresponding adsorption isotherm. For an evenly nonuniform surface, (123) and (147) give Y+

= (n/sin an)(K+/f)Pe-"f'

(149)

or

r + = k+Pe-"f'

( 150)

where k + = (n/sin an)(lc+O/f).

(151)


219

Equation (150) is the well-known equation for adsorption rate found by Zel'dovich and Roginskii (45) (sometimes it is erroneously called the "Elovich equation"). This equation was experimentally confirmed for many cases of chemisorption. It follows from (135) and (146) that on a surface where the Freundlich isotherm is valid, the adsorption rate is proportional to P/Om. Inverse proportionality between r + and a fractional power of 0 was found by Bangham (46). Equations for desorption rate, r - , could be obtained by calculations similar to those used for adsorption rate, r + , but there is a simpler method based on the following considerations that are analogous to those developed in Section VIII. On a uniform surface at gas pressure P and fugacity of the adsorbed layer p , adsorption rate is r+

=

+

(152)

+ up).

(153)

K + P (-~0) = K + P / ( ~ up)

and desorption rate is r- = K - 0

=

K-ap/(l

Hence

r-b+ Since at equilibrium p

=

= (K-/K+)(UP/P).

(154)

=u

(155)

P, K+/K-

and, consequently, r-/r+

=

p/P.

(156)

This equation holds for the contributions of each site of a nonuniform surface to adsorption and desorption rates, and hence, for total r+ and r values on a nonuniform surface. Therefore, (148) with n taken from (145) gives directly r-

=

kip".

(157)

It may seem strange that (148) and (157) contain the same constant k , since the comparison of (148) and (146) shows that k,

=

(n/sin mn)y/(eYf - l)[~+'/(a~)'"]

(158)

(i.e., the expression for k , contains constants K+' and m that describe the process of adsorption). But since m + n = 1, sin mn = sin nn and, accord= uo,so (158) is equivalent to the following expression: ing to (155),K+'/K-'

k,

=

(n/sin nn)y/(eYf- l)lc-'(u,)".

(159)

220

M. I. TEMKIN

Equation (159) contains constants rc-' and n that describe the process of desorption. For an evenly nonuniform surface we express p in terms of 0 by means of (123). Introducing the frequently used symbol j?=l-Cr,

(160)

we have in this case n = B and r-

=

(n/sin fin)(.- '/f)ePfe

(161)

or r- = k-esfe

where

k-

= (n/sin Bn)(rc-'/f).

Desorption rate depends exponentially on surface coverage [i.e., according to (162)], as Becker and Langmuir observed experimentally. It can be easily seen that if adsorption equilibrium follows the Freundlich isotherm, then r- will be proportional to 8". According to (143), (145), and (160), n=j?+y.

(164)

Let us consider now the case (47) when there are molecules of two different gases on the surface, the ability to be adsorbed not differing greatly for the two kinds of molecules; we shall make use of Assumption 3 formulated in Section IX (i.e., assume that the change in the standard Gibbs energy of adsorption at passing from one surface site to another is the same for both kinds of molecules). If gas A is being adsorbed from a mixture with gas A , the probability that at adsorption equilibrium a certain surface site is occupied by a particle A is a=

UP

1

+ up + a'p'

where p and p' are the equilibrium pressures of gases A and A , a and a' are their adsorption coefficients for this site ; i.e., the constants of the equilibria Z+A=ZA

(166)

and Z

+ A' = ZA'.

(167)


22 1

We shall denote the equilibrium constant of adsorption displacement ZA 4- A = Z A i- A

(168)

as c. Since (168) is the difference between (167) and (166),

aria = c .

(169)

According to the assumption made, c is identical for all surface sites. Thus, adsorption coefficients of gases A and A are mutually proportional. Equilibrium surface coverage by gas A is apds

1

+ up + a’p’

or, since in the case of proportionality of adsorption coefficients c is independent of s, P

The comparison of this equation with (107) shows that the integrand of (171) can be obtained from the integrand of (107) by substitution of p cp’ for p . Therefore, if (107) leads to some adsorption isotherm of a pure gas A,

+

0 = F(p),

(172)

where F is a function determined by the form of dependence of a on s, then for the adsorption of A from a mixture with A we have

It can be easily seen that the equilibrium surface coverage by gas A

These results can be generalized naturally for the case of adsorption equilibrium of more than two gases with proportional adsorption coefficients. Thus, the passage from an adsorption isotherm of a pure substance to the corresponding adsorption isotherm of a mixture is very easy, supposing that the model of a nonuniform surface is applicable and adsorption coefficients are proportional. If, for instance, adsorption of pure gas A is described by the Freundlich isotherm (135), then for adsorption of A from mixture

222

M. I. TEMKIN

with A', according to (173), we have

P

e=c

(p

+ cp')' -

y'

(175)

Rate of adsorption, r + , of gas A from mixture with gas A will be considered on the assumption of high surface mobility. This assumption allows us to consider the probability for a site to be free as determined by the fugacities, p and p', of the substances A and A in the adsorbed state, the fugacities being identical for all sites. Then

u+Pds 1 u p u'p'

+ +

where P is the pressure of gas A, K + is the adsorption rate constant of A, which is a function of s. In accordance with (169) r+ =

jo+ 1

Ic+Pds u(p cp')'

+

(1771

If adsorption coefficients are proportional (i.e., if c is independent of s), this equation differs from the equation for the rate of adsorption of gas A in the absence of A only by substituting ( p + cp') for p . It follows that if for pure gas A r+

=

m P )

(178)

where G(p)is some function of p ; then for adsorption from a mixture with A' r + = PG(p

+ cp').

( 179)

If, for example, (148) holds for r+ in the case of adsorption of pure gas A, then for adsorption of A from a mixture with A' r+

=

k,P/(p

+ cp')".

(180)

In a similar manner from r- =

'

K-upds

1

+ u p + alp"

if the rate of desorption from a surface containing only particles of A is r- = H(p)

(182)

where H ( p ) is a function of p , then for desorption from the same surface


223

containing particles of A and A‘

It follows from (173) and (174) that

e + el = ~ ( +p cp’). Hence, p

+ cp’ =

+ el)

~ - 1 ( e

(185)

where F - ’ is a function inverse to F. Substituting this into (179), we have r+ = PG[F-’(d

+ el)].

(186)

Equation (186) means that the rate of adsorption of A from a mixture of A and A depends on the total coverage of the surface by both gases in the same manner as it depends on the coverage of the surface by gas A when only A is being adsorbed. Thus, in the case of adsorption of A and A on an evenly nonuniform surface we have for the rate of adsorption of A, instead of (150), = k+pe-am+u (187)

Xi.

The Kinetics of Reactions on Nonuniform Surfaces

The larger the decrease in the Gibbs energy at a chemical reaction, the larger the accompanying loss in the ability to perform a useful work, and hence the less perfect the process is from the viewpoint of thermodynamics. Therefore, it seems not to be accidental that in commercial large-scale processes reactions are utilized that occur at conditions close to equilibrium [i.e., reactions with small (-AG) values]. Such are the industrial processes of ammonia synthesis, sulfur dioxide oxidation, hydrogen production by reactions of methane and carbon monoxide with steam, methanol synthesis, ethylene hydration, stepwise butane dehydrogenation to butadiene, and a number of other reactions. If a system that is in equilibrium with respect to all stages of a complex reaction is being shifted to states more and more distant from equilibrium by gradual variation of concentrations of substances participating in the reaction, then, since the rates of elementary reactions may differ to any

224

M. I. TEMKIN

extent, it is natural to expect that not all the stages will become nonequilibrium simultaneously. First of all, it will happen to the stage whose elementary reactions are the slowest. It is not surprising, therefore, that the rate of single-route reactions at conditions near equilibrium is determined, as a rule, only by one stage; other stages are at equilibrium. In such cases, as the numeration of the stages of a cyclic steady-state reaction is arbitrary, we shall ascribe the number one to the rate-determining stage; equilibrium stages, if there are several of them, will be united into one common complex stage, the number of which will be two. We shall confine ourselves to the consideratiofi of a simple two-stage mechanism when the surface contains in a significant amount only intermediate particles of one kind, the degrees of coverage by other particles, if there are any, being small. Such mechanism may be described by the following scheme (48): I. 2.

+ Z $ B1 + ZI a2A2+ a2’A2’+ ZI = b2B2 + b2’B2‘ + Z A, + a2A2+ a2’A2’= B, + b2Bz + b2’Bz‘ A,

1 1

(188)

Here A,, A,, etc. are substances, and a,, etc. are stoichiometric coefficients. The sign = denotes equilibrium (or quasi-equilibri~m).~ The reaction direction, taken as forward, in particular cases may correspond to either the forward or reverse direction of stage 1 of scheme (1 88). Therefore, we distinguish the directions of stage 1 as follows. The direction that results in the occupation of a free site on the surface is called the “adsorption direction,” and the direction that results in a site becoming unoccupied is called the “desorption direction.” The rate of stage 1 in adsorption direction [i.e., in the forward direction in our record (188)] is denoted as r, and the rate in desorption direction [i.e., in the reverse direction in scheme (188)] is denoted as rB. When applied to concrete reactions one of these values will stand for the forward reaction rate, r , , and the other will stand for the reverse reaction rate, r - . Transfer coefficients for adsorption and desorption directions will be denoted as a and p, respectively; so a+P=l. Adsorption and desorption processes are particular cases of stage 1; namely, when substance B, is absent, and I coincides with A , , r, is the rate of adsorption and r, is the rate of desorption. Since the equilibrium of stage 2 is maintained as a result of mutually reverse elementary reaction, the particles I very frequently leave some surface sites and appear on others. (the sign of constellation Previously the author had proposed a more expressive sign, Libra or balance), for this purpose. It turned out, however, that its use caused difficulties in typesetting.


'

225

Consequently, they are distributed on the surface randomly even if there is no surface migration. Therefore, in contrast to the discussion of adsorption and desorption processes in Section X,the mechanism (188) needs no assumed rapid surface migration (it is evident, of course, that if the migration occurs, it does not affect the results). The analysis of adsorption and desorption rates given in Section X needs only minor alterations to its application to stage 1 ; namely, products KBPB, would be substituted for rate constants of desorption on separate surface sites, IC- , and the fugacity of adsorbed particles I, p I for p . Therefore, in analogy to (148) and (1 57), we obtain

where k , and k , are constants and m and n, as before, are defined by (143) and (164) (i.e., m = CI - y and n = /3 y, so that m n = 1). At equilibrium r,, = r,; since m n = 1 , this gives k A / k , = pB&/PA1. Hence

+ +

kA/kB

=

+

KA

(191)

where k Ais the equilibrium constant of the reaction of formation of particles I from A in the gas phase :

,

A,

= B,

+ I.

(192)

The equilibrium of stage 2 can be considered as the difference of gas phase chemical equilibrium

I

+ a2A2 + a2'A2' = b2B2 + b,'B,'

( 193)

and adsorption equilibrium

I

+ z = ZI.

( 1 94)

The term "adsorption chemical equilibria" has been suggested for equilibria such as those shown by stage 2; they can be observed directly by experimentation (49-52). Denoting the equilibrium constant of (193) as K,,we have whence

226

M. I. TEMKIN

and rB = (kBIK?) P, ,(Pg2Pg2',/p"A',PA';, )".

( 197)

Equations (196) and (197) are valid under the condition that 0 I m 4 1 or under the equivalent condition that 0 I n 5 1. The results obtained indicate the interpretation of fractional reaction orders frequently obtained in experimental studies of heterogeneous catalytic reactions. If in the course of a reaction, not one, but several adsorption chemical equilibria exist, it is always possible to write the chemical equations of the equilibria in such a way that the stoichiometric number of one of these equations be l and that of other equations be zero. Confining ourselves to two equilibria, we have the following mechanism instead of (1 88) (48): 1.

+ Z 3 B, + ZI a2A2 + a2'A2' + ZI = b2B2 + b2'B2' + Z a3A3 + a,'A,' + Z = b3B3 + b,'B,' + ZJ A,

1

'0

2. 3.

A,

1

0

+ a2Az + a2'A2' = B, + b2B2 + b2'Bz'

We suppose that only substances I and J occupy essential parts of the surface; then the form of equations for r, and rB results from the equation for the simultaneous adsorption of two gases (1 80). Let us consider now the kinetics of a reaction following the simple twostage mechanism (53) on a nonuniform surface with none of the stages of the reaction being assumed to be at equilibrium (53).Since, in contradistinction to scheme (188), both stages of scheme (53) are equivalent, we may always ascribe number one to the stage whose positive direction is of the adsorption type. Therefore, without loss of the generality of the analysis, we can distinguish the directions of the elementary reactions in scheme (53) as forward and reverse. Accordingly, the rate constants of the elementary reactions of stage 1 will be denoted as K~ and K - ~and those of stage 2 as K~ and K - ~ For . brevity we introduce A = - t o .Its largest value is 5 0 ="f We accept the simplifying Assumption 4 of Section IX (i.e., assume that the transfer coefficient c1 is identical for both stages). Then, from the linear relationship between AG+ and AG,' for each stage and bearing in mind that rate constants K , and K - ~refer to adsorption directions, and K - ~ and K~ refer to desorption directions, we obtain, in a manner similar to

Ni > Co > Pd > Cu > Fe,O, z Ag > Au.

For nickel, activation energy determining the temperature dependence of = 12 kcal/mol; at 150°C k, x 2 x lo6 mol/atm m2 sec. The kinetics of tritium exchange between steam and hydrogen

k,, E ,

HTO

+ H2 = H2O + HT

(364)

was investigated on porous nickel at 1.65-200°C and nickel supported on chromium oxide at 100- 132°C (Z35). An equation corresponding to (359), i.e., r = k + PnTo(Pn2/PH2oIrn - k -Pn~(Pn201Pn2 m, was verified with m = 0.5.

(365)

270

M. 1. TEMKIN

In handling experimental results the constant of equilibrium (364) was calculated on the basis of the literature data (136) substantiated by the measurements made by the authors (135). The stages of isotopic exchange mechanism (357), if we do not distinguish between isotopes H and D, coincide with the forward and reverse directions of stage 1 of mechanism (343) of carbon monoxide conversion. The reactions of isotopic exchange corresponding to stage 2 or this mechanism

co,

+ COi8 = co + coo'8

(366)

and

c1402+ co = c140 + co,

(367)

can also be observed. The kinetics of reaction (366) was investigated on Fe,O, (137) and that of reaction (367) on other catalysts as well (138). On Fe,O, the rates of isotopic exchange reactions (356), (366), and (367) are close to the. rate of carbon monoxide conversion, as should be expected from mechanism (343).

XIX.

Phosgene Synthesis

Phosgene is used in the production of diisocyanates, starting materials for polyurethanes, and in other organic syntheses. Commercially, phosgene is obtained from carbon monoxide and chlorine

co + c1,

=

COCI,

(368)

with activated charcoal as a catalyst. The reaction is conducted at 80- 150°C (temperature in various parts of the commercial reactor is different) and at pressures near to atmospheric. An experimental study of kinetics of reaction (368) on activated charcoal (139) preceding our investigation was carried out at 30-65°C (i.e., at temperatures below those used in industrial reactors). This promoted us to reinvestigate the kinetics of this reaction (140). The utilization of a flow circulation system for laboratory experiments was made difficult by the toxicity of all three gases, CO, Clz, and COCl,. For this reason, in contrast to other reactions, we studied reaction (368) in a closed system. The experimental apparatus was equipped with a circulation pump so that the gas mixture continuously passed through the catalyst bed in the reactor. The reaction rate was determined from the change in total


27 1

pressure, P,with time, t. The following formula was used in the calculations :

+

dP(t) - - - -2P(t - 2h) - P(t - h) P(t dt 1Oh

+ h) + 2P(t + 2h)

(369)

where h is a small time interval between successive measurements of P. This formula corresponds to the least-squares polynomial of the second degree interpolation of five experimental points (141). The experiments were done at 70, 100, and 130°C and at pressures somewhat lower than atmospheric. Under these conditions reaction (368) is practically irreversible. Activated charcoal of the trademark "Bayer AKT-4" ground to grain size 0.25-0.5 mm served as a catalyst. Estimation of the efficiency factor on the basis of the determination of the effective difusion coefficient of hydrogen in nitrogen or helium has shown that for this grain size the results of reaction rate measurements refer to the kinetic region. Estimation of relaxation time of the reaction rate from (67) showed the reaction to be quasi-steady at the condition of our experiments in the closed system. Activated charcoal is a hydrocarbon rather than an elemental carbon because it contains hydrogen chemically bound to carbon. We observed that when exposed to chlorine, activated charcoal evolves HC1; the substitution of hydrogen by chlorine, i.e., the reaction of the type RCH

+ CI,

=

RCCl

+ HC1

(370)

probably occurs. Analysis after experiments in which charcoal served as a catalyst of reaction (368) demonstrated that this material contained 23 wt. % C1 corresponding approximately to the composition CloC1. Specific surface of the charcoal measured by low temperature adsorption of nitrogen was 1000 m2/g before the kinetic experiments and only 200 m2/g after the experiments. It may be conceived that reaction (370) causes the closing of a part of the micropores. The decrease in the catalytic activity of fresh charcoal that is observed during reaction (368) may also be related to the process (370). After 15 hours' exposure of charcoal in chlorine at 13O"C, the catalytic activity becomes constant. Kinetic measurements were carried out with a catalyst prestabilized in this way. Experimental data on the reaction kinetics agree with the kinetic equation for a reaction on a nonuniform surface in the region of medium coverages if the mechanism is as follows: 1.

2.

z+coF?zco zco + c1, F? z + coc1, co + CI,

=

COCI,

1 I

(37 1)

272

M. I. TEMKIN

Scheme(53) Scheme(371)

A, CO

A,

B,

C1,

-

Bz

co

I

co

Since B, is absent, I coincides with A,. At first sight, scheme (371) does not agree with the results of our adsorption experiments; these experiments showed that activated charcoal does not chemisorb CO at 100°C. It should, however, be taken into consideration that the surface of charcoal subjected to activation or even simply after storage in contact with air is covered with chemisorbed oxygen. The studies of the reactions of carbon with C 0 2 and steam (see Section XX) have demonstrated that oxygen chemisorbed on carbon is indistinguishable from chemisorbed carbon monoxide. So it may be reckoned that activated charcoal is already covered with carbon monoxide before the contact with this gas. Desorption of CO from the surface of such charcoal occurs only at temperatures much higher than those employed in phosgene catalytic synthesis; therefore, stage 1 of scheme (371) is irreversible under the usual conditions of the reaction. Kinetic measurements show that the reaction is retarded by its product, phosgene. To account for this in terms of scheme (371), stage 2 must be considered as reversible. In such a case (214) is directly applicable. After the substitution specified above we obtain r = k P c o [ ~ c , , / ( A P c+ o ~coc~,)lm.

(372)

Experimental data gave m = 1/4 at the three temperatures (i.e., 70, 100, and 130°C). According to (99) and (143), m = u - (T/O)

(373)

where a and 0 are temperature-independent; therefore, if m # u, m depends on T. However, this dependence is hardly noticeable. Let us, for instance, accept for the transfer coefficient a the most usual value, 0.5, and assume that m = 0.25 at 100°C. Then we find that 0 = 1490°K; hence, at 70"C, m = 0.27 and at 130"C, m = 0.23. The deviation of m from the average value 0.25, by k0.02 is difficult to detect. Taking m = 0.25, we found numerical values of constants in (372) from experimental data for activated charcoal Bayer. These constants are described by the equations

+ 3.40 = - 1875/T + 3.80.

logA = -1900/T

(374)

log k

(375)

273


Values of k from (375) give the phosgene formation rate, in moles per second per gram (200 m2) of charcoal at pressures expressed in atmospheres. According to (379, the apparent activation energy, E = 8.6 kcal/mol. Equation (372) with rn = 0.25 and A found from (374) agrees well with experiments not only for the charcoal Bayer, but also for the Soviet activated charcoal AR-3. In the temperature range covered by the experiments by Potter and Baron (139)one may have A = 0 as follows from (374). This means that stage 2 of scheme (37 1) is virtually at equilibrium under these conditions. Then, instead of (372), we have the following equation corresponding to (21 5) : r = kPc,(Pc,,/Pc,,,,)".

(376)

Experimental results of the referenced work agree with (376) ;namely, points fall into straight lines being plotted in the log(r/Pco) vs. log(PcI2/Pc~~,) coordinates. Slopes of the lines for most of experimental series give m values near to 1/4. XX.

Reactions of Carbon with Carbon Dioxide and Steam

The reactions of coal gasification by carbon dioxide

c + co2= 2co

(377)

and by steam

C + (1

+ s ) H ~ O= (1 - s)CO + s C O ~+ (1 + s ) H ~ ;

0 < s < 1 (378) in combination with carbon monoxide conversion (the reaction discussed in Section XVII) were widely used as a source of hydrogen in the production of synthetic ammonia. At present this method is excluded by a more economical process of methane reforming (considered in Section XIV). The exhaustion of natural gas resources may, however, restore the importance of coal gasification. In addition, reaction (377) is a constituent part of the blast furnace process. Reaction (378) is employed in the production of activated charcoal. Reactions (377) and (378) are usually not ranked among catalytic processes, although they are, as a matter of fact, autocatalytic since they occur on the surface of carbon, one of the participants of the reactions. True, it is more usual to speak of autocatalysis for reactions catalyzed by products rather than by reactants, but it is not essential in our case. Equations (377) and (378) may be read from right to left as the reactions are reversible. Coal gasification is accelerated by the components of ash such as potash and iron oxide; sometimes, to promote gasification, catalytically active

274

M. I. TEMKIN

substances are added. Here we will consider only the reactions of pure carbon with carbon dioxide and steam. The kinetics of these reactions is closely related to the kinetics of some of the reactions discussed above. As a result of their investigations of reaction (377), Frank-Kamenetskii and Semechkova (142, 143) advanced a mechanism that can be presented as follows : 1. 2.

z c + c o 2 # z c o +co zco # z + co c + co,

=

1 1

(379)

2CO

In scheme (379) ZC as well as Z denotes a site on the surface of carbon; in the symbol ZC we only accentuate one of the C atoms; viz., the atom that binds chemically to the 0 atom. Surface formation ZCO, according to Frank-Kamenetskii, can be equally well considered as a chemisorbed 0 atom (on ZC) or as a chemisorbed CO molecule (on Z ) . The equivalency of ZC and Z being taken into account, scheme (379) becomes a particular case of a simple two-stage scheme (53). The correspondence is specified by the following table: ~~

Scheme(53) Scheme(379)

A, C02

B, CO

A, -

~

B,

co

I

co

In deriving a kinetic equation for mechanism (379), Frank-Kamenetskii assumed the surface to be uniform. The kinetic equation answering to this assumption is obtained immediately from (61):

If the reaction is conducted under such conditions that it is practically irreversible, but is not too far from equilibrium, then in accordance with considerations in Section XI, one of its stages must be irreversible; namely the slowest one. Therefore, one of the rate constants of backward reactions, k - , or k-2, should be taken as equal zero. Assuming k-2 = 0, we obtain from (380)

and assuming k - l

=

0, we obtain


275

Equations (381) and (382) are kinetically indistinguishable. Dividing the numerator and the denominator by k,, we reduce them both to the equation

where I,, I,, and I, are constants. An equation of the form of (383) was proposed after the works by Frank-Kamenetskii also by Gadsby et al. ( 1 4 4 , who derived it from notions about the mechanism of the reaction somewhat different from those of Frank-Kamenetskii, but still assuming the surface to be uniform. Equation (383) is in qualitative agreement with the experiment. Reaction (377) is indeed retarded by its product, carbon monoxide, even when the concentration of the product is smaller than equilibrium concentration to such an extent that the reverse reaction is virtually absent. This was verified also by our experiments described below. According to the meaning of (381), the retardation of the reaction by its product results from the reversibility of stage 1 ; (382) implies that the reaction is retarded by adsorption of the product. The correct choice can be made by the comparison of the reaction rate and the rate of exchange of Oi8 between CO, and CO, since of the two reaction stages in scheme (379), only the stage 1 (when it proceeds forward and backward) results in this isotopic exchange. Our investigation of the kinetics of reaction (377) (145) was done on a practically ash-free carbon sample. The carbon was obtained by kilning Bakelite prepared from chemically pure phenol and formaldehyde with subsequent activation with carbon dioxide. Specific surface area of the carbon measured by adsorption of methylene blue was 2.2 mz/g for the initial sample and 3.7 m2/g for the sample that reacted with COz for 155 hours. Surface area measurements by low-temperature adsorption of nitrogen showed a much greater increase in specific surface with the burning off of carbon; namely, from 24 m2/g for fresh carbon to 300 m2/g for carbon after 155 hours of reaction with CO, ,presumably, as a result of the opening of the orifices of micropores. Kinetic experiments demonstrated that the rate of the reaction (377) is proportional to the specific surface measured by adsorption of methylene blue. This means that the reaction occurs mainly on the surface of wide pores accessible to the molecules of methylene blue and that the surface of micropores takes virtually no part in the reaction. The experiments were conducted in the temperature range of 700-780°C. The degree of CO, conversion in these experiments did not exceed 3%; CO concentration at the outlet comprised only 5-6% of its equilibrium content under the conditions of the experiment. Therefore, in discussing the results, the reaction could be considered irreversible.

276

M. 1. TEMKIN

The size of carbon grains in the kinetic experiments was 1-2 mm. Preliminary tests with grains of different sizes had shown that this size range ensured the reaction course in the kinetic region. The same carbon that was used for reaction (377) was employed in the measurements of the rates of the reactions COO'S

+ co = coo + C0'8

(384)

coo'8+ co

(385)

and

coo + CO'8

=

in the temperature range of 550-570°C. The concentrations of carbon monoxide and carbon dioxide corresponded in this case to equilibrium with respect to reaction (377). The measurements were conducted in a closed circulation system. The rate of isotopic exchange is proportional to the difference of contents of labeled atoms in the exchanging molecules. The factor of proportionality was proposed (145) to be named the "total rate of exchange" since it is equal to the rate of the continuous exchange of atoms that cannot be detected unless the atoms are labeled. The rate of isotopic exchange would be equal to the total rate of exchange if all atoms of an element in one of the species were labeled and those in the other species were normal. Total rates of oxygen exchange between CO, and CO calculated from the rates of reactions (384) and (385) were compared to the forward rate of reaction (377) in the equilibrium state found by extrapolation of measured rates by means of an empirical kinetic'equation. It turned out that at 750°C the total rate of exchange exceeds the rate of reaction by a factor of 50 and at 700°C more than of 100. The difference in rates is so large that it cannot depend on the inexactness of measurement of the rates of isotopic exchange or on the uncertainty associated with the extrapolation of the reaction rate to equilibrium. It follows that stage 1 of scheme (379) is much faster than stage 2. Therefore, if one of the reaction stages becomes irreversible as a result of going away from equilibrium, it is stage 2 and not stage 1. This means that (381), but not (382), is qualitatively correct. The comparison of our measurements of the rate of reaction (377) with (383) showed that this equation is not obeyed quantitatively. The disagreement of (383) with experimentation already noted in preceding works can be explained by the nonuniformity of the surface not being taken into account. To describe the rate on a nonuniform surface under the conditions of practical irreversibility of the reaction, (214) can be applied. It should, however, be remembered that this equation corresponds to the mechanism with the irreversible first stage, and in our case it is the second stage that is irreversible. Therefore, in order to make use of (214), the numbering of


277

stages should be altered; stage 2 of scheme (379) should be regarded as the first stage and stage 1 as the second stage. Since A = ic1/x2,A - should be substituted for A . We obtain r = "[Pco,/(A

+ PC0)I"

(386)

or where k = [ k ] A m . Equation (387) at rn = 0.5 agrees quantitatively with experimentation (145). Constants in (387) are described as functions of temperature by the equations k = 3.25 x 105e-48*000/RT (388) A

=

5.3

10-11eS8.S00/RT

(389)

Here the rate of reaction, r, is expressed in units (mole CO,)/mZ hr (specific surface area of carbon being measured by adsorption of methylene blue); pressures are in atmospheres. A number of authors (246-150) who studied the reaction (378) concluded that the kinetics of reaction (378) follows the equation r(l,

=

'lPH20

1

(390)

+ i2PH2+ i3PHzO

where r(') is the rate of entering of carbon into the reaction; i , , i, , and i, are constants. This equation is analogous to (383). As in (383), it corresponds to the assumption that the surface is uniform. It is natural to assume on the basis of mechanism (379) of the reaction of carbon with carbon dioxide and the mechanism of carbon monoxide conversion, (343), that the reaction of carbon with steam occurs as follows: 1. 2. 3.

+

Arc')

+ H2 zco ?3z + co z c o + c o * c o 2 +zc ZC

MI); C

H 2 0

P ZCO

1 1. 0

"2'

1 0 1

(391)

+ H 2 0 = CO + H2 + H 2 0 = C 0 2 + H,

N'"; CO

Stage 1 of this mechanism is analogous to stage 1 of mechanism (379), stage 2 of both mechanisms are identical, and stage 3 coincides with stage 1 of mechanism (379) written in the reverse direction. On the other hand, stages 1 and 3 conform with the mechanism of carbon monoxide conversion (343). Such a mechanism of reaction (378) was proposed by Key (15Z).

278

M. I. TEMKIN

The reaction is described by two basic routes; in (391) they are chosen so that carbon enters the reaction only in route Mi), and route N”)results in the conversion of a part of CO formed in route N(’)into CO, . With such a choice of the basis of routes, r(l), the rate of the reaction along route N(’)is the rate of carbon gasification; i.e., corresponds to symbol r(l) in (390). Our study of the kinetics of the reaction of carbon with steam (152) was conducted by the circulation flow method under atmospheric pressure in the temperature range of 900- 1000°C. Dilution with helium was employed to vary the sum of partial pressures of the reaction participants. The experiments were carried out with nonporous graphite of high purity (the content of admixtures did not exceed lo-’%). The roughness factor of graphite was found to be 2-2.5 (from electrochemical measurements). Equation (390) proved not to be obeyed quantitatively; the results of the variation of PH2 in a broad range by addition of H, to the gas mixture at the inlet do not form a straight line in the plot l/r(” vs. P,, . Experimental data answer the equation r(” = k(1)(PH20/PH2)m

(392)

with rn = 0.5. In order to obtain (392) from the scheme of mechanism (391), the surface of carbon should be regarded as nonuniform, as in the derivation of (387) for reaction (377). Besides that, it should I& assumed that stage 1 of mechanism (391) is at equilibrium (ie., occurs much faster than stages 2 and 3). In this case stage 3 virtually does not change the surface coverage with oxygen, the coverage being determined by the equilibrium of stage 1. Since, on the other hand, carbon does not participate in stage 3, this stage may be ignored in determining r(”. Thus, we have a simple two-stage reaction con-. sisting of stages 1 and 2, stage 1 being at equilibrium. In scheme (188) stage 2 is at equilibrium. Hence, to apply the kinetic equation of this scheme, the numbering of stages should be altered. Then we obtain (392) from (196). According to the results of experiments at 950, 1000, 1050, and 1110°C k(1) = 4.2 1014e-76.800/RT (393) Reaction rate is expressed here in units cm3 (CO + CO,)NTP/hr m2. The rate of formation of C 0 2 , r“), is found in a similar manner. Since stage 1 is at equilibrium, stage 2 does not change the surface coverage and it is sufficient to take into account only stages 1 and 3 in determining r(2). Thus we obtain mechanism (343), only stage 2 of this mechanism is now designated as stage 3. Since stage 1 is at equilibrium, we have the mechanism of type (188), stages 3 and 1 of mechanism (391) corresponding to stages 1


279

and 2 of mechanism (188), respectively. Therefore, in accordance with (196) and (197) r ( 2 )=

k~2~pCO(pH~O/pHz)m[1 - ~ 1 / K ' 2 ~ ~ ~ p ~ O ~ p H ~ / p ~ (394) OpHzO~~

where k(') is a constant and K")is the equilibrium constant of route N"). It follows from (392) and (394) that (153) P)/r(') = k(z)/k(l)[l- (I/@ z, )(pco2pH2/pcopH20)I.

(395)

Equation (395) is confirmed by experimental data (154). Constants k(') and k(') are related to constants k and A in (386); it can be shown that A = k(2)/k(') and [k]= k(')[fi2)]-". Therefore the measurements of the rates of formation of CO and COz in steam gasification of carbon give the possibility of calculating theoretically the rate of gasification of the same carbon with COz . The comparison of the results of such a calculation with the direct measurement of the rate of reaction (377) on pyrographite at 1050°C demonstrated a good agreement (154). It shows that the homogeneous reaction CO

+ H20= COz + Hz

(396) did not play any essential role under the conditions of these experiments. It can be explained by the homogeneous reaction (396) being a chain one with large chain length (155). On the surface of carbon the chains terminate and reaction (396) is strongly retarded. XXI.

Ammonia Oxidation

Catalytic oxidation of ammonia to nitric oxide is the basis of production of nitric acid. It is also used in other processes, of which may be mentioned hydroxylamine synthesis (Section XIII) and the chamber process for the production of sulfuric acid. The products of the reaction are nitric oxide, water, and nitrogen, so that the reaction can be described by the equation 4NH3

+ (3 + 2~)Oz= 4sNO + 2(1 - s)N2 + 6HZO;

0 < s < 1. (397)

In this equation s is the fraction of the whole amount of ammonia reacted that is converted into nitric oxide; i.e., the selectivity. If the process is carried out properly, selectivity is near to unity. In the production of nitric acid the reaction is accomplished with mixtures of ammonia and air or sometimes with mixtures of ammonia, air, and oxygen, containing about 10 vol % NH, at pressures of 1-10 atm (below

280

M. I. TEMKIN

the concentrations of ammonia in gas mixtures will also be stated in volume percent). Elevated pressure and the enrichment of gas mixture with oxygen accelerate the subsequent oxidation of NO into NOz. Gauzes made of platinum with several percent rhodium are used as a catalyst. Rhodium reduces the wear of platinum during the process and improves selectivity; under favorable conditions s x 0.97. Together with rhodium, several percent of palladium (which is cheaper than platinum) are introduced without any noticeable adverse effect. The temperature of the reaction on platinum gauzes is near 850°C. Besides platinum, certain metal oxides and their mixtures may serve as catalysts of reaction (397). Cobaltous-cobaltic oxide, Co304, is the best among nonplatinum catalysts. The reaction on this catalyst is carried out at 7O0-75O0C, selectivity is near 0.95. Until now we limited ourselves to the discussion of reaction rates in the kinetic region even though in the commercial realization of the reactions the effect of transfer processes could have been essential. This was justified since the allowance for transfer effects would mostly be of the character of corrections. The situation is different in the case of ammonia oxidation. Both on platinum (156) and nonplatinum (157) catalysts under the conditions of a commercial process, the reaction occurs in the external diffusion region. Diffusion of ammonia rather than of oxygen is determining the rate since the reaction is conducted with oxygen in excess with respect to stoichiometry, as given by (397). Concentration of ammonia at the surface of the catalyst is so small as compared to its concentration in the gas flow that the difference of concentrations that determines the rate of diffusion virtually coincides with the ammonia content in the flow. It may be supposed that it is the low concentration of NH3 as compared to concentration of Oz (at the surface of the catalyst) that results in high selectivity. Taking this into consideration, it is appropriate in this case to first consider the course of the reaction in the external diffusion region. A characteristic feature of reaction (397) arising from the external diffusion character of its kinetics is a large temperature difference between the surface of the catalyst and the gas stream. Ammonia-air mixture fed to the catalyst is preheated only to approximately 150°Cin the case of platinum gauzes, and is not preheated at all in the case of Co304. Converters for ammonia oxidation contain no arrangements for heat exchange. The reaction is brought virtually to the end ;thus neglecting heat losses the temperature of the gas mixture leaving the catalyst layer may be taken as equal to the isoenthalpic temperature of the reaction, TH.This is the h a 1 temperature of the gas mixture (calculated from the reaction heat) after the reaction is completed without heat exchange with the surrounding medium. Often


28 1

THis called “adiabatic temperature of the reaction”; this is not rigorous. Indeed, since the reaction is carried out at constant pressure, and not at constant volume, the enthalpy H (as in the Joule-Thomson experiment) is constant but not the energy. However, the difference is not large. At 10% NH, in the mixture isoenthalpic temperature increment is about 700°K. It is not far from the factual temperature increment since, owing to the exceptionally high reaction rate and therefore large intensity of evolution of heat, heat losses with radiation comprise only a small fraction of heat transferred to the gas flow by convection. Whereas the temperature of the gas mixture increases as the mixture passes through the catalyst bed, the temperature of the whole catalyst surface is almost uniform and is near to the isoenthalpic temperature. Thus, at the beginning of the catalyst layer the temperature difference between the catalyst and the gas mixture is near to 700°K and at the end of the layer it nearly vanishes. Heat conductivity of the catalyst is of no consequence in this. The description given is equally applicable to platinum gauzes with high heat conductivity and to Co,O, with low heat conductivity. Such temperature regime results from the coincidence of the form of equations determining the distribution of concentrations and of temperature in a moving fluid, the only difference being that in the first case the equation cantains diffusion coefficientD and in the second case heat dausivity a, defined by a = I / c p p where A is heat conductivity, cp is specific heat capacity at constant pressure, and p is density of the fluid. There is a corollary of the kinetic theory of gases verified by experiments that states in gas mixtures D and a are numerically close. Thus, at 0°C and 1 atm heat diffusivity of air, a = 0.172 cm2/sec and diffusion coefficient of NH, in air, D = 0.198 cm2/sec. Therefore, the thickness of the temperature layer is approximately equal to that of the diffusion layer.g Let us assume for simplicity that the heat capacity of the gas mixture is constant, and neglect heat losses of the catalyst by radiation and heat conduction. We shall denote the isoenthalpic temperature increment at the completion of the reaction as ATH.Let the temperature of the catzlyst be higher than the initial temperature of the gas mixture by ATH.When a certain fraction, x,of the initial amount of ammonia undergoes oxidation, the initial concentration of ammonia in the gas stream, c, decreases by xc. At this the amount of heat is liberated equal to the fraction x of the complete heat of reaction, i.e, the amount capable of increasing the temperature of the gas mixture by XAT,. The initial concentration difference that determines the rate of diffusion, c, is diminished by xc. Since the rates of the The thickness of the diffusion layer, 6,, is defined by j = (D/G,)hc wherej is the flow of a substance to the unit surface, D is its diffusion coefficient, and Ac is the difference of concentrations in the main gas stream and at the surface. The thickness of the temperature layer is defined in a similar manner.

282

M. I. TEMKIN

equalization of concentrations and temperatures coincide, the gas stream removes from the surface such an amount of heat that the initial temperature difference, ATH,is diminished by xAT,. But the amount of heat liberated on the surface is just this. Thus, heat liberation is equal to heat removal, and the temperature regime is steady. The steady state is stable because if the surface temperature is varied the rate of diffusion and with it the rate of liberation of heat will change but slightly while the removal of heat will increase if the surface temperature is increased or decreased in proportion to the temperature difference. In both cases surface temperature will return to the initial value. Since the thickness of the diffusion layer decreases with increase in the linear velocity of the gas stream, it does make a difference whether the gas mixture passes through the catalyst bed of thickness h at a velocity w or through the bed of thickness 2h at a velocity 2w. In the second case the rate of the reaction is higher. Therefore, in contradistinction to the reactions in the kinetic or internal diffusion region, in ammonia oxidation time of contact (or space velocity) does not determine the fraction of ammonia reacted. The characteristics of the process conditions must include the thickness of the catalyst bed and the linear velocity of the gas stream or some other equivalent values. For the process on gauzes it is convenient to state the number of gauzes with definite characteristics (thickness of wires and density of weaving) instead of the thickness of the catalyst bed. Instead of linear velocity of the stream, it is convenient to employ a value proportional to it that will be called the “load.” The load is defined as the ratio of the amount of ammonia-air mixture fed into the reactor per unit time, to the cross-section of the catalyst bed. Since the content of ammonia in the feedstock is almost constant, it is allowable to state not the amount of ammonia-air mixture but the amount of ammonia. This is more convenient from the viewpoint of process engineering, for the load so defined determines more directly the capacity of the reactor. With increased load the thickness of the diffusion layer and that of the temperature layer decrease equally and the balance is maintained as long as the reaction rate is determined by external diffusion. But at a certain, high enough load, owing to the limitation imposed by the rate of the reaction on the surface of the catalyst, the concentration of ammonia at the surface ceases to be small and the reaction starts to pass from external diffusion region to the kinetic region. Beginning with this load, the growth of heat liberation rate falls behind the growth of heat removal. The catalyst gets cooler and this further slows down the reaction; the cooling then grows, and the reaction stops. One may now say that the catalyst is extinguished. The load that extinguishes the catalyst will be called the “limiting load.” It is a function of the reaction rate on the surface in the kinetic region, and hence depends on the nature of the catalyst.


283

A more complete discussion of the heat regime in the external diffusion region which takes into account the fact that a and D values are not exactly equal can be found in the literature (9).Our presentation aimed to graphically demonstrate the main features of the phenomenon. Because of the external diffusional character of kinetics, the degree of ammonia conversion can be predicted theoretically as a function of geometric parameters of the catalyst bed and hydrodynamic characteristics of the gas stream. For ammonia oxidation on platinum gauzes under atmospheric pressure, using the analogy between diffusion and heat transfer, an equation was obtained (156) equivalent to the following:

+

lg(co/cl) = (Sn/d~)[0.43 0.274(d~)O.~~].

(398)

Here co is NH, concentration in the initial mixture, c1 is NH, concentration in the mixture behind the gauzes, S is geometric surface of a gauze per its unit area, n is the number of gauzes, d is the diameter of wire in centimeters, and w is the load of the ammonia-air mixture in liters of mixture (NTP)/hr cm2. It is to be noted that in the mixture at 10% NH,, the value of w in 1 mixture (NTP)/hr cm2 coincides numerically with the value in m3 NH, (NTP)/hr m2. Equation (398) was confirmed experimentally (156). The gauzes used in these experiments were of d = 9 x lo-, cm, S = 1.87. The number of gauzes varied between 1 and 8, the load was changed in the range of from 55 to 4000 1 mixture (NTP)/hr cm2. In the commercial atmospheric pressure reactors with three gauzes, the load is about 200 m3 NH, (NTP)/hr m2. It follows from (398) that under these conditions the fraction of unreacted ammonia is practically nil. c0304 pellets used in practice are of 4-5 mm in size. Thus, they are much larger than the diameter of wires in platinum gauzes. For this reason, in contrast to the reaction on gauzes, the reaction on c0304 pellets under atmospheric pressure is characterized by the Reynolds number much larger than 1, the Reynolds number being defined by Re = ul/v where u is the linear velocity of the stream, 1 is the characteristic dimension, v is the kinetic viscosity coefficient. The thickness of the diffusion layer for such pellets is

6, 11~~0.5 (399) (since in gases 6, is close to the thickness of the Prandtl layer, 6). Hence it can be deduced that Ig(c,/c,) = A(h/PW0.5) (400) where h is the thickness of the catalyst bed and A is a constant depending on the shape of pellets. If h and I are expressed in centimeters and w is measured in 1 mixture (NTP)/hr cm2, then, according to laboratory measurements

284

M. I. TEMKIN

under atmospheric pressure (150, A = 4.2 for cylindrical pellets with the diameter of base, 1, equal to the height, and A = 2.5 for spherical pellets with the diameter 1. These A values refer to the downward stream of the gas mixture usually employed in the reactors with the nonplatinum catalyst. Taking a co/cl ratio such that there would be practically no diffusion of ammonia, e.g., co/cl = lo4, we obtain from (400) the required thickness of the bed of pellets. Let us consider now the effect on the process of passage from atmospheric pressure to a higher pressure P, with load, w ,and the size of catalyst pellets, I, unchanged. The Sherwood number, Sh = l / h D , is determined by the Reynolds, Re = ul/v, and the Schmidt, Sc = v/D, numbers. The value of w being given, v, D, and u values are inversely proportional to pressure, P. Therefore, 6, is independent of P at constant w and 1. Since u and D are inversely proportional to P, the increase in the time of contact of the gas mixture with the catalyst by a factor of P as a result of passage from 1 atm to P atm is compensated by the P times diminition of the diffusion coefficient, and the fraction of ammonia reacted remains unchanged. We conclude that (398) and (400) may also be used without any modification in the case of reaction under pressures differing from atmospheric. In ammonia oxidation, under elevated pressures higher loads are used than under atmospheric pressure, possibly because there is no difficulty in overcoming the resistance to the flow. For instance, at 7 atm the load is 5000 m3NH, (NTP)/hr m2 and the number of gauzes is 16. Although the load is 25 times higher than the above-mentioned load for the atmospheric pressure reactors, the number of gauzes is only 5.3 times larger. Therefore, the productivity of gauzes in the reactors under elevated pressure is much higher. It would be a mistake to think that this increase in productivity results from the increase in pressure; it results from the increase of load making the diffusion layer thinner. The possibility of such increase in productivity is restricted by the value of limiting load, but for platinum gauzes this value is very high. Experiments with gauzes characterized above (wire thickness 0.09 mm) done without preheating of the ammonia-air mixture under atmospheric pressure and at 10% NH, gave a limiting load about 30000 1 mixture (NTP)/hrcm2 (157). The limiting loads for nonplatinum catalysts are smaller than that for platinum. The largest value was found for Co,O, (158). Pertinent data will be given and discussed below. To observe ammonia oxidation on platinum in the kinetic region, the reaction must be carried out at low pressures since the decrease in pressure increases the diffusion coefficient and slows down the reaction, both factors favoring the transition to the kinetic region. mm Hg in a static We studied the reaction under pressures about


285

system on an electrically heated platinum wire (159). The temperature of the wire was 900°C (160)." Since the rate of the homogeneous reaction 2N0

+ O2 = 2N02

obeys an equation of third order, the extent of this reaction under these low pressures is negligible for the duration of the experiment. This allows one to judge whether NO is the primary reaction product. Earlier Bodenstein (161, 162) and Krauss (263, Z64 oxidized NH3 on Pt at low pressures with a freezing off of the products on a surface cooled with liquid air. They observed the formation of nitrous acid H N 0 2 and hydroxylamine N H 2 0 H . The authors concluded that under usual conditions (i.e., at pressures of the order of 1 atm) NO is formed as a result of the decomposition of primary products such as HNO, in the hot gas stream. In order to determine NO we applied the previously discovered (165)fast heterogeneous reaction of NO with oxygen on a glass surface cooled with liquid air. The product of this reaction has a composition intermediate between N 2 0 3and N204.The reaction is characterized by negative activation energy; namely, its rate increases twofold at lowering the temperature by 2 K. At pressures about lo-' mm Hg when a small glass appendix, jointed to the system of 0.5 1 volume, is cooled with liquid air, the reaction proceeds virtually to completion in several minutes. With the help of this reaction it has been definitely ascertained that NO is the main product of the reaction of NH, with O2 on Pt ; besides NO, N, is the only nitrogen containing species formed. Evidently, HNO, in Bodenstein's experiments resulted from the interaction of water with the product of the heterogeneous reaction of NO with O2 that occurred on the wall cooled with liquid air. The formation of NH,OH is more difficult to explain. In Bodenstein's experiments NHzOH was observed at platinum temperatures above 1100°C. It may be supposed that platinum underwent dispersion. On its particles in the condensate containing dissolved NO and H, (the product of decomposition of NH, on Pt) NH,OH could be formed by the reaction discussed in Section XIII. Experiments have shown that the reaction rate is approximately proportional to PNH3 and is independent of Po,. A comparison of the reaction rate with the number of impacts of NH, molecules upon the surface shows that approximately one of 20 impacts results in the reaction. The simplest explanation is that the rate-determining elementary reaction occurs at the impact of the NH, molecule upon the surface nearly completely covered with oxygen. l o The previously given figure (820°C) (159) was calculated without correction for the cooling effect of the leads.

286

M. I. TEMKIN

Ammonia oxidation under low pressures, as a method of transferring this reaction into the kinetic region, is inapplicable in the case of the c0304 catalyst since, at such temperatures as 700°C and low O2 pressures, Co,O, decomposes with the formation of COO.In order to obtain information on the kinetic of NH, oxidation on Co304,we studied limiting loads at which the catalyst is extinguished (166, 167). The experiments were performed at pressures from 1 to 9 atm. A catalyst pellet was placed in a vertical tube of a diameter such that the cross-section occupied by the pellet comprised onehalf of the cross-section of the tube. This was an imitation of conditions in the bed of pellets. The gas mixture at the inlet was at room temperature; the stream of the mixture was directed downward. Cylindrical pellets of the catalyst were used with the diameter 4 mm, equal to the height. In the experiments for determination of the dependence of limiting load on the size of pellets, spheres with a diameter of from 2 to 5 mm were employed. The limiting load for cylindrical pellets 4 mm in size, which worked for a short time under atmospheric pressure and at 10% NH, in the mixture with air, was near to 10,000 1 mixture (NTP)/cm2 hr. The measurements of the temperature of pellets with an optical pyrometer showed that it is nearly independent of load up to the limiting load. The limiting load was found to be proportional to the total pressure of the ammonia-air mixture and to the size of the pellets. The limiting load is increased twofold with the increase in NH, content in the mixture with air from 9 to 10%. It is also increased twofold at constant NH, content and the substitution of oxygen for air. From these results and applying the theory of critical ignition and extinction phenomena of strongly exothermic heterogeneous reactions developed by Frank-Kamenetskii and Buben (9),it was found that the following form of kinetic equation is the most probable: r = kP$zp Pz;'4

(401)

where k is a constant, PNH3 and Po, are partial pressures of NH, and O2 at the surface of the catalyst. The industrial application of the Co,O, catalyst for ammonia oxidation is complicated by its sensitivity to poisoning action of small amounts of sulfur compounds in the ammonia-air mixture. This phenomenon was studied with the use of radioactive sulfur containing S3' that made it possible to measure very low concentrations of sulfur (168). Poisoning results in decrease both of selectivity and limiting load. A noticeable decrease in selectivity starts at sulfur concentrations in the gas mixture from 0.05 mg/m3. This concentration is many times lower than minimum HzS and SOz concentrations in air detected by smell.


287

REFZRENCE.9

1. Kiperman, S . L., “Vvedenie v kinetiku geterogennykh kataliticheskikh reaktsii.” Nauka, Moscow, 1964. 2. Snagovskii, Yu. S., and Ostrovskii, G. M., “Modelirovanie kinetiki geterogennykh kataliticheskikh protsessov.” Khimiya, Moscow, 1976. 3. McGlashan, M. L., Annu. Rev. Phys. Chem. 24,51 (1973). 4. Temkin, M. I., Kiperman, S. L., and Luk‘yanova, L. I., Dokl. Akad. Nauk SSSR 74,763 (1950). 5. Temkin, M. I., Kinet. Katal. 3, 509 (1962). 6 . Sidorov, I. P., Shishkova, V. V., and Temkin, M. I., Tr. GIAP 6,323 (1956). 7 Korneichuk, G. P., and Strel’tsov, 0. A., Probl. Teor. Prakt. Issled. Obl. Katal. Chapter 3 (1973). 8. Ionov, Yu. V., Rastaturin, V. A., and Kurochkin, Yu. Yu., Zh. Prikl. Khim. 64,606 (1971). 9 . Frank-Kamenetskii, D. A., “Diffuziya i teploperedacha v khimicheskoi kinetike,” 2nd ed. Nauka, Moscow, 1967. 10. Satterfield, C . N., “Mass Transfer in Heterogeneous Catalysis.” 1970. 11. Temkin, M. I., Kinet. Katal. 16, 504 (1975). 12. Thiele, E. W. 2nd. Eng. Chem. 31,916 (1939). 13. Smith, N. L., and Amundsen, N. R., Ind. Eng. Chem. 43,2156 (1951). 14. Graham, T., London Edinburgh Philos. Mag. J. Sci. [3] 2, 175, 269, and 351 (1833). 15. Herzfeld, K. F., in “Lerbuch der Physik” (Miiller and Pouilets, eds.), 11th ed., Vol. 3, 2nd half. Vieweg, Braunshweig, 1925. 16. Hoogschagen, H.,Ind. Eng. Chem. 47,906 (1955). 17. Scott, D. S., and Dullien, F. A., A2ChE J. 8, 113 (1962). 18. Evans, R. B., Watson, G. M., and Mason, E. A,, J. Chem. Phys. 35,2076 (1961). 19. Maxwell, J. C., London Edinburgh Philos. Mag. J. Sci. [4] u),21 (1860). 20. Langmuir, I., J. Am. Chem. Soc. 40, 1361 (1918). 21. Temkin, M., Zh. Fiz. Khim. 11, 169 (1938); Acta Physicochim. URSS8,141 (1938). 22. Langmuir, I. J. Am. Chem. SOC.17,621 (1921-1922). 23. Ternkin, M. I., Bulg. Acad. Sci., Commun. Dep. Chem. 1, No 3, 65 (1968) (in Russian). 24. Temkin, M. I., Zh. Fiz. Khim. 4, 573 (1933); Acta Physicochim. URSS 1,36 (1934). 25. Kuchaev, V. L., and Temkin, M. I., Kinet. Katul. 13, 719 and 1024 (1972). 26. Horiuti, J., and Nakamura, T., Z . Phys. Chem. (Frankfurt am Main) [N.S.] 11,358 (1957). 27. Temkin, M. I., in “Mekhanizm i kinetika slozhnych kataliticheskich reaktsii” (S. Z. Roginskii, G. V. Isagulyants, and I. I. Tret’yakov, eds.), p. 57. Nauka, Moscow, 1970; Int. Chem. Eng. 11,709 (1971). 28. Slin’ko, M. G., Bykov, V. I., Yablonskii, G. S., and Akramov, T. A., Dokl. Akad. Nauk SSSR 226,876 (1976). 29. Noyes, R. M., and Field, R. J., Apnu. Rev. Phys. Chem. 25,95 (1974). 30. Semenov, N. N., “Tsepnye reaktsii.” Goskhimtechizdat, Leningrad, 1934. 31. Temkin, M. I., Dokl. Akad. Nauk SSSR 152, 156 (1963). 32. Temkin, M. I., Nauchn. Om.Podbora Proizvod. Katal., Dokl. Vses. Soveshch., 1962 p. 46 (1964). 33. Temkin, M. I., J. Res. Inst. Catal., Hokkaido Uniu. 16, 355 (1968) (in Russian). 34: Harary, F., “Graph Theory.” Addison-Wesley, Reading, Massachusetts, 1969. 35. Temkin, M. I., Zh. Fiz. Khim. 31,3 (1957). 36. Temkin, M. I., Kinet. Katal. 17, 1095 (1976). 37. Temkin, M. I., Ann. N . Y . Acad. Sci. 213,79 (1973). 38. Khomenko, A. A., Apel’baum, L. O., Shub, F. S., and Temkin, M. I., Kinet. Katal. 13,251 (1972).

288

M. I. TEMKIN

39. Temkin, M. I., Dokl. Akad. Nauk S S S R 165, 615 (1965). 40. Temkin, M. I., Zh. Fir. Khim. 15,296 (1941). 41. Temkin, M. I., Kinet. Katal. 13, 555 (1972). 42. Evans, M. G . , and Polanyi, M., Trans. Faraday SOC.32,1333 (1936). 43. Zel’dovich, J., Acta Physicochim. URSS 1,961 (1935). 44. Temkin, M. I., Probl. Kiriet. Katal. 6, 54 (1949). 45. Zel’dovich, J., Acta Physicochim. URSS 1, 449 (1934); see also Roginskii, S. Z., Probl. Kinet. Katal. 3, 356 (1937). 46. Bangham, D. H., and Burt, F. P., Proc. R. SOC.London Ser. A 105,481 (1924). 47. Temkin, M. I., Kinet. Karal. 16, 1461 (1975). 48. Temkin, M. I., Kinet. Katal. 5, 1005 (1967). 49. Romanushkina, A. E., Kiperman, S . L., and Temkin, M. I., Zh. Fiz. Khim. 27,1181 (1953). 50. Kul’kova, N. V., and Temkin, M. I., Zh. Fiz. Khim. 21,2017 (1957). 51. Kil’kova, N. V.,and Temkin, M. I., Zh. Fiz. Khim. 36, 1731 (1962). 52. Kul’kova, N. V . ,Levchenko, L. P., and Temkin, M. I., Zh. Fiz. Khim. 42,2688 (1968). 53. Temkin, M. I., Dokl. Akad. Nauk S S S R 161, 160 (1965). 54. Kurilenko, A. I . , Kul’kova, N. V., Ostrovskii, V. E., and Temkin, M. I., Dokl. Akad. Nauk S S S R 123,878 (1958). 55. Ostrovskii, V. E., Kul’kova, N. V., Lopatin, V. L., and Temkin, M. I., Kinet. Katul. 3,189 (1962). 56. Kayumov, R. P.,Kul’kova, N. V., and Temkin, M. I., React. Kinet. Catul. Lett. 1,29 (1 974) (in Russian). 57. Kayumov, R. P., Kul’kova, N. V., and Temkin, M. I., Kinet. Katal. 15, 157 (1974). 58. Kayumov, R. P., Kul’kova, N. V., and Temkin, M. I., Kinet: Katul. 15, 1349 (1974). 59. Kurilenko, A. I., Kul’kova, N. V., Rybakova, N. A., and Temkin, M. I., Zh. Fiz. Khim. 32, 797 and 1043 (1958). 60. Kurilenko, A. I., Kul’kova, N. V., Baranova, L. P., and Temkin, M. I., Kinet. Katal. 3,208 (1962). 61. Ostrovskii, V. E., Kul’kova, N. V., Kharson, M. S.,and Temkin, M. I., Kinet. Katal. 5,469 (1964). 62. Ionov, Yu.V . ,Zhidkova, L. K., Kul’kova, N. V., and Temkin, M. I., Kinet. Katal. 14,642 (1973). 63. Temkin, M. I., and Kul’kova, N. V., Dokl. Akad. Nuuk S S S R 105, 1021 (1955). 64. Rudnitskii, L. A., Kul’kova, N. V., and Temkin, M. I., Metody Issled. Katal. Katal. Reakt. 3, 145 (1965). 65. Rudnitskii, L. A,, Kul’kova, N. V., andTemkin, M. I., Probl. Kinet. Katul. 12, 113 (1967). 66. Rudnitskii, L. A,, Shakhovskaya, L. I., Kul’kova, N. V., and Temkin, M. I., Doki. Akad. Nuuk S S S R 182, 1358 (1968). 67. Shakovskaya, L. I . , Rudnitskii, L. A,, Kul’kova, N. V., and Temkin, M. I., Kine,. Katal. 11, 467 (1970). 68. Ostrovskii, V. E., and Temkin, M. I., Kinet. Katal. 7, 529 (1966). 69. Twigg, G . H . , Trans. Faraduy SOC.42,284 (1946). 70. Orzechowski, A., and MacCormack, K. E., Can. J. Chem. 32,388, 415, and 433 (1954). 71. Temkin, M. I., Zh. Vses. Khim. 0-ua. u),7 (1975). 72. Temkin, M. I., and Kul’kova, N. V., Kinet. Katal. 16, 1211 (1975). 73. Margolis, L. Ya., “Geterogennoe kataliticheskoe okislenie uglevodorodov (sintez monomerov).” Gostoptekhizdat, Moscow, 1962. 74. Lyubarskii, G. D., Dokl. Akad. Nauk S S S R 110, 112 (1956). 75. Kilty, P. A,, and Sachtler, W. M. H., Catal. Re#.-Sci. Eng. 10, 1 (1974). 76. Kenson, R. E., and Lapkin, M., 1. Phys. Chem. 74, 1943 (1970).


289

77. Savodnik, N. V., Kul’kova, N. V., Dokholov, D. M., Lopatin, V. L., and Temkin, M. I., Kinet. Katal. 13, 1520 (1972). 78. Brian, P. L. T., Hales, H. B., and Sherwood, T. K., AIChE J. 15, 727 (1969). 79. Horiuti, J., and Toya, T., Solid State Surf Sci. 1 (1969). 80. Shlygin, A. I., and Frumkin, A. N., Acta Physicochim. URSS 3,791 (1935); 4,911 (1936). 81. Temkin, M. I., Vopr. Khim. Kinet., Katal. Reakt. Sposobn., Dokl. Vses. Soveshch., 1955 p. 484 (1955). 82. Bodrov, I. M., Apel’baum, L. O., and Temkin, M. I., Kinet. Katal. 5, 696 (1964). 83. Khomenko, A. A,, Apel’baum, L. O., Shub, F. S., andTemkin, M. I., Kinet. Katal. 11,1480 (1970). 84. Khomenko, A. A., Apel’baum, L. O., Shub, F. S., Snagovskii, Yu. S., and Temkin, M. I., Kinet. Katal. 12,423 (1971); Temkin, M. I., Shub, F. S., Khomenko, A. A., and Apel’baum, L. O., Nauchn. O m . Katal. Konvers. Uglevodorodov, 1977. p. 3 (1977). 85. Bodrov, I. M., Apel’baum, L. O., and Temkin, M. I., Kinet. Katal. 8,821 (1967). 86. Eidus, Ya. T., and Zelinskii, N. D., Izu. Akad. Nauk SSSR, Ser. Khim p. 289 (1940). 87. Eidus, Ya. T., Izv. Akad. Nauk SSSR, Ser. Khim. p 65 (1943). 88. Fisher, F., and Tropsch, H., Brennst.-Chem. 7,97 (1926). 89. Craxford, S. R., and Rideal, E. K., J . Chem. Sor. p. 1604 (1939). 90. Emmett, P. H., ed., “Catalysis,” 3, pp. 171 and 265. New York, 1955. 91. Nielsen, A,, “An Investigation on Promoted Iron Catalysts for the Synthesis of Ammonia,” 3rd ed. Giellerups, Copenhagen, 1968. 92. Malina, I. K., “Razvitie issledovanii v.oblasti sinteza ammiaka.” Nauka, Moscow, 1973. 93. Temkin, M., and Pyzhev, V., Zh. Fir. Khim. 13,851 (1939); Acta Physicochim. URSS 12, 327 (1940). 94. Temkin, M. I., in “Catalysis,” Proc. Int. Congr. 5th, 1972, (J. W. Hightower, ed.), Vol. 1, p. F-I 13. North-Holland Publ., Amsterdam, 1973. 95. Morozov, N. M., Luk’yanova, L. I., and Temkin, M. I., Kinet. Katal. 6, 82 (1965). 96. Morozov, N. M., Shapatina, E. N., Luk’yanova, L. I., and Temkin, M. I., Kinet. Katal. 7, 688 (1966). 97. Scholten, J. J. F., “Chemosorption of Nitrogen on Iron Catalists in Connection with Ammonia Synthesis.” Croniger, Amsterdam, 1959. 98. Temkin, M. I., Khim. Nauka Promst. 2, 219 (1957). 99. Shishkova, V. V., Sidorov, I. P., and Temkin, M. I., Tr. GIAP 7, 62 (1957). 100. Sokolova, D. F., Morozov, N. M., and Temkin, M. I., Zh. Fiz. Khim. 33,471 (1959). 101. Temkin, M. I., Morozov, N. M., and Shapatina, E. N., Kinet. Katal. 4,260 and 565 (1963). 102. Smirnov, I. A., Morozov, N. M., and Temkin, M. I., Dokl. Akad. Nauk SSSR 153, 386 (1963). 103. Smirnov, I. A,, Morozov, N. M., and Temkin, M. I., Kinet. Katal. 6, 351 (1965). 104. Nielsen, A,, Kjer, J., and Hansen, B., J. Catal. 3, 68 (19642. 105. Kiperman, S., and Temkin, M., Zh. Fiz. Khim. 20,369 (1946). 106. Kiperman, S., and Temkin, M., Zh. Fiz. Khim. 20,623 (1946). 107. Kiperman, S. L., and Temkin, M. I., Zh. Fiz. Khim. 21,927 (1947). 108. Ozaki, A., Taylor, H., and Boudart, M., Proc. R. Soc. London, Ser. A 258,47 (1960). 109. Schulz, G., and Schaefer, H., Ber, Bunsenges. Phys. Chem. 70,21 (1966). 110. Shapatina, E. N., Kuchaev, V. L., and Temkin, M. I., Kinet. Katal. 12, 1476 (1971). I l l . Boreskova, E. G., Kuchaev, V. L., Pen’kovoi, B. E., and Temkin, M. I., Kinet. Katal. 13, 358 (1972). 112. Igranova, E. G.,Ostrovskii, V. E., and Temkin, M. I., “Kinet.-2,” Vtor. Vses. Konf. Kinet. Katal. Reakts., Vol. 1, p. 135 (1975). 113. Emmett, P. H., and Kummer, J., Ind. Eng. Chem. 35, 677 (1943).

290

M. I. TEMKIN

114. Temkin, M. I., Zh. Fiz. Khim. 24, 1312 (1950). 115. Sidorov, I. P., and Livshits, V. D., Zh. Fiz. Khim. 21, 1177 (1947).

116. Sokolova, D.F., Morozov, N. M., and Temkin, M. I., Zh. Fiz. Khim. 33,471 (1959). 117. Kjm, J., ”Measurement and Calculation of Temperature and Conversion in Fixed-Bed Catalytic Reactors.” Copenhagen, 1958. 118. Fauser, G., Acta Technol. Chim. Reaz. Reattori, 6th, 1962 p. 1 (1962). 119. Bridger, G. W., and Snowdon, C. B., in “Catalyst Handbook,” pp. 142 and 145. Wolfe, London, 1970. 120. Tamaru, K., Bull. Chem. SOC.Jpn. 37,771 (1964); Proc. Int. Congr. Catal., 3rd, 1%4(1965). 121. Sastri, M. V. S., Catal., Proc. Int. Congr. 5th, 1972 Vol. 2, p. 1273 (1973). 122. Smirnov, I. A., Kinet. Katal. I, 107 (1966). 123. Bulatnikova, Yu. I., Apel’baum, L. O., and Temkin, M. I., Zh. Fiz. Khim. 32,2717 (1958). 124. Kul’kova, N. V., and Temkin, M. I., Zh. Fiz. Khim. 23, 695 (1939). 125. Boudart, M., AIChE J. 18,465 (1972). 126. Shchibrya, G. G., Morozov, N. M., and Temkin, M. I., Kinet. Katal. 6, 1057 and 1115 (1956).

127. Kodama, S., Fukui, K., Tame, T., and Kinoshita, M., Shokubai 8,50 (1952). 128. Kodama, S., Mazume, A., Fukuba, K., and Fukui, K., Bull. Chem. SOC.Jpn. 28,318 (1955). 129. Stelling, O., and Krusenstierna, O., Acta Chem. Scmd. 12, 1095 (1 958). 130. Cherednik, E. M., Morozov, N. M., and Temkin, M. I., Kinet. Katal. 10, 603 (1969). 131. Temkin, M. I., Nakhmanovich, M. L., and Morozov, N. M., Kinet. Katal. 2, 722 (1961). 132. Nakhmanovich, M. L., Morozov, N. M., Buadze, L. G., and Temkin, M. I., Dokl. Akad. Nauk SSSR 148,1346 (1963). 133. Weston, R. E., Tetrahedron 6, 37 (1959). 134. Nakhmanovich, M. L., Thesis, Moscow (1963). 135. Taikhert, A. M., Morozov, N. M., and Temkin, M. I., Kinet. Katal. 4,904 (1963). 136. Black, J. F., and Taylor, H. S., J . Chem. Phys. 11, 395 (1943). 137. Kul’kova, N. V., Kuznets, E. D., and Temkin, M. I., Dokl. Akad. Nauk SSSR 90, 1067 (1953).

138. Stroeva, S.S.,Kul’kova, N. V., andTemkin, M. I. Dokl. Akad. Nauk SSSR 124,628 (1959). 139. Potter, C., and Baron, S., Chem. Eng. Prog. 47,473 (1951). 140. Shapatina, E. N., Kuchaev, V. L., Pen’kovoi, B. E., and Temkin, M. I., Kinet. Katal. 17, 644 (1976); Shapatina, E. N., Kuchaev, V. L., and Temkin, M. I., ibid. 18,968 (1977). 141. Lanczos, C., “Applied Analysis.” Prentice-Hall, Englewood Cliffs, New Jersey, 1956. 142. Frank-Kamenetskii, D. A., Dokl. Akad. Nauk SSSR 23,662 (1939). 143. Semechkova, A. F., and Frank-Kamenetski!, D. A., Zh. Fiz. Khim. 19,291 (1940). I44. Gadsby, J., Long, F. J., Sleightholm, P., and Sykes, K. W., Proc. R. SOC.London, Ser. A 193, 357 (1948). 145. Evropin, V. A,, Kul’kova, N. V., and Temkin, M. I., Zh. Fiz. Khim. 30,348 (1956). 146. Gadsby, J., Hinshelwood, C. N., and Sykes, K. W., Proc. R. SOC.London, Ser. A 187,128 (1946). 147. Long, F. J., and Sykes, K. W., Proc. R . Soc. London, Ser. A 193, 377 (1948). 148. Long, F. J., and Sykes, K. W., J . Chim. Phys. 47, 361 (1950). 149. Johnstone, H. F., Chen, C. J., and Scott, D. S., Ind. Eng. Chem. 44, 1564 (1952). 150. Jolley, L. J., and Poll, A., J. Inst. Fuel 26, 33 (1953). 151. Key, A., Gas Res. Board, Commun. 40,(1948); Strickland-Constable, R. F., J . Chim. Phys. 47, 356 (1950). 152. Cherednik, E. M., Apel’baum, L. O., and Temkin, M. I., Dokl. Akad. Nauk SSSR 174,891 (1 967). 153. Temkin, M. I., Cherednik, E. M., and Apel’baum, L. O., Kinet. Katal. 9,95 (1968).


29 1

154. Cherednik, E. M., Temkin, M. I., and Apel’baum, L. O., Kinet. Katal. 9,824 (1968). 155. Shub, F. S., Apel’baum, L. O., and Temkin, M. I., Kinet. Katal. 11,566 and 11 19 (1970). 156. Apel’baum, L. O., and Temkin, M., Zh. Fiz. Khim. 22, 179 and 195 (1948). 157. Temkin, M. I., Morozov, N. M., Pyshev, V. M., Apel’baum, L. O., Luk’yanova, L. I., and Demidkin, V. A,, Probl. Fiz. Khim. 2, 14 (1959). 158. Morozov, N. M., Lukyanova, L. I., and Temkin, M. I., Kinet. Katal. 7, 172 (1966). 159. Apel’baum, L. O., and Temkin, M., Dokl. Akad. Nauk SSSR 74,963 (1950). 160. Apel’baum, L. O., and Temkin, M. I., Zh. Fiz. Khim. 33,2697 (1959). 161. Bodenstein, M., and Biittner, G., Z. Angew. Chem. 47, 364 (1934). 162. Bodenstein, M., Z . Angew. Chem. 48,327 (1935); Trans. Electrochem. SOC. 71,353 (1937); Z.,Elektrochem. 47, 501 (1941); Z . Phys. Chem., Abt. B 50, 333 (1941). 163. Krauss, W . , Z . Phys. Chem., Abt. B 39,83 (1938). 164. Krauss, W . , and Schuleit, H., Z. Phys. Chem., Abt. B 45, 1. (1940). 165. Temkin, M., and Pyzhev, V., Acta Physicochim. URSS 1, 177 (1934); 2,473 (1935). 166. Shub, F. S., Khomenko, A. A., Apel’baum, L. O., and Temkin, M. I., “Kinet.-2,” Vtor. Vses. Konf. Kinet. Katal. Reakts., Vol. 2, p. 13 (1975). 167. Shub, F. S., Khomenko, A. A., Apel’baum, L. O., andTemkin, M. I., Kinet. Katal. 17,1586 (1976). 168. Bulatnikova, Yu. I., Apel’baum, L. O., and Temkin, M. I., Zh. Fiz. Khim. 32,2717 (1958).


ADVANCES IN CATALYSIS, VOLUME 28

Metal = Catalyzed Dehyd rocyc Iizat ion of Alkylaromatics SIGMUND M. CSICSERY Chevron Research Company Richmond, California

I. Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . .

. . . . . . . . . . . . . 111. The Dehydrocyclization of C,, and Higher Alkylbenzenes . . . . . . . A. Metal-Catalyzed Cyclization . . . . . . . . . . . . . . . . . . B. Cyclization over Dual-Function Catalysts . . . . . . . . . . . . . C. Reactions that Accompany Dehydrocyclization . . . . . . . . . . IV. Double Cyclization of Cs and Higher Paraffins . . . . . . . . . . . . 11. The Dehydrocyclization of C, Alkylbenzenes

293 295 296 299 306 309 312

V. The Dehydrocyclization of Alkylbenzenes Over Chromia-Alumina Catalysts . . . . . . . . . . . . . . . . . . . . 314 VI. The Dehydrocyclization of Alkylnaphthalenes . . . . . . . . . . . . 3 15 VII. Dehydrocyclization of Diphenylalkanes . . . . . . . . . . . . . . . 3 18 VIII. Conclusions . . . . . . . . . . . . . . . . . . . . . . . . . . . 3 19 References . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3 19

I.

Introduction

Catalytic cyclization is a very important reaction, both commercially and theoretically. Here we review metal-catalyzed cyclization of alkylaromatic hydrocarbons, and we present reaction mechanisms that govern these reactions. Monoalkylaromatics and ortho-substituted dialkylaromatics containing at least three carbon atoms in the side chains can easily undergo dehydrocyclization to form an additional ring. One can classify cyclizations of . alkylaromatics in many different ways.

1. The new ring can have five or six (or very rarely more or fewer) carbon atoms. 2. The new bond can be formed between two sp3 hybridized carbon atoms (the cyclization of ortho-substituted dialkylaromatics) or between an sp3 and an sp2 hybridized carbon atom (the cyclization of monoalkylaro293

Copyright

0 1979 by Academic Press, Inc.

All rights of reproduction in any form reserved. ISBN 0-12-007828-7

294

SIGMUND M. CSICSERY

matics). A new ring can also be formed between two sp2 hybridized carbon atoms (the formation of fluorene from diphenylmethane). 3. Formation of the new bond can involve two primary carbon atoms, or a primary and a secondary carbon atom, or two secondary carbon atoms, or even tertiary carbon atoms. 4. The new ring can be saturated (e.g., indan, tetralin) or aromatic (e.g., indene, naphthalene). 5. If the alkyl side chain has five or more carbon atoms, the second ring can form entirely on the side chain, not involving the aromatic ring; for example, phenylcyclopentane, phenylcyclopentene, and phenylcyclopentadiene can be obtained from n-pentylbenzene.

n-Pentylbenzene

Phenylcyclopentane

Phenylcyclopentene

Phenylcyclopentadiene

Cyclization can be catalyzed by acids and by metals. Whereas the mechansim of acid-catalyzed cyclization is well understood, many questions remain about the nature of metal-catalyzed cyclization. It is, however, generally observed that the presence of an aromatic ring enhances the rate of cyclization : alkylaromatics cyclize at a much higher rate than aliphatic hydrocarbons. The dehydrocyclization of alkylaromatics was first described more than four decades ago. In 1936 Moldavskii and Kamusher reported the formation of naphthalene from n-butylbenzene on chromia at 475°C ( I , 2). In 1945 Herington and Rideal reported the formation of indene from n-propylbenzene over chromia-alumina (3). Platinum-containing catalysts were first used for these reactions in 1956 by Kazanskii and co-workers (4-6).

DEHYDROCYCLIZATION OF ALKYLAROMATICS

295

II. The Dehydrocyclization of C, Alkylbenzenes

The chemistry and kinetics of the cyclization reactions of 1-methyl-2ethylbenzene and n-propylbenzene were studied in great detail by Shephard and Rooney (7) and by Liberman, Bragin, and co-workers (6,8,9).Shephard and Rooney studied the reactions of n-propylbenzene, o-methylstyrene, and 1-methyl-2-ethylbenzene over a 0.5% platinum-on-?-alumina catalyst between 337°C and 490°C, in the presence of excess hydrogen. They used the microreactor pulse technique. Hydrogenation and dehydrogenation are very rapid over this catalyst. Equilibrium is obtained between the olefins and the corresponding paraffins at all the temperatures investigated. Cyclization gives indan. Subsequent hydrogenolysis of the five-membered ring of the indan yields n-propylbenzene from 1-methyl-2-ethylbenzene, and 1methyl2-ethylbenzene from n-propylbenzene. This reaction shows that platinum can catalyze isomerization, via cyclic intermediates. The mechanism i s quite general: it could involve normal and branched paraffins, and alkylaromatics as well (20-22).Shephard and Rooney believe that multiple x-bonded transition states are involved in the C,-cyclization of alkylbenzenes, in ring opening and in isomerizations involving cyclic intermediates. Barron, Gault, and co-workers favor an a, p, y-triadsorbed intermediate (20).Pdl, Dobrovolszky, and TttCnyi propose a dual-site mechanism for C, -cyclizations and isomerizations involving five-membered cyclic intermediates (224. Bragin and co-workers found that over platinum-on-carbon catalysts, both paraffins and alkylaromatics follow zero-order kinetics. Activation energy for C,-dehydrocyclization in which the new bond is formed between two sp3 hybridized atoms is substantially less than the activation energy of cyclization in which the new bond is formed between one sp3hybridized atom and the sp2 hybridized carbon atom of the aromatic ring. Over one batch of platinum-on-carbon catalyst, B r a e and co-workers obtained 20 kcal/mol and 27.5 kcal/mol activation energies for the dehydrocyclization of paraffins and monoalkylbenzenes, respectively (6). Another batch of platinum on carbon (which differed only in some minor details of preparation from the first batch), gave 14 kcal/mol for the cyclization of 1-methyl-2-ethylbenzene and isooctane, and 29 kcal/mol for the cyclization of secondary butylbenzene (8) (Fig. 1). The results of Shephard and Rooney support these observations. The activation energy for n-propylbenzene over platinum-on-alumina catalyst is twice that of the cyclization of 1-methyl-2-ethylbenzene (1 1.6 versus 5.8 kcal/mol; Fig. 2). However, these last values are based on pulse-reactor results; therefore, they should be used only for qualitative comparisons.

296

SIGMUND M. CSICSEXY 1.3-

0.9 1.1

0) 0

0.5 0.1

I

I

I

I

I

1.5

1.6

1.7

0.3

Bragin el al. (8).

s

rn

-0 1

1.0

-

0.8 1.4

10001T"K

FIG. 2. C,-Dehydrocyclization of 1-methyl-2-ethylbenzene (A) and n-propylbenzene ( 0 ) over Pt-y-alumina catalyst (0.5% Pt) in the presence of hydrogen. Plotted from the data of Shephard and Rooney (7) by Bragin et al. (8).

111.

The Dehydrocyclization of C,, and Higher Alkylbenzenes

Commercial reforming catalysts have both metal and acid sites. Both could contribute to cyclization. If there are four or more carbon atoms in the side chain of a mono-alkylaromatic or ortho-substituted dialkylaromatic hydrocarbon, cyclization can yield either five- or six-membered rings. This multiplicity of reaction pathways helps to clarify the roles of the metal and acid components in dehydrocyclization and other reactions. The most important reactions of alkylbenzenes over dual-functional catalysts exhibiting acidic and dehydrogenation activities are hydrogenation, dehydrogenation, isomerization, cyclization, hydrogenolysis, and cracking. To elucidate the mechanisms of these reactions, Csicsery reacted n-butylbenzene (13), n-pentylbenzene (14, and 2-phenylpentane (14) over catalysts covering a wide range of dehydrogenation and acid activities (Tables 1-111).

297


Catalyst

Platinum Surface Area, Micromole CO/g

Dehydrogenation Activity

Acid Activity

High

None

High

Intermediate

25

Intermediate

Strong

66

None

Strong

None

2% platinum on silica gel 0.75% platinum on alumina 2% platinum on silica-alumina Silica-alumina

TABLE I Reactions of n-Butylbenzene over Different Catalystfib Catalyst

Platinum (2%) on silica gel

Platinum (0.75%) on alumina

Platinum (2%) on silicaalumina

Silicaalumina

Reaction temperature (“C) Products (moles per 100 moles of feed)

427

371

427

482

~

Methane Ethane Propane, propylene Butanes and butenes 1-Phenylbutenes 1-Methylindan Methylindenes Naphthalene sec-Butylbenzene Isobutylbenzene Benzene Toluene Ethyl benzene n-Propylbenzene Other Total n-butylbenzene converted (mole%) Csicsery (13, 29).

0.91 1.57 1.07 0.58 1.99 8.04 2.89 11.87 0.36 0.31 0.61 1.12 1.37 1.02 0.34 29.93

0.46 0.57 0.87 1s o 0.40 2.12 0.21 1.29 0.10 0.12 1.57 0.53 0.44 0.41 0.06 1.24

0.42 0.39 0.82 0.37 1.85 7.28 2.38 0.91 0.73 0.54 0.63 0.28 0.26 0.17 2.64 17.67

0.53 -

0.41 7.54 -

0.77 0.06 0.02 0.08 0.03 7.94 0.15

0.05 0.01 0.46 9.57

’Reaction conditions: Atmospheric total pressure, an LHSV of 9, and an initial hydrogen-ton-butylbenzene mole ratio of 3.

298

SIGMUND M. CSICSERY

Dehydrocyclization of n-butylbenzene produces 1-methylindan, methylindenes, and naphthalene:

j-

Naphthalene

H*

1 -Methylindan

1 -Methylindene

3-Methylindene

The primary products of n-pentylbenzene cyclization are 1 -ethylindan, 1-ethylindenes, and 1 -methylnaphthalene.Secondary (consecutive) isomerization may produce dimethylindans, 2-ethylindan, the corresponding indenes, and 2-methylnaphthalene:

n-Pentylbenzene

1 -Methylnaphthalene

2-Methylnaphthalene

1 -Ethylindan

1 -Ethylindene

3-Ethylindene

2-Ethylindan

Dimethylindans

Dimethylindenes

2-Ethylindene


299

The primary products of the cyclization of 2-phenylpentane are cis- and truns-l,3-dimethylindans,1,3-dimethylindene,and 1-methylnaphthalene :

2-Phenylpentane

1-Methylnaphthalene

Cyclization to five-membered rings forms the alkylindans and indenes ; cyclization to six-memberedrings gives tetralins and naphthalenes. Tetralins and decalins, however, were not observed in any of the experiments, because of unfavorable equilibrium. For example, less than 0.02-0.1% tetralin and less than 0.001% decalins would be expected if they were in equilibrium with the naphthalene formed in the n-butylbenzene experiments. [Equilibrium conversions calculated from the data of Egan (IS),Allam and Vlugter ( I d ) , and Frye and Weitkamp (IQ.1 In the cases of n-butylbenzene and 2-phenylpentane, there is an additional difference (other than ring size) between cyclization to naphthalenes and alkylindans. The reaction forming naphthalenes involves the addition of a primary carbon atom to the aromatic ring, while formation of alkylindans (and alkylindenes) involves secondary carbon atoms. With n-pentylbenzene, however, secondary carbon atoms are involved in both five- and six-ring cyclizations.

A. METAL-CATALYZED CYCLIZATION Cyclization to both five- and six-membered rings occurs over platinum on silica gel. Cyclization to six-membered rings has higher activation energy than cyclization to five-membered ring products. The (alkylindan plus alkylindenes)/naphthalene ratios therefore decrease with increasing temperature (Fig. 3). The activation energy difference is about 8-9 kcal/mol, in good agreement with the results of Kazanskii and Liberman (18). The dehydrocyclization rate constants (k, for five-membered ring cyclization,

300

SIGMUND M. CSICSERY

TABLE I1 Reactions of n-Pentylbenzene over Different Catalysts"*b Catalyst


Platinum (0.75%) on alumina


Silicaalumina

Reaction temperature ("C) Product composition (moles per 100 moles of feed) Unreacted n-Pentylbenzene Benzene Toluene Ethyl benzene n-Propyl benzene Indan n-Butylbenzene 1-Methylindan and 1-methylindene Naphthalene Phenylpentanes Phenylpentenes I-Ethylindan and 1-ethylindene Dimethylindans and dimethylindenes 1-Methylnaphthalene 2-Methylnaphthalene Methane Ethane, ethylene Propane, propylene Butanes, butenes Pentanes, pentenes

427

371

427

427

77.06 0.43 0.38 0.68 0.33 0.25 0.34 0.53 0.23 1.02 2.89 6.85 3.32 4.43 0.10 0.97 0.46 0.45 0.35 0.39

89.29 0.92 0.42 0.56 0.18 0.04 0.20 0.15 0.03 0.28 0.58 3.40 1.oo 2.70 0.12 0.38 0.17 0.50 0.36 0.61

35.34 1.07 0.25 0.35 0.10 0.14 0.08 0.26 0.32 0.55 0.60 1.50 6.95 29.42 21.54 0.70 0.24 0.33 0.25 0.62

94.15 4.00 0.06 0.10 0.01 -

0.006 0.02 -

0.20 0.03 0.33 0.8 1 0.05 0.09 0.01 0.09 0.10 0.22 3.80

Csicsery (14,29). Reaction conditions: Atmospheric total pressure, an LHSV of 6.8, an initial hydrogen-tofeed ratio of 3.0, and an initial hydrocarbon partial pressure of 0.25 atm.

k, for six-ring cyclization) of three model hydrocarbons are compared in Table IV, assuming first-order kinetics. A comparison of dehydrocyclization rates over platinum on silica and other nonacidic platinum catalysts suggests the following. 1. Platinum-catalyzed cyclization of alkylaromatics is faster than the cyclization of paraffins because the presence of the aromatic ring enhances the rate. The rate of dehydrocyclization further increases with the number of aromatic rings in the feed molecule. A comparative study of the de-

301

DEHYDROCYCLIZATION OF ALKYLAROMATICS TABLE I11 Reactions of 2-Phenylpentane over Different CatalystfPb Catalyst



Silicaalumina

Reaction temperature (“C) Product Composition (moles per 100 moles of feed) Unreacted 2-Phenylpentane Benzene Toluene Ethyl benzene n-Propylbenzene Cumene n-Butylbenzene Phenyl-Methylbutanes 3-Phenylpentane Dimethylindans and dimethylindenes 1-Methylnaphthalene 2-Methylnaphthalene Methane Ethane, ethylene Propane, propylene Butanes, butenes Pentanes Pentenes ~~

371

37 1

427

96.27 0.32 0.03 0.19 0.02 0.10 0.03 0.69 1.45 0.20 0.06 0.1 0.2

44.66 29.20 0.20 4.06 0.08 0.16 0.10 3.71 2.16 11.96 0.27 0.25 0.9 0.2 2.4 24.5 1.2

13.73 80.80 0.29 0.72 0.07 0.34 0.07 0.33 1.11 1.03 0.03 0.04 0.2 3.8 5.5 68.7

0.26 0.02

~

Csicsery (14,29). Reaction conditions: Atmospheric total pressure, an LHSV of 6.8, an initial hydrogen-tofeed ratio of 3.0, and an initial hydrocarbon partial pressure of 0.25 atm. a

’

hydrocyclization of n-octyl- and dodecylbenzenes, and of 2-n-butyl- and 2-n-octylnaphthalene, shows that alkylnaphthalenes cyclize faster than alkylbenzenes (19). 2. The rate of dehydrocyclization is probably often inhibited by (bicyclic) product desorption. Apparent first-order rate constants in the dehydrocyclization of n-butylbenzene over platinum-on-silica-gel decreased with increasing levels of bicyclic aromatic product (Fig. 4). Sinfelt and co-workers also found that product desorption is rate-controlling in methylcyclohexane aromatization (20). 3. Relative cyclization rates of paraffins and alkylbenzenes may be

302

SIGMUND

22

-

20

-

y

M. CSICSERY

18-

W -I

a

$ 16n

a z

>l4v)

W

z

s+

z

8-

c3 4

P

F' I

6-

I-

W

2 4-

1000/T°K I

300

350

I

I

I

400

450

500

'C

TEMPERATURE FIG. 3. Cyclization of n-butylbenzene. Methylindan/naphthalene ratios. According to Csicsery (I3,29).

influenced by electronic, factors. Davis and co-workers modified catalytic selectivity for the formation of bicyclic aromatics from n-nonane and n-decane by adding either tin or sulfur compounds to platinum on neutralized alumina (21). While total dehydrocyclization rates are not affected in the case of nonane, and are increased in the case of decane, both sulfur and tin decrease the formation of bicyclic aromatics by more than a factor of ten (Tables V and VI). This selectivity change parallels a change in the ethylbenzenelortho-xylene ratio in n-octane dehydrocyclization. The ethylbenzenelortho-xylene ratio changes from 1.0 over pure platinum to 0.5 over platinum-tin (1 :4) or thiophene-doped catalyst. Davis and co-workers suggest that tin or sulfur modify the electronic properties of platinum. As a

303


TABLE IV Comparison of the Cyclization of Alkylbenzenef n-Butylbenzene First-order cyclization rate constants

n-Pentylbenzene

2-Phenylpentane

371

421

371

421

311

0.12 0.08 1.46

0.48 0.52 0.92

0.14 0.05 2.8

0.31 0.13 2.3

0.034 0.0047 7.3

0.18 0.008 22.6

0.39 0.04 10.5

0.14 1.17 0.12

0.36 2.08 0.17

Platinum on silica gel

k, k6 kdk6 Platinum on silica-alumina

a

0.45 0.025 18

Csicsery (13, 14,29).

2:

0.20

v)

2-

-

0

0.16

N -I > V

u

LL

0

0.12

t\ \

a 0.04 U W

c3

m

0 Iu)

LLL z o

NAPHTHALENE + METHYLINDENES FORMED, MOLE %

FIG.4. Dehydrocyclization of n-butylbenzene over 2% platinum on silica gel catalyst. (-) Cyclization to methylindan and methylindenes and (---) cyclization to naphthalene. According to Csicsery (23).

304

SIGMUND M. CSICSERY TABLE V Efjkcis of Tin and Thiophene on Bicyclic Aromatic Formation over Platinum CatalysPpb n-Nonane

Feed hydrocarbon Reagent added to platinum catalyst Total aromatics formed (%) Total bicyclic aromatics (%) Bicyclics (% of total aromatics)

None Thiophene 42 21 50

45 2 4

n-Decane Tin 66 1.2 2

None Thiophene 23 15 40

45 4.3 10

Tin 12 4.8 7

Davis et al. (24. Reaction conditions: Temperature, 482°C; Onstream time, 50 minutes.

TABLE VI The Effect of Thiophene on the Aromatization of n-Nonane over Nonacidic Platinum on Alumina-K Catalyst in the Presence of HZEsb Thiophene in n-Nonane feed (mol %) Time onstream (minutes) Thiophene in product (mol %) Total conversion to aromatics ( m o l x ) C, aromatics produced (mol% of total aromatics produced) n-Propylbenzene 1-Methyl-2-ethylbenzene 1-Methyl-3-ethylbenzene 1-Methyl-4-ethylbenzene Indan Indene

0 37 0 40

25.6 47.1 6.1 0.6 15.2 5.5

15.1

39 1.2 45

30.3 65.5 Trace Trace 4.2 Trace

* Davis et al. (21). Reaction conditions: Temperature, 482°C; LHSV, 0.3.

result, the monocyclic aromatics produced from nonane or decane desorb before they can undergo a second cyclization. 4. C5- and C,- cyclizations are parallel reactions. Csicsery has shown that isomerization of tetralin to methylindan over platinum-alumina at 371°C is extremely slow (22). Davis and Venuto provided further evidence by showing that methylindan is also not converted to tetralin or naphthalene over platinum on silica-alumina (23). This behavior is similar to that observed in the cyclization of aliphatic hydrocarbons. Davis and Venuto also reported that the major aromatic products obtained from ten C8-C, paraffins and olefins at 482°C are only formed by direct six-membered ring


305

closure (24). This also proves that C, C, ring interconversion is very limited in the absence of an acidic catalyst. 5. Other things being equal (such as in the case of n-pentylbenzene, where secondary carbons are involved in both five- and six-ring cyclizations), platinum-catalyzed cyclization favors five-membered ring products over six- membered rings (14). However, the difference between the two reactions is small and may be inverted by changing process conditions (e.g., by increasing the reaction temperature or decreasing the partial pressure of H2). 6 . Over platinum-on-carbon catalyst at relatively low temperature (31OOC), C,-cyclization of alkylbenzenesprobably proceeds by direct closure of the ring between the carbon atoms of the side-chain and the benzene ring, bypassing dehydrogenation to olefins (25-27). However, at higher temperatures and on platinum-alumina or platinum-on-silica C,-dehydrocyclization could involve olefinic intermediates (7, 13, 28). 7. Platinum catalyzes at least two types of c6-dehydrocyclization, one of which involves olefinic intermediates (13, 28, 29). In the case of paraffins, this latter reaction involves the ring-closure of hexatrienes (30, 31). In the C,-dehydrocyclization of n-butylbenzene and n-pentylbenzene, phenylbutadiene and phenylpentadiene could correspond to these triene intermediates (13, 14). The second C,-dehydrocyclization mechanism is similar to C,-dehydrocyclization, and may not involve olefinic intermediates. 8. The methyl group in the y-position of the side-chain interferes with cyclization. The rate of C,-cyclization of n-butylbenzene at 371"C over platinum-on-silica gel is 3.5 times higher than that of 2-phenylpentane (Table IV) (14). The difference in C,-cyclization is even larger. Now, the side-chain carbon atoms involved in the cyclizations of n-butylbenzene and 2-phenylpentane have identical natures (i.e., secondary in five-membered ring closure and primary in six-membered ring closure). The difference between the two molecules (the extra methyl group) is far removed from the two carbon atoms involved in the formation of the new bond:

Adsorption geometry could cause the observed rate difference. Crawford and Kemball have found that deuterium exchange of the two methyl groups of the isopropyl side-chain in cumene over nickel films occurs in two steps (32). Both methyl groups cannot exchange at the same time. The cumene molecule probably must leave the surface and re-adsorb before the second

306

SIGMUND M. CSICSERY

methyl group can exchange. A similar situation might exist over platinum. The phenyl ring is probably held by n-bonding parallel to the surface of the metal (32).This necessarily limits the configurations in which the side-chain can adsorb to two: with either the methyl group or the propyl group pointing away from the surface. Cyclization is possible only in the first case. If adsorption could happen both ways, the rate of cyclization of 2-phenylpentane would be half that of n-butylbenzene. A slight preference for adsorption of the methyl group will further decrease the rate of cyclization. 9. Alkyl side-chains may be removed from the new ring of the bicyclic product during dehydrocyclization. The relative proportion of bicyclics dealkylated increases with rising temperature. For example, in the dehydrocyclization of n-pentylbenzene over platinum-on-silica at 482"C, 12% of the cyclization product is dealkylated (14). Hydrogenolysis is the probable mechanism that removes short (one- or two-carbon-long) side-chains. There is more such dealkylation over platinum-on-silica catalyst than over platinum on (the very acidic) silica-alumina. Isomerization of alkyl side-chains around the ring, however, is very slow over platinum-on-silica and other nonacidic platinum-containing catalysts. For example, in the dehydrocyclization product of n-pentylbenzene, the amount of 2-methylnaphthalene is only 2-3% of the total methylnaphthalenes formed. Of course, over platinum on acidic carriers, isomerization of the side-chain around the ring is very fast. 10. Over platinum-on-silica catalysts, different alkylindans are at equilibrium with the corresponding alkylindenes. Similarly, I -methyIindene is at equilibrium with 3-methylindene (13, 14). 11. Neither C,-nor C,-cyclization involve carbonium-ion intermediates over platinum metal. The rates of the n-propylbenzene + indan reaction (where the new bond is formed between a primary carbon atom and the aromatic ring) and the n-butylbenzene + 1-methylindan reaction (which involves a secondary carbon atom) are quite similar (13). Furthermore, comparison of the C,-cyclization rates of n-butylbenzene and n-pentylbenzene (forming naphthalene and methylnaphthalene, respectively) over platinum-on-silica catalyst shows that in this reaction a primary carbon has higher reactivity than a secondary carbon (Table IV) (29). Lester postulated that platinum acts as a weak Lewis acid for adsorbed cyclopentenes,creating electron-deficient species that can rearrange like carbonium ions (33). The relative cyclization rates discussed above strongly contradict Lester's cyclization mechanism for platinum metal. B. CYCLIZATION OVER DUAL-FUNCTION CATALYSTS Cyclization selectivities are very different over platinum on silica-alumina than over platinum on silica gel (Table IV). In the case of n-butylbenzene, for example, methylindan/naphthalene ratios differ by about an order of


307

magnitude. Over platinum-on-silica-alumina, rates of cyclization to sixmembered rings for n-pentylbenzene are 50 to 100 times higher than for n-butylbenzene. As a consequence alkylindan/methylnaphthalene ratios in the reaction product of n-butylbenzene are 60 to 190 times higher than in the product of n-pentylbenzene reactions. Cyclization rates and k,/k, ratios of 2-phenylpentane are similar to those of n-butylbenzene, but differ significantly from those of n-pentylbenzene (13, 14). These rate differences show that, over platinum on silica-alumina, two cyclization reactions occur simultaneously. One is catalyzed by the platinum metal; it is the mechanism observed over platinum on silica. Acid catalyzed self-alkylation is the second reaction. The following steps are involved. Dehydrogenation of alkylbenzene over the platinum component yields phenylalkenes. Protonation of the phenylalkene over the acid component forms a carbonium ion. (A phenylallyl cation may be produced by proton addition to phenylbutadiene or by hydride ion removal from a phenylalkene.) Attack of this carbonium ion on the aromatic ring closes either a five- or six-membered ring. Stabilization of the product occurs by proton elimination or hydride abstraction. This step may be followed by dehydrogenation to the thermodynamically most-stable species (e.g., to an alkylnaphthalene in the case of six-membered ring closure). (See p. 308.) The stability of the intermediate carbonium ion determines whether cyclization forms five- or six-membered rings. In the case of n-butylbenzene and 2-phenylpentane, acid-catalyzed six-ring cyclization would involve very unstable primary carbonium ions: C6H5-CH2-CH,-CH2-CH,+ or C,H5CH(CH3)-CH,-CH,-CH, +. Therefore, only five-membered rings could be formed in acid-catalyzed cyclization from these two hydrocarbons. On the other hand, the n-pentylbenzene + methylnaphthalene reaction proceeds through C,H,-CH2-CH,-CH2-C+H-CH3, a secondary carbonium ion. As a consequence, acid-catalyzed cyclization produces both five- and six-membered rings. The possible carbonium ion intermediates leading to five- or six-membered ring closure may have similar structures (e.g., both are secondary carbonium ions, as in the case of n-pentylbenzene). If so, acid-catalyzed cyclization favors six-membered ring products, as shown by the k5/k6 ratios (Table IV). About half of the 1-methylnaphthalene formed from n-pentylbenzene and 2-phenylpentane isomerizes to 2-methylnaphthalene over platinum on silica-alumina (while over platinum on silica less than 3% of the methylnaphthalene isomerizes to 2-methylnaphthalene). Alkylindan (and alkylindene) isomerization is also considerable over platinum on silica-alumina (13, 14). Platinum on alumina has properties between those of platinum on silica gel and platinum on silica-alumina. Only about one-fifth of the methylindan is produced by the acid-catalyzed route (13). Also, the isomerization of


309

1-methylnaphthalene is substantially slower over the alumina-supported catalyst than over platinum on silica-alumina (14). Cyclization of alkylbenzenesis much slower over silica-alumina than over the platinum-containing catalysts. To clarify the successive steps of cyclization, Csicsery performed a set of experiments with 4-phenyl-l-butene, a dehydrogenation product of n-butylbenzene (28). Cyclization of phenylbutenes over the platinum-on-silica gel catalyst almost exactly parallels that of n-butylbenzene. Since rapid hydrogenation of the phenylbutenes results in about the same n-butylbenzenelphenylbutene ratio as observed with the n-butylbenzene feed, this is hardly unexpected. Over silica-alumina, the rate of cyclization of 4-phenyl-1-butene is about 2000 times higher than that of n-butylbenzene. Methylindan is the product in both cases. This reconfirms the mechanism proposed above for the acidcatalyzed cyclization of alkylbenzenes: a self-alkylation process involving carbonium-ion intermediates. Thus, alkylbenzenes cyclize over silicaalumina by first dehydrogenating to phenylalkenes (by thermal dehydrogenation or by acid-catalyzed hydrogen transfer). Carbonium ions are formed when the olefinic bonds are protonated. Attack on the aromatic ring by the electron-deficientcarbon atom of the side-chain completes the cyclization reaction. In the absence of a dehydrogenation component (such as for pure silica-alumina), the slow formation of phenylalkenes is the rate-limiting step. This mechanistic interpretation is based on the assumption that, once formed, five- or six-membered products of dehydrocyclization do not undergo interconversion. As discussed above, isomerizations are extremely slow at 317°C for tetralin to methylindan and methylindan to tetralin over alumina, silica-alumina, platinum-on-alumina, and platinum-on-silicaalumina catalysts (22, 23). C. REACTIONS THAT ACCOMPANY DEHYDROCYCLIZATION The cylization of alkylaromatics over platinum catalysts is usually accompanied by isomerization, hydrogenation and dehydrogenation, fragmentation (i.e., hydrogenolysis and cracking), and other reactions. Isomerization over the neutral platinum-on-silica gel catalyst proceeds by two different mechanisms. 1. C,-cyclic intermediates are involved in the 1-methyl-2-ethylbenzeneP n-propylbenzene isomerization :

310

SIGMUND M. CSICSERY

Similarly, n-butylbenzene may be converted to sec-butylbenzene (13), and n-pentylbenzene to 2- and 3-phenylpentanes (14):

I

I

and

\ CNC

I

c

2-Phenylpentane

ac;c CI

c,C

C c\ 3-Phenylpentane

2. Noncarbonium-ion-type 1-2-methyl shifts have been described by Barron et al. ( I ] ) , and by Anderson and Avery (34). The reaction proceeds through a-y-diasorbed intermediates over platinum on neutral supports and does not involve carbonium-ion intermediates. According to this mechanism, the n-butylbenzene i$isobutylbenzene reaction involves the following steps (surface sites are represented by *) :

All side-chain isomers are formed in acid-catalyzed isomerization. Carbonium ions are the intermediates here. Over dual-function catalysts, such as platinum-on-alumina and platinum-on-silica-alumina, platinum increases the rate of isomerization by dehydrogenating alkanes to olefins. This facilitates the formation of carbonium ions.


DEHYDROGENATION

n-PROPYLBENZENE

ETHYLBENZENE

>0

TOLUENE

.r

I

+'4'10

OR '4%

BENZENE

/

SECONMRFEL~TYLEENZENE FIG.5. Catalytic reactions of n-butylbenzene.According to Csicsery (I3,29).

31 1

312

SIGMUND M. CSICSERY

There are two catalytic fragmentation procedures : hydrogenolysis and cracking. In the former a molecule of hydrogen is added, and no new unsaturated bond is formed. In cracking, one of the fragments is formed with an additional unsaturated bond. However, differentiating between the two reactions is not always simple. In the presence of hydrogen over a hydrogenating catalyst, the olefinic products of cracking could become saturated ; or vice versa, the paraffinic products of hydrogenolysis could be dehydrogenated. Hydrogenolysis of the different side-chain bonds of n-alkylbenzenes occurs at an approximately equal rate over platinum on neutral supports. On secondary alkylbenzenes, the bond next to the phenyl ring is cleaved preferentially. Fragmentation of alkylbenzenes over silica-alumina occurs exclusively by acid-catalyzed cracking. The reaction selectively cleaves the bond between the phenyl ring and the a-carbon of the side-chain. This occurs more than 100 times more often than the cracking of all the other bonds combined. Cracking rates of secondary alkylbenzenes are about an order of magnitude higher than those of n-alkylbenzenes. Acid-catalyzed cracking and platinum-catal yzed hydrogen01ysis proceed simultaneously over dual-function catalysts. The distribution of the scission products is determined by the relative strengths of the acidic and metal-type catalytic components. These parallel and consecutive reactions for n-butylbenzene are shown in Fig. 5.

IV.

Double Cyclization of C, and Higher Paraffins

Paraffins with more than eight carbon atoms can dehydrocyclize to form bicyclic products. According to Shuikin and Bekauri, bicyclic products can be formed from paraffins by either successive dehydrocyclization or by simultaneous closure of several carbon-carbon bonds (35). The second possibility follows Balandin’s “sextet” model (36). A large number of hydrocarbons follow the consecutive mechanism (27).Thus far there is no evidence for simultaneous closure. Il’in and Usov have shown that over acidic platinum-alumina catalyst there are at least two consecutive pathways from n-nonane to indan (37): either through alkylaromatic intermediates, by first closing the six-membered ring; or through alkylcyclopentane intermediates, by first closing a fivemembered ring (Table VII and Fig. 6).

313


TABLE VII Some Cyclic Products Formed in the Aromatization of n-Nonant? ~

~

~~

Temperature ("C) Concentration in Product (wt. %) Unreacted n-nonane 1 -Methyl-2-n-propylcyclopentane 1,2-DiethyIcyclopentane n-Bu tylcyclopentane n-Propylbenzene 1-Methyl-2-ethylbenzene Indan

340

400

460

490

73.4 2.7

41.7 2.5

17.7 0.9

0.7 0.8

1.1 3.0

25.4 1.3 0.4 1.1

2.3 5.0

5.6

9.3

0.4 9.8

11.8

17.4

0.5

4.9

12.6

Il'in and Usov (37).

FIG.6. The conversion of n-nonane to indan.

0.3

20.4 14.5

314

SIGMUND M. CSICSERY

A third possible pathway could yield indan through cyclononane intermediate. We know that cyclononane undergoes transannular dehydrocyclization over platinum-on-charcoal catalyst at 300°C, to perhydroindan and then to indan (38);but so far there is no evidence for the direct cyclization of n-nonane to cyclononane. Unfortunately, Il’in and Usov used an acidic catalyst and we cannot separate the contributions of acid and metal catalysis to the two mechanisms. Experiments over nonacidic platinum catalysts could show the relative importances of the platinum metal in the two cyclization pathways. Kazanskii and co-workers have described an interesting special case, double cyclization of n-octane at 310°C over platinum-on-charcoal catalyst at 0.2 liquid hourly space velocity. The reaction product contains about 0.25% cis-octahydropentalane and 2.2-4.5% alkylcyclopentanes (an approximately 1 : 1mixture of trans-1-methyl-2-ethylcyclopentaneand n-propylcyclopentane) (39, 40). Indirect evidence suggests that most of the octahydropentalane is formed from l-methyl-2-ethylcyclopentane,which cyclizes significantly faster than n-propylcyclopentane.

l-Methyl-2ethylcyclopentane

cis -0ctahydropentalane (3,3,0-bicyclooctane)

n-Octane

n-Propylcyclopentane

V.

The Dehydrocyclization of Alkylbenzenes Over Chromia-Alumina Catalysts

It is interesting to compare the dehydrocyclization activity of platinum with that of chromia-alumina. Pines and Goetschel reacted different butylbenzene isomers over acidic and nonacidic chromia-alumina catalysts between 480°C and 492°C (44.Dehydrocyclization is much slower over

315

DEHYDROCYCLIZATION OF ALKnAROMATICS

chromia-alumina than over platinum-containing catalysts. No methylindan is produced from n-butylbenzene over the nonacidic chromia-alumina ; the only bicyclic products formed are tetralin and naphthalene, showing that nonacidic chromia-alumina does catalyse C,-dehydrOCyCliZatiOn. A moderate amount of methylindan is formed from secondary- and isobutylbenzenes over this catalyst. Apparently, nonacidic chromia-alumina catalyzes C,-dehydrocyclization only when the new bond is formed between a primary carbon atom and the aromatic ring. We know that the isomerization of alkylbenzenes over nonacidic chromia-alumina involves free-radical intermediates and proceeds by phenyl or vinyl migration (41,42).One can speculate that dehydrocyclization also has a free-radical mechanism here. Methylindan and smaller amounts of tetralin and naphthalene are formed from all three butylbenzene isomers over the acidic chromia-alumina catalyst. These reactions proceed by a cationic mechanism. VI.

The Dehydrocyclizationof Alkylnaphthalenes

The rate of dehydrocyclization increase with the number of aromatic rings in the molecule (29). The dehydrocyclization of alkylnaphthalenes can follow the same pathways as the cyclization of alkylbenzenes: C,-dehydrocyclization gives benzindans and benzindenes, while C,-dehydrocyclization yields anthracenes and phenanthrenes. In addition to these two pathways, a-substituted alkylnaphthalenes can cyclize to acenaphthenes and acenaphthylenes :

&*Q&-&&!) CH3

CH2

1 -Ethylnaphthalene

1 -Vinylnaphthalene

Acenaphthylene

Acenaphthene

This reaction was first observed by Plate, Erivanskaya, and KhalimaMansur over platinum-on-carbon and platinum-on-alumina catalysts (4348). Platinum-on-carbon catalyzes this reaction between 310°C and 390°C (above which the catalyst is poisoned) (44). Over an acidic platinum-alumina catalyst containing 0.5 wt% platinum and 0.1 wt% sodium, 16.7% acenaphthenes and 1.5% acenaphthylene are obtained at 460°C and at 0.4 liquid hourly space velocity in hydrogen diluent. Conversions are considerably lower in helium.

316

SIGMUND M. CSICSERY

Kinetic studies at short residence times fist suggested the following reaction sequence: ethylnaphthalene dehydrogenates to vinylnaphthalene ; vinylnaphthalene dehydrocyclizes to acenaphthylene ; and finally acenaphthylene is hydrogenated to acenaphthene. However, further work by Isagulyants and co-workers, using ''C-labeled 1-vinylnaphthylene, shows that over platinum on alumina at 470"C, acenaphthene and acenaphthylene are formed from both 1-ethylnaphthalene and 1-vinylnaphthalene. Vinylnaphthalene dehydrocyclizes about three times faster than ethylnaphthalene. The vinylnaphthalene intermediate remains adsorbed on the catalyst surface during the reaction (48). 1-Propylnaphthalene can give two types of products by C5-dehydrocyclization: dehydrocyclization at the peri-carbon atom of naphthalene gives 1-methylacenaphthene and 1-methyl-acenaphthylene, while dehydrocyclization involving the b-carbon atom of naphthalene gives 4,Sbenzindan and 45benzindene : 7H3

b-b+d+& CH2

Peri-dehydrocyclization

p-dehydrocyclization

Hydrogen has an important directing effect here. In the presence of hydrogen over 0.5 wt% platinum on y-alumina, Erivanskaya and co-workers found that the rate of peri-dehydrocyclization of 1-propenylnaphthalene is about four times faster than the rate of P-C5-dehydrocyclization (49). However, if the catalyst, after the usual reduction, is treated with helium for three hours and the experiment is also carried out in helium, the rate of peri-dehydrocyclization decreases by about a factor of seven. The rate of j-dehydrocyclization does not change significantly. The rates for both types of C,-dehydrocyclization increase with the acidity of the alumina (50,51).

No investigator has observed C,-dehydrocyclization involving the pericarbon atom of naphthalene. Phenalene and 2,3-dehydrophenalene have not been detected over any of the catalysts investigated. The C,-dehydrocyclization of 2-n-butylnaphthalene can give 4 5 or 5,6-methylbenzindans and benzindenes. Similarly, C,-dehydrocyclization can yield either phenanthrene or anthracene. The product distribution depends on reaction temperature and catalyst type. The C,-dehydrocyclization of 1-(2-naphthyl)-butenedepends on the acidity of the alumina, but


317

C,-dehydrocyclization does not (50). As a consequence, over nonacidic platinum catalysts above 400"C, C,-dehydrocyclization predominates over C,-dehydrocyclization (27). Furthermore, the phenanthrene/anthracene ratio is independent of catalyst acidity (52).The effect of reaction temperature is, however, very interesting. Over platinum-on-carbon catalyst between 350°C and 400"C, more anthracene is formed than phenanthrene. Above 450°C phenanthrene is the main product (53).Phenanthrene is also the main product over chromia-alumina between 360°C and 440°C; whereas, as seen above, anthracene is formed in this temperature range over platinum-oncarbon catalyst (54). We know that C,-cyclization of l-(naphthyl-2)-butene is possible without metal catalysts. The products are dihydrophenanthrene over quartz and 1,2,3,4-tetrahydrophenanthreneplus phenanthrene over alumina (50). The latter apparently catalyzes the redistribution of hydrogen in dihydrophenanthrene. Neither anthracene nor dihydro- or tetrahydroanthracene are formed over quartz or alumina from l-(naphthy1-2)-butene. Plate and Erivanskaya concluded from this that the 2-alkylnaphthalene + anthracene reaction does not involve naphthylbutene intermediate (27). The a-position in naphthalene (and other condensed polycyclic aromatics) is sterically hindered. Hodges and Garnett have shown, for example, that at 100°C and in the presence of Pt(I1)-salts the p-hydrogen exchanges with deuterium 28 times faster than the a-hydrogen (55). This could suggest that Pt-catalyzed direct C,-cyclization will also favor the P-position. Such steric effects decrease with increasing temperature. On the other hand, thermal ring closure of 2-alkenylnaphthalenes occurs exclusively with the a-carbon, giving phenanthrene. More 2-alkenylnaphthalene is formed at higher temperatures. This, combined with the change of steric effect, determines anthracene/phenanthrene ratios at different temperatures (27).

If this is true, the simultaneous formation of anthracene and phenanthrene from 2-n-butylnaphthalene gives us an extraordinary and fortunate opportunity to differentiate between two types of C,-dehydrocyclization (27). Anthracene might be the product of direct cyclization, a mechanism related

318

SIGMUND M. CSICSERY

to C,-dehydrocyclization. Phenanthrene is probably formed through dehydrogenated intermediates. This mechanism probably corresponds to the cyclization of hexatriene, as described by P d l and Tetenyi (30, 31). The significance of the dependence of the anthracene/phenanthrene ratio on temperature and catalyst type cannot be overemphasized; one wishes that there were more experimental data available on the effects of hydrogen partial pressure and other variables, to give more support to this hypothesis. Erivanskaya and co-workers also studied the dehydrocyclization of 2-nbutylnaphthalene over supported palladium, rhodium, and iridium catalysts (56-58). Palladium-alumina showed the lowest C,-dehydrocyclization activity, but was the most active for the C,-dehydrocyclization of 2-n-butylnaphthalene. A later study showed, however, that this enhanced activity was due to the high chlorine content of the palladium-alumina catalyst and not to some mysterious inherent catalytic activity of palladium (56). The dehydrocyclization activity of rhodium-alumina is lower than that of platinum-alumina. Hydrogenolysispredominates over all the other reactions with this catalyst (57). The effect of temperature on the anthracenelphenanthrene ratio in the product from 2-n-butylnaphthalene is the same over iridium-alumina catalyst as that observed over platinum-alumina : more phenanthrene and less anthracene are formed at high temperatures (58). VII.

Dehydrocyclization of Diphenylalkanes

Scola studied the dehydrocyclization of diphenylmethane, bibenzyl, trans-stilbene, 1,4-diphenylbutadiene, and 1,l'-binaphthyl, over 0.6% platinum-on-silica gel catalyst, at around 550"C, and in the presence of hydrogen (59). Stable, fused polycyclic aromatics were formed through the abstraction of ortho-hydrogens from one or two of the phenyl rings. Diphenylmethane was thus converted to fluorene, with excellent yield (45% at 550°C and 12 seconds contact time). 1,ZDiphenylethane (bibenzyl) was converted to phenanthrene with 92% yield. On the other hand, trans-stilbene gave only 33% phenanthrene at the same conditions. 1,4-Diphenylbutadiene gave a nonvolatile liquid fraction which contained 53% 2-phenylnaphthalene, 16% 1-phenylnaphthalene, 10% fluoranthene, and 21% unidentified other products. 1,l '-Binaphthyl gave a mixture containing 42% benzofluoranthene and 39% perylene. Scola believes that carbonium ion intermediates are involved in all of these dehydrocyclization reactions. This is unlikely, as platinum-silica usually does not have sufficient acidity to generate carbonium ions. Dehydrocyclization follows molecular rearrangement when 1,l-diphenylethane reacts over platinum alumina between 420°C and 540°C in the


319

presence of steam. The product contains 1,l dphenylethylene, trans-stilbene, and phenanthrene. Small amounts of 9-methylfluorene and fluorene are also formed (60). The phenanthrene is probably formed through stilbene intermediates. In the absence of steam, the catalyst is deactivated. VIII.

Conclusions

1. The presence of the aromatic ring enhances the rate of dehydrocyclization; alkylaromatics dehydrocyclize faster than paraffins. Furthermore, the rate of dehydrocyclization increases with the number of aromatic rings within the reacting molecule. 2. Platinum can catalyze C5- and C,-dehydrocyclizations. The two reactions are parallel. Interconversion is very limited between five- and sixmembered ring products over nonacidic platinum catalysts. Platinumcatalyzed dehydrocyclization does not involve carbonium ions. 3. There are at least two C,-dehydrocyclization mechanisms; one of these proceeds through arylalkene intermediates and corresponds to the hexati-iene-type C,-dehydrocyclization of paraffins. The other pathway is direct ring closure. It is probably related to C5-dehydrocyclization. 2Butylnaphthalene may differentiate between the two mechanisms; phenanthrene is probably formed by the first reaction, anthracene by the second. 4. There appear to be two C5-dehydrocyclizations over platinum-oncarbon catalyst. Activation energy differences suggest that the reaction involving an sp2 and an sp3 carbon atom (a cyclization in which the new bond is formed between an aliphatic and an aromatic carbon atom) is different from cyclizations involving two sp3 carbon atoms (in which the new bond is formed between aliphatic carbon atoms of two side-chains). REFEXENCB 1. Moldavskii, B. L., and Kamusher, G. D., Dokl. Akad. Nu& SSSR I, 355 (1936). 2. Moldavskii, B. L., Kamusher, G. D., and Kobylskaya, M. V., Zh. Obshch. Khim. 7, 169 (1937). 3. Herington, E. F. G., and Rideal, E. K., Proc. R. SOC.London, Ser. A 184,447 (1945). 4. Liberman, A. L., Bragin, 0. V., and Kazanskii, B. A., Dokl. Akad. Nauk SSSR 111,1039 (1956). 5 . Liberman, A. L., Bragin, 0. V., and Kazanskii, B. A., Izu. Akad. Nauk SSSR, Ser. Khim. p. 879 (1956). 6. Liberman, A. L., Bragin, 0. V., Ming-Nan, C., and Kazanskii, B. A., Dokl. Akad. Nauk SSSR 129,578 (1959). 7. Shephard, F. E., and Rooney, J. J., J . Catal. 3, 129 (1964). 8. Bragin, 0.V., Gur’yanova, G. K., and Liberman, A. L., Dokl. Akad. Nauk SSSR 160,823 (1 965).

320

SIGMUND M. CSICSERY

9. Bragin, 0. V., and Liberman, A. L., Dokl. Akud. Nuuk SSSR 148,338 (1963). 10. Barron, Y.,Maire, G., Comet, D., and Gault, F. G., J. Cuful.2, 152 (1963). 11. Barron, Y., Maire, G., Muller, J. M., and Gault, F. G., J . Cutal. 5, 428 (1966). 12. Csicsery, S. M., and Burnett, R. L. J. Catal. 8, 75 (1967). 12u. P d l , Z., Dobrovolszky, M., and Tetenyi, P., J . Catal. 45, 189 (1976). 13. Csicsery, S. M., J . Cutal. 9, 336 (1967). 14. Csicsery, S. M., J . Cutal. 15, 111 (1969). 15. Egan, C. J., J. Chem. Eng. Data 8, 532 (1963). 16. Allam, M. I., and Vlugter, J. C., J . Inst. Per., London 52, 385 (1966). 17. Frye, C. G., and Weitkamp, A. W., J. Chem. Eng. Data 14, 372 (1969). 18. Kazanskii, B. A,, and Liberman, A. L., World Pet Congr., Proc., 5th, 1959 Sect. 4, Paper 3, p. 29 (1960). 19. Bekauri, N. G., Shuikin, N. I., and Shakarashvili, T. S., Kinet. Kurd. 8, 1275 (1967). 20. Sinfelt, J. H., Hurwitz, H., and Shulman, R. A,, J. Phys. Chem. 64, 1559 (1960). 21. Davis, B. H., Westfall, G. A., Watkins, J., and Pezzanite, J., Jr., J. Cutul. 42, 247 (1976). 22. Csicsery, S. M., J . Cutal. 12, 212 (1968). 23. Davis, B. H., and Venuto, P. B., J. Cutal. 17, 274 (1970). 24. Davis, B. H., and Venuto, P. B., J . Cutul. 15, 363 (1969). 25. Liberman, A. L., Kinet. Katal. 5, 128 (1964). 26. Bragin, 0. V., Preobrazenskii, A. V., Liberman, A. L., and Kazanskii, B. A., Kinet. K d a l . 16,472 (1975). 27. Plate, A. F., and Erivanskaya, L. A,, Neftekhimiya 16, No. 3, 348 (1976). 28. Csicsery, S. M., J. Curd. 12, 183 (1968). 29. Csicsery, S. M., Intra-Sci. Chem. Rep. 6, No. 2, 43 (1972). 30. PaaI, Z., and Tetenyi, P., J. Catal. 30, 350 (1973). 31. Pail, Z., and Tetenyi, P., Actu Chim. Acad. Sci. Hung. 58, 105 (1968). 32. Crawford, E., and Kemball, C., Trans. Furuduy Soc. 58,2452 (1962). 33. Lester, G. R., J. Cutal. 13, 187 (1969). 34. Anderson, J. R., and Avery, N. R., J. Cutal. 5,446 (1966). 35. Shuikin, N. I., and Bekauri, N. G., Dokl. Akud. Nuuk SSSR 126, 103 (1959). 36. Balandin, A. A,, “Present State of the Multiplet Theory of Heterogeneous Catalysis.” Nauka, Moscow (1968). 37. Win, V. F., and Usov, Yu. N., Neftekhimiya 13, No. 3, 387 (1973). 38. Khromov, S. I., Balenkova, E. S., Lishenok, 0. E., and Kazanskii, B. A., Dokl. Akad. Nuuk SSSR 135, 110 (1960). 39. Kazanskii, B. A,, Liberman, A. L., Loza, G. V., Kuznetsova, I. M., Aleksanyan, V. T., and Sterin, Kh. E., Izv. Akad. Nuuk SSSR, Ser. Khim. p. 1071 (1959). 40. Kazanskii, B. A., Liberman, A. L., Kuznetsova, I. M., Aleksanyan, V. T., Sterin, Kh. E., and Loza, G. V., Dokl. Akad. Nauk SSSR 133,364 (1960); C . R. Acad. Sci. USSR 133,785 (1960). 41. Pines, H., and Goetschel, C. T., J . Cutul. 6,371 (1966). 42. Pines, H., and Goetschel, C. T., J . C u r d 6, 380 (1966). 43. Khalima-Mansur, A., Erivanskaya, L. A., and Plate, A. F., Vestn. Mosk. Univ. Khim. 13. No. 1, 115 (1972). 44. Erivanskaya, L. A,, Khalima-Mansur, A., Grishin, Yu.K., and Plate, A. F., ~ e f t e k ~ i m ~ ~ a 12, No. 2, 183 (1972). 45. Erivanskaya, L. A,, Khalima-Mansur, A., and Plate, A. F., Neftekhimiya 13, No. 1, 27 (1973). 46. Khalima-Mansur, A,, Erivanskaya, L. A,, and Plate, A. F., Neftekhimiya 13, No. 1, 32 (1973).


321

47. Plate, A. F., Erivanskaya, L. A. and Khalima-Mansur, A., Znr. Chem. Eng. 14,83 (1974). 48. Isagulyants, G. V., Greish, A. A., Korovina, L. M., Erivanskaya, L. A,, and Plate, A. F., Neftekhimiya 17, 390 (1977). 49. Erivanskaya, L. A., Khalima-Mansur, A,, and Plate, A. F., Kinet. Karal. 15, No. 3, 810 (1974). 50. Erivanskaya, L. A., Shevtsova, G, A., Khalima-Mansur, A., and Plate, A. F., Kinet. Katal. 15, No. 3, 722 (1974). 51. Erivanskaya, L. A,, Lozhkina, V. P., Korovina, L. M., Antipina, T. V., and Plate, A. F., Nefrekhimiya 15, No. 3, 341 (1975). 52. Erivanskaya, L. A., Shevtsova, G. A. Komissarova, N. L., and Plate, A. F., Neftekhimiya 8, No. 2, 192 (1968). 53. Erivanskaya, L. A., Komissarova, N. L., Shevtsova, G. A,, and Plate, A. F., Neftekihimiya 9, No. 4, 513 (1969). 54. Komissarova, N. L., Erivanskaya, L. A., and Plate, A. F., Neftekhimiya 7, No. 5,725 (1967). 55. Hodges, J. L., and Garnett, J. L., J. Phys. Chem. 73, 1525 (1969). 56. Erivanskaya, L. A., Shevtsova, G. A., and Plate, A. F., Neftekhimiya 14, No. 3,373 (1974). 57. Erivanskaya, L. A,, Shevtsova, G. A,, and Plate, A. F., Neftekhimiya 12, No. 3,329 (1972). 58. Shevtsova, G. A., Erivanskaya, L. A,, and Plate, A. F., Neftekhimiya 13, No. 4,509 (1973). 59. Scola, D. A,, Ind. Eng. Chem., Prod. Res. Dev. 4, No. 2, 136 (1965). 60. Nametkin, N. S., Ryabov, V. D., and Bykov, V. I., Neftekhimiya 17, No. 1, 31 (1977).


.


Meta Iloenzyme Catalysis JOSEPH J . VILLAFRANCA*

AND

FRANK M . RAUSHELt

Department of Chemistry Pennsylvania State University University Park. Pennsylvania

I. Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . .

324

I1. Thermolysin . . . . . . . . . . . . . . . . . . . . . . . . . . . . 326 A. Background . . . . . . . . . . . . . . . . . . . . . . . . . . 326 B. Kinetic Studies . . . . . . . . . . . . . . . . . . . . . . . . . 327 C. Chemical Modification . . . . . . . . . . . . . . . . . . . . . . 328 D. X-Ray Structure . . . . . . . . . . . . . . . . . . . . . . . . 329 E. Physical Studies of the Metal Ion Environment . . . . . . . . . . . 331 F. Conclusions . . . . . . . . . . . . . . . . . . . . . . . . . . 336 111. Yeast Hexokinase . . . . . . . . . . . . . . . . . . . . . . . . . 336 A. Background . . . . . . . . . . . . . . . . . . . . . . . . . . 336 B. Substrate Specificity . . . . . . . . . . . . . . . . . . . . . . . 337 C. Stereochemistryof the Active Complex of MgATP . . . . . . . . . . 339 D. Kinetic Mechanism . . . . . . . . . . . . . . . . . . . . . . . 341 E. X-Ray Structure . . . . . . . . . . . . . . . . . . . . . . . . 344 F. Chemical Mechanism . . . . . . . . . . . . . . . . . . . . . . 348 G . Conclusions . . . . . . . . . . . . . . . . . . . . . . . . . . 349 IV . Glutamine Synthetase . . . . . . . . . . . . . . . . . . . . . . . . 349 A. Background . . . . . . . . . . . . . . . . . . . . . . . . . . . 349 B . Kinetic Mechanism . . . . . . . . . . . . . . . . . . . . . . . 350 C. Analogs of the Transition State or Intermediates in the Reaction . . . . . . . . . . . . . . . . . . . . . . . . 355 D. EPR and NMR Studies of Conformational Changes at the n1 Metal Ion Site . . . . . . . . . . .- . . . . . . . . . . . . . 358 E. Metal-Metal Distances between then, and n, Sites . . . . . . . . . 361 F. Distances between the Adenylyl Site and the Metal Ion Sites . . . . . . . . . . . . . . . . . . . . . . . . . . . 363 G . NMR Studies of Distances to the Feedback Modifier Regulatory Sites and the Glutamate Site . . . . . . . . . . . . . 364 H. Conclusions . . . . . . . . . . . . . . . . . . . . . . . . . . 366 References . . . . . . . . . . . . . . . . . . . . . . . . . . . . 366

* Established Investigator of the American Heart Association.

t National Research Service Awardee (AM.05966) . 323

Copyright 8 1979 by Academic Ress. Inc. All rights of reproduction in any form ICSCNUI . ISBN 012-007828-7

324

JOSEPH J. VILLAFRANCA A N D FRANK M. RAUSHEL

I.

Introduction

Catalysis by metalloenzymes occupies a relatively unique position when placed in the context of heterogeneous and homogeneous catalysis. For heterogeneous catalysis, it is difficult many times to obtain information on the number and types of atoms at the “active site” of the catalyst. For enzymic catalysis, a distinguishing feature is that even if one tends to focus on the metal ion as essential to the catalytic center, unique information can be obtained about changes at this one center in spite of the fact that the metal ion is bound to a large protein molecule of molecular weight from 6000 in the simplest cases to over 600,000 in the more complex metalloenzymes. The primary, secondary, and tertiary structures of the macromolecule surrounding the metal ion, however, make possible an enormous variation in the microenvironment of the metal ion. The microenvironment also consists of amino acids whose side chains (and also perhaps the peptide backbone) can assume a role in a given catalytic reaction in addition to the metal ion. Such variation is not readily achieved in small molecular systems, but many variations in solvent polarity, pH, etc., can be applied to homogeneous catalysts. The chemistry of the metalloenzymes must be considered as a special case of enzymic catalysis since most active sites of enzymes are stereospecific for only one molecule or class of molecules and many do not involve metal ions in catalysis. Since the metal ion is absolutely essential for catalysis in the examples chosen for this review, the mechanisms undoubtedly involve the metal ion and a particular protein microenvironment or reactive group(s) as joint participants in the catalytic event. It is our belief that studies of catalysis by metalloenzymes will have as many, if not more, features characteristic of protein catalysis in general, in a fashion similar to metal ion catalysis, and these studies will be directly applicable to heterogeneous and homogeneous catalytic chemical systems where the metal ion carries most of the catalytic function. Numerous experimental techniques have been applied to study enzymes that possess a metal ion at the catalytic region in the protein molecule (1-3). These techniques for determining structure include nuclear magnetic resonance (NMR), electron paramagnetic resonance (EPR), fluorescence, and x-ray crystallography. Kinetic approaches have been used extensively to study the order of binding of substrates and the release of products. Rapid kinetic techniques are useful in determining individual rate constants. In this review we have chosen enzymes for which many independent experimental approaches have been utilized to obtain the coordinates in space of substrate atoms, metal ions, and various atoms of the protein that are involved in binding and catalysis. With these examples both structure and function have

325

METALLOENZYME CATALYSIS

been studied at the atomic level and provide the insight necessary for a complete understanding of the tremendous catalytic behavior of enzyme molecules. The distinction between metalloenzymes and metal ion-activated enzymes is an empirical one. If an isolated enzyme contains a tightly bound metal ion that remains bound to the protein through purification, then the enzyme is designated a metalloenzyme. Metal ion-activated enzymes are often inactive after purification, but regain activity when the appropriate metal ion is added back. It is possible in many metalloenzymes to substitute the native metal ion with another first transition ion either by removing the original ion with chelating agents or by exchange dialysis. In addition to certain interesting physical properties that can be studied by this method, an opportunity is provided to discover what kind of specificity for catalysis resides in the electronic properties of the metal ion. Studies of this type have been conducted with several metalloenzymes. Metalloenzymes or metal ion-activated enzymes catalyze an enormous variety of organic reactions that are not restricted to any particular reaction class, but appear as catalysts for all types of reactions. Thus neither the presence of the metal ion nor the reaction type seems to be restrictive as far as metal-assisted enzyme catalysis is concerned. In some cases the metal ion appears to function as an electron acceptor or donor, but flavin cofactors have substituted as redox centers during evolution in some enzymes. There is now considerable evidence that indicates that metal ion-activated enzymes and metalloenzymes catalyze reactions via ternary complexes consisting of a 1 : 1 :1 ratio of enzyme: metal ion: substrate. Three coordination schemes are possible : E-S-M

M-E-S

I

I1

I11

These involve the substrate bridge complex (I) wherein the metal ion is not bound to enzyme residues, the enzyme bridge complex (11) in which the metal ion is not involved in bond-making or -breaking steps (or for that matter substrate binding), and the metal ion bridge complex (111). For metalloenzymes the substrate bridge complex is not possible by definition. The metal ion bridge complex must have the substrate bound to the enzyme because all enzymic reactions are stereospecific and the chirality of a reaction could only reside in the substrate molecule and the enzyme groups. To understand an enzyme mechanism thoroughly, which involves a

326

JOSEPH J. VILLAFRANCA AND FRANK M. RAUSHEL

metalloenzyme or a metal ion-activated enzyme, the following points must be known: (1) the affinity of substrates, coenzymes, and cofactors, (2) the rate constants for each step, (3) the three-dimensionalgeometric relationship among the substrates, coenzymes, cofactors, and catalytically important residues of the enzyme, and (4) the chemical mechanism for each step; i.e., the electronic and atomic rearrangements. Many aspects of points (1)-(3) can be answered by magnetic resonance techniques (NMR, EPR, ENDOR), magnetic circular dichroism, rapid flow kinetics, and x-ray crystallography. Thus, a wide variety of physical techniques can be utilized to understand structure-function relationships in enzymology. The nature of the chemical mechanism (bond making and breaking steps) can be investigated in select cases by studying kinetic and thermodynamic isotope effects when "heavy" isotopes replace normal isotopes in substrate molecules (3H or 'H for 'H, 3C for "C, "N for 14N, l8O for l6O, etc.). Chemical modifications of specific amino acids can be useful in determining if these residues are involved in binding substrates or metal ions. Enzymes can therefore be studied in macroscopic and microscopic detail by almost every technique available to scientists. The following detailed discussion of three enzymes that have metal ions at their active sites will point out the current state of the art of enzymologists' understanding of enzymic catalysis. The examples have been chosen to include the most advanced use of stereochemical techniques, kinetic methodology, solution structural data (NMR, EPR, fluorescenceenergy transfer), and x-ray crystallographic structures.

'

II. Thermolysin A. BACKGROUND

Thermolysin is a metalloenzyme isolated from Bacillus thermoproteolyticus. It is a heat-stable extracellular endopeptidase of molecular weight 34,600. The enzyme catalyzes the hydrolysis of peptide bonds that have the amino group as part of hydrophobic residues such as phenylalanine, isoleucine, or leucine. -CH-C-NH-CH-

I

II

I

+--CH--CO,-

I

Ri 0 R,' Rl R, = phenylalanine(best) R,' = phenylalanine, isoleucine, leucine

+ H3N-CHI

R,'

Hydrolysis is fastest when both R, and R; are hydrophobic residues. Latt et al. (4) showed that the isolated enzyme has one Zn" and four Ca2+.


327

The Zn2+ is required for catalysis, while the Ca2+ ions are necessary for thermostability. This enzyme has been studied extensively by x-ray, kinetic, NMR, optical, circular dichroic, and fluorescence techniques. Thus, many approaches have been used to explore the role of the metal ions in catalysis and of other protein residues in substrate binding and catalysis. The review of this enzyme will serve to point out the information to be gained from using multiple biophysical approaches in understanding metalloenzyme catalysis. B. KINETICSTUDIES Thermolysin belongs to a class of proteases (called neutral proteases) which are distinct from the serine proteases, sulfhydryl proteases, metalloexopeptidases, and acid proteases. Neutral proteases A and B from Bacillus subtilis resemble thermolysin in molecular weight, substrate specificity, amino acid content, and metal ion dependence. Since physiological substrates are most likely proteins, it is difficult to design simple experiments that can be interpreted in terms of substrate specificity and relative velocities. Therefore, studies of substrate specificity and other kinetic parameters must be carried out on di- and tripeptides so that details of the mechanism of catalysis can be obtained and interpreted simply. Most kinetic studies of enzymes are conducted under conditions where substrate concentration spans the upper and lower sides of the Michaelis constant, &.The most general form of kinetic expression for a one-substrate enzyme is given as E

+ S * ES

+

product

In kinetic studies with neutral proteases, the substrates are relatively insoluble and X, values are generally high. For these enzymes, pseudo-firstorder kinetic data can be obtained where (5') 4 K,, leading to a rate expression, u = kcat(E)(S)/Km.Values of kcat/& and X, are measured and data compared for various substrates. Table I presents data at pH = 8 for a neutral protease from B. subtilis (5). From these data one can see that the variation in f& is small for the various dipeptide substrates that are amides (-NH,), but that k,,,/K, [proportional to k/(E)]values differ by almost lo3. This indicates that the major change in catalysis is in the catalytic rate constant, k,,,, and that the side chains must bind differently to the enzymes to produce this effect. The dipeptide Gly-Leu-NH, has a free a-amino group and k/(E)is reduced about 10' over Z-Gly-Leu-NH,. A free carboxyl in the substrate (Z-Gly-Leu)

328


TABLE I Hydrolysis of Various Dipeptides by Neutral Protease Substrate"

k/(E) x lo4, S-'(mg/ml)-'

K,,, x 10, (M)

Z-Gly-Ala-NH, Z-Gly-Val-NH, Z-Gly-Leu-NH, Z-Gly-Phe-NH, Z-Ala-Leu-NH, Z-Thr-Leu-NH, GIy-Leu-NH, Z-Gly-Leu

0.43 6.95 54.6 3.47 147.7 406.5 10.51 2 1.38

1.04 2.16 2.94 0.32 1.01 1.22

Z is benzyloxycarbonyl. The -NH, represents the amide of the dipeptide. When -NH, is absent the dipeptide has a free carboxyl group, and when Z is absent the peptide has a free a-amino group.

lowers this value about 40-fold. This study establishes that the enzyme functions best with a bulky nonpolar substituent as amino acid side chain and that the active site will not accommodate charged groups in the peptide backbone or in the peptide chain. The preceding study is directly applicable to understanding the active site of thermolysin, since a recent kinetic study (6)comparing neutral protease from B. subtilis with thermolysin showed that the two enzymes have identical pH rate profiles that peak near pH 7.

-

C.

CHEMICAL MODIFICATION

Chemical modification studies of Pangburn and Walsh (6) and dipeptide inhibition studies of Feder et al. (7) led to the implication in catalysis of an enzymic group or groups with pK values around 7-8. To understand what group may be involved in binding and/or catalysis, one must know the primary sequence of the protein and then solve its x-ray structure. The amino acid sequences of neutral protease A and thermolysin have been reported (8, 9) and a high degree of homology is evident. As will be evident in the following section, histidine-231 is implicated in thermolysin catalysis by x-ray crystallographic studies. A plot of kc,,/& vs. pH for the substrate furylacryloylglycyl-L-leucinamideshows two breaks at pH 5.9 and 7.5, thus implicating two ionizable enzymic groups in catalysis. When histidine-231 is chemically modified by ethoxyformic anhydride, the enzyme is inactivated, suggesting that histidine is the enzymic group with pK = 7.5. The enzymic group with pK = 5.9 is most likely glutamic-143, which serves as a general base to aid in the deprotonation of a water molecule


329

for nucleophilic attack on the carbonyl of a peptide substrate (see the following section). D. X-RAY STRUCTURE Matthews and his collaborators have undertaken an x-ray study of crystalline thermolysin. Using the amino acid sequence and crystallographic data on the enzyme taken to a resolution of 2.3 A (lo),they have elucidated the native enzyme structure. ZnZ is tetrahedrally coordinated by the enzyme and the four ligands are ring nitrogens of histidine-142 and -146, an oxygen of the y carboxyl of glutamic-166, and an oxygen of a water molecule (Fig 1) (10).When the structure is determined in the presence of the dipeptide inhibitor 8-phenylpropionyl-L-phenylalanineand the x-ray difference map is determined, the water molecule on the Zn2+ is displaced by the carbonyl of the dipeptide inhibitor. The structure of an enzyme-substrate complex is shown in Fig. 2, inferred from the inhibitor binding studies. One salient feature of this model of the active site with a bound substrate is that the carbonyl of the peptide bond that is to be broken is coordinated to Zn2+.This feature implies that the metal ion may serve as an electrophile to polarize the carbon-oxygen bond, thereby rendering the carbon succeptible to nucleophilic attack by water. In the x-ray structure a water molecule is +

THERMOLYSIN ACTIVE SITE FIG. 1 . The zinc site of thermolysin. Data adapted from Kester and Matthews (10).

330


FIG.2. Thermolysinactive site. The structure shown is of a substrate bound to the active site Zn2+and amino acid residues. From Kessler and Matthews (10). Reprinted with permission of The American Chemical Society.

HC

' 'CH I

0,

OH

1

,CH,

H

CHZ I

FIG.3. Structure of the inhibitor phosphoramidon.The tetrahedral phosphorus atom binds at the active site of thermolysin and mimics the transition-state complex.

hydrogen-bonded to the y-carboxyl of glutamic acid 143. Glutamic-143 can remove a proton from water, thereby aiding the nucleophilic attack. Another enzyme residue, histidine-231, is near the peptide nitrogen and can protonate this nitrogen during catalysis. Further x-ray studies were conducted with an inhibitor that has a tetrahedral phosphorus atom where the carbonyl of the peptide bond would be located (22). The structure of this inhibitor, phosphoramidon, is given in Fig. 3. This compound binds to the enzyme about 10'-fold better than other dipeptide inhibitors that have been studied (phosphoramidon, K, =


331

NH

f FIG. 4. Proposed transition-state structure for an enzyme-substrate complex with thermolysin. From Kessler and Matthews (10). Reprinted with permission of The American Chemical Society.

2.8 x lo-* M ; P-phenylpropionyl-L-phenylalanine,KI = 1.6 x M). This inhibitor can be considered a good “transition-state” analog of the thermolysin catalyzed reaction. The x-ray difference map shows that the tetrahedral phosphorus binds in the pocket occupied by the carbonyl of a dipeptide and that one phosphorus oxygen binds to the Zn2+while another oxygen displaces the water bound to glutamic-143. The proposed transition-state structure for the catalytic reaction is given in Fig. 4. Overall the mechanism of the thermolysin reaction including the role of the Zn2+in catalysis is given in Fig. 5. This proposed mechanism is based on inhibitor data and x-ray structure determination.

E.

PHYSICAL

STUDIESOF

THE

METALION ENVIRONMENT

1. Absorption Studies

Since Zn2+ has a full 3d shell of electrons (d”), it is a colorless, diamagnetic metal ion. Replacement of ZnZ+by Co2+ (d’) does not produce large

332

JOSEPH J . VILLAFRANCA AND FRANK M. RAUSHEL

H "O\

p I

Asp

226

FIG. 5. Proposed mechanism of action of thermolysin showing the catalytically important residues.

---

Co-corboxypeptidose Co- thermolysin

WAVELENGTH (rim)

FIG.6 . Absorption and magnetic circular dichroism spectra of Coz+ thermolysin. (A) MCD spectra. (B) Absorption spectra.


333

alterations in the protein conformation or structure of thermolysin since the Co2+-substituted enzyme is still catalytically active (12).With Co2+one can study the absorption, circular dichroism (CD), magnetic circular dichroism (MCD), and EPR properties of enzyme-bound C o z + .All of these techniques will provide information on the coordination environment of the metal ion and the changes in the environment upon addition of substrates or inhibitors. Figure 6 shows the absorption and MCD spectra of Co2+-thermolysin (13). The addition of the inhibitor P-phenylpropionyl-L-phenylalanineproduces a change in intensity of both the optical and MCD spectra. These changes suggest alteration of the coordination environment of bound C o 2 + ,and this would be consistent with displacement of the water molecule bound to Co2+by oxygen of the carbonyl, as demonstrated in the x-ray study. 2. Models for Metal Ion Site

One would like to be able to deduce structural information from studies on the enzyme in solution rather than having to rely on the resolution of the complete x-ray structure of the enzyme. An approach to this problem is to synthesize model complexes, to measure their spectra, and to compare these to spectra from the enzyme complex. Ishly and Horrocks (14) have studied

H-

R =C,H,

R=CH,

FIG.7 . CoZf complexes with two oxygen and two nitrogen ligands that are models for the metal ion site of thermolysin.

334


the two complexes shown in Fig. 7 with two oxygen and two nitrogen ligands to Coz+.One complex has two propionic acid residues, while the other has two acetate residues as oxygen ligands to Coz+. The metal-ligand bond angles are given in the figure to emphasize the differences in these complexes. Figure 8 shows the MCD spectra for these two complexes. The spectra of the model complexes resemble well the Co'+-thermolysin spectra. Substitution of 2-methylimidazole for imidazole produces a spectrum similar to comthat for the Coz+-thermolysin-/?-phenylpropionyl-6-phenylalanine plex, while the spectrum with imidazole resembles closely that of the Co2+thermolysin spectrum without inhibitor. Additional differences are noted in the EPR spectra of these two complexes (Fig. 9). The g anisotropy in the EPR spectra is different for the two complexes in spite of very small differences in their tetrahedral bond angles (Fig. 7). The EPR spectrum for the protein (15) does not have the excellent resolution of that of the model complexes and further studies may be necessary to demonstrate quantitative similarities between the model complexes and the enzyme-Co2 complex. +

3. Other Studies of the Metal Ion Environments

A novel approach for determining metal-metal distances in proteins in solution was applied to thermolysin by Horrocks et al. (16). As mentioned earlier, thermolysin has four Ca2+ binding sites. Three of the four can be replaced by trivalent lanthanide ions. Tb3+ fluorescence is enhanced when Tb3+ is bound to the so-called # 1 Ca2+ site of thermolysin. When Coz+ is added to the Znz+active site, the fluorescenceof bound Tb3+is quenched by I

1

I

I

a

-I- Methylimidorole - 0.1-

2- Methylimidozole

--

I

1

500

600

r.7.7=

lmidozole I

500 WAVELENGTH,nm

I

-

600

FIG.8. Magnetic circular dichroism spectra of Co2+ complexes. The oxygen and nitrogen ligands are given for each complex.


335

FIG.9. Electron paramagnetic resonance spectra of Coz+ complexes in Fig. 7.

89%. Using the Forster energy transfer equations ( I d ) , one can calculate that Coz+ and Tb3+ are 13.7 A apart, in excellent agreement with the value of 13.7 A from the protein x-ray structure. This technique should be valuable for future applications to other multimetal ion binding proteins and enzymes. Another technique that uses the fluorescence properties of trivalent lanthanides is that of the detection of fluorescenceemission decay induced by pulsed dye laser excitation. Horrocks and Sudnick (17) have applied this technique to the study of water molecules bound to metal ions in small complexes and proteins. In one study they found that the exponential decay of Tb3+ fluorescence is altered when H 2 0 is replaced by D20 and that this change can be used to determine the number of coordinated water molecules on the metal ion. With thermolysin, bound Tb3+ had 1-2 water molecules in the first coordination shell. This number is consistent with the x-ray structure. NMR studies of protons of water also can be used to determine the number

336


of water molecules in enzyme-metal ion complexes, and this approach has been used with thermolysin. Bigbee and Dahlquist (18) found that there was one exchangable water molecule in a thermolysin-Mn2+ complex and that this water molecule appeared to be displaced (at least partially) by inhibitors. These results also are consistent with the x-ray data These authors also studied inhibitor binding by I9F-NMR (19). They concluded that the inhibitor N-trifluoroacetyl-D-phenylalaninebinds to the enzyme at two sites, one of which is perhaps in the coordination sphere near the metal ion.

F. CONCLUSIONS Overall, many biophysical approaches have been applied to the study of catalysis by the metalloenzyme thermolysin. The metal ion was found by all studies to be involved in binding substrates and inhibitors. In addition, the x-ray structure of the protein has been useful in reaching the conclusion that the metal ion probably functions in catalysis as an electrophilic “sink.” 111.

Yeast Hexokinase

A. BACKGROUND

Yeast hexokinase catalyzes the following reaction : hexose

+ MgATP % hexose-6P + MgADP + H’

This enzyme plays a key role in the metabolism of glucose and other related sugars. The physical and kinetic properties of yeast hexokinase have been extensively studied. Numerous recent studies have been made of its role in the phosphoryl transfer reaction. Hexokinase has a molecular weight of 102,000 and is composed of two identical subunits of 51,000 molecular weight each (20-22). In yeast, hexokinase exists as a mixture of two isoenzymes (23). These forms have been named A and B (23).These isoenzymes can be separated by chromatography and have been found to be chemically different (23, 24). Some of the earlier work with hexokinase was done with preparations that contained a mixture of isoenzymes. Work also was done with enzyme that was proteolytically modified. These variations undoubtedly have been responsible for some of the controversy concerning the properties of this enzyme. Much of the earlier work on hexokinase has been summarized in two reviews that appeared in 1973 (25,26).The emphasis here will be on information that has appeared since that time. Significant advances have been made with respect to kinetic mechanism, stereospecificity of phosphoryl transfer, and x-ray structure.

337


B. SUBSTRATE SPECIFICITY 1. Sugar Site

Unlike many enzymes, hexokinase has a broad substrate specificity. A partial list of substrates for the sugar site is presented in Table 11. Structures for some of these compounds are given in Fig. 10. It should be pointed out that kinetic constants for the various substrates were obtained under a variety of conditions with different forms and modifications of the enzyme. From the list of substrates in Table 11 (27-33) it is apparent that the orientations of the hydroxyls on carbons 3 and 4 of glucose(I) are important for binding and activity, since neither D-allose(1I) nor mgalactose(II1) are very good substrates. The position of the hydroxyl at carbon 2, however, is relatively unimportant, since D-mannose(1V)and other compounds modified at this carbon are very good substrates. 1,5-Anhydro-~-glucitol(V) and 1,5anhydro-D-mannitol(V1) are substrates that lack the anomeric hydroxyl and thus are locked into a pyranose configuration. The. low activity seen with these compounds suggests that the anomeric hydroxyl is important for binding to the enzyme. However, the anomeric hydroxyl can be in either TABLE I1 Kinetic Constantsfor the Sugar Substrates of"Yeast Hexokinase Relative Substrate D-Glucose D-Galactose D-Allose D-Mannose 2-Deoxy-~-glucose D-Glucosamine D-Glucosone 5-Thio-~-glucose 2-Deoxy-2-fluoro-~-glucose 2-Deoxy-2-fluoro-D-mannose 3-Deoxy-3-fluoro-D-glucose 4-Deoxy-4-fluoro-~-glucose 1,5-Anhydro-~-glucitol 1,5-Anhydro-~-mannitol D-Fructose 2,5-Anhydro-~-mannitol 2,5-Anhydro-~-glucitol 2,5-Anhydro-~-mannose 1-Deoxy-D-fructose 5-Thio-D-fructose a

Relative to fructose.

Km,mM 0.01

> 50

> 100 0.05 0.30 1.50 0.02 4.0 0.19 0.41 70 84 3 7 7 6.3 47 0.31 614 1.7

Vmax

100 20.2

> 10 80 100 70 20 1

50 85 10 10 1 6 180 155 109 109 2" 2.6

Reference 27 27 27 27 27 27 27 28 29 29 29 29 27 30 27 31 31 31 32 33

338


"b I

Vlll

VII

IX

I1

X

XI

FIG.10. Structures of sugars and sugar analogs that have been tested as substrates for yeast hexokinase.

the a or p configuration since both a-D-glucoseand B-D-glucose are substrates (34-36). The rate with a-D-glucose is 1.2-1.5 times faster than with p-Dglucose (34-36). Since neither 3-deoxy-3-fluoro-~-g~ucose(VII) nor 4-deoxy4-fluoro-~-glucose(VIII) are good substrates, this suggests that the hydroxyls at carbons 3 and 4 are acting as proton donors rather than proton acceptors in the formation of hydrogen bonds to enzyme (10). Hexokinase will also phosphorylate D-fructose(1X) and fructose analogs. This shows that the enzyme also will accept furanose rings at the active site. The same hydroxyls appear to be important for binding and activity since Dtagatose(X) and D-psicose(XI), the C-4 and C-3 epimers of fructose, are not substrates for hexokinase (37).

339


2. Nucleotide Site

A partial list of substrates that serve as phosphoryl donors are presented in Table 111 (38-42). Alterations have been made in the base, sugar, and phosphate portions of ATP. The amino group at position 6 appears to be relatively unimportant since a large number of bulky substituents can be attached here with little effect on binding or activity (38-40). It can also be removed since purine riboside-5’-triphosphateis a good alternate substrate (38). Replacement of the amino group with a hydroxyl group (ITP) greatly increases the K,, but only reduces the V,,, about 50% (40). The high activity of 7-deazaadenosine-5’-triphosphateindicates that the nitrogen at position 7 is not in strong interaction with the enzyme (40).Both of the ribose hydroxyls appear important for good activity since removal of either increases the K, about sixfold and decreases the V,, by a factor of 16 (40). C . STEREOCHEMISTRY OF THE ACTIVE COMPLEX OF MgATP

Cornelius and Cleland have recently determined the absolute stereochemistry of the Mg2+complex of ATP that is active with yeast hexokinase (43). MgATP can exist in solution as a pair of bidentate diastereomers (monodentate and tridentate complexes can also exist). The two possible bidentate isomers are shown schematically in Fig. 11. The two configurations have been labeled A and A (44). There are also four possible tridentate isomers. TABLE 111 Kinetic Constantsfor the Nucleotide Substrates of Yeast Hexokinase Relative Substrate ~~~

Kln

Vm,x

Reference

0.100 0.046 0.13 0.34 0.60 0.34 7.46 0.032 0.57 0.59 0.18 1.2 ND“

100 56 83 2 53 2 11 40 6 6 94 28 ND

38 39 40

~

Adenosine-5’-triphosphate N6-Benzoyladenosine-5’-triphosphate N6-Monomethyladenosine-5’-triphosphate Guanosine-5’-triphosphate Purine riboside-5’-triphosphate 2-Aminopurine riboside-5’-triphosphate 2,dDiaminopurine riboside-5’-triphosphate 8-Bromoadenosine-5’-triphosphate 2’-Deoxyadenosine-5’-triphosphate 3’-deoxyadenosine-5’-trip hosphate 7-Deaza-adenosine-5’-triphosphate Adenosine-5’-(1-thiotriphosphate) Adenosine-5’-(2-thiotriphosphate)

ND, not determined.

40

38 40 40

38 40 40 40 41

42

340


0, ,0-AMP

A

A

FIG.1 1 . Structures of the two fl,y-bidentatediastereomers of MgATP.

Which of the two possible isomers is active with yeast hexokinase? Since Mgz+exchanges ligands rapidly, it is not possible to separate the two species (45). Cornelius and Cleland synthesized Co(NH,),ATP, which is bidentate and, since Co3+ exchanges ligands very slowly, the individual isomers should be stable to isomerization (46). Co(NH,),ATP was found to be a substrate for hexokinase in the following reaction (43): D-Glc

+ Co(NH,),ATP

-P

Co(NH3),(Glc-6-P)ADP

The glucose-6-P remains in the coordination sphere of the cobalt. Only one-half of the mixture of the two isomers was a substrate; the part remaining was inert, although it did inhibit. The inert material and the product were separated and a degradation product of the inert material was crystallized. The absolute stereochemistry was determined by x-ray crystallography (44). The active isomer of Co(NH,),ATP used by hexokinase was found to correspond to the A isomer of P,y-bidentate MgATP (43). Jaffee and Cohn have used the above information to assign the absolute configuration to the two isomers of adenosine-5’-(2-0-thiotriphosphate) (42). Since the p-P of ATP is prochiral, the introduction of a sulfur as a replacement for one of the nonbridge oxygens results in a pair of diastereomers. These diastereomers have been labeled A and B by Eckstein and Goody (47). Jaffee and Cohn found that only the B isomer of adenosine-5’-(2-O-thiotriphosphate) was a substrate for hexokinase and thus, by analogy with the results of Cornelius and Cleland, the B isomer has the structure shown in Fig. 12 (42). The Mg2+ is presumed to be binding to oxygen rather than sulfur because of the known preference of Mgz+ for oxygen over sulfur (48). When the Mg” was replaced by Cd2+,hexokinase was specific for the A isomer of adenosine-5’-(2-0-thiotriphosphate)(42). This occurs because cadmium greatly prefers to coordinate through the sulfur atom rather than the oxygen atom (48, 49). The active species still corresponds to the A isomer of p,y-bidentate MgATP, as shown in Fig. 13. Hexokinase has also been found to use predominantly the A isomer of adenosine-5’-(1-0-thiotriphosphate) (41).The absolute configuration of the two isomers recently has


0 1 I /P-O,

0,o

Mg-0

34 1

,0-AMP

,p\

s

FIG. 12. Structure of the Mg2+ complex of the B isomer of adenosine-5’-(2-O-thiotriphosphate).

0

FIG. 13. Structure of the Cd2+ complex of the A isomer of adenosine-5’-(2-O-thotriphosphate).

been determined. The A isomer has the S configuration and the B isomer has the R configuration (50, 51).

D. KINETICMECHANISM 1. Introduction

In the last 15 years a large number of publications have been concerned with the determination of the kinetic mechanism of yeast hexokinase. Unfortunately, there has been much disagreement between various authors on the conclusions to be drawn from such studies. In a two-substrate reaction similar to that catalyzed by hexokinase, two basic mechanisms may be at work. First, a “ping-pong” reaction may be occurring in which the enzyme shuttles between a stable enzyme intermediate, such as a phosphorylated enzyme, and a free enzyme. Second, the reaction may be sequential, in which case no reaction occurs until both substrates are on the enzyme. There are two types of sequential mechanisms. If one substrate cannot bind until after the addition of the other substrate the mechanism is said to be ordered. However, if they can combine in any order the mechanism is said to be random. The various kinetic methods for distinguishing between these mechanistic forms have been summarized by Cleland (52). The evidence for and against these possible kinetic schemes will now be summarized for yeast hexokinase.

2. Initial Velocity Studies A number of initial velocity studies have been made with yeast hexokinase. All results to date indicate an intersecting initial velocity pattern in

342


both the forward and reverse reactions (53-55). This indicates that hexokinase has a sequential kinetic mechanism. Thus, the possibility that ATP phosphorylates the enzyme, which in turn phosphorylates glucose, is excluded as being kinetically important. 3. Product Inhibition Studies

Since the mechanism is sequential, product inhibition studies have been used to try to determine whether the mechanism is ordered or random. These results have been contradictory and have varied from laboratory to laboratory. Rudolf and F r o m (56) found that glucose-6-P was a noncompetitive inhibitor versus either glucose or MgATP. MgADP was also noncompetitive versus MgATP or glucose. However, Noat et al. found that glucose-6-P was competitive versus glucose and noncompetitive versus MgADP, while MgADP was noncompetitive versus both glucose and MgATP (55). Kosow and Rose have shown that MgADP is noncompetitivewith respect to MgATP at low MgADP concentrations, but competitive at high concentrations (57). Noat et al. argue that their data support an ordered reaction mechanism in which glucose adds before MgATP, followed by the ordered release of MgADP and Glc-6-P (55).Fromm and co-workers argue that the mechanism is actually random, but they must postulate a number of abortive complexes with certain sets of rate constants to make their data compatible with the usual results seen for a random mechanism (54, 56). On the basis of their results, Kosow and Rose have argued that the release of products is random, but that there is a preferred pathway in which MgADP leaves before Glc-6-P. They also state that the release of Glc-6-P from the enzyme-(Glc-6-P) complex is at least partially rate limiting for the reaction (57).

4. Isotope Exchange at Equilibrium Fromm et al. measured the ATPct ADP and Glc-Glc-6-P isotope exchange at equilibrium (58). They found that both exchange rates increased and then leveled off to a plateau as the level of ATPIADP or GlclGlc-6-P was raised (58). No inhibition was seen of either exchange (58). These results are in agreement with a random mechanism and not with an ordered one. However, both exchange rates were not the same. The MgATP t)MgADP exchange was about 50% faster than the Glc c)Glc-6-P exchange (58). This indicates that the exchange is partially limited by the release of glucose and/or glucose-6-P from the enzyme. If the mechanism was strictly ordered as postulated by Noat et al., the Glc ct Glc-6-P exchange would have been inhibited at high concentration of MgATPlMgADP (52).


343

5. ATPase Reaction

In the absence of added glucose, hexokinase was found to catalyze the very slow hydrolysis of MgATP (59). This has been explained by assuming that water has replaced glucose at the active site of the enzyme. This ATPase activity can be inhibited by compounds that inhibit the hexokinase activity (60) and can be stimulated by compounds such as D-xylose or D-lyxose which lack the terminal -CH20H of glucose (61).The ATPase reaction has been used to support evidence that hexokinase has a random kinetic mechanism, since it shows that ATP can bind to hexokinase in the absence of glucose (62). 6. Binding Studies Binding studies have shown that hexokinase will bind one glucose molecule per subunit (63). Direct binding studies with ATP could not be determined since its binding constant is apparently too high (64). 7. Isotope Partitioning Experiments Rose et al. (65) have recently introduced a method for determining the rates of release of substrates from the various enzyme forms relative to the overall turnover of the enzyme. This is done by incubating hexokinase with enough radioactive glucose so that all of the enzyme is complexed with substrates. This solution is then added to a solution containing a large excess of unlabeled glucose and various amounts of MgATP and then rapidly quenched with acid. The glucose-6-P that is formed is isolated and, from its specific activity, the amount of initially bound glucose that proceeds to form the product, as opposed to that dissociating into the pool of unlabeled glucose is determined. From the percentage of initially bound glucose that is trapped, the kinetic constant for MgATP, and using the appropriately derived equations, the rate of release (Koff)of glucose from the E-glucose and E-glucose-MgATP complexes can be determined (65). They found that glucose did not dissociate from the E-glucose-MgATP complex to any extent and Koff from E-glucose was 0.3 times the turnover rate (65). Similar experiments in the back reaction of hexokinase have shown that glucose-6-P dissociates from E-glucose-6-P about 160 times faster than the enzyme turnover in the reverse reaction, and very slowly from ~-glucose-6-PMgADP (66). In a preliminary report, Solheim and Fromm also have tried the isotope partitioning technique with MgATP in the forward reaction. They find that some of the initially bound radioactive ATP is converted to radioactive

344

JOSEPH .I. VILLAFRANCA AND FRANK M. RAUSHEL

product (67).This is excellent evidence that hexokinase can operate through a random mechanism. 8. Conclusions

When taken collectively, the overall evidence indicates that hexokinase can both add and release substrates and products in a random mechanism. However, the mechanism cannot be described as rapid equilibrium random. The evidence also indicates that the preferred pathway is the ordered addition of glucose followed by ATP, then the release of ADP followed by glucose-6-P. Danenberg and Cleland have recently attempted to assign relative rate constants to a general random mechanism for hexokinase as shown in Fig. 14 (30). In this model the unimolecular constants are relative to the turnover number and the bimolecular constants are chosen to yield equilibrium constants in units of millimolar. The model is primarily based on dead-end inhibition by CrATP, the Michaelis constant for ATP in the ATPase reaction, the isotope partitioning experiments of Rose et al. (65), and various binding and kinetic constants found in the literature. The final model was based on a computer simulation study attempting to discover what combination of rate constants would fit the isotope partition data and the observed kinetic and binding constants. This model predicts that 98% of the reaction flux will go through the pathway with glucose adding first when both glucose and MgATP are present at I&, levels. At higher MgATP levels both pathways would carry nearly equal flux (30).

E. X-RAYSTRUCTURE 1. Crystal Forms The x-ray structure determination of yeast hexokinase has been undertaken by the Steitz group at Yale University (68). Although the complete amino acid sequence of hexokinase has not been determined, much information about the structure and function of this enzyme has been determined from the x-ray picture presented thus far. Work has progressed on three different crystal forms of hexokinase. These have been labeled BI, BII, and BIII. Form BI is in the space group P21 221 with 4 molecules per unit cell (68).An electron density map has been calculated to 6 A resolution. The two subunits of hexokinase in this crystal form are related by a rotation of 180" plus a translation of 3.5 A along the sym-

FIG. 14. Model for the kinetic mechanism of yeast hexokinase.

346


BI $180"

BII

& 160'

FrG. IS. A schematic drawing showing the difference in the arrangement of subunits found in the asymmetric unit of BI and BII crystals. From Anderson et al. (70).

metry axis (Fig. 15) (69).Each subunit appears to be identical and is bilobal with a narrow bridge of density between the lobes (69).The dimension of the dimer is 150 A x 45 A x 55 A and each subunit is 80 A x 40 A x 50 A (69). Unfortunately, neither glucose nor MgADP will bind to these crystals without disintegration of the protein crystals. This form has not been characterized further. The BII crystals, grown in potassium phosphate, are also dimeric and will bind sugars and nucleotides (70). BIII hexokinase has been proteolytically modified. It is monomeric in solution and in the crystalline state (71).

2. Quaternary and Tertiary Structure In the BII crystals the dimer has the overall dimensions of 90 A x 85 x 60 A. In addition, each subunit has the overall dimensions of 80 A x 50 A x 40 A (70). Like the BI crystals, the two subunits in the BII crystals are not related by a simple 180" rotation axis. The subunits, however, are related by a 160" rotation axis coupled with a translation by 13 A along the symmetry axis (Fig. 15) (70). The tertiary structure of the individual subunits in the BI and BII crystals appear to be identical. Glucose and O-toluoylglucosamine bind to these crystals in the deep cleft that separates the lobes of each subunit (70).ADP or AMP-PNP only bind in the presence of a sugar substrate or inhibitor. Only one nucleotide binds per dimer, and its binding site is located at the point of contact between the two subunits (72).Parts of the binding site are on each subunit. This site has been labeled the I site (72).A schematic drawing of the BII structure with the location of the various binding sites is shown in Fig. 16 (72). The third crystal form, BIII, contains one monomer per asymmetric unit and also binds glucose and AMP (52). Like the BII dimer, glucose binds in the deep cleft that separates the two lobes. AMP does not bind at the nucleotide site of the BII crystals since each monomer has only part of the complete

A


P \

347

/

W

FIG.16. A schematic drawing of the hexokinase dimer showing the location of observed glucose and nucleotide binding sites. The two sugar binding sites are represented by S and S . The intersubunit AMP-PNP site I is formed by regions Iu and I d of each subunit. The symmetry related sites are indicated by Iu‘ and Id. The location of the AMP binding sites in the monomeric BIII crystal form are labeled A and A . From Anderson et al. (72).

site. AMP binds at another site close to the glucose site (71). This form has been determined to a 2.7 A resolution. The tertiary structure of the BII is almost identical with the BII enzyme. A schematic drawing of the polypeptide backbone with the location of the various binding sites is shown in Fig. 17 (71). 3. Intersubunit Site

As mentioned, AMP-PNP or ADP in the presence of glucose will bind only to the BII crystals at a site between the two subunits. Nucleotides bound at this site appear to be in a fully extended conformation (73). ATP analogs bound at this site make contact with amino acid residues from both subunits. The y-phosphate of ATP bound at this site is 20 A from the 6hydroxyl of bound glucose on one subunit and 30 A from the glucose on the other subunit (73). It has been proposed that this site is an allosteric regulatory site for hexokinase and not the substrate site for ATP where phosphoryl transfer occurs (73). 4. ATP Site The intersubunit site does not exist in the monomeric BIII crystals. AMP binds to these crystals at another site labeled the “A site.” Model building,

348

JOSEPH

J.

VILLAFRANCA AND FRANK

M. RAUSHEL

V FIG.17. A schematic drawing of the course of the polypeptide backbone of yeast hexokinase. From Anderson et al. (74). Reprinted with permission of Academic Press.

based on AMP and sulfate difference maps, with ATP at this site show it to be in an extended conformation (73). When placed in this manner, the yphosphate of the ATP is only 6 A from the 6-hydroxyl of the bound glucose molecule. It has therefore been proposed that this be the active site for ATP (73). 5 . Glucose Site

Glucose binds in both the BII and BIII crystal forms in the deep cleft separating the two lobes of each subunit. Glucose appears to bind in the C-1 (chair equatorial) a-D-glucopyranose conformation (73). All of the hydroxyls, except the 1-hydroxyl, are thus in an equatorial position. Protein side chains are hydrogen bonded to the 3-, 4-, and 6-hydroxylsand possibly to the 1-hydroxyl of glucose. The 2-hydroxyl appears to be pointing toward open space (73). This corresponds nicely with the observed substrate specificity. Although the amino acid sequence is not known, the specific amino acid residues at the active site of hexokinase have been attempted to be determined through a crystallographic refinement of the data to 2.1 A (74, 75).Steitz and co-workers have concluded that Asp-189 is hydrogen bonded both to the 6and 4-hydroxyl groups of glucose. The 4-hydroxyl is also hydrogen bonded to Asx-188 and Asx-215. The 3-hydroxyl interacts with Asx-245 and Asx-188 (76). The appearance of an aspartic acid residue at the active site hydrogen bonded to the 6-hydroxyl is consistent with recent kinetic evidence. F. CHEMICAL MECHANISM From pH kinetic studies, Viola and Cleland have proposed that hexokinase requires a group on the enzyme that must be unprotonated for the forward


349

FIG.18. A schematic drawing of the chemical mechanism of yeast hexokinase.

reaction and protonated for the reverse reaction (77). This group has been identified as a carboxylic acid from its enthalpy of ionization and behavior in organic solvents (77). In agreement with this are results from chemical modification studies which show that all of the histidine groups can be modified without complete loss of activity (78).Steitz et al. have also shown that there appears to be an aspartic acid residue hydrogen bonded to the 6-hydroxyl of glucose in the x-ray picture of hexokinase (76). The pH kinetic studies and the x-ray data suggest that during catalysis this carboxyl is acting as a general base in accepting a proton from the glucose in the forward reaction, and acting as a general acid in the back reaction by donating a proton to glucose6-P. This would tend to increase the nucleophilicity of the 6-hydroxyl and facilitate its attack on the y-phosphoryl of ATP (76). This is shown schematically in Fig. 18. G. CONCLUSIONS The role of the metal ion in hexokinase seems to be to coordinate the p,y phosphate groups of ATP. In this manner the metal ion serves as an electrophilic center to neutralize the negative charges on the ionized phosphate groups of ATP. This then leads to facilitated catalytic attack by the 6hydroxyl group of substrate on the y-phosphoryl of ATP.

IV.

Glutamine Synthetase

A. BACKGROUND The reaction that glutamine synthetase catalyzes is the following: g glut am ate

2M” + ATP + NH,+L-glutamine

+ ADP + Pi

Divalent cations are required for activity (Mg”, Mn2+,Co2+)regardless of the origin of the enzyme. The role of these metal ions in catalysis, as well as regulation of catalysis, has been the subject of numerous reviews (79-84). Homogeneous preparations of glutamine synthetase are available from many sources including Escherichia coli, Salmonella typhimurium, peas, sheep brain, and rat liver. Glutamine synthetases from bacteria have 12

350

JOSEPH J. VILLAFRANCA AND FRANK M. RAUSHEI

subunits, while the enzymes from mammals have eight. The subunit sizes of all glutamine synthetases are 44,000-50,000. Glutamine synthetase has a central role in nitrogen metabolism and presently there is substantial interest in all aspects of catalysis and regulation of this key enzyme. The synthesis of glutamine is probably the most central reaction in nitrogen metabolism. The amide nitrogen of glutamine is involved in production of amino acids for protein biosynthesis and DNA biosynthesis, as well as other important metabolic compounds (including several enzyme cofactors). The regulation of glutamine synthetase activity in E. cofi has been reviewed (79, 82-83). Glutamine is an important intermediate in the assimilation of ammonia by E. cofi and, consequently, a rigorous cellular control has evolved. The synthesis of the enzyme is decreased when E. coli are grown on media high in ammonia, whereas growth conditions on limiting ammonia produce a 20-fold increase of the enzyme. Additionally, feedback inhibition by multiple end products of glutamine metabolism (L-alanine, glycine, L-histidine, L-tryptophan, cytidine triphosphate, adenosine monophosphate, glucosamine-6-P)also regulate the activity of existing glutamine synthetase molecules. The enzyme presumably contains separate sites for each of the above-mentioned inhibitors. The main emphasis of this section is to summarize the kinetics, NMR, and EPR work with the enzyme and to speculate on the role of the metal ions in catalysis. Escherichia coli have also developed an elegant method to control enzyme catalysis that occurs by covalent modification of each subunit. In this latter reaction a single tyrosyl residue per subunit is adenylylated to produce a stable 5’-adenylyl-U-tyrosylderivative. Recent NMR and fluorescence data will be reviewed concerning the nature of this adenylyl site and its spatial relationship to the metal ions at the catalytic site. The enzymes responsible for the covalent adenylylation reaction comprise a “cascade system” for amplifying the activation or inactivation of glutamine synthetase molecules (81)-

Experimental techniques for determining distances must be employed to establish the structure of the active site components of glutamine synthetase. Techniques that are available for these studies are x-ray crystallography, EPR, NMR, and fluorescence energy transfer. All approaches are currently being employed to study the structure and function of this metalloenzyme.

B. KINETICMECHANISM Since the reaction catalyzed by glutamine synthetase has three substrates (L-glutamate, NH,, and metal-ATP) and three products (L-glutamine, Pi, and metal-ADP), the kinetic mechanism as deduced by steady-state kinetics

351

METALLOENZYME CATALYSIS ATP

Glu

NH,

Gln

Pi

ADP

1 1 1 T T T I

Scheme Glu I

I

I

II

I

UI

I

Gln

+ I

+

4

Pi

I I

ADP

I I

I

Scheme II FIG.19. Two proposed kinetic schemes for the glutamine synthetase reaction.

presents a challenging problem. To simplify the situation we will consider only two alternative mechanisms (Fig. 19) and review the supporting or refuting data. One mechanism involves the sequential ordered addition of ATP (this is always metal-ATP), L-glutamate, and ammonia prior to release of any products. Another mechanism has a random (but sequential) order of addition of all substrates, and perhaps a random order of product release. The dashed lines indicate L-Glu binding at any of the three points in the sequence. For a completely random addition of substrates, the positions of binding of ATP and NH, could also be reversed. 1. Isotope Exchange at Equilibrium

The kinetic approach that has been most extensively applied to glutamine synthetase is that of equilibrium isotope exchange (85-88). Experimental procedures involve preparing a series of reaction solutions at chemical equilibrium and monitoring the following exchanges:

* ['"C]glutamine [15N]NH3* ['5N]glutamine [32P]Pi* ["PIATP ['4C]ADP * [14C]ATP

['"C]gIutamate

Scheme I

Results of such experiments with enzymes from the sheep brain, pea seed, and E. coli (87) show that, without exception, the relative rates of exchange are (glutamate s glutamine) > (NH, ~t glutamine) > (Pi P ATP) N (ADP e ATP). If the substrate interconversion were the rate-limiting step

352


in catalysis all exchanges would be equal at equilibrium. Since this is not the case and since ADP P ATP is the slowest exchange rate, the overall rate of the enzymic reaction is limited (partially at least) by nucleotide release. However, a distinction in mechanisms exists for the sheep brain enzyme as contrasted with the enzyme from the pea seed and E. coli. The latter two enzymes exhibit kinetic isotope exchange patterns consistent with random substrate addition and release (but still catalysis is limited by nucleotide release), whereas the enzyme from sheep brain has a preferred order of addition of substrates consistent with the preceding scheme I. 2 . Isotope Exchange During Net Enzymatic Reaction Recent experiments with ['80]Pi and E. coli enzyme by Stokes and Boyer (89) and Balakrishnan et al. (YO) show that transfer of oxygens from Pi to glutamine is the most rapid of the measured isotopic exchanges. The reaction for net formation of ATP follows. ['*O]P,

+ ADP + L-Gin -,L-GIU + NH, + ATP

The partial reaction of the l80exchange from [180]Pito L-glutamine was 5-7 times faster than net ATP formation. This establishes (along with the equilibrium isotope exchange data) the following order of events with the E. coli enzyme. ATP release

iNH,

release < L-glu release < substrate interconversion

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