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INORGANIC REACTION MECHANISMS

Edited by JOHN 0. EDWARDS DEPARTMENT OF CHEMISTRY BROWN UNIVERSITY PROVIDENCE, RHODE ISLAND

INTERSCIENCE PUBLISHERS a division of JOHN WILEY & SONS New York * London * Sydney * Toronto


Progress In Inorganic Chemistry

Volume 13

Progress in Inorganic Chemistry

Editor: STEPHEZ J. LIPP-iRD DEPARTMEST OF CHEMISTRY, COLUMBIA UNIVERSITY, NEW YORK, NEW YORK

Advisory Board

THEODORE L. BROWS KXIVERSITY OF ILLISOIS, URBANA, ILLINOIS

J-UIES P. COLLMAN STASFORD TNIVERSITY, STASFORD. CALIFORNIA

F. ALBERT COTTOS M.I.T., CAMBRIDGE, MASSACHUSETTS

RILEY SCHAEFFER ISDIAXA UNIVERSITY, BLOO?rIISGTOX, INDIANA

GEOFFREY WILKISSOX IMPERIAL COLLEGE, LOA-DOS, ESGLAND


Edited by JOHN 0. EDWARDS DEPARTMENT OF CHEMISTRY BROWN UNIVERSITY PROVIDENCE, RHODE ISLAND

INTERSCIENCE PUBLISHERS a division of JOHN WILEY & SONS New York * London * Sydney * Toronto

The paper used in this book has a pH of 6.5 or higher. It has been used because the best information now available indicates that this will contribute to its longevity.

Copyright 0 1970 by John Wiley & Sons, Inc.

All Rights Reserved. No part of this book may be reproduced by any means, nor transmitted, nor translated into a machine language without the written permission of the publisher.

Library of Congress Catalog Card Number 72-105386

SBN 471 23315 3

Printed in the United States of America

10 9 8 7 6 5 4 3 2 1

Preface

The need for surveys in all areas of chemistry remains clear. This volume is the first product of a desire for a series in the area of inorganic reaction mechanisms. It was decided to collect reviews by active inorganic chemists working in the area of mechanisms and to have a group of these reviews appear together as one volume in the series “Progress in Inorganic Chemistry.”

Two points deserve mention. First, I believe that understanding of the inorganic compounds and reactions is important to the understanding of the mechanisms. Thus considerable latitude of subject occurs in this volume. Second, attempts were made to include as authors chemists in other countries whose work on mechanisms may be less familiar to the American chemists than is the work carried out in the United States.

I would appreciate comments and criticisms. Thanks are due to the many persons who helped make this volume

possible.

Providence, Rhode Island October 23, 1969

John 0. Edwards

V

Contents

The Formation, Structure, and Reactions of Binudear Complexes of Cobalt

By A. G. SYKES, School of Chemistry, The University, Leeds, England, AND J. A. WEIL, Chemistry Division, Argonne National Laboratory, Argonne, Illinois . . 1

Fast Metal Complex Reactions By KENNETH KUSTIN, Department of Chemistry, Brandeis University, Waltham, Massachusetts, AND JAMES SWINEHART, Department of Chemistry, University of California, Davis, California . . 107

Recent Developments in the Redox Chemistry of Peroxides By S. B. BROWN, PETER JONES, AND A. SUGGETT, Department of Physical Chemistry, The University, Newcastle upon Tyne, England . . 159

Replacement as a Prerequisite to Redox Processes By ELEANOR CHAFFEE AND JOHN 0. EDWARDS, Department of Chemistry, Brown Uniuersity, Providence, Rhode Island . , 205

Nonbridging Ligands in Electron-Transfer Reactions By JOSEPH E. EARLEY, Department of Chemistry, Georgetown University, Washington, D.C. . . 243

The Intimate Mechanism of Replacement in d8 Square-Planar Complexes

By L. CATTALINI, Institute of General and Inorganic Chemistry, University of Padua, Italy . . 263

Subject Index 34 9

Cumulative Index

vii


Progress In Inorganic Chemistry

Volume 13

The Formation, Structure. and Reactions

of Binuclear Complexes of Cobalt

By A . G . SYKES

School of Chemistry. The University. Leeds. England

and J . A . WmL

Chemistry Division. Argonne National Laboratory. Argonne. Illinois

I . Introduction . . . . . . . . . . . . . . . . . . I1 . The p-Peroxo Dicobalt Complexes

A . Single-Bridged Species . . . . . . . . . . . . . . . . . . . . . . . . .

1 . Formation and Physical Properties . . . . . . . . . 2 . Stability of the p-Peroxo Complexes . . . . . . . . .

B . Double-Bridged Species . . . . . . . . . . . . . 1 . Formation and Physical Properties of the p-Amido-p-Peroxo

Dicobalt Complexes . . . . . . . . . . . . . 2 . Stability . . . . . . . . . . . . . . . . .

. . . . . . . . . . 111 . The p-Superoxo Dicobalt Complexes A . Formation and Stability B . Nature of the Superoxo Bridge . . . . . . . . . . . C . Photochemical Decomposition . . . . . . . . . . .

A . Reactions of p-Peroxo Complexes B . Reactions of p-Superoxo Complexes . . . . . . . . . .

V . The Interconversion of p-Amido Dicobalt Complexes . . . . . . A . Preparation and Properties

1 . The Ammine Series . . . . . . . . . . . . . . 2 . The Polyamine Series . . . . . . . . . . . . .

B . Kinetic Studies of Redox and Substitution Reactions . . . . . VI . The p-Hydroxo Dicobalt Complexes . . . . . . . . . . .

A . Preparation and Properties B . Stability and Redox Reactions

VII . Formation and Properties of Other Dicobalt Complexes VIII . Summary and Further Discussion

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

IV . Redox Reactions of p-Peroxo and p-Superoxo Dicobalt Systems . . . . . . . . . . . .

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

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

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

References . . . . . . . . . . . . . . . . .

2 4 4 4

16 23

23 27 28 28 32 39 41 42 46 56 56 57 66 69 76 76 81 85 91 94

1

2 A. G . SYKES AND J. A. WEIL

GLOSSARY OF SYMBOLS

.I1 AH* IS* h

Concentration, in moles liter ~

Enthalpy of activation = E(Arr1ienius) - RT, in kcal mole-' Entropy of activation, in eu = cal mole- ' ( 'K). ' Wavelength, in nm ( = l o - ' meter) Molar absorbance coefficient, in liter mole- ' c m - ' (for binuclear cobalt complexes, this will refer to moles of complex rather than moles of cobalt atoms) Ionic strength, in moles liter-'. Symbol p also denotes bridging groups in the complexes.

f

Temperatures cited are in 'C.

Redox potentials cited are used \\ ith sign convention recommended by IUPAC, that is, the sign of the potential is the same as the charge on the electrode. Configurations of metal complex stereoisomers are designated herein by D and L. Ne\s ly proposed nomenclature denotes these by -1 and A , respectively [see Information Bulletin No. 33, IUPAC, Dec. 1968, and hrorg. Chew., 9, 1 (1970)l.

I. INTRODUCTION

During the last ten years, binuclear complexes of cobalt have received a good deal of attention, and significant advances have been made in understanding the chemistry of these substances. Although the main concern of the present review is with the mechanisms of reactions leading to the formation, interconversion, and breakdown of dicobalt complexes, a reappraisal of the work concerned with the structures of the complexes is obviously relevant and will form a substantial part of what follows. It is fitting once again to give credit to the work of Werner and his students, which continues to dominate the background of the field despite the corrections and elaborations which have now been made.

The first dicobalt complex \{as reported by Fremy in 1852 (143,144). The chemistry of the binuclear cobalt complexes was further explored by Gibbs (159), Vortmann (420-423), Jorgensen (218-220), Maquenne (253), and Mascetti (258). A number of papers by Werner (435-444) systematized and greatly broadened the scope of the field. One of these (338) , published in 1910, remains a classic of its kind, summarizing the work of some tuenty-four doctoral students and assistants. I n it are reported the preparation and properties of a wide range of dicobalt complexes, with as many as three bridging groups, many of which have not been studied in any detail since this time. Ligands which are known to bridge two cobalt(ll1) atoms include NH,- , OH-, 02-, O Z 2 - , 0 2 H - , SO4"-, Se0,2-, N O 2 - , N,OZ2 . HP04 ' - , and CH,CO,-. With the excep- tion of the hydroxo-bridged species, complexes having more than one

BINUCLEAR COMPLEXES OF COBALT 3

bridge of the same kind are rare. The most common nonbridging ligands are ammonia, polyamines, amino acids, and dipeptides. Other ligands which may be present in dicobalt complexes include H,O, NO,-, CN-, F- , CI-, Br-, and I- . With chloride and bromide, there is now some evidence (23,370) that in certain circumstances the halide ion can form a bridging group. In the complexes considered, six ligands are attached to each cobalt atom in an approximately octahedral configuration. The geom- etry of complexes having one, two, or three bridging groups can best be visualized as two octahedra sharing a corner, edge, or face, respectively.

Much but by no means all of the recent work has centered around the complexes having an 0, bridge. The single-bridged peroxo complexes are important as intermediates in the preparation of other dicobalt complexes and, on the biological side, as models of oxygen carriers. Important electronic and isomeric variations which can occur within the series have proved to be an absorbing and now much better understood field of research. Thus it is now clear that there are two main types of dicobalt complexes having an 0, bridge. In one, the bridge may be thought of as an 02,- peroxide linkage, and in the other, as an 0,- superoxide group. We use the peroxo and superoxo terminology throughout this review to distinguish between the two, and delineate the basis for this terminology in subsequent sections. This terminology has only become generally accepted during the last few years, and a certain amount of discrimination is called for in reading earlier papers. The peroxo cobalt(II1,IV) description used by Werner (438) (implying nonequivalent cobalt atoms in the superoxo complexes) is no longer acceptable, nor does the statement that the two atoms are equivalent because of equal contributions of the resonance forms cobalt(II1,IV) and cobalt(IV,III) have useful chemical significance. Other names which have been used (90,260,262,284) are “diamagnetic peroxo ” and “paramagnetic peroxo ” complexes, respectively. These names were for a time desirable because they left open the question of assigning formal oxidation states to the two cobalt atoms. There now seems to be little doubt that both cobalt atoms are in oxidation state (111) in both types of complexes.

A very complete but not always discriminating summary of the vast array of dicobalt ammine complexes reported in the literature up to ca. 1960 may be found in Gmelin’s Handbook (165). Many of these compounds are also discussed in the compendium Nouveau Trait6 de Chimie Minkrale (304). Previous reviews covering oxygen-bridging species (but not including decisive new developments, particularly with regard to the peroxo and superoxo systems) include “Molecular Oxygen as a Ligand in Metal Porphyrins and other Metal-Complex Compounds ” by Michaelis

4 A. G. SYKES AND J. A. WEIL

(274), chapter 8 of Martell and Calvin’s book Chemistry of the Metal Chelate Compounds (257), “The Nature of Peroxo-Bridged Dicobalt Complexes” by Goodman, Hecht, and Weil (1 67) , “Synthetic Reversible Oxygen-Carrying Chelates” by Vogt, Faigenbaum, and Wiberley (41 S), a section of “Peroxy Compounds of Transition Metals” by Connor and Ebsworth (90), “Reactions with Molecular Oxygen” by Fallab (132), and “ Reversible Oxygenierung von Metallkomplexen ” by Bayer and Schretz- mann (29). Basolo and Pearson (25) in Mechanisnis of Inorganic Reactions (Chapter 8) include a short discussion of cobalt-containing synthetic oxygen carriers. Dicobalt complexes with other bridging groups have been considered in Sidgwick’s text Chet?iical Elements and Their Compounds (360) and recently in the review ‘‘ Polynuclear Complexes of Cobalt(TI1) Ammines” by Chester (85).

Procedures for the preparation of some of the dicobalt complexes are to be found in Palmer’s Experituental Inorganic Chemistry (299), in Handbook of Preparatice Inorgarlic Chemistry by Brauer (49), and in Inorganic Syntheses, Volume XI1 (104).

11. THE p-PEROXO DICOBALT COMPLEXES

A. Single-Bridged Species

I . Fortnation and Phpical Properties

I t is now well established that many cobalt(I1) complexes will take up molecular oxygen very readily in aqueous solutions to give binuclear p- peroxo complexes. The best known example is with cobalt(II)/ammonia solutions. The reversibility of the reaction and rate of oxygen uptake have been investigated by numerous authors in the past (82,164,221,224,236,275, 425), and the possible use of these complexes for oxygen-fixation has been explored (1 64).

Recent kinetic studies by Simplicia and Wilkins (365) are consistent with the reaction scheme

hi

h - 1 CO(NH,)~(H,O)~ * + 0, CO(NH~)~(OZ)’+ + H,O (1)

h z

k - z CO(NH,)5(Oz)2t + C O ( N H J ) ~ ( H ~ ~ ) ~ ~ ( N H , ) , C O . O ~ . C O ( N H ~ ) ~ ~ + +HzO ( 2 )

The reaction with oxygen is rapid, and over short periods (- 10 min) completely reversible, as can be demonstrated by passing nitrogen through the


solution (275,330). The reversibility over many cycles is limited in that side reactions tend eventually to destroy the primary reaction products, with the formation of mononuclear cobalt(lI1) complexes. The binuclear p- peroxo complex can be oxidized further by a number of reagents to yield the corresponding p-superoxo complex, as is discussed in detail later (Sec. III-A). Because there is as yet no satisfactory way of recrystallizing salts of the p-peroxobis-[pentaamminecobalt(III)] ion, contamination by other complexes is best minimized in the synthesis. For example, in the preparation of (NH3),Co.02.C~(NH3)64+, it is useful to work with nitrate solutions since the nitrate salt is fairly insoluble, with the result that the complex is effectively removed from the system (104). Use of oxygen rather than air and care in maintaining the correct ammonia concentration during the reaction are also beneficial (228). Proofthat the bridging 0, group comes from the molecular oxygen rather than from the oxygen of the water molecules has been obtained from isotopic (oxygen-18) studies (284). No defin- itive evidence has yet been obtained for the existence of CO~+(NH~)~(O,) , for which one may also write the resonance form CO~+(NH~)~(O, - ) . Of the various aquo-ammine complexes of cobalt (II), the aquo-pentaammine ion appears most reactive toward oxygen addition (221,365). Fallab and co- workers have suggested (37,132) that at least three nitrogen ligands are needed for combination of the mononuclear complex with molecular oxygen. On the other hand, cobalt(I1) complexes with six nitrogen ligands appear to be much less efficient in absorbing oxygen (132,276). These various effects are discussed more fully below.

The pentacyanocobaltate(I1) ion is an extremely good reducing agent and in aqueous solutions reacts rapidly with oxygen to give a binuclear p-peroxo complex

2Co(CN)S3- + 0 2 + (CN),Co.O,-Co(CN),'- (3)

This complex was first described by Haim and Wilmarth (179), who found that it can be oxidized further to yield a superoxo complex. The 6- ion is brown in color and stable in fairly strong alkaline solution. The equilibrium with respect to cobalt(I1) and 0, strongly favors the p-peroxo species. The reactions analogous to (1) and (2) have been discussed in conjunction with the development of a sensitive analytical scheme for measurements of 0, concentrations (255). Some evidence has been obtained from EPR spectroscopy (31) for the possible existence of CO(CN),(O,)~- as a stable entity; further details regarding the nature of this ion and its role in reaction (3) would be of interest.

Information regarding the nature of peroxo-bridged complexes can be obtained from magnetic, x-ray crystallographic, and optical studies.

TA

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by

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119"

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BINUCLEAR COMPLEXES OF COBALT I

Static magnetic susceptibility measurements of the ion (NH,),Co. 0,. Co(NH,),*+ have shown that it is diamagnetic (39,134,211,264) and therefore that it does not contain cobalt(I1) ions or loosely bound (triplet state) oxygen molecules. From this it can be concluded that an electron has been transferred from each cobalt to the 0, group at the time of its formation. Similarly, the ion (CN),Co-O,. CO(CN),~- has been shown to be diamagnetic (21 1).

The best evidence for the peroxo nature of the 0, bridge comes from a recent structural determination by Schaefer (342) of the 4+ decaammine ion in the disulfate tetrahydrate salt. It demonstrated the presence of a classical a-bonded system in a four-atom nonlinear chain

0

co' \ co /

0

with structure similar to that of hydrogen peroxide. The 0-0 distance of 1.47 A is typical of the OZ2- group (see Tables I and 11; the first of these tables includes data for complexes referred to later in the text). The four atoms are not coplanar, being situated to form a torsional angle of 146" about the 0-0 direction. The six ligands around each cobalt atom form nearly regular octahedra. An earlier structure of the same ion (412) in the tetrathiocyanate salt appears to be unsatisfactory.

As a general rule, the optical spectra of the peroxo complexes are almost featureless in the 400-700 nm range. This is because intense (oxygen to cobalt) charge-transfer bands centered in the UV region overlap into the visible, giving typical brown-yellow colors and in many cases obscuring

TABLE I1 Bond Distances d for Some Diatomic Oxygen Species

Species 4 8, Refs.

0 2 + 1.12 84 0 2 1.20 84 0 2 - 1.26 0.02 84

"1.32-1.35 182 0 2 2 - 1.49 * 0.02 84

"An allowance has been made for librational motions of Oz- in the crystal Iattice.

8 A. G. SYKES A N D J. A. WEIL

the less intense d - t d bands which are expected for cobalt(II1). In ammoniacal solution, the brown decaammine complex has a characteristic absorption band at 360 nm (6 = 2900 i 300M-' cm-') (365). In more acidic media, there is evidence for protonation of the complex with an accompanying change in color from br0u.n to red (196,383) and development of a peak at 500 nm (< - 200) (196,283). Salts containing this protonated form of the complex were first prepared by Vortmann (422). For example, on slow addition of the [(NH,),CO.O,.CO(NH,),]~+ sulfate dihydrate to ice-cold 3'2 sulfuric acid, a red salt of the ion (NH3)iCo.(OnH).Co(NH3)a5', which appears to contain the counter ions ( H S 0 4 ) 3 S 0 4 , is obtained. The complex is rather unstable and difficult to recrystallize, and the structure has not as yet been determined. However, from what follows (pp. 24-26), i t is seen that such a hydroperoxo complex could have either structure (11) or (111):

Whether the proton in structure (11) should be shown centered between oxygen atoms (as indicated) or bonded exclusively to one of the oxygen atoms is not yet known. The protonated complex, like the unprotonated species, is readily oxidized to the corresponding green p-superoxo (5 +) ion (270,283). The situation with respect to protonation is very similar for the p-peroxo decacyano complex, as will be discussed (p. 22).

Many salts of the p-peroxobis[pentaarnminec~balt(III)]~ + ion have been prepared, and these have been summarized in Gmelin (165). Num- erous other monobridged p-peroxo dicobalt complexes have been reported and more or less well characterized. One such series contains various polyamines as nonbridging ligands. In a number of papers, Fallab (37,123,131) has discussed the solution chemistry, optical spectra, and re\ersible oxygen uptake of cobalt(I1) complexes of ethylenediamine (en), diethylenetriamine (dien), and triethylenetetraamine (trien). Formal names for dien and trien are 3-azapentylene-l,5-diamine and 3,6-diazaoctylene- I ,8-diamine, respectively. Jonassen and co-workers (213,214) have reported oxygen addition to cobalt(I1) complexes of tetraethylenepentaamine (tetren) and trien. It is possible (131) to follow the formation of the trien peroxo complexes by making p H measurements, since there is appreciable protonation of the free ligand in solutions up to p H - 9. The formation of the trien p-peroxo complex has been studied by a variety of methods


( I ,276a) including amperometric, conductimetric, and potentiometric titra- tions, spectrophotometry, and measurements of (oxygen) pressure changes ; static susceptibility measurements have shown that it is diamagnetic.

The presence of oxygenic groups as ligands attached to the cobalt(I1) reduces the readiness with which these complexes form stable adducts with molecular oxygen. Thus aminopolycarboxylic acids such as nitrilotriacetic acid, ethylenediamine-N,N-diacetic acid, and ethylenediaminetetraacetic acid (EDTA) form complexes which do not appear to combine with 0, (132). On the other hand, it has been shown (37,131) by combining polyamine and oxygenic ligands in various ways that three nitrogen ligands together with two oxygen ligands suffice to allow the cobalt(I1) complex to absorb oxygen. For example, the complex Co"(H,O)(dien)(C,O,) is capable of this. Presumably the difference between nitrogen and oxygen ligands is the result of the somewhat greater a-electron donating ability of the former, which in turn allows the cobalt(I1) to transfer an electron to the oxygen molecule. The effect can also be rationalized by making comparisons with other cobalt(II1) complexes. For example, hexaaquo- cobalt(II1) (potential ca. 1.9 volts) is a much better oxidizing agent than hexaamminecobalt(II1) (potential 0.1 volt), and in mixed aquo-ammine complexes there is an increase in oxidizing power of the cobalt(l1I) center as the number of water molecules is increased, This is consistent with the above observations regarding the stability and ease of formation of peroxo dicobalt complexes. A more quantitative understanding of this effect, and the relative effectiveness of different ligands in positions cis and trans to the 0, group, must await further studies.

In the oxidation of solutions of Co"-EDTA with hydrogen peroxide (pH 6.5-8.5), there is evidence (450) for the formation of an intermediate which has been assigned the formula (EDTA)CO.O,.CO(EDTA)~-. The preparation of such a complex in this way is of particular interest since the same complex does not appear to be formed when 0, is allowed to interact with CoII-EDTA solutions (1 32,364).

A number of p-peroxo complexes containing mixtures of ammonia, ethylenediamine, diethylenetriamine, and triethylenetetraamine, as well as the single pentadentate ligand tetraethylenepentaamine have recently been isolated and characterized, as have the corresponding p-superoxo salts (34,115). These p-peroxo complexes are listed in Table 111. Also included are p-peroxo complexes with trimethylenediamine and dipro- pylenetriamine as ligands (148).

Other cobalt(I1) amine systems which have been studied (276) include that with the tridentate ligand diethylenetriamine (dien). In this case, one of the amine-cobalt bonds in the complex CoT1(dien), must dissociate before it can form a stable adduct with oxygen. By studying the rate of

10 A. G. SYKES A N D J. A. WEIL

TABLE 111 Peroxo Dicobalt Complexes [(L)gCo.Op. Co(L),](ClO,),~2Hz0 with Various Multidentate Polyaniine Ligands L. Spectral Data for the Visible Region Are

Included, where Available

Complex" A, nmb 6, M - ' ~ L T - ' ~ Refs.

(en),(NH,)Co.0z.Co(NH,)(en)2ii (dien)(NH,),Co.02.Co(NH3)2(dien)i (dien)(en)Co 02. Co(en)(dien)* * (dpt)(tmd)Co. 02.Co(tnid)(dpt)4 (dien)(tnid)Co' 0,. Co(tnid)(dien)* +

(dpt)(cn)Co.02.Co(en)(dpt)*' (trien)(NH,)Co 0,. Co(NH3)(trien)i + (tnid)dNH,)Co. 0,. Co(NH,)(tnid)24 - (tetren)Co. O2 .Co( t e t r e r~ )~

d

d

550 537 420

-

d 562

437

~-

d

241 115 212 115 551 115

148 148 148

230 115 148

619 115

-

__ ~

-

a Abbre\,iations: en = ethylenedianiine: dien = diethylenetriamine; trien = triethvlenetetraaniine; tetren = tetraeth~lenepentaamine; dpt = dipropylenetetra- aniine; tnid = trimethylenedianiine = 1,3-dianiinopropane.

Wavelength at shoulder of absorption band. Molar absorption coefficient, measured in water at room temperature. The corresponding superoxo 5 + complexes have also been prepared. Reference 34 reports maxima at 305, 356, 452, and 602 nm.

oxygen uptake, it has been concluded that protonation of the free arm of the dien ligand is appreciable even at low, pH ( - 7.5). A very similar situation occurs with the hexadentate ligand pentaethylenehexaamine (NH,CH,CH,),NCH,CH,N(CH2CH2NHz)z. The cobalt(I1) complexes of histamine and 2-(aminomethyl)pyridine are also oxygen carriers (276). Peroxo complexes of the type (L)(cyclam)Co. 0, Co(cyclam)(L) have been prepared and isolated (48), where cyclam is the cyclic quadridentate ligand I ,4,8.11 -tetraazacyclotetradecane, and the ligand L which is trans to the oxygen may be either NCS- . N 3 - . NOz- , HzO, or CI- (see Table IV). The complexes are diamagnetic, and have optical spectra in which d + d bands characteristic of cobalt( 111) r i 6 complexes can be identified. These tend to be overlaid by intense charge-transfer bands (centered at ca. 350 nm), as in other peroxodicobalt complexes. More detailed studies of the relative stabilities and rates of reactions as a function of the trans ligand L v~ould be of considerable interest.

A brown oxygen adduct with the cobalt(I1) complex of 2,2'-bipyridine probably having formula [(H,O)(bip),Co~Oz~Co(bip),(H,O)](NO3),, has been prepared and studied in solution by polarographic (64) and spectro- photometric (327) means. The complexes of dimethylglyoxime and other


TABLE IV Spectral Details of Cyclarn p-Peroxo Dicobalt Complexes" in Aqueous Solution (48)

Complex A, nrn E, M - l cm-l A, nrn E, M - l crn-l ~ ~~~ ~ ~~

[(SCN)(cyclam)Co . 0,. Co(cy~lam)(SCN)]~ + 530 246 - 420 703

[(N02)(cyclarn)Co. 0,. Co(cyclam)(NOz)lZ + - 500 270 [(H,O)(cyclarn)Co. 0,. Co(cy~larn)(H~O)]~ + 530 170 - 430 470 [(Cl)(cyclarn)Co . 02. C~(cyclarn)(Cl)]~ +

[(N,)(cyclam)Co. 02. Co(cy~larn)(N,)]~ + - 550 390 -440 1000 - -

- b445 - b570

a Cyclam is the cyclic quadridentate ligand 1,4,8,1 I-tetraazacyclotetradecane. Diffuse reflectance spectrum.

related oximes with cobalt(I1) also are known to form adducts with oxygen and are very sensitive to irreversible oxidation to mononuclear cobalt(II1) complexes (71,171,353). Oxygen adducts appear also to have been observed in phthalocyanine pigments (356).

Oxygen adducts of various Schiffs base cobalt(I1) complexes have been known for many years (400). These were extensively studied in the period 1940-1950 by Calvin and co-workers (17,19,69,70,184,202,446), and by Diehl and his students (106,107), with some success toward develop- ing reversible oxygen-fixing compounds. Further work, including rate studies by oxygen isotope exchange, has been reported by Panchenkov and Tolmachev (300,301). The best known compound of this general type is the square-planar (125,292) complex N,N'-ethylenebis(salicy1idene- iminato)cobalt(II), abbreviated to Co(SaEn) and depicted in Figure 1. An x-ray structure determination of the 1: l chloroform adduct of this complex has recently been completed (344). The effect of substituents in ring positions 3 and 5 has been studied. In aprotic solvents such as dimethylformamide and dimethylsulfoxide, complexes of the type (L)(SaEn)Co .O,. Co(SaEn)(L) are formed and have been isolated (67). Here L represents a molecule of solvent. These complexes are diamagnetic and have the structure shown in Figure 2. The x-ray structural determination (68) of the adduct with dimethylformamide has shown the Co-0- 0-Co system to be nonplanar, with a torsional angle of 108.7" about the 0-0 bond and with angle Co-0-0 of 120.4 f 0.3". The bond distances are 1.90 f 0.01 A for Co-0, and 1.35 & 0.01 A for 0-0. The latter distance is small for a peroxide group (see Tables I and II), presumably due to the presence of three other oxygen atoms on each cobalt. As was already mentioned, these tend to be weaker u donors than nitrogen atoms, so that less efficient electron transfer from the cobalt atoms into the oxygen bridge is to be expected. The complexes are thermally unstable in


Co ( S o En)

( L H 2 ) 3 - N - - ( b H 2 ) 3

H

Co ( S o Pr)

Fig. 1. The mononuclear cobalt(l1) complexes with ligands N,N'-ethylenebis- (salicylideneimine) and 1V.N'-di-r~propylaminobis(salicylideneimine), abbreviated to Co(SaEn) and Co(SaPr), respectively.

the solid phase at about 80", releasing their oxygen and solvent molecules L quantitatively when heated in zmuo (67,139). The so-called active form of Co"(SaEn) is produced in this way. The oxygen is also released by dissolving the adducts in various inert solvents such as chloroform, dichloromethane, and benzene, but in this case the active form of CoT1(SaEn) is not recoverable from the solutions. Pyridine oxide as well as certain anionic ligands (namely, L = SCN-, N3- , and CH3C0,-) can promote oxygenation even in these solvents, possibly by acting as strong u donors (and poor T acceptors) in the position trans to the oxygen.

Solid samples of active Co(SaEn) can absorb oxygen directly, with evolution of heat. The exact form of the oxygen-containing species in the

Fig. 2. An oxygen-carrying dicobalt complex with amino acids and various other groups L as ligands attached to the cobalt(II1) ions.

BINUCLEAR COMPLEXES OF COBALT 1 3

solid is unknown since no crystals suitable for x-ray analysis have yet been obtained. Presumably L in the solid is absent or is a group from a neigh- boring molecule. The absorption of 0, by such solids is highly reversible at sufficiently high temperatures, and has been extensively studied. The form of the crystals appears to be an important factor in that the rate of absorption and release of 0, depends both on the spaciousness of the lattice and the relative orientation of the monomer molecules. An infrared study of the solid dispersed in KBr has been reported (406). A summary of the solid-state properties of these systems, as well as a study by polarography of the reduction of 0, in the presence of various metal chelates, can be found in the book by Martell and Calvin (257). The preparation of Co(SaEn) has been published in detail (109). Diehl and Henn (108) have suggested and Stewart et al. (374) agreed that water is a necessary com- ponent for oxygen absorption and have postulated that it may form a bridging group holding the cobalt monomers together prior to oxygen uptake. However, evidence against this viewpoint has been given (190), so that the role of water remains controversial, particularly as there is no convincing evidence for H,O bridges in other binuclear complexes. If water is involved, it may be that it or possibly a hydroxyl group serves as the ligand L trans to the oxygen.

A structure determination (426) of an oxygen-carrying cobalt complex of 3-fluoroSaEn (see Fig. 1) has shown this substance to be tetrameric, [(H,O)(F-SaEn)Co .O,. Co(F-SaEn)],, with the two units linked by the sharing of two oxygen atoms from F-SaEn between the central two cobalt atoms. Here, again, the peroxo 0-0 distance is short (1.31 k 0.03 A).

Active forms of CoII(SaEn) have been used to fractionate the oxygen isotopes. The method makes use of the difference in the isotopic abundances in the gas and solid phases when molecular oxygen is equilibrated with the complex. The single-stage enrichment factor (1aO/160),,,/(1aO/'60)solid found by Brown and Drury (55) was 1.012 f 0.002, considerably lower than the earlier value reported by Panchenkov et al. (302,303).

With other related Schiffs base complexes, a tendency to give 1 : 1 and not 2:l (cobalt to oxygen) adducts has been observed. Thus the complex N,N'-di-n-propylaminobis(salicylideneiminato)cobalt(II), included in Figure 1 and abbreviated to Co(SaPr), absorbs oxygen in the solid state to give a 1:1 (or slightly higher) adduct. On the other hand, Fritz and Gretner (145) have obtained evidence for formation of a 2:l adduct with Co(SaPr) in chloroform solution. More recently, Floriani and Calderazzo (1 39) have confirmed that over long reaction times variable amounts of oxygen are in fact absorbed, depending on the solvent used. With pyridine solutions of the complex Co(3-CH30-SaEn), a 1 : l


adduct is found with 0, and has been successfully isolated and characterized by these authors, who suggest from the infrared spectrum of the adduct (lack of an 0-0 stretch) that the 0, group may be symmetrically bonded to the cobalt atom, as in metal-olefin complexes. The adduct, formulated as Co( 3-CH3O-SaEn)(O,)(py), was found to be paramagnetic and appears to have one unpaired electron ( 1 39.279a). This is consistent with presence of cobalt(ll1) with an 0,- ligand or, alternatively, with a lou-spin cobalt(l1) and a spin-paired O2 ligand. Kon and Sharpless (239a) in paramagnetic resonance studies of oxygenated Co(SaPr) (see Figure 1) observed the 1 : 1 adduct and found results consistent with the presence of an equilibrium analogous to reaction scheme ( 1 ) and (2). Similarly, Crunibliss and Basolo (96.97) have isolated and characterized 1 : 1 adducts of the Schiff's base complex bis(acetylacetone)ethylenediiminecobalt(II) with 0,. These adducts also contain one molecule of a Lewis base per cobalt atom, presumably attached trntis to the 0,. They are unstable above 0' and in the absence of the base. Observation of infrared bands in the region 1130-1 140 cni- ' tentatively assigned to 0-0 stretching vibrations suggests unsymmetric bonding Co-0-0 in these complexes. Para- magnetic resonance studies (197,198) confirm that the adducts are indeed mononuclear and best thought of as containing the superoxo group, probably with angle C o o 0 somewhere between 0 and 180". Similarly, crystalline Cob(I1)alamin (vitamin BIZr) and its solutions react reversibly with oxygen to yield a complex shown by magnetic resonance to be mononuclear (33); cob(I1)inamides show analogous behavior. Earlier work (210,240) reporting that vitamin B,,, combines reversibly with oxygen gas to form a 2:1 adduct ivill need to be reexamined, particularly since B,,, is a cobalt(II1) complex. Studies of Co(l1) complexes of N , N ' - bis(o-hydroxybenzylidene)-l,2-diaminoethane have provided evidence for 1 : 1 complexes with 0, in dimethylformamide, in contrast with 2 : l complexes deposited from such solutions ( 1 71). The relative tendencies to form 1 : I and 2:l adducts in all these systems are not yet quantitatively understood, but presumably depend on the effectiveness of electron transfer from the cobalt atom to the attached oxygen molecule.

The combination of molecular oxygen with cobalt(I1) complexes carrying amino acid ligands is well known. A detailed study of the p- peroxobis[dihistidinecobalt(III)] system has been carried out by Hearon et al. (59-61,185-188) and a review of the various a: and /3 amino acid ligands capable of taking up oxygen, and a summary of the relevant literature, has been published (188). The structure for the bis(L-histidine)- cobalt(I1) monomer, determined by x-ray diffraction (183), is shown in Figure 3. The octahedral configuration of ligands around each cobalt in


H

H

Fig. 3. The mononuclear cobalt(I1) ion surrounded by two L-histidine molecules in octahedral coordination.

the mononuclear complex has also been deduced from its effective magnetic moment (4.81 BM at 291"K), visible spectrum (118), and its nuclear magnetic resonance spectrum (266). The optical spectrum (peak at 385 nm, E = 1620) (188) and diamagnetic nature (273) of the oxygenated histidine complex are consistent with the existence of a p-peroxo bridge. Polarographic studies (66,294,295,322,363) are also consistent with this interpretation. It seems probable that the octahedral configuration (Fig. 2; L = imidazole group from one histidine) is maintained after combination with oxygen. The p-peroxo complex has been isolated as brown needles, and an infrared spectral study is in agreement with this proposed structure (336); these authors have also suggested that at each cobalt atom the carboxylate group from one of the histidine ligands is free and the carboxylate group from the other is bonded to the cobalt atom. Kinetic studies relating to the formation and stability of this complex will be discussed in the next section.

The absorption of O2 by solutions containing the cobalt(I1) complex with glycylglycine (NH,CH,CONHCH,CO,H) has been studied in some detail (160,389). Here the brown complex ion formed initially is most probably a p-peroxo complex, and has been assigned the formula [(HO)(GG),Co . 0,. Co(GG),(OH)I2 - . This brown form is increasingly less stable as the pH is decreased, and transforms irreversib1y.into a red species. A brown intermediate with optical spectrum very similar to that of the above has been observed (451) during the oxidation of the cobalt(I1)


glycinate ion with H,O,. Some authors (389,390,445) considered the red complex to be binuclear with p-peroxo-p-hydroxo bridging ligands. However, on the basis of polarographic studies and optical evidence, it was also suggested (35,36,65,66,362) that the red product is really a monomeric cobalt(II1j complex, possibly Co"'(GG),(OH), or Co"'. (GGj,(OH),- or similar species containing H,O ligands. This is consistent with the general decomposition reaction of p-peroxo complexes at low pH. The reaction may proceed via a protonated binuclear intermediate (162). An x-ray structure of the red product has revealed that this complex has a mononuclear structure with two terdentate glycylglycinate rings (161). Recent studies of the acid-base properties have helped elucidate the chemistry of these systems (1 61,267,275a). Comparative studies of the reaction of oxygen with the cobalt(I1) complexes of the diastereomeric ligands DL- and LL-leucyltyrosines have been reported (277,278); the latter in solution turns brown and then red more rapidly than the former. Other cobalt(I1) dipeptides have been investigated (95,160,161,267) and were found to take up oxygen: i t seems likely that the initial products are p-peroxo dicobalt complexes. Red complexes are formed subsequently, as in the above-mentioned systems.

2. Stabilit-y of the p-Peroxo Complexes

In ammoniacal solution, the decaammine complex (NH3),Co.02. Co(NH,),*+ is most stable in a narrow range near pH = 12 ([NH,] - 7:W). This is presumably the range of pH in which the aquo-pentaammine cobalt(I1) complex predominates in the absence of oxygen. The significance of this is clear, in view of the equilibrium

(which is the overall reaction corresponding to eqs. (1) and (2). Thus, for example, the instability of the 4 + ion at pH greater than 12 in aqueous ammonia can be explained by the removal of C O ( N H ~ ) ~ ( H ~ O ) ~ + from the equilibrium in (4) as CO(NH,I, '~ is formed. At pH values lower than 12, the instability is likewise the result of the removal of CO(NH,),(H,O)~+.

The use which Simplicio and Wilkins (365) have made of an oxygen sensor probe to measure oxygen concentrations in aqueous ammonia solutions has enabled them to make a very thorough study of the Co(II)/NH,/O, system. Thus they were able to measure the equilibrium constant K for reaction (4) under a variety of conditions. At 25" in 15M NH,, they found K = 6.3 x 105M -,, and the reaction enthalpy A H N 30 kcal mole-'. In further studies, with stopped-flow techniques, they were able to follow the rate of formation of the peroxo complex, using in most

ZCO(NH~),(H,O)~ + + 0 2 ( N H ~ ) ~ C O . 0 2 . C O ( N H ~ ) ~ ~ + + 2HzO (4)


cases the absorption peak at 360 nm. The data could be fitted to the rate law

d - [O,-adduct] = (~,[CO(NH,),(H,O)~ + I + ~;[CO(NH,),~+])[O,] (5) dt

where the second term on the right proved to be negligible for [NH,] < 9M. The rate constant k , for the reaction of CO(NH,),(H,O)~+ is 2.5 x 104M-' sec-I at 25" and presumably corresponds to the forward reaction in (1); the activation parameters are AH': = 4 kcal mole-, and AS: = - 25 eu. The upper limit for the rate constant k; for the reaction of C O ( N H , ) ~ ~ + is 1.3 x 103M-1 sec-' at 25"; i.e., it is 20 times smaller than kl.

Simplicio and Wilkins have also studied the kinetics of the decomposition of the peroxo complex in ammonia solution, by using ethylenediaminetetraacetic acid (EDTA) to shift the equilibrium (4). This technique is feasible because EDTA reacts with the mononuclear cobalt(I1) ammine complexes rapidly and essentially completely. There appears to be no interference from possible oxygen adducts of the cobalt(I1)-EDTA complexes. The decomposition was found to be independent of [EDTA] and [NH,], and the same whether the p-peroxo complex was produced in situ or by dissolving solid p-peroxo complex. The rate-determining step in the decomposition is thought to be the reverse reaction in eq. (2). At %5", the measured parameters (365) are k - , = 56 sec-l, A H ? , = 18 kcal mole-', and AS?, = +9 eu.

Fremy (143,144), as well as Werner and Mylius (444), reported that the decomposition of the peroxo complex (NH,),Co. 0,. C O ( N H , ) ~ ~ +

in acidic solution is rapid and proceeds according to the equation

( N H ~ ) ~ C O . O ~ . C O ( N H ~ ) ~ ~ + + 2C02+ + 0 2 + ION& (6)

Experiments with oxygen-1 8 have confirmed that the molecular oxygen evolved originates from the p-peroxo complex (330). As has already been mentioned, the p-peroxo complex is at least partially protonated in acidic solutions. Two other studies which indicated that cobalt(II1) complexes are formed appear to be in error because the samples of (NH,),Co.O,. CO(NH,),~ + used either were not pure to start with or had already under- gone some decomposition. Thus Charles and Barnartt (81) reported that only half of the peroxo oxygen was liberated as oxygen gas and that at least eight cobalt species were produced, all but one of these being cobalt(II1) complexes. Jakob and Ogorzalek (208) have also reported the formation of cobalt(II1) complexes. However, work by Rohm (330) has indicated that if sufficient care is taken with the preparation of (NH,),Co. O,.CO(NH,),~ + (i.e., in the preparation, oxygen is bubbled through


solutions for only - 10 min, and care is taken in storing the sample), then the decomposition proceeds as in (6). This is true with most acids, for example, with aqueous solutions of sulfuric acid (0.2-6.OM), perchloric acid (0. I-2.0), hydrochloric acid (O.25), and hydrobromic acid (0.1-1 .O). With nitric acid (> lM), some oxidation of (NH, )5Co .02 .C~(NH3)54+ to the superoxo 5+ complex occurs.

Consistent with these results, the decomposition of the y-peroxo 4+ ion prepared i j 7 situ by reduction in acidic solution of the p-superoxobis- [pentaamminecobalt(III)] 5 + ion has also been shown to proceed according to the reaction (6), with formation of Co(I1) and 0,. This has been observed n i th a number of reducing agents. Gleu and Rehm (163) reported such a result for reduction i\ith As,O, in H,SO, solution. The reactions ith 1 - (lOo,358,385), Fe2 + (101,383), Sn2 + (383), S,032- (383), Cr2+ (196), V 2 + (196), Eu2- (196), and powdered silver (20) and the electrolytic reduction a t platinum electrodes (20) all give Co(I1) and 0, i n more or less quantitative yield. The p-peroxo 4+ ion, present as an intermediate in these reactions, does not seem to be attacked by the reducing agents used. A quantitative study (196) of the decomposition in acidic solution indicates that the p-peroxo complex is rapidly and essentially completely protonated in ca. 1M HCIO, and that the complex loses its proton prior to decomposition. Thus the decomposition is believed to occur via the reactions

h’ ( N H , ) ~ C O . O ~ H . C O ( N H ~ ) ~ ’ ~ ( N H ~ ) ~ C O . O ~ . C O ( N H ~ ) ~ ~ + + H + (7)

(NH&Co.O, C O ( N H & ~ ~ A ZCo*- + IONH, + O,, (8)

\+here Kk = 5.27M-I sec-’ at 25”. Hoffman and Taube (196) suggested that the complex undergoing decomposition has bonding as in formula (111) (p. 8). If this is so, isomerization of (11) to (111) must occur much more rapidly than in the related (en),Co.y(NH2,02).Co(en)23+ system (283).

The conditions in which the ion (NH,),Co.O,. Co(NH,),, + yields mononuclear Co(II1) species are not clearly understood. In acidic solution, it appears that reaction (6) predominates in the decomposition, but other paths seem also to be possible. For example, Khakham and Reibel (226,227) have studied the formation and decomposition of this ion in ammoniacal solution, in the presence of activated charcoal. In this case, the products are CO(NH,),~- and some CO(NH, )~(H,O)~+ ; only half the molecular oxygen absorbed during formation of the complex is liberated. The reaction may be written as

2(NH3)jCo.0z.Co(NH,)j4- + 4NH4’ * 4Co(NH,)S3+ + 2HZO + 0 2 (9)

inorganic reaction mechanismsdownload.e-bookshelf.de/download/0000/5701/08/l-g... · 2013-07-23 ·...

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