Mechanisms of Reactions of Metal Complexes in Solution 9781839161865

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Mechanisms of Reactions of Metal Complexes in Solution

Mechanisms of Reactions of Metal Complexes in Solution By

Debabrata Banerjea

Calcutta University, India Email: [email protected] and

M. K. Bharty

Banaras Hindu University, India Email: [email protected]

Print ISBN: 978-­1-­83916-­186-­5 PDF ISBN: 978-­1-­83916-­360-­9 EPUB ISBN: 978-­1-­83916-­361-­6 A catalogue record for this book is available from the British Library © Debabrata Banerjea and M. K. Bharty 2023 All rights reserved Apart from fair dealing for the purposes of research for non-­commercial purposes or for private study, criticism or review, as permitted under the Copyright, Designs and Patents Act 1988 and the Copyright and Related Rights Regulations 2003, this publication may not be reproduced, stored or transmitted, in any form or by any means, without the prior permission in writing of The Royal Society of Chemistry or the copyright owner, or in the case of reproduction in accordance with the terms of licences issued by the Copyright Licensing Agency in the UK, or in accordance with the terms of the licences issued by the appropriate Reproduction Rights Organization outside the UK. Enquiries concerning reproduction outside the terms stated here should be sent to The Royal Society of Chemistry at the address printed on this page. Whilst this material has been produced with all due care, The Royal Society of Chemistry cannot be held responsible or liable for its accuracy and completeness, nor for any consequences arising from any errors or the use of the information contained in this publication. The publication of advertisements does not constitute any endorsement by The Royal Society of Chemistry or Authors of any products advertised. The views and opinions advanced by contributors do not necessarily reflect those of The Royal Society of Chemistry which shall not be liable for any resulting loss or damage arising as a result of reliance upon this material. The Royal Society of Chemistry is a charity, registered in England and Wales, Number 207890, and a company incorporated in England by Royal Charter (Registered No. RC000524), registered office: Burlington House, Piccadilly, London W1J 0BA, UK, Telephone: +44 (0) 20 7437 8656. Visit our website at www.rsc.org/books Printed in the United Kingdom by CPI Group (UK) Ltd, Croydon, CR0 4YY, UK

Preface Mechanisms of reactions of metal complexes attracted the attention of Alfred Werner, who, based on qualitative observations, proposed an intramolecular mechanism of racemization of tris-chelate complexes of metals that involves opening and closing of a chelate ring, but studies on the kinetics of reactions of metal complexes were first started by Lamb and co-­workers early in the last century. However, during the first half of the twentieth century, many stalwarts in inorganic chemistry did not consider such investigations worth pursuing. In the early 1940s, a young Fred Basolo, after listening to a lecture delivered by the German chemist Walter Hieber on some reactions of metal carbonyls, asked him if he had any idea as to how such reactions occur, namely their mechanisms. Professor Hieber replied, “Young man, we are interested in chemistry, not the philosophy of chemistry”. However, since the late 1940s, the field has started to attract the attention of several groups, mainly in the UK and USA, and in other countries research schools in this field developed in the following years. During the period 1980–2010, publications on the mechanisms of reactions of metal complexes comprised about 15% of the total publications in inorganic chemistry. In the six decades from the 1950s, there has been an explosive growth in this field, significantly more than in any other area of inorganic chemistry, and particularly since the 1960s with the development of several techniques for following fast reactions. Most of the early developments were well covered in the book Mechanisms of Inorganic Reactions: A Study of Metal Complexes in Solution by Fred Basolo and Ralph G. Pearson, the first edition of which was published in 1958 and the second edition in 1967, with a reprint in 1972. A few other books and reviews covering the field, specifically of different individual classes of reactions, have been published, some of them in the latter part of the last century, such as the book The Study of Kinetics and Mechanism of Reactions of   Mechanisms of Reactions of Metal Complexes in Solution By Debabrata Banerjea and M. K. Bharty © Debabrata Banerjea and M. K. Bharty 2023 Published by the Royal Society of Chemistry, www.rsc.org

v

vi

Preface

Transition Metal Complexes by Ralph G. Wilkins, the first edition being published in 1974 and the second edition in 1991. In this book, we have attempted to highlight the major progress and developments in this area, up to the second decade of the present century. The coverage is much more exhaustive than in previously published books, and includes reactions of some s and p block elements also. However, profuse literature compelled us to be selective according to our discretion and we apologise for any omissions. We acknowledge the help received from Dr V. K. Jain, former Head of the Chemistry Division, Bhabha Atomic Research Centre, Mumbai, and Dr S. Goswami, Assistant Professor of Chemistry, Calcutta University, in preparing the drafts of some of the chapters in Word format from the PDF versions prepared by the senior author, and the co-­author prepared one chapter in Word format and made all necessary corrections in the drafts of all the chapters before submission to the publisher. Our thanks are also due to Drew Gwilliams and staff at the Royal Society of Chemistry for the publication of this book. D. Banerjea M. K. Bharty

Contents Chapter 1

 eactions of Metal Complexes  R 1.1 Introduction  1.2 Molecularity and Order of a Reaction  1.3 General Mechanism of Ligand Replacement Reactions  1.4 Experimental Evidence for Mechanisms  1.5 Methods for the Characterization of Reactive Intermediates  1.6 Order of Reaction and Reaction Mechanisms  1.7 Lability of Complexes  1.7.1 Ligand Field Theory  References 

Chapter 2 T  echniques for Following Reactions and Factors that Affect Rates  2.1 Techniques for Evaluating Rates of Reactions  2.1.1 Direct Chemical Analysis  2.1.2 Spectrophotometric Methods  2.1.3 Electrometric Methods  2.1.4 Polarimetric Methods  2.1.5 Use of Isotropic Tracers  2.2 Special Techniques for Studying Fast Reactions  2.2.1 Flow Methods  2.2.2 Electrochemical Methods  2.2.3 Relaxation Methods  2.2.4 Flash Photolysis 

  Mechanisms of Reactions of Metal Complexes in Solution By Debabrata Banerjea and M. K. Bharty © Debabrata Banerjea and M. K. Bharty 2023 Published by the Royal Society of Chemistry, www.rsc.org

vii

1 1 5 8 16 18 21 26 30 52 59 59 59 60 63 66 67 67 67 69 70 71

Contents

viii



2.2.5 Nuclear Magnetic Resonance (NMR) and Electron Paramagnetic Resonance (EPR) Methods  2.2.6 Studies of Exchange Rates by NMR Spectroscopy  2.2.7 Electron Paramagnetic Resonance (EPR) Spectroscopy  2.2.8 Application of Ultrasonic Absorption and Pulse Radiolysis  2.3 Rate Constants in Some Complex Systems  2.4 Factors that Affect Rates of Reactions  2.4.1 Effect of Temperature  2.4.2 Effect of External Pressure  2.4.3 Effect of Ionic Strength  2.4.4 Influence of Solvent  2.5 Nucleophilicity and Rate  2.6 Relative Nucleophilicities  2.7 Hammett Relationship  2.8 Taft Relationship  2.9 Linear Free Energy Relationship (LFER)  2.10 Isotope Effects  References 

Chapter 3 L  igand Replacement Reactions of Metal Complexes of Coordination Number Four and Higher  3.1 Square-­planar Complexes  3.1.1 Complexes of Platinum(ii)  3.1.2 trans Effect Theories  3.1.3 π-­Bonding Theory of the trans Effect  3.1.4 Mechanism of Reaction  3.1.5 Energy Profile for Reactions of Square-­planar Complexes  3.2 trans Effect in Platinum(iv) Complexes  3.3 Other Square-­planar Metal Complexes  3.3.1 Palladium(ii) Complexes  3.3.2 Gold(iii) Complexes  3.4 Further Comments on trans-­and cis Effects  3.4.1 trans Effect in Terms of Discrimination  3.4.2 cis-­Labilizing Effect  3.4.3 cis Effect in Terms of Discrimination  3.5 Square-­planar Complexes of Nickel(ii) and Copper(ii)  3.6 Reactions of Tetrahedral Complexes  3.7 Complexes of Coordination Number Five  3.8 Complexes of Higher Coordination Number  References 

71 72 75 76 76 80 80 84 87 89 89 90 94 96 97 97 99 105 105 105 107 121 125 154 155 155 156 161 163 164 165 165 167 172 177 182 183

Contents

Chapter 4 L  igand Replacement Reactions of Octahedral Complexes  4.1 Aquation/Solvolysis, Anation/Formation and Ligand Exchange Reactions  4.1.1 Effect of Leaving Ligand  4.1.2 Effect of Charge on Reaction Rate  4.1.3 Steric and Structural Effects of Spectator Ligands  4.1.4 Electronic Effects of Spectator (Non-­leaving) Ligands  4.1.5 Stereochemical Change Accompanying a Ligand Replacement Process  4.1.6 Other Evidence  4.1.7 Activation Parameters and Reaction Mechanism  4.1.8 Solvent Effect  4.1.9 Comparison of the Rate of Replacement of a Metal-­bound Aqua Ligand in a Metal Complex With the Rate of Exchange of the Bound Water With the Water in the Solvent  4.1.10 Stoichiometric Mechanisms  4.2 Base Hydrolysis  4.3 Ligand Replacement Reactions of [M(CO)6]  4.4 Reactions of s-­ and p-­Block Metals  References  Chapter 5 Catalysed Reactions and Formation Reactions  5.1 Electrophilic and Nucleophilic Catalysis  5.1.1 Acid Catalysis  5.1.2 Electrophilic Catalysis by Metal Ions  5.1.3 Nucleophilic Catalysis  5.1.4 Electron Transfer Mechanism of Reactions of Metal Complexes  5.2 Formation Reactions  References  Chapter 6 I somerization, Optical Inversion and Racemization Reactions  6.1 Linkage Isomerization  6.2 Geometrical Isomerization  6.2.1 Square-­planar Complexes  6.2.2 Octahedral Complexes  6.3 Other Types of Structural Isomerization  6.4 Optical Inversion  6.5 Optical Isomerization (Racemization)  6.5.1 Intermolecular Mechanism 

ix

194 194 194 198 200 206 207 211 218 223

226 228 231 246 247 248 255 255 256 276 282 285 286 335 347 347 350 351 353 360 362 364 365

x



Contents

6.5.2 Intramolecular Mechanism  6.5.3 Isomerization and Racemization of Tris Chelates of Unsymmetrical Chelating Ligands  6.6 Structural Changes in Complexes of Terdentate Ligands of the Type M(L′)(L″)  6.7 Structural Changes in Four-­coordinate Complexes  6.8 Optical Isomerism in Tetrahedral Complexes  6.9 Configurational Changes in Some Planar Complexes  References 

Chapter 7 E  lectron Transfer Reactions  7.1 Introduction  7.2 O  uter-­sphere Mechanism  7.3 The Marcus Equation: Marcus Cross-­relation and Its Applications  7.4 Inner-­sphere Mechanism  7.4.1 Atom (or Group) Transfer Processes  7.5 Comproportionation  7.6 Mixed Outer-­ and Inner-­sphere Reactions  7.7 Estimation of Redox Rate Constants for Inner-­sphere Reactions  7.8 Electron Transfer Reactions in Heterogeneous Systems  7.9 Solvated Electrons  7.10 Oxidative Addition Reactions  7.10.1 Mechanisms of Oxidative Addition  7.10.2 Five-­coordinate Eighteen-­electron Substrates  7.10.3 Four-­coordinate Sixteen-­electron Substrates  7.10.4 Four-­coordinate Eighteen-­electron Substrates  7.11 Reductive Elimination  References  Chapter 8 A  ctivation of Molecules by Coordination and Reactivity of Coordinated Ligands  8.1 Introduction  8.2 Activation of Some Diatomic Molecules  8.2.1 Activation of Dihydrogen by Coordination  8.2.2 Activation of Dioxygen by Coordination  8.2.3 Activation of Dinitrogen by Coordination  8.3 Reactivity of Coordinated Ligands  8.3.1 Reaction of Metal-­bound CO Ligand  8.3.2 Reactions of Coordinated CO2 and SO2 

367 372 376 377 377 378 380 386 386 388 388 394 395 411 412 412 413 416 419 422 424 426 426 427 429 435 435 435 435 441 445 449 449 450

Contents



xi

8.3.3 Reactions of Coordinated NO in Nitrosyl Complexes  8.3.4 Reactions of Some Coordinated Organic Ligands  8.3.5 Catalysed Oxidation of Coordinated Ligands  8.3.6 Amino Acid Ester Hydrolysis  8.3.7 Decarboxylation of β-­Keto Acids  8.3.8 Wacker Process  8.4 Insertion Reactions  8.4.1 Insertion of CO  8.4.2 Insertion of Sulfur Dioxide  8.4.3 Insertion of Carbon Dioxide  8.4.4 Insertion of Carbon Disulfide  8.4.5 Insertion of Olefins (Alkenes)  8.4.6 Olefin (Alkene) Polymerization  8.4.7 Acetylene (Alkyne) Polymerization  8.4.8 Asymmetric Synthesis Catalysed by Coordination Compounds  References 

451 452 457 460 461 467 467 470 476 478 480 480 482 484 487 488

Chapter 9 P  hotochemical Reactions of Metal Complexes  9.1 Introduction  9.2 Experimental Results  References 

497 497 501 517

Subject Index

522

Chapter 1

Reactions of Metal Complexes 1.1  Introduction Mechanisms of reactions of metal complexes attracted the attention of the early investigators of coordination compounds, particularly Alfred Werner who, based mainly on qualitative observations proposed an intramolecular mechanism for racemization of tris-chelate complexes of metals (see Chapter 6). But extensive studies began in different laboratories in the late 1940s, and since then there has been an almost explosive growth in the field. The development of methods and techniques for studying fast reactions added a new dimension since the 1960s. Most of the early developments are well documented in the authoritative treatise by Basolo and Pearson,1 the first edition of which appeared in 1958. Since then, a number of books and reviews have appeared on various aspects of the subject.2–9 Studies on the rates of reactions of metal complexes for the elucidation of reaction mechanisms have been a fascinating field of research of coordination chemists, which has already reached a high level of sophistication and maturity. Here (in contrast to organic compounds), we come across a variety of reaction centres, viz., metal ions, with widely varying properties that are primarily dependent on their electronic configurations and their oxidation states, and several different types of stereochemistry of their complexes and innumerable types of ligands, all of which have a profound influence on the dynamics (rate) and mechanisms of their reactions. Such knowledge has been well exploited in the synthesis of many novel compounds and various applications of metal complexes.

  Mechanisms of Reactions of Metal Complexes in Solution By Debabrata Banerjea and M. K. Bharty © Debabrata Banerjea and M. K. Bharty 2023 Published by the Royal Society of Chemistry, www.rsc.org

1

Chapter 1

2

A reaction A + B → Products can occur only through an encounter of the reactant species. A more complex reaction involving several reactant species obviously cannot proceed in one step as the simultaneous encounter of several species is extremely unlikely. Such a reaction therefore occurs through a number of simple (elementary) steps, each involving the encounter of generally not more than two and rarely of three species; the sum total of all these steps accounts for the stoichiometry for the overall process:

Overall: Each of these steps will proceed at a specific rate, but of these, generally, but not always, the slowest will be the rate of the overall process, provided that the others are all sufficiently faster; this slowest step is regarded as the rate-­determining step. The objective in the elucidation of the mechanism of a reaction is to identify, on the basis of kinetics and other evidence, the various elementary steps involved and the nature of the transition state for the rate-­determining step. Although it is often possible to make an unequivocal assignment of a mechanism, there are many situations where one can at best propose a plausible mechanism that adequately explains all the experimental observations. According to the principle of microscopic reversibility, for the process kf

A  B#C  D kb

the value of the equilibrium constant K shows how far the reaction proceeds at equilibrium, while kf indicates how fast the change takes place, where K = kf/kb. Hence, one cannot fully understand the nature and behaviour of the system based on either the K value or one of the k values alone. This is because a very large value of K merely indicates kf ≫ kb, irrespective of their actual magnitudes, and a similar argument holds good also for small K values (i.e. kb ≫ kf ). Although a reaction may be thermodynamically feasible, it may not proceed at an observable rate because of a lack of a suitable reaction path, and it is of interest to understand the factors that govern this behaviour. Studies on the kinetics of reactions provide a means for the quantitative comparison of various reactions and also help to provide a fuller understanding of the nature of a system. From systematic kinetic studies, it is also possible to ascertain the relative importance of the many factors that may influence reaction rates, and also to infer about the different steps in the overall change, i.e. to gain knowledge about the reaction mechanism.

Reactions of Metal Complexes

3

A very large number of the reactions encountered in inorganic systems (including those of metal complexes) may be classified as follows:    (i) Reactions involving structural changes, such as isomerization, e.g. interconversion of geometric (cis ⇌ trans, fac ⇌ mer) and configurational isomers (e.g. square planar ⇌ tetrahedral), interconversion of linkage isomers, such as the conversion of nitritopentaamminecobalt(iii) ion into nitropentaamminecobalt(iii) ion, and of optical isomers (i.e. racemization). Another type is conformational isomerism, which arises from δ and λ conformations of chelate rings of diamines such as en, pn, etc., in complexes. (ii) Nucleophilic and electrophilic substitutions in metal complexes, i.e. replacement of a ligand by another (nucleophile) or replacement of a metal by another (electrophile), which are broadly speaking acid–base reactions involving Lewis acids and bases. Ligand exchange and metal exchange reactions of metal complexes also belong in this category. Formations of metal complexes from aqua metal ions are also ligand substitution reactions, involving replacement of an aqua ligand by another ligand. Replacement of a simple anionic monodentate ligand by a solvent molecule in solution is conventionally called solvolysis; the term aquation is used if the solvent is water. Such reactions are carried out in aqueous solution at pH < 5 to avoid what is known as base hydrolysis (replacement of the anionic ligand by OH−) and is also referred to as acid hydrolysis. The corresponding reaction involving loss of an uncharged ligand such as NH3, py, etc., or of a bidentate or polydentate ligand (irrespective of whether this is charged or uncharged) is conventionally referred to as a dissociation reaction. The reverse of the aquation reaction, i.e. replacement of an aqua ligand in a complex by a simple anionic monodentate ligand, is referred to as an anation reaction, and is termed a formation reaction if the entering ligand is an uncharged monodentate ligand, or a bidentate or a polydentate ligand (charged or uncharged).    (iii) Electron transfer reactions either involving a net chemical change (such as a redox reaction), or involving no net chemical change, such as the electron exchange reaction in the [Fe(CN)6]4−–[Fe(CN)6]3− system, viz. *FeII + FeIII ⇌ *FeIII + FeII (iv) Reactions of coordinated ligands. The reactivity of a species is often profoundly altered on binding to a metal and as a result it may undergo novel types of reactions that are not possible for the free species. Some reactions of the free species that occur under drastic conditions occur in a facile manner under far less drastic or even ambient conditions in the metal-­bound species. Thus, the N2 ligand in many complexes can be converted to NH3 and/or N2H4 under mild conditions.

Chapter 1

4

(v) Insertion reactions. These are reactions in which a small molecule inserts into a metal–ligand bond, as in the following examples (see Chapter 8, Section 8.4): M–H + O2 → M–OOH M–R + SO2 → M–OSOR M–CR3 + CO → M–C(O)CR3 M–R + CO2 → M–C(O)OR Also known are migration–insertion reactions, as in the following example (see Chapter 8, Section 8.4.1): Mn(CO)5Me + *CO → Mn(CO)4(*CO){C(O)Me} The product is formed by migration of the Me to a cis-­CO and entry of *CO into the site vacated by Me. (vi)  Oxidative addition reactions. Some examples of these are the following: Ir(PPh3)2(CO)X + O2 → Ir(PPh3)2(CO)(O2)X 2[Co(CN)5]3− + RX → [RCo(CN)5]3− + [Co(CN)5X]3− [Pt(en)2]2+ + Cl2 → trans-­[Pt(en)2Cl2]2+ The reverse of oxidative addition is reductive elimination; many oxidative additions are reversible (see Chapter 7).    Cluster formation reactions have also attracted attention in recent years. Studies on MoFe310 and Rh411 cluster formation reactions and mechanistic aspects of some reactions of Mo3M (M = Pd, Pt) clusters12 and of W3Pd13 have been reported. Photochemical (light-­induced) reactions of metal complexes also have interesting features. Photoracemization, photosubstitution and photoredox reactions of metal complexes, and also photochemical reactions of metal carbonyls and complexes of other π acid ligands, have been studied. In many cases, thermal and photochemical reactions have different pathways, leading to different products (see Chapter 9). Ligand replacement in metal complexes and electron transfer reactions are by far the most important types of reactions in inorganic systems. Ligand replacement reactions generally involve the rupture of the metal– ligand bond, but there are cases in which replacement of one ligand by another may occur without any fission of the primary metal–ligand bond. i.e. the bond between the metal and the donor atom of the ligand, and such a processes is called pseudo-­substitution, as in the case of the dissociation of [Co(CO3)(NH3)5]+ in an acidic medium to form [Co(NH3)5(OH2)]3+, and

Reactions of Metal Complexes

5 2+

3+

the formation of [Co(ONO)(NH3)5] from [Co(NH3)5(OH2)] and nitrite ion in weakly acidic buffered medium at low temperature. Another interesting case is ligand replacement by an electron-­transfer mechanism (see Chapters 3 and 5). There has been a profuse growth in the literature on inorganic reaction mechanisms, and developments in the field during 2009–2012 have been highlighted in the Annual Reports14 published by The Royal Society of Chemistry. A review on ligand substitution reactions has been published15 and also a broader review on mechanistic aspects of reactions of metal complexes16 (for earlier reviews, see ref. 9). An entire issue of Dalton Transactions contains a series of papers on “Inorganic Reaction Mechanisms: An Insight into Chemical Challenges”.17

1.2  Molecularity and Order of a Reaction Molecularity expresses the total number of species (n) involved in the rate-­ determining step of the overall reaction, and order expresses the sum of the powers of the concentrations of the species that appear in the rate law expression. Thus we have unimolecular (n = 1) and bimolecular (n = 2) reactions; termolecular (n = 3) reactions are very rare, since the simultaneous encounter of three species is likely to be very infrequent, and reactions of higher molecularity are unknown, being almost impossible; molecularity is always integral. If for a reaction A+B→C+D the rate law (which expresses the dependence of the rate on the concentrations of various species) is of the form Rate = k[A]x[B]y where k is the rate constant, the order of the reaction is x + y. For a first-­order reaction, x = 1 and y = 0, or x = 0 and y = 1. For a second-­order reaction, x = y = 1, or x = 2 and y = 0 or x = 0 and y = 2. One often speaks of the order with respect to a particular reacting species, and in the above example the order with respect to A is x and that with respect to B is y. The order can be integral, fractional or even zero. If the rate law has a concentration term for a particular species in both the numerator and the denominator [as in eqn (1.54)], then it is not possible to mention a physically significant order for the reaction or order with respect to that species. However, for a rate law of the form Rate = k[A][B]/[H+] we may say that the order with respect to both A and B is 1 and the reaction is inverse first order in H+, but the overall order for the reaction will have no significance here also.

Chapter 1

6

A zeroth-­order reaction is of the type A → B and the rate is independent of the concentration of A. Heterogeneous reactions such as the thermal decomposition of CaCO3(s) are of zeroth order. For a zeroth-­order A → B transformation, the rate is given by –d[A]/dt = k hence 18

  



[A]t = [A]0 − kt

(1.1a)

  

and a plot of [A]t versus t is straight line with intercept [A]0 and slope −k. The units of k are obviously M s−1, i.e. mol dm−3 s−1. For a reaction having a rate law Rate = k[A]a[B]b[C]c if a = 0 and the species A is used in deficiency, then –d[A]/dt will not vary with change in concentration of A and hence a pseudo-­zeroth-­order rate behaviour will be observed. For a first-­order reaction of the type A → B, the rate expression is –d[A]/dt = k[A] which leads to

  



log[A]t = log[A]0 − kt/2.3

(1.1b)

  

Hence a plot of log[A]t versus t is linear with slope –k/2.3 (k in s−1). For a reaction A → B that is second order in A, we have –d[A]/dt = k[A]2 which leads to   



1/[A]t = 1/[A]0 + kt

(1.1c)

  

Hence a plot of 1/[A]t versus t is linear with slope k (k in M−1 s−1) (see ref. 7, 1st edn, p. 8 for zeroth-­order and p. 10 for first-­and second-­order reactions). However, the most common form of second-­order reaction is of the type A + B → Product(s) In this case, if the reaction is carried out with equal concentrations of the reactants A and B, then –d[A]/dt = –d[B]/dt = k[A][B] = k[A]2 = k[B]2 Hence in this situation eqn (1.1c) is valid, but if the initial concentrations of A and B are different then the following expression for the k value can be derived:19   

  

log([B]t/[A]t) = {[B]0 − [A]0}kt/2.3 + log([B]0/[A]0)

(1.1d)

Reactions of Metal Complexes

7

and the k value can be evaluated graphically using this equation. However, it is more convenient to carry out the reaction under pseudo-­first-­order conditions by using a large excess of one of the reactants, e.g. with [B]0 at least 20 times (but preferably 100 times) [A]0, and evaluate the pseudo-­first-­order rate constant kobs graphically using eqn (1.1b), where   

kobs = k[B]0



(1.1e)

  

This allows the evaluation of the second-­order rate constant k graphically using experimentally evaluated kobs values at different (large excess) concentrations of B. Both of the above cases are examples of bimolecular second-­order reactions. However, the reaction conforming to eqn (1.48) will have the rate law expressed in eqn (1.55a) if K is small such that the concentration of I is negligible; the reaction will be of second order but unimolecular, since the rate-­ determining step is the conversion of I to P. For a termolecular third-­order reaction between A, B and C forming product P, if the reaction is carried out with equal concentrations of A, B and C, then it can be shown20 that the rate constant k is given by   

k = (1/2t)/{(1/[A]t2) − (1/[A]02)}



(1.1f)

  

Although termolecular third-­order reactions are extremely rare, bimolecular third-­order reactions are well known. Such a situation is met in the following reactions: K

A  B# I (rapid equilibrium forming I, equilibrium constant K ) k  P (rate-determining step (bimolecular), rate constant k ) I  C 

For this, the rate of formation of the product P is   



d[P]/dt = k[I][C] = kK[A][B][C]

(1.1g)

  

eqn (1.1g) will be strictly valid if K is small such that concentration of I is negligible compared with those of A and B. A third-­order rate will also be observed if all three reactants A, B and C form the intermediate I followed by its (unimolecular) transformation into P. Well-­known examples of this behaviour are the Pt(ii)-­catalysed ligand replacement reactions of Pt(iv) complexes21 (see Chapter 3). For a reaction of three reactants A, B and C forming the product P, if the rate conforms to Rate = d[A]/dt = k[A]a[B]b[C]c

Chapter 1

8

the order with respect to A can be ascertained by carrying out the reaction in the presence of fairly high concentrations of B and C compared with that of A, which is at rather low concentration, and establishing how the concentration of A changes with time during the course of the reaction. Under this condition, Rate = kobs[A]a The experimentally evaluated rate constant (kobs) under this condition is related to the specific rate constant k as follows: kobs = k[B]b[C]c To ascertain the order with respect to B, kobs values are obtained for different (excess) concentrations of B but a fixed concentration (excess) of C; for this, the value of kobs is kobs = kC[B]b The slope of the linear plot of log kobs versus log[B] is b. The order with respect to C is obtained similarly by determining kobs at different (excess) concentrations of C using a fixed concentration of B: kobs = kB[C]c

1.3  G  eneral Mechanism of Ligand Replacement Reactions Taking the case of a metal complex having a replaceable ligand X with all the other ligands not taking part in the reaction (and hence called spectator ligands) and denoting this M–X (omitting the spectator ligands), we may consider the following as the most general mechanisms for ligand replacement (i.e. replacement of X in M–X by Y forming M–Y):   

M–X + Y → M–Y + X



  

(1.2a)

1. Dissociative mechanism (D):   

  

slow ( k )

1 M – X  M  X

fast M  Y  M– Y



For such a process, the rate of formation of M–Y is given by    d[M–Y]/dt = −d[M–X]/dt = k1[M–X]   

(1.2b)

(1.2c)

Generally, there will be a competition for recapture of X, and taking that into consideration we have   



  

k1

M – X#M  X; k1

k2 M  Y   M – Y

(1.3)

Reactions of Metal Complexes

9

Using the steady-­state principle d[M]/dt = 0 to evaluate the concentration of M, we obtain the rate law   

d[M–Y]/dt = k1k2[M–X][Y]/{k−1[X] + k2[Y]}



(1.4)

  

This under limiting situations leads to the following:   

(a) k2[Y] ≫ k−1[X]: 

d[M–Y]/dt ≈ k1[M–X]

(1.5)

(b) k−1[X] ≫ k2[Y]:  

d[M–Y]/dt ≈ (k1k2/k−1)[M–X][Y]/[X]

(1.6)

  

  

2. Associative mechanism (A):   

slow ( k )



fast 2 Y  M – X   Y – M – X  Y – M  X

(1.7)

d[M–Y]/dt = k2[M–X][Y]

(1.8)

  

For this process   

  

However, in general the rate-­determining step will be preceded by a fast pre-­equilibrium forming an outer-­sphere complex M–X·Y, as discussed below for the I mechanism, and the rate law will be of the form shown in eqn (1.12). 3. Examples: Examples of true D and A mechanisms are scarce in metal complexes, and in most cases the replacement takes place by an interchange process (I) in which both dissociation of the M–X bond and formation of the M–Y bond are synchronous, leading to a transition state. However, in such a case one can make a further subdivision into Id and Ia processes, depending on whether the M–X bond dissociation or M–Y bond formation is relatively more predominant in generation of the transition state. The operation of this mechanism in its simplest form may by represented by

  



k Y  M – X   Y  M X  Y – M  X

(1.9)

  

The rate law for this will be similar to that of the A process:   

d[M–Y]/dt = k[M–X][Y]



(1.10)

  

In this, however, it is more appropriate to consider a differential encounter as a fast pre-­equilibrium forming the encounter (outer-­sphere) complex, M–X·Y, in which Y replaces a solvent molecule in the first solvation layer of M–X, followed by its transformation in the rate-­determining step:   



  

K OS

k M – X  Y #M – X  Y   Y  M X  Y – M  X



(1.11)

Chapter 1

10

This leads (when [Y] ≫ [M–X]) to   

d[M–Y]/dt = kKos[M–X][Y]/(1 + Kos[Y])



(1.12)

  

In limiting situations, this becomes   

(a) Kos[Y] ≪ 1: 

d[M–Y]/dt ≈ kKos[M–X][Y] ≈ k′[M–X][Y]

(1.13)

(b) Kos[Y] ≫ 1: 

d[M–Y]/dt ≈ k[M–X]

(1.14)

     

   Scheme 1.1 illustrates the broad classification of intimate and stoichiometric mechanisms. In a situation where bond making and bond dissociation are of comparable importance, the interchange process is designated I. These designations, now universally used by inorganic chemists, were first proposed by Langford and Gray,3 and they correspond to the conventional classical nomenclature (according to Ingold's notations for organic substitution reactions) as follows:    D: SN1 (limiting); Id: SN1 A: SN2 (limiting); Ia: SN2   

Scheme 1.1

Reactions of Metal Complexes

11

We have just seen that the rate law as such and without other evidence does not enable us to postulate the mechanism even in the broadest sense, A or Ia, or D or Id. This will further be obvious if we consider the following few typical situations often encountered in reactions of metal complexes in a good coordinating solvent:   

k fast  M – S  X; M – S  Y  M – Y  S (S  solvent module) (a)M – X  S  (1.15)

  

Irrespective of whether the rate-­determining step (with rate constant k) is associative or dissociative, the rate law for the process will be   

d[M–Y]/dt = k[M–X]



(1.16)

  

The concentration of the solvent, [S], is too large and is constant and so even in an associative process the rate constant (gross) will be k[S] = constant. Hence, in such a situation, irrespective of whether the mechanism is associative or dissociative, we obtain a rate law corresponding to that for a D process:   

k1

k2 (b)M – S#M  S; M  Y   M – Y



k1

(1.17)

  

Applying the conventional steady-­state principle, the concentration of the reactive intermediate M can be evaluated and used to deduce the rate expression [eqn (1.20)]:   



d[M]/dt = k1[M–S] − k−1[M][S] − k2[M][Y] = 0

(1.18)

[M] = k1[M–S]/(k−1[S] + k2[Y])

(1.19)

d[M–Y]/dt = k2[M][Y] = k1k2[M–S][Y]/(k−1[S] + k2[Y])

(1.20)

  

Hence,   

  

Therefore,   

  

Two limiting situations are possible:    (i) If k2[Y] ≫ k−1[S], eqn (1.20) transforms into   

d[M–Y]dt = k1[M–S]



(1.21)

     

(ii) If, however, k−1[S] ≫ k2[Y], then eqn (1.20) takes the form

  

  

d[M–Y]/dt = (k1k2/k−1)([M–S][Y]/[S])

(1.22)

Chapter 1

12

However, since the concentration of the solvent, [S], in dilute solution is extremely high and hence virtually constant, eqn (1.22) is of the form   

d[M–Y]/dt = k′[M–S][Y]



  

(1.23)

where k′ = k1k2/k−1[S] is the experimentally observed pseudo-­first-­order rate constant (kobs). Hence the same mechanism may lead to a rate law similar to that of a dissociative or an associative mechanism, depending on the relative values of the component rate constants. Therefore, the rate law, which shows the order of the reaction with respect to the various reactants, does not by itself lead to any definite conclusion regarding the reaction mechanism – hence the need to carry out investigations to gather other evidence to enable a plausible mechanism to be proposed, i.e. the one that can best explain all the observed facts. This is often not straightforward and involves a lot of logical speculations and arguments based on reasoning. However, one may encounter a situation where two alternatives may almost equally well appear to be acceptable. It is of further interest to note in this connection that for a reversible process, with Y in excess to have a pseudo-­first-­order condition, [Y] ≫ [M–S], in the reaction   

k1

M – S  Y#M – Y  S



  

the rate law is

k1



(1.24)

†

  

Rate = (k1[Y] + k−1)[M–S]   

(1.25)

Hence,   

kobs = k−1 + k1[Y]



(1.26)

  

An identical rate law is observed if the reaction occurs by the following concurrent paths (i) and (ii):   

k1 (i) M – S   M  S;

  



  

fast M  Y  M – Y

k2 (ii) M – S+Y   M – Y + S

(1.27) (1.28)

for which   

Rate = k1[M–S] + k2[M–S][Y] = (k1 + k2[Y])[M–S]

(1.29)

  

Hence   

kobs = k1 + k2[Y]



(1.30)

  

†

This is because for A ⇌ B, the rate of attainment of equilibrium is (kf + kb)[A], where kf and kb are the forward and reverse rate constants, respectively.

Reactions of Metal Complexes

13

Thus, additional evidence is needed to choose between the two possible alternatives.    The replacement of X in M–X with Y generally follows the Eigen–Wilkins mechanism (see Chapter 5) proposed (during 1960–1965) in connection with detailed investigations on the formation of metal complexes from aqua metal ions and ligands in solution. It is logical to believe that in order to enable Y to take part in the product formation it must be suitably paced in a position in the first solvation shell of M–X, forming what is known as an outer-­sphere complex. For species that are oppositely charged, this outer-­ sphere complex will be an ion pair. However, such outer-­sphere association is possible even between a charged and an uncharged species, and also between species of the same charge type, say an anionic complex and an anionic entering ligand, due to various types of weak electrostatic interactions, including hydrogen bonding interactions in suitable cases (see Chapter 5). The outer-­sphere complex may transform into the product by either a D, an I or an A process, as shown in Scheme 1.2, but commonly Id or Ia instead of pure D or pure A, since L is favourably disposed in the outer-­sphere complex to enter into a bonding interaction with M to a certain extent in the transition state. Scheme 1.2 may be represented in a general way by   

  

K OS

kos M – X  L#(M – X) – L  M – L  X



(1.31)

This leads, under pseudo-­first-­order conditions attained using excess L, to   



kobs = kosKos[L]/(1 + Kos[L])

  

Scheme 1.2

(1.32)

Chapter 1

14

(a) If the conditions are such that Kos[L] ≫ 1, then kobs ≈ kos and this limiting value will be reached at a high concentration of L. (b) However, if the conditions for the system are such that Kos[L] ≪ 1, the above relation leads to   



kobs ≈ kosKos[L] = kf[L]

(1.33)

  

where kobs is the observed pseudo-­first-­order rate constant and kf is the second-­order formation rate constant, the reaction being first order with respect to each of the reactants, viz. M–X and L. Eqn (1.32) can be used for the graphical evaluation of Kos and kos; a plot of 1/kobs versus 1/[L] gives a straight line with intercept 1/kos and the intercept/slope ratio is equal to Kos. From eqn (1.33),   



kobs/Kos = kf/Kos = kos

(1.34)

  

If X = H2O (i.e. an aqua complex), the k value so obtained in (a) or (b) may be compared with kex (the rate constant for water exchange of M–OH2) to infer about the mechanism. Theoretically, kos may be expected to show some dependence on the nature of L. However, there is evidence22 that outer-­ sphere association of M–OH2 with L does not markedly alter its kex value, i.e. the lability of M–OH2. Hence for a dissociative activation kos will be comparable to kex; it may be even less but it will never exceed kex. A lower value is possible owing to competition between L and H2O for the vacated coordination position of the metal centre, and here, as the concentration of H2O is very high in the solution, it will have an overwhelming effect, particularly when the entering ligand is of low nucleophilicity. However, for associative activation kos may be much larger than kex for a good entering nucleophile. In order to apply this criterion one may evaluate Kos by analysis of kinetic data in some cases23 and experimentally by various techniques such as spectrophotometry, potentiometry, conductometry, polarography, etc., in favourable cases,24 but it can always be estimated by theoretical calculations using the Fuoss–Eigen equation.25 Kos is appreciably large for interaction of SO42−, a 2− anion, with 2+ and 3+ cationic complexes, e.g. 2.9 × 102 for [Co(Cl) (NH3)5]2+, 2 × 103 for [Co(NH3)6]3+ and 2.8 × 103 for [Co(en)3]3+ (all values at 25 °C, I ≈ 0).   



Kos = (4πNa3/3 × 103)exp[–U(a)/kT]

(1.35)

  

where U(a) is the Debye–Hückel interionic potential (coulomb energy):   



U(a) = (z1z2e2/aD) − [z1z2e2κ/D(1 + κa)]

(1.36a)

κ2 = 8πNe2I/103DkT

(1.36b)

  

and   

  

Reactions of Metal Complexes

15

with N = Avogadro's number, a = distance of closet approach of two ions bearing the charges z1 and z2 (a is usually taken as 5 Å), k = Boltzmann's constant, e = charge carried by an electron (in esu units), D = bulk dielectric constant, I = ionic strength and z1, z2 = number of units of charge on the reactants. Eqn (1.35) is applicable only for ligands bearing a charge; a similar equation has been derived for uncharged (neutral) ligands.26 The calculation has been applied to cationic–anionic27 and cationic–zero charged species28 and also cationic–cationic29 interactions. In most cases, the agreement between the calculated and experimentally observed values is fair. Some typical Kos values obtained by such calculations are given in Table 1.1. If, however, the formation of M–L from M–OH2 takes place by a D mechanism, without the formation of an outer-­sphere complex in a pre-­equilibrium, the reaction scheme is   



k1

k2 M – OH2 #M  H2 O, M  L  M – L k1

(1.37a)

  

This leads to [cf. eqn (1.20)]   



d[M–L]/dt = k1k2[M–OH2][L]/{k−1[H2O] + k2[L]}

(1.37b)

  

From this, it follows that (c) when k−1[H2O] ≫ k2[L],   



d[M–L]/dt ≈ (k1k2/k−1)[M–OH2][L]/[H2O]

(1.37c)

  

Since in aqueous solution [H2O] is virtually constant,   



d[M–L]/dt ≈ kf[M–OH2][L]

(1.37d)

  

where kf = (k1k2/k−1[H2O]) If H2O is X (as in M–X), a rate retardation by X in solution will be observed [see eqn (1.6)]. Table 1.1  Typical  Kos values obtained using the Fuoss–Eigen equation. System

Kos/M−1, at 25 °C (I ≈ 0)

[Ni(OH2)6]2+·MePO42− [Ni(OH2)6]2+·CH3CO2− [Ni(OH2)6]2+·SCN− [Ni(OH2)6]2+·NH3 [Ni(OH2)6]+·enH+ [Fe(OH2)5(OH)]2+·[Fe(OH2)5(OH)]2+

14 3 1 0.15 0.02 1.3 × 10−3

Chapter 1

16

(d) When k2[L] ≫ k−1[H2O], as may happen when [L] is very large and L is a much better nucleophile than H2O,   



−d[M–L]/dt ≈ k1[M–OH2]

(1.37e)

  

This means that for different entering nucleophiles, L, the same limiting rate will be reached at a fairly high concentration of L in the solution and here the formation rate constant is   



kf ≈ k1 ≈ kex

(1.37f)

  

Hence it is not easy to distinguish between a D and an Id(D–OS) pathway. However, compared with reactions of [Co(NH3)5(OH2)]3+ the reactions of cobalamin with various entering ligands are far less discriminating, which seems to suggest the Id and D mechanism, respectively, for the two complexes (see Chapter 4).

1.4  Experimental Evidence for Mechanisms From the foregoing discussion, it is amply clear that a reaction mechanism can hardly be “proven”; rather, one can merely suggest the most plausible mechanism that is in accord with most of the evidence, with the least evidence against. The first task in elucidating the mechanism is to establish the intimate mechanism, viz., if the mechanism is dissociative or associative; when referring to an intimate mechanism these are denoted d and a, respectively, and D, A, Id and Ia are used to denote the stoichiometric mechanisms (see Section 1.3). Most of the earlier evidence was based on data on the reactions of inert complexes of metal ions, such as those of Co(iii), Cr(iii), Rh(iii), Ir(iii) and Pt(ii), and also of Ni(ii), which is in the borderline for inert and labile systems. This distinction is based on Taube's original proposal that complexes for which t½ is greater than 1 min are inert. Most of the earlier data are on inert systems that are amenable to studies by conventional techniques, and UV–visible spectroscopic methods have largely been used for the complexes of transition metal ions, in addition to techniques such as conductometry, potentiometry, polarography, polarimetry and even titrimetry (for very slow reactions), etc. However, the fast reactions of labile systems require the use of special techniques, depending on the range of t½, such as stopped-­flow spectrophotometry, NMR spectroscopy and T-­jump and P-­jump techniques (see Chapter 2). Figure 1.1 shows the energy versus reaction coordinate profiles for A and D processes. In an A process, both leaving and entering groups are participants in the transition state and here the entering group plays an important role in determining the activation energy of the process. In a D process, the entering group does not participate in the transition state for the reaction and in the intermediate. Hence the absence of a role of the entering group in the activation energy of the process is a characteristic feature of

Reactions of Metal Complexes

17

Figure 1.1  Energy  profiles for reaction A and D processes [free energy (or enthalpy)

versus reaction coordinate diagram]. The reaction coordinate is a parameter representing the continuous stages in the transformation of the reactants to the products. T.S. = transition state; ΔG = free energy of reaction (ΔG = ΔH − TΔS); ΔG‡ = free energy of activation (ΔG‡ = ΔH‡ − TΔS‡); ΔH, ΔH‡ = enthalpy changes; ΔS, ΔS‡ = entropy changes; ΔH − ΔH‡ = ΔHT (enthalpy of transition); ΔG − ΔG‡ = ΔGT (free-­energy of transition).

a D process. The transition state in an Ia process will display substantial bonding to both the entering and leaving groups, and naturally, therefore, the entering group will play an important role in determining the energy of the transition state and hence the activation energy of the process. In the transition state for an Id reaction, there is only weak bonding to both the entering and leaving groups and hence the effect of the entering group on the rate will be small. For reaction energy profiles for stoichiometric mechanisms for ligand replacement reactions of square-­planar complexes and for inner-­sphere mechanisms of electron-­transfer reactions, see Chapters 3 and 7, respectively; for other reaction systems, see ref. 7 (1st edn, pp. 83 and 287). The valleys in Figure 1.1 represent the existence of intermediates, but this need not always be true, and the change from the transition state to the product may be a smooth one in many cases (i.e. without passing through a reactive intermediate of sufficient stability). In the assignment of a plausible reaction mechanism, detailed knowledge of the stoichiometry and the exact nature of the product and the reaction intermediates provide very useful information. In a typical investigation, the reaction products and the stoichiometry have to be determined first. The stoichiometry should be determined with each of the reactants in excess. A definite stoichiometry may suggest a single pathway, but a variable stoichiometry with variable products certainly

Chapter 1

18

suggests more than one pathway for the reaction. Thus, in the oxidation of HSO3− by Fe(iii), the products formed are SO42− and S2O62− and the overall reactions are   

HSO3− + 2Fe(iii) + H2O → 2Fe(ii) + 3H+ + SO42−

(1.38)

  

and   



2HSO3− + 2Fe(iii) → 2Fe(ii) + 2H+ + S2O62−

(1.39)

  

In the presence of large excess of Fe(iii), the stoichiometry corresponds to eqn (1.38) and the mechanism is Fe3+ + HSO3− ⇌ Fe(HSO3)2+ Fe(HSO3)2+ → Fe2+ + HSO3 Fe3+ + HSO3* + H2O → Fe2+ + SO42− + 3H+ However, in the absence of a large excess of Fe(iii), transformation of HSO3·occurs as follows: 2HSO3 → 2H+ + S2O62− hence the stoichiometry corresponds to eqn (1.39). The situation is similar with all other one-­electron oxidants, while two-­electron oxidants do not yield S2O62− and form only SO42−: HSO3− → H+ + SO3 + 2e− SO3 + H2O → SO42− + 2H+

1.5  M  ethods for the Characterization of Reactive Intermediates The existence of unstable intermediates, which never accumulate in any appreciable concentration, cannot be easily demonstrated by any simple analytical or isolation method. However, the existence of intermediates may be inferred by indirect methods. It is sometimes possible to react (trap or scavenge) all or part of the reactive intermediates with an added substrate, thereby allowing indirect observation of the intermediate. For this method to be successful, the rate of reaction of the intermediate with the added substrate must be faster than or at least comparable to the other reactions of the intermediate. Furthermore, the added substrate must be unreactive to the reactants and the products, at least on the time scale required for the experiment. Hence the method has not had widespread application, but has proved successful in some cases. In

Reactions of Metal Complexes

19

the reduction of various oxidants by Sn(ii) there is evidence for Sn(iii) as a reactive intermediate.30 Thus, in the oxidation of Sn(ii) by Ce(iv): Sn(ii) + 2Ce(iv) → Sn(iv) + 2Ce(iii) in the presence of [Co(C2O4)3]3− reduction of the Co(iii) complex to Co(ii) has been observed. Since Ce(iv), Ce(iii), Sn(iv) and Sn(ii) do not react with [Co(C2O4)3]3− on the time scale of interest, the observation suggests the generation of a stronger reductant as an intermediate, which is logically believed to be Sn(iii). Similar behaviour has been observed in some other oxidations of Sn(ii), all of which are thus assumed to involve Sn(iii) as an intermediate: Sn(ii) + Ce(iv) → Sn(iii) + Ce(iii) Sn(iii) + Ce(iv) → Sn(iv) + Ce(iii) Evidence for Pt(iii) in the [PtCl6]2−–Cu(i) reaction31 and of Au(ii) in the [AuCl4]−–Fe(ii) system32 has been furnished. In the reduction of [Pt(Cl) (NH3)5]3+ and [Pt(OH)(NH3)5]3+ by [Cr(OH2)6]2+, there is evidence for Cr(iv) as an intermediate.33 In the solvolysis of halido complexes also, the situation is favourable for the experimental demonstration of the existence of a reactive intermediate:   



LN-­1MX + S → LN-­1MS+ + X−

(1.40)

  

where S is a solvent molecule (such as H2O), which may function as a ligand. If the transition state for this reaction is reached by a considerable extent of rupture of the M–X bond, there are two possibilities: (i) the substitution may proceed smoothly in a concerted single step, or (ii) an intermediate of reduced coordination number (LN-­1M+) may be formed prior to the entry of the solvent molecule. In the latter case, introduction of the free X− into the solution would lead to a reduction in the observed rate of solvolysis and this would serve as evidence for the intermediate (LN-­1M+), since the free X− would enter into a competition with S for this intermediate, to retard the rate of solvolysis by transforming some of the intermediate back to the initial complex. To apply this criterion, however, care must be taken to exclude other possible effects of X−, such as the effect of ionic strength on the rate. Similarly, other competitors, Y, for the intermediate might lead to a change in the product distribution, without affecting the rate of disappearance of the initial complex:   



L N-1 MX  L N-1 M   X  L N-1 M   Y  L N-1 MY  

  

L N-1 M  S  L N-1 MS

(1.41)



where Y stands for other ligands present. For a series of different Y (such as Y′, Y″, Y‴, …, etc.), the relative amounts of LN-­1MY′, LN-­1MY″, LN-­1MY‴, …, etc.

Chapter 1

20

formed would depend on the relative affinities of Y′, Y″, Y‴, …, etc., for the intermediate LN-­1M+. A still more subtle way to observe selective reactivity of an intermediate follows from the idea of Winstein and co-­workers.34 Consider the effects of introducing the non-­coordinating anion perchlorate to the LN-­1MX solvolysis reaction system in a solvent that suppresses ionic dissociation. If the intermediate is formed in this solvent medium, it may remain for some time as an ion pair, in which the leaving group will be able to re-­enter the inner sphere of the complex, and there may not be any overall reaction. Addition of excess ClO4− can assist in the replacement of the leaving group by replacing X− in the ion pair by ClO4−, which thus preserves the intermediate for solvolysis. The net result is a special acceleration of solvolysis on the addition of perchlorate, which is not produced by any coordinating anion:   





(1.42)

  

Finally, if LN-­1MX has properties of geometric or optical isomerism and LN-­1M+ is a highly symmetrical species, it may be possible to infer the existence of the intermediate from an isomerization process, that occurs as fast as (or even faster than) the solvolysis reaction. Here again, formation of an ion pair and re-­entry of the leaving group are involved, but the isomerization is explained by assuming that the ion pair survives at least long enough for LN-­1M+ to rotate with respect to its immediate environment, so that re-­entry of X− may occur from a new direction to form the isomer. All these different tests for identification of an intermediate would succeed only when the intermediate survives long enough for rearrangement of its outer coordination sphere. A competitor for the intermediate must find its way into the outer coordination sphere of the intermediate. In the minimal case, the intermediate must survive long enough for rotation with respect to its outer coordination sphere. There must, therefore, be significant activation energy for reaction of the intermediate if it is to survive long enough to be detected. On a time scale, the intermediate must survive longer than a minimum near 10−10 s.

Reactions of Metal Complexes

21

1.6  Order of Reaction and Reaction Mechanisms It may be mentioned in this connection that, in the case of a reaction of the type shown in eqn (1.2a), when studied in solution in a good coordinating solvent such as water, the independence of the rate on the concentration of Y does not conclusively prove an SN1 (or D) mechanism, since the reaction may involve a bimolecular (A process) replacement of X in the complex by a molecule of the solvent, S, followed by fast replacement of S in the intermediate product by Y (shown below, with charges on species omitted):   

  



Slow ( A )

L N-1 MX  S  L N-1 MS  X

(1.43)

Fast L N-1 MS  Y  L N-1 MY  S

(1.44)

  

This will obviously be a pseudo-­first-­order reaction, despite being a bimolecular process in the rate-­determining step, owing to the concentration of S in the reaction system being very high and hence remaining virtually unchanged during the reaction. Conversely, the first-­order dependence of the rate of the reaction shown in eqn (1.2a) on the concentration of Y does not necessarily mean an SN2 (A) reaction, since the substrate LN-­1MX and Y may enter into a fast pre-­equilibrium reaction to form an intermediate, followed by a purely dissociative (D) transformation of the intermediate in the slow (rate-­determining) step to form the product. Thus, if LN-­1MX is a cation and Y an anion, an outer-­sphere association leading to ion-­pair formation may occur, followed by transformation of the ion pair into the product:   





(1.45)

  

The slow step shown in eqn (1.45) is unimolecular involving the ion pair, but may involve a dissociative (D), associative (A) or interchange (I) mechanism (and hence D-­IP, A-­IP or I-­IP process), irrespective of which the overall change would be second order, first order in Y, because of the pre-­equilibrium. In many other systems, where ion-­pair formation is precluded (as in cases where at least one of the reactants in uncharged, or in reactions of anionic complexes with anionic ligands), outer-­sphere association may still occur due to ion–dipole interactions, hydrogen bonding interactions, stacking interactions (see Chapter 5), etc. The anionic ligand H2Y2− (where H4Y = ethylenediaminetetraacetic acid) is a zwitterion containing >NH+, which enters into electrostatic interactions with anionic complexes, leading to significant outer-­sphere association in such cases.35 In the reaction of

Chapter 1

22 −

2−

3−

cis-­[Cr(ox)2(H2O)2] with oxalate (ox ) to form [Cr(ox)3] , hydrogen bonding interactions lead to significant outer-­sphere association,36 although ion pairing is precluded. Apart from outer-­sphere association, the pre-­equilibrium between the reactants may involve an acid–base reaction in favourable cases. Thus, the hydrolysis of [Cr(N3)](NH3)5]2+ is catalysed both by H+ and OH− ions, due to the formation of the conjugate acid and conjugate base forms of the complex, respectively:37   

  



[Cr(N 3 )(NH3 )5 ]2   H # [Cr(HN 3 )(NH3 )5 ]3

(1.46)

[Cr(N 3 )(NH3 )5 ]2   OH # [Cr(N 3 )(NH3 )4 (NH2 )]  H2 O 

(1.47)

( Conjugate acid, CA )

(Conjugate base, CB)

  

The acid-­catalysed reactions of many complexes involve the conjugate acid mechanism, SN1CA, i.e. D-­CA, or SN2CA, i.e. A-­CA, depending on the role of water, while the base hydrolysis of many complexes having NH3 or RNH2 as ligands involves the conjugate base mechanism, SN1CB, i.e. D-­CB, or SN2CB, i.e. A-­CB, depending on the role of water, which determines whether the transformation of the conjugate acid and conjugate base into the product is dissociative or associative in nature. For any system where the species undergoing the change in the rate-­ determining step is generated in a pre-­equilibrium, the general rate expression may be derived as follows:   

  

K

S  R #

Substrate

Reagent

I

Intermediate

Slow, k

 P Product

(1.48)

The rate of formation of P is given by   



Rate 

d[P]  k[I] dt

(1.49)

[I] [S][R]

(1.50)

  



K

  

where K is the equilibrium constant of the pre-­equilibrium step of eqn (1.48). For material balance,   



TS = [S] + [I]

(1.51)

  

where TS is the total concentration of the substrate in the system. Hence, when TR ≫ TS, [R] ≈ TR and, combining eqn (1.50) and (1.51), we obtain   

  



  

K

[I] {TS  [I]}TR

(1.52)

KTSTR 1  KTR

(1.53)

[I] 

Reactions of Metal Complexes

23

Now combining eqn (1.49) and (1.53), we obtain   

  

Rate 

d[P] KkTSTR  dt 1  KTR

(1.54)

Two limiting situations are to be considered:    (i) I f the value of K is quite small, so that at all feasible concentrations of R in the solution KTR ≪ 1, eqn (1.54) transforms into   

Rate ≈ (KkTSTR) = k′TSTR

(1.55a)

  

i.e. under this condition the reaction obeys the rate law for a typical second-­ order reaction (first order with respect to each of the reactants), although the change in the slow (rate-­determining) step is truly unimolecular. (ii) I f K is sufficiently large, so that even at rather low concentrations of R in the solution KTR ≫ 1, eqn (1.54) reduces to the following form:   

Rate ≈ kTS

(1.55b)

  

Hence, under this condition, the reaction obeys the rate law for a typical first-­ order reaction which is independent of the reagent (R) concentration.    In between the aforesaid limits, there will be a situation where at relatively low concentrations of R the experimentally observed pseudo-­first-­order rate constant, kobs = kKTR/(1 + KTR), would increase linearly with the concentration of R (up to point ‘a’ in Figure 1.2) and will tend to attain a limiting value at high concentrations of R (beyond point ‘b’ in Figure 1.2). Various types of reactions (such as ligand replacement and redox reactions) may show such behaviour as expressed by eqn (1.54). A typical example is the

Figure 1.2  Variation  of rate constant (kobs) with the reagent concentration TR in a system involving a pre-­equilibrium forming the reactive species (see text).

Chapter 1

24 2−

38

39

oxidation of [Co(edta)] by periodate, and another is the base hydrolysis of [Cr(NCS)(NH3)5]2+. The reaction shown in eqn (1.48) should be expressed more generally as follows (since the reversible formation of I may not be instantaneous):   

k1



k3 S  R # I   P

(1.56a)

d[P]  k3 [I] dt

(1.56b)

k2

  

This, as before, leads to   

  

Applying the steady state principle:   



  

Hence,

d[I]  k1 [S][R]  ( k2  k3  )[I] 0 dt

  

  

[I] 

k1 [S][R] k2  k3

(1.56c)

(1.56d)

From eqn (1.56b) and (1.56d), we obtain   

  

d[P] k1 k3 [S][R]  dt k2  k3

(1.56e)

Hence, if k2 ≫ k3, then eqn (1.56e) becomes   

  

d[P] k1 k3  [S][R ]  Kk3 [S][R ] dt k2

(1.56f)

However, if [S] ≫ [I], and R is present in large excess, i.e. TR ≫ TS, then [S] ≈ TS and [R] ≈ TR and eqn (1.56f) changes to eqn (1.56g) [cf. eqn (1.55a)]:   

  

d[P]  Kk k kfTR 3TSTR f TSTR dt

(1.56g)

where K is the equilibrium constant for the formation of I from S and R [eqn (1.56a)]. A formation reaction of the type   



LN-­1M(OH2) + Y → LN-­1MY + H2O

(1.57)

  

often obeys the second-­order rate law (first order with respect to each of the reacting species). However, from this information alone it cannot be concluded unequivocally that the reaction is bimolecular. The same second-­ order kinetics would be observed for a unimolecular process of the following type:   

Reactions of Metal Complexes



25



(1.58)

  

The system is similar to that of eqn (1.17) and the reaction is unimolecular since the first step is rate determining. However, it will have a first-­order rate law only if the conversion of the intermediate to the product is enormously favourable compared with the back reaction [see eqn (1.21)]. Since the concentration of the solvent in the solution is exceedingly high compared with that of the reactants, the back reaction might be more favourable than the second step, particularly if Y is a poor nucleophile, and in this case a second-­ order rate law will be observed [first order in each of the reactants, see eqn (1.23)]. The foregoing discussion clearly demonstrates that the observed order of a reaction hardly informs us unambiguously about the molecularity of the reaction and hence also the mechanism. Various other criteria need to be applied to arrive at a plausible mechanism; the one proposed should explain adequately all the observed facts (or at least most of them). Using simple electrostatic arguments, it is possible to speculate on the influence of the changes in sizes and charges of the central metal ion, the entering group, the group being displaced and the non-­replaceable groups (spectator ligands) on the rate of the three general processes, viz. D, I and A. The summary of the predictions given in Table 1.2 is based on broad generalizations and considerations that the important features of D and A mechanisms are bond breaking and bond formation, respectively, whereas in an interchange (I) process bond breaking and bond formation are of comparable importance (concerted process). Table 1.2  Effects  of sizes and charges on rates of reactions involving dissociative (D) interchange (I) and associative (A) mechanisms.1

Effect on rate Change

D process

I process

A process

Increase in positive charge of central atom Increase in size of central atom Increase in negative charge of entering group Increase in size of entering group Increase in negative charge of leaving group Increase in size of leaving group Increase in negative charge of non-­ replaceable groups

Decrease Increase No effect

Little change Increase Increase

Increase Increase Increase

No effect Decrease Increase Increase

Decrease Decrease Little change Little change

Decrease Decrease Decrease Decrease

Chapter 1

26

Figure 1.3  Comparison  of the rates of water exchange for various metal aqua M  OH2 complexes. The figures along the x-­axis are log kex ; kex values (s−1) at 25 °C. Owing to the high instability of [Co(OH2)6]3+, its kex value is uncertain.

The effect of the charge and size of the central metal ion is demonstrated elegantly in the results of studies on the fast water exchange rates of aqua-­ metal ions (see Figure 1.3).40 This shows that the rate of exchange decreases in the sequence Na+ > Mg2+ > Al3+, with increasing cationic charge. When the charge of the metal ion remains the same, the rate decreases with decrease in size of the cation in the following sequences: Cs+ > Rb+ > K+ > Na+ > Li+ and Ba2+ > Sr2+ > Mg2+ > Be2+. These observations are consistent with a dissociative mechanism.

1.7  Lability of Complexes The stability of a complex in the thermodynamic sense is expressed in terms of the formation constant βn, which is the equilibrium constant for the formation of the complex from the constituents in solution:   



  

M + nL ⇌ MLn

(1.59)

for which   



  

n 

[ML n ] [M][L]n

(1.60)

Reactions of Metal Complexes

27

The higher the value of βn, the greater is the stability of the complex. This thermodynamic stability of the complex, however, is not the only factor influencing the observed behaviour of a complex: the rate of reaction of the complex also plays a vital role. Although it is often true that thermodynamically stable systems react slowly, whereas unstable systems react rapidly, there is no absolute correlation between the stability and the rate of reaction of the species, and complexes are no exception in this respect. Thus, [Ni(CN)4]2– (β4 ≈ 1022), [Hg(CN)4]2− (β4 ≈ 1042) and [Fe(CN)6]4− (β6 ≈ 1037) are all very stable and yet the first two complexes exchange the bound cyanide with free cyanide in solution (studied with isotopically labelled *CN−) extremely rapidly (almost instantaneously), whereas the corresponding exchange in the case of [Fe(CN)6]4− is very slow, although it is much less stable than [Hg(CN)4]2−. Again, the complex ion [Co(NH3)6]3+ is thermodynamically very unstable in an acidic medium, having a considerable tendency to dissociate:   

[Co(NH3)6]3+ + 6H3O+ ⇌ [Co(OH2)6]3+ + 6NH4+



(1.61)

  

for which the equilibrium constant is of the order of 1023, and yet [Co(NH3)6]3+ persists unchanged for weeks in such a medium. In fact, the salt [Co(NH3)6] Cl3 can be crystallized from hot dilute hydrochloric acid solution without any decomposition. This is because the rate of the forward reaction of eqn (1.61) is extremely slow under ordinary conditions. Thus, the stability of a complex in the ordinary sense (i.e. its failure to undergo any observable change under a particular set of conditions) may be due either to its thermodynamic stability or to its kinetic nature (extremely slow reaction), for which latter the term inert was suggested by Taube;41 the term labile is used for a rapidly reacting system. Thus, in the examples cited, [Co(NH3)6]3+ is unstable but inert, [Ni(CN)4]2− and [Hg(CN)4]2− are stable but labile and [Fe(CN)6]4− is stable and inert; inert with respect to the mentioned reactions. Also interesting in this connection is the observation reported by Grinberg and Nikol'skoya42 that the rate of exchange in the system [PtX4]2−– X−, studied using X− labelled with a radioactive isotope of X, increases sharply and paradoxically in the same order as that of increasing stability of the complex: Complex 2−

t½/min‡ β4

[PtCl4]

[PtBr4]

∼840 ∼1017

∼6 ∼1020

2−

[PtI4]2−

[Pt(CN)4]2−

∼4 ∼1030

∼1 ∼1041

‡

t½ is the half-­time for exchange and this is related to the ligand exchange rate constant. From eqn (1.1b), it follows that k = 0.69/t½.

Again, in a series of complexes of Co(iii) and Cr(iii) with some biguanides, it has been observed that the rate of dissociation in acidic medium increases in the order of increasing stability of the complex, in the sequence phenylbiguanide < n-­hexylbiguanide < biguanide.43

Chapter 1

28

Consideration of the general kinetic behaviour of the octahedral complexes (the most common type) of elements of the first transition series indicates that the complexes of an ion having the d3 configuration and the low-­spin complexes of d4, d5, d6 and d7 ions are generally inert. All the complexes of d0, d1, d2, d9 and d10 are labile, as are also the high-­spin complexes of d4, d5, d6 and d7. A large number of complexes of Ni(ii) d8 ion are labile and a few are also fairly inert, and the lability of complexes of Ni(ii) d8 is much less than that of analogous d7 (high-­spin) Co(ii) and d9 Cu(ii) systems. From consideration of such general observations, Taube41 suggested that most of the data on the lability of the coordination compounds can be explained on the basis of their electronic configuration, as proposed in the valence bond theory of Pauling. In general, all the complexes of the outer orbital type, and those of the inner orbital type having at least one vacant d orbital in the penultimate valence shell of the central ion, are labile. Similar considerations also apply to complexes of elements of the second and the third transition series. This seems to suggest the operation of the displacement SN2 mechanism for the reactions of inner orbital complexes, in which a seventh group, the reagent, must be bonded in the coordination sphere of the central ion. Accordingly, such a process will be facilitated if any empty d orbital of lower energy is available for bond formation. If such an orbital is not available, then the seventh group may be added by using an outer d orbital of higher energy, and in such a case the reaction is consequently slower, because a higher activation energy is required for the process. For outer orbital complexes, which are generally labile, the lability decreases with increasing charge on the central ion, more precisely the oxidation state, as seen in the observed sequence of lability AlF63− > SiF62− ≫ PF6− > SF6. Hence outer orbital complexes can also exhibit inertness, as is the case particularly for SF6, and PF6− is also somewhat similar owing to the high charge and small size of the central “ion” (formally S6+ and P5+ due to the high ionic character of the bonds with the strongly electronegative F in these covalent species). The effect of ion size and charge on lability is shown in Table 1.3. The bonding in [Be(OH2)4]2+ uses sp3 of Be2+, and in all the other cases of [M(OH2)6]m+ (m = 2, 3) the bonding is of outer orbital type involving sp3d2 hybrid orbitals of the central “ion”. Since the outer d orbitals are less stable (higher in energy) than the inner d orbitals, it follows that bonds involving Table 1.3  Effect  of ion size and charge on lability in the case of water exchange rates for the aqua ions of M2+ and M3+ metal ions in aqueous solution (kex values at 25 °C). Mm+ Parameter

Be2+

Mg2+

Ca2+

Ba2+

Al3+

Ga3+

In3+

La3+

rMm /Å kex/s−1

0.41 ∼102

0.86 ∼105

1.14 ∼108

1.49 ∼109

0.68 ∼10

0.76 ∼103

0.94 ∼106

1.17 ∼108

Reactions of Metal Complexes

29

them are less stable. Therefore, it is likely that such complexes will react by a dissociation (SN1) mechanism through the intermediate formation of a complex of coordination number five. However, as the formal positive charge on the central atom increases, the bond strength also increases so that with a higher value of charge the complex eventually becomes inert. However, it should be mentioned that the assumption implied in the valence bond theory that the bonds in the outer orbital complexes are weaker is not really justified. The fact that the hydration energies for the formation of outer orbital hexaaqua complexes of the ions d6, d7, d8, d9 and d10 are much higher than those for the formation of inner orbital hexaaqua complexes of d0, d1, d2 and d3 ions44 (see Figure 1.4 and the corresponding diagram for [M(OH2)6]3+ in ref. 44) suggest that the bonds in the outer orbital complexes are not necessarily weaker. (It must be realized that valence bond theory suggests that an inner orbital complex will always be formed if a sufficient number of vacant inner d orbitals are available, as is true for the formation of octahedral complexes of ions having the configuration d0, d1, d2 and d3.) This therefore invalidates the explanation given above for the lability of the outer orbital complexes. The other serious limitation of the valence bond approach is that it can make only qualitative predictions regarding the lability of complexes without being able to make even semiquantitative predictions. Hence, although the valence bond method can account for the inertness of the inner orbital complexes of d3, d4, d5 and d6 ions, it cannot give any explanation for the order of reactivities of these inert complexes and the observed differences in the energies of activation in a series of similar complexes. In these respects, as will be shown below, the ligand field theory has a decided advantage.

Figure 1.4  Enthalpy  of hydration of M2+(g) forming [M(OH2)6]2+ in aqueous solution plotted against n of the dn configuration of the M2+ ion. For this and the corresponding plot for M3+ (see ref. 44), the ∆H values for Mn+(g) + 6H2O(g) → [M(OH2)6]n+(g) (n = 2, 3) show similar trends.45

Chapter 1

30

1.7.1  Ligand Field Theory46 This is a sophisticated form of the electrostatic crystal field theory. The basic postulate of the crystal field theory is that the five d orbitals which are degenerate (equal in energy), in the free (gaseous) ion of the metal become differentiated, i.e. split up, into levels of different energy, under the influence of the electrostatic field of the ligands. In an octahedral complex, the central ion (or atom) is surrounded by the six ligands that are arranged at the corners of a regular octahedron with the metal ion (or atom) at its centre (Figure 1.5). The ligands, which are negative ions or polar molecules, are thus arranged with their negative ends pointed towards the central ion along the x-­, y-­and z-­axes. The electrostatic repulsion between these ligands and the electrons of the central ion (or atom) will be much greater if the electrons are in the d x2  y2 and d z2 orbitals, which point directly towards the ligands, than if they are in the dxy, dxz and dyz orbitals, which lie in between the metal–ligand bond axes and thus maintain maximum separation from the ligands (Figure 1.6a).§ Thus, the five degenerate d orbitals of the free ion, which is the centre of the coordination complex, are split into two groups, the energy of the d x2  y2 and d z2 orbitals being increased far more by the presence of the ligands in the octahedral configuration. It should be noted that because of the negative charge on the ligands, the energy of all the d orbitals in an octahedral field should be represented as in Figure 1.7, i.e. higher than in the field-­free gaseous Mm+. However, when the ligands are added, the total energy of the system decreases. Hence it is customary to represent the splitting of the d orbitals

Figure 1.5  Arrangement  of six ligands around a central metal ion in an octahedral complex.

§

This conclusion is strictly valid only if we consider the ligands as point charges (or point dipoles). However, if we are to represent the ligands in a realistic manner, i.e. as finite spheres of negative charge with a positive charge located at the centre of the sphere, it can be shown that the stability will still be slightly in favour of orbitals having lobes pointing in between the metal–ligand bond axes,48 hence this will not alter the final conclusions arrived at assuming the ligands as only point charges.

Reactions of Metal Complexes

31

Figure 1.6  Complete  set of d orbitals in the metal in (a) an octahedral complex,

where black circles are ligands occupying the centres of faces of a cube (the torus of the orbital in the x–y plane has been omitted for clarity), and (b) a tetrahedral complex, the ligands (black circles) occupying alternate corners of a cube. The metal atom is at the centre of the cube in both cases.

Figure 1.7  Diagrammatic  representation of loss in degeneracy of the five d orbitals in a ligand field of octahedral symmetry 10Dq = Δ. As usual, the splitting obeys the centre of gravity rule and as such the doubly degenerate eg is raised in energy by 6Dq whereas the triply degenerate t2g is lowered in energy by 4Dq.

32

Chapter 1

in the manner shown in Figure 1.8, where the t2g orbitals are represented as having lower energy than the mean energy of the d orbitals of the metal ion in the complex. On similar theoretical considerations, it is possible to arrive at the types of splitting observed in complexes of different stereochemistry, and some common examples are illustrated in Figure 1.8. The splitting in a cubic field is qualitatively similar to that in a tetrahedral field but ΔK ≈ 2ΔT ≈ 0.9ΔO. The crystal field splitting for a few other configurations (linear, trigonal planar, etc.) have also been discussed, for which the original references should be consulted.47 The tetragonal configuration included in Figure 1.8 is a distorted octahedron in which the two groups along the z-­axis are at a greater distance from the central ion than the other four groups along the x-­ and y-­axes; the same ordering is valid for the square-­pyramidal structure also. However, Bosnich et al.49 proposed the ordering to interpret the spectra of low-­spin square-­ pyramidal [Ni(tetars)X]+. The ordering of the d orbitals in a square-­planar configuration is somewhat uncertain, but the ordering shown in Figure 1.8 has been established for the 5d8 configuration from both spectroscopic data and theoretical calculations.50 The energy separation between the split d orbitals can be estimated from spectroscopic data, and can also be calculated

Figure 1.8  Crystal  field splitting of the d orbitals of a central metal ion in complexes of different structures. (a) Energy level of the five degenerate d orbitals of the metal in a hypothetical field of spherical symmetry. This represents a hypothetical case where the negative charge of the ligands is assumed to be uniformly smeared over a spherical surface around the central metal. The energy of these degenerate d orbitals represents the mean energy of the d orbitals in the complex, (b) octahedral, (c) tetrahedral, (d) tetragonal (octahedron elongated along the z-­axis) and similarly for square pyramidal, (e) square planar, (f) trigonal bipyramidal and (g) cubic. In a strong square-­planar field d z 2 may be even lower than dxz and dyz, and according to some authors this is likely for Pt(ii) complexes (see text), ∆K ≈ (8/9)∆O.

Reactions of Metal Complexes

33 51

by a rigorous theoretical approach. The low-­intensity absorption bands (εmax < 100) in the visible spectrum of a complex result from transitions of electron between the various energy levels which the d electron can occupy. Hence it is possible to estimate the energy difference between the different d orbitals in electrostatic fields of different geometry from spectroscopic data on the d–d absorption bands of complexes. For octahedral and tetrahedral structures, where the splitting of the d orbitals is of a rather simple nature, the Δ or 10Dq value for the octahedral field, Δoct, is much larger than for the tetrahedral field, Δtet. From theoretical calculations, it has been shown52 that when the central ion, the ligands and the metal–ligand bond distances are the same in both the octahedral and tetrahedral structures, then ΔT ≈ (4/9)ΔO. In other words, all the factors being equal, splitting of the d orbitals in the tetrahedral structure will be less than half the splitting in the octahedral structure. In [Co(NH3)6]2+ Δoct = 10.1kK and in [Co(NCS)4]2− Δtet = 4.9kK;53 both have Co–N bonds, but bond distances are not the same. For tetrahedral [CoCl4]2−, Dq = 312 cm−1, but for Co(ii) in an octahedral field of six Cl−, as in solid CoCl2, Dq = 690 cm−1. The extent of splitting in a square-­planar structure would depend on the nature of the central metal and the ligands. Semiquantitative calculations with parameters appropriate for square-­planar complexes of Co(ii), Ni(ii) and Cu(ii) lead to Δ1 ≈ ΔO, Δ2 ≈ 2/3ΔO and Δ3 ≈ 1/12ΔO, i.e. Δ1 + Δ2 + Δ3 ≈ 1.75ΔO (see data in Table 1.2). Experimental (spectroscopic) results indicate54 that for [Ni(CN)4]2−, Δ2 ≈ 2/5Δ1 and Δ3 ≈ 1/38Δ1, and for Pd(ii) and Pt(ii), in [MCl4]2−, Δ1 + Δ2 + Δ3 ≈ 1.3ΔO. Thus, the three spin-­allowed absorption bands in the spectrum of [PtCl4]2− are due to either of the following sets of transitions:    (I)53 v−1: 1A1g → 1A2g; v−2: 1A1g → 1Eg; v−3: 1A1g → 1B1g (II)58   v−1: 1A1g → 1A2g; v−2: 1A1g → 1B1g; v−3: 1A1g → 1Eg    For clarity, the electronic configurations of the four singlet states are as follows: 1 1

A1g: eg4a1g2b2g2; 1A2g: eg4a1g2b2g1b1g1 (⇅)

B1g: eg4a1g1b2g2b1g1 (⇅); 1Eg: eg3a1g2b2g2b1g1 (⇅)

Of the three transitions mentioned above, the first is the lowest in energy and its wavenumber corresponds to the energy separation Δ1 (Figure 1.8e), while the wavenumber of the third transition similarly corresponds to (Δ1 + Δ2 + Δ3) (Figure 1.8e) irrespective of the assignments (I) and (II). In Figure 1.8e, it should be noted that in the transformation of the octahedral structure to the square-­planar structure, through the intermediate tetragonally distorted octahedral structure, the orbitals dxy and dx2−y2 are raised in energy to the same extent and therefore their separation remains the same throughout (see also data in Table 1.2). Hence Δ1 = Δ0 (of the octahedral structure). For the complex ion [PtCl4]2−, the theoretically predicted and experimentally

Chapter 1

34

observed values of the wavenumbers of the above-­stated three spin-­allowed transitions are as follows:55 Theoretical v− values: 25 970, 30 160 and 34 200 cm−1 Observed v− values: 26 300, 29 600 and 36 500 cm−1 Based on these, the value of R = (Δ1 + Δ2 +Δ3)/Δ1 = ∑Δ/Δ0 = ν3/ν1 is 1.32 (theoretical) and 1.39 (experimental). In a similar manner, the corresponding R values for [PtBr4]2−, [PdCl4]2− and [PdBr4]2− can be evaluated using reported spectral data55 and are given in Table 1.4. For the trigonal bipyramidal structure ΔTB ≈ 0.98ΔO and for cubic (eight-­ coordinate) ΔK ≈ 2ΔT ≈ 8/9ΔO. Since E = hv = hc(1/λ), where h is Planck's constant (6.626 × 10−27 erg s), v is the frequency, λ is the wavelength of the maximum of the appropriate absorption band and c is the velocity of light (∼3 × 1010 cm s−1), it is customary to express the value of Δ or 10Dq in terms of the value of 1/λ, i.e. the wavenumber (v−), which is expressed in cm−1. Hence E (cal mol−1) = [hcN/(4.2 × 107)] × 1/λ = 2.86(1/λ) = 2.86 v−. Therefore, by multiplying the value of the wavenumber in cm−1 by 2.86, one obtains the value of the energy in calories per gram ion (the factor N in the above relation is Avogadro's number, 6.022 × 1023 mol−1 and 4.2 × 107 erg ≈ 1 cal). Hence 1000 cm−1 (i.e. 1kK) is 2.86 kcal mol−1, i.e. ∼2.9 kcal mol−1 ≈ 12 kJ mol−1 (1 kcal = 4.184 kJ). As usual, the splitting shown in Figure 1.7 obeys the centre of gravity rule and as such the doubly degenerate eg is raised in energy by 6Dq while the triply degenerate t2g is lowered in energy by 4Dq so that when all the five d orbitals are singly or doubly occupied by electrons there is no net change in energy. The notations eg and t2g (and correspondingly e and t2 in a tetrahedral field, see Figure 1.8) were introduced by Mulliken,56 while Van Vleck57 proposed for these dγ and dε and Bethe58 proposed γ3 and γ5, respectively. Apart from the ordering of the d orbitals in a square-­planar field shown50 in Figure 1.8, it has been proposed that in a strong square-­planar field d z2 may be even lower than dxz and dyz, and according to some authors55,59a,b this is likely even for [PtCl4]2−, and reasons for this have been offered.59c,d For the ligand field splitting parameter, several different symbols have been proposed, the two most common being Δ60 and 10Dq.61 Like the d orbitals, the f orbitals are also expected to be split in a ligand field. Thus, in an octahedral field, the f orbitals are expected to be split into three sets. However, in the case of the lanthanides, the 4f orbitals are well Table 1.4  R  values for [PtBr4]2−, [PdCl4]2− and [PdBr4]2−. Complex, [MX4]2−

R value

M

X

Theoretical

Experimental

Pt Pt Pd Pd

Cl Br Cl Br

1.32 1.27 1.27 1.24

1.39 1.25 1.33 1.34

Reactions of Metal Complexes

35

buried in the core and well shielded by the outer electrons, and hence are little affected by the ligands. Therefore, the splitting of the 4f orbitals is small enough not to be of any significance. However, this is not so for the actinides, where the 5f orbitals are well diffused, but not so well shielded as in the case of the 4f orbitals in the lanthanides, but are also not so well exposed as are the d orbitals of transition metals; as a result, the ligand field effect on the 5f orbitals in the actinides is perceptible compared with that in the lanthanides, and more so for the higher valent actinide ions (but less so than for transition metal ions) since the ligand field splitting increases with increasing positive charge on the central metal ion. In the fundamental postulates of the crystal field theory, as was developed originally and which has been discussed above, the interaction between the metal ion and the ligands was believed to be purely electrostatic in nature. However, such an assumption can never be true for any real combination between the metal and the ligands in a complex, and there is already sufficient experimental evidence to indicate that with the closeness of approach of the metal ion and ligands implied in the formation of a complex, there is some extent of overlap of the metal and ligand orbitals, leading to an interaction other than a purely electrostatic interaction. A certain degree of covalency of the bonds between the metal and the ligands in a coordination complex is thereby indicated. This, however, in no way affects the fundamental approach of the theory, with regard to the splitting of the d orbitals. The term ligand field theory is now used for the viewpoint (first introduced by Van Vleck57) that utilizes the basic concepts of the electrostatic crystal field approach, with the necessary modifications to take into account a certain degree of covalency of the bonds between the metal and ligands, causing deviations from the electrostatic idealization, which, however, does not completely invalidate the usefulness of the basic concepts and methods of calculation of the electrostatic crystal field theory.62 It is known that even on the basis of the MO viewpoint, we arrive at a similar result with regard to the splitting of the d orbitals in a ligand field. Owing to the splitting of the d orbitals in a complex, the system gains extra stability due to the rearrangement of the d electrons, preferentially filling the low-­lying d levels. The consequent gain in bonding energy (energy lost by the system) is called the crystal field stabilization energy (ligand field stabilization energy), generally abbreviated as CFSE (LFSE). This can be easily calculated for the octahedral and tetrahedral complexes of different ions in terms of the appropriate Dq values. In the case of a tetrahedral complex, each electron in the t2 level has the energy +4Dq and each electron in the e level has the energy – 6q, and therefore the total energy for the configuration (t2)x(e)y is (4x − 6y) Dq and the LFSE value is expressed as −(4x − 6y)Dq = (6y − 4x)Dq, where Dq is the ligand field splitting parameter appropriate for the tetrahedral field (DqT); DqT ≈ (4/9)DqO. If the ground state of an octahedral complex is (t2g)x(eg)y, its ligand field stabilization energy will similarly be (4x − 6y)Dq, where Dq is the parameter appropriate for the octahedral field (DqO). The appropriate theoretical single

Chapter 1

36

electron energies of the different d orbitals of a central atom for a number of important geometrics found in coordination compounds have been calculated by Ballhausen52 and Jørgensen.63 These values, in terms of the Dq value for the octahedral geometry, are given in Table 1.5 and are useful for calculating the LFSE for different types of complexes. Such values are given in Table 1.6 for complexes of the commonest types of structures. LFSE values for complexes of other structures can also be calculated using the data in Table 1.5. In Table 1.6, the charge on species such as Cr5+, Mn7+, etc. is merely formal. The numbers in parentheses are the number of electrons that must be paired in the d level, on going from the weak field to the strong field case. In all the strong field cases, where an additional n electron pairs result on complex formation, the net stabilization energy would be obtained by subtracting nP from the listed LFSE values, where P denotes the electron pairing energy. The contributory role of LFSE in rates of reactions of transition metal complexes has been qualitatively discussed by Orgel64 and Jørgensen,65 but Basolo and Pearson made a more sophisticated approach (see ref. 1, 2nd edn, Chapter 3). Table 1.5  Energy  (in DqOh) of an electron in different d orbitals in ligand fields of different symmetries.52,53

Coordination number Structure 1 2 3 4 4 5 5 6 6 7 7 8 8 9 10 a

a

— Lineara Trigonal planarb Tetrahedral Square planarb Trigonal bipyramidalc Square pyramidalc Octahedral Trigonal prismatic Pentagonal bipyramidalc Trapezoidal octahedrald Cubic Square antiprismatic Tricapped trigonal prismatic (as in ReH92−) Icosahedralg

d x 2  y2

d z2

dxy

dxz

dyz

−3.14 −6.28 5.46 −2.67 12.28 −0.82 9.14 6.00 −5.84 2.82 8.79 −5.34 −0.89 −0.38

5.14 10.28 −3.21 −2.67 −4.28 7.07 0.86 6.00 0.96 4.93 1.39e −5.34 −5.34 −2.25

−3.14 −6.28 5.46 1.78 2.28 −0.82 −0.86 −4.00 −5.84 2.82 −1.51e 3.56 −0.89 −0.38

0.57 1.14 −3.86 1.78 −5.14 −2.72 −4.57 −4.00 5.36 −5.28 −2.60f 3.56 3.56 1.51

0.57 1.14 −3.86 1.78 −5.14 −2.72 −4.57 −4.00 5.36 −5.28 −6.08f 3.56 3.56 1.51

0.00

0.00

0.00

0.00

0.00

 ond/bonds along z-­axis. B Bonds in the x–y-­plane. c Pyramid base in the x–y-­plane. d Square pyramid (inverted) having the base in the x–y plane with two more groups above the base and bonded through two opposite faces of octahedron (see Figure 1.9c); it can also be viewed as a square pyramid with two more groups below the base, i.e. inverted Figure 1.9c, for which the bond length is ∼25% larger, also known as an octahedral wedge structure. ed  z2 + dxy hybrids. f dxz + dyz hybrids. g No splitting of d orbitals. b

Reactions of Metal Complexes

37

Table 1.6  Ligand  field stabilization energies (LFSE) in units of Dqoct (DqO) for octahedral, tetrahedral and square-­planar complexes of various dn ions in weak and strong fields (WF and SF).a LFSE/Dqoct Free ion configuration 0

d

d1 d2 d3 d4 d5 d6 d7 d8 d9 d10

Octahedral

Ions (typical examples) 2+

3+

WF

3+

4+

Ca , Al , Sc , Ti , 0 V5+, Cr5+, Mn7+ Ti3+, V4+ 4 Ti2+, V3+, Mo4+, Os6+ 8 V2+, Cr3+, Re4+ 12 Cr2+, Mn3+, Ru4+, Os4+ 6 Mn2+, Fe3+, Ru3+, Os3+, 0 Ir4+ 2+ Fe , Co3+, Ni4+, Ru2+, 4 Rh3+, Os2+, Ir3+, Pt4+ Co2+, Ni3+ 8 Ni2+, Cu3+, Pd2+, Pt2+, 12 Au3+ Cu2+, Ag2+ 6 Ni0, Cu+, Ag+, Au+, 0 Zn2+, Cd2+, Hg2+, Ga3+

Tetrahedral

Square planar

SF

WF

SF

WF

SF

0

0

0

0

0

4 8 12 16(1) 20(2)

2.67 5.34 3.56 1.78 0

2.67 5.34 8.01(1) 10.68(2) 8.9(2)

5.14 10.28 14.56 12.28 0

5.14 10.28 15.42(1) 20.56(2) 24.84

24(2) 2.67

7.12(1)

18(1) 5.34 12 3.56

5.34 3.56

10.28 26.84(1) 14.56 24.56

6 0

1.78 0

12.28 12.28 0 0

1.78 0

5.14 29.12(2)

a

 alculated using the values in Table 1.2; see ref. 1, 2nd edn, p. 70, Table 2.5, where the values C for strong field square-­planar d3 and d4 are incorrect.

In weak field cases, all of the d orbitals, in the sequence of increasing energy are singly filled up before any electron pairing takes place. Hence they have the maximum number of unpaired electrons, and this happens in systems where 10Dq is less than the electron pairing energy, P. In strong field systems, electron pairing occurs in the d orbitals in the sequence of increasing energy before any orbital is singly occupied. Hence they have the minimum number of unpaired electrons (or no unpaired electrons, as in the case of strong field octahedral d6), and this happens in systems where 10Dq is larger than P. In cases where P = 10Dq, both configurations will have identical energy, as shown below for octahedral d6 cases: Weak field 6

Configuration for d Energy

4

2

t2g eg –4Dq + P

Strong field t2g6 −24Dq + 3P

When P = 10Dq, both configurations will have the energy 6Dq. For octahedral complexes, it is obvious that both types of configurations are possible only in d4 to d7 cases. The configurations of weak field and strong field octahedral d4, d5 and d7 are given in Table 1.7.

Chapter 1

38

Table 1.7  Configurations  of weak field and strong field octahedral d , d and d7. 4

5

Configuration n

d

4

d d5 d7

Weak field 3

1

t2g eg t2g3eg2 t2g5eg2

Strong field t2g4 t2g5 t2g6eg1

Figure 1.9  Intermediates  in SN1 (a) (square pyramidal) and SN2 (b) (pentagonal bipyramidal) and (c) (trapezoidal octahedral) mechanisms of ligand substitution in octahedral complexes:

ML5X + Y → ML5Y + X In order to explain the kinetic features of reactions of complexes on the basis of the ligand field theory, it is necessary to calculate with the help of the data in Table 1.2 the LFSE values for various d electron configurations in both weak field and strong field cases, for a regular octahedral structure (the reacting species for which a regular octahedral structure may be assumed as an approximation, even if the all the ligands are not the same) and the intermediates expected to be formed in the transition states for the two possible mechanisms SN1¶ and SN2.∥ For the SN1 mechanism, the intermediate of coordination number 5 resulting from the loss of ligand may reasonably be assumed to be one approximating to a square pyramidal structure** (Figure 1.9a), whereas for the SN2 mechanism the structure of the intermediate of coordination number 7 formed by the addition of the incoming ligand Y may be viewed as approximating to a pentagonal bipyramid (see Figure 1.9b) or a trapezoidal octahedral (Figure 1.9c) assuming possible bonding by the entering ligand through the edge or a triangular face of the octahedron. The difference between the LFSE of the original octahedral structure and that of the intermediate formed in the transition state may be considered as the ligand field contribution, ΔEa, to the total activation energy (see ref. 1, 2nd ¶

SN1 = substitution nucleophilic unimolecular (dissociative, D, mechanism).  N2 = substitution nucleophilic bimolecular (associative, A, mechanism). S **Under certain conditions, the square-­pyramidal species expected to be formed by loss of a ligand from the octahedral complex may change into a trigonal bipyramidal species. However, for our present discussion this is not of concern, as the ΔEa value will be determined by the species of coordination number 5 formed initially by loss of a ligand from the original octahedral structure. ∥

Reactions of Metal Complexes

39

edn, Chapter 3) and is often referred to as the ligand field activation energy (LFAE); some prefer to call it the crystal field activation energy (CFAE) for the reaction.†† A large positive value of ΔEa (i.e. a considerable loss in LFSE on forming the transition-­state intermediate) implies a slow rate of reaction by that particular path. A negative value of ΔEa implies that the original octahedral structure is less stable than the transition state, which is just improbable. Hence, in all such cases, it appears probable that distortion is present in the original structure (i.e. it is not purely octahedral), such that it is further stabilized [cf. extra stabilization due to Jahn–Teller distortion in the case of Cu(ii)], so that ΔEa becomes at least zero. Indeed, this situation arises in systems where Jahn–Teller distortion is expected to occur in the original structure. Hence a negative ΔEa value may also be considered as essentially zero. Based on consideration of ΔEa values (Table 1.8), the following orders of labilities are predicted for reactions of octahedral complexes of strong and weak field ligands:    1. SN1 pathway (square pyramidal intermediate):

Table 1.8  Ligand  field activation energies (LFAE) (ΔEa) for reactions of octahedral complexesa.1

Calculated values of ΔEa (in Dqoct) Strong-­field cases S N1

SN2

Weak-­field cases SN1

SN2

Square Pentagonal Trapezoidal Square Pentagonal Trapezoidal System pyramidal bipyramidal octahedral pyramidal bipyramidal octahedral d0 d1 d2 d3 d4 d5 d6 d7 d8 d9 d10

0 −0.57 −1.14 2.00 1.43 0.86 4.00 −1.14 2.00 −3.14 0

0 −1.28 −2.56 4.26 2.98 1.70 8.52 5.34 4.26 1.07 0

0 −2.08 −0.68 1.80 −0.26 1.14 3.63 −0.98 1.80 −2.79 0

0 −0.57 −1.14 2.00 −3.14 0 −0.57 −1.14 2.00 −3.14 0

a

0 −1.28 −2.56 4.26 1.07 0 −1.28 −2.56 4.26 1.07 0

0 −2.08 −0.68 1.80 −2.79 0 −2.08 −0.68 1.80 −2.79 0

In modern terminology, SN1 = D and SN2 = A, meaning dissociative and associative, respectively (see Section 1.3). All negative values are to considered ∼0 (see text).

††

An improved method for estimating LFAE has been reported.66

40

Chapter 1

2. SN2 pathway (pentagonal bipyramidal intermediate):

3. SN2 pathway (trapezoidal octahedral (octahedral wedge) intermediate):

Hence, irrespective of the mechanism, octahedral complexes of d3, d8 and d (low-­spin) are all inert; low-­spin octahedral complexes of d4 and d5 may also be expected to react slowly, whereas all the octahedral complexes of d0, d1, d2, d9 and d10, and possibly also of d7, and the high-­spin octahedral complexes of d4, d5 and d6 are predicted to be labile. All these are borne out by experimental results on the dissociation of [M(AA)3]2+ (AA = phen, bipy),67 replacement of water in [M(OH2)6]2+ by other ligands,68 dissociation, racemization and oxalate exchange rates69 of [M(ox)3]3− (ox = oxalate), etc. It is interesting that the observed lability order d4 > d5 > d3 > d6 for the exchange of cyanide ligand in [M(CN)6]n– (n = 3, 4) with free cyanide in solution70 suggests the operation of the SN2 wedge mechanism, since it is only this mechanism that predicts a higher rate of reaction for d4 compared with d5 in octahedral complexes of strong field ligands. From the above-­mentioned results, it is clear that there is reasonable agreement between the valence bond and ligand field theories in predicting rapid reactions for complexes in which a lower d orbital is unoccupied, although the actual explanations are very different. The valence bond theory stresses an SN2 mechanism of reaction, the operation of which is favoured energetically if a lower d orbital is available (empty). In the ligand field theory, the reaction would be facile and may take place by either an SN1 or SN2 mechanism, provided that the transition-­state intermediate of lower or higher coordination number can be formed without any loss in LFSE (i.e. ΔEa = 0). Accordingly, it is not necessary for a complex having an empty inner d orbital in the metal to react by an SN2 process since energetically a dissociation mechanism is equally feasible. There is also agreement between the two theories in predicting the lability of complexes of non-­transition elements and the lanthanides (d0 configuration) and those of post-­transition elements (d10 configuration) such as Cu(i), Zn(ii), Ga(iii), etc. Both theories can predict the lability of all octahedral complexes of d9 and high-­spin d4, d5 and d6 systems. A major point of disagreement is for outer-­orbital (weak field) octahedral 6

Reactions of Metal Complexes

41

8

complexes of d , for which the valence bond theory predicts the same lability as for the similar complexes of d4 to d7 and d9 systems, whereas the ligand field theory predicts inertness comparable to that of a corresponding complex of a d3 ion. Where the theories agree, the theoretical predictions are perfectly in agreement with experimental observations. In fact, a complex that is not expected to be inert on LFSE considerations rarely reacts slowly. The exceptions are complexes of some polydentate ligands, such as the naturally occurring chlorophyll (which is extremely inert to exchange with Mg2+ ion in solution71) and [Fe(edta)]− (which exchanges with Fe3+ ion at an ordinarily measurable rate72). In agreement with the theoretical predictions, all the complexes of Cr(iii) are inert, and this is true even for [Cr(OH2)6]3+, which exchanges the ligand with 18OH2 in solvent extremely slowly (t½ ≈ 80 h at 25 °C).73 All the low-­spin complexes of Fe(ii), Fe(iii), Co(iii), Rh(iii), Os(iii) and Ir(iii) are very inert. In general, the low-­spin complexes of Co(iii) are more inert than the Cr(iii) complexes (see Table 1.5), which is in agreement with predictions of the ligand field theory. There are some apparent expectations of this relative order of reactivity of Cr(iii) and Co(iii) complexes. It has already been mentioned that the rate of exchange of H2O in [Cr(OH2)6]3+ with 18OH2 in solvent is extremely slow (kex = 2.4 × 106 s−1 at 25 °C);73 in contrast, however, the corresponding exchange in the case of [Co(OH2)6]3+ is extremely rapid.74 In this case, exchange takes place by a process involving catalysis by [Co(OH2)6]2+, which is invariably present in the system, due to the strongly oxidizing character and instability of [Co(OH2)6]3+, and no cleavage of the Co–O bond takes place:   

[Co(OH2)6]2+ + 6H2O* ⇌ [Co(*OH2)6]2+ + 6H2O (Rapid exchange)

(1.62)

  

  

[Co(OH2)6]3+ + [Co(*OH2)6]2+ ⇌ [Co(OH2)6]2+ + [Co(*OH2)6]3+ (Rapid electron transfer) (1.63)

A similar mechanism operates75 in the rapid exchange of the free ligand BB with its Co(iii) complex, [Co(BB)3]3+, where BB = 1,10-­phenanthroline, 2,2′-­bipyridine or ethylenediamine (Table 1.9). However, the inertness of these complexes is indicated by the fact that [Co(phen)3]3+ does not dissociate even on refluxing for 3 h in 2.5 M sulfuric acid, and no dissociation of [Co(en)3]3+ has been observed even in 1 M perchloric acid over a period of 60 days at 30 °C.43 In fact, in all other known cases where low-­spin Co(iii) complexes behave as labile, the reaction does not involve any cleavage of the metal–ligand bond, which is really the process predicted to be inert. Therefore, the extremely rapid aquation78 of [Co(CO3)(NH3)5]+ (instantaneous liberation of CO2 on adding acid to a solution of the carbonato complex) is due to the operation of the following mechanism (pseudo-­substitution process):   

Chapter 1

42

Table 1.9  Rate  constants for comparable reactions of some Co(III) and Cr(III) complexes.

Reaction Ligand exchange: [M(CN)6]3− + 6*CN− ⇌ [M(*CN6)]3− + 6CN− Aquation: [M(NH3)5X]2+ + H2O → [M(NH3)5(OH2)]3+ + X− X− = Cl− X− = Br− X− = NCS− Dissociation: [M(BB)3]3+ + H3O+ + H2O → [M(BB)2(H2O)2]3+ + BBH+ BB = ethylenediamine BB = biguanide BB = phenylbiguanide

Temperature/ °C Co(iii)

k or t½ Cr(iii)

Ref.

24

t½ = ∞

t½ = 986 h

70

25 25 25

1.7 × 10−6 s−1 6.3 × 10−6 s−1 2.8 × 10−9 s−1

9.7 × 10−3 s−1 6.3 × 10−5 s−1 9.3 × 10−8 s−1

76 76 76

30

No dissociation 1.6 × 10−6 s−1 43 and in 1 M HClO4 (at pH ≈ 1)80 77 in 60 days46 1.5 × 10−2 M−1 s−1 0.12 M−1 s−1 43 1.4 × 10−3 M−1 s−1 4.3 × 10−2 M−1 s−1 43

35.5 35.5



(1.64)

  

The “Rapid” step is the rate-­determining slowest step in the reaction sequence. The reaction involves protonation of the carbonate ligand forming a bicarbonate ligand, followed by O–C bond rupture, hence it is a decarboxylation process and not normal aquation. On carrying out the reaction in 18 O-­enriched water as solvent, no uptake of 18O in the product aqua complex

Reactions of Metal Complexes

43

79a,b

was detected, which proves that the reaction occurs without any rupture of the Co–O bond. Similar studies have been carried out80a,b with the complex [Co(CO3)(NH3)4]+, which has a chelated carbonate ligand, and it has been observed that in the diaqua complex formed only half of the aqua ligands are 18 OH2. This shows that the chelate ring opening involves normal Co–O bond breaking with entry of 18OH2 into the vacated position, followed by rapid loss of CO2 by O–C bond breaking and hence no entry of the 18OH2 in this case. The formation of [Co(ONO)(NH3)5]2+ in the reaction of the corresponding aqua complex with NO2− in HNO2–NO2− aqueous buffer solution at low temperature (ice bath) also involves a pseudo-­substitution process.81 This has been demonstrated by carrying out the reaction with 18O-­labelled aqua complex, which led to the formation of the 18O-­enriched nitrito complex.82a However, the reported anation of cis-­[Co(NH3)4(OH2)2]3+ by NO2− in HNO2 + NO2− buffer occurs in two consecutive steps, both being first order with respect to each of the reactants, but astonishingly with no evidence of the formation of any nitrito complex even as an intermediate in the reactions.82b Similar reactions leading to the formation of [M(ONO)(NH3)5]n+ [M = Rh(iii),83,84 Ir(iii),83,84 Pt(iv)84] under similar conditions have been reported. The formation of [Co(OSO2)(NH3)5]+ by the action of SO2 on [Co(OH) (NH3)5]2+ and its aquation on acidification leading to instantaneous loss of SO2 is very fast, hence this reaction also occurs without Co–O bond rupture.85a,b This has been demonstrated by 17O exchange studies followed by 17 O NMR spectroscopy.85c Detailed kinetic studies on the formation of this sulfito complex from the aqua complex and sulfite in aqueous ammonia buffer solution (pH > 9) have also been reported.85d Acidification of the sulfito complex causes instantaneous loss of SO2, forming the aqua complex, which fully retains the 17O. This result rules out linkage rotation, Co–17OSO2 ⇌ Co–OSO17O, although this is known to occur with the corresponding nitrito complex (see Chapter 6). In a similar manner, it has been shown that the formation of the carbonato complex, similarly to the sulfito complex, involves addition of CO2 without Co–O bond rupture.86 Similar studies87 based on 17O exchange using 17O NMR spectroscopy corroborated earlier evidence82 for pseudo-­substitution in the formation of [Co(ONO)(NH3)5]2+ from [Co(NH3)5(OH2)]3+. For the overall process:   

kf

[(NH3 )5 M – OH2 ]3  HCO3 #[(NH3 )5 M – CO3 ]  H  H2 O kb

  

(1.65)

the values of kf (s−1) and kb (M−1 s−1) at 25 °C for the complexes having M = Co, Rh and Ir are comparable: kf kb

Co88

Rh89

Ir89

1.10 220

1.13 470

1.45 590

Chapter 1

44

In contrast, complexes of Co(iii), Rh(iii) and Ir(iii), which dissociate by rupture of M–ligand bonds, have rates that vary widely. Thus, for the aquation of [M(Br)(NH3)5]2+ the kaq (s−1, 25 °C) values are 6.3 × 10−6 (Co),90 ca. 1 × 10−8 (Rh) and ca. 2 × 10−10 (Ir). The nearly identical rates for the formation and dissociation of the carbonato complexes of these metals is due to operation of the pseudo-­substitution mechanism, which does not involve rupture the metal– ligand bond (M–O bond in these cases). In each case, the rate is very much faster than that for a normal ligand replacement process. Thus, aquation of [Co(OSO3)(NH3)5]+,91 [Co(ONO2)(NH3)5]2+,92 and [Co(O2CCH3)(NH3)5]2+,93 all of which involve Co–O bond rupture, proceed very much slower (by a factor of ca. 107–109), the values of 106kaq being 1.2, 27 and 0.26 s−1 (at 25 °C), respectively. As mentioned above, the formation of the thermodynamically less stable red nitrito isomer of the pentaammine series, [Co(ONO)(NH3)5]2+, in over 80% yield by the action of HNO2–NO2− buffer (pH ≈ 4.1) on the aquapentaamminecobalt(iii) ion at low temperature (ice bath)81 is a pseudo-­substitution process. An apparent possibility is the following reaction scheme:   



  



Fast

(H3 N)5 Co3 – OH2 #(H3 N)5 Co2  – OH  H

(1.66)

H  NO2 #HNO2

Fast ( K eq ) 0 C HNO2  H # NO   H2 O ( K eq  2.6  105 )

(1.67)

  



(1.68)

  

A similar scheme involving attack by HNO2 instead of NO+ can also explain the formation of the nitrito complex without any rupture of the C–O bond. However, none of these are in accord with the experimental observation that in the weakly acidic (pH 4.1) buffered solution of NO2− and HNO2, in which the rate of formation has been studied, the rate law is   

Rate = k[Co(NH3)5(OH2)3+][HNO2]2

(1.69)

  

This rate law agrees with Scheme 1.3, proposed by Pearson, Basolo and co-­workers.81 Based on this, the rate is   

Rate = k′(K1K2K3/[H+])[(H3N)5Co3+–OH2][HNO2]2   

(1.70)

Reactions of Metal Complexes

45

Scheme 1.3 hence at a fixed pH the observed rate constant is k′(K1K2K3/[H+]) and   

Rate = k[Co(NH3)5(OH2)3+][HNO2]2

(1.71)

  

which has been observed81 experimentally.‡‡ The system illustrates that a knowledge of the rate law is necessary to arrive at a plausible mechanism when more than one scheme may explain formation of the reaction products. Since the process does not involve dissociation of the metal–ligand bond, formation of this nitrito isomer is favoured kinetically over that of the nitro isomer (formation of which would require M–OH2 bond dissociation), despite the fact that the nitrito isomer is thermodynamically less stable. The analogous complexes of Pt(iv), Rh(iii) and Ir(iii) have been made similarly.84 The rate law for their formation is similar to that for formation of the Co(iii) complex, hence the same mechanism is valid. A pseudo-­substitution mechanism also operates for aquation of [M(ONO) (NH3)5]2+ (M = Co, Rh, Ir) in acidic medium (see Chapter 5). Conversion of the red nitrito isomer [Co(ONO)(NH3)5]2+ into the yellow and thermodynamically more stable N-­bonded nitro isomer [Co(NO2) (NH3)5]2+ (linkage isomerization) is a slow process both in the solid state and in solution, owing to fission of the Co–O bond involved in this change (see Chapter 6). ‡‡

In their book (ref. 1, 2nd edn, p. 231), but not in the original publication,81 Basolo and Pearson appropriately showed the site of N2O3 involved in the electrophilic attack on the oxygen of the OH group of the hydroxo complex. However, they expressed the rate law as Rate = k[hydroxo complex][HNO2]2 (1.72) Although this is not incorrect, it is more appropriate to use the concentration of the hydroxo complex in terms of the concentration of the aqua complex for which the pKa value is 6.55 at 25 °C; a literature value of 5.7 mentioned by them is at 15 °C (see Stability Constants, Special Publication No. 17, Chemical Society, London, 1964, Table 2). Hence, at the pH of their experimental solutions of ca. 4.1, the concentration of the hydroxo complex was negligible and its concentration could be expressed as K1[aqua complex]/[H+]. Substitution of this value in eqn (1.72) will lead to eqn (1.71) at a fixed pH.

Chapter 1

46

Table 1.10  Base  hydrolysis of Co(iii) and Cr(iii) complexes (25 °C). kOH/M−1 s−1 Complex

Co(iii)

Ref.

Cr(iii)

Ref.

0.87 6 2.8 × 10−3 1 × 103 3 × 103

89 89 96 97 97

1.7 × 10−3 6.9 × 10−2 2.5 × 10−6 2.8 × 10−2 3.7 × 10−2

98 98 23f 99 99

2+

[M(X)(NH3)5] X− = Cl− X− = Br− X− = NCS− cis-­[MCl2(en)2]+ trans-­[MCl2(en)2]+

In contrast to the rate of aquation, the rate of displacement of an acidic group by OH− (called base hydrolysis) is usually very much faster for Co(iii) than for Cr(iii), as shown by the data in Table 1.10. This suggests special mechanisms of reaction in these cases, other than a simple SN1 or SN2 mechanism, and possibly different mechanisms for Co(iii) and Cr(iii). Furthermore, it is known94 that for similar complexes of Co(iii), Rh(iii) and Ir(iii), the rate for a comparable reaction decreases in the sequence Co(iii) to Rh(iii) to Ir(iii), in accord with the fact that the value of Dq increases in the same sequence. However, other factors are also of importance. In contrast to the behaviour of [Fe(CN)6]3+ (low-­spin d5), the exchange of [Fe(C2O4)3]3− (high-­spin) with C2O42− is extremely rapid,95a as is to be expected. The corresponding Cr(iii) (d3 system) and Co(iii) (low-­spin d6 system) complexes are, however, inert.95b Consideration of the experimental values of the kinetic parameters (rate constant k, activation energy Ea and entropy of activation ∆S‡) for the dissociation of chelate complexes of d-­block metal ions and the calculated values of ∆Ea (ligand field contribution to the activation energy, assuming a dissociation mechanism) as given in Table 1.11 is interesting. From the data in Table 1.11, it appears that for all cases where ∆Ea = 0 the rates are very fast. Again, for the highest complexes of any particular ligand, the activation energy increases by ca. 5–8 kcal mol−1 in each step, in the sequence d7 < d8 < d6 (low-­spin), corresponding to an increase in ∆Ea of 2Dq. Since the value of Dq for these complexes is ca. 3–4 kcal mol−1, the observed increase in the activation energy is nicely accounted for. For the lower complexes also, Ea increases by ca. 4–7 kcal mol−1 on passing from d7 to d8, corresponding to an increase in ∆Ea of 2Dq. Similarly, the increase in ∆Ea from Cu(bipy)2+ to Ni(bipy)2+ is 9.6 kcal mol−1, corresponding to an increase in ∆Ea of 2Dq. However, the activation energy for the dissociation of Fe(terpy)2+, which is of the high-­spin type, is ca. 2 kcal mol−1 less than that for Co(terpy)2+, for both of which ∆Ea = 0, whereas ∆Ea for Ni(terpy)2+ is ca. 6 kcal mol−1 higher than that for Fe(terpy)2+, corresponding to an increase in ∆Ea of 2Dq. Fe(terpy)2+ is of high-­spin type, in contrast to [Fe(terpy)2]2+, which is of the low-­spin type, and in agreement with this the former is very much labile and the activation energy of the latter is ca. 11 kcal mol−1 greater than that of the former. This is an advantage of the ligand field theory over the valence

Reactions of Metal Complexes

47

Table 1.11  Kinetic  parameters and LFSE losses (∆Ea, LFAE) in the formation of a

transition state for the dissociation of some complexes in aqueous solutionb (Ea = ΔH‡ + RT; see Chapter 2, Section 2.13).a

Electronic configuration

ΔEa

Complex ionb

k/s−1 at 25 °C

Ea/kcal mol−1

∆S‡/cal K−1 mol−1

t2g3eg0 t2g4eg0 t2g3eg2 t2g6eg0

2 1.43 0 4

t2g4eg2

0

t2g5eg2

0

t2g6eg2

2

V(phen)32+ Cr(bipy)32+ Mn(phen)2+ Fe(phen)32+ Fe(bipy)32+ Fe(terpy)22+ Fe(phen)2+ Fe(terpy)2+ Co(phen)22+ Co(phen)2+ Co(terpy)22+ Co(terpy)2+ Ni(phen)32+ Ni(phen)22+ Ni(phen)2+ Ni(bipy)32+ Ni(bipy)2+ Ni(terpy)22+ Ni(terpy)2+ Cu(phen)22+ Cu(phen)2+ Cu(bipy)2+ Zn(phen)2+ Zn(bipy)2+ Cd(phen)2+ Hg(bipy)2+ Ag(phen)2+

8.8 × 10−5 c 0.38d 1.5 × 102 8 × 10−5 1.6 × 10−4 2.2 × 10−5 2.7 0.47 1.44 0.2 5.5 × 10−2 7 × 10−3 2.2 × 10−6 4 × 10−5 1.15 × 10−5 3 × 10−3 4.8 × 10−3 1.3 × 10−4 1.8 × 10−5 2 × 102 63.3 0.3 6.0 23.3 43.3 32.5 5.2 × 102

21.3 22.6 10.4 32.1 28.4 28.7 12.8 18.0 19.4 20.6 14.8 20.2 26.2 23.1 25.2 22.2 23.7 20.8 24.2 — — 14.1 12.3 12.1 14.4 10.7 18.0

−8 13 −16 28 17 14 −16 −2 5 5 −17 −3 1 −4.6 1 2 8 −9 −6 — — −16 −16 −14 −5 −18 12

t2g6eg3 t2g6eg4

0

a

 ata from ref. 100a–d except as specified in footnotes c and d. D phen, 1.10-­phenanthroline; bipy, 2,2′-­bipyridine; terpy, 2,2′,2″-­terpyridine. Calculated value using Ea and ∆S‡ values quoted in ref. 1 (2nd edn), Chapter 3, Table 3.10. d Data from ref. 101. b c

bond method, as the former can usually account for the observed differences in activation energies for the reactions of a series of similar complexes, often in a semiquantitative manner. In the cases of the Co(ii) and Ni(ii) complexes, however, we find that the activation energy of the mono species is ca. 3–5 kcal mol−1 greater than that for the bis species. Hence the effective ligand field contribution to the activation energy for the bis species of Fe(ii) is ca. 14–16 kcal mol−1 relative to that for the mono species. This is in very good agreement with the predicted value of ∆Ea = 4Dq ≈ 12–16 kcal mol−1. It should also be mentioned that although the rate of reaction generally decreases in the same sequence as the increase in activation energy for similar complexes of any of the ligands, the case of [Fe(phen)3]2+ is exceptional, as this reacts much faster than the corresponding Ni(ii) complex. This is due, however, to the exceptionally high value of

Chapter 1

48 ‡

the entropy of activation, ∆S , in the case of the Fe(ii) complex, which more than compensates for the high activation energy, thereby making the Fe(ii) complex react faster. The high value of ∆S‡ in this case is believed to be due to the fact that dissociation of the Fe(ii) complex causes a changeover from the low-­spin to the high-­spin form and consequently involves an expanded form of the complex in the transition state, in which the metal–ligand bonds are possibly stretched to become larger (free ligand). The same also holds good for [Fe(bipy)3]2+, but in this case the value of ∆S‡ is not sufficiently high, hence the effect is less on the relative rates of Fe(ii) and Ni(ii) complexes of bipy. This incidentally shows that in order to test the applicability of the ligand field theory in reaction kinetics, we have to take into consideration the relative values of the activation energies rather than relative rates. However, in most analogous systems, the rate decreases in order of increasing activation energy, without the entropy of activation making any significant contribution to change this order, but this may not always be true. In aqueous solution, where the dissociation reactions have been studied, a sufficient number of water molecules will invariably be present in the coordination sphere to keep the metal ion six-­coordinate, except in the cases of Hg(ii) and Ag(i). Complexes of d5 (high-­spin), d9 and d10 ions react very rapidly, and this lability is to be expected from consideration of the LFSE losses (∆Ea) (see Table 1.4). The case of [Cr(bipy)3]2+ is interesting because of its low-­spin character (∆Ea = 1.43Dq) and its activation energy is comparable to that of [Ni(bipy)3]2+, but the considerably higher rate of the Cr(ii) complex is an entropy effect. However, more common cases of Cr(ii) complexes are of high-­ spin type and these are all very labile like those of Cu(ii). In agreement with the predictions, [V(phen)3]2+ (d3) is also inert like [Ni(phen)3]2+ (d8). However, the smaller Ea value for [V(phen)3]2+ may be due to the lower effective nuclear charge on the metal, and M–L π-­bonding relatively weaker for d3 V(ii) than for d8 Ni(ii), resulting in a lower M–L bond energy in [V(phen)3]2+. However, kex of [M(OH2)6]2+ with H2O* in aqueous solution at 25 °C is 90 s−1 for V(ii) (Dq = 1200 cm−1) and 3 × 104 s−1 for Ni(ii) (Dq = 850 cm−1); the experimentally determined Ea values are ca. 17 kcal mol−1 for [V(OH2)6]2+ and ca. 12.2 kcal mol−1 for [Ni(OH2)6]2+, and the difference is much larger than the predicted value based on the actual ∆Ea values in the two cases. Hence not too much emphasis should be placed on the satisfactory agreement in some cases and the predictions of the theory are at best only qualitative. From consideration of the LFSE losses (Table 1.8), it may appear that octahedral complexes of a d8 ion will have the same degree of inertness as those of a d3 ion. However, ∆Ea values predict only the relative rates and not the absolute rates. The absolute rate is dependent on several other factors, such as the charge on the central atom, the nature of the ligands, etc., and in any application of the ∆Ea value to interpret the lability of a complex it is necessary to compare the behaviours of similar complexes of different dn ions. Certainly, ligand replacement involving M–L bond dissociation will be more favourable for Ni(ii) than for Cr(iii). Hence the reactions of Ni(ii)

Reactions of Metal Complexes

49

will be faster than those of Cr(iii). For octahedral complexes of Ni(ii) and Cr(iii) of any ligand, the Dq value is much higher for Cr(iii) than for Ni(ii), hence actual values of ∆Ea are much higher for the Cr(iii) complexes. However, in comparison with the high-­spin complexes of Mn(ii), Fe(ii) and Co(ii) and the complexes of Cu(ii) and Zn(ii), the reactions of Ni(ii) are definitely slower (Table 1.7). Even the rates of formation and dissociation of the labile [Ni(en)3]2+, [Ni(en)2(OH2)2]2+ and [Ni(en)(OH2)4]2+ are less than those of the corresponding Cu(ii) complexes;102 the rates of formation of [Ni(phen)(OH2)4]2+ is also quite conveniently measurable.103 Again, the exchange of *Ni2+ with [Ni(edta)]2− is slow,104 but the corresponding reaction of Co(ii) is faster105 and that of Fe(ii) is too fast.106 In a large number of cases of octahedral complexes showing normal behaviour, the rates of formation and dissociation of Cu(ii) complexes are 104–105 times faster than those of the corresponding Ni(ii) complexes (see Chapter 5).107 Complexes of transition metal ions for which ∆Ea = 0 are all labile, as expected, but in these cases, unlike in the cases of complexes of d0 and d10 ions, the lability is not a simple function of the size and charge of the ion as seen in cases of complexes of M2+ and M3+ ions of this latter groups, as mentioned earlier. Data for the aqua ions of Ti3+, V3+ and Fe3+ (high spin) are given in Table 1.12 (see Figure 1.3). Thus, despite the same charge and nearly identical sizes of these ions, we observe large variations in the lability of their octahedral hexaaqua ions; the lability increases with decreasing electron density in the d orbital and consequently less hindrance due to electrostatic repulsion to the incoming nucleophile in some sort of associative activation.§§ Accordingly, d0 [Sc(OH2)6]3+ ( rSc3 = 0.89 Å) is much more labile, kex = 3.7 × 107 s−1 (25 °C), than [Ti(OH2)6]3+. There are other interesting applications of the ligand field theory in the elucidation of mechanisms of reactions of complexes.96,108 If part of the energy of activation in reactions of complexes were to be due to the energy required in the cleavage of a metal–ligand bond, the activation energies for a series of complexes of the same metal ion should be related to the ligand field strengths of the ligands. It is known that the longest wavelength absorption peak in the visible spectrum of a complex is generally a Laporte-­ forbidden, spin-­allowed transition, which for octahedral complexes gives the Table 1.12  Data  for the aqua ions of Ti3+, V3+ and Fe3+. [M(OH2)6]3+ Parameter

M = Ti

M=V

M = Fe

rM3 /Å Dq/cm−1 kex/s−1 a

0.81 2000 4.8 × 105

0.78 1995 1.3 × 103

0.79 1400 1.6 × 102

a

For H2O* exchange (25 °C).

§§

These reactions presumably occur by an Ia mechanism (see Chapter 4).

Chapter 1

50

value of Dq after making corrections for parameters of interelectronic repulsion (which amount to at most 20%).53 The acidopentaammine complexes of Co(iii) do not have perfect octahedral symmetry and the extent of distortion will depend on how far away the sixth ligand is from ammonia in its position in the spectrochemical series. Neglecting this factor, however, and assuming an average environment as an approximation, one may set up a crude spectrochemical series for the acidopentaamminecobalt(iii) complexes by considering the νmax for the first ligand field band of each of the complexes. It has been found that the activation energy (Ea) for the following aquation reaction:   

  

[Co(NH3)5X]2+ + H2O → [Co(NH3)5(OH2)3+] + X

(1.73)

increases almost linearly with increase in νmax, i.e. increases with increase in the ligand field strength in the following order:108 RBr2+ < RCl2+ < R(NO3)2+ < R(NCS)2+ < R(NO2)2+ [where R = Co(NH3)53+]. Similar results have been reported108 for base hydrolysis, RBr2+ < RCl2+ < RN32+, but the position of the nitropentaamine complex is anomalous, possibly due to a π-­bonding effect, which is believed to play an important role in this case.109 From detailed consideration of the manner in which the enthalpy of activation, ∆H‡, for base hydrolysis of a series of cobalt(iii) complexes of the type [Co(X)(NH3)5]2+ (where X = Cl−, Br−, N3−, NO2− and NCS−) changes with the ligand field strength, Banerjea and Das Gupta concluded96 that whereas the chlorido, bromido and azido complexes react by the SN1-­IP (D-­IP) mechanism, the nitro and isothiocyanato complexes react by the SN2-­IP (A-­IP) (IP = ion pair) process (see Chapter 4). The difference in behaviour has been explained as being due to the ability of NO2− and NCS− to form a π-­bond of the metal-­to-­ligand type. Lalor and Carrington110 reported a linear relationship between the activation energies and the ligand field parameters 10Dq for base hydrolysis of [M(i)(NH3)5]2+ complexes (M = Cr, Co, Rh, Ir), which suggests operation of a common mechanism (D-­CB) in these cases. The difference in rates is obviously due to significant variation with the metal of the labilizing effect of the amido ligand in the conjugate base forms of the complexes. Activation energies for base hydrolysis of the halidopentaammine series of complexes of Rh(iii)111a–c differ from those of the corresponding Co(iii) series by 1–4 kcal mol−1, which is in fair agreement with the predicted value of 5 kcal mol−1 based on the ligand field theory.111d The kinetic features of the reactions of four-­coordinate square-­planar complexes have also been investigated in a number of cases (see Chapter 3). Detailed investigations on complexes of Pt(ii)112 and Pd(ii)113 seem to indicate that these are generally inert, although complexes of Pd(ii) are generally much more labile than the analogous complexes of Pt(ii), and the rate of reaction depends considerably on the nature of the incoming group, the group being displaced and the other groups present in the complex. In the case of Pt(ii) complexes, reaction proceeds fairly rapidly when the incoming group and/or the group in the trans position to the group being displaced is

Reactions of Metal Complexes

51

capable of forming a strong π-­bond of the metal-­to-­ligand type, while even the most powerful nucleophilic reagent, OH− ion, is no better than H2O, NH3, pyridine or Cl− ion. These results have provided much valuable information on the well-­known114 trans effect phenomenon in Pt(ii) complexes (see Chapter 3). In the case of Pd(ii) complexes, however, the rate of reaction is very much dependent on the nucleophilic character of the incoming group,113 including OH−. Exchange of *Cl− with AuCl4− also proceeds slowly,115 but the rate is ca. 200 times faster than the corresponding exchange in the Pt(ii) system (see Chapter 3).116 The general inertness of complexes of Pd(ii), Pt(ii) and Au(iii) is due to the ligand field contribution to the activation energy, ΔEa, which can be shown to be fairly high for the complexes of d8 ions (see Table 1.13). In general, an associative mechanism operates for ligand replacement in square-­planar complexes (see Chapter 3) and the reaction intermediate is a species of coordination number five, for which the two possible alternative structures are trigonal bipyramidal and square pyramidal. The loss in LFSE in the formation of these transition states for the reactions of the square-­ planar complexes of d7, d8, d9 and d10 ions can be calculated with the help of the data in Table 1.2 as usual and the results arrived at in this way are given in Table 1.13. We need only consider strong field cases, where the square-­planar structure is most likely to be stabilized (see Table 1.6), and the observed facts are in agreement with this prediction. From the data in Table 1.13, the trigonal bipyramid is expected to be energetically less favoured than the square pyramid. However, it can be shown that the former is more likely to be formed, if stabilized by π-­bonding.112 The other point is that the order of reactivity should be d10 ≫ d9 > d8 > d7. However, data on ligand replacement or ligand exchange in square-­planar complexes of all these M(ii) complexes are not available to verify this prediction. But in accordance with the LFAE values, unlike in the cases of octahedral complexes, the square-­planar complexes of Cu(ii) are also predicted to be inert but less so than the analogous complexes of Ni(ii). In agreement with this are the results of dissociation of [M(baen)] complexes (M = Ni, Cu) and of square-­planar complexes of these metals formed by macrocyclic ligands of the cyclam type (see Chapter 5). Studies on the exchange of M2+ ions (M = Co, Ni, Cu and Zn) with their chelate complexes (square-­planar) of quadridentate ligands (two such) of the Schiff Table 1.13  Ligand  field contribution (ΔEa) to the activation energy for reactions of square-­planar complexes (strong field cases). LFSE/Dq System 7

d d8 d9 d10

ΔEa/Dq

Square planar (1)

Trigonal bipyramidal (2)

Square pyramidal (3)

(1) − (2)

(1) − (3)

26.84 24.56 12.28 0

13.34 14.16 7.09 0

19.14 18.28 9.14 0

13.50 10.40 5.19 0

7.70 6.28 3.14 0

Chapter 1

52 117

base in solution in pyridine have shown that the rate increases in the order Ni(ii) < Cu(ii) < Co(ii) < Zn(ii); the Ni(ii) complexes are so inert that they show no exchange even after 48 h at room temperature, whereas for the Zn(ii) complexes exchange is complete in ca. 0.5 min, under similar conditions. These exchange reactions must be due to reversible dissociation of the complexes and hence their rates are expected to follow the trend which is the reverse of the LFSE values. The observed order is in agreement, except for Co(ii), with this prediction (see Table 1.13). Although from magnetic and X-­ray crystallographic data these complexes are known to be square-­planar in the solid state, there is no authentic evidence that this structure is retained in solution in a coordinating solvent without expansion of coordination number due to bonding of one or two solvent molecules. Such an increase in coordination number is more likely for square-­planar stable complexes of d7 Co(ii) than those of Ni(ii) having a strong ligand field (see LFSE data in Table 1.3), which may account for the anomalous behaviour of Co(ii). It is worth noting that whereas in a solution of cyanide Co(ii) exists as [Co(CN)5]3−, Ni(ii), Cu(ii) and Zn(ii) are present as [M(CN)4]2− {some [Ni(CN)5]3− is formed in high concentrations of CN− since its formation constant for formation from [Ni(CN)4]2− is 0.19 M−1 at 25 °C (see Chapter 3, Section 3.5}.

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Reactions of Metal Complexes

55

37. C. Chatterjee and P. Chaudhuri, Indian J. Chem., 1971, 9, 1132. 38. A. Y. Kasim and Y. Sulfab, Inorg. Chim. Acta, 1977, 24, 247. 39. D. Banerjea and C. Chatterjee, Z. Anorg. Allg. Chem., 1968, 361, 99. 40. (a) M. Eigen, Pure Appl. Chem., 1965, 6, 105; (b) M. Eigen and R. G. Wilkins, Adv. Chem. Ser., 1965, 49, 55; (c) A. McAuley and J. Hill, Q. Rev., 1969, 23, 18; (d) D. J. Hewkin and R. H. Price, Coord. Chem. Rev., 1970, 5, 45; (e) A. Cusanelli, U. Frey, D. T. Richens and A. E. Merbach, J. Am. Chem. Soc., 1996, 118, 5265. 41. H. Taube, Chem. Rev., 1952, 50, 69. 42. (a) A. A. Grinberg and L. E. Nikol’skoya, Zhur. Priklad. Khim., 1949, 22, 542; (b) A. A. Grinberg and L. E. Nikol’skoya, Zhur. Priklad. Khim., 1951, 24, 893; (c) D. R. Stranks and R. G. Wilkins, Chem. Rev., 1957, 57, 743. (data in Table 10). 43. D. Banerjea and B. Chakravarty, J. Inorg. Nucl. Chem., 1964, 26, 1233. 44. P. George and D. S. McClure, in Progress in Inorganic Chemistry, ed. F. A. Cotton, Interscience, New York, 1959, vol. 1, p. 428. 45. O. G. Holmes and D. S. McClure, J. Chem. Phys., 1957, 26, 1686. 46. (a) C. J. Ballhausen, Introduction to Ligand Field Theory, McGraw-­Hill, New York, 1962; (b) B. N. Figgis, Introduction to Ligand Fields, Interscience, New York, 1965; (c) T. M. Dunn, D. S. McClure and R. G. Pearson, Some Aspects of Crystal Field Theory, Harper and Row, New York, 1966; (d) C. K. Jørgensen, Recent Progress in Ligand Field Theory, Springer, New York, 1966; (e) C. K. Jørgensen, Modern Aspects of Ligand Field Theory, North Holland, Amsterdam, 1971; (f) B. N. Figgis, Ligand Field Theory, in Comprehensive Coordination Chemistry, ed. G. Wilkinson et al., Pergamon Press, New York, 1987, ch. 3, vol. 6. 47. (a) A. L. Companion and M. A. Komarynsky, J. Chem. Educ., 1964, 41, 257; (b) R. F. W. Bader and A. D. Westland, Can. J. Chem., 1961, 39, 2306; (c) C. Jørgensen and H. H. Schimidke, Z. Phys. Chem., 1963, 38, 118; (d) D. R. Loyd and E. W. Schlag, Inorg. Chem., 1969, 8, 2544. 48. (a) A. J. Freeman and R. E. Watson, Phys. Rev., 1960, 120, 1254; (b) F. A. Cotton, J. Chem. Educ., 1964, 41, 466. 49. B. Bosnich, W. G. Jackson and S. T. D. Lo, Inorg. Chem., 1975, 14, 2998. 50. (a) R. F. Fenske, D. S. Martin Jr and K. Ruedenberg, Inorg. Chem., 1962, 1, 441; (b) H. B. Gray and C. J. Ballhausen, J. Am. Chem. Soc., 1963, 85, 260. 51. (a) Von F. E. Ilse and H. Hartmann, Z. Phys. Chem., 1951, 197, 239; (b) Von F. E. Ilse and H. Hartmann, Z. Naturforsch., 1951, 6A, 751. 52. (a) C. J. Ballhausen, Kgl. Danske Vidensk. Selsk. Mat.-­Fys. Medd., 1954, 29(4); (b) C. J. Ballhausen, Introduction to Ligand Field Theory, McGraw-­ Hill, New York, 1962, pp. 108–110. 53. C. K. Jørgensen, Absorption Spectra and Chemical Bonding in Complexes, Pergamon Press, Oxford, 1962. 54. (a) G. Maki, J. Chem. Phys., 1958, 29, 162; (b) G. Maki, J. Chem. Phys., 1958, 29, 1129; (c) C. J. Ballhausen and A. D. Liehr, J. Am. Chem. Soc., 1959, 81, 538; (d) S. C. Nyburg and J. S. Wood, Inorg. Chem., 1964, 3, 468. 55. L. G. Vanquickenborne and A. Ceulemans, Inorg. Chem., 1981, 20, 796. 56. R. S. Mulliken, Phys. Rev., 1933, 43, 279.



56

Chapter 1

57. J. H. Van Vleck, J. Chem. Phys., 1935, 3, 803–807. 58. H. Bethe, Ann. Phys., 1929, 395, 133. 59. (a) J. Chatt, G. A. Gamlen and L. E. Orgel, J. Chem. Soc., 1958, 486; (b) D. S. Martin Jr, M. A. Tucker and A. J. Kassman, Inorg. Chem., 1965, 4, 1682; (c) F. A. Cotton and C. B. Harris, Inorg. Chem., 1967, 6, 369; (d) M. A. Hitchman and H. B. Gray, Inorg. Chem., 1979, 18, 1745. 60. (a) J. Owen, Proc. R. Soc. A, 1955, 22, 183; (b) J. S. Grigffith, J. Inorg. Nucl. Chem., 1966, 22, 1229. 61. W. G. Penney and W. G. Schlapp, Phys. Rev., 1932, 41, 194. 62. (a) A. J. Freeman and R. E. Watson, Phys. Rev., 1960, 120, 1254; (b) F. A. Cotton, J. Chem. Educ., 1964, 41, 466. 63. (a) C. J. Ballhausen and C. K. Jørgensen, Kgl. Danske Vidensk. Selsk. Mat.-­ Fys. Medd., 1955, 29(14); (b) N. S. Hush, Aust. J. Chem., 1962, 15, 378; (c) Ref. 1 (2nd edn), p. 69; (d) J. J. Zuckerman, J. Chem. Educ., 1965, 42, 315; (e) R. Krishnamurty and W. B. Schaap, J. Chem. Educ., 1969, 46, 799. 64. (a) L. E. Orgel, J. Chem. Soc., 1952, 4756; (b) L. E. Orgel, J. Chem. Phys., 1955, 23, 1004; (c) L. E. Orgel, J. Chem. Phys., 1955, 23, 1819; (d) P. George, D. S. McClure, J. S. Griffith and L. E. Orgel, J. Chem. Phys., 1956, 24, 1269. 65. C. K. Jørgensen, Acta Chem. Scand., 1955, 9, 605. 66. A. L. Companion, J. Phys. Chem., 1969, 73, 739. 67. (a) P. Ellis and R. G. Wilkins, J. Chem. Soc., 1959, 299; (b) R. Hogg and R. G. Wilkins, J. Chem. Soc., 1962, 341. 68. (a) M. Eigen, Z. Elektrochem., 1960, 64, 115; (b) M. Eigen and K. Tamm, Z. Elektrochem., 1962, 66, 93; (c) M. Eigen and K. Tamm, Z. Elektrochem., 1962, 66, 107; (d) M. Eigen, Ber. Bunsenges. Phys. Chem., 1964, 67, 753. 69. (a) K. V. Krishnamurty and G. M. Harris, Chem. Rev., 1961, 61, 213; (b) R. W. Oliff and A. L. Odell, J. Chem. Soc., 1964, 2417. 70. (a) A. W. Adamson, J. P. Welker and W. B. Wright, J. Am. Chem. Soc., 1951, 73, 4789; (b) A. G. MacDiarmid and N. F. Hall, J. Am. Chem. Soc., 1954, 76, 4222; (c) J. J. Christensen, R. M. Izatt, J. D. Hale, R. T. Pack and G. D. Watt, Inorg. Chem., 1963, 2, 337. 71. (a) S. Ruben, G. T. Seaborg and J. W. Kennedy, J. Appl. Chem., 1941, 12, 308; (b) S. Ruben, A. W. Frinkel and M. D. Kamen, J. Phys. Chem., 1942, 46, 710; (c) R. S. Becker and R. K. Sheline, J. Chem. Phys., 1943, 21, 946. 72. (a) C. M. Cook Jr and F. A. Long, J. Am. Chem. Soc., 1951, 73, 4119; (b) S. S. Jones and F. A. Long, J. Chem. Phys., 1952, 56, 25. 73. (a) cf. A. E. McAuley and J. Hill, Q. Rev., 1969, 23, 18; (b) A. Cusanelli, U. Frey, D. T. Richens and A. E. Merbach, J. Am. Chem. Soc., 1996, 118, 5265. 74. H. L. Friedman, H. Taube and J. P. Hunt, J. Chem. Phys., 1950, 18, 757. 75. P. Ellis, R. G. Wilkins and M. J. G. Williams, J. Chem. Soc., 1957, 4456. 76. D. R. Stranks, The Reaction Rates of Transitional Metal Complexes, in Modern Coordination Chemistry, ed. J. Lewis and R. G. Wilkins, Interscience, New York, 1960, ch. 2, Table IX and references cited therein. 77. H. L. Schlafer and O. Kling, Z. Phys. Chem., 1958, 16, 14. 78. A. B. Lamb and K. J. Mysels, J. Am. Chem. Soc., 1945, 67, 468.

Reactions of Metal Complexes

57

79. (a) J. P. Hunt, A. C. Rutenberg and H. Taube, J. Am. Chem. Soc., 1952, 74, 268; (b) C. A. Bunton and D. R. Llewellyn, J. Chem. Soc., 1953, 1692. 80. (a) F. A. Posey and H. Taube, J. Am. Chem. Soc., 1953, 75, 4099; (b) H. Scheidegger and G. Schwarzenbach, Chimia, 1965, 19, 166. 81. R. G. Pearson, P. M. Henry, J. G. Bergmann and F. Basolo, J. Am. Chem. Soc., 1954, 76, 5920. 82. (a) R. K. Murmann and H. Taube, J. Am. Chem. Soc., 1956, 78, 4886; (b) M. C. Ghosh and P. Banerjee, Bull. Chem. Soc. Jpn., 1983, 56, 2871. 83. F. Basolo and G. S. Hammaker, J. Am. Chem. Soc., 1960, 82, 100. 84. F. Basolo and G. S. Hammaker, Inorg. Chem., 1962, 1, 1. 85. (a) R. van Eldik and G. M. Harris, Inorg. Chem., 1980, 19, 880; (b) R. van Eldik, Inorg. Chim. Acta, 1980, 42, 49; (c) R. Van Eldik, J. Von Jouanne and H. Kelm, Inorg. Chem., 1982, 21, 2818; (d) U. Spitzer and R. Van Eldik, Inorg. Chem., 1982, 21, 4008. 86. R. Van Eldik, D. A. Palmer, H. Kelm and G. M. Harris, Inorg. Chem., 1980, 19, 3679 and literature cited therein. 87. W. G. Jackson, G. A. Lawrence, P. A. Lay and A. M. Sargeson, J. Chem. Soc., Chem. Commun., 1982, 70. 88. E. Chaffee, T. P. Dasgupta and G. M. Harris, J. Am. Chem. Soc., 1973, 95, 4169. 89. D. A. Palmer and G. M. Harris, Inorg. Chem., 1964, 13, 965. 90. A. W. Adamson and F. Basolo, Acta Chem. Scand., 1955, 9, 1261. 91. (a) B. Adell, Z. Anorg. Chem., 1942, 249, 251; (b) H. Taube and F. A. Posey, J. Am. Chem. Soc., 1953, 75, 1463. 92. J. N. Bronsted, J. Phys. Chem., 1926, 122, 383. 93. F. Basolo, J. G. Bergmann and R. G. Pearson, J. Phys. Chem., 1952, 56, 22. 94. S. A. Johnson, F. Basolo and R. G. Pearson, J. Am. Chem. Soc., 1963, 85, 1741. 95. (a) F. A. Long, J. Am. Chem. Soc., 1941, 63, 1353; (b) H. C. Clark, N. F. Curtis and A. L. Odell, J. Chem. Soc., 1954, 63. 96. D. Banerjea and T. P. Das Gupta, J. Inorg. Nucl. Chem., 1966, 28, 1667. 97. (a) R. G. Pearson, R. E. Meeker and F. Basolo, J. Inorg. Nucl. Chem., 1955, 1, 342; (b) R. G. Pearson, R. E. Meeker and F. Basolo, J. Am. Chem. Soc., 1956, 78, 2673. 98. M. A. Levine, T. P. Jones, W. E. Harris and W. J. Wallace, J. Am. Chem. Soc., 1961, 83, 2453. 99. R. G. Pearson, R. A. Munson and F. Basolo, J. Am. Chem. Soc., 1958, 80, 504. 100. (a) P. Ellis and R. G. Wilkins, J. Chem. Soc., 1959, 299; (b) R. Hogg and R. G. Wilkins, J. Chem. Soc., 1962, 341; (c) R. S. Bell and N. Sutin, Inorg. Chem., 1962, 1, 359; (d) R. H. Holyer, C. D. Hubbard, S. F. A. Kettle and R. G. Wilkins, Inorg. Chem., 1965, 4, 929. 101. B. R. Baker and B. D. Mehta, Inorg. Chem., 1965, 4, 848. 102. (a) J. Bjerrum and K. G. Poulsen, Nature, 1952, 169, 463; (b) A. K. S. Ahmed and R. G. Wilkins, J. Chem. Soc., 1960, 2091.

58

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103. D. W. Margerum, R. I. Bystroff and C. V. Banks, J. Am. Chem. Soc., 1956, 78, 4211. 104. C. M. Cook Jr and F. A. Long, J. Am. Chem. Soc., 1958, 80, 33. 105. R. Barbieri, U. Belluco and M. Bruno, Gazz. Chim. Ital., 1957, 87, 1393. 106. Ref. 77, page 143. 107. T. S. Roche and R. G. Wilkins, J. Am. Chem. Soc., 1974, 96, 5082. 108. G. C. Lalor and J. Long, J. Chem. Soc., 1963, 5620. 109. P. J. Staples and M. L. Tobe, J. Chem. Soc., 1960, 4812. 110. G. C. Lalor and T. Carrington, J. Chem. Soc. A, 1969, 2328. 111. (a) G. W. Bushnell and G. C. Lalor, J. Inorg. Nucl. Chem., 1968, 30, 219; (b) C. S. Davies and G. C. Lalor, J. Chem. Soc. A, 1968, 2328; (c) G. C. Lalor, J. Inorg. Nucl. Chem., 1969, 31, 1206; (d) C. H. Langford, Mechanism and Steric Course of Octahedral Substitution, in MTP International Review of Science: Inorganic Chemistry, Series One, ed. M. L. Tobe, Butterworths, London, 1972, ch. 6, vol. 9, p. 243. 112. D. Banerjea, F. Basolo and R. G. Pearson, J. Am. Chem. Soc., 1957, 79, 4055. 113. D. Banerjea and K. K. Tripathi, J. Inorg. Nucl. Chem., 1958, 7, 78. 114. J. V. Quagliano and L. Schubert, Chem. Rev., 1952, 50, 201. 115. R. L. Rich and H. Taube, J. Phys. Chem., 1954, 58, 1. 116. (a) L. F. Grantham, T. S. Elleman and D. S. Martin Jr, J. Am. Chem. Soc., 1955, 77, 2965; (b) T. S. Elleman, J. W. Reishus and D. S. Martin Jr, J. Am. Chem. Soc., 1958, 80, 536. 117. B. O. West, J. Chem. Soc., 1954, 395.

Chapter 2

Techniques for Following Reactions and Factors that Affect Rates 2.1  Techniques for Evaluating Rates of Reactions The rate of a chemical reaction is the rate of change with time of the concentration of either a reactant or a product, and any property that is a function of the concentration of a reactant or a product species can be used to evaluate a reaction rate. One important consideration in choosing an experimental technique is that it must be readily adaptable for the rapid determination of concentrations at different time intervals in the course of a reaction, with high precision, ideally of ±1% or so. The different techniques applicable for this purpose are briefly described in the following.

2.1.1  Direct Chemical Analysis The direct analysis of either a reactant or a product is the most obvious method for evaluating the rate of a reaction involving a net chemical change. Exchange processes and also intramolecular processes and configurational changes are inaccessible with this approach. Titrimetric procedures are often convenient, for reasonably slow reactions, as in the replacement of halides and thiocyanate of some metal ions by nucleophilic reagents:   



[PtCl(NH3)3]+ + H2O → [Pt(NH3)3(OH2)]2+ + Cl−

(2.1)

[Co(NCS)(NH3)5]2+ + OH− → [Co(OH)(NH3)5]2+ + SCN−

(2.2)

  



  

  Mechanisms of Reactions of Metal Complexes in Solution By Debabrata Banerjea and M. K. Bharty © Debabrata Banerjea and M. K. Bharty 2023 Published by the Royal Society of Chemistry, www.rsc.org

59

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60

Potentiometric titration of the released chloride with AgNO3 solution, using an Ag/AgCl indicator electrode and a suitable reference electrode (such as a glass electrode) in acetone–water mixture, acidified with nitric acid at ice temperature, is fairly sensitive and convenient. The low temperature and the solvent mixture suppress the direct attack of Ag+ on the chloride bound in the complex to prevent overestimation and a sharper end-­point is obtained owing to the low solubility of AgCl in the acetone–water medium. This method has been widely used in studying the rate of halide replacement reactions of many metal complexes.1 Similarly, in the hydrolysis of a thiocyanato complex, amperometric determination2 of the SCN− released by AgNO3 may be used. The rate of replacement of S2O32− in [Co(S2O3)(NH3)5]+ by various nucleophiles has been followed by microtitration of the released S2O32− iodimetrically, using starch indicator;3 pH–potentiometric titrations using a glass electrode are also applicable in some cases, as in the aquation and base hydrolysis of some metal complexes, the product aqua complexes being weakly acidic (pKa values ranging from 4 to 6 in many cases).†

2.1.2  Spectrophotometric Methods Spectrophotometric methods are very convenient and have been widely used to study solvolysis and other types of ligand replacement reactions and also isomerization and redox reactions of metal complexes. This technique makes use of the characteristic absorptions of the reacting complex and the product. Basically, a spectrophotometer has a means of rendering an essentially monochromatic beam of light of definite intensity and a photoelectric device for measuring the amount of monochromatic light absorbed by the sample placed in the path of the monochromatic beam. Instruments covering the range 200–1000 nm (2000–10 000 Å) are commonly available; the characteristic d–d absorption bands of transition metal complexes lie within this region, whereas normally the common reagents reacting with a complex do not absorb above 250 nm. Wavelengths below 300 nm are generally avoided, since ion association (ion-­pair formation) often causes the appearance of charge-­ transfer bands and consequent ambiguities in the observations below 300 nm. The molar concentration c of an absorbing species can be calculated from Beer's law, which relates the ratio of the intensities of the incident (I0) and the transmitted (I) light to the molar extinction coefficient (εM, cm−1 M−1) at a definite wavelength, its molar concentration c and the pathlength (l, cm) of light in the sample:   

  

I  A log  0  I

  lc  M 

(2.3)

where A is the optical density (molar absorbance) of the sample, which is obtainable directly from the absorption scales of the spectrophotometer. †

 Ka values: [Pt(NH3)5(OH2)]4+, 4; [Pt(NH3)2(OH2)2]2+, 5.6 (cis), 4.3 (trans); [Rh(NH3)5(OH2)]3+, 5.9; p [Co(NH3)5(OH2)]3+, 6.6; cis-­[Co(NH3)4(OH2)2]3+, 6.0; [Cr(NH3)5(OH2)]3+, 5.3; [Ru(NH3)5(OH2)]3+, 4.2.

Techniques for Following Reactions and Factors that Affect Rates

61

To minimize errors, concentrations are so adjusted, or an absorption cell of appropriate pathlength is used, to have optical densities between 0.7 and 0.2 (i.e. ca. 20–60% transmission) against a suitable reference blank (which is generally the solvent placed in an identical cell). Within this range of A values, the error in a measured concentration generally does not exceed ±1%. Before applying this method, it is necessary to find an optimum wavelength for measurement (and determine the extinction coefficients of the reactant and product complexes at that wavelength) by measuring the absorption of the complexes over a range of wavelengths. The absorption spectra thus found will usually differ for the reactant and product complexes having different ligands and also the geometrical isomers of the same complexes. From comparison of the absorption spectra, it is possible to select a suitable wavelength (generally an absorption maximum of either the reactant or the product), for following the reaction, where the extinction coefficients of the reactant and product complexes differ most. Not infrequently both the reactant and products absorb to some extent at the selected wavelength, but provided that no interactions occur between the different species 1, 2, …, etc. the measured absorbance of their mixture will be given by   

A = (ε1c1 + ε2c2 + …)l   

(2.4)

where c1, c2, …, etc. are the respective concentrations of the species 1, 2, …, etc. and ε1, ε2, … etc. are their molar extinction coefficients. Zero-­time and infinite-­time A values can thus be predicted and checked experimentally. For a precision of ±5% in measured rates, the differences in ε values of reactant and product complexes should be at least 50 at 10−3 M concentration, using an absorption cell of 1 cm light path for the measurements. In order for the method to be applicable, Beer's law must hold good for the reaction system at the selected wavelength, which is true in the absence of molecular interactions, such as association, dissociation, changes in ionic solvation, etc. This can easily be tested by measuring the absorbance of synthetic mixtures containing different concentrations of the reactant and product complexes and all other electrolytes in the medium to be used in the actual kinetic study. An especially useful application of spectrophotometry is in the measurement of the rate of change in H+ concentration accompanying a reaction, as a means of measuring the rate of reaction, by incorporating a suitable indicator (for H+) in the reaction mixture, and this is suitable for rapid reactions. Recording the spectra of a reaction mixture over a wavelength region at different time intervals often provides useful information. The collection of a family of spectra as a reaction proceeds may show isosbestic points, indicating that at these points the absorbance remains unchanged despite changes in concentrations of the reactant and product (i.e. composition of the reaction mixture) with time.‡ The occurrence of such isosbestic points during a ‡

 or a reaction X → Y, an isosbestic point will correspond to the wavelength where X and Y have F the same molar absorbance value; hence during the reaction the absorbance does not change with time at this wavelength.

Chapter 2

62

reaction suggests that the original reactant is being replaced by one product (or if more than one product is formed, these are present in the solution, and are formed, in a constant ratio). The absence of appreciable amounts of reaction intermediates is thus implied. An illustration of this is the occurrence of seven isosbestic points in the fairly fast reaction of Hg(TPP) with Zn2+ in pyridine medium forming Zn(TPP) (TPPH2 = meso-­tetraphenylporphyrin) without the formation of a detectable amount of free TPPH2, which has a different spectrum from either of the two complexes.4 In the aquation of cis-­[CrCl(NH3)4(OH2)]2+ in 1 M HClO4, three isosbestic points are observed, for ca. 85% reaction, at 503 nm (εM, 35.5), 433 nm (εM, 10.5) and 368 nm (εM, 28.1). This suggests that the primary product (≥95%) is cis-­[Cr(NH3)4(OH2)2]3+, which has εM values of 35.2, 10.0 and 27.5 at 501, 432 and 368 nm, respectively, indicating that no significant amounts of trans-­ [CrCl(NH3)4(OH2)]2+ or trans-­[Cr(NH3)4(OH2)2]3+ are formed.5 For a first-­order reaction (or a reaction followed under pseudo-­first-­order conditions), the rate constant (k) is obtained by graphical evaluation based on the relation   



 A  A0 log    A  At

 kt   2.303

(2.5)

  

where A0, At and A∞ are the measured absorbance of the solution where the reaction is occurring, initially (at t = 0), after time t and at t = ∞, i.e. on completion of reaction. Evaluation of k ± δk by the method of least squares is of advantage in providing the best value of k with the standard deviation δk. However, if A∞ cannot be evaluated or estimated, evaluation of k by Guggenheim's method6 is useful and this has been used in many of the studies in the senior author's laboratory. This method does not require knowledge of the final equilibrium parameters and also avoids dependence on a single equilibrium value. It is therefore useful for following a first-­order (or pseudo-­ first-­order) reaction in which complications may set in towards the end of the primary reaction, which does not permit observation of the final equilibrium value of the parameter (such as absorbance, conductance, etc.) used in following the reaction. It can be shown that a plot of log[At + ∆t − At] (using A values at times t + Δt and t, respectively, with Δt a constant time interval) versus time (t) is linear with slope −k/2.303. However, for obtaining accurate k values, Δt (arbitrarily chosen) should be equal to at least twice the t½ for the reaction. Roseveare7 gave an excellent account of the mathematical analysis for evaluating rate constants by different methods and a procedure similar to Guggenheim's method for evaluating rate constants for second-­order reactions. This was subsequently generalized by Sturtevant,8 who simplified the calculations. A reaction may be followed spectrophotometrically, either by carrying out the reaction in the absorption cell (thermostated) of the spectrophotometer or by carrying out the reaction in a reaction vessel kept in a thermostat and

Techniques for Following Reactions and Factors that Affect Rates

63

periodically removing samples for subsequent spectrophotometric analysis. For this it is necessary to quench the samples suitably to arrest the reaction (either by cooling or by a suitable chemical change to destroy a reactant, such as acidification to destroy free base when following a base hydrolysis reaction). For studies in situ, a specially designed cell holder is used, in which the cell is kept in a hollow metal block, through which water from a thermostat is pumped rapidly to maintain it at the desired experimental temperature. This technique has been widely used for following the rates of various ligand replacement, redox and isomerization reactions of metal complexes. For such studies, the existing literature is vast, but some typical examples reported from the senior author's laboratory will suffice to illustrate the principles.9–26 Of particular interest is the application of this method and appropriate graphical evaluation procedures to study rates of consecutive reactions for which the rates are comparable and hence the two steps are not well separated.26,27 Spectrophotometry is also particularly useful to obtain valuable information about reaction intermediates.27a,28 Of great interest for labile systems are rapid scanning methods, for which commercial instruments are available, which can scan the spectrum of the reaction mixture every 10 ms or so.27b,29 Some reactions can be conveniently followed by infrared spectrophotometry. Thus, the linkage isomerization reaction  Co  ONO   NH3 5  Cl2  Co   NO2   NH3 5  Cl2 (nitrito isomer) (nitro isomer)

was studied in the solid state by following the absorption at 9.5 µ of the sample in a KCl disc.30 At this wavelength, the O-­bonded nitrito isomer absorbs strongly but the N-­bonded nitro isomer does not.

2.1.3  Electrometric Methods Measurements of conductance, emf, pH and polarographic diffusion current can be used with advantage to follow the rates of reactions in suitable cases. Thus, the rate of the reaction   



[PtCl(NH3)3]+ + H2O → [Pt(NH3)3(OH2)]2+ + Cl−

(2.6)

  

was followed by measuring the increase in conductance of the solution with time.1b The method is not applicable, however, in the case of the reaction   



[PtCl4]2− + H2O → [PtCl3(OH2)]− + Cl−

(2.7)

  

which occurs with no significant change in conductance. However, in the reaction   

  

[PtCl4]2− + OH− → [PtCl3(OH)]2− + Cl−

(2.8)

Chapter 2

64

there is a considerable decrease in the conductance during the reaction, due to replacement of OH− in solution by Cl−, since the former has an ionic mobility ca. 2.6 times that of Cl− (at 25 °C, λOH− = 196 and λCl− = 76), and this has been used with advantage for measuring the rate of this reaction.1b However, for satisfactory results, observations have to be made in a solution that has equimolar concentrations of the complex and OH−, so that the conductance change is significant. Data are then analysed by the usual relations for first-­and second-­order reactions, to see which gives the best fit and thus derive the order of the reaction. The rate of Cl− release can also be followed potentiometrically by measuring with a sensitive potentiometer the emf of a concentration cell of the following type at different intervals of time: Ag/AgCl electrode    + Standard solution of Cl− (medium same as for the experimental solution)

  Salt bridge

Ag/AgCl electrode   + Experimental solution of the complex

This method was used by Selbin and Bailar31 for following the aquation of cis-­[CrCl2(en)2]+. In order for the method to be applicable, the reference electrode must be reversible with respect to the ion whose change in concentration is to be followed. This limits the applicability of this procedure to elimination and substitution reactions involving Cl−, Br−, I−, SCN−, etc. However, in some cases, spurious results may be obtained due to interaction of the reacting complex with the reference electrode. For following reactions of aquated cations, redox potentials may be employed with advantage. Thus, measurement of 2+ the redox potential of an Fe3+ aq − Feaq couple was employed for following the 3+ rate of substitution in Feaq by fluoride ion32 forming Fe2+ aq. Many reactions of complex ions, especially those involving aqua complexes, are accompanied by substantial changes in the pH of the reactant solutions. The aquation of [Co(CO3)(NH3)5]+ is accompanied by an increase in pH due to release of CO2− 3 . The rates of such reactions can be followed by measuring with a sensitive pH meter the change in pH with time, using a glass electrode coupled with a calomel reference electrode. The accuracy of the pH readings should be ±0.001 pH units for precise evaluation of the rate constant. Such precision pH meters are commercially available. The commercial instrument known as a ‘pH-­stat’ automatically adds acid or alkali to a solution so as to maintain the pH constant (to at least ±0.01 unit) during the progress of a reaction, which generates a base or an acid, respectively; here also a glass electrode is used as a pH indicator. Thus, for the reaction   



  

Ni(tmen)22+ → Ni(tmen)2+ + tmen

(2.9)

where tmen = N,N,N'N'-­tetramethylethylenediamine, Me2NCH2CH2NMe2, the rate of dissociation was measured by the rate of consumption of acid

Techniques for Following Reactions and Factors that Affect Rates

65

(corresponding to protonation of the amine being released in the reaction), as the pH is automatically maintained at a constant predetermined value.33 In this particular case, the measured rate was identical with that measured spectrophotometrically. The pH-­stat has considerable promise for the study of reactions of this type. The polarographic method is also suitable in many cases.34–36 In a reaction of the type   

S + R → P



  

(2.10)

the rate can be followed by following the change in the diffusion current (id) due to either the reactant or the product, at a suitable potential, with time, since id is proportional to the concentration of the electroactive species for any particular electrode (i.e. m and t are constants) in accordance with the well-­known Ilkovic equation:   



1

2 1

id  607nD 2 cm 3 t 6

(2.11)

  

where id (µA) is the diffusion current (i.e. the current corresponding to the plateau of the polarographic wave), n is the number of electrons involved in the electrode process, which gives rise to the polarographic wave, c (mM) is the concentration of the electroactive species, D is the diffusion coefficient of the electroactive species under the experimental conditions, m (mg) is the mass of mercury flowing out through the capillary of the dropping mercury electrode (DME) per second and t (s) is the drop time (i.e. the time for formation of one drop of mercury flowing out of the capillary of the DME). The polarographic method is also of utility in evaluating the rate of generation of an electroactive species (B), that gives rise to a polarographic wave, from an inactive species (A) in solution:   



k1



 ne A#B   Product fast k2

(2.12)

  

Under this condition, the polarographic wave is controlled by the rate of generation of B and is a kinetic wave. There are various criteria for a kinetic wave that help in its identification. For such a system, the limiting current, ik, is related to k1 by the following relation derived by Koutecky:37   



1 ik  0.886  Kk1t  2 id  ik

(2.13)

  

where id is the theoretical value of the limiting current (as given by the Ilkovic equation) had there been a direct electroreduction of A at the DME by a purely diffusion-­controlled process, K is the equilibrium constant for transformation of A into B (K = [B]/[A] = k1/k2) and t is the drop time.

Chapter 2

66

A more general situation would be the following, where all the higher species MXN to MXN−n are electroinactive, whereas some lower species are electroactive:   





(2.14)

  

From an analysis of the kinetic current, it is of course not possible to ascertain which of the lower species actually undergoes reduction. For this situation and under conditions that Kj[X] ≫ 1 for all j > 0 (i.e. j = 1, 2, …, N):   



1 1 ik  0.886  kdt  2 K N K N 1  K N  n  X N ( N  n )  2 id  ik

(2.15)

  

where KN, KN−1, …, etc. denote the stepwise stability (formation) constants for the consecutive complexation stages. The species whose dissociation controls the overall rate may then be determined from a plot of log[ik/(id – ik)] versus log[X], since from eqn (2.15) the slope of this log–log plot is (N − n) − N − ½ = −(n + ½). Cadmium(ii) in cyanide medium has been found to produce a kinetic wave for which −(n + ½) is −1.5. Hence, since N is known to be 4 in this case and n = 1, the species whose dissociation controls the current is [Cd(CN)3]−, and therefore from the analysis of the kinetic wave the rate constant kd for the dissociation of [Cd(CN)3]− to [Cd(CN)2] could be evaluated.38 At 25 °C (ionic strength 0.1 M), kd was found to be 8.0 × 105 s−1. This polarographic technique thus affords a means of evaluating the rate constants of fast reactions. Other techniques for studying fast reactions are discussed in Section 2.2.

2.1.4  Polarimetric Methods Optical rotation measurements enable the rate of racemization of an optically active complex to be studied.39 The same technique has also been used to study the rate of the following electron transfer reaction:   

l-­[Os(bipy)3]2+ + d-­[Os(bipy)3]3+ ⇌ l-­[Os(bipy)3]3+ + d-­[Os(bipy)3]2+

(2.16)

  

Here, advantage was taken of the fact that the molar rotations are different for the complexes corresponding to the two different oxidation states of osmium.40 A possible source of error in the polarimetric method is that the light used in optical rotation measurements may accelerate isomerization processes. Thus, the Na D-­line (5890 Å) accelerates the loss of optical activity of l-­cis-­ [CoCl2(en)2]+ in methanol.41

Techniques for Following Reactions and Factors that Affect Rates

67

2.1.5  Use of Isotropic Tracers Apart from their routine use, in following metal exchange and ligand exchange reactions in metal complexes, and also in following the rate of electron transfer between two different oxidation states of a metal {such as in the [Fe(CN)6]4− − [Fe(CN)6]3− system}, isotopic tracer techniques yield much additional and valuable information concerning the fine details of the reaction mehanism.42 A number of radioactive isotopes of different elements and also inactive isotopes such as 15N, 17O, 18O, etc., are useful for this purpose. Examples of some typical applications can be found in several of the following chapters.

2.2  Special Techniques for Studying Fast Reactions Several methods are available for following the rates of fast reactions and a few commonly used methods are discussed below.

2.2.1  Flow Methods43 A reaction rate cannot be measured with the techniques mentioned in the preceding section if most of the reaction occurs while the reactants are being mixed. In order to be able to follow the progress of a reaction, only an inappreciable amount of reaction should occur during the time needed for an observation of the extent of reaction. There are procedures for overcoming these limitations and reactions that go to completion even in a few milliseconds can be studied using flow methods, and this technique is useful for studying reactions having t½ in the millisecond range. In the constant-­flow method, equal volumes of solutions of the two reactants are forced at high speed into a mixing chamber and from there the mixed solution flows through an observation tube (Figure 2.1). With proper design, the two solutions can be thoroughly mixed in the mixing chamber in times as short as 10−3 s. With a constant flow rate, each position along the observation tube corresponds to the lapse of some definite time interval after mixing, i.e. after initiation of the reaction. Hence the composition of the mixed solution at each position along the observation tube corresponds to a particular extent of reaction corresponding to the appropriate time interval and therefore, with a constant flow rate, steady concentrations of reactants and products are established along the observation tube, i.e. concentrations at each point along the observation tube are held constant. Hence in this arrangement, one can detect the steady-­state concentration of either one of the reactants, or of the products, by using any suitable device making use of optical,44 conductance45 or potentiometric46 measurements and no fast response device is needed. For following the rate of the reaction, one has to make observations at different fixed points along the observation tube, so as to have concentration data for different reaction times. Alternatively, observations can be made at

Chapter 2

68

Figure 2.1  Schematic  diagram of the constant-­flow apparatus with photometric

observation. A, B, reservoirs for storing the solutions of the two reactants; S1, S2, syringes connected to reservoirs A and B through stopcocks; P1, P2, pistons, tightly fitting the syringes S1 and S2 on which pressure can be applied to drive the solutions at a constant flow rate; C, mixing chamber to ensure rapid mixing of the two solutions; T, observation tube.

a fixed point in the observation tube, to have concentration value for a fixed reaction time corresponding to the flow rate used. Data for different reaction times are then obtained by changing the velocity of flow, by changing the speed at which the reactants are forced through the observation tube. Since reactions often occur with appreciable evolution or absorption of heat, the temperature of the solution along the observation tube of a constant-­flow apparatus will vary. Hence accurate measurements of temperature at various positions along the observation tube by using an appropriately designed sensitive probe containing a thermocouple also allow the evaluation of the rate of the reaction, as has actually been used47 in the following case, for which the rate constant thus obtained was 1.04 × 103 M−1 s−1 at 25 °C:   



CO2(aq) + OH−(aq) → HOCO−2(aq)

(2.17)

  

A diagrammatic representation of the equipment with photometric detection is shown in Figure 2.1. In the stopped-­flow method, reactants are mixed as in the constant-­flow method, but the flow of the reacting liquid is suddenly and instantaneously arrested. At this instant, the concentration change of a reactant or a product

Techniques for Following Reactions and Factors that Affect Rates

69

Figure 2.2  Schematic  representation of a stopped-­flow assembly. The flow stops with the stop plate striking the microswitch trigger and thereby activates the detector/recorder to monitor the change in absorbance with time.

is measured in the observation cell with a fast-­response device. Fast-­response spectrophotometric devices are very convenient for this purpose48 and many such sensitive instruments are commercially available. A schematic representation of such an instrument is shown in Figure 2.2. In the quenched flow method, after rapid mixing the flowing reactants are discharged into a quenching solution to arrest the reaction, and this solution is then analysed by a convenient accurate standard analytical procedure. The limiting feature of this method is the rate at which quenching can be accomplished. For studies using isotopic tracers the use of the quenched flow method is necessary, as in the MnO−4 − MnO2− 4 and Ag(i)–Ag(ii) exchange systems.49

2.2.2  Electrochemical Methods In these methods, an equilibrium existing in the solution is disturbed by removing one of the components of the system in equilibrium using a suitable electrode process, and a steady state is thereby established, in which the measured current or voltage is determined by the rate at which the equilibrium is re-­established. The two electrochemical techniques

Chapter 2

70

most commonly used for this purpose are polarography and voltammetry at constant current. The latter uses platinum as the working electrode, so that processes that occur at potentials at which mercury (which is used in polarography) is oxidized can also be studied. First-­order rate processes can be studied conveniently by these techniques and hence a pseudo-­first-­ order condition is maintained in the reacting solution, using a large excess of one reactant as in the oxidation of Ti(iii) by a large excess of NH2OH studied polarographically.50 For applications of the polarographic method, see also Section 2.1.3. An automatic coulometric apparatus for determining the rate of base hydrolysis has been described.51 Here OH− is generated by an electrochemical process involving the reaction   



2Ag + 2Br− + 2H2O → 2AgBr + H2 + 2OH−

(2.18)

  

The OH− is generated at a rate which, with a known excess of the other reactant present, equals the rate of base hydrolysis. Hence the rate constant for the base hydrolysis can be calculated from the steady-­state current that is used to generate the OH−.

2.2.3  Relaxation Methods52 In these methods, the position of a chemical equilibrium is distributed by a very rapid change of a physical parameter, such as temperature or pressure, which affects the position of equilibrium and the change (relaxation) of the system to the new equilibrium position is followed. For this the photometric method is often used and the photomultiplier response of the instrument is traced on an oscilloscope, which is photographed using a camera for a permanent trace or recorded otherwise. A temperature jump (T-­jump) of 5–10 °C can be achieved rapidly (within 10−10 s) by discharging a capacitor through the reacting solution. Similarly, a pressure jump (P-­jump) of ∼100 atm within a very short time allows the measurement of the rate of the reaction H+ + OH− → H2O, for which the rate constant is 1.5 × 1011 M−1 s−1 (25 °C), and this is the fastest reaction known in solution. The significance of a rate constant of this magnitude is that 0.1 M solutions of strong acid and strong base, if mixed instantaneously, would be 50% reacted in ca. 10−8 s and ca. 99% reacted in ca. 10−6 s. The P-­jump method is obviously applicable for reactions that occur with a change in volume generally due to a change in the number of species in the reaction. However, the P-­jump technique has also been used for the reaction (where there is a change in molar volume) using a conductance method for monitoring the change:   



[Fe(OH2)6]3+ + Cl− → [FeCl(OH2)5]2+ + H2O

(2.19)

  

Relaxation methods are useful for reactions having rate constants in the range 106–1010 M−1 s−1 (t½ in the µs range), and applications of relaxation

Techniques for Following Reactions and Factors that Affect Rates

71

techniques to study complex formation in labile aqua metal ion–ligand systems were described by Eigen,53 who was awarded the Nobel Prize in Chemistry in 1967 for this work. An ultra-­fast T-­jump method with a time resolution of the order of 5 ps for the study of different types of reactions in aqueous solution, including water exchange reactions of aqua (H2O) ligands in metal complexes, has been reported.54

2.2.4  Flash Photolysis55 When iodine vapour is illuminated with light of appropriate wavelength, the dissociation reaction I2 + hν → 2I occurs. The absorbed photon is represented by hν, where h is Planck's constant and ν is the frequency of light in s−1. Rapid recombination of iodine atoms occurs in a non-­photochemical reaction, 2I → I2, which consumes the iodine atoms produced by the photochemical dissociation reaction. Under steady illumination with light of moderate intensity, a steady state is established with only a very small fraction of iodine remaining dissociated to atoms. This steady state results from equality of rates of the photodissociation of I2 and non-­photochemical recombination of I atoms. If illumination occurs for a very brief period with a very intense flash, an appreciable number of iodine atoms can be produced, and their recombination can be directly observed after the flash. In actual experiments, the peak intensity of a flash may be built up in 10−5 s and then decay over a period of ∼10−4 s. After light intensity from the flash has decayed, light of low intensity can be used in a spectrophotometric arrangement, to follow the increase in concentration of molecular iodine with time, which can be traced on an oscilloscope. This technique, called flash photolysis, has provided much valuable information regarding the rates of very fast reactions involving free radicals, which can be generated by photolysis both in the gas phase and in solution.

2.2.5  N  uclear Magnetic Resonance (NMR) and Electron Paramagnetic Resonance (EPR) Methods56 These techniques make use of the Zeeman splitting of nuclear and electron energy levels, respectively, which depend on the chemical environment of the magnetic moments. If the magnetic field is in the kilogauss region, the resonance frequency is of the order of megacycles per second for nuclear moments and kilomegacycles per second for electron spin moments. Hence NMR is observed in the radiofrequency range and EPR at microwave frequencies, and commercial instruments are available for both techniques. Techniques based on these phenomena are extremely useful in following reactions having second-­order rate constants ranging from ca. 104 to 1010 M−1 s−1. NMR is exhibited by an isotope of any element having a non-­zero nuclear spin, i.e. its nuclear spin quantum number I is not zero, hydrogen (I = ½) being the most important example. Proton magnetic resonance (PMR) techniques are very useful, not only in proton exchange rate studies but also in

Chapter 2

72 31

structural elucidation. Other usable isotopes of non-­metals are P (I = 1/2), 19 F (I = 1/2), 15N (I = 1/2), 14N (I = 1) and 17O (I = 5/2). 12C and 16O, both having I = 0, are not NMR active. NMR spectroscopy is particularly suitable for studies of exchange reactions, such as using 17OH2 to study the exchange of aqua metal ions with the bulk solvent (water). Electron paramagnetic resonance or electron spin resonance (EPR or ESR) is exhibited by any species having an unpaired electron. For experimental investigations of complex ions in solution, a useful NMR signal is observed with conventional apparatus if the isotope of the element under study has a large nuclear magnetic moment, i.e. a nuclear spin of 3/2 or greater, and an isotopic abundance greater than 10%. The characteristic absorption frequency should also be high. With transition elements, usable isotopes include 51V (I = 3/2), 53Cr (I = 3/2), 57Fe (I = 1/2), 59Co (I = 7/2), 61Ni (I = 3/2), 63Cu (I = 3/2), 65Cu (I = 3/2), 93Nb (I = 9/2), 185Re (I = 5/2), 187Re (I = 5/2) and 193Ir (I = 3/2). Since hydrogen is readily detected by NMR spectroscopy and is a common constituent of most ligands, it is useful not only in studying proton dissociation (or proton exchange) rates, but also for studying rapid redox reactions, provided that suitable chemical shifts occur in the complexes. However, to obtain an adequate signal response, an NMR study of a solution reaction will require the use of concentrated solutions often exceeding 1 M, and this obviously leads to difficulties in interpretation. ESR spectroscopy is a more sensitive technique and concentrations lower than 0.1 M can usually be employed. However, experimental difficulties limit the application of the ESR technique to the study of paramagnetic ions in solutions. Redox reactions can be studied by this technique provided that one complex is paramagnetic and the other diamagnetic. The advantage of the NMR and ESR techniques is that the measurements are performed on systems in chemical equilibrium. Moreover, it is often possible to identify directly the atoms that are actually involved in the chemical process.

2.2.6  Studies of Exchange Rates by NMR Spectroscopy57 The principle of this technique can be illustrated by considering a hypothetical case of proton exchange between two sites that are equally populated and have identical lifetimes as in a mixture of A–H and B–H molecules. The position of the NMR line due to the proton would occur in different frequency regions in A–H and B–H owing to the different environments of the proton in the two molecules. In the absence of any exchange of proton between the molecules A–H and B–H, two separate proton NMR lines (signals) would appear in a mixture of A–H and B–H (see Figure 2.3a) which are identical in position and width as for pure A–H and B–H. In the case that there is exchange, however, the two signals would increase in width and overlap gradually more and more with increasing rate of exchange (see Figure 2.3b–d) followed by their coalescing (Figure 2.3e) and finally, at high exchange rate, narrowing of the single signal occurs (Figure 2.3f). This behaviour has been observed in the typical example of [Pt{P(OEt)3}4]–P(OEt)3 exchange.58

Techniques for Following Reactions and Factors that Affect Rates

73

Figure 2.3  Change  in the nature of NMR spectra with change in the rate of exchange

of an atom between two sites (such as proton exchange between A–H and B–H). (a) Extremely slow exchange (no exchange); (b)–(e) increasing rate of exchange; (f) extremely rapid exchange.

Under the conditions in Figure 2.3b–d, the lifetime of the proton in A, τA, which is equal to the lifetime of the proton in B, τB, can be obtained from

  

1



  



2  A  B obs  1    1 2 2 0 2  A0  B0  2π   A  B0  

 

(2.20)

1

2 1  0 2  A  B0    A  B obs  2    2

(2.21)

  

where ν0A − ν0B is the separation of the proton NMR peaks of A–H and B–H in the absence of exchange (as in Figure 2.3a) and (νA − νB)obs is the separation of the peaks that have begun to coalesce (as in Figure 2.3b–d) due to exchange and τ = τA/2 (values of νA and νB are in cps). Thus, from observations of the NMR spectra, it is possible to evaluate τA and from this the value of the rate constant, since τA is related to the rate constant. For a first-­order rate equation, d[A]/dt = k[A],   

  

1/τA = k or 1/τ = 2k

(2.22)

whereas for a second-­order rate equation, d[A]/dt = k[A][B],   



  

[A]/τA = [B]/τB = k[A][B]

(2.23)

1/τA = k[B]

(2.24)

1/τB = k[A]

(2.25)

Hence   



  

and   

  

Chapter 2

74

In a typical study, the NMR spectrum of a pure substance A–H is first recorded, followed by a series of observations of mixtures of A–H and B–H containing increasing amounts of B–H, so as to have observations on a series of mixtures having increasingly faster rates of exchange, corresponding to increasing concentration of B–H in the mixture. From eqn (2.21), it follows that with increasing rate of exchange when the two peaks just merge (as in Figure 2.3e), (νA − νB)obs = 0 and eqn (2.21) transforms to   





  

2 2π  A0  B0 

(2.26)

From eqn (2.21) and (2.26), it follows that for two separate peaks to be observed at frequencies νA and νB the lifetimes of the two states (A–H and B–H) must be greater than 1/(νA − νB). For smaller values of τ, resulting in the merged peaks (as in Figure 2.3e), eqn (2.26) can be used to calculate the lifetime. Under this condition of a single broad signal, the width of the signal at the peak half-­height also allows the evaluation of τ and hence of k (see ref. 28, p. 166). For extremely fast exchange (very small τ value), a single sharp peak is observed, as shown in Figure 2.3f, and measurement of such a fast rate is outside the range of the NMR method. However, when the exchange rate falls outside the limit of measurement by this technique, the rate may be measured if conditions (temperature, concentration, etc.) are changed to alter the rate sufficiently to bring it within the range of the NMR method. Under the condition of a slow rate of exchange, the signals due to A–H and B–H may appear separately but are broader than in the absence of exchange (exceedingly slow rate of exchange). For this situation of NMR line broadening, the difference in the width of the signal at the peak half-­height between the exchange-­broadened signal WEA–H and the signal W0A–H in the absence of exchange (i.e. in A–H without added B–H), and the broadening is still much smaller than their separation and the width expressed in cps or hertz, then in this slow exchange situation, WEA–H − W0A–H = 1/πτA where 1/τA = kA (kA being the first-­order rate constant for the exchange of H in A–H); rate = kA[A–H]. Similarly, for the signal due to B–H,

WEB–H − W0B–H = 1/πτB;  1/τB = kB; rate = kB[B–H]

 B

A 

[A  H] kB  [B  H] kA

This slow exchange region has been most useful in studying the rates of complex-­ion reactions (see ref. 28, p. 165). There have been numerous applications of the NMR technique. The electron exchange rate in the Cu(i)–Cu(ii) system in HCl medium has been studied59 (making use of the 63Cu NMR signal), for which the rate constant

Techniques for Following Reactions and Factors that Affect Rates

75

−

Table 2.1  Kinetic  parameters for exchange of Cl and H2O in oxovanadium(iv) species using 17O and 35Cl NMR spectroscopy.

Complex

Exchanging ligand

k/s−1 (27 °C)

Ea/kJ mol−1

[VO(OH2)4]2+ [VO(Cl)(OH2)3]+ [VO(Cl)2(OH2)2] [VO(Cl)(OH2)3]+ [VO(Cl)2(OH2)2]

H217O H217O H217O 35 − Cl 35 − Cl

(5.9 ± 0.2) × 102 (2.3 ± 0.6) × 104 (8.5 ± 0.3) × 105 (2.0 ± 1.6) × 103 (2.3 ± 0.2) × 105

62.8 ± 4.2 46 ± 4.2 41.8 ± 4.2 50.2 ± 8.3 29.3 ± 4.2

is of the order of 108 M−1 s−1.17O NMR spectroscopy has been used to study the rate of water exchange in [M(OH2)6]n+ [see Chapter 1, Figure 1.3] and [M(OH2)4]2+ (M = Pd, Pt),60 and 195Pt NMR spectroscopy has been used to study the formation of [Pt(OH2)4−xA](2−x)+ in aqueous solutions of Pt(ii) con− − 61 taining various anions, Ax− (Cl−, NO−3, SO2− Using 4 , H2PO 4, CH3CO 2, etc.). 17 35 O and Cl NMR spectroscopy, the kinetic parameters for exchange of Cl− and H2O in oxovanadium(iv) species have been obtained and the data are given in Table 2.1.62 NMR line broadening has been used to measure the kinetics of exchange of glycine in [VO(Hgly)]2+ with free glycine in aqueous solution, and also the kinetics of formation of [VO(Hgly)]2+ (the zwitterion of Hgly behaving as a monodentate ligand) by the stopped-­flow method.63 [M(nta)2]3−–nta3− exchange in lanthanides (MIII = La, Ce, Pr, Nd, Sm, Eu, Tm, Yb, Lu and also Y) were studied64a as a function of pH by 1H NMR spectroscopy. Three kinetic pathways were observed (see Chapter 4). Ligand exchange in metal β-­diketonates has been studied by NMR spectroscopy.64b

2.2.7  Electron Paramagnetic Resonance (EPR) Spectroscopy EPR line broadening is a technique that is useful in limited cases to study the rates of reactions with t1/2 ≈ 10−4 – 10−10 s, i.e. almost up to the limit for a diffusion-­controlled process in solution. However, its most useful application is for the detection of free radical intermediates in oxidation– reduction reactions. This technique has been employed in the following illustrative case for evaluating rate constants for reactions of pyridine and pyridine derivatives:65 k VO  acac 2  py   VO  acac 2  py 

k = 1.0 × 109 M−1 s−1 (25 °C). The method has also been used to study the rate of solvent exchange in the [VO(DMF)5]2+–DMF (solvent) system and the consecutive steps in the acid-­ catalysed aquation of [Cr(CN)2(NO)(OH2)3]+ (CN− lost as HCN).66 The following electron transfer reaction has also been studied by this method:67a

[61Ni(cyclam)]2+ + [58Ni(cyclam)]3+ ⇌ [61Ni(cyclam)]3+ + [58Ni(cyclam)]2+

Chapter 2

76

This makes use of the fact that the Ni(iii) complexes display EPR signals; those of 61Ni show a splitting of the g∥ feature due to its nuclear spin I = 3/2, but this is missing in the corresponding Ni(ii) complex. Kinetic studies on the oxidation of VO2+ by H2O2 using the ESR technique have been reported:67b k1

k2 VO2   H2 O2 #OVOOH ;OVOOH   VO2  OH k1

Fast VO2+ +HO  VO2 +H+

At 25 °C, k1 = 52 M−1 s−1, K1k2 = 0.21 s−1 (K1 = k1/k−1).

2.2.8  Application of Ultrasonic Absorption and Pulse Radiolysis Application of ultrasonic absorption in studies of the kinetics of complex formation has been reviewed.68a This is particularly useful for determining the characteristic time of each of the elementary steps in the overall formation of the complex. Another useful technique is pulse radiolysis, of use in the generation of solvated electrons, free radicals and metal ions in unstable oxidation states to study their reactions that include many redox reactions.68b,c

2.3  Rate Constants in Some Complex Systems69 .   Concurrent irreversible reactions: 1   

For this,   

  

 d A    Rate  k  A   k1  k2  A     dt 

(2.27)

If the amounts of B and C produced can be estimated experimentally, then evaluation of k1 and k2 is possible from eqn (2.27) and the following relation:   



  

k1/k2 = [B]/[C]

(2.28)

2.  Monophasic reversible reactions:    k1  B A   k-1 For this, the observed rate constant for approach to the equilibrium state will be kobs = (k1 + k−1). By evaluating the equilibrium constant K = k1/k−1, it is possible to evaluate the values of k1 and k−1. In a case such as

Techniques for Following Reactions and Factors that Affect Rates

77

k1

  AX A  X  k1

the system will be pseudo-­first order in the forward direction also if [X] ≫ [A], and kobs = k1[X] + k−1. Hence, by plotting kobs versus [X], k−1 and k1 are obtained from the intercept and slope, respectively, of the linear plot. The following reaction is an example:70 3

1    V  SCN   OH2    V  OH2 6   SCN    5  k1

k

2

 H2 O

   3.  Biphasic irreversible system:    This corresponds to two consecutive irreversible reactions, viz. k1 k2 A   B  C

Although for the concurrent process the ratio [B]/[C] is constant, independent of reaction time [eqn (2.28)], for the consecutive process as also for the mixed process A → C along with concurrent A → B → C, the ratio [B]/[C] is a reaction time-­dependent variable. Radioactive decay series are examples of multi-­step irreversible consecutive reactions. If k1 ≫ k2, then both steps A → B and B → C can be analysed separately using, e.g., the usual absorbance versus time data for the initial and final stages of the reaction. If, however, k2 ≫ k1, then   





d  A  d  C   k1  A  dt dt

(2.29)

  

permitting the evaluation of k1 only. In such a case, k2 can be evaluated only if B can be prepared and studied. This is also helpful for systems in which k1 ≈ k2, as in the following case:71 3

H3 O  Cr  en   OH2 4     Cr  enH   OH2 5  k1 

4

3

H3 O  Cr  OH2 6   enH22   k2 3

Here k1 and k2 are within a factor of 2–4 under the experimental conditions. If B is not isolable (or cannot be prepared in solution), a difficult situation arises when k2 ≫ k1, unless one of the steps is multi-­order, since then the associated first-­order rate constant is modified by changes in reagent concentration. Such a situation in the latter case arises in the acid-­catalysed hydrolysis of [Co(CO3)L4]+ [L4 = (NH3)4, (en)2, tren, etc.]:72 

k  k  H 

0 H   CoL 4  HCO3   OH2    CoL 4  CO3    H3 O

3

2

k2    CoL 4  OH2 2   HCO3 H2 O

Chapter 2

78 +

+

Here k1 = k0 + kH[H ] and is rate determining, except at [H ] ≥ 2 M for L4 = tren, when k0 + kH[H+] ≫ k2.73 Methods for evaluation of k1 and k2 for the system k1 k2 A   B  C

where k1 > k2 have been reported extensively.25,26,71,74 Maximum time and time-­ratio methods have also been used to determine rate constants for such consecutive (successive) reactions.75 Another system of interest is the following:76 k1 k2 A   C  B

Denoting initial concentrations of these species as A0, B0 and C0 and analogously at time t and on completion of reaction (t = ∞), we have  k1t  A A , B B0 e k2t 0e

C C  A0 e k1t  B0 e k2t t

where C∞ = A0 + B0, and   



  

log  C  Ct   log  A0 e k1t  B0 e k2t 

(2.30)

If k1 = k2 or only one reactant A or B is present, then a plot of log(C∞ − Ct) versus t would be linear. If both are present and k1 ≠ k2, then the plot is curved, becoming linear at longer time (see Figure 2.4) when the more reactive species, say A, has disappeared. Now, log(C∞ − Ct) = log Bt = log B0 − (k2t/2.303)

Figure 2.4  Graphical  evaluation of k2 (see text).

Techniques for Following Reactions and Factors that Affect Rates

79

Hence, from the slope and intercept of the plot shown in Figure 2.4, k2 and B0 are evaluated, hence Bt can be calculated at different times t: log Bt = log B0 − (k2t/2.303) With this information At can be calculated: At = C∞ − Ct − Bt A plot of log At versus t permits the evaluation of A0 and k1: log At = log A0 − (k1t/2.303) For both concurrent and consecutive reactions mentioned above, there may be more complex situations with one or both the steps being reversible.    4. Competition methods (see ref. 28, 1st edn, Chapter 3, p. 165):    It is often useful to have relative rate constants for a series of reactants acting on a common substrate and this can be achieved by competition methods avoiding the kinetic approach needed for evaluating individual rate constants. This is a decided advantage, particularly for fast-­reacting systems for which evaluation of individual rate constants requires the use of sophisticated instruments. An illustration of the method is the second-­ order reactions of the two Co(iii) complexes [Co(X)(NH3)5]2+ (X = Cl, Br) with [Cr(OH2)6]2+ forming [Co(OH2)6]2+ and [Cr(X)(OH2)5]2+.77a,b In the actual procedure, a solution containing a mixture of the two Co(iii) complexes at known concentrations is added very quickly to a vigorously III stirred solution of the Cr(ii) present in deficiency, i.e. [CrII]0 < [CoIII A ]0 + [Co B ]0 (the subscripts A and B denote the two complexes and the subscript 0 refers to the initial concentrations of the reactants). Owing to the use of Cr(ii) in deficiency there will be some unreacted Co(iii) complexes in the solution on III completion of the reaction and their concentrations [CoIII A ]c and [Co B ]c can be estimated by a suitable procedure. The rate constant ratio kB/kA can then be evaluated using the following relation: III III III log([CoIII B ]c/[Co B ]0) = (kB/kA)log([Co A ]c/[Co A ]0)

This can be easily derived from the following rate expressions: III II −d[CoIII A ]/dt = kA[Co A ][Cr ] III II −d[CoIII B ]/dt = kB[Co B ][Cr ]

The rate constant ratios obtained by this method for such reactions agree fairly well with the values obtained by the direct evaluation of the individual rate constants.77c–e

Chapter 2

80

2.4  Factors that Affect Rates of Reactions 2.4.1  Effect of Temperature The value of the rate constant k is dependent on the Arrhenius frequency factor, A, and the energy of activation, Ea, according to the well-­known Arrhenius relation:78a   

k  Ae Ea / RT



(2.31)

  

Since the value of k depends on both A and Ea, knowledge of these two factors is necessary in any comparison of k values for different systems. A and Ea can be evaluated from measured values of k at several (three or more) temperatures. According to the collision theory of reaction rates, A = PZ, where P is the probability or steric factor and Z is the collision number. Since P cannot be related to a recognizable physical quantity, the collision theory has largely been replaced by the transition state theory,78b especially when considering reactions in solution. According to this theory, the reactants in a chemical reaction are in equilibrium with a species of sufficient energy, called the activated complex, in which the reactants are in a state for formation of the products. The region of occurrence of the activated complex in the change from reactants to products along the reaction coordinate is known as the transition state (see Figure 2.1), which represents the region of highest energy in the path from one stable system to another. Quantitatively, the rate of passage from one stable form to another can be derived from a statistical thermodynamic approach. For a general bimolecular reaction,   

A + B ⇌ X‡ → C + D

(2.32)

  

where X‡ represents the activated complex, its equilibrium concentration is given by§   

[X‡] = K‡[A][B]



(2.33)

  

where K‡ is the equilibrium constant for the formation of the activated complex. Hence  f f  X ‡   K ‡  A B A B  f ‡  X

where a-­is activity and f is activity coefficient. §

Actually,

 K‡

aX ‡ f ‡ [X]‡   X aA aB [A][B] f A fB

  

Techniques for Following Reactions and Factors that Affect Rates

81 ‡

‡

In very dilute solutions (I ≤ 0.01 M), the f values are ∼1, hence [X ] = K [A][B]. At constant ionic strength of the solution, f is constant and this relation is applicable with the fAfB/fX‡ term incorporated in K‡. If the solution is not very dilute (I > 0.01 M), then by incorporating the f values (activity coefficient terms) we obtain

  



 f f kT  Rate d P  /d t   K ‡  A B A B  h   f X‡

  

(2.34)

  

and accordingly, at any particular temperature,   

kT   f A fB k2  K ‡    h   f X ‡



  f f   k0  A B   f X‡

  

(2.35)

  

where k0 = k value at I ≈ 0. The rate of reaction will be given by the rate of formation of a product, d[P]/ dt, and this is proportional to the concentration (not “activity”) of X‡ and can be shown to be equal to (kT/h)[X‡]. Substituting in this the value of [X‡] from eqn (2.33), we obtain   



d[P]/dt = (kT/h)[X‡] = (kT/h)K‡[A][B]

(2.36)



= k[A][B]

(2.37)

k = K‡(kT/h)

(2.38)

  

  

  

where k is the second-­order rate constant (in M−1 s−1), k is Boltzmann's constant (= R/N) and h is Planck's constant. K‡ in‡ eqn (2.36) can be replaced by the well-­known thermodynamic function eG / RT , leading to   



  

d P  /d t   kT h  eG

‡

  kT h  eH

  

‡

/ RT

/ RT

[A][B] eS

‡

/R

 A B

(2.39) (2.40)

where ΔG‡, ΔH‡ and ΔS‡ are the standard free-­energy of activation, enthalpy of activation and entropy of activation, respectively, and represent the change in free energy, enthalpy and entropy, respectively, in the formation of the activated complex (transition state) from the reactants. ΔS‡ is related to the pre-­exponential term A of the Arrhenius relation, eqn (2.31); correct ‡ A   ekT / h  eS / R {hence ΔS‡ = R[ln A + ln(h/kT) − 1]}¶. The energy of activation, Ea, is greater than ΔH‡ and for a reaction in solution Ea = ΔH‡ + RT. However, ΔH‡ is usually of the order of 40–120 kJ mol−1, compared with which RT

At 25 °C, ΔS‡ = 4.6(log A – 13.2) cal K−1 mol−1.

¶

Chapter 2

82 −1

is negligible (at 25 °C, RT = 2.5 kJ mol ) and is commonly ignored. From eqn (2.37) and (2.40), it follows that   



  

k2   kT h  eH

‡

/ RT

eS

‡

/R



(2.41a)

Hence   



  

−log(kh/kt) = ΔH‡/2.3RT + ΔS‡/2.3R

(2.41b)

This allows the graphical evaluation of the ΔH‡ and ΔS‡ values. Using appropriate values of h (6.626 × 10−34 J s), k (1.381 × 10−23 J K−1), R (2 cal K−1 mol−1) and k (s−1, M−1 s−1 or M−2 s−1 as applicable), the ΔH‡ and ΔS‡ values obtained are in cal mol−1 and cal K−1 mol−1, respectively. The enthalpy change in the formation of the products from the transition state, called enthalpy of transition, ΔHT, which is equal to ΔH − ΔH‡, where ΔH is the enthalpy of reaction, i.e. the enthalpy change in the transformation of reactants to products. This ΔHT is an important parameter the value of which is diagnostic of the mechanism (see Chapter 4). The enthalpy of activation will generally include a significant electrostatic contribution due to the energy required in bringing charged reactants up to each other, and also the energy required for internal reorganization of the bonds within the reactant molecules, solvation energies of the reactants and products, etc. A substantial difference in activation energies is generally observed for comparable reactions of analogous complexes of different transition elements, due to differences in their bond energies. However, for similar complexes of the same element, different rates usually arise due to difference in ΔS‡ values. If a pre-­equilibrium precedes the rate-­determining step, then the ΔH value for this equilibrium will also be included in the observed ΔH‡ value for the reaction. Since similar rates are often observed in different systems due to changes in ΔH‡ and ΔS‡ balancing each other's effect on the rate, knowledge of ΔH‡ and ΔS‡ is necessary in any comparison of the behaviours of different systems. For analogous reactions in different comparable systems, proceeding by a similar mechanism, a linear relationship is observed between ΔH‡ and ΔS‡ (called the isokinetic trend):78c   



ΔH‡ = α + βΔS‡

(2.42)

  

A significant deviation from this linear relation indicates a difference in mechanism. In eqn (2.42), β is the isokinetic temperature (in K) at which all the reactions, which conform to the particular linear relationship, would occur at the same rate and α is the corresponding free energy of activation ΔG‡. Below this temperature, the relative rates of the reactions are controlled by ΔH‡, and above this temperature they are controlled by the ΔS‡ values. Isokinetic trends have been observed in redox reactions such as the oxidation of Ru(ii) complexes of bipyridine, 1,10-­phenanthroline and their derivatives by Tl(iii)79 and redox reactions of actinide ions,80 in ligand replacement reactions of complexes of Pt(ii)81 (see Chapter 3), in anation reactions of [Cr(OH)(OH2)5]2+ (involving replacement of an H2O ligand by an anion, see

Techniques for Following Reactions and Factors that Affect Rates 82

+ and replacement of S2O2− 3 in [Co(S2O3)(NH3)5]

83 −

Figure 2.5) by OH (base hydrolysis) in aqueous solution and solutions in mixed solvents (H2O + ethanol and H2O + acetone);83 the rate of replacement of the S2O32− in the complex by Cl− and NH3 in aqueous solution also conforms to the same trend (see Chapter 4). An isokinetic relationship has also been made use of in elucidating whether in the dissociation of a chelating ligand from a metal complex the chelate ring opening or the subsequent loss of the chelating ligand is rate determining24 (see Chapter 5, Figure 5.2). For reactions in solution, the entropy of activation measures the total change in entropy of the reactants and the solvent on formation of the activated complex (transition state). The sign and magnitude of ΔS‡ depend very much on the charge of the activated complex relative to the charges of the reactants, since the charge on a species determines its extent of solvation. For reactions between oppositely charged ions, the activated complex will have a lower charge than the reactants and is thus less solvated, hence its formation will be accompanied by an increase in disorder; consequently, ΔS‡ will be positive. For a reaction between ions of like charge, ΔS‡ will be negative due to increased solvation of the transition state, having a charge greater than that of the separate reactants. This electrostatic contribution to ΔS‡ approximately equals −10 × ZAZB eu (1 eu = 1 cal K−1 mol−1 = 4.18 J K−1 mol−1), where ZA and ZB are the number of units of charges carried by the ionic reactants A and B, respectively. For dissociation reactions, involving separation of ions (such as loss of halide from a cationic complex), the activated state approximates to an ion pair which is more solvated than the original substrate, hence ΔS‡ is negative. However, the electrostatic contributions of the type mentioned above are not always dominant and other factors may be important. Thus, for the dissociation reaction   

[Fe(phen)3]2+ + H3O+ + H2O → [Fe(OH2)2(phen)2]2+ + phenH+

(2.43)

  

Figure 2.5  Isokinetic  trend in the anation (replacement of one H2O by an anion) reactions of [Cr(OH)(OH2)5]2+. Anions: (1) Cl−, (2) SCN−, (3) NCS−, (4) Br−, (5) NO−3 and (6) HF (entry of F− with loss of H+).

Chapter 2

84 ‡

the ΔS value is +28 eu, which may be due to much freer ligands in the transition state, and arguments have been put forth84a regarding the sign and magnitude of ΔS‡. In the replacement of one ligand by another in a metal complex, there is often a significant contribution to ΔS‡ due to losses and gains in the entropy of the entering and departing groups being markedly different. The entropy of crystallization of water being −6.0 eu, an entropy change of this magnitude might be expected for fixing a water molecule in the transition state.84b However, from consideration of the fact that the entropy loss per mole of hydration (i.e. entropy of hydration/hydration number) for Cs+, Li+, Ba2+ and Mg2+ ions is −3.85, −2.95, −2.25 and −2.80 eu, respectively, Powell84c concluded that an entropy change of −3.0 eu will result in fixing one molecule of water in the transition state. Also, in a bimolecular reaction in which one mole of reactant disappears on formation of the activated complex, there will be a decrease in entropy, which for water is equal to −Rln(1000/55.5) = −5.75 eu, and this will contribute to the overall ΔS‡ for the reaction. Hence it is apparent that many factors must be considered for correlating an observed ΔS‡ value with a particular reaction mechanism. However, many important applications of ΔS‡ in the inference of a reaction mechanism are known. Thus, for the reaction   



[CoCl(NH3)5]2+ + OH− → [Co(OH)(NH3)5]2+ + Cl−

(2.44)

  

which is first order with respect to each of the reactants, the observed ΔS‡ is +37 eu, whereas ΔS for the overall reaction is −1.0 eu. The decided lack of correlation is convincing evidence against a simple SN2 mechanism.85 Also, for the aquations of cis-­[CrCl2(en)2]+, trans-­[CoCl2(dmgH)2]− and trans-­ [CoCl(dmgH)2(OH2)] (where en = ethylenediamine and dmgH2 = dimethylglyoxime), involving replacement of a chloride ligand by a molecule of water, the observed ΔS‡ values are nearly the same as the entropy of hydration of chloride ion (−22.8 eu); this is a convincing evidence for a dissociation mechanism, where separation of the chloride ligand from M(iii) is virtually complete in the transition state.74h,i. Measured ΔS‡ values are particularly useful when analogous reactions are compared, since here many common factors are eliminated. It is should be mentioned in this connection, however, that too much reliance should not be placed on small differences in ΔS‡ values (≤2 eu), since the uncertainty in the evaluation of ΔS‡ is usually at least ±0.5 eu.

2.4.2  Effect of External Pressure86 The effect of external pressure on the rate of reaction is expressed by the relation   

  

V  P2  P1  k  log  2    2.303RT k  1 T ‡

(2.45)

Techniques for Following Reactions and Factors that Affect Rates

85

‡

where ΔV (volume of activation) is the molar volume change that takes place in the conversion of the reactants to the activated complex; k2 and k1 are the rate constants at pressures P2 and P1, respectively, at a constant temperature T. Here P is expressed in dyne cm−2 (= 10−6 bar) and R = 83.1 × 106 dyne cm mol−1 K−1; ΔV‡ is obtained in cm3 mol−1. More usually pressure is expressed in atmospheres (1 atm = 1.01 bar). A fairly large pressure change (2000–10 000 atm) must be used to obtain a sufficiently marked effect on the rate (and hence k) to allow the evaluation of ΔV‡. However, for the relation in eqn (2.45) to be strictly valid, the pressure interval P2 – P1 should not exceed 1000 atm, since ΔV‡ usually changes rapidly with pressure. Because of the dependence of ΔV‡ on pressure (since the compressibilities of the reactants and transition state are different), it is customary to use ΔV‡ values as obtained at relatively lower pressures, or better to use values at low (normal) pressure (or even at P = 0, giving ΔV‡0) by extrapolation using the following relation:   

ln k = a + bP + cP2



(2.46)

  

Up to a pressure of ca. 1000 atm the (log k)T versus P plot is generally linear87 and its slope is −(ΔV‡)/2.303RT. However, for the precise evaluation of ΔV‡ it is usually necessary to work at a much higher pressure range where the (log k)T versus P plot is non-­linear. In such cases the slope at any point of the curve is the value of −(ΔV‡)/2.303RT at the corresponding P. ΔV‡ values, like other activation parameters (ΔH‡, ΔS‡), are additive, as illustrated below with two reaction schemes:    (a) A + B ⇌ C, ΔV; C → Products, ΔV‡1    ΔV‡ = ΔV + ΔV‡1   

(b)    

 V ‡

k1 k2 V1‡  V2‡ k1  k2 k1  k2

Hence a change of reaction path by a change of pressure is quite feasible. Expected ΔV‡ values for different types of processes are available.88 Since ΔV‡ can have either positive or negative values, reactions may be accelerated or retarded by an increase in pressure. The magnitude of ΔV‡ is often diagnostic of the mechanism89 and many such applications will be mentioned in some of the following chapters. Although the assumptions used in the detailed analysis of ΔV‡ values to diagnose reaction mechanisms have been questioned by Langford,90 this has been firmly countered by Newman and Merbach91 and Swaddle.92 It is further of interest that for the aquation

Chapter 2

86 3

and water exchange reactions of several octahedral d [Cr(iii)] and low-­spin d6 [Fe(ii), Co(iii), Rh(iii), Ir(iii)] complexes, a fairly good linear correlation between ΔV‡ and ΔS was reported:93   

  

ΔV‡ (cm3 mol−1) = (1.04 ± 0.10)ΔS‡ (cal K−1 mol−1) − (4.4 ± 0.3) (2.47)

Similar linear relationships between ΔV‡ and ΔS‡ were demonstrated earlier for several reactions.94 This is of significance, since knowledge of ΔV‡ is of considerable value for the elucidation of mechanisms,95 but the experimental evaluation of ΔS‡ is much easier. However, incorporation of additional data from later work into the ΔV‡ versus ΔS‡ plot for the d3 and low-­spin d6 octahedral complexes has shown too much scattering of the data points from linearity, hence the relationship of eqn (2.47) is not quantitative.89d ‡ Two components compose ΔV‡, an intrinsic component ΔVint , which arises from changes in the internuclear distances of the reactants during the formation of the transition state, and the electrostriction component ‡ ΔVelec , which arises from changes in the electrostriction of solvent in the solvation sphere of the species as the charge distribution changes in the formation of the transition state.96–98 In a polar solvent such as water, a charged species, owing to its electrostatic influence, attracts the solvent molecules and holds them around itself firmly in its primary solvation sphere, and there will be a similar but somewhat milder effect in the secondary solvation sphere also, leading to a contraction in volume (solvent ‡ electrostriction). When charged reactants are involved, ΔVelec may domi‡ ‡ nate ΔV so that ΔVint, which contains most of the mechanistic information, cannot be readily determined. However, solvent exchange reactions such as water exchange of [M(OH2)n]m+ involve negligible differences in electrostriction between the reactants and the transition state such that ‡ ‡ ≈ 0 and hence ΔV‡ ≈ ΔVint . In these cases, the activation volume is parΔVelec ticularly useful and can be a direct indicator of mechanism through comparison;97 this topic has been reviewed.98 A necessary presumption for this argument is that the volume change within the coordination sphere of the complex itself can be largely neglected, allowing the sign of the activation volume to be directly diagnostic of the type of activation. This follows the proposal originally made by Stranks99 that an associatively activated process will have a negative activation volume corresponding to the loss, from the initial state, of what will be the effective partial molar volume of the entering ligand on forming the transition state. Conversely, a dissociatively activated process will involve the liberation, at the transition state, of the equivalent effective partial molar volume of the leaving group, thus leading to a positive activation volume. Because of the usefulness of ΔV‡ values in the assignment of reaction mechanisms, the compilation of ΔV‡ values for various types of reactions of metal complexes and organometallic compounds100 is useful. For earlier literature data, several reviews are available.101

Techniques for Following Reactions and Factors that Affect Rates

87

2.4.3  Effect of Ionic Strength102 The simple Debye–Hückel theory expresses the dependence of the activity coefficient of an ion on the ionic strength of the solution as follows:   

1



 Zi2 I 2

 log fi  1 1   i I 2

(2.48)

  

where αi (Å) is the diameter of the solvated ion present in the solution. For H2O at 25 °C α = 0.509 and β = 0.33. Eqn (2.48) is valid for I up to 0.2 M. At I ≪ 1,  i I 1 / 2 ≪ 1 and therefore   



1

 log fi   Zi2 I 2

(2.49)

  

Based on eqn (2.35) and (2.49), and assuming that the diameters of the solvated reactant ions and of the activated complex are the same, this leads to the prediction that the rate of a reaction between ions of charge ZA and ZB should vary with the ionic strength, I, of the solution in the following manner: 1   2 Z A ZB I 2    log k1  log k0  1   2    1 I i     



1

 log k0  2 Z A ZB I 2 (when I  1)

(2.50)

  

where α and β are constants for a particular solvent and temperature, αi (Å) is the diameter of the solvated ions in the solution, k1 and k0 are the rate constants at ionic strengths I and 0, respectively, the constant α is proportional to the −3/2 power of the solvent dielectric constant and is equal to 0.059 and β = 0.33 for water at 25 °C. The ionic strength, I, is defined following Lewis by the equation    1 I   ci Zi 2 (2.51) 2   

where ci is the concentration of the ion i and Zi is the number of units of charge of the ion; the summation being taken of all of the ions in the solution. Hence, for a uni-­univalent electrolyte, I is equal to the molar concentration of the electrolyte, but it is greater than the concentration when the salt contains ions of higher charges. Eqn (2.50) predicts that the plot of log k1 versus I1/2 should be a straight line of slope 2αZAZB (≈ZAZB for reactions in aqueous solution at 25 °C, since α = 0.509) and intercept equal to log k0, hence k0 can be evaluated by extrapolation to I = 0. This relationship is strictly valid, however, only for dilute solutions (up to 0.01 M for uni-­univalent electrolytes)

Chapter 2

88

and in the absence of complex formation or any reaction between reactant and/or products, with the ions of the electrolyte used for maintaining the ionic strength; hence the inert salt sodium perchlorate is generally very convenient for this purpose. As a precaution against unwanted ionic strength effects, the kinetics of reactions in solution are studied in the presence of a constant and fairly high concentration of an inert salt. The use of a constant excess of an inert salt effectively swamps variations in ionic strength, that would otherwise occur due to changes in concentrations of reactants, acidity, etc. In the event that one of the reactants is a non-­ionic species, ZAZB is zero and in such a case the rate constant is independent of ionic strength. However, this is true only in dilute solutions (particularly when I ≤ 0.01 M). In solutions of higher ionic strength, a slight dependence of rate on ionic strength is observed even when one of the two reacting species is uncharged. Under such conditions, the activity coefficient of an ion is given by the modified Debye–Hückel equation (which is valid up to I ≈ 1 M)∥:   1  2 2  Z I i  log  fi  1  2  1   i I



   I  i 

(2.52)

  

For an uncharged species under similar conditions, the f0 value is given by102b   

−log f0 = γ0I



(2.53)

  

Based on these and eqn (2.52), it can be shown that for a reaction involving an ion and a non-­ionic species in solutions that are not very dilute (I > 0.01 M), the following relation is valid:   

log kI = log k0 + mI



(2.54)

  

where m is a constant. Hence, for such a reaction system the plot of log kI versus I is linear. Apart from a primary salt effect of the type mentioned above, secondary salt effects are often observed due to ion-­pair formation or other specific interactions. Thus, in the reaction of [CoBr(NH3)5]2+ and Hg2+ in aqueous solution leading to the formation of [Co(NH3)5(OH2)]3+ and HgBr+, the rate depends not on the ionic strength but on the concentration and character of negative ions of the electrolyte added to maintain the ionic strength; the effect is ∥

See also the Davies modification (C. W. Davies, Ion Association, Butterworth, London, 1962): 1    Zi2 I 2  2  log  fi  1    Zi I  2   1  I i  

valid at I up to 0.6 M; γ = 0.10 (for H2O at 25 °C).

Techniques for Following Reactions and Factors that Affect Rates

89

known to be due to ion-­pair formation by the anion with reactant ions, each ion pair reacts at a characteristic rate. At I > 0.01, serious deviations from the predicted behaviour [eqn (2.50)] are observed owing to inapplicability of the Debye–Hückel theory and also because many of the complex ions, being highly charged, are substantially associated as ion pairs. It may be significant to mention that even in 0.001 M solution, ca. 18% of hexaamminecobalt(iii) chloride exists as the ion pair [Co(NH3)6]3+ Cl−. Ion association is generally least in perchlorate salt solutions, hence the desirability of their use in rate studies. In many cases nitrates such as KNO3 are also suitable.

2.4.4  Influence of Solvent103 See also Chapter 4, Section 4.1.8, and Chapter 5, Table 5.4. If the interacting ions have charges ZA and ZB and they are at a distance dAB apart in the transition state, then it can be shown that   



 Z A ZB e 2 log  k log k    2.303 dAB kT

  

(2.55)

  

where e is the electronic charge, k is Boltzmann's constant, T is the absolute temperature, kε is the rate constant in a medium of dielectric constant ε and k∞ is the value of the rate constant in a medium of infinite dielectric constant, where the second term on the right-­hand side of eqn (2.55) becomes zero. Hence log kε should vary linearly with 1/ε, and this has been verified experimentally in a number of cases of reactions carried out in a series of mixed solvents of varying ε; from the slope of the line dAB can be calculated and reasonable values (a few Å) have always been obtained. Deviations from this linear relationship are often observed, particularly at low dielectric constant, due to failure of the assumption involved in deriving eqn (2.55) and in some cases due to changes in the reaction mechanism as the solvent is varied (see Chapter 4). From the dependence of the rate constant on the dielectric constant of the medium and other well-­known empirical solvent parameters, such as Y,104 Z105 and ET,106 which express the solvating and ionizing power of the solvent, important conclusions can be drawn regarding reaction mechanisms.107

2.5  Nucleophilicity and Rate Ingold defined nucleophiles as “reagents which act by donating their electrons to, or sharing them with, a foreign atomic nucleus”. In this broad concept, not only bases and ligands but also all reducing agents are nucleophiles. The generally accepted concept is, therefore, that “both bases and nucleophiles are reagents, which have a tendency to form new covalent bonds by sharing their electron pairs”. However, “basicity” is used in the thermodynamic sense, whereas nucleophilicity is used in describing the kinetic

Chapter 2

90

behaviour. Thus, the basicity of B in aqueous solution is measured by the equilibrium constant, Kb, for the reaction   



B + H3O+ ⇌ BH+ + H2O

(2.56)

  

whereas the rate constant k of the following reaction is a measure of the nucleophilic reactivity of B with respect to that of X:   



  

k RX  B   RB  X

(2.57)

In the system represented by eqn (2.56), H+ is acidic, and R in the system represented by eqn (2.57) is an electrophile. For the equilibrium represented by eqn (2.56) in aqueous solution, we have   



BH  H2 O BH     K b  , i . e . K K b b  H2 O  B H3O   B H3O  

(2.58)

  

and this Kb is related to Ka, the acid dissociation constant of BH+, as follows:   

  

 Ka

B H3O  

1   Kb BH 

(2.59)

For a given group of nucleophiles containing the same nucleophilic atom, and the structural features of which in the immediate vicinity of the nucleophilic atom are similar, nucleophilicities are well correlated with their basicities.

2.6  Relative Nucleophilicities General base catalysis is observed when all proton acceptors catalyse a reaction by removing a proton from the substrate, SH, in the transition state:   



slow SH  B    S  BH

(2.60)



fast S  Products

(2.61)

     

where the first step is rate determining. For this change,   

Rate = kb[B][SH]

(2.62)

  

Brønsted108 showed that the rate constant kB is related to the basicity constant Kb of B as follows:   



log kb = α + βlog Kb

(2.63)

  

where α and β are constants for the particular system. Similarly, for acid catalysis of the type   



  

slow S  AH   SH  A 

(2.64)

Techniques for Following Reactions and Factors that Affect Rates

91



fast SH  Products

(2.65)



Rate  kA [AH][S]

(2.66)

     

for this system the Brønsted relationship is   



log kA = α′ + β′log Ka

(2.67)

  

where α′ and β′ are constants for the particular system and Ka is the acid dissociation constant of AH. In both general base catalysis and acid catalysis, the proton is only partially transferred in the activated complex. If transfers were complete, the values of the constants α (or α′) and β (or β′) would be unity, since ionization (as measured by Kb and Ka) involves complete dissociation. If, however, there is negligible transfer, then the values of α (or α′) and β (or β′) would be almost zero and there would be no observable acid or base catalysis in this case. In actual practice, the values of α (or α′) and β (or β′) lie between 0 and 1, and the magnitude of the value may be taken as a measure of the degree of proton transfer in the activated complex. Detailed studies on the base-­catalysed hydrolysis of Cr2O2− 7 have been reported; a reasonable correlation of log kb versus log Kb was observed for a fairly large number of bases, which is in agreement with eqn (2.63).109 The Brønsted equation for general base catalysis is a correlation of the rates of nucleophilic displacement on hydrogen, with the basicities of the nucleophiles as expressed by the pKa values of the conjugate acid forms of the nucleophiles (see Table 2.2). Rates of nucleophilic displacement on some other substrates have also been found to follow these basicities. For SN2 reactions, Swain and Scott110 suggested   



  

log  kY / kH2 O   n

(2.68)

where kY is the rate constant of the reaction CH3Br + Y → CH3Y+ + Br− kH2 O is the rate constant for corresponding reaction with the solvent H2O and n is the nucleophilic constant for the nucleophile Y. Hence, by definition, n = 0.00 for H2O. For any other substrate,   



  

log  kY / kH2 O   sn

(2.69)

where s is the substrate constant; for the standard substrate (CH3Br), s = 1.00 by definition. Values of n for a variety of nucleophiles are also given in Table 2.2. Edwards111 proposed another nucleophilic scale. This was based on the assumption that the reactivity of a nucleophile in a displacement reaction is governed by a linear combination of the factors that govern its ability to be oxidized (i.e. by its electrode potential) and its thermodynamic affinity for a proton (i.e. by its basicity). The following relationship was suggested:   



  

log  kY / kH2 aEn  bH O

(2.70)

Chapter 2

92

Table 2.2  Nucleophilic  strengths of various nucleophiles in different scales. Nucleophile

pKaa

S2− SO32− S2O32− SC(NH2)2 I− CN− SCN− C6H5NH2 NO2− OH− N3− Br− NH3 Cl− C5H5N CH3CO2− SO42− F− H2O

12.9 7.2 [9.1] 1.7 [1.9] −0.96 [0.4] (−10) 9.2 [9.1] (−1.8) [−0.7] 4.5 3.3 [3.4] 15.7 4.7 (−7) 9.2 [9.5] (−5) [−4] 5.2 [5.3] 4.7 2.0 3.2 −1.7

nb

6.36 (4.1) 5.04 4.77 4.49 4.20 4.00 3.89 3.04 (3.6) 2.72 (2.5) 0.00

Enc 3.08 2.57 2.52 2.18 2.06 2.02 1.83 1.78 1.73 1.65 1.58 1.51 1.36 1.24 1.20 0.95 0.59 −0.27 0.00

Pd 0.611

0.718 0.373

0.143 0.539 0.184 0.389

−0.150 0.00

a

pKa values of conjugate acids of nucleophiles at 25 °C in aqueous solution; parentheses ( ) denote approximate values. In several cases the values agree with those quoted by Edwards in his book (see ref. 5a in Chapter 1), but there are difference in several cases, as given in brackets [ ]. b Nucleophilic scale of Swain and Scott (1953).110 c Electrode potential scale of Edwards (1954).111 d Polarizability scale of Edwards (1956).112

where kY and kH2 O are the rate constants of the reaction of a substrate with the nucleophile Y and H2O, respectively, and a and b are substrate constants that measure the sensitivity of the substrate to changes in En and H, respectively. By definition, H = pKa + 1.74 (where Ka is the ionization constant for YH+ in water at 25 °C and 1.74 is a correction term for the pKa of H3O+) and En = E0 + 2.60 (where E0 is the standard electrode potential for oxidation of the nucleophile Y and 2.60 is a correction term for the corresponding value for H2O; hence En is the standard oxidation potential of Y relative to that for H2O). Both H and En are characteristics of the nucleophile. In this double scale, nucleophilicity is correlated with two independent properties of the nucleophile, viz. electrode potential and basicity, and their relative contributions may vary from one substrate to another. Allowance is thus made for the observed variation of nucleophilic order with varying substrates. The constants a and b are determined empirically for each substrate; with their values suitably chosen, this equation can be made to fit a large amount of rate data. Once a and b are known for a substrate, the nucleophilicity of any nucleophile can be estimated knowing En and H. Values of nucleophilicity of several nucleophiles on this scale are also given in Table 2.2.

Techniques for Following Reactions and Factors that Affect Rates

93

112

In a later correlation, Edwards hypothesized that En is dependent on the polarizability of the nucleophile and its basicity to proton, H+:   



En = a'P + b'H

(2.71)

  

where P is a measure of the polarizability of the nucleophile Y relative to H2O and is by definition P  log  RY / RH2 O , R being molar refractivity (at infinite wavelength), a' = 3.60 and b' = 0.0624. Hence, combining eqn (2.70) and (2.71), the following is obtained:   

  

log  kY / kH2 O   a  aP  bH   bH

 aa ' P  (ab ' b) H  AP  BH

(2.72)

where A and B are constants for the substrate. With their values suitably chosen, eqn (2.72) can be made to fit a large amount of rate data. Nucleophilicity values on this scale are also given in Table 2.2. Belluco, Pearson and co-­workers113 proposed a nucleophilicity scale, 0 n Pt, with trans-­[PtCl2(py)2] as a reference substrate in methanol solution (at 30 °C):   



  





log kY / kCH3OH  nPt0

(2.73)

For reaction of any other Pt(ii) substrate in any solvent S,   



logkY = sn0Pt + logkS

(2.74)

  

where s is the substrate constant (nucleophilic discrimination factor), being (by definition) 1.00 for trans-­[PtCl2(py)2]. In eqn (2.73), the solvent rate constant kCH3OH has been divided by 26 M, the molar concentration of methanol, so that n0Pt is dimensionless. In Table 2.3, n0Pt values are compared for dif0 0 ferent nucleophiles with their pKa and nCH values. Values of nCH3I are also 3I defined by eqn (2.74), but only with respect to CH3I as the substrate (i.e. displacement of I− from CH3I by different nucleophiles). It can be seen that 0 there is a rough correlation of the n0Pt and nCH3I values, but many exceptions 0 are noticeable. A similar n Pd scale has been proposed.114 kS and kY in eqn (2.74) are the solvent-­dependent and reagent (entering nucleophile)-­dependent paths, respectively (see Chapter 3). For any Pt(ii) complex, reacting with different nucleophiles (Y) in a solvent (S), the plot of log kY versus n0Pt of the nucleophile is a straight line (see Figure 2.6) of slope s of the complex of Pt(ii) and intercept = log kS. The kS value so obtained agrees reasonably well with kS determined experimentally. However, for systems for which kS is exceedingly small and cannot be determined directly, the intercept of the log kY versus n0Pt plot is highly negative, as expected. The s and kS values for several complexes of Pt(ii) reacting with different nucleophiles in solution in a solvent are given in Table 2.4. The data show

Chapter 2

94

0 Table 2.3  Nucleophilicity  values in n0Pt scale compared with pKa and nCH values for 3I

different nucleophiles.

Nucleophile

pKaa

n0Pt

0 nCH 3I

CH3COO− CH3O− C6H5NH2 Cl− NH3 C5H5N NO2− N3− NH2OH N2H4 C6H5SH Br− (CH3)2S I− (CH3)2Se SCN− SO32− CN− (CH3O)3P SeCN− C6H5S− SC(NH2)2 S2O32− (C2H5)3As (C6H5)3P (C4H9)3P (C2H5)3P

4.75 15.8 4.52 (−5.7) 9.25 5.23 3.33 4.74 5.82 7.93 — (−7.7) −5.3 (−10.7) — −1.8 7.26 9.2 — — 6.52 −0.96 1.7 Br− > Cl− > NH3 > OH− This sequence is similar to the trans effect sequence, with some striking exceptions, notably for CO and CN−, which have a high trans effect owing to their π-­acid character and hence a strong tendency for M → L π-­bonding; the π-­bonding ability of different ligands is H2C=CH2 >> CO > CN− > NO2− > SCN− > I− > Br− > Cl−> NH3 > OH− when both this sequence and the trans influence sequence are considered, the trans effect sequence can be rationalised fairly well. Based on 119Sn NMR studies on a series of SnCl3− complexes of platinum(ii), the following trans influence series was proposed:17d H− > PR3 > AsR3 > SnCl3− > olefin > Cl− Syrkin18 used a different approach to explain the trans effect, based on the concept of resonance, and suggested that the greatest labilization occurs to a group trans to the ligand with the maximum degree of covalent bonding in the complex (Figure 3.7). In [PtCl(NH3)3]+, the two principal structures

Chapter 3

112

Figure 3.7  Syrkin's  resonance theory of the trans effect. Resonance in (a) [PtCl(NH3)3]+ and (b) [PtCl3(NH3)]−.

Figure 3.8  (a)  s and d x 2  y 2 orbitals of Pt(ii); (b) s–d hybrid orbitals of Pt(ii). between which resonance occurs would lead to Pt–NH3 bonds in cis positions to the Pt–Cl bond being 50% ionic and 50% covalent, while the bonding of NH3 in the trans position to Cl− is essentially 100% ionic (Figure 3.7a), hence this NH3 (trans to Cl−) is readily replaced. Similar arguments hold good for [PtCl3(NH3]− (Figure 3.7b). For the metals of the third transition series in their bivalent state, the relative energies of the bonding atomic orbitals are 5d ≈ 6s < 6p. Hence the strongest covalent bond will make greater use of the d and s orbitals of the metal. Use of an s + d hybrid orbital for one bond prohibits the use of the s + d hybrid orbital for the ligand in the trans position to the first bond, but allows an s–d hybrid for use of the ligand in its cis position (Figure 3.8). Hence the trans position has less covalent bonding and therefore the ligand in this position is more readily replaced.

Ligand Replacement Reactions of Metal Complexes

113

Thus, in both Grinberg's and Syrkin's theories, the trans effect is believed to be due to a weakening of the bond trans to a ligand having a strong trans effect. X-­ray crystallographic data revealed that in the complexes of the type trans-­PLA2LX, the increase in Pt–X bond length parallels the increase in the trans effect of L. Although the difference between L = NH3 and L = Cl− is small, that for L = NH3 and L = Br− is significant. The behaviour is similar for [PtX3(NH3)]− (X = Cl, Br)19a and [PtX3(H2C=CH2)]− (X = Cl,19b Br19c), as shown in Table 3.2. Again, in the following dinuclear Pd(ii) complex (1) the Pd–Cl bond opposite to terminal Cl is shorter (2.32 Å) than the other Pd–Cl bond opposite to H2C=CH2 (2.42 Å). Similar results have been observed in the case of the corresponding styrene complex. Thus, on the basis of commonly used criterion of bond length as related to bond strength in Pt(ii) and Pd(ii) complexes, an olefin trans to a Cl− or Br− results in a larger Pt–X distance and hence a weaker Pt–X bond. However, the Pt–Br bond length increases in the following sequence of L = Br− > H2C=CH2 > Cl− ≈ NH3, which is not quite the same as the trans effect order of these ligands. Also, the Pt–N bond trans to H2C=CH2 or SCN− is of almost normal length. It has now been established that the high trans effect of SCN− and H2C=CH2 is due to entirely different reasons (see below).

Infrared data obtained by Chatt et al.19d on the N–H stretching frequency (νN–H) in complexes of the type trans-­[PtCl2L(NHRR′)] have shown that νN–H decreases (i.e. the N–H bond is weakened and hence the Pt–N bond strength increases) in the following order of the trans ligand, L, which (except for H2C=CH2) is also roughly in the order of decreasing trans effect of L: PR3 > SbR3 > P(OR)3 > AsR3 > TeR2 > H2C=CH2 > SeR2 > SR2 > piperidine > 4-­n-­pentylpyridine. Table 3.2  Some  Pt–X bond lengths in [PtX3(L)]−. K[PtX3(L)]

X

trans partner of X (L or X)

Pt–X bond length/Å

K[PtCl3(NH3)]

Cl Cl Br Br Cl Cl Br Br

Cl NH3 Br NH3 Cl H2C=CH2 Br H2C=CH2

2.35 2.32 2.70 2.42 2.32 2.42 2.42 2.50

K[PtBr3(NH3)] K[PtCl3(H2C=CH2)] K[PtBr3(H2C=CH2)]

114

Chapter 3

Hence, except for H2C=CH2, the greater the trans effect of L, the weaker is the trans-­Pt–N bond, and the same behaviour has been observed in the cases of the corresponding Pd(ii) complexes. Also of significance is the observation that for L = R2S, R2Se, R2Te, PR3, AsR3 and SbR3 (R = n-­alkyl group) νN–H increases with decreasing electronegativity of the ligand (the sequence of electronegativity is R2S > R2Se > R2Te > R3P > R3As > R3Sb), except for PR3, and this anomalous behaviour of PR3 is greater in Pt(ii) than in Pd(ii) complexes. The effect can be simply understood as discussed below. The decrease in νN–H with increasing electronegativity of L can be explained as follows. An increase in the electronegativity of L causes this to withhold its electrons from the Pt atom in the L–Pt bond and so increase the attraction of Pt for electrons in the Pt–N bond as shown in 2. Thus N becomes less negative and the proton less strongly bound, resulting in the observed decreases in νN–H. The exceptional behaviour of H2C=CH2 is that, although it has a very high trans effect, it does not give the weakest Pt–N bond, and in support of this are the X-­ray data mentioned earlier, that a Pt–N bond trans to H2C=CH2 has almost the normal length. Measurements by Powell20 of νPt–N in trans-­ [PtCl2L(NH3)] (L = NH3, SEt2, H2C=CH2) of 507, 493 and 481 cm−1, respectively, are also in accord with the above order.

The anomalous behaviour of H2C=CH2 and of PR3 has been explained as being due to their π-­bonding with the d electrons of the metal and the effect of this on the N–H vibration. This is presumed to decrease the direct interaction of the N-­bonded hydrogen atom with the filled d or dp hybrid orbital of Pt(ii). Like the trans influence, there is also a cis influence, and the stereochemical effects of mutual influence of ligands have been reviewed.21 Illustrative examples of trans and cis influence are given in Tables 3.3–3.6. The data in Table 3.3 show explicitly that in square-­planar complexes of Pt(ii), the trans ligand strongly influences the Pt–Cl bond length. Substitution of the trans ligand changes the Pt–Cl bond length from 2.28 to 2.45 Å. In comparison, the cis influence (Table 3.4) is much weaker than the trans influence. The data in Table 3.5 allow us to follow the variation of one of the most informative parameters of the trans influence, the trans elongation ∆ = LM–Xtrans − LM–Xcis in a series of MOX5 complexes with d0, d1 and d2 configuration.22 It is seen that there is an essential decrease in trans influence along the series d0 > d1 > d2. Note, however, that in Table 3.5 the change of the dn configuration is associated with a change of the M atom. Table 3.6 summarizes some other structural characteristics for trans-­ and cis-­[PtCl2(PR3)2] (R is mainly Et and sometimes Me): bond lengths LPt–Cl and LPt–P, vibrational frequencies

Ligand Replacement Reactions of Metal Complexes

115

Table 3.3  trans  influence on bond length LPt–Cl in some Pt(ii) complexes.

22

Complex

trans atom (donor of ligand)

LPt–Cl/Å

K[PtCl(acac)2] trans-­[PtCl2(PEt3)2] cis-­[PtCl2(4-­C6H4S)2] trans-­[PtCl(CO)(PEt3)2] cis-­[PtCl2(CNEt)(PEt2H)] K2[PtCl4] [PtCl2(l-­methionine)] K[PtCl3(NH3)]·H2O K[PtCl3(H2C=CH2)]·H2O [PtCl3(H2NCH2CH=CHCH2NH2)] cis-­[PtCl2(PMe3)2] cis-­[PtCl2(CNEt)(PEt2Ph)] trans-­[PtCl(CH2SiMe3)(PMe2Ph)2] trans-­[PtCl(H)(PEtPh2)2] trans-­[PtCl(SiMePh2)(PMe2Ph)2]

O Cl S of C6H4S− CO C of EtNC Cl S of l-­methionine N C=C C=C P P C of CH2SiMe3 H Si

2.28 2.30 2.30 2.30 2.314 2.316 2.32 2.321 2.327 2.342 2.37 2.39 2.415 2.42 2.45

Table 3.4  cis  influence on the bond length LPt–Cl in some Pt(ii) complexes.22 Complex

cis atom (donor of ligand)

LPt–Cl/Å

trans-­[PtCl2(PEt3)2] [PtCl2(cis-­MeCH=CHCH2NH2)] K[PtCl3(H2C=CH2)]·H2O K[PtCl3(NH3)]·H2O K2[PtCl4] trans-­[PtCl2(NH3)2]

P C=C C=C N Cl N

2.29 2.297 2.305 2.315 2.316 2.32

Table 3.5  Comparison  of cis and trans bond lengths, LM–Xcis and LM–Xtrans, and trans

elongation (∆ = LM–Xtrans − LM–Xcis) in some MOX5 complexes of Mn+ of d0, d1 and d2 configurations.22 dn configuration LM–Xtrans/Å

Compound K2[NbOF5] K[MoOF5] K2[MoOF5]·H2O K2[MoOCl5] [NH4]2[MoOBr5] K[OsOF5] K2[ReOCl5]

0

d d0 d1 d1 d1 d2 d2

2.06 2.29 2.03 2.587 2.83 1.72 2.47

LM–Xcis/Å

∆/Å

1.84 1.86 1.87 2.39 2.55 1.78 2.39

0.22 0.43 0.16 0.20 0.28 −0.06 0.08

Table 3.6  Comparison  of structural data for cis-­and trans-­[PtCl2(PR3)2] (R = mostly Et, Me in a few cases).22

Parameter LPt–Cl/Å LPt–P/Å 1 JPt–P/Hz 2 JP–P/Hz ν(35Cl)/MHz

Cis 2.37 2.25 3520 −18.7 ∼18.0

trans 2.29 2.30 2400 510 20.99

Parameter −1

νPt–Cl/cm νPt–P/cm−1−1 δ31P/ppm Eb/(Cl, 2p)/eV

cis

trans

294 435 24.0 198.2

340 419 15.8 198.0

Chapter 3

116 1

2

νPt–Cl and νPt–P, constants of nuclear spin–spin interaction JPt–P and JP–P, NMR chemical shifts δ(31P) with respect to H3PO4, NQR frequency ν(35Cl) and the 2p electron bonding energy Eb in Cl determined from X-­ray photoelectron spectra. A comparison of these data shows that the Pt–Cl bonds in the trans complexes and Pt–P in the cis complexes are stronger than the same bonds in the cis and trans complexes, respectively. This shows that PR3 has a stronger trans influence than Cl, and this is reflected in both bond length and the bond stretching frequency. The first attempt to explain the origin of the trans influence was based on a comparison of the σ-­donor and π-­acceptor properties of the ligands. In a square-­planar complex MX4, replacement of one X by a better σ-­donor Y will weaken the Pt–X bond trans to Y as they use the same set of AOs of Pt for σ-­bonding. A simple expression of trans influence was made by the angular overlap model (AOM).23 Based on this, it was reasonably concluded that trans influence (Pt–X bond elongation) will increase with increasing σ-­donor character (and hence a decrease in Pauling electronegativity) of the donor Y in the trans position to X in Y–Pt–X. Figure 3.9 shows that the experimental data, in general, confirm this trend;23 note that the σ-­donor effect on the trans influence had been suggested much earlier.24 The presence of π-­bonds is crucial for the mutual influence of ligands. The role of π donation in the trans effect was first reported by Chatt et al.25 and independently by Orgel.26 In connection with the role of π-­bonding on the trans influence, it is of interest that in cis-­K2[PtCl2(NO2)2] the two NO2− ligands which can form strong π-­bonds with Pt(ii) are relatively weakly trans influencing; the Pt–Cl bond length (trans to NO2) is 2.34 Å, which is the same as for the Cl trans to C=C in square-­planar complexes of Pt(ii) (see Table 1.2), although the π-­acceptor property of NO2− is much stronger.

Figure 3.9  Pt–Cl  bond length (Å) in square-­planar Pt(ii) complexes (Table 3.3) versus Pauling electronegativities, χp, of the donors of the trans ligands.

Ligand Replacement Reactions of Metal Complexes

117 −

This discrepancy has been explained as being due to the fact that NO2 is bent with respect to the position in which it could form strong π-­bond; X-­ray data confirmed the bent position of NO2− in the dichlorodinitro complex (see ref. 22, p. 478). Experimental data suggest that generally ligands having a high trans influence have a weak cis influence and vice versa. In this connection, it is worth comparing the Pt–Cl bond lengths in the two complexes trans-­[PtCl2(PEt3)2] and trans-­[PtCl2(NH3)2]. They differ in the trans linear fragment P–Pt–P and N–Pt–N, respectively, in the cis position of the trans linear fragment Cl–Pt–Cl in both complexes. Hence the difference in the Pt–Cl bond length is entirely due to the cis influence of the PEt3 and NH3 ligands. The data in Table 3.5 show that whereas the trans influence of P is stronger than that of N, the cis influence of P is weaker than that of N, since the Pt–Cl bond length in the trans-­bis(phosphine) complex is 2.29 Å and that in the trans-­diammine complex is 2.32 Å. Grinberg and others27a,b studied the equilibrium shown in eqn (3.5) and reported the equilibrium constant values for different X−: I−, 5 × 10−7; SCN−, 2.8 × 10−7; Br−, 2.8 × 10−9; Cl−, 1 × 10−10; i.e. the affinity of Pt(ii) for X− is I− ≈ SCN− > Br− > Cl−, which is also the order of decreasing trans effect of these ligands. Similar behaviour has been reported for [PtX4]2− (see Chapter 1, Section 1.7).   



trans-­[Pt(OH)2(NH3)2] + 2X− ⇌ trans-­[PtX2(NH3)2] + 2OH−

(3.5)

  

Chatt and Leden27c determined the equilibrium constant (K) for the system shown in eqn (3.6):   



[PtCl3(H2C=CH2)]− + L ⇌ trans-­[PtCl2(H2C=CH2)L] + Cl−

(3.6)

  

and observed that K increases in the following sequence of L (K values are relative): H2O < F− S2O32− (R = alkyl group). In fact, log k for replacement of the aqua ligand trans to L versus Co–OH2 bond length is strictly linear (Figure 3.6).17b Cardwell30 was the first to offer a non-­thermodynamic explanation for the trans effect phenomenon. He suggested that the square-­planar complexes of Pt(ii) undergo substitution by a displacement SN2 mechanism. The original complex having a d x 2  y 2 spx p y configuration has a vacant pz orbital in Pt(ii), close in energy to the bonding orbitals. Addition of the incoming ligand to Pt(ii), utilizing a pz orbital, is highly likely to give rise to an intermediate species of coordination number 5 with d x 2  y 2 sp x p y pz bonding. In order to account for most of the observed facts regarding trans substitution in Pt(ii) complexes, Cardwell suggested that owing to the electrostatic repulsion from the approaching ligand, the two most electron-­repelling trans groups are forced slightly out-­of-­plane, to form a sort of trigonal bipyramid structure for the intermediate, in which these electron-­repelling groups along with the approaching ligand occupy the positions in the trigonal plane (Figure 3.10). However, to account for the difference in the behaviour of [PtCl3 (H2C=CH2)]− and [PtCl3(NH3)]− (substitution trans to H2C=CH2 in the former and cis to NH3 in the latter), one has to make the rather unusual suggestion that H2C=CH2 is more electron repelling than Cl−. Hence this mechanistic theory is also not satisfactory in the case of H2C=CH2, which has a high trans effect. Another objection to this theory is that the hybridization scheme proposed by Cardwell, i.e. d x 2  y 2 sp x p y pz is believed to lead to a square-­pyramidal configuration, whereas the trigonal bipyramid is formed as a result of dz 2 spx p y pz hybridization. If this is so, formation of the trigonal bipyramid

Figure 3.10  Cardwell's  mechanism30 to explain the trans effect: [PtCl(NH3)3]+

reacting with Cl− forms trans-­[PtCl2(NH3)2]; in this case L = Cl−, X = NH3, A = NH3, Y = Cl−. Again, [PtCl3(NH3)]− reacting with NH3 gives cis-­ [PtCl2(NH3)2]; in this case L = Cl−, X = Cl−, one A = NH3, other A = Cl−, Y = NH3.

Ligand Replacement Reactions of Metal Complexes

119

necessitates a considerable change in the hybridization scheme and cannot occur merely by accommodating the ligand through the pz orbital. Moreover, transformation of the square-­planar complex to the trigonal bipyramidal intermediate requires considerable distortion of a trans-­bonded pair, from an angle of 180° to 120°. Both processes would require considerable energy and hence kinetically would be less favoured normally than the formation of the square pyramid, unless under special conditions (see Section 3.1.4). However, such a transformation is not an impossibility (see below) and may occur concurrently during entry of the ligand tending to form a square-­ pyramidal intermediate. Dyatkina and Syrkin31 proposed this geometry for the transition state for the reactions of Pt(ii) complexes on such energy considerations. According to them, substitution occurs by a series of changes as shown in Figure 3.11, path (a), while Haake32 suggested changes as shown in path (b). Chan and Wong33 proposed a similar scheme, but the transformation of the square-­pyramidal intermediate PtA2LXY having Y in an apical position formed is believed to take place through a trigonal bipyramid (Berry pseudo-­rotation) and then a square pyramid, in which X occupies an apical position and Y becomes equatorial, as shown in Figure 3.12. Incidentally, it should be mentioned that the theory may be extended to account for the catalytic action of Pt(ii) in the ligand substitution reactions of Pt(iv) complexes. For this it was first postulated that the Pt(ii) accepts a ligand from the Pt(iv), so that each becomes square pyramidal, followed by the addition of the reagent in the vacated coordination position of Pt(iv) (Scheme 3.2). However, actually, according to subsequent evidence, ligand replacement and ligand exchange in Pt(iv) complexes [PtIVX2A4]2+ catalysed by [PtIIA4]2+ proceed by an electron transfer mechanism of inner-­sphere type34 (see Chapter 7), as shown in Scheme 3.3 (there is simultaneous Pt exchange also).35 Ligand replacement can also occur similarly (in this case X* = Y). Hence the Pt(iv) product contains Pt which was originally Pt(ii), hence this is not really a

Figure 3.11  Mechanisms  of ligand substitution in Pt(ii) complexes: trans-­[PtA2LX] + Y → trans-­[PtA2LY] + X. (a) Dyatkina and Syrkin's mechanism;31 (b) Haake's mechanism.32

Chapter 3

120

Figure 3.12  Chan  and Wong's mechanism showing the steric course of substitution in Pt(ii) complexes.33

Scheme 3.2

Scheme 3.3 ligand replacement reaction at the Pt(iv) centre. The proposal of Basolo et al. that during breakdown of the dinuclear intermediate the bridging Cl− is transferred to the Pt(ii) centre as Cl+ and a consequent transfer of 2e to Pt(iv) that changes it to Pt(ii) is not acceptable on the basis of available evidence that the bridging ligand's role is to facilitate electron transfer from the metal in a lower oxidation state to the metal in a higher oxidation state followed by bond rupture at the less inert site (see Chapter 7). Basolo et al.34a showed that in the system shown in Scheme 3.4, where progress of the reaction could be followed by observing the change in optical rotation with time [since the optical rotation of l-1,2-diaminopropane (l-­pn) is different in Pt(ii) and Pt(iv) species], the t½ value is (within limits of experimental error) equal to that of 36Cl− exchange in trans-­[PtCl2(en)2]2+ in the presence of [Pt(en)2]2+. Basolo et al.34b further showed by using 13C-­labelled en in one of the platinum complexes in the [PtCl2(en)2]2+‒[Pt(en)2]2+‒*Cl− system that exchange of en occurs at the same rate as *Cl− exchange at 25 °C: k Cl– = 15 M−2 s−1, ken = 16 M−2 s−1, kPt = 12.4 M−2 s−1.

Ligand Replacement Reactions of Metal Complexes

121

Scheme 3.4 The reaction of [PtCl(NH3)5]3+ with Cl− catalysed by [Pt(NH3)4]2+, leading to the formation of trans-­[PtCl2(NH3)4]2+ and NH4+, in acidic medium containing Cl− can also be explained in this manner, as also can the catalysis of Cl− exchange in trans-­[PtCl2(en)2]2+ by [Pt(en)2]2+. In agreement with this mechanism is the observation that Pt exchange occurs at the same rate as Cl− exchange in the latter system.35

3.1.3  π-­Bonding Theory of the trans Effect Chatt et al.25 pointed out that any hypothesis of trans-­directed substitution must not only explain the high trans effect of a ligand such as H2C=CH2 but also the fact that the rate of substitution in the presence of a strongly trans-­ directing group is very much faster than the corresponding reaction in the presence of a group having a low trans effect [eqn (3.8) and (3.9)].   



  

slow [PtCl3 (NH3 )]  NH3   cis -[PtCl2 (NH3 )2 ]  Cl 

(3.8)

 fast [PtCl3 (H2 C CH = CH2 )]  Cl   (3.9) 2 )]  NH3  trans -[PtCl 2 (NH3 )(H2 C

  

From consideration of the fact that many ligands that have a high trans effect have suitable unfilled orbitals to overlap with the filled d orbitals of Pt(ii) to form π-­bonds of the metal-­to-­ligand type, Chatt et al.25 and Orgel26 independently suggested that as a result of π-­bond formation of this type, there is a considerable reduction in the electron density in the antinodal regions, lobes a, a′ (Figure 3.13), thereby facilitating nucleophilic attack by the incoming ligand Y in this region, which is close to the trans group X, and eventually therefore it is this group that is replaced with Y. In the case of an olefin, say H2C=CH2, a Pt‒olefin σ-­bond results from the overlap of an unfilled dsp2 hybrid bonding orbital of Pt(ii) with the filled bonding π-­MO (bonding) of H2C=CH2; π-­bonding occurs through overlap of the filled d orbital [dxz or dyz, assuming that the four ligands are disposed around Pt(ii) in the XY-­plane along the X-­ and Y-­axes] of Pt(ii) with the unfilled antibondπ   olefin [Dewar,36 Chatt and Duncanson37 (DCD) ing π-­MO of H2C=CH2, Pt    model for metal‒olefin bonding]. In agreement with this π-­bonding is the well-­known fact that the ion [PtCl3(H2C=CH2)]−, unlike [PtCl3(NH3)]−, shows the same resistance to oxidation as [PtCl5(NH3)]−.38 In the case of ligands such as PR3, AsR3, SR2, etc., a vacant d orbital of the donor atom of the ligand is presumably involved in π-­bond formation. Chatt et al.25 further argued that hybridization of 5d and 6p orbitals of Pt(ii) produces dp orbitals, which have an orientation similar to that of the d orbital

122

Chapter 3

Figure 3.13  π-­  Bonding trans effect theory of Chatt et al.25 and Orgel26 [the shaded

portion represents the distribution of electronic charge in the lobes of the d orbital of Pt(ii) due to π-­bonding by L; in the absence of such π-­bonding there would have been a uniform charge distribution in all the four lobes (including a and a′ of the d orbital)].46 L = PR3: the d orbital of P is unfilled and this is used for π-­bonding (Pt → P). L = H2C=CH2: antibonding π* MO of the H2C=CH2 is unfilled and is used in π-­bonding with Pt(ii).

but are especially favourable for π-­bond formation with suitable ligands, because of more extensive overlap. On the basis of the utilization of hybrid dp orbitals, one can nicely account for the fact that in the series Ni(ii), Pd(ii) and Pt(ii) only Pt(ii) shows a significant trans effect since in case of Pt(ii) the 5d and 6p orbitals are sufficiently close in energy to allow for strong hybridization and strong π-­bond formation, whereas in the cases of Ni(ii) and Pd(ii) the energy difference between the (n − 1)d and np orbitals is appreciable. It is also known that in the series of amines, phosphines and olefins, the tendency to form π-­bonds increases in the sequence amines < phosphines < olefins, which is also the sequence of increasing trans effect of these ligands. On the basis of the π-­bonding theory, the increased rate of trans substitution has been explained as being due to lowering of the activation energy for the formation of the transition state by stabilizing the activated complex,39 rather than any effect of the trans-­directing group on the strength of the trans bond. The lowering of activation energy results from the fact that the trigonal bipyramid configuration in the transition state is more stabilized by π-­bond formation. π-­Bonding by L will reduce considerably the electron density along the directions of the Pt–X and Pt–Y bonds (Figure 3.14) and thus stabilizes the transition state. π-­Bonding by Y will have a similar influence in reducing the electron density along the Pt–L bond; in fact L, X and Y occupy equivalent positions in the trigonal bipyramid, hence even if the entering nucleophile Y is a good π-­withdrawing ligand it will lead to faster rate of substitution. This nicely accounts for the observations of Banerjea et al.47 discussed in Section 3.1.4.

Ligand Replacement Reactions of Metal Complexes

123

Figure 3.14  Orientation  of the dxy and d x 2  y 2 orbitals of Pt(ii) in the trigonal bipyr-

amid transition state [Pt(X)(Y)A2(L)] [for the reaction PtXA2(L) + Y → PtYA2(L) + X], where ligands L, X and Y are in the trigonal plane; the two ligands A, A are in the Z-­axis of the tbp and are not shown. (a) The high electron density due to the dxy orbital along the Pt–X and Pt–Y bond directions will be reduced due to π-­withdrawal of electrons by L from the dxy orbital of Pt(ii). (b) Similarly, π-­bonding by Y (or X) using the d x 2  y 2 orbital of Pt(ii) will reduce the electron density due to filled d x 2  y 2 orbital along the Pt–L bond direction (see Banerjea et al.47 in Section 3.1.4).

However, if the group X in the trans position to L is essentially σ-­bonded to Pt(ii), the strength of this bond will be increased by the strong Pt → L π-­bonding by the trans group, L. The theory further predicts that the substitution in the presence of a strong trans-­directing group must be essentially associative (SN2, A process) in character. However, according to available evidence (Section 3.3), such ligand replacement reactions involve rupture of the Pt–X bond and simultaneous formation of a Pt–Y bond, the latter being predominant. Zvyagintsev and Karandasheva40 observed that for the reaction shown in eqn (3.10), where X− is substituted under the trans influence of Z− [X− = Cl−, Br−; Z− = Cl−, Br−, NO2−; Y− = Cl−, Br−, NO2−; the compounds studied were X− = Y− = Z− = Cl−; X− = Y− = Cl−, Z− = Br−; X− = Cl−, Y− = Z− = Br−; X− = Z− =NO2−, Y− = Cl−, etc.].   

[Pt(NH3)XYZ]− + py → [Pt(NH3)(py)YZ] + Z−



(3.10)

  

The rate was first order in each of the reactants and for X− = Y− = Cl− the rate decreases in the following sequence of Z Z:

H2C=CH2 >>

NO2− >

Br− >

Cl−

k/M−1 s−1 (25 °C): E0/kcal mol−1:

—

56 × 10−3

18 × 10−3

6.3 × 10−3

—

11

17

19

Chapter 3

124 −

For [PtCl3(H2C=CH2)] , the rate was too fast to be measurable by the ordinary techniques used in all other cases. Hence the observed order of decreasing rates is also the order of decreasing trans effect of Z. Again, for [PtCl3(NH3)]− and [PtCl3(py)]− the rate is slightly faster (at 25 °C) for the py complex (8.5 × 10−3 M−1 s−1) than for the NH3 complex (6.3 × 10−3 M−1 s−1), in which the py and NH3 are cis to Cl− being replaced. Grinberg41a called this the cis effect. It has also been reported that Cl− has a greater cis effect than NO2−; replacement of Cl− cis to NH3 is also ca. 20% faster than when cis to NO2−. Hence the cis effect of NH3 is greater than that of NO2− and Cl− > NO2−, but the reverse is true for their trans effect. Kinetic studies suggested a cis effect order py > NH3 > NO2−, which prompted the suggestion that a good trans-­activating ligand will be a poor cis activator.41a,b However, this is not a good generalization, since data are available to show parallel cis and trans effects of ligands in some systems. Thus, the rate of reaction of cis-­[PtCl(L)(PEt3)2] with pyridine increases by a factor of 3 with changes in L in the order11 Cl− < C6H5− < CH3−. The same order is found for the trans effect of these L in the corresponding reaction of trans-­[PtCl(L)(PEt3)2] and the spread in rates is very much larger (170 at 25 °C in ethanol).11 Based on investigations42a,b on the rates of hydrolysis of series of complexes[PtCl4]2− and [PtCl4‒x(NH3)x]2−x (x = 1–3), it was concluded that in such systems having weakly trans-­directing groups, the cis neighbour has a greater influence on the kinetics than the trans neighbour to the leaving group. However, whereas the cis effect is generally small, the trans effect can be very large. Thus, kinetic studies reported by Elding and Groning42c on the replacement of Cl− by H2O in cis and trans positions of L in Pt(ii) complexes [PtCl3L]n− led to the following data regarding the cis and trans effects of different ligands (relative k values are given in parentheses):42d cis effect:

 2C=CH2 (0.05) < Br− (0.3) ≈ Cl− (0.4) < NH3 (1) ≈ H2O (1) < I− (2) H < DMSO (5)

trans effect: H  2O (1) < NH3 (200) < Cl− (330) < Br− (3 × 103) < I− (4 × 104) < SCN− (1 × 105) < DMSO (2 × 106) [PtCl4]2− > trans-­[PtCl(OH)(NH3)2]. At 25 °C, trans-­[PtCl2(NH3)2] reacts only ca. three times faster than the corresponding cis isomer. However, this is not due to the higher stability of the cis isomer. That the trans isomer is more stable than the cis isomer is indicated by the estimated value of ΔH° (ca. −3 kcal mol−1)42a for isomerization of cis-­ [PtCl2(NH3)2] to trans-­[PtCl2(NH3)2]. Results similar to those of Banerjea et al.47 were also reported by Martin et al.42b As opposed to this, [CoCl(NH3)5]2+ reacts more slowly than trans-­[CoCl2(NH3)4]+ by a factor of 103 (see Chapter 4, Table 4.5). The insensitivity of the rates in the case of the Pt(ii) complexes is incompatible with a dissociative (SN1) mechanism, which occurs predominantly by bond breaking. It is instead what one would expect if both bond making and bond breaking are of comparable importance. Moreover, one cannot ignore the well-­known fact that a square-­planar complex will invariably have two groups coordinated from above and below the plane, presumably at a greater distance than the four primary ones; these may be the donor solvent molecules

126

Chapter 3

in solution. A mechanism involving the solvated tetragonal complex had therefore been proposed.47 This can accommodate all the observed facts, for reactions of both categories (i) and (ii) (Figure 3.15). It is evident from the proposed mechanism that in the formation of the reaction intermediate of coordination number five, both bond making and bond breaking are important, irrespective of whether the reaction occurs by path I or path II (Figure 3.15). The investigations reported above established for the first time that ligand replacement reactions of square-­planar complexes of Pt(ii) may proceed by a path that is zero order or first order in the nucleophile. These conclusions

Figure 3.15  General  mechanism of ligand substitution in Pt(ii) complexes.47

Ligand Replacement Reactions of Metal Complexes

127

were based on the observation that on carrying out the reaction in the presence of a high concentration of nucleophile, the observed pseudo-­first-­order rate constant kobs remained practically unchanged on doubling the concentration of the nucleophile in some cases, but in other cases such a change nearly doubled the kobs value. The effect of nucleophile concentration on kobs was studied at only two (in a few cases three) concentrations of the nucleophile. With such limited data, it is not possible to conclude that a reaction that appears to proceed by path I does not have a small contribution of a concurrent path II and vice versa. It is now known that such reactions of all square-­planar complexes in general proceed by paths I and II concurrently; but their relative contributions may vary within wide limits; in some cases the contribution of either of the two paths may be small or even insignificant compared with the other. Based on the results,47 the nucleophilicity order for the nucleophiles which react by path II is SC(NH2)2 > NO2− > NH3 > H2C=CHCH2OH > C2O42−. However, OH−, which is a powerful nucleophile for organic halides and for also for the octahedral complexes of Cr(iii), Co(iii), etc., reacts by path I, and so also do Cl− and py, whereas allyl alcohol reacts with [PtCl4]2− and even trans-­ [PtCl2(NH3)2], and no reaction was observed with [PtCl(NH3)3]+. This is not unexpected because electron donation from a filled orbital of the metal to the antibonding vacant orbital of the olefin (metal-­to-­olefin π-­bonding) is an important component of the metal‒olefin bond; this is surely expected to be much less favoured when the metal is in a cationic complex. This is in fact the reason why no cationic complex having an olefin ligand is known. Analysis of EXAFS spectral features of Ptaq2+ in acidic (HClO4) aqueous solution showed that the Pt2+ has four H2O bonded square planarly with a Pt–O distance of 2.012 Å and has one or two axially bonded H2O with a Pt–O distance of 2.392 Å. cis-­[Pt(NH3)2(OH2)2]aq is similar with two Pt–NH3 and two Pt–OH2 with Pt–O/N = 2.012 Å with one or two axially bonded H2O with a Pt–O distance of 2.372 Å.48 These findings corroborate what Banerjea et al.47 proposed earlier to explain the kinetic features of ligand substitution in platinum(ii) complexes in solution. It has been suggested that both square-­pyramidal and trigonal bipyramidal geometries for the intermediates are involved in the course of the ligand replacement processes in square-­planar complexes.49 Consideration of the decrease in ligand field stabilization energy in the formation of trigonal bipyramidal and square-­pyramidal intermediates (see Chapter 1, Table 1.8) predicts that the square pyramid is energetically more likely to be formed than the trigonal bipyramid. Opposing this is the smaller repulsion between the ligands in the latter configuration. An estimate of the ligand–ligand repulsion has been made for the two structures (ref. 1h, p. 394). Thus, in the case of [PtCl5]3−, the mutual repulsion in the trigonal bipyramid structure is ca. 250 kcal mol−1 and in the square-­pyramidal structure ca. 276 kcal mol−1, i.e. a difference of ca. 26 kcal mol−1. However, since the Dq value for such platinum(ii) complexes is ca. 5 kcal mol−1, the square pyramid is ca. 21 kcal mol−1 more stable than the trigonal bipyramid based on the LFSE values given in

Chapter 3

128

Table 1.8 (Chapter 1). Hence, on such considerations, it is not possible to make a definite prediction regarding the preferred structure of the intermediate. However, π-­bonding will have the effect of reducing the decrease in LFSE, particularly in the case of the trigonal bipyramid, and the superiority of this structure becomes evident when L, X or Y is a good π-­bonding group. In the trigonal plane, two π-­bonds can be formed simultaneously with any two ligands L, X and Y. Also, π-­bonding by any of these groups will stabilize the trigonal bipyramid (Figure 3.14)47 If the five-­coordinate species has a sufficient lifetime to permit its detection, as in the reaction of trans-­ [PtCl(mesityl)(PEt3)2] with py in methanol,50 then it should be considered as a reactive intermediate, otherwise {as in the reaction of [Pt(NO2)(dien)]+ with I−}51 as a transition state. The possibility and consequences of pseudo-­ rotation leading to square pyramid‒trigonal bipyramid interconversion have been reviewed.52 There is ample evidence that the low-­spin square-­planar complexes of Ni(ii), Pd(ii), Pt(ii) and Au(iii) have a definite tendency to add additional ligands to form complexes of coordination number five and even six.53 Thus, in the presence of a large excess of CN− in solution [Ni(CN)4]2− forms [Ni(CN)5]3− (at 25 °C, the formation constant K5 is 0.19 M),54a which was isolated as a salt, [Cr(en)3][Ni(CN)5]·1.5H2O.54b Crystal structure determination of the latter showed that half of the [Ni(CN)5]2− is distorted square pyramidal and the rest is distorted trigonal bipyramidal.54c Some square-­planar low-­spin complexes of Ni(ii) formed by strong-­field ligands can add two ligands in axial positions, forming six-­coordinate high-­spin complexes.54d,e Sacconi et al.54f reported thermodynamic studies on the formation of a number of six-­coordinate complexes from four-­coordinate square-­planar complexes of Ni(ii) by addition of two monodentate ligands in the axial positions. Syntheses of salts of several five-­coordinate anionic Pt(ii) complexes having monodentate ligands have been reported, viz. [Pt(SnCl3)5]3−, [PtH(SnCl3)4]3− and [PtH(SnCl3)2(PEt3)2]− (ref. 1h, p. 377); [Pt(SnCl3)5]3− has been shown to have a trigonal bipyramidal structure.55 An example of a six-­coordinate complex of Pt(ii)56a is [Pt(NH3)4(NCMe)2]Cl2. Also known are [MX2(diars)2] [M = Ni(ii), Pd(ii), Pt(ii); diars = o-­C6H4(AsMe2)2].56b Crystal structure determination of iodo complexes (X = I) showed that these have distorted octahedral structures.56b However, in nitrobenzene solution these complexes behave as uni‒univalent electrolytes, i.e. [MX(diars)2]X.56c Trigonal bipyramidal [PtCl(PEt3)2(phen)]+ and [M(F3CCOCHCOCF3)2L] [M = Pd, L = P(o-­tolyl)3; M = Pt, L = P(cyclohexyl)3], which are square-­pyramidal (having L in an equatorial position with an oxygen donor of a chelated β-­ diketonate in the apical position) are known.57a,b Convincing experimental evidence for a square-­pyramidal intermediate in the reaction of a Pt(ii) complex was obtained in the following radiochloride exchange studies in nitromethane:58   

  

trans-­[Pt(36Cl)2(py)2]+ 2 36Cl− ⇌ trans-­[Pt(36Cl)2(py)2] + 2 Cl−

(3.11)

Ligand Replacement Reactions of Metal Complexes −

129

The rate is independent of Cl in this solvent, but substances such as acetic acid and boric acid increase the rate of exchange greatly. The reaction evidently occurs through path II(a) in Figure 3.15, where Y = OAc− or BO2− and Z− = Cl−. Formation of a trigonal bipyramidal intermediate is expected to lead to a product containing Y (OAc− or BO2− in this case). Such is the case, however, in the reaction of [PtCl(dien)]+ with a solution containing I− and OH−, where reaction occurs in two stages:59 first there is a fairly rapid reaction to form the iodo complex, followed by a slow reaction to form the hydroxo complex by further action on the iodo complex. The initial formation of the iodo complex suggests that here the reaction occurs by path II(b) in Figure 3.15 (Y = I−) through the formation of a trigonal bipyramidal intermediate. Haake32 obtained kinetic and spectral evidence in favour of the formation of an intermediate of higher coordination number [associative nucleophilic displacement in a Pt(ii) complex] in the reaction of cis-­[PtCl2(NH3)2] with NO2−. By potentiometric titration, both stages of substitution (one Cl− first and then the other) could be followed. The second stage showed a non-­linear dependence on NO2− concentration, falling off at high concentration and tending to reach a limiting value. This is indicative of a mechanism shown in Scheme 3.5. Also in agreement with the pre-­equilibrium step was the immediate enhancement of the absorbance of the solution on adding NO2− to a solution of cis-­[PtCl(NO2)(NH3)2]; the increase in absorbance was too fast to be due to the formation of a dinitro product. The magnitude of the increase in absorbance again showed a non-­linear dependence on the concentration of NO2−. The structural changes suggested are shown in Figure 3.11, path (b) (A = NH3, L = NO2−, X = Cl−, Y = NO2−). Evidence for the formation of five-­coordinate intermediates in ligand replacement reactions of complexes of Pt(ii) was also furnished by Chan and Tong.60 Reactions of bis(β-­diketonates) of Pt(ii) [and also of Pd(ii)] with PR3 occur through association forming a square-­pyramidal five-­coordinate intermediate, having PR3 in the equatorial position, many of which have been characterized spectroscopically in solution and some have also been isolated and examined crystallographically.32,57,61 Evidence for an intermediate of coordination number five has also been reported in the replacement of Cl− by py in trans-­[PtCl(mesityl)(PEt3)2].62 A reaction proceeding through path I in Figure 3.15 will naturally show zero order with respect to the entering ligand Y, whereas in path II the rate will be first order with respect to both the complex and the entering ligand. Naturally, which path will predominate will depend on a particular complex, particular the entering ligand and the solvent. This accounts for the observation that NH3 reacts essentially by path I in the case of [PtCl(NH3)3]+ and by path II in the case of trans-­[PtCl2(NH3)2]. It has further been observed

Scheme 3.5

Chapter 3

130

that reagents such as thiourea, nitrite ion and allyl alcohol, which have a high trans effect, always react by path II, the observed order being thiourea > NO2− > NH3 > allyl alcohol > C2O42−, while even the very strong nucleophile OH− reacts by path I (zero order with respect to OH−). The difference in behaviour of OH− and allyl alcohol is very striking and shows that the π-­bonding ability of the reagent is much more important than its nucleophilic strength. This is expected, since, as has been stated earlier, the activation energy for the formation of the transition state (trigonal bipyramid) will be reduced owing to π-­bonding by either L, Y or even X. Based on this π-­bonding concept, it is possible to account for the observation of Grinberg and Nikol'skaya63 that for the exchange of X− with PtX42− the rate decreases in the sequence CN− > I− > Br− > Cl−, which is also the order of decreasing trans effect of these ligands and thermodynamic stability of the complex (see Chapter 1, Section 1.7). Similar observations have also been reported by Grinberg and co-­workers for ligand exchange in the systems [Pt(tu)4]2+–tu(S*) and [Pt(EtNH2)4]2+–EtNH2(C*). The more stable thiourea (tu) complex undergoes exchange very rapidly,64 whereas the exchange in the less stable amine is extremely slow.65 In the original square-­planar configuration, only three π-­bonds can be formed at a time, utilizing dxz, dyz and dxy orbitals of Pt(ii), whereas in the trigonal bipyramidal configuration four π-­bonds can be formed simultaneously, utilizing d x 2  y 2, dxz, dyz and dyy orbitals of Pt(ii). Hence, in the presence of four potential π-­bonding ligands, there will be a pronounced tendency to add a fifth ligand, which will lower the activation energy for the formation of the transition state, leading to rapid exchange; in such a case, the rate increase will be in the order of increasing affinity for π-­bonding, as observed. Data are available that indicate that the reactivity of a π-­bonding ligand depends on the net charge on the complex. This is to be expected, since π-­bonding by the incoming ligand in the transition state will be more favoured with increasing overall negative charge on the complex. The data of Gray and Olcott59 are significant (Table 3.7). It is obvious that the efficiency of NO2− compared with Cl− as a reagent is much greater for the neutral and anionic complexes trans-­[PtCl2(NH3)2] and [PtCl4]2– than it is for the cationic complexes. Indeed, for [Pt(dien)(OH2)]2+ the reactivity order is even inverted. Hence it is justified to call reagents such as NO2−, PR3, thiourea, etc., biphilic reagents since both the ability to donate electrons to the metal and to accept electrons from the metal are of importance.58 Several examples are now known where both paths I and II contribute to the total change, but their relative contributions vary depending on the complex and also Y and the solvent. For the reaction in eqn (3.12):   

  

[PtX(dien)]+ + py → [PtX(dien)(py)]2+ + X−

(3.12)

where dien = H2N(CH2)2NH(CH2)2NH2, Basolo et al.69 observed the rate to conform to the equation   

Ligand Replacement Reactions of Metal Complexes

131

Table 3.7  Second-­  order rate constants (kY) for ligand replacement reactions of complexes of Pt(ii) with Cl− and NO2− as entering nucleophiles, Y.a 104kY/M−1 s−1 Entry 1 2 3 4 5 6 7 8

Y = Clb

Complex 2+

[Pt(dien)(OH2)] [Pt(Et4dien)(OH2)]2+ [PtBr(dien)]+ [PtCl(NH3)3]+ [PtCl4]2– [PtCl2(py)2] [PtCl2(pip)2] [PtCl2(PEt3)2]

4

1 × 10 7.4 × 104 8.8 0.7 ∼0 4.5 9.3 0.29

Y = NO2−

k NO /kCl

Ref.

0.56 1.76 4.2 124.3 Very large 1.5 2.15 0.93

59a 59b 59c 66 67 68 68 68

2

4

0.56 × 10 13 × 104 37 87 104 6.8 20 0.27

a

 or entries 1–5 an aqueous solution at 25 °C and for entries 6–8 a methanol solution at 30 °C F were used.  Cl− was used to study exchange reactions of chloro complexes. Cl− exchange of [PtCl4]2− occurs almost exclusively by the kS path with hardly any contribution from the kY path67 (see Table 3.20).

b36



Rate = k0[complex] + kY[complex][py]

(3.13)

  

where k0 and kY are the rate constants for the reagent (Y)-­independent path (path I) and reagent-­dependent path (path II), respectively. In fact, this is a general behaviour for ligand replacement in square-­planar complexes and the general rate law in eqn (3.14) has been established; however, the k0 or kY path may be virtually non-­contributory in different systems.   



Rate = (k0 + kY[Y])[complex] = (kS + kY[Y])[complex]

(3.14)

  

Eqn (3.14) is a two-­term rate law, one rate independent of the incoming ligand concentration and the other showing a first-­order dependence on the incoming ligand (Y) concentration. Various evidence, including solvent dependence of k0 and absence of a k0 path in non-­coordinating solvents, indicates that the term independent of Y concentration is a path involving the solvent as the incoming nucleophile (hence k0 is denoted kS), and the overall reaction is represented in its simplest form as shown in Scheme 3.6 (the complex is represented as M–X for convenience, omitting the other ligands).70 kS is a pseudo-­first-­order rate constant observed experimentally. For any comparison of nucleophilic efficiency of a solvent (S) with that of any other nucleophile, it is necessary to compare the second-­order rate constant kS/[S] with kY. Based on Scheme 3.6, under pseudo-­first-­order conditions in the presence of excess Y, the observed rate constant is given by   



  

k obs 

k 1 k 2 [Y]  k 3 [Y] k 1  X   k 2 [Y]

(3.15)

Chapter 3

132

Scheme 3.6 If k2[Y] >> k−1[X] (i.e. the k−1 path is virtually absent), which is most likely, particularly in the absence of any X added to the reacting solution, eqn (6.15) leads to   



kobs = k1 + k3[Y]

(3.16)

  

and this is what is observed experimentally in the majority of cases of substitution not only in Pt(ii) complexes but also in square-­planar complexes in general. Eqn (3.16) is identical with eqn (3.14) with k1 = k0, i.e. kS and k3 = kY; the relative contributions of the kS and kY paths vary from one system to another (one of these paths may even be insignificant) for complexes of any particular metal ion such as Pt(ii) and with changes in the nature of the solvent, and also in the case of analogous complexes as one changes from one metal ion to another such as Pt(ii) to Au(iii) (Table 3.10). A typical illustration71 is the ligand (Cl) replacement reaction of trans-­[PtCl2(py)2] with various nucleophiles (Y) studied in methanol at 30 °C. For this, plots of kobs versus [Y] (Figure 3.16) for different Y are linear with different slopes (kY) but the same intercept (kS). Another system in which similar behaviour has been observed is in the replacement of Br− in [PtBr(dien)]+ with Y (Y− = *Br−, Cl−)72 (Figure 3.17), for which kS at 25 °C is 1.3 × 10−4 s−1, being identical with the rate constant for aquation of [PtBr(dien)]+ determined independently under identical conditions.73 It is of interest that even for the following alkyne insertion reaction: trans-­[PtH(MeCN)(PPh3)2]+ + RCCR → trans-­[Pt(MeCN)(RC=CHR)(PPh3)2]+ where R = MeC(O)O−, studied in CDCl3 by 31P NMR spectroscopy, a two-­term rate law was reported.272 There are also known cases where the kS path is absent (kS = 0)74 or insignificant, or kY = 0 46 or insignificant. Examples are known for which in good coordinating solvents the kS path predominates, but in poor coordinating solvents kS becomes insignificant, as in 36Cl− exchange with trans-­[PtCl2(py)2] in different solvents (Table 3.8).54 From the solvent effect, the steric effect, the effect of overall charge of the complex on the rate of ligand replacement in Pt(ii) complexes, and a variety of other evidence, it has been concluded that the kS and kY paths involve an associative mechanism. The most important to note in this connection are the ∆V‡ values for the kS and kY paths for a number of complexes of Pt(ii). For [PtX(dien)]+ (X− = Cl−, Br−, I−, N3−) reacting with Y− (Y− = OH−, I−, N3−, NO2−, SCN− and also py) the ∆V‡ values are −15 to −8

Ligand Replacement Reactions of Metal Complexes

133

Figure 3.16  Plots  of pseudo-­first-­order rate constants (k/s−1 at 30 °C) versus [nucleophile] for reactions of trans-­[PtCl2(py)2] in methanol.71 Intercepts (kS) are identical.

cm3 mol−1 for the kS path and −6 to −12 cm3 mol−1 for the kY path.75,76 The values are in conformity with a mechanism in which Pt–X bond dissociation and Pt–Y (or S) bond formation are synchronous in the transition state: X⋯PtII ⋯Y (or S). This is also evident from the fact that the rates of such reactions of complexes of Pt(ii) are fairly insensitive to the overall charge on the complex.42b,47 Hence a more appropriate reaction scheme is shown in Figure 3.15, which can be represented in a simpler fashion (Scheme 3.7)77 taking into consideration the fact that platinum in the complex has two weakly bonded axial ligands (H2O) in solution47 (see earlier in Section 3.1.4). In this scheme, as the Pt–X bond breaks two bonds are formed synchronously with the axial ligands, so the process is associative but not of the limiting type. The highly negative ∆V‡ for kS is due to considerable solvation (electrostriction effect) of the leaving group X− in the transition state. However, in the kY path desolvation of the Y− as it becomes partly bonded to Pt(ii) in the transition state will make a positive contribution to the overall ∆V‡, making this less negative than for the corresponding kS path. For the reaction

Chapter 3

134

Figure 3.17  Plots  of pseudo-­first-­order rate constants (k/s−1 at 25 °C) versus [anion]

for reactions of [PtBr(dien)]+ with Y (Y = Cl−, *Br−) in water. Intercepts (kS) are identical; different slopes are values for kY.72

Table 3.8  Effect  of solvent on the rate of exchange of 36Cl− with trans-­[PtCl2(py)2] at 25 °C.54

Solvents in which rate is independent of Cl− Solvent DMSO H2O MeNO2 EtOH n PrOH

105kS/s−1 380 3.5 3.2 1.4 0.42

Solvents in which rate is dependent on Cl− Solvent CCl4 C6H6 t BuOH EtOAc Me2CO DMF

k Cl /mol−1 s−1 

4

10 102 0.1 0.01 0.01 1 × 10−3

Scheme 3.7 of the sterically hindered [PdX(Et4dien)]+ (X− = Cl−, Br−, I−, N3−, NCS−) with Y− (Y− = N3−, Br−, I−) reacting by the kS path, only highly negative values of ∆V‡ have been observed, which for any particular X− is independent of Y−, with ∆V‡ varying linearly78 with ∆V for the change

Ligand Replacement Reactions of Metal Complexes

135 −

[PdX(Et4dien)] + H2O → [Pd(Et4dien)(OH2)] + X +

2+

The slope of this plot suggests ca. 50% bond breaking in the transition state. The reaction of [PtCl2(bipy)] with py proceeds exclusively by the kY path.70 Even for the same complex, the relative contributions of the kS and kY paths depend on the nature of the solvent (S), which strongly suggests that kS represents a path comparable to the kY path with S acting as the attacking nucleophile. Thus, the ligand exchange reaction of trans-­[PtCl2(PPr3)(*NHEt2)] with NHEt2 occurs exclusively by the kS path in methanol and exclusively by the kY path in the non-­coordinating n-­hexane.80 A comparison of the reactions in eqn (3.17) and (3.18) also provides information concerning the nature of the kS and kY paths (quoted k values are at 25 °C). Since on going from the reaction in eqn (3.17) to that in eqn (3.18) both kS and kY decrease by a comparable magnitude (∼103), both paths must be occurring by similar mechanisms.81   



cis-­[PtCl2(PPr3)2] + py → cis-­[PtCl(PPr3)2(py)]+ + Cl−

(3.17)

  

(kY = 1.66 M−1 s−1, kS = 0.83 × 10−2 s−1 in ethanol);   



trans-­[PtCl2(PPr3)NHEt2] + py → trans-­[PtCl(PPr3)(py)] + NHEt2 (3.18)

  

(kY = 0.63 × 10−3 M−1 s−1, kS = 0.86 × 10−5 s−1 in methanol). Taking either kS or kY, the following order of decreasing ease of removal of X− was observed:59,69 NO3− > H2O > Cl− > Br− > I− >> N3− > SCN− > NO2− > CN−. A factor of ca. 105 (at 25 °C) was observed between the rate constants for H2O and CN− replacement, with the replacement of NO3− very much faster than that of H2O. These observations indicate that trans-­activating groups are themselves difficult to dislodge. The order of decreasing lability is also the order of increasing stability, and hence also more or less the order of increasing strength of the Pt–X bond. In the ease of replacement of X in [PtX(dien)]+ by Y the order of decreasing ease of replacement of X has been reported71b to be Cl− ≈ Br− ≈ I− >> N3− > SCN− > NO2− > CN−. From observations on the rate of approach to the following equilibrium:   



trans-­[PtCl(PEt3)2(L)]+ + py ⇌ trans-­[Pt(PEt3)2L(py)]2+ + Cl−

(3.19)

  

followed conductometrically in ethanol solution, Basolo et al.11 obtained the following rate law for the approach to equilibrium:   



Rate = k1[complex] + k2[complex][py]

(3.20)

  

where both k1 and k2 are composite quantities; the reactions are reversible, hence both forward and reverse rates contribute to the overall observed

Chapter 3

136

rates. From the values of k1 and k2, they derived the following relative trans-­ labilizing influence of L: Cl− < biphenyl < p-­methoxyphenyl ≈ p-­chlorophenyl < C6H5− < CH3− < PEt3 ≈ H−. Furthermore, for L = C6H5− the rate constant (at 25 °C) is ca. 35 times the value for L = 2,6-­Me2C6H3− in which two methyl groups occupy positions above and below the plane of the complex. However, for complexes of the cis series, the corresponding factor is ca. 105. This effect of increased steric hindrance in reducing the rate is more compatible with a sort of SN2 mechanism rather than a simple SN1 mechanism. Similar conclusions were arrived at from the observation that rates of reactions of Pt(ii) complexes are rather insensitive to the overall charge of the complex, so that both bond making and bond breaking must be of comparable importance (see earlier in Section 3.1.4). The observation that cis blocking is much more effective than trans blocking is more compatible with a trigonal bipyramidal geometry for the transition state. From studies on the rate of the replacement of Br− in [PtBr(dien)]+ with different reagents, Gray82 arrived at the following sequence of the reactivity of Y (based on kY values, kobs = kS + kY[Y]): thiourea > SCN− > I− >> N3− > NO2− ≈ py > Cl−. Here also no OH−-­dependent path was observed for reaction with OH−. From kinetic studies on the replacement of a Cl− in trans-­[PtCl2(py)2] in methanol solution by different reagents, Belluco et al.71 established the following sequence of reactivity of the reagents (based on kY values): MeO− < Cl− ≈ NH3 < py < NO2− < N3− < NH2OH ≈ N2H4 < Br− < PhSH < I− < SCN− < SO32− Cl−, but for Y = I− the order is Me > Cl− > Ph. In view of this, it was suggested that trans effect can be defined as the influence of the trans ligand on the ability of the substrate to discriminate between entering nucleophiles. The following order was arrived at from studies83b–d,84 on replacement in aqua complexes: [PtClx(OH2)4−x]2−x+ (x = 4–0) (relative k values): H2O (∼10−3) < DMSO (0.3) ≈ H2C=CH2 (0.4) < Cl− (1) < Br− (4) < SCN− (40) < I− (100). In the Cl− exchange in trans-­[PtCl2(py)2] in various solvents72 (studied using 36 Cl), it was found that the rate is independent of Cl− in good coordinating solvents (i.e. here the kS term only is important) and for such solvents (S) the observed order is EtOH < MeNO2 ≈ H2O Br− > Cl−, is maintained in all of the solvents although for both Br− and Cl− the reactivity increases on changing the solvent from MeOH to Me2CO, but the opposite is true for I−. Hence, although the solvating power of the solvent (S) is important, the polarizability of the nucleophile Y also makes a significant contribution to its reactivity towards Pt(ii) in the complex. All these results suggest associative mechanisms for both the kS and kY paths.85 Values for the activation parameters, ∆H‡ and ∆S‡, for ligand replacement reactions of platinum(ii) complexes, such as the aquation of [PtCl4]2−, [PtCl3(NH3)]−, cis-­[PtCl2(NH3)2], trans-­[PtCl2(NH3)2] and [PtCl(NH3)3]+, reported by Martin et al.,42b and replacement of X− in [PtX(dien)]+ (X− = Cl−, Br−, I−, N3−, SCN−, NO2−) with different reagents (H2O, Br−, I−, N3−, SCN−, SeCN−, S2O32−, Table 3.9  Values  of rate constant kY in different solvents (S) at 25 °C for the reaction.85

trans-­[PtCl2(pip2)] + Y → trans-­[PtClY(pip2)] + Cl− 103kYa/M−1 s−1 Y

MeOH

Me2CO

DMSO

DMF

MeCN

MeNO2

S Cl− Br− I− SCN− SC(NH2)2

1.2 0.9 6.2 300 400 3500

8.0 3.5 80 165 375 10 600

70 1.0 5.2 18 15 480

4 0.56 6.7 27 — 780

3 1.1 13 86 — 910

2.5 1.5 15 78 — 710

105kS/s−1 for Y = S.

a

Chapter 3

138

thiourea) and similar reactions of trans-­[PtCl2L2] (L = PEt3, AsEt3, piperidine) in methanol solution and several other such systems, reported by Belluco et al.,71b have shown that the increase in rate parallels the decrease in ∆H‡, with ∆S‡ strongly negative, ranging from −16 to −34 eu (1 eu = 1 cal K−1 mol−1) for the cases mentioned in Table 3.10. Large negative ∆S‡ and rather small ∆H‡ values (4–21 kcal mol−1 in all cases in Table 3.10) suggest that the formation of the transition state is accompanied by a net increase in bonding, thus indicating an associative mechanism for the reactions. According to the proposed mechanism (Scheme 3.7), as a Pt–X bond breaks, two Pt‒ligand (ligand = S or Y) bonds are formed and this accounts for large negative ∆S‡ and small ∆H‡ values observed in these reactions. On plotting ΔH‡ versus ΔS‡ for the reactions in Table 3.10 (Figure 3.18), six distinct isokinetic trends (see Chapter 2) are observed, with a few exceptions. Reactions of each such group presumably differ from the others with respect to the relative bond-­making and bond-­breaking contributions in reaching the transition state of the reaction. Values of the activation volumes, ∆VS‡ and ∆VY‡ for the kS and kY paths, respectively, for the following reactions:   



[PtX(dien)]+ + Y → [PtY(dien)]+ + X−

(3.22)

  

where X− = Cl−, Br−, I−, N3−; Y− = OH−, I−, N3−, SCN−, NO2−, SCN− (and also py) are significantly negative (∆VS‡ = −15 to −18 cm3 mol−1; ∆VY‡ = −6 to −12 cm3 mol−1),75,76 and this is also in accord with the proposed mechanism (Scheme 3.7) with significant net bond making.77 The proposed mechanism implies that the coordination number of Pt(ii) increases in the transition state and available thermodynamic data show that such a process is always accompanied by a significant negative ΔS° value. Thus, Sacconi et al.54f reported studies on the following equilibrium: [Ni(DBH)] + 2L ⇋ [Ni(DBH)L2] where DBH = diacetylbisbenzoylhydrazine-­2H, a quadridentate 2N,2O donor ligand(2−), and L is a phosphine or an amine, for several L and reported for these equilibria ΔH° values of 13–17 kcal mol−1 and ΔS° values of −18 to −42 cal K−1 mol−1. However, for [Pt(OH2)4]2+ it is believed86 that there is a gradual change in the mechanism from Id to Ia as the entering ligand changes in the order H2O, DMSO, Cl−, Br−, I− and SCN−; for the last four ligands the rate constants for replacement of H2O are much larger than that for replacement by DMSO, for which the rate is comparable to that of water exchange, suggesting an Id process, but Ia for Cl−, Br−, I− and SCN−. Belluco and co-­workers87 studied the reaction of [PtCl3(olefin)]− with bipy in aqueous 95% methanol. They observed a two-­step reaction when the olefin is ethylene (Scheme 3.8).87a

Ligand Replacement Reactions of Metal Complexes

139

Table 3.10  Activation  parameters for reactions of some complexes of Pt(ii) in aqueous solutions, except those where methanol is mentioned.a

Complex 2−

Nucleophile, Y ∆H‡/kcal mol−1 ∆S‡/cal K−1 mol−1

H2O H2O H2O H2O H2O Br− [PtCl(dien)]+ H2O Br− N3− I− SCN− SC(NH2)2 [PtBr(dien)]+ H2O N3− I− SCN− SC(NH2)2 [PtI(dien)]+ N3− SCN− SC(NH2)2 [PtN3(dien)]+ I− SCN− SeCN− SC(NH2)2 S2O32– + [Pt(SCN)(dien)] H2O [Pt(NO2)(dien)]+ SC(NH2)2 trans-­[PtCl2(PEt3)2] NO2− (in MeOH) N3− SCN− SeCN− trans-­[PtBr2(PEt3)2] NO2− (in MeOH) N3− NO2− trans-­[PtI2(PEt3)2] (in MeOH) N3− trans-­[PtCl2(pip)2] *Cl− (in MeOH)b Br− NO2− N3− SCN− SeCN− trans-­[PtCl2(AsEt3)2] SCN− (in MeOH) SeCN− PtCl4 [PtCl3(NH3)]− cis-­[PtCl2(NH3)2] trans-­[PtCl2(NH3)2] [PtCl(NH3)3]+

a

21.0 19.0 20.0 20.0 18.0 17.0 20.0 13.0 16.0 11.0 10.0 8.5 19.5 15.5 11.0 9.5 8.5 15.5 10.5 9.0 14.5 14.0 13.0 13.0 9.5 15.5 13.5 15.1 15.5 10.0 8.5 12.4 13.0 11.0 12.5 14.7 13.8 13.7 10.8 9.1 8.4 11.2 7.9

−16 −23 −22 −19 −26 −16 −18 −25 −17 −25 −28 −31 −17 −17 −25 −27 −29 −16 −25 −27 −19 −22 −21 −24 −29 −14 −30 −31 −24 −28 −27 −34 −24 −30 −16 −21 −23 −25 −33 −30 −28 −23 −28

Entry no. 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43

 ata from ref. 71b in which data of Martin et al.42b for reactions (1)–(6) are also quoted. The D ∆H‡ and ∆S‡ values correspond to kY (M−1 s−1), hence for reactions with the solvent H2O the observed pseudo-­first-­order rate constants kS (s−1) were divided by the molar concentration of the H2O (55.5 M) to obtain the second-­order rate constants (in M−1 s−1), and these values were used for the evaluation of ∆H‡ and ∆S‡ in these cases. b pip = piperidine.

Chapter 3

140

Figure 3.18  Plot  of ∆H‡ versus ∆S‡ for the data in Table 3.10.

Scheme 3.8 However, for olefins such as allyl-­NH3+, allyl-­PEt3+, allyl-­SO3− and pent-­4-­ enyl-­NH3+, the experimental data conform to the following equation:87b

  

k obs 



  

k 1 k 2 [bipy] k 1 Cl    k 2 [bipy]

(3.23)

Hence   

1



  

k obs

[Cl  ] 1  k 1  [Cl  ]    k 1  k 2  k 1  k 1 k 2  [bipy] bipy 

(3.24)

conforming to the mechanistic scheme shown in Scheme 3.9 for the overall change (S = solvent).

Ligand Replacement Reactions of Metal Complexes

141

Scheme 3.9 Table 3.11  Values  of rate constants (see text). Olefin +

Allyl-­NH3 Allyl-­PEt3+ Pent-­4-­enyl-­NH3+ Allyl-­SO3−

102k1/M−1 s−1

k2/k−1

0.33 0.32 2.7 1.1

7.4 × 104 5.4 × 104 1.4 × 104 8.2 × 104

The k′1 and k2− values were evaluated graphically using eqn (3.24) and from there the k1 = 1/k′1 and k2/k−1 = k′1/k2− values were obtained (at 25 °C) (Table 3.11). A wide range of rates have been reported in nucleophilic substitution reactions of [PtCl4]2− with a variety of nucleophiles, such as H2O and 36Cl−,42b OH− and allyl alcohol,47 NH3,88a diethylenetriamine (dien),88b 2,2′-­bipyridine (bipy),88c 1,10-­phenanthroline (phen),89 a series of olefins,90 sulfoxides,91 thiourea,92 and anions such as NO3−, Br−, I−, SCN−, SeCN−, S2O32−,92 and SnCl3− 93 (see also Table 3.13). It is worth mentioning that while Cattalini et al.92 reported that in the reaction with SCN− there are both kS and kY paths, according to later work94 the reaction proceeds by the kY path only (kS = 0). Reactions with other nucleophiles, DMSO, thioxan [cyclo-­S(CH2CH2)2O] and Et2S,95a propane-­1,3-­diamine (pn)95b and acetonitrile95c have been reported. Exchange of [PtCl4]2− with 36Cl− in solution is very slow, but is catalysed by 1e− oxidants such as Ce(iv) but not by 2e− oxidants such as Cl2. Addition of a very small amount of Ce(iv) to a solution of [PtCl4]2− in chloride medium generates the Pt(iii) complex [PtCl5]2−, which serves as the catalyst. Pt(iii) being d7, the [PtCl5]2− is very labile and undergoes rapid exchange with 36Cl− in the solution to attain an equilibrium distribution of 36Cl−. This forms a Cl−-­bridged dinuclear species with [PtCl4]2− in which inner-­sphere electron transfer from Pt(ii) to Pt(iii) followed by breakdown of the dinuclear species occurs, forming [PtCl4]2− with the 36Cl− and the [PtCl5]2− and catalyst, and the process continues.96a In aqueous solution, [PtCl3(H2C=CH2)]− aquates to an equilibrium mixture of [PtCl3(H2C=CH2)]− and trans-­[PtCl2(H2C=CH2)(OH2)]. This exchanges with 36 − Cl in solution in which there is a second-­order path in the aqua complex obviously involving a bridged dinuclear species, trans-­[(H2C=CH2)ClPt(µ-­Cl)2 PtCl(H2C=CH2)].96b Similarly, in *Br− exchange of [PtBr4]2− there is a path that is second order in the Pt(ii) complex and involves a bridged dinuclear Pt(ii) species (ref. 1h, p. 410).

Chapter 3

142

It is worth noting that there are some known cases of reactions of square-­ planar complexes where the k0 path does not correspond to kS path, as in the reaction shown in eqn (3.25).97 The first step (rate determining) involves intramolecular transformation of the substrate into the three-­coordinate [Pt(COMe)(I)(PPh3)], which then adds the L in a fast step forming the product.   



[PtMe(I)(CO)(PPh3)] + L → [Pt(COMe)(I)(L)(PPh3)]

(3.25)

  

where L = AsPh3, As(4-­MeC6H4)3, MeAs(2-­MeC6H4)2, SbPh3. In the ligand exchange of Pd(acac)2, the k0 path corresponds to one-­ended dissociation of a chelate ring in which there is no participation of the solvent.98 The mechanism of substitution in square-­planar complexes is modified by steric parameters.88b,89a Kinetic studies on the reaction shown in eqn (3.26) have been published;99 this illustrates how it may often be possible to evaluate the component rate constants in a complex system.   



[PtCl4]2− + L → [PtCl2(L)] + 2Cl−

(3.26)

  

where L is amidinothiourea (atu), H2NC(S)NHC(NH)NH2, which is a bidentate ligand that binds to Pd, Pt, etc., through sulfur and nitrogen of =NH.100– 102 In the presence of fairly high concentrations of H+, Cl− and L, compared with [PtCl4]2−, the following rate law, which is typical for square-­planar complexes, has been observed:   



  

Rate = k0[(PtCl4)2−] + kL[(PtCl4)2−] [L]

(3.27a)

kobs = k0 + kL[L]

(3.27b)

Hence   

  

The results of detailed studies are in conformity with Scheme 3.10. In Scheme 3.10, the transformation of II to III involves replacement of the H2O ligand because the Pt(ii) being a ‘soft’ centre has a greater affinity for Cl− than H2O; the cis effect of the Cl− in a cis position to the H2O also favours this process. However, the transformation of V to VI involves replacement of Cl− cis to the OH− ligand because of the greater cis effect of OH− as it is a better π-­donor than Cl− (for such a cis effect, see Chapter 4). Under the experimental conditions, the log KH value evaluated by the usual pH-­metric method is 5.66 ± 0.03. This means that in the range of acid concentration used in the experiments (0.08–0.25 mol dm−3), virtually all the amidinothiourea will exist in the monoprotonated form (LH+). Assuming a steady-­state concentration for the intermediate species II, i.e. [PtCl3(OH2)]−, its concentration will be as shown in eqn (3.28):   



  

  II 

k 1 [I]

k 2 [LH ]

 k 1 [Cl  ]



(3.28)

Ligand Replacement Reactions of Metal Complexes

143

Scheme 3.10 Therefore,   

V 

  

k 1K a [I ] k 2 LH   k 1 Cl   [H ]

(3.29)



Based on Scheme 3.10, under pseudo-­first-­order conditions, i.e. [PtCl4]2− > k−1[Cl−], as is likely, then eqn (3.30) transforms to   



  

 k k K  k obs k 3 LH   k 3 Cl    k 1 1 2  a  k H  

2



 

At fixed [H+] and [Cl−], eqn (3.31) corresponds to eqn (3.27b) with    k k K   k L k 3 and k 0 k 3 Cl    k 1  1 2  a k 2 H    

(3.31)

(3.32)

Chapter 3

144 −

+

+

when [Cl ] and [H ] are held constant, a plot of kobs versus [LH ] is linear with slope S1 = k3 and intercept I1:    k k K  I 1 k 3 Cl    k 1  1 2  a (3.33) k 2 H    

A plot of I1 versus [Cl−] at fixed [H+] is also linear with slope S2 = k−3 and intercept I2:    k k K I k1  1 2  a (3.34) 2 k 2 H    

At fixed [Cl−], a plot of I1 versus 1/[H+] is also linear with slope S3 = k1k′2Ka/k2 and intercept 13 = k1 + k−3/[Cl−]. From the linear plot of 13 versus [Cl−] also the k1 value is obtained as the intercept (I4) and the k−3 value from the slope (S4). In this way, values of k3 and a set of two values for k−3, k1 and k1k′2Ka/k2 were obtained; in each case the two values were in good agreement. From the k1k′2Ka/k2 and k1 values, the value of k′2/k2 could be obtained from a knowledge of Ka determined by measurement of H+ concentration (from measured pH and referring to a pH versus [H+] calibration curve) of equilibrated solutions of PtCl42− having different concentrations of Cl− under the experimental conditions used in measurements of the rates of reaction; in this procedure, both K1 and Ka could be evaluated graphically using the following equation, which can be easily derived:94   



{TPt − [H+]}/[H+]2 = TCl/K1Ka + 1/Ka

(3.35)

  

where TPt is the sum of the concentrations of Pt(ii) complexes (hence equal to the concentration of PtCl42− used) and TCl is the concentration of the added Cl−. The values of KH, Ka and K1 are given in Table 3.12 and the rate constants in Table 3.13. From the experimentally determined values of k1 and K1, the k−1 values are 10−2. The aqua species [PtCl3(OH2)]− is a very weak acid; at 35 °C and an ionic strength of 1 M, its pKa is 7.18. Values of the rate constants k3, k−3 and k1 and the corresponding ∆H‡ and ∆S‡ values are given in Table 3.12 along with the values of k′2/k2. The ∆H‡ and ∆S‡ values in Table 3.12 are based on the reported99 values in SI units, viz. kJ mol−1 and J K−1 mol−1, respectively. The k1 value of 2.19 × 10−4 s−1 at 35 °C (at ionic strength 1 M) is in fair agreement with the calculated value of 1.78 × 10−4 s−1 (at ionic strength 0.3 M) Table 3.12  Values  of the kinetic parameters (see text); ionic strength 1 M. Temperature/°C

k3/M−1 s−1

102k−3/M−1 s−1

104k1/s−1

10−7k′2/k2

35 40 45 ∆H‡/kcal mol−1 ∆S‡/cal K−1 mol−1

0.14 0.29 0.55 26.4 ± 0.3 23 ± 0.8

1.03 1.73 2.66 17.97 ± 0.5 −11.16 ± 1.7

2.19 3.96 6.59 20.98 ± 0.4 −7.6 ± 1.4

5.45 7.65 9.05

Ligand Replacement Reactions of Metal Complexes

145

Table 3.13  [PtCl  4] –L reactions, kL values (for L = atu, kL is k3, Scheme 3.10). 2−

L Amidinothiourea (atu) H2O H2C(CH2NH2)2 Et2S MeCN (H2N)2CS

Temperature/°C 25 55 25 25 25 50 55

kL/M−1 s−1 a

0.032 2.08a 6.85 × 10−7b 2.4 × 10−3 0.12 4.5 × 10−3 2.0

Ref. 99 99 67 103 104 105 92

a b

 alculated values from the reported ∆H‡ and ∆S‡ values.99 C This is the pseudo-­first-­order rate constant divided by 55.5 (molar concentration of H2O in the dilute solution) to convert to the second-­order form to be compatible with the kL values for other nucleophiles.

based on the ∆H‡ and ∆S‡ values of 21 kcal mol−1 and ‒8 cal K−1 mol−1, respectively, reported by Martin and co-­workers67 from studies on the exchange of [PtCl4]2− with 36Cl−. Since the ∆H‡ values are almost identical, the observed difference in k1 values is due to small differences in ∆S‡ values due to differences in ionic strength. The data in Table 3.13 show that the reactions of [PtCl4]2− with S donor ligands proceed very much faster than those with N donors. Hence the kL value for reaction with amidinothiourea (atu) suggests that the rate-­determining step involves the formation of a Pt–S bond followed by a much faster ring closure forming a Pt–N bond (through N of the =NH) leading to the product [PtCl2(atu)]. Since the first bond formation (rate determining) is with S of the thiourea moiety of amidinothiourea, the calculated value of kL (using the ∆H‡ and ∆S‡ values) is almost the same as that of the reaction with thiourea, being 2.08 and 2.0 M−1 s−1, respectively, at 55 °C. The observed ratio of ∼107 for k′2/k2 for formation of [PtCl2(atu)] shows that the species [PtCl3(OH)]2− is very much more reactive than [PtCl3(OH2)]−. The reaction discussed above is an example of complex formation by a bidentate ligand and it has been proposed with reasoning that the first bond formation is rate determining followed by a faster ring closure. This is the usual feature of such reactions according to available data106–109 and more similar information is available in the cases of octahedral complexes (see Chapter 5). Thus, in the following reaction in methanol, containing sufficient HCl to prevent interference from hydrolysis products, the first step is rate determining:87a   

   

(3.36)

Chapter 3

146

It is possible to compare the ring closure with a process involving a unidentate ligand [eqn (3.37) and (3.38); values on the arrows are rate constants at 25 °C]. Ring closures are over 103 times faster than unidentate binding in 1 M NH3. Considering the greater nucleophilicity of ethylenediamine compared with ammonia, the ring closure is ca. 103 times faster and this is partly due to the higher effective concentration of free –NH2 of the unidentately bonded ethylenediamine nearer the Cl− replaced in the ring closing step [eqn (3.37)]. However, other specific effects may have a role.   



(3.37)

  



(3.38)

  

Such a specific influence has been described in comparing the rate characteristics of the following reactions in methanol.109 Only one slow stage is seen for the reaction shown in eqn (3.40), so the formation of Pt–L–L is rate determining. This is further supported by similar rate constants for the reactions shown in eqn (3.39) and (3.40). However, the activation parameters (Table 3.14) are different, ranging from 15 to 17 kcal mol−1 (∆H‡) and from −12 to −17 eu (∆S‡) for different L in the reaction shown in eqn (3.39), and Table 3.14  Rate  constants (k) and activation parameters for the reaction of [PtCl2(bipy)] with unidentate and bidentate ligands in methanol at 25 °C.109

Ligand

103k/M−1 s−1

∆H‡/kcal mol−1

∆S‡/eu

py 4-­Mepy n BuNH2 en tn Piperazine TMEDA (Me2NCH2CH2NMe2)

5.3 5.4 25 60 53 15 0.14

16.3 15.6 15.7 11.2 11.7 12.2 17.9

−15 −17 −14 −28 −27 −28 −18

Ligand Replacement Reactions of Metal Complexes

147

−1

from 11 to 12 kcal mol and ‒27 to −28 eu, respectively, for different L–L in the reaction shown in eqn (3.40). These differences are ascribed to increased bonding and compactness in the species in the transition state 3 associated with the reaction shown in eqn (3.40).   



(3.39)

  



(3.40)

  

In the first step of a number of chelation reactions, the k0 (i.e. kS) path of the two-­term rate law is missing or negligible compared with kY; this may be due to the comparatively rapid solvolytic path compared with the chelation or to unreactivity of the species formed by solvolysis.110 Ring opening of chelates appears to occur at only a slightly slower rate than cleavage of unidentate ligands from the metal.108 Comparison of data for [Pt(en)2]2+ and [Pt(NH3)4]2+ indicates that the enhanced stability of chelated en over NH3 complexes resides largely (103) in enhanced ring closure and only slightly (10) in decreased cleavage (ring opening).108 The rate-­determining step in the replacement of PhSCH2CH2SPh from Pd(ii) with py is ring opening by the entering nucleophile. When the py has two ortho substituents, the rate-­determining step changes to loss of the PhSCH2CH2SPh by rupture of the second Pd–S bond.111 A proton may assist the removal of a chelating ligand in a manner similar to that known for octahedral complexes (see Chapter 5), but second-­order reaction paths are also possible (Scheme 3.11; X and Y may be halide ligands or even H2O). Reactions of some Pd(ii) complexes follow this scheme.112 If the chelated ligand has a free basic site, its protonation, prior to ring opening, will assist the ring-­opening step, as seen in the reactions of [M(ox)2]2− (M = Pt113 or Pd114), [Pd(biguanide)2]2+,115 [PdCl2(biguanide)],116 [Ni(biguanide)2]2+ and [M(ethylenedibiguanide)]2+ (M = Ni, Cu).117 Electrophiles such as metal ions can have a similar role. In some such cases the ring opening may be fast followed by a rate-­determining loss of the chelating ligand (see Chapter 5). Teggins and Milburn118 reported results of studies on the exchange of [Pt(ox)2]2− with 14C-­labelled oxalate in solution. The rate of exchange is

Chapter 3

148

Scheme 3.11

Scheme 3.12 independent of pH in the range 3.6–6.8 and occurs slowly according to the following two-­term rate law:   



Rate = k0[Pt(ox)2] + kox[Pt(ox)2−2][ox2−]

(3.41)

  

At 25 °C (ionic strength 0.17 M) in oxalate buffer (pH 4.7), k0 = 1.6 (± 0.2) × 10−9 s−1 and kox = 4.6 (± 0.3) × 10−7 M−1 s−1. The proposed mechanism is shown in Scheme 3.12. The observed kinetic behaviour can be explained if the intermediate I in Scheme 3.12 is formed only in very low concentration in the pre-­equilibrium

Ligand Replacement Reactions of Metal Complexes

149

reaction 1 (k″1 > k′1) and I reacts through the rate-­determining reactions 2 and 3 to form very low concentrations of the intermediates II and III. Reactions 2 and 4 together contribute to the kox path. Teggins and Milburn119 also reported on the exchange of O of [Pt(ox)2]2− with that of the solvent (studied in H2O enriched with 18OH2); all eight oxygens exchange at the same rate, which is much faster than the exchange of [Pt(ox)2]2− with free ox2− in solution (mentioned above). At pH 3–8 they observed the results in conformity with the following rate law:   



Rate = 8(ka[H+] + kb[OH−])

(3.42)

  

At pH 3.74 and 45 °C (ionic strength 0.035 M), kex = 1.7 × 10−9 s−1; at 45 °C (ionic strength 0.05 M), ka = 1.4 × 10−3 M−1 s−1 and kb = 1.0 M−1 s−1. The relative values of the rate constants are such that below pH 5.5 the ka path predominates. Scheme 3.13 has been suggested for the exchange. This implies that exchange takes place on a free carboxylate resulting from one-­ended dissociation of the protonated oxalate ligand. Since all the four Pt–O bonds have equal chances to undergo dissociation, all eight oxygens are kinetically equivalent, hence all these oxygen atoms exchange at the same rate. In agreement with this is the fact that oxygen exchange in free Hox− occurs at a comparable rate.120 Similar observations have been reported121 for oxygen exchange in [Cr(ox)3]3− (see Chapter 6). Formation of [PtX4]2− in the reaction of [Pt(ox)2]2− with X− (X = Cl, Br) in acidic solution has been reported.113 The following rate law explains the experimental observations:   

  

Rate = (k1 + k2[H+])[X−][Pt(ox)22−]

(3.43)

A mechanistic scheme for the overall reaction has been proposed that agrees with the rate law. Essentially this involves association, forming

Scheme 3.13

Chapter 3

150 −

X ⋯Pt(ox)2 , which suffers rate-­determining one-­ended dissociation of a chelated ox2− by concurrent [H+]-­independent and [H+]-­dependent paths (the latter involving the protonated carbonyl group of the chelated ox2−) with simultaneous entry of X− into the vacated position; the products [Pt(ox)(X) (–ox)]3− and [Pt(ox)(X)(–oxH)]2− so formed suffer fast loss of the unidentate ox2− and oxH−, respectively, with simultaneous entry of an X− into the vacated position forming [Pt(ox)X2]2−, which, owing to the trans effect of the bound halide (X−) ions, suffers fast dissociation with simultaneous entry of X− into the two vacated positions, forming [PtX4]2−. Values of k1 and k2 at several temperatures and the corresponding ∆H‡ and ∆S‡ values have been reported.113 It has already been mentioned (see earlier in Section 3.1.4) that steric hindrance caused by bulky ligands that spill over and block the apical positions of a square-­planar complex will hinder the approach of the incoming nucleophile, causing significant retardation of the rate of ligand replacement in such complexes. This has been illustrated taking the reactions of cis-­ and trans-­[PtCl(L) (PEt3)2] (L = phenyl, mesityl) as examples. Such a steric effect of a bulky ligand can therefore affect the contribution of the kY path to the overall process, and in some cases the kY path may be absent and the reaction rate is independent of the concentration of Y, as in the case of [MX(Et4dien)]n+ (M = PdII, PtII, AuIII), which is typical behaviour of octahedral complexes of Co(iii) (see Chapter 4), hence such square-­planar complexes are referred as pseudo-­octahedral.122 For [PdX(Et4dien)]+ (X = Cl or Br), the rate of replacement of X− by various nucleophiles Y is independent of [Y] (hence kY ≈ 0),122a,b and the same is true for the analogous Pt(ii) complex.122c [AuCl(Et4dien)]2+ easily forms the conjugate base [AuCl(Et4dien-­H)]+ from which Cl− can be displaced by various nucleophiles almost exclusively by the k0 path, and this is also true for [AuCl(Et4dien)]2+; however, at 25 °C the former reacts ca. 70 times faster than the latter.122d In the Cl− replacement reactions of trans-­[PtCl(Me)(PEt3)2] (no steric blocking) with the nucleophiles N3−, Br− and I− in methanol, both k0 and kY paths have been observed, with kY increasing in the order N3− < Br− < I−, which is the order of the nucleophilicity of these ligands, but in the case of the weak nucleophile NO3− the reaction proceeds exclusively by the k0 path.123 However, for comparable reactions of the sterically hindered trans-­[PtCl(mesityl) (PEt3)2] in methanol, only very strong nucleophiles such as CN−, SeCN− and SC(NH2)2 exhibit the two-­term rate law, with kY increasing in the sequence CN− < SeCN− < SC(NH2)2, but NO3−, N3−, Br− and I− react exclusively by the k0 path.124 Since the associative path is suppressed by steric hindrance in the case of trans-­[PtCl(mesityl)(PEt3)2], it is most likely that the k0 path involving replacement of Cl− with methanol (solvent) proceeds by an Id mechanism, and this is followed by fast replacement of MeOH with the nucleophile by an interchange process, forming the product. In support of this is the fact that the k0 value is very much reduced if the reaction medium is changed from a protic solvent to a dipolar aprotic solvent.125 A steric effect is also seen in cases of bulky entering nucleophiles. Thus, in the reaction shown in eqn (3.44) in methanol, using radioactivity monitoring 2−

Ligand Replacement Reactions of Metal Complexes

151

it was observed that both k0 and kY of the two-­term rate law were significantly reduced in the case of a hindered amine, kY decreasing in the sequence primary amine > secondary amine > tertiary amine (in this case kY = 0).126   

trans-­[PtCl2{NH(14C2H5)2}(PnPr3)] + amine ⇌ trans-­[PtCl2(amine)(PnPr3)] + NH(14C2H5)2 (3.44)

  

Anation of [Pt(Me5dien)(OH2)]2+ [Me5dien = Me2NC2H4N(Me)C2H4NMe2] by Cl , Br−, I−, SCN− and S2O32−, and also reaction with SC(NH2)2, have been studied in aqueous solution and it was observed that despite steric hindrance, the reactions [replacement of the H2O ligand (a weak nucleophile) with the nucleophiles] proceed via the kY path only (k0 = 0).127 The observed kY values are in the order of increasing nucleophilicity of the nucleophiles: Y = Cl− < Br− < I− < SCN− < S2O32−; the corresponding Pd(ii) complex also behaves similarly. All the results mentioned above suggest that among the factors that determine pseudo-­octahedral behaviour are the bulkiness of the ligands in the Pt(ii) complex, the nucleophilicity of the incoming ligand and of the ligand being replaced and also the nature of the solvent. Thus, for trans-­ [PtCl(mesityl)(PEt3)2], Cl− exchange and Cl− replacement with N3−, SCN− and PhS− proceed by the k0 path only (kY = 0), whereas stronger nucleophiles such as CN− and SeCN− react by the k0 + kY paths in methanol solution. However, on changing the solvent to DMSO, not only strong nucleophiles such as thiourea and SeCN−, but even weaker nucleophiles such as SCN− and I− (which react only by the k0 path in methanol), react by the k0 + kY paths128 An important point to note in the role of the solvent is the absence of a large effect on the rate of the ligand replacement reaction and the discriminating ability of the substrate in Pt(ii) complexes, as reported by Belluco and co-­workers.129,130 For reaction between a particular complex of Pt(ii) and a particular nucleophile in different solvents, the contribution of k0 (which is kS) is generally in the following order of solvents,129 which is also the order of their nucleophilicity towards Pt(ii): −

DMSO > MeNO2 > H2O > ROH The following order was reported from k0 values in the replacement of Cl− in trans-­[PtCl2(py)2]:86 DMSO > MeNO2 > EtOH > nPrOH However, with trans-­[PtCl2(pip)2] the following order was reported (Table 3.9): DMSO >> Me2CO > DMF >> MeCN >> MeNO2 > MeOH In the case of trans-­[PtCl(NO2)(pip)2], the order is DMF > DMSO > MeOH The reported results further show that the order is also substrate dependent to some extent. Thus, the value of the nucleophilic discrimination

Chapter 3

152

parameters [see eqn (2.74), Chapter 2) changes with the nature of the solvent in reactions of trans-­[PtCl2(pip)2] in the following order (values in s in parentheses):86   



DMSO (0.61) < MeNO2 (0.64) < MeCN (0.74) >> DMF (0.74) < Me2CO (0.78) < MeOH (0.91)

  

but kY for displacement of Cl− by Br− changes in the following order (kY values in M−1 s−1 in parentheses):86 DMSO (5.2 × 10−3) < MeOH (6.16 × 10−3) < DMF (6.7 × 10−3) < MeCN (1.28 × 10−2) < MeNO2 (1.5 × 10−2) < Me2CO (8 × 10−2) Isomerization reactions of square-­planar complexes are discussed in Chapter 6. In addition to nucleophilic substitution, ligand replacement in square-­ planar complexes of Pt(ii) may also take place by electrophilic substitution, or oxidative addition followed by reductive elimination. All three types are illustrated in Scheme 3.14 using the example of [PtCl(Me)(PMe2Ph)2].131 The reaction with HgCl2 can be thought of as an electrophilic attack by Hg(ii) on the Pt–C bond. The oxidative addition reaction shows oxidation of Pt(ii) to Pt(iv) with the simultaneous expansion of coordination number from 4 to 6 for Pt. This is followed by elimination of MeCl with the Pt returning to its +2 oxidation state and coordination number 4 with a net substitution of Cl− for Me− (methanide). Bridge-­splitting reactions of the type shown in eqn (3.45) have been used to prepare some complexes of Pt(ii).132 Studies on the kinetics of such reactions have been reported, and some of the data are given in Table 3.15.133a   



(3.45)

  

It is significant that the rates of reactions of bridged complexes are about two to three orders of magnitude faster than those for the corresponding

Scheme 3.14

Ligand Replacement Reactions of Metal Complexes

153

Table 3.15  Rates  of reactions of some platinum(ii) complexes with py at 25 °C. Complex

kS/s−1

Kpy/M−1 s−1

[Pt2(µ-­Br)2Br4]2− [Pt2(µ-­I)2I4]2− [Pt2(µ-­Cl)2Cl2(pip)2] [PtBr4]2− [PtCl3(pip)]−

8 × 10−4 1 × 10−4 7 × 10−4 2 × 10−6 2 × 10−6

8 × 10−2 2 × 10−2 2 × 10−2 3 × 10−5 2 × 10−5

Figure 3.19  Proposed  mechanism for the symmetrical cleavage of halogen-­bridged Pt(ii) complexes.133a

monomeric complexes. This supports the view that there is considerable strain in the four-­membered ring system, resulting in a weaker Pt–X bond for the bridging groups than for the terminal groups. However, a chelate bridge requires the aid of a nucleophile to open, and the usual two-­term rate law is observed for such reactions. The plausible mechanism proposed is shown in Figure 3.19. This is essentially the same in its rate-­determining step as that for substitution in mononuclear square-­planar complexes. Thus, one of the bridging ligands is displaced by either the solvent or the reagent to give as the initial product a mono-­bridged complex. The latter is not expected to be very stable, because compounds of this type

Chapter 3

154

are generally not known, and the final reaction is presumably fast. Studies on bridge-­opening reactions of [Pt2(µ-­Br)2Br4]2− with alkenes and en in acetone solution have been reported.133b Kinetic studies on some dimerization reactions have been reported. An example is the transformation shown in eqn (3.46a) and (3.46b), studied by UV spectrophotometry in aqueous solution (pH 4–5).134 The product is formed mainly from cis-­[Pt(NH3)2(OH)(OH2)]+ with a small contribution of cis-­[Pt(NH3)2(OH2)2]2+.   



cis-[Pt(NH3)2(OH2)2]2+ ⇌ cis-[Pt(NH3)2(OH)(OH2)]+ + H+

(3.46a)

  

2cis-[Pt(NH3)2(OH)(OH2)]+ ⇌ [(H3N)2Pt(µ-OH)2Pt(NH3)2]2+ + 2H2O (3.46b)

3.1.5  Energy Profile for Reactions of Square-­planar Complexes The tendency for an enhanced rate of substitution generally observed in square-­ planar complexes of metal ions other than Pt(ii) is presumably due to the greater ability of these metal ions to form five-­coordinate species. In their cases, bond making by the entering nucleophile is more predominant than bond breaking by the departing ligand, a characteristic seen in Ni(ii) where reaction intermediates have been detected and the k0 term is unimportant and reaching a limit in complexes of Rh(i). In reactions of Rh(i) of the type shown in eqn (3.47) the rate is independent of concentration of amine but is dependent on its nature.135   



[RhX(diolefin)(SbR3)] + amine → [RhX(diolefin)(amine)] + SbR3 (3.47)

  

In this case, bimolecular attack of amine leads to rapid formation of a five-­coordinate complex (Figure 3.20c), which slowly eliminates SbR3. The behaviour of Rh(i) differs from those of Pt(ii) and Au(iii) complexes.

Figure 3.20  Reaction  profiles for different plausible situations in the mechanism of ligand replacement reactions of square-­planar complexes (see text).

Ligand Replacement Reactions of Metal Complexes

155

The energy profile for various stoichiometric mechanisms for ligand replacement in square-­planar complexes is shown in Figure 3.20. The case of Rh(i) complexes is represented in Figure 3.20c. Here the energy of the five-­ coordinate intermediate will be lower than that of the reactants but higher than that of the products. The case of Ni(ii) will be similar. In these cases, the major transition state (second peak) will follow the minor transition state. In the situation shown in Figure 3.20a the M–X bond breaking is rate determining, but in Figure 3.20b M–Y bond making is rate determining. In the cases in Figure 3.20a and b there is the formation of an unstable intermediate in which both X and Y are bonded to the metal.

3.2  trans Effect in Platinum(iv) Complexes Unlike platinum(ii), platinum(iv) forms octahedral complexes. From the results of investigations reported by Zvyagintsev and Karandasheva136 on the reactions Pt(iv) complexes, the order of the trans effect in Pt(iv) complexes is I− > Br− > Cl− > OH− > NH3 ≈ NO2−. The unusual position of NO2− (its exact position is uncertain, but it is definitely weaker than Cl−) is significant; and so also that of OH−. However, many reactions of Pt(iv) complexes are complicated owing to their photosensitivity and catalysis by Pt(ii)46 present in trace amounts, the nature of which has already been discussed (see Section 3.1.2). Direct substitutions are, however, not unknown, as in the base hydrolysis of [PtCl(NH3)5]3+, cis-­[PtCl2(NH3)4]2+ and trans-­[PtCl2(NH3)4]2+, leading to the formation of hydroxo complexes in each case.137 The reactions of the first two are not catalysed by Pt(ii), and the third one reacts directly with OH− slowly and requires Pt(ii) catalysis for rapid reaction. For such substitutions, the order of the trans effect is NH3 > Cl−. All these observations indicate that trans activation of the type known in Pt(ii) is of no significance in Pt(iv). Because of its higher positive charge, Pt(iv) is definitely less polarizable and a much weaker π-­donor than Pt(ii), and this explains the observations. X-­ray data have shown that, unlike in the case of Pt(ii) complexes, the Pt–X bond length does not change noticeably with the nature of L in the system (trans) L–PtIV–X. Replacement of X− (X = Cl or Br) with I− in trans-­[PtX2(ox)2]2− proceeds by reduction to [Pt(ox)2]2− by I− followed by oxidative addition of I2 to form trans-­[PtI2(ox)2]2−.138a However, in methanol exchange of [Co(MeOH)5(py)]2+ in methanol solution, a kinetic trans effect has been reported.138b At −30 °C the rate constant for exchange of MeOH trans to py is 1.2 × 103 s−1, which is ca. 3.5 times faster than for exchange of MeOH in the cis sites (k = 4.1 × 102 s−1).

3.3  Other Square-­planar Metal Complexes Although Pt(ii) complexes are the best known and most extensively investigated square-­planar complexes, other d8 systems also form such stable planar complexes. Many such complexes of Ni(ii), Pd(ii), Au(iii), Rh(i) and Ir(i),

Chapter 3

156

which are thermodynamically stable, are generally fairly labile and hence their structural isomers are less common than for Pt(ii). Hence stereospecific reactions, if they occur, may go unnoticed because an unstable isomeric product formed in a favoured kinetic step may readily rearrange to yield a more stable or less soluble isomer. Thus, for the reaction of [Pd(NO2)4]2− with NH3, since NO2− has a higher trans effect than NH3, the reaction product is expected to be the cis-­diammine, but instead the trans isomer is isolated and appreciable amounts of the cis isomer are produced under conditions of lower temperature (ca. 10 °C).139 It has been suggested that the initial product is the cis isomer, which rearranges to give the less soluble trans isomer. Similarly, reaction between [PdCl4]2− and NH3 in solution leads to trans-­[PdCl2(NH3)2]. The cis isomer has, however, been made by strongly acidifying a solution of [Pd(NH3)4]2+ with HClO4, generating cis-­[Pd(NH3)2(OH2)2]2+, followed by addition of a saturated solution of NaCl to precipitate cis-­[PdCl2(NH3)2]. The products of these reactions are just the opposite of what are expected on the basis of the trans effect as known in the reactions of the corresponding complexes of Pt(ii) (see Section 3.1.1). Hence the trans effect is of no significance in complexes of Pd(ii), for which an explanation was offered by Chatt et al.25 (see Section 3.1.3).

3.3.1  Palladium(ii) Complexes Kinetic data seem to indicate that the trans effect of the type so well known in Pt(ii) complexes is not of much significance in Pd(ii) complexes. Thus, the rate of hydrolysis of cis-­[PdX2(NH3)2] (X = Cl or Br) is reported to be greater than that of the corresponding trans isomer. More comprehensive kinetic studies on substitution reactions in Pd(ii) complexes were first reported by Banerjea and Tripathi.139 Their results indicated that the rates vary from quite fast to slow, comparable to what is known in Pt(ii) series, but the factors that influence the rates are different in the two series. Thus, a comparison of the rates of replacement of a Cl− by H2O in the following Pt(ii) and Pd(ii) complexes is of significance: trans-­[PtCl2(NH3)2]: cis-­[PtCl2(NH3)2]: [PtCl2(en)]: trans-­[PdCl2(NH3)2]: [PdCl2(en)]:

k1 = 1.0 × 10−4 s−1 at 25 °C k1 = 3.8 × 10−5 s−1 at 25 °C k1 = 5.3 × 10−5 s−1 at 25 °C k1 = 3.8 × 10−4 s−1 at 36 °C k1 = very fast at 36 °C

Thus, whereas in the Pt(ii) series the trans isomer reacts about two to three times faster than the cis isomer, in the Pd(ii) series the cis isomer reacts very much faster than the trans isomer, despite the fact that the rate of reaction of the latter is comparable to that of the trans-­Pt(ii) compound. Again, trans-­ [Pd(NO2)2(NH3)2] reacts with aniline at a measurable rate in contrast to that of the dichloro complex, which reacts immeasurably fast,140 although an opposite behaviour is expected because NO2− has a greater trans effect than Cl−. All these results suggest that the trans effect of the type known in Pt(ii) is not observed in Pd(ii), and whereas the trans effect is of no significance in

Ligand Replacement Reactions of Metal Complexes

157

Pd(ii) complexes, the cis effect is dominant. In view of this, the report based on studies on the replacement of Cl− ligands in [PdCl4]2− with NH3 that both the trans effect and cis effect are larger for Pd(ii) than in Pt(ii) complexes141 needs to be looked into critically with further investigations. Interestingly, for any particular complex, the rate is dependent on the nucleophilic character of the incoming group. Hence OH−, which is not at all a good reagent for Pt(ii), is extremely powerful for Pd(ii), and all the complexes examined reacted immeasurably fast with OH−. Also, all the reacting ligands studied react by an associative process, with almost all the complexes, and the rates were found to decrease more or less in the following sequence of the entering ligands:140 OH− >> SC(NH2)2 > py > glycine > PhNH2 > H2O It may be argued that the difference in behaviour of the cis and trans isomers in Pt(ii) and Pd(ii) complexes occurs because in the Pt(ii) complexes the trans isomer is thermodynamically less stable than the cis isomer, whereas the reverse is true for Pd(ii) complexes. Experimental data are available, however, that indicate that this is not so and that for both Pd(ii) and Pt(ii) the trans isomer is thermodynamically more stable than the cis isomer [see also Section 3.1.4 for Pt(ii)]. Experimental data on the relative stabilities of cis and trans isomers of [MX2A2] [M = Pd(ii), Pt(ii); X = Cl−, NO2−] were reported by Chakravarty and Banerjea.142 From observations on the half-­wave potentials (E½) for the polarographic reductions of the complexes under comparable conditions, it was concluded that for both Pt(ii) and Pd(ii) trans isomers are more stable than the corresponding cis isomers for the complexes examined (Table 3.16). It should be noted that the greater the thermodynamic stability of the complex, the more negative is E½. Since in the case of Pt(ii) the dichloroethylenediamine complex has almost the same E½ as the cis-­dichlorodiammine complex, it is not unreasonable to presume that the same will be true for the corresponding Pd(ii) complexes also. Cattalini and co-­workers143 studied the kinetics of replacement of X− in [PdX(dien)]+ with pyridine, which obeys a second-­order rate law, being first order with respect to each of the reactants. A comparison of their data with the results for the corresponding Pt(ii) complexes69 (Tables 3.17 and 3.18) is interesting. Table 3.16  Polarographic  reduction data for some Pd(ii) and Pt(ii) complexes. Complex

Medium

E½/V vs. SCE

trans-­[PtCl2(NH3)2] cis-­[PtCl2(NH3)2] [PtCl2(en)] trans-­[PdCl2(NH3)2] [PdCl2(en)] trans-­[Pd(NO2)2(NH3)2] cis-­[Pd(NO2)2(NH3)2]

0.1 M KCl 0.1 M KCl 0.1 M KCl 1 M KCl 1 M KCl 0.1 M NaClO4 0.1 M NaClO4

−0.165 −0.10 −0.11 −1.37 −1.25 −0.33 −0.27

Chapter 3

158

Table 3.17  Comparison  of rates of reaction of analogous complexes of Pd(ii) and Pt(ii).

[MX(dien)]+ + py → [M(dien)(py)]2+ + X− Kpy/M−1 s−1 at 25 °C X−

Pd(ii)143

Pt(ii)69

Cl− Br− I− N3− SCN− NO2− CN−

Fast Fast 3.55 × 10−2 1.30 × 10−2 3.10 × 10−1 6.4 —

3.48 × 10−5 2.30 × 10−5 1.00 × 10−5 8.33 × 10−7 3.20 × 10−7 5.00 × 10−8 1.67 × 10−8

Table 3.18  Comparison  of the rate data for analogous complexes of Pd(ii) and Pt(ii): replacement of X− with py1h.

kobs/s−1 at 25 °C Complex

Pd(ii)

Pt(ii)

[MI(dien)]+ [M(SCN)(dien)]+ [M(NO2)(dien)]+ [MCl(Et4dien)]+ trans-­[MCl(o-­tolyl)(PEt3)2]

3.2 × 10−2 4.3 × 10−2 3.3 × 10−2 2.1 × 10−3 5.8 × 10−1

1.0 × 10−5 3.0 × 10−7 2.5 × 10−7 8.5 × 10−6 6.7 × 10−6

The data in Tables 3.17 and 3.18 indicate that, as reported by Banerjea and Tripathi,140 Pd(ii) complexes are much more labile than those of Pt(ii), which is generally found and is the expected trend for 4d8 and 5d8 systems. However, the results reported by Banerjea and Tripathi140 also showed that for aquation trans-­dichlorodiamminepalladium(ii) is an exception. A 17O NMR spectroscopic study144 of water exchange in [Pd(OH2)4]2+ showed that the rate is ca. 106 times faster than the corresponding process in [Pt(OH2)4]2+. A low negative ∆V‡ value (−3.0 ± 0.2 cm3 mol−1) for the reaction in eqn (3.48) suggests an interchange associative (Ia) mechanism145 rather than an associative (A) process.   

  

[Pd(Et4dien)(OH2)]2+ + Cl− → [PdCl(Et4dien)]+ + H2O

(3.48)

Comprehensive studies146 of the aquation and halide anation reactions of palladium halide and aqua halide complexes, including reactions of cis-­ and trans-­[PdX2(OH2)2] (X = Cl, Br), showed that the Pd(ii) complexes react ∼104–105 times faster than the analogous Pt(ii) complexes147 and the activation enthalpies are 8–10 kcal mol−1 smaller for Pd(ii) (Table 3.19). Also, the relative trans effects (Cl/H2O and Br/H2O) are ca. six times smaller for Pd(ii) than for Pt(ii), but the cis effects (H2O/Cl and H2O/Br) are somewhat greater for Pd(ii) than for Pt(ii).

plexes; ionic strength 1.00 and 0.50 M for Pd and Pt, respectively. Values of ΔH‡/kcal mol−1, with standard deviation ±1) and ΔS‡/cal K−1 mol−1, with standard deviation ±3)a .146 M = Pd ΔH

‡

M = Pt ‡

ΔS

−1

Aquation reactions cis-­MCl2 → MCl+ + Cl− trans-­MCl2 → MCl+ + Cl− trans-­MBr2 → MBr+ + Br− MCl3− → trans-­MCl2 + Cl− MBr3− → trans-­MBr2 + Br− MCl3− → cis-­MCl2 + Cl− MBr3− → cis-­MBr2 + Br−

k/s (3.9 ± 1.6) × 10−1 (5.6 ± 0.5) × 10 (2.6 ± 0.2) × 102 (2.7 ± 0.9) × 10−2 (5.9 ± 1.5) × 10−3 (1.4 ± 0.3) × 10 >1 × 102

Anation reactions MCl+ + Cl− → cis-­MCl2 MBr+ + Br− → cis-­MBr2 MCl+ + Cl− → trans-­MCl2 MBr+ + Br− → trans-­MBr2 trans-­MCl2+ Cl− → MCl3− trans-­MBr2+ Br− → MBr3− cis-­MCl2+ Cl− → MCl3− cis-­MBr2+ Br− → MBr3−

k/M−1 s−1 (5.1 ± 1.3) × 102 (4 ± 3) × 103 (3.5 ± 0.9) × 104 (6.9 ± 1.6) × 105 21.7 ± 0.2 (8.3 ± 0.4) × 10 (5.4 ± 0.6) × 103 >2 × 105

ΔH‡

ΔS‡

kPd/kPt

−1

— 12 9 — — — — — — — — 13 11 10 —

— −9 −17 — — — —

k/s ∼3 × 10−7 ∼1 × 10−4 — 2.8 × 10−8 1.4 × 10−8 6 × 10−5 6 × 10−4

— — — 24 26 20 21

— — — −12 −8 −11 −3

∼1 × 106 ∼6 × 105 — 1 × 106 4 × 105 2 × 105 >1 × 105

— — — — −9 −14 −7 —

k/M−1 s−1 ∼2 × 10−8 — ∼5 × 10−1 — 4.6 × 10−5 1.8 × 10−4 7.5 × 10−2 3.3

— — — — 23 20 18 14

— — — — −2 −10 −4 −9

∼3 × 105 — ∼7 × 104 — 5 × 105 5 × 105 7 × 104 >6 × 104

Ligand Replacement Reactions of Metal Complexes

Table 3.19  Rate  constants (25 °C) and activation parameters for aquation and anation reactions of palladium(ii) and platinum(ii) com-

a

Water molecules have been omitted from complex species for clarity.

159

Chapter 3

160

0 Nucleophilic reactivity constants  nPd  were determined for some ligands in substitution reactions of Pd(ii) complexes149 following a procedure similar to that used for evaluating nPt0 values (see Chapter 2), except that trans-­ 0 [Pd(NO2)2(PnPr3)2] was used as the standard substrate. For nPd , the sequence − − – − − is Cl < N3 < Br < I < SCN < (NH2)2CS, i.e. the same as for nPt0 ; the Pd(ii) substrates also obey the linear free-­energy relationship that allows the calculation of nucleophilic discrimination factors (sPd) for Pd(ii) complexes (cf. sPt, see Chapter 2). Rate constants and activation parameters for the reactions

  



[PdX(Et4dien)]+ + Y− → [PdY(Et4dien)]+ + X−

(3.49)

  

where X = Cl, Br or I and Y = Cl, Br or I were determined in a range of solvents in order to gain more detailed knowledge about the reaction mechanism for these pseudo-­octahedral systems.149 Ancillary NMR evidence indicated the following order for the ability of the solvents to interact with Pd(ii): H2O > MeOH > EtOH ≈ DMSO > CH2Cl2. The results of the kinetic studies showed that for a given complex in a given solvent, the kinetic parameters are independent of Y, as in octahedral Co(iii) complexes (see Chapter 4), hence these complexes behave as pseudo-­ octahedral substrates. In protic solvents, the reactivity of [PdX(Et4dien)]+ decreases in the order X = Cl > Br > I. This decrease stems from the variation in ∆H‡ rather than ∆S‡. These observations suggest an associative type of mechanism, with a transition state [PdX(Et4dien)(solvent)]+. Reactivity trends in aprotic solvents (DMSO and DMF), where the fastest substitution occurs with [PdX(Et4dien)]+, are better interpreted in terms of a dissociative mechanism, in which solvation of the leaving halide is a significant factor, but solvation at the Pd(ii) is not.149 In another investigation,150 it was shown that replacement of X− with Br− in [PdX(Et4dien)]+ is independent of X− for X− = SCN−, NCS− or NCSe−, but for X− = SeCN− the rate law is of the usual two-­term type, observed for complexes containing ligands less bulky than Et4dien. Such a two-­term rate law operates for the reactions of [Pd(SeCN)(Et4dien)]+ with I−, N3− or CN−. It is noteworthy that the reactivity order (entering ligand-­dependent path) is N3− > Br− > I−, with the reactivity trend apparently controlled by size rather than by nucleophilicity towards Pd(ii). Such a dependence on size of the incoming group is just what one would expect for associative reactions with a complex containing the bulky and sterically obstructional ligand Et4dien. Although in cases of analogous complexes of Pt(ii) and Pd(ii) the reactions of Pd(ii) are generally much faster, by ca. 104–106 times,144,146 these reactions show similarities in several respects, such as conformity with the two-­term rate law, effect of nucleophilicity of the entering nucleophile, effect of steric hindrance, solvent effect, etc. A ratio of the order of 104 has also been reported for the reactions of [MCl4]2− (M = Pd151 and Pt99) with amidinothiourea. However, the reactions of [PdCl4]2− with amidinothiourea, as also with thiosemicarbazide

Ligand Replacement Reactions of Metal Complexes

161

and biguanide, seem to be of the interchange dissociative (Id) type, but the reaction of [PtCl4]2− with amidinothiourea appears to be Ia.151 Ligand replacement reactions of [PdCl4]2− with several other nucleophiles (NH3, en, bipy, phen and allyl alcohol)88,152 have been reported. Exchange of acac with [Pd(acac)2] follows the two-­term rate law, which is a common feature of square-­planar complexes. A novel feature of this investigation was the use of non-­polar solvents, in several of which the value of k0 is the same. This k0 path therefore is not solvent-­assisted dissociation (kS) as applies for reactions in polar solvents, but is simple dissociation.98 Reaction of [Pd(biguanide)2]2+ in aqueous solution containing HCl and NaCl occurs in two stages:   

  

[Pd  bigH 2 ]2    PdCl2  bigH    [PdCl4 ]2  HCl, NaCl

HCl, NaCl



(3.50)

Both stages of the reaction were followed by spectrophotometry and experimental observations led to the rate law given in eqn (3.51) for the first stage. The rate law for the second stage is identical. Values of the rate constants and corresponding activation parameters have been reported and plausible reaction schemes have been proposed.153,   



−d[Pd(bigH)2]2+/dt = {k1[H+][Cl−] + k2[H+]2[Cl−]2}[Pd(bigH)2]2+ (3.51)

  

Reactions of [Pd(AA)2]2− (AA = ox2−, mal2−) in aqueous solution in the presence of HCl and NaCl forming [PdCl4]2− have been followed by stopped-­flow spectrophotometry; there are two consecutive stages and for both, under pseudo-­first order conditions, kobs = k0 + k2[H+][Cl−]. Values of the rate constants and activation parameters have been reported and a plausible reaction scheme has been proposed.154 However, a trans effect of the type seen in Pt(ii) is almost absent in Pd(ii). Chatt et al.25 explained this on the basis of their π-­bonding trans effect theory. In forming a Pt(ii) to ligand π-­bond, the metal uses its hybrid d–p orbital, which gives a better overlap than a pure d orbital. In the case of Pt(ii), this involves a 5d–6p hybrid, formation of which is energetically not unfavourable because of the closeness in energy of the 5d and 6p in Pt(ii). However, the corresponding hybrid for Pd(ii) would be 4d–5p. However, owing to a fairly large difference in energy of 4d and 5p orbitals, this hybridization is not favoured energetically, hence π-­bonding responsible for the trans effect does not operate for Pd(ii), which, therefore, does not exhibit a trans effect as seen in Pt(ii). Moreover, M → L donation from 5d would be more favoured than from 4d. It should also be mentioned that the reactions of Au(iii), due to a higher charge on the metal ion, are significantly more associative than those of Pt(ii), which is one of the principle reasons for the predominance of the kY path in the reactions of Au(iii) complexes (Table 3.20).

3.3.2  Gold(iii) Complexes A comparison of the behaviours of square-­planar complexes of Au(iii) with those of Pt(ii) is also interesting (Table 3.20).

Chapter 3

162

Table 3.20  Comparison  of rates of reactions of analogous complexes of Au(iii) and Pt(ii) in aqueous solution [20 °C for Au(iii) and 25 °C for Pt(ii)]. kS/s−1

Reaction − 36

−

AuCl4 – Cl exchange PtCl42−–36Cl− exchange [AuCl(dien)]2+ + Br− → [AuBr(dien)]2+ + Cl− [PtCl(dien)]+ + Br− → [PtBr(dien)]+ + Cl−

−3

6.0 × 10 3.8 × 10−5 0.5 8.0 × 10−5

kY/M−1 s−1

Ref.

1.47 −0 154 5.3 × 10−3

155 67 156 and 157 156 and 157

Table 3.21  Relative  rates of reactions of anionic complexes of Au(iii) and Pt(ii). Reaction [n = 1 for Au(iii) and 2 for Pt(ii)]

kAu/kPt at 25 °C

[MCl4]n– + Br−→ [MCl3Br]n– + Cl− [MBr4]n– + Cl−→ [MBr3Cl]n− + Br− [MCl4]n− + H2O →[MCl3(OH2)]n−1– + Cl− [MBr4]n− + H2O → [MBr3(OH2)]n−1− + Br−

1.3 × 106 4.0 × 105 6.0 × 102 1.2 × 103

The results indicate that Au(iii) complexes are more labile than those of Pt(ii); the higher the overall positive charge on the complex, the greater is the lability {[AuCl(dien)]2+ >> AuCl4−}. These results indicate that the reactions of Au(iii) complexes proceed by an associative mechanism, where bond making by the incoming nucleophile is more important than bond breaking by the departing ligand, i.e. the reactions of Au(iii) have significantly more associative character, and it is appropriate to consider these as an associative (A) process, whereas for analogous cases of Pt(ii) and Pd(ii) the processes are Ia and Id, respectively. This is also true for anionic complexes158 under comparable conditions (Table 3.21). An associative character for reactions of complexes of Au(iii) was also inferred from the observed solvent dependence of the rates of these reactions.159 For the following displacement process:   



[AuCl(dien-­H)]+ + Y− → [AuY(dien-­H)]+ + Cl−

(3.52)

  

where dien-­H = conjugate base of (H2NCH2CH2)2NH, the order of reactivity for Y− (kY/M−1 s−1, values at 25 °C given in parentheses)157 is I− (6100) > SCN− (1300) > Br− (190) > N3− (80) >> OH− (very small). This behaviour is comparable to that observed in the Pt(ii) system, with the exception of the order I− > SCN− being the reverse of that generally observed in Pt(ii). On going from Pt(ii) to Au(iii) complexes, increased reactivity is accompanied by an increase in all factors that favour the formation of new bonds, and in addition the nature of the leaving group is also, as a rule, important in determining the kinetic behaviour. Furthermore, the effects due to the nature of the entering and leaving groups are more dependent on one another than in the case of Pt(ii) complexes. This can be seen in the fact that the nature of the leaving group influences the ability of the substrate to discriminate between different nucleophiles.160 The trans-­labilizing effect Br− >

Ligand Replacement Reactions of Metal Complexes −

163 −

161

Cl has been reported in the displacement of the two Br ligands in trans-­ [Au(CN)2Br2]− by Cl−. It is of interest to note how the reactivity of an entering ligand Y in ligand replacement reactions of square-­planar complexes depends on the basicity of the nucleophile (Y) as measured by pKa value of LH+. For a group of nucleophiles, such as Cl−, Br−, I−, OH−, NO2−, N3−, SCN− etc., there is no relationship between their pKa values (of the protonated nucleophile) and their reactivity towards d8 square-­planar complexes. Moreover, the reactivity sequence is often the reverse of the basicity sequence. Thus, OH− and MeO−, despite being strong bases, do not react with any contribution by the kY path with complexes of Pt(ii), and for the halides the reactivity order is always I− > Br− > Cl−. For a class of reagents of the same type, such as heterocyclic amines having no (or comparable) steric hindrance and different basicities, the following general type of free energy relationship has been observed in a number of cases in various solvents:   



log kY = A(pKa) + a

(3.53)

  

where α is a constant. The parameter A is a measure of the ability of the substrate to discriminate between various amines, and is therefore a measure of sensitivity of the system with respect to proton basicity. For complexes of Pt(ii) the A values are fairly small. Thus, for the reaction shown in eqn (3.54) A = 0.06 79 and for that in eqn (3.55) A =0.05 (see ref. 1f, p. 272).   



MeOH [PtCl  bipy   am ]  Cl  PtCl2  bipy    am 

(3.54)



MeOH trans -[PtCl2 (py)2  am  [PtCl(py)2  am ]  Cl 

(3.55)

     

However, in cases of other d8 complexes, A values in the range 0–0.9 have been reported, such as the following values for complexes of Au(iii) in acetone: AuCl4− 0.15, [AuCl2(phen)]+ 0.22, [AuCl2(bipy)]+ 0.46, [AuCl2(5-­NO2phen)]+ 0.89 and [AuCl2(tmen)]+ 0.14. The A values for some complexes of Pd(ii) are [PdCl4]2− 0.24 and [PdX(dien)]+ Me > Ph > 4-­ClC6H4 > 4-­MeOC6H4 > 4-­PhC6H4 > Cl−, with a change in reactivity of four orders of magnitude along the sequence.   



trans-­[PtCl(L)(PEt3)2] + py → trans-­[Pt(L)(PEt3)2(py)]+ + Cl−

(3.56)

  

Studies of the isotopic exchange processes of trans-­[PtCl2(L)(NHEt2)] with [14C]NHEt2 demonstrated the following trans-­labilizing sequence in various solvents:166 H2C=CH2 > iPrCH=CHMe > Me2(OH)CC=CC(OH)Me2 > Et3Sb > Ph3Sb > Me3P > Et3P > nPr3P > Ph3P > Et3As > Ph3As > nPr2S Based on NMR studies of some Pt(ii) complexes,167 it has been shown that H2C=CH2 is a better trans-­labilizing olefin than either cis-­or trans-­2-­butene. There is a similarity between this trans-­labilizing sequence and the order of nucleophilicity towards Pt(ii) substrates; this is not surprising since in the transition state both the entering ligand Y and the trans group L occupy equivalent positions (Figure 3.14). However, one must also consider that, as far as the ligand L is concerned, the effect on the rate is related to the differing ability of the ligand to stabilize the ground state in relation to the transition state, but for Y comparison must be made between the transition state and the free Y in the reaction mixture. The reasons for the observed sequence of trans-­labilizing influence are certainly related to various factors, such as micro-­polarizability and σ-­and π-­bonding. The relative σ and π contributions were estimated in the ground state by a 19F NMR spectroscopic study168 of complexes of the type trans-­[Pt(L)(R)(PEt3)2], where R = 4-­FC6H4 or 3-­FC6H4. A comparison of the role of R in the ground and transition states was also made169 by means of the isotopic effect in the reactions of trans-­[PtCl(H) (PEt3)2] and trans-­[PtCl(D)(PEt3)2].

3.4.1  trans Effect in Terms of Discrimination The concept that the trans effect is the power to labilize a ligand in the trans position is generally accepted and is useful. However, data are available that show that care must be taken when discussing and interpreting experimental observations. Thus, in the following process:   



trans-­[PtCl(T)(PEt3)2] + Y− → trans-­[Pt(T)(Y)(PEt3)2] + Cl−

(3.57)

  

the rate is largely dependent on the nature of T; hence the usual trans-­ labilizing sequence Me− > Ph−> Cl− was observed for the entry of a poor nucleophile such as NO3−, N3− and Br−, but the entry of I− and thiourea led to the sequence Me− > Cl− > Ph−. Apparently the order of trans-­labilizing effect is also dependent on the nature of the entering group. Actually, when the nature of T is changed this affects the ability of the substrate to discriminate between various reagents (s parameter, see Chapter 2) and an inversion of reactivity is

Ligand Replacement Reactions of Metal Complexes

165

observed depending on the nature of the entering group. Because of this it was suggested that it might be possible to define the trans effect as the influence of the trans ligand on the ability of the substrate to discriminate between entering nucleophiles.

3.4.2  cis-­Labilizing Effect It is possible to consider the cis effect from two angles, viz. (a) the cis-­labilizing effect and (b) the effect of the nature of the cis group on the discriminating ability of the substrate. The rate is usually not as much affected by cis groups as it is by trans groups. Data are available that indicate that the cis-­labilizing sequence can be the same (qualitatively) as the trans-­labilizing sequence. Thus, in the reactions of cis-­[PtCl(L)(PEt3)2] with py, the rate changes with the nature of L in the sequence Me− > Ph− > Cl−, as reported for the trans effect.165 However, for the aquation of various chloroamine complexes of Pt(ii) the cis ligand has a somewhat greater influence on the reactivity than the trans ligand.165 In a study on the isotopic exchange of X− with the complexes trans-­ [PtX2(am)(PEt3)2] (X = Cl, Br), no systematic difference was observed between the effect of phosphine or an amine in the cis position.170

3.4.3  cis Effect in Terms of Discrimination The role of the cis ligand cannot be discussed in terms of the reaction rate alone, since inversion of reactivity has often been observed on changing the nature of the entering nucleophile. Comparison of discrimination parameters such as s [see Chapter 2, eqn (2.74)] and A (eqn (3.53)) seems to be the best way to discuss the cis effect of various ligands. In reactions of the following type carried out in methanol at 30 °C:171   



trans-­[PtCl2L2] + Y− → trans-­[PtCl(Y)L2] + Cl−

(3.58)

  

the nucleophilic discrimination parameter s changes with the nature of L in the following sequence:    (a) PEt3 (1.62) > AsEt3 (1.3) > SeEt2 (1.1) > py (1.0) > piperidine (0.95) but for the following reaction in methanol at 25 °C:   



[PtCl(L)(bipy)] + Y− → [Pt(Y)(L)(bipy)] + Cl−

(3.59)

  

the corresponding sequence of s, depending on the nature of L, is (b)  NCS− (1.3) > N3− (0.95) > NO2− (0.87) > Cl− (0.75)    Both sequences (a) and (b) can be interpreted as being due to one or two cis groups which are able to delocalize negative charge away from the metal either because of π interactions as in the case of PEt3 in sequence (a) or because they are high in the scale of micro-­polarizability as in the case of

Chapter 3

166 −

NCS in sequence (b), and thus enhances the electrophilicity of the reaction centre leading to relatively easier attack by the nucleophile. The actual rates depend on both bond formation and bond rupture, hence inversion of reactivity sequence can be observed in some cases. For the following reaction:   



trans-­[PtCl2(am)2] + Y− → trans-­[PtCl(Y)(am)2] + Cl−

(3.60)

  

the s value increases as the basicity of the amine (am) decreases.172 The reactions of Pd(ii) complexes of the type shown in eqn (3.61) carried out in 1,2-­dimethoxyethane at 25 °C show that the cis effect of X− on the discrimination is I− > Br− > Cl−, i.e. the same as the order of micro-­ polarizability. The same is true (NCS− > I− > Br− > N3−) for the cis effect of X− in the reactions shown in eqn (3.62), which is believed to be the rate-­determining step in the cis–trans isomerization of these complexes catalysed by am.173   



trans-­[PdX2(4-­Clpy)(am)] + 4-­Clpy

(3.61)

cis-­[PdX2(am)2] + am → [PdX(am)3]X

(3.62)

  

  

For the reactions shown in eqn (3.63) carried out in 1,2-­dimethoxyethane at 25 °C, the order of the cis effect of X− on discrimination is Cl− > Br− > I−. This anomaly was attributed to stereoelectronic interaction between the lone pair of electrons of the sulfur atom, which is not used in bonding, and the cis ligand X−. Another such anomaly was observed (N3− > Cl−) for the entry of thioethers in [PtCl(L)(bipy)] complexes (ref. 1f, p. 287).   



trans-­[PdX2(am)2] + RSR′ → trans-­[PdX2(am)(RSR′)]X + am

(3.63)

  

However, for opening of the chelate ring in complexes of the type [PdX2(PhSCH2CH2SPh)] by amines [eqn (3.64)], where the ligand X plays the role of both cis and trans partners,174 the sequence of the dependence of discrimination on the nature of X− is normal, viz.

SCN− (0.42) > I− (0.36) > N3− (0.3) > Br− (0.28) > Cl− (0.23)

[A values in parentheses, where A is the slope of the linear plot of log k2 versus pKa of amH+, (eqn (3.53)].   



(3.64)

  

In the case of Au(iii), a comparison of the following two reactions is useful:175,176   

Ligand Replacement Reactions of Metal Complexes −



167 −

[AuCl3(am) + Y  → [AuCl3Y ] + am

(3.65)

trans-­[Au(CN)2Cl(am)]Y− → trans-­[Au(CN)2Cl(Y)]− + am

(3.66)

  

  

In the latter reactions [eqn (3.66)], discrimination between entering nucleophiles is decreased to some extent compared with the former [eqn (3.65)], and this can be interpreted as a decrease in the electrophilicity of the metal on changing from the chloro to the cyano complex.

3.5  S  quare-­planar Complexes of Nickel(ii) and Copper(ii) Peloso177 published a review of kinetic studies on complexes of Ni, Pd and Pt in this common oxidation state. Octahedral complexes of nickel(ii) have been investigated extensively, but square-­planar complexes to a much lesser extent. Exchange of *CN− (14C) with [Ni(CN)4]2− is completed within 30 s at room temperature178 and is believed to occur by a bimolecular displacement process {it is known that [Ni(CN)5]3− is formed179 in a solution of [Ni(CN)4]2− having a high concentration of CN−}. In this connection, the data on the rate of exchange of *CN− with [M(CN)4]2− (Table 3.22), studied by 13C NMR spectroscopy in D2O, are worth noting181 (all these exchanges take place by an associative process). The exchange of [Ni(CN)4]2− with Ni(ii) (studied using a radioactive isotope of nickel) has also been investigated.181 Exchange of Ni(ii) is very fast with [Ni(OH2)6]2+, but slow exchange occurs with anionic complex species such as [Ni(C2O4)2]2−; probably the latter complex must dissociate for exchange to occur. From an investigation of the kinetics of the reaction shown in eqn (3.67),11 it is known that the rate increases with increase in concentration of pyridine and follows the typical two-­term rate law for M = Ni(ii), Pd(ii) and Pt(ii), and the rate constant for the solvent path increases in the sequence (kS values at 25 °C are given in parentheses): Pt(ii) (6.7 × 10−6 s−1) < Pd(ii) (0.58 s−1) < Ni(ii) (33 s−1).   

trans-­[PtCl(o-­tolyl)(PEt3)2 + py → trans-­[Pt(o-­tolyl)(PEt3)2(py)]+ + Cl−

(3.67)

  

Table 3.22  Ligand  exchange rate constants for some tetracyano complexes of d8 metal ions.

Complex

kex/M−1 s−1 at 24 °C

ΔH‡/kJ mol−1

ΔS‡/J K−1 mol−1

[Ni(CN)4]2− [Pd(CN)4]2− [Pt(CN)4]2− [Au(CN)4]−

5 × 105 1.2 × 102 26 3.9 × 103

— 17 ± 2 26 ± 3 28 ± 1

— −178 ± 7 −148 ± 8 −199 ± 3

Chapter 3

168

Thus, the rate decreases with increase in stability of the complex. Also, as expected for a bimolecular displacement process, a marked steric retardation (ca. 1000 times at 25 °C) has been found on changing L from o-­tolyl to mesityl in trans-­[NiL(Cl)(PEt3)2]. A systematic study of exchange of 14C-­labelled α-­amineoximes with their bis-­complexes of Ni(ii) have been reported, and also the reactions of these complexes with EDTA.182 Here also, steric crowding retards the rate significantly. For dissociation of C-­substituted bis(ethylenediamine) complexes of nickel(ii) in acidic solution, the rate decreases with increase in steric hindrance,183 as expected for an associative process. Ligand replacement reactions of square-­planar Ni(ii) dithiolate complexes indicate that the formation of a five-­coordinate intermediate is rapid, followed by subsequent slow steps.184 The reaction systems studied were bis-­complexes of Ni(ii) with several dithiolate ligands and their stepwise transformations (Scheme 3.15; charges have been omitted for clarity). These reactions occur through reversible formation of five-­coordinate intermediates [Ni(S–S)2(L–L)] and [Ni(S–S)2X], respectively, and in a few cases the rate of formation of [Ni(S–S)2X] could be followed. For the reactions with a bidentate entering ligand (a dithiolate or en) Scheme 3.16 has been proposed, which accounts for the observed facts. The five-­coordinate intermediates (for which there is evidence in a few cases) have been shown as square pyramidal but could also be trigonal bipyramidal.

Scheme 3.15

Scheme 3.16

Ligand Replacement Reactions of Metal Complexes

169

Based on this scheme and applying the steady-­state principle to evaluate the concentrations of the intermediates C and D, the following expression for the pseudo-­first-­order rate constant could be deduced:   



kobs = k′K[L–L]/(1 + K[L–L])

(3.68)

  

where k′ = k1k3k5/[k2(k4 + k5) + k3k5]. If K[L–L] is negligible compared with 1, then eqn (3.68) is reduced to   



kobs = k′K[L–L]

(3.69)

  

All the reactions studied, except one, were found to be first order in [L–L] and hence in conformity with eqn (3.69), and for the exceptional case eqn (3.68) was valid. For reaction with CN− Scheme 3.17 has been proposed. In Scheme 3.17, K1 and K2 are equilibrium constants for rapid equilibria and concentrations of (C + D) in the steady state, and assuming that k2 >> k3K2[CN] and with large [CN] (for pseudo-­first-­order conditions), kobs is given by   



kobs = k″K1[CN]2/(1 + K1[CN])

(3.70)

  

where k″ = k1k5K2/k2. A reasonable value of K2[CN] for the system at the concentrations of CN used has been shown to be less than 1. For this reason, and because k2 >> k5 (expected), the assumption k2 >> k5K2[CN] is reasonable. In some cases it has been possible to determine the rate of the reversible formation of the five-­coordinate species B in Scheme 3.17; in these cases, the pseudo-­first-­order rate constant for formation of B is given by   



kobs = kf[CN−] + kb

  

(3.71)

Hence the linear plot of kobs versus [CN−] permitted the evaluation of kf and kb, which are the forward and reverse rate constants, respectively, for the formation of B. It was found in this way that kb = 0 in some systems.

Scheme 3.17

Chapter 3

170 184b,185

Similar studies on [M(S–S)2] (M = Pd and Pt) complexes showed similar mechanisms and the relative reactivity (reaction rate) of Ni(ii) : Pd(ii) : Pt(ii) was found to be 105 : 102 : 1, which is the usual reactivity trend, but the relative rates were not; the relative rates are usually 106 : 105 : 1 for Ni(ii) : Pd(ii) : Pt(ii).186 As mentioned before, for the reaction of [MCl4]2− (M = Pd, Pt) with amidinothiourea the ratio kPd/kPt is of the order of 104.151 Based on available literature data, it was concluded185 that depending on the reaction system, the values of the kPd/kPt ratio show a wide variation, ranging between 102 and 105. Thus, for Cl− replacement with NH3 in Pd and Pt complexes trans-­ [MCl2(NH3)2] and [MCl(NH3)3]+, the kPd/kPt values are of the order of 105. Redistribution reactions of the type shown in eqn (3.72) have been studied in C2Cl4 solution. The kinetic pattern corresponds to a reversible second-­ order process, with significant negative ΔS‡ values; the rate law and activation parameters are consistent with an associative mechanism.187a   



[Ni(L1L222] + [Ni(L1L322] → 2[Ni(L1L2L1L3]

(3.72)

  

Bidentate ligands L–L (L–L = S–S and Se–Se) have been used in kinetic studies of the formation of mixed-­ligand complexes in reversible reactions of the type shown in eqn (3.73) [M = Ni(ii), Pd(ii) and Pt(ii)].187b The reactions were followed by observing the change in the EPR spectrum of the Cu(ii) species. The exchange is inhibited by the free ligand and accelerated by Cuaq2+ ions, which is just the opposite of what has been observed189a in ligand exchange between [Cu(edta)]2− and [Ni(trien)]2+ for which a chain mechanism was proposed. A chain mechanism also operates for the reaction of eqn (3.73) but the chain carriers are the mono complexes [Cu(S–S)] and [M(S′–S′)], Similar studies on the [Cu(mnt)2]2−–[Ni(Et2dsc)2] system have been reported (mnt = maleonitriledithiolate(2−); Et2dsc = diethyldithiocarbamate(1−)).189b   



[Cu(S–S)2]2− + [M(S′–S′)2]2− ⇌ [Cu(S–S)(S′–S′)]2− + [M(S′–S′)(S–S)]2−

(3.73)

  

The planar complexes [NiX2L2] (X = Cl, Br, I, NCS; L = 2-­ and 6-­methyl-­ substituted pyridine derivatives) are uniquely inert to substitution and are not even attacked by strong mineral acids.189 This lack of reactivity is presumably due to steric hindrance. In studies on the replacement of various coordinated bidentate ligands containing a sulfur donor atom (ethyl xanthate, dialkyldithiophosphates, diphenyldithiocarbamate, monothioacetylacetonate) with dithiocarbamate in acetone solution in square-­planar complexes of Ni(ii), Pd(ii) and Pt(ii), several different stoichiometries and types of kinetic behaviour were observed.190 In several cases the substrate [M(S–S)2] reacted with a nucleophile, (S′–S′), to give the product [M(S′–S′)2] in a second-­order process, but no intermediates were detected. In other cases, a mixed-­ligand intermediate [M(S–S) (S′–S′)] was observed and the total reaction forming [M(S′–S′)2] was followed in two stages. In a few cases, an intermediate containing three ligand molecules was detected, which could be square-­pyramidal [M(S–S)2(S′–S′)] having one of the ligands, either S–S or S′–S′, bonded axially unidentately, or even

Ligand Replacement Reactions of Metal Complexes

171

a square-­planar [M(S–S)2(S′–S′)] having one S–S and the S′–S′ both bonded unidentately, but no definite conclusion could be drawn. The representative results given in Table 3.23 were obtained in the reactions of bis(dithiophosphate) complexes with N,N-­diethyldithiocarbamate. The trend in the activation parameters is consistent with a greater degree of bond formation for Ni(ii) and a more significant contribution of bond breaking for Pt(ii); this is rather unusual and contrary to the evidence from observations on many other types of complexes of these metals. Ray191 reported the synthesis of bis(biguanide) complexes of Ni, Cu and Pd; structure determinations by X-­ray crystallography established a square-­ planar structure for [M(bigH)2]2+ (M = Ni, Cu).192a,b Stepwise reactions (shown in Scheme 3.18) of [Cu(bigH)2]2+ with bipy and phen in buffered (pH 8.0) aqueous solution were investigated193 and the rate constants for reactions I–VI were evaluated. The use of alkaline conditions ensured that there were no complications due to acid assisted dissociation of [Cu(bigH)2]2+. In each case, the pseudo-­first-­order rate constant was found to be a composite one (kobs= k0 + kL[L]; L = bipy, phen), which indicates a two-­term rate law as is generally observed for ligand replacement reactions of square-­planar complexes. These k0 and kL values at 25 °C (I = 0.1 M) are given in Table 3.24. Activation parameters ΔH‡ and ΔS‡ corresponding to each of these k values have also been determined; the values corresponding to all six kL conform to a good isokinetic trend.193 The result k0V = k0VI > k0III = k0IV shows that compared Table 3.23  Kinetic  parameters for some analogous complexes of d8 metal ions. Metal ion

Pt(ii)

Relative second-­order rate constants at 25 °C (I = 0.1 M, in acetone solution) ΔH‡/kJ mol−1 ΔS‡/J K−1 mol−1

Pd(ii)

1

1.9 × 10

76 ± 2 ‒12 ± 7

42 ± 1 ‒62 ± 3

Scheme 3.18

Ni(ii) 3

9 × 105 20 ± 1 ‒86 ± 5

Chapter 3

172

Table 3.24  Component  rate constants for steps I‒VI in Scheme 3.18 at 25 °C in borax buffer (pH 8 ± 0.1) and at ionic strength 0.1 M.

Complex

L

Path

k0/s−1

10−3kL/M−1 s−1

[Cu(bigH)2]2+

bipy phen bipy phen bipy phen

I II III IV V VI

6.2 6.2 0.29 0.29 7.25 7.1

1.04 0.73 0.31 1.65 0.14 0.97

[Cu(bigH)(bipy)]2+ [Cu(bigH)(phen)]2+

with bipy, phen causes greater labilization of bigH in [Cu(bigH)(L)]2+. The ΔH‡ value corresponding to any kL is significantly lower than the value for the corresponding k0, whereas the ΔS‡ value for kL is significantly more negative than that corresponding to k0, suggesting much more associative character of the kL path compared with the corresponding k0 path. Dissociation of square-­planar complexes of Ni(ii) and Cu(ii) formed by biguanides, dibiguanides, bis(acetylacetonato)ethylenediamine (baen), cyclam-­t ype macrocyclic ligands, etc., in acidic solution, which are mostly acid-­assisted reactions, are discussed in Chapter 5.

3.6  Reactions of Tetrahedral Complexes Substitutions in tetrahedral complexes have been studied less extensively. A tetrahedral structure is generally found in systems that are not stabilized by crystal field effects (see Chapter 1). Their reactions are therefore generally fairly rapid and hence techniques for fast reactions are required for their investigation. A few typical examples of such reactions are discussed below. One of the earliest studies was on the system shown in eqn (3.74):   



[Ni(CO)2(PR3)2] +PR′3 ⇌ [Ni(CO)2(PR3)(PPR′3)] + PR3

(3.74)

  

for which Rate = k1[Ni(CO)2(PR3)2] and a dissociative mechanism was proposed.194 Studies on fast 14CO exchange and ligand replacement rates of Ni(CO)4 established a dissociative mechanism for these reactions [eqn (3.75)]:195   

(3.75) 

  

For such reactions the rate law is Rate = k1[Ni(CO)4]

Ligand Replacement Reactions of Metal Complexes

173 196

The behaviour of [Co(CO)3(NO)] is similar. A dissociation mechanism is to be expected on different considerations. First, filled d orbitals of Ni(0), a d10 system, should discourage nucleophilic attack. Second, calculations show that an increase in π-­bond character of the metal‒ligand bond occurs on changing from an sp3 tetrahedral to an sp2 triangular planar structure, and this would stabilize the transition state, Ni(CO)3; infrared spectroscopic evidence for this exists.197 For the reaction of Ni(CO)4 with P(OEt)3 in heptane solvent, where electrostriction effects cannot be of significance, the reported ΔV‡ value of 8 ± 1 cm3 mol−1 suggests a dissociative mechanism.198 Exchange of CO in [Co(CO)2(NO)L] (where L = phosphine or phosphite) [eqn (3.76a) and (3.76b)], however, follows a two-­term rate law: Rate = (k1 + k2[CO])[Co(CO)2(NO)L] where k1 and k2 correspond to dissociative and associative paths, respectively.199   



[Co(CO2)(NO)L] ⇌ [Co(CO)(NO)L] + CO

(3.76a)

[Co(CO2)(NO)L] + *CO ⇌ [Co(CO2)*CO(NO)L]

(3.76b)

  

  

Like Ni, Pd and Pt in zero oxidation state also form tetrahedral complexes. Ligand exchange reactions of [M{P(OEt)3}4] with P(OEt)3 in toluene were studied by 1H NMR spectroscopy and a first-­order rate law was observed independent of the free ligand concentration.272a Since the solvent is inert (non-­coordinating) this cannot be a kS path observed in cases of reactions of square-­planar complexes of Pt(ii) (mentioned in a preceding section) and hence indicates a dissociative (D) process. The same kinetic behaviour was reported272a for the ligand replacement reaction of [Ni{P(OEt)3)}4] with cyclohexyl isocyanide in n-­hexane and in benzene solutions, studied by IR spectroscopy. The enthalpy of activation values for these reactions involving M–P bond rupture are in the order Ni > Pd < Pt, which is rather unusual; but a similar unusual trend was reported for ligand replacement reactions of M(CO)6, viz. Cr < Mo > W,272b and for CO replacement by PPh3 in (Cp)M(CO)2, viz. Co < Rh > Ir,272c for which an explanation has been offered (see ref. 200 and citations therein) on the basis of the lower electronegativity of M of the 3d series as one factor and good matching of the orbital energies of M of the 5d series and CO. It is generally true that for analogous compounds a 4d metal is more reactive than its congeners in the 3d and 5d series.272d Reactions of bidentate chelating ligands such as diarsine and phen with [Ni(CO)4], [Co(CO)3(NO)] and [Fe(CO)2(NO)2] have been reported. [Ni(CO)4] reacts in two steps, the first being a rapid dissociation to form a three-­coordinate intermediate, which reacts fast to form a species having unidentately bonded chelating ligand that undergoes ring closure forming the product with loss of a CO. [Fe(CO)2(NO)2] reacts with phen by an associative process only. Reaction of [Co(CO)3(NO)] involves both first-­and second-­order kinetics.

174

Chapter 3

The earliest study on the kinetics of complex formation by [Be(OH2)4]2+ was its anation by SO42− studied by an ultrasound absorption technique.201 Water exchange of [Be(OH2)4]2+ was studied by 17O NMR spectroscopy,202,203 and a ΔV‡ value of −13.6 ± 0.5 cm3 mol−1 was reported,203 and based on this value a limiting A process was proposed, the theoretically estimated value204 for the A process being −12.9 cm3 mol−1. Results of the anation reaction of [Be(OH2)4]2+ with F− (and also the corresponding reaction with HF) have been reported.205 At 25 °C the kf values are 720 M−1H sO−1 (F−) and 73.6 M−1 s−1 (HF). 2 The value for F− is 13.6 times higher than k ex (52.8 M−1 s−1),203 which suggests an associative process. The reaction with HF is ca. 10 times slower than reaction with F−, as is usually observed for reaction of LH+ compared with L (see Chapter 4). A 19F NMR spectroscopic study of F− exchange of BeF42− has been reported.206 The results are in conformity with a rate law having two parallel paths, one dissociative and the other involving an F−-­bridged dinuclear intermediate [Be2F7]3–. Studies on ligand redistribution reactions207 of a series of [Be(β-­diketonate)2] compounds with [B(β-­diketonate)2]+ species have been reported; the β-­diketones used included various alkyl, aryl and fluoroalkyl derivatives.207 In the case of [Be(acac)2], ligand exchange occurs by parallel dissociative (k1) and associative (k2) paths. The relative importance of these two paths depends on the solvent. In CHCl3 or CH2Cl2 the dissociative path is of negligible importance, but in THF the associative path accounts for ca. 75% of the reaction, hence in this case there is a substantial contribution from the dissociative path also.208 In a study of ligand exchange reactions of [Be(TMP)4]2+ with TMP [TMP = trimethyl phosphate, (MeO)3PO] in dichloromethane solvent , the rate is first order in the complex and zero order in TMP; hence a dissociative mechanism was proposed and the following values were reported:209 kex (25 °C) = 3.6 s−1, ΔH‡ = 55.9 kJ mol−1 and ΔS‡ = −43.9 J K−1 mol−1. In a later investigation by another group,203 exchange of [BeL4]2+ with L in an inert non-­coordinating solvent, CD3NO2, was studied for different L [L = dimethyl sulfoxide (DMSO), dimethylformanide (DMF), trimethyl phosphate (TMP), tetramethylurea (TMU) and dimethylpropylurea (DMPU)]. In the case of DMF the observed rate law was kex = k1 + k2[DMF]. With DMSO and TMP, k1 = 0 and exchange proceeds exclusively by the k2 path, which implies an A or Ia mechanism. With TMU and DMPU, owing to steric hindrance, k2 = 0 and hence for these bulky ligands exchange is exclusively by a dissociative (D) process. For the k2 path the ΔV‡ values are −2.5, −4.1 and −3.1 cm3 mol–l for DMSO, TMP and DMF, respectively, which also suggest an A or Ia mechanism in these cases. For the k1 path, the ΔV‡ values are +10.5 and +10.3 cm3 mol−l for TMU and DMPU, respectively, suggesting a D mechanism. The steric course of ligand replacement reactions of tetrahedral complexes can only be followed if optical isomers of asymmetric complexes such as [M(L1)(L2)(L3)(L4)] can be separated. One such study was on [Co(Cp)(CO) (NO)(PPh3)] and another on [Fe(Cp)(CO)(I)(PPh3)], as these are sufficiently inert to permit resolution.210

Ligand Replacement Reactions of Metal Complexes

175

Kinetic studies have also been carried out on the exchange of ligand L with [NiX2L2], where L = PMe2Ph, PPh2(OEt), PPh(OEt)2 or P(OEt)3 and X = Cl, Br or CN. Results with the cyano complexes show the mechanism to be dissociative [eqn (3.77)]:211   



[NiX2L2] ⇌ NiX2L + L

(3.77)

  

A qualitative comparison shows that for L= PMe2Ph, the rate constant for the dissociation step increases in the following sequence of X−: CN− < Br− < Cl−. A study of the reaction shown in eqn (3.78) in different solvents has been reported.212 Both the forward and reverse reactions are associative.   



[CoCl3(S)]− + Cl− ⇌ [CoCl4]2− +S

(3.78)

  

Exchange of [VO(OR)3] with ROH (R = nPr, iPr, nBu and n-­pentyl) was studied by 1H NMR spectroscopy. Highly negative ΔS‡ values (−129 to −150 J K−1 mol−1) and rather low ΔH‡ values (21.4–32.7 kJ mol−1) suggest an associative reaction.213 Reaction of [M(PEt3)4] with C6H4(F)X involves oxidative addition and substitution (ligand replacement), for which the rate varies as follows:214 M = Ni > Pd > Pt and X = I > Br > Cl > CN >> F. Enthalpies of activation for ligand exchange in [M{P(OEt)3}4] follow the trend Ni < Pd > Pt {the same trend as for thermal stabilities of [M(PF3)4]}. This is unlike the trend observed for ligand replacement reactions of square-­planar complexes of these metals in their +2 oxidation state. A delicate balance of σ-­ and π-­bonding can account for this trend.214b A similar trend for the ligand replacement reaction of [M(PF3)4] (M = Ni and Pt) with C6H11NC was reported.214c For exchange reactions of [MBr2(PR3)2] (M = Fe, Co, Ni and R = Ph, o-­tolyl) with PR3, studied in CDCl3 by NMR spectroscopy, the lability trend is Fe > Ni > Co.214d,e The rate law for chloride exchange215 of [CoCl4]2− and [CoCl3(py)]− in the presence of pyridine in nitromethane, determined from 35Cl NMR studies, is Rate = (k1[Cl−] + k2[py])[CoCl3(py)−] which bears a close resemblance to the two-­term rate law for substitution in square-­planar complexes of d8 metal ions. The simplest mechanism consistent with the rate law is parallel associative reactions involving Cl− and pyridine, respectively; slow formation of [CoCl3py]− and its rapid chloride exchange account for the py-­dependent path. Associative reactions should be relatively easy for cobalt(ii) with a low coordination number of four. Formation of [VO2(H2edta)]− from vanadate and H2edta2− in aqueous solution (pH 1.5–2.0) was studied by stopped-­flow spectrophotometry and the results led to a rate law that is consistent with Scheme 3.19.216 All the k values were reported. Similar studies have been reported for complex formation in the reactions of molybdate with catechol217 and catechol derivatives218 (HMoO4− reacts

Chapter 3

176

Scheme 3.19 faster than MoO42−) of vanadate with alizarin,219 and of boric acid with alizarin and quinalizarin in weakly alkaline (pH 7–8) aqueous solution.220 In the latter case, the reaction presumably occurs as shown in eqn (3.79). Values of kf and kb and the corresponding activation parameters (ΔH‡ and ΔS‡) were reported.   



(3.79)

  

Rate constants for the reversible formation of [Cu(AH2)aq]+ from Cuaq+ and AH2 (AH2 = maleic and fumaric acids, i.e. cis-­ and trans-­HOCH2CH=CHCO2H, respectively) were determined by pulse radiolysis.221 In the product complexes, AH2 is bonded to Cu+ as M-­alkene by a µ-­bond. At 25 °C (I = 0.04 M) the following values were reported; the kf values are in the range expected for d10 Cuaq+. −1

−1

kf/M s kb/s−1

Maleic acid

Fumaric acid

9

1.7 ± 0.4 ×109 2.4 ± 0.4 ×105

2 ± 0.4 ×10 1.8 ± 0.4 ×105

In aqueous solution, [Cu(phen)]+ presumably exists as tetrahedral [Cu(phen)(OH2)2]+ species, which will be distorted because the phen, having a rigid structure, has a bite distance that cannot span a pair of tetrahedral sites having an angular separation of 109.5°; the N–Cu–N bond angle in [Cu(phen) (OH2)2]+ will therefore be less than the normal tetrahedral bond angle without any major change in the H2O–Cu–OH2 bond angle. The [Cu(phen)aq]+ reacts with phen to form [Cu(phen)2]+ reversibly, for which the kf value is 6.2 × 104 M−1 s−1 at 25 °C (pH 6.0; I = 0.2 M),222 which is lower than expected by a factor of 105. This is presumably because of the necessity for further distortion of the substrate to permit binding of the second phen that raises the activation enthalpy and hence lowers the rate; the kb value is 0.2 s−1, which is negligible compared with kf. In the case of Ag(i), the following transformations [eqn (3.80) and (3.81)] have k values of ∼109 M−1 s−1 at 25 °C,223a but for complexation of Tlaq+ by I3− the rate constant is ∼104 M−1 s−1 at 25 °C.223b   



[Ag(S2O3)]− + S2O2−3 → [Ag(S2O3)2]3−

(3.80)

[Ag(phen)]+ + phen → [Ag(phen)2]+

(3.81)

  

  

Studies on the kinetics of oxygen exchange of tetraoxometallates [MO4]n− with H2O in aqueous alkaline solution have been reported for M = V(v),224

Ligand Replacement Reactions of Metal Complexes 225

226

227

177 228a,b

Cr(vi), Mo(vi) and W(vi), Fe(vi), Mn(vi) and Mn(vii) and Re(vii).228c 3− − In the case of [VO4] , the rate is not dependent on OH concentration, but with [MO4]2− (M = Mo or W) the reaction occurs in two parallel paths, one dependent on and the other independent of OH− concentration. In the case of [ReO4]−, a dissociative mechanism was proposed since the kex value varies very little with variation of H2O concentration for the reaction carried out in a water–methanol solvent medium. Depolymerization of decavanadate, [V10O28]6−, to mononuclear vanadate in aqueous medium of pH 8–10 is independent of pH, but in strongly alkaline solution [0.05–0.2.0 M, adjusted with MOH (M = Li, Na, K)] a path dependent on OH− concentration also exists (kobs = k1 + k2[OH−]).229a,b The value of k2 depends on the M+ of MOH used, increasing in the order Li+< Na+< K+ (k2 = 0 for Me4N+). A linear dependence of k2 on [Na+] suggests that k2 = k3[Na+]. For k1 the value varies as Li+ > Na+ > K+. The observed effect of M+ is due to ion-­pair formation. For the k2 path ΔS‡ = −80 J K−1 mol−1; this high negative value suggests an associative (A) process. For the base hydrolysis of Cr2O72−, ΔS‡ = −126 J K−1 mol−1, and an associative mechanism was proposed.229c In similar studies230 on paramolybdate, [Mo7O24]6−, kobs = k1 + k2[OH−] but, unlike in the case with decavanadate, the effect of M+ shows the same trend for k1 and k2 (Li+ > Na+ > K+). In contrast to the case with decavanadate, the ΔS‡ value is positive, +21 J K−1 mol−1. Hence, despite similarities in the structural features of these two isopolyanions, the mechanisms of their degradation in alkaline solution must be different. Furthermore, it is worth noting that the reaction of paramolybdate is very fast, whereas that of decavanadate is exceedingly slower. Kinetic studies on the aquation of Cr2O72− to HCrO4− have been reported;231 the reaction is acid catalysed.232 Kinetic studies on the aquation of the peroxo complex[Cr(O2)4]3− have been reported,233 and also on the following transformation in solution:234   



[SiMo12O40]4− → [SiMo11O39]8− + [Mo7O24]6−

(3.82)

  

Mechanisms of the formation of isopoly anions of Mo and W have been reviewed.235

3.7  Complexes of Coordination Number Five The behaviour of such species is of fundamental importance for an understanding of associative (A) substitution mechanisms of four-­coordinate square-­planar complexes as discussed in previous sections and a dissociative (D) mechanism for reactions of six-­coordinate octahedral complexes (see Chapter 4). There is good evidence that all of the reactions of Pd(ii) and Pt(ii) bis-­β-­diketonates with tertiary phosphines proceed via an associative pathway forming a five-­coordinate species, which in some cases cannot be observed at all, others can be detected in solution spectroscopically and in a few cases can even be isolated and examined crystallographically.57b,81

Chapter 3

178

A simple example of ligand replacement in a complex of coordination number five is shown in eqn (3.83):   



Ru(CO)5 + PR3 → [Ru(CO)4(PR3)] + CO

(3.83)

  

For various PR3 the rate of the reaction in cyclohexane is independent of the nature and concentration of PR3, which suggests a dissociative mechanism;236 the following lability orders were reported: Fe(CO)5 < Ru(CO)5 and Mo(CO)6 < Ru(CO)5 < Pd(CO)4. A 17e radical Mn(CO)5 was generated by laser photolysis of a solution of Mn2(CO)10 in cyclohexane. Its reaction with P(tBu)3 was shown to be associative.237 The 17e radicals [Mn(CO)3L2] [L = P(nBu)3 or P(tBu)3] are reasonably stable. Their reaction with CO forming [Mn(CO)4L] in hexane is also associative.238 Molecular structures of [M(L–L)2(PCy3)] and [M(L–L)2{P(o-­tolyl)3}] [M = Pd, Pt; L–L = (CF3COCHCOCF3)−] are square pyramidal with apical (chelate) oxygen. 1H and 13C NMR spectra of these complexes in CDCl3 or CD2Cl2 solutions show intramolecular fluxional behaviour of the Berry pseudo-­rotation type and two such mechanisms (A and B, shown below) have been identified, resulting in site exchange of different groups (Scheme 3.20). The anion [Pt(SnCl3)5]3− has a trigonal bipyramidal structure, shown by X-­ray crystallography of its [Ph3PMe]+ salt; the apical Pt–Sn bonds are shorter than the equatorial bonds.239 The anion is non-­rigid in acetone solution in the temperature range −90 to + 90 °C, as shown by 195Pt and 119Sn NMR spectroscopic investigation, but coupling is preserved, again indicating a non-­ dissociative process, presumably Berry pseudo-­rotation. A series of complexes [PdX(Me6tren)]Y [X = Y = Br, I or SCN; X = Cl, Br, I or SCN; Y = PF6− or BF4−; Me6tren = N(CH2CH2NMe2)3] were examined by 1H NMR spectroscopy in different solvents. In protic and aprotic solvents, there

Scheme 3.20

Ligand Replacement Reactions of Metal Complexes

179

is a rapid (NMR time scale) intramolecular rearrangement, as shown in Figure 3.21, involving four-­and five-­coordinate complexes.240 At low temperature the rate is significantly retarded and the equilibrium shifts to the right. In aqueous solution the free arm of the tridentate Me6tren (5) is readily protonated. Anation of trigonal bipyramidal [Cu(Me6tren)(OH2)]2+ by NCO−, Cl− and Br− has been reported; these reactions are considerably slower (kan is in the range 1–60 s−1 at 10–35 °C) than normally encountered for reactions of six-­coordinate complexes of copper, and are believed to proceed by an Id mechanism.241 X-­ray crystallographic studies on a tetradentate N2S2 macrocyclic complex of palladium, [Pd(tetraL)]Cl2·2H2O, showed the existence of two five-­ coordinate isomers (6 and 7, Figure 3.22), one with a Pd–Cl distance of 3.20 Å and the other with a Pd–Cl distance of 3.68 Å.242 This supports the results of previous solution studies, which indicated that as Cl− approaches [Pd (tetraL)]2+, the square folds back about the N–Pd–N axis to form a trigonal bipyramidal [PdCl(tetraL)]+.243 A variable-­temperature and variable-­pressure 13C NMR spectroscopic study in acetonitrile was carried out of solvent exchange on the five-­coordinate nickel(ii) complex 8.244 The activation volume of 2.3 ± 1.3 cm3 mol−1 suggests a dissociative mechanism. The UV-­visible spectra of 8 are compatible with a five-­ or four-­coordinate equilibrium, so it is reasonable that the solvent exchange on the five-­coordinate species proceeds dissociatively via the four-­ coordinate complex

Figure 3.21  Intramolecular  rearrangement of a Pd(ii) complex (see text).

Figure 3.22  Two  isomers of a Pd(ii) complex (see text).

Chapter 3

180

  



  

Kinetic studies on the following fast reactions have also been reported:245 methanol [Ni  X   QAS ]  Y    [Ni  Y   QAS ]  X 

(3.84)

where QAS = tris(o-­diphenylarsinophenyl), X = Br for Y = NO2, I, NCS, N3 or CN and X = I, NO2, Cl, N3, NCS or Br for Y = CN. The exchange of monoolefins with [Fe(CO)4(PhCH=CH2)] is slow and occurs by a dissociative path, whereas the exchange of diolefins with [Fe(CO)3(diolefin)] is predominantly associative.246 Several quadridentate arsenic donor ligands are effective in stabilizing five-­ coordinate complexes, e.g. Ni(ii) complexes of the type [M(LLLL)(X)]+. Complexes of this type (9), with M = Ni, Pd or Pt, LLLL = tetars and X = halide, are thermodynamically stable in non-­polar solvents but are fairly labile kinetically. A variable-­temperature NMR spectroscopic study of solutions of such complexes showed that exchange processes are complicated, owing to the involvement of three different bimolecular mechanisms. The stereochemical characteristics of this ligand are important in elucidating the possible reaction pathways.247 The kinetics of reaction of the [NiCl(QAS)]+ cation 9, where QAS is a tripodal ligand [10, tris(o-­diphenylarsenophenyl)arsine], with cyanide or thiocyanate in methanol have been reported,248 but no firm assignment of mechanism was possible owing to inherent difficulties.

Complexes of the ligand 1,5-­diazacyclooctane-­N,N′-­diacetate (dacoda) are five-­coordinate since one of the ligand alkyl protons blocks the potential sixth coordination position.249 The kinetics of formation and dissociation of the complexes [M(dacoda)(OH2)] (M = Co, Ni) have been studied.250 Formation may involve an internal conjugate-­base mechanism (ICB), but a claim that the ICB mechanism is a superfluous postulate (see Chapter 5) should be noted.251 Dissociation rates are acid dependent, with kobs = k1 + k2[H+]. Rate constants and activation parameters have been reported.250

Ligand Replacement Reactions of Metal Complexes

181

Studies on the substitution kinetics of [M(S∩S)2L]-­t ype complexes have been reported. Reactions of the tri-­n-­butylphosphine adducts of bis (O-­dialkyldithiophosphato)nickel(ii) complexes with bidentate nitrogen or phosphorus nucleophiles take place by an SN1(lim.), i.e. D mechanism.252 It now appears that the substitution mechanism operating for this type of complex depends on the charge on the complex, i.e. on the formal oxidation state of the metal. Reactions of the following type:   



[M{S2C2(CF3)2}L]z + L′ → [M{S2C2(CF3)2}L′]z + L

(3.85)

  

with M = Fe or Co and L and L′ = phosphines or phosphites, proceed by a dissociative (D) mechanism when z = −1, but by an associative (A) mechanism when z = 0.253 It seems that substitution reactions of 14-­or 16-­electron systems of this type may proceed by dissociative or associative mechanisms, whereas substitution reactions of 18-­electron systems always take place by a dissociative mechanism.253,254 Water exchange is relatively slow for the Co(ii) complex [Co(LLLL)(OH2)]2+ (LLLL = 1,4,8,11-­tetramethyl-­1,4,8,11-­tetraazacyclotetradecane). Studies by 17 O NMR spectroscopy gave a value of 4.2 × 104 s−1 for kex at 25 °C, with ∆H‡ = 36.5 kJ mol−1 and ∆S‡ = −34 J K−1 mol−1. The markedly negative value of ∆S‡ was indicative of an associative mechanism for the water exchange.255 A single kinetic process was observed in a stopped-­flow investigation of the dissociation of [Co(Me6tren)(OH2)]2+; the process was ascribed to dissociation of the first primary amine group.256 The kinetics and the likely mechanism of reaction of the five-­coordinate Rh(ii) species [RhCl2(H) (PR3)2]− with di-­µ-­chloroplatinum(ii) and -­palladium(ii) complexes have been reported.257 The dipalladium complexes trans-­[Pd2(µ-­Cl)2Cl2(PR3)2] react with the rhodium complex to give di-­µ-­chloro species containing one Rh and one Pd [and similarly for reaction of the Pt(ii) complex]. The rates of formation of the products, trans-­ [MCl2(PR′3)(PR″3)], depend strongly on the nature of the phosphines PR′3 and PR″3.257 The mechanism of the reaction of [Ir(X)(COD)(phen)] (X = Cl, I or NCS) with en involves the [Ir(cod)(en)(phen)]+ cation, presumably containing unidentate en, as an intermediate as shown in eqn (3.86).   

 (3.86)   

The structures in solution in dichloromethane of [NiX2(PMe3)3] (X = Cl, Br, I, CN) and their ligand exchange reactions were investigated by 1H and 31P NMR spectroscopy. These compounds are rare amongst five-­coordinate compounds in exhibiting stereochemical rigidity at accessible temperatures, in this case at and below 200 K. At such temperatures they have trigonal bipyramidal structures; at higher temperatures they undergo intramolecular

Chapter 3

182

exchange, which is fastest for the iodo compound. Even at 300 K, the following equilibrium can be observed:   



[Ni(CN)2(PMe3)3] ⇌ [Ni(CN)2] + PMe3

(3.87)

  

The rate law for exchange of PMe3 indicates a dissociative mechanism.258 A dissociative mechanism with rate-­limiting breaking of an Ni to terminal As bond also operates for the reactions of trigonal bipyramidal [NiX(QAS)]+ with the nucleophiles CN−, NO2−, N3−, SCN−, I−, thiourea and PPh3.259 For the similar reaction in methanol solution:   

  

kf

[NiBr(diars)2 ]  SCN  #[Ni(SCN)(diars)2 ]  Br  kb

(3.88)

the kf and kb values for approach to equilibrium have been reported and the results are consistent with a limiting dissociative (D) process. The authors expressed the view that all the substitutions at low-­spin five-­coordinate nickel(ii) (and other d8 complexes) will be dissociative in character, and that associative substitution will be found only for five-­coordinate centres with fewer than 18 electrons.254,260

3.8  Complexes of Higher Coordination Number Six-­coordinate complexes are discussed fairly elaborately in Chapter 4, but relatively few studies have been made on complexes of coordination number higher than six. A few representative examples are discussed below. Investigations on the photochemical aquation261–263 of [Mo(CN)8]4− (I) substitution and photochemical reduction264 of [Mo(CN)8]3− (II) have been reported. The primary step in the photochemical aquation of I is reversible CN− loss,261 whereas the photochemical reduction of II involves aquation (CN− loss) followed by electron transfer within the ion-­pair intermediate, [Mo (CN)7(OH2)]2−[CN]−. Studies on the photochemical aquation262 of [W(CN)8]4− have been reported.262 Complexes [W(LL)4] (LL = oxine or its substituted derivatives) are inert to substitution like eight-­coordinate cyano complexes.265 Ligand exchange in [Nd(tren)2]3+ has been reported.266 BH4− exchange in 12-­coordinate [Zr(BH4)4] with LiBH4 in ether solution is associative, but ligand exchange of [Zr(BH4)4] in the gaseous state is dissociative.267 Merbach and co-­workers268 reported studies on water exchange reactions of eight-­coordinate Lnaq3+ ions; kinetic parameters of these reactions are given in Table 3.25. Small negative values of ∆V‡ may be an indication of an associative interchange process (Ia) in which both Ln3+–*OH2 bond formation and Ln3+⋯OH2 bond dissociation are involved with some predominance of the former in the transition state. Scaq3+ and Yaq3+ are also extremely labile (for k ex H2 O see Chapter 1, Figure 1.3). Rate constants kf for complex formation by [Sc(OH2)7]3+ are ∼107 M−1 s−1.269 The axial oxygens of [UO2(OH2)5]2+ are quite inert; their exchange with O of H2O solvent has a rate constant of 9.9 × 10−9 s−1 at 25 °C, but the equatorial H2O are vary labile ( k ex H2 O = 7.6 × 105 s−1 at 25 °C).270

Ligand Replacement Reactions of Metal Complexes

183

Table 3.25  Kinetic  parameters for water exchange of [Ln(OH2)8] ions in aqueous 3+

solutions (kex values at 25 °C).

Ln in [Ln(OH2)8]3+

10−7kex/s−1

∆H‡/kJ mol−1

∆S‡/J K−1 mol−1

∆V‡/cm3 mol−1

Gd Tb Dy Ho Er Tm Yb

83.0 55.8 43.4 21.4 13.3 9.1 4.7

12.0 12.1 16.6 16.4 18.4 22.7 23.3

−30.9 −36.9 −24.0 −30.5 −27.8 −16.4 −21.0

−3.3 −5.7 −6.0 −6.6 −6.9 −6.0 —

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187. (a) J. C. Lockhart and W. S. Mossop, Chem. Commun., 1971, 61; (b) I. N. Marov, M. N. Vargaftik, V. K. Belyaeva, G. A. Evnkova, E. Hoyer, R. Kirmse and W. Dietzsch, Russ. J. Inorg. Chem., 1980, 25, 101. 188. D. C. Olson and D. W. Margerum, J. Am. Chem. Soc., 1963, 85, 297. 189. W. L. Darby and L. M. Vallarino, Inorg. Chim. Acta, 1979, 36, 253. 190. M. U. Fayyaz and M. W. Grant, Aust. J. Chem., 1979, 32, 2159. 191. P. Ray, Chem. Rev., 1961, 61, 313. 192. (a) T. C. Creitz, R. Guell and D. L. Wampler, J. Chem. Soc., Chem. Commun., 1969, 1371; (b) C. H. L. Kennard, G. Smith and E. J. O. Reilly, Inorg. Chim. Acta, 1983, 7, L113. 193. (a) D. Banerjea, Pure Appl. Chem., 1988, 60, 1337; (b) T. Paramanik, A. Das and D. Banerjea, Transition Met. Chem., 1991, 16, 332. 194. L. S. Meriwether and M. L. Fiene, J. Am. Chem. Soc., 1959, 81, 4200. 195. J. P. Day, F. Basolo, R. G. Pearson, L. F. Kangas and P. M. Henry, J. Am. Chem. Soc., 1968, 90, 1925. 196. J. P. Day, D. Dlemente and F. Basolo, Inorg. Chim. Acta, 1969, 3, 363. 197. A. J. Rest and J. J. Turner, J. Chem. Soc., Chem. Commun., 1960, 1026. 198. K. L. R. Brower and T.-­S. Chen, Inorg. Chem., 1973, 12, 2198. 199. G. Reichenbach, G. Innorta and A. Foffani, Inorg. Chim. Acta, 1969, 3, 139. 200. B. J. Plankey and J. V. Rund, Inorg. Chem., 1979, 18, 957. 201. M. Eigen and K. Tamm, Z. Electrochem., 1962, 66(93), 107. 202. R. E. Connick and D. N. Fiat, J. Chem. Phys., 1963, 39, 1349. 203. P.-­A. Pittet, G. Elbaza, L. Helm and A. E. Meerbach, Inorg. Chem., 1990, 29, 1936. 204. T. W. Swaddle, Inorg. Chem., 1983, 22, 2663. 205. W. G. Baldwin and D. R. Stranks, Aust. J. Chem., 1968, 21, 2161. 206. S. F. Lincoln, A. C. Sandercock and D. R. Stranks, Aust. J. Chem., 1977, 30, 271. 207. A. Barabas, J. Coord. Chem., 1973, 3, 91. 208. A. Barabas, Inorg. Nucl. Chem. Lett., 1970, 6, 775. 209. J. Crea and S. F. Lincoln, J. Chem. Soc., Dalton Trans., 1973, 2075. 210. H. Brunner, Angew. Chem., Int. Ed., 1971, 10, 249. 211. C. G. Grimes and R. G. Pearson, Inorg. Chem., 1974, 13, 971. 212. U. Mayer, Pure Appl. Chem., 1975, 41, 291. 213. P. J. White, M. J. Kaus, J. O. Edrwards and P. H. Rieger, J. Chem. Soc., Chem. Commun., 1976, 429. 214. (a) G. W. Parshall, J. Am. Chem. Soc., 1974, 96, 2360; (b) M. Meier, F. Basolo and R. G. Pearson, Inorg. Chem., 1969, 8, 793; (c) R. D. Johmston, F. Basolo and R. G. Pearson, Inorg. Chem., 1971, 10, 247; (d) L. H. Pignolett, D. Forster and W. deW. Horrocks Jr., Inorg. Chem., 1968, 7, 828; (e) L. H. Pignolett and W. deW. Horrocks Jr., J. Am. Chem. Soc., 1968, 90, 822. 215. R. E. Gentzler, T. R. Stengle and C. H. Langford, J. Chem. Soc., Chem. Commun., 1970, 1257. 216. J. Lagrange and P. Lagrange, Bull. Soc. Chim. Fr., 1972, 19. 217. K. Kustin and S.-­T. Lin, J. Am. Chem. Soc., 1973, 95, 2487. 218. K. Gilbert and K. Kustin, J. Am. Chem. Soc., 1976, 98, 5502. 219. K. Kustin and D. L. Toppen, J. Am. Chem. Soc., 1973, 95, 3564.

192

Chapter 3

220. A. Bhoumik and D. Banerjea, Indian J. Chem., 1990, 29A, 1208. 221. D. Meyerstein, Inorg. Chem., 1975, 14, 1716. 222. H. L. Hodges and M. A. de Araugo, Inorg. Chem., 1982, 21, 3236. 223. (a) M. M. Farrow, N. Purdie and E. M. Eyring, Inorg. Chem., 1975, 14, 1584; (b) N. Purdie, M. M. Farrow, M. Steggall and E. M. Eyring, J. Am. Chem. Soc., 1975, 97, 1078. 224. (a) M. R. Baloga and J. E. Early, J. Phys. Chem., 1963, 67, 96; (b) R. H. Hllyer and H. W. Baldwin, Can. J. Chem., 1967, 45, 418; (c) R. K. Murmann, Inorg. Chem., 1977, 16, 46. 225. A. Okumura, M. Kitani, Y. Toyomi and N. Okazaki, Bull. Chem. Soc. Jpn., 1980, 53, 3143. 226. H. von Felton, B. Wernli, H. Gamsjager and P. Baertschi, J. Chem. Soc., Dalton Trans., 1978, 496. 227. H. Goff and R. K. Murmann, J. Am. Chem. Soc., 1971, 93, 4184. 228. (a) H. C. McDonald, Diss. Abstr., 1960, 21, 454; (b) K. H. Heckner and R. Landsberg, J. Inorg. Nucl. Chem., 1967, 29, 413; (c) R. K. Murmann, J. Phys. Chem., 1967, 71, 97; R. K. Murmann, J. Am. Chem. Soc., 1971, 93, 4184. 229. (a) J. B. Goodard and A. M. Gonas, Inorg. Chem., 1973, 12, 574; (b) D. M. Druskovich and D. L. Kepert, J. Chem. Soc., Dalton Trans., 1975, 947; (c) P. Moore, S. F. Kettle and R. G. Wilkins, Inorg. Chem., 1966, 5, 220. 230. (a) P. Lagrange and J.-­P. Schwing, Bull. Soc. Chim. Fr., 1970, 1340; (b) J.-­P. Collin, P. Lasgrange and I.-­P. Schwing, J. Less-­Common Met., 1974, 36, 117; (c) J. P. Collin and P. Lagrange, Bull. Soc. Chim. Fr., 1074, 773, 777; (d) D. M. Druscovich and D. L. Kepert, Aust. J. Chem., 1975, 28, 2365. 231. J. R. Pladziewicz and J. H. Espenson, Inorg. Chem., 1971, 10, 634. 232. R. Banarad, B. Perlmutter-­Hayman and M. A. Wolf, J. Phys. Chem., 1969, 73, 4301. 233. S. B. Brown, J. O. Edwards, I. J. Herman, P. Jones, R. J. Mills and J. E. Earley, Inorg. Chim. Acta, 1969, 3, 351. 234. M. Biquard and M. Lamache, Bull. Soc. Chim. Fr., 1971, 32. 235. K.-­H. Tytko and O. Glemser, Adv. Inorg. Chem. Radiochem., 1976, 39, 239. 236. R. Huq, A. J. Poe and S. Chawla, Inorg. Chim. Acta, 1980, 38, 121. 237. H. Yesaka, T. K. K. Yasufuku and S. Nagakura, J. Am. Chem. Soc., 1983, 105, 6249. 238. S. B. McCullen, H. W. Walker and T. L. Brown, J. Am. Chem. Soc., 1982, 104, 4007. 239. J. H. Nelson and N. W. Alcock, Inorg. Chem., 1982, 21, 1196. 240. S. N. Bhattacharya and C. B. Senoff, Inorg. Chem., 1983, 22, 1607. 241. J. F. Coates, P. R. Collins and S. F. Lincoln, Aust. J. Chem., 1980, 33, 1381. 242. G. Ferguson, R. McCrindle, A. J. McAlees, M. Parvez and D. K. Stephenson, J. Chem. Soc., Dalon Trans., 1983, 1865. 243. R. McCrindle, G. Ferguson, A. J. McAlees, M. Parvez and D. K. Stephenson, J. Chem. Soc., Dalton Trans., 1982, 1291. 244. L. Helm, P. Meier, A. E. Merbach and P. A. Tregloan, Inorg. Chim. Acta, 1983, 73, 1.

Ligand Replacement Reactions of Metal Complexes

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245. H. L. Collier and E. Grimley, Inorg. Chem., 1980, 19, 511. 246. P. M. Burkinshaw, D. T. Dixon and J. A. S. Howell, J. Chem. Soc., Dalton Trans., 1980, 999. 247. B. Bosnich, W. G. Jackson and S. T. D. Lo, Inorg. Chem., 1975, 14, 2998. 248. E. Grimley, J. M. Grimley, T.-­D. Li and D. Emerich, Inorg. Chem., 1976, 15, 1716. 249. D. F. Averill, J. I. Legg and D. L. Smith, Inorg. Chem., 1972, 11, 2344. 250. C. Chatterjee and Th. A. Kaden, Helv. Chim. Acta, 1975, 58, 1881. 251. R. B. Jordan, Inorg. Chem., 1976, 15, 748. 252. D. A. Sweigart and P. Heidtmann, J. Chem. Soc., Dalton Trans., 1975, 1686. 253. D. A. Sweigart, J. Chem. Soc., Chem. Commun., 1975, 688; Inorg. Chim. Acta,1976, 18, 179. 254. C. A. Tolman, Chem. Soc. Rev., 1972, 1, 337. 255. P. Meier, A. Merbach, S. Bürki and T. A. Kaden, J. Chem. Soc., Chem. Commun., 1977, 36. 256. S. F. Lincoln and C. D. Hubbard, Proc. 16th Int. Confc. Coord. Chem. (ICCC), Dublin, 1974, Abs. No. 3.30, see Chem. Abstr., 1976, 85, 37.623. 257. R. Huis and C. Masters, J. Chem. Soc., Dalton Trans., 1976, 1796. 258. P. Meier, A. E. Merbach, M. Dartiguenave and Y. Dartiguenave, J. Am. Chem. Soc., 1976, 98, 6402. 259. E. B. Grimley and H. L. Collier, Proc. 173rd Meeting of the American Chemical Society (ACS), Abs. No. Inorg. 222. 260. D. A. Sweigart, Inorg. Chim. Acta, 1977, 23, L13; C. A. Tolman, Chem. Soc. Rev., 1972, 1, 337. 261. R. D. Wilson, V. S. Sastri and C. H. Langford, Can. J. Chem., 1971, 49, 679. 262. R. P. Mitra, B. K. Sharma and H. Mohan, Aust. J. Chem., 1972, 25, 499. 263. M. Shirom and Y. Siderar, J. Chem. Phys., 1972, 57, 1013. 264. G. W. Gray and J. T. Spence, Inorg. Chem., 1971, 10, 2751. 265. W. D. Bonds and R. D. Archer, Inorg. Chem., 1971, 10, 2057. 266. M. F. Johnson and J. H. Forsberg, Inorg. Chem., 1972, 11, 2683. 267. N. Davies, D. Saunders and M. G. H. Wallbridge, J. Chem. Soc. A, 1970, 29. 268. (a) C. Cossy, L. Helm and A. E. Merbach, Inorg. Chem., 1988, 27, 1973; 1989, 28, 2699; (b) K. Micskei, H. Powell, L. Helm, E. Brucher and A. E. Merbach, Magn. Reson. Chem., 1993, 31, 1011; (c) D. H. Powell and E. Merbach, Magn. Reson. Chem., 1994, 32, 739. 269. G. Geier, Ber. Bunsenges. Phys. Chem., 1965, 69, 617. 270. W.-­S. Jumg, M. Narada and H. Fukutomi, Bull. Chem. Soc. Jpn., 1984, 57, 2317; 1985, 58, 938; 1986, 59, 3761; 1987, 60, 489; 1988, 61, 3895. 271. L. M. Venanzi, in Coordination Chemistry-­20, ed. D. Banerjea, IUPAC, Pergamon Press, Oxford, 1980, see pages 102–103. 272. (a) M. Meier, F. Basolo and R. G. Pearson, Inorg. Chem., 1969, 8, 795; (b) R. J. Angelici, Organomet. Chem. Rev., 1968, 3, 173; (c) H. G. Schuster-­ Wolden and F. Basolo, J. Am. Chem. Soc., 1966, 88, 1657; (d) R. B. King, Adv. Organomet. Chem., 1964, 2, 177.

Chapter 4

Ligand Replacement Reactions of Octahedral Complexes Octahedral complexes of chromium(iii), cobalt(iii), ruthenium(ii), rhodium(iii), iridium(iii), iridium(iv), platinum(iv), etc., are fairly inert and are well known. A very large volume of work has been reported on the kinetics of ligand substitution in octahedral complexes, which has led to an understanding of the mechanistic features of these processes.1 Various pieces of evidence point to the conclusion that in general, the reactions of Co(iii) complexes proceed by a dissociative process, whereas the corresponding reactions of chromium(iii) and more particularly those of rhodium(iii) are associative in character. Octahedral complexes of Co(iii) (low-­spin type) and Cr(iii) have been extensively investigated and hence discussion in this chapter will be focused on these complexes. Ligand replacement reactions of Pt(iv) complexes are almost invariably catalysed by Pt(ii) and are discussed in Chapter 3, although uncatalysed (direct) reactions are also known.2

4.1  A  quation/Solvolysis, Anation/Formation and Ligand Exchange Reactions 4.1.1  Effect of Leaving Ligand A comparison of the rates of replacement of ligand L in cobalt(iii) complexes of the type [Co(NH3)5L] with a molecule of water (aquation reaction) forming [Co(NH3)5(OH2)]3+ in acidic solution (pH < 5) has shown that the rate decreases in the following sequence of L, with increasing thermodynamic stability of the complex, i.e. increasing affinity of Co(iii) for L: NO3− > I− >   Mechanisms of Reactions of Metal Complexes in Solution By Debabrata Banerjea and M. K. Bharty © Debabrata Banerjea and M. K. Bharty 2023 Published by the Royal Society of Chemistry, www.rsc.org

194

Ligand Replacement Reactions of Octahedral Complexes −

−

2−

−

−

−

195 −

−

3−

Br > H2O > Cl > SO4 > H2PO4 > F > MeCO2 > NH3 > NCS > N3 > PO4 ≈ NO2− ≈ OH−. Rate data are given in Table 4.1. It can be seen that the rate depends significantly on the nature of the leaving group, L. The aquation rates of halidopentaammine complexes of Cr(iii) also decrease in the sequence I− > Br− > Cl− > F− > NCS−,4 and this is in keeping with the corresponding Cr–X bond strength. The aquation rates also show a greater dependence on the nature of the leaving group than in the analogous complexes of Co(iii).5 These results suggest a dissociative mechanism; a ∆V‡ of −8.6 cm3 mol−1 for aquation of [Cr(NCS)(NH3)5]2+ supports the dissociative pathway for the reaction.6 In contrast to the significant dependence of the rate of aquation on the leaving group, the formation/anation rate for Co(iii) complexes is fairly insensitive to the nature of the entering ligand, as shown by some representative data in Table 4.2. Based on these observations, it is logical to conclude that for ligand replacement reactions of Co(iii) (Oh, low spin) the rate is more sensitive to the kind of bond being broken than those formed, i.e. the mechanism is dissociative. Further support comes from the observation that in [CoX(NH3)5]2+ no direct replacement of X− with Y− takes place, and the reaction always occurs by aquation followed by anation of the aqua complex by Y−. This indicates that the energetically significant process is breaking of the CoIII–X bond, and the bond making has so little significance that the species present in greatest concentration in the solution (H2O in aqueous solution) acts as the initial entering group, and the nucleophilicity of the entering group (Y−) is of little significance. The limiting rate constant values in Table 4.2 refer to the situation expressed in eqn (1.32) in Chapter 1, when Kos[L] ≫ 1 (Yn− = L), giving kobs ≈ kos attained at large [Yn−]. It should be noted that the value of the ratio klim/kex Table 4.1  Rate  constants for aquation of complexes of Co(iii), [Co(NH3)5L]n+, including solvent (H2O) exchange rate of [Co(NH3)5(OH2)]3+ at 25 °C.3,4

L

k/s−1

L

k/s−1

H2O (MeO)3PO NO3− I− Cl− Br−

5.9 × 10−6 2.5 × 10−4 2.7 × 10−5 8.3 × 10−6 1.7 × 10−6 6.3 × 10−6

SO42− H2PO4− MeCO2− F− N3− NCS− (N-­bonded)

1.2 × 10−6 2.6 × 10−7 1.2 × 10−7 8.6 × 10−8 2.1 × 10−9 5.0 × 10−10

Table 4.2  Limiting  rate constants for anation by Yn− and water exchange (kex) for [Co(NH3)5(OH2)]3+ at 45 °C.7

Yn−

Klim/s−1

H2O N3− SO42− Cl− NCS−

1.0 × 10−4 (kex value) 1.0 × 10−4 2.4 × 10−5 2.1 × 10−5 1.6 × 10−5

Chapter 4

196

Table 4.3  Values  of Koskos (i.e. kobs), Kos and kos for the formation of nickel(ii) complexes from [Ni(OH2)6]2+ and unidentate ligands at 25 °C.8

Ln−

10−3Koskos/M−1 s−1

Kos/M−1

10−4kos/s−1

MePO42− MeCO2− NCS− F− HF H2O NH3 Pyridine H2N(CH2)2NMe3+

290 100 6 8 3 — 5 ∼4 0.4

40 3 1 1 0.15 — 0.15 0.15 0.02

0.7 3 0.6 0.8 2 3 3 ∼3 2

is 1.0 (N3−), 0.24 (SO42−), 0.21 (C1−) and 0.16 (NCS−). The fact that the ratio is fairly insensitive to the entering ligand further supports a dissociative mechanism, likely to be Id, and this is much more true for [Co(CN)5(OH2)]2− (see Section 4.1.10). Results reported so far by different investigators indicate that kobs/Kos = kos [see eqn (1.34) in Chapter 1] values for the formation of [Ni(OH2)5L]n+ from [Ni(OH2)6]2+ and a variety of L (uncharged and anionic) are also fairly insensitive to the nature of the entering ligand L, suggesting a dissociative activation for these reactions; for L of different charge type, kobs/Kos = kos values were logically considered for comparison (Table 4.3).8 Replacement of the aqua ligand in [Ru(NH3)5(OH2)]2+ by a wide variety of ligands, L, including NH3, py and substituted pyridines, MeCN, PhCN and imidazole in weakly acidic solution follows the rate law8 rate = k[complex][L]. The k value varies only slightly with variation of L; a dissociative (D) mechanism was proposed. The observed effects of steric hindrance on reactivity also support the D mechanism. Thus, the very much slower rate of formation of the 2,6-­lutidine complex compared with that of the pyridine complex (the ratio of rates is 5 × 10−4 ) can be explained in terms of very unequal competition between 2,6-­lutidine and H2O, but more equal competition between py and H2O for the [Ru(NH3)5]2+ intermediate.9a The activation parameters for these reactions are in a narrow range of ∆H‡ = 15.1–16.2 kcal mol−1 and ∆S‡ = −7 to −13 cal K−1 mol−1. The values for the various entering nucleophiles are equal within the limits of experimental errors and moreover are very close to the ∆H‡ and ∆S‡ values reported9b for the reactions of [Ru(NH3)5(OH2)]2+ with N2 and N2O. Similarity of the activation parameters for a wide range of entering nucleophiles provides further convincing evidence for the proposed D mechanism. Similar behaviour has been reported10 for the entry of several other ligands L in reactions with [Ru(NH3)5(OH2)]2+ (data at 25 °C): L −

103 kL/M−1 s−1 a

[Ru(NH3)5(N2)]2+a

NO2

Isonicotinamide

py

CO

N2

7.2

6.0

11.8

12

8.0 3.6

This forms the well-­known [(H3N)5Ru–N2–Ru(NH3)5]4+.

Ligand Replacement Reactions of Octahedral Complexes

197

An interchange dissociative (Id) mechanism has been proposed for these reactions. The hexaaqua complex [Ru(OH2)6]2+ has also been reported11 to react with anions (anation) including ClO4− by a dissociative (Id) process, since kan is fairly insensitive to the nature of the anion (data at 25 °C): X− 103kan/M−1 s−1

Cl−

Br−

I−

ClO4−

8.5

9.7

9.0

3.2

Similar dissociative anation has been reported12 for the [Ru(NH3)5(OH2)]3+ anion (data at 54.7 °C): X− 103kan/M−1 s−1

Cl−

Br−

I−

2.1

1.32

0.74

In contrast to the above-­mentioned behaviour of complexes of Co(iii), Ru(ii), Ni(ii) and even Ru(iii), the rate constants for anation of [Ti(OH2)6]3+ indicate a very pronounced variation of rate with the nature of the entering ligand (Table 4.4), pointing to an associative mechanism, presumably Ia. [V(OH2)6]3+ shows similar behaviour and so also [M(NH3)5(OH2)]3+ (M = Rh, Ir) (see Table 5.5, Chapter 5). For the anation of [Cr(OH2)6]3+ by Cl−, Br−, I−, NO3−, SCN− and NCS−, the rate constants at 25 °C span the wide range 8 × 10−10–7.3 × 10−7 M−1 s−1, indicating the associative character of the reaction14,15 as for [Ti(OH2)6]3+. A significant observation was made by Lalor and Long16 that for the aquation of several Co(iii) complexes of the type [Co(NH3)5L]2+, the activation energy, Ea, increases linearly with the wavenumber, in cm−1, of the first d–d absorption band of the complex (which is related to the average Dq value of the complex; see Chapter 2) in the following sequence of L: Br− < NO3− < NCS− (N-­bonded) < NO2− (N-­bonded). This also suggests that Co–L bond breaking Table 4.4  Rate  constants for anation of [Ti(OH2)6]3+and its water exchange rate constant (kex) at 13 °C (entering ligand Yn−).13a

Yn−

k/M−1 s−1a

H2O NCS− ClCH2CO2− MeCO2−

8.6 × 103b 8.0 × 103c 2.1 × 105 1.8 × 106d

I n modern SI convention, M−1 s−1 is expressed as dm3 mol−1 s−1. This is kex (s−1)/55.5 M, where 55.5 M is the molar concentration of water in dilute aqueous solution (kex = 4.77 × 105 s−1).13b c Data at 9 °C. d For Tiaq3+ + HC2O4− → [Ti(C2O4)aq]+ + H+; k10 °C = 3.9 × 105 M−1 s−1.13c a b

Chapter 4

198

is important in the transition state, but gives no information as to whether or not bond making by the incoming water molecule is of any importance. Based on rate constant values for aquation of [MX(NH3)5]2+ complexes (M = CrIII, CoIII, RuIII, RhIII and IrIII; X = Cl or Br)17 mentioned below the reactivity order for these MIII complexes is Cr > Co > Ru > Rh > Ir: MIII Cr −1

kaq/s at 25 °C

Cl Br

Co −6

9.3 × 10 6.8 × 10−5

Ru −6

1.7 × 10 6.3 × 10−6

Rh −7

7.1 × 10 8.7 × 10−7

Ir −7

1.5 × 10 6.3 × 10−8

1.1 × 10−9 1.1 × 10−9

This order, except for the relative positions of Co(iii) and Ru(iii), is what is expected on the basis of ligand field activation energy (LFAE) values taking into consideration the fact that the value of Dq increases in the sequence 3d < 4d < 5d (see Chapter 1). Evidence based on the linear free energy relationship (LFER) (see Section 4.1.6) and other evidence suggest a dissociative mechanism for the Co(iii) complexes in which there is almost complete dissociation of the Co–X bond in the transition state with insignificant bond formation by the entering nucleophile H2O. Hence for the Co(iii) complexes, the value of LFAE is 4DqCo (see Chapter 1). However, the reactions of Ru(iii), Rh(iii) and Ir(iii) are associative (A) processes based on ∆V‡ values and other evidence mentioned in the following sections, hence their LFAE values are 1.7Dq, assuming a pentagonal bipyramidal species in the transition state. Although DqRu will be different from DqCo, it is possible to make an estimate of the ratio DqRu/DqCo from available spectroscopic Dq values.18 Thus, the ratio is 1.4 for [Fe(OH2)6]3+/2+of the 3d series. Assuming the same ratio to be valid (approximately) for [Ru(NH3)6]3+/2+(both H2O and NH3 are σ-­bonding ligands) and the reported Dq value of 2810 cm−1 for [Ru(NH3)6]2+, the estimated Dq value for [Ru(NH3)6]3+ is 3930 cm−1. Since the Dq value of [Co(NH3)6]2+ is 2290 cm−1, the value of DqRu ≈ 1.7DqCo, and this may be true, albeit approximately, for [MCl(NH3)5]2+ complexes also. Hence the LFAE value of 1.7DqRu is ca. 2.9DqCo, which is perceptibly less than the LFAE value of 4DqCo for the Co(iii) complex, which predicts the reactivity order Ru > Co, but the reverse order is observed, presumably owing to the higher bond energy of the RuIII−X bond.

4.1.2  Effect of Charge on Reaction Rate With everything else being the same, an increased positive charge (overall) of the complex should make bond breaking between the metal and a ligand more difficult, hence we expect a decrease in rate with increasing positive charge in a dissociative process. The rates for water exchange for the aqua ions (see Figure 1.3, Chapter 1) of the main group elements decrease in any period of the Periodic Table with increase in positive charge of the metal, e.g. Na+ ≫ Mg2+ ≫ A13+. This, coupled with similarity of kex with kf (rate constant for the formation of their complexes from the aqua ions), suggests a dissociative activation. In the case of transition metal complexes, LESE

Ligand Replacement Reactions of Octahedral Complexes

199

considerations (see Chapter 1) are superimposed on those of charge. Moreover, for a metal ion in a given oxidation state, the overall charge can be varied only by introducing differently charged ligands, which may lead to changes in σ-­ and π-­bonding effects. Despite this, a substantial increase in rate with decrease in overall positive charge is strong evidence for dissociative activation, as in the examples of aquation of the complex of Co(iii) (Table 4.5); data showing the contrasting behaviour of Cr(iii) are also included in Table 4.5 for comparison (compare the behaviour of the first two complexes). The data indicate that in the case of Co(iii) a decrease in the overall positive charge has a very significant effect in increasing the rate of aquation of the Cl− ligand, whereas the corresponding effect in the case of Cr(iii) is far less significant, as indicated by the ratio of kaq values at 25 °C (Table 4.6). These findings suggest that the Co(iii) reactions are dissociative with extensive bond breaking in the transition state, whereas the analogous reactions of Cr(iii) are interchange associative processes with less significant bond breaking in the transition state. The extreme sensitivity of the rate to the overall charge for the Co(iii) complexes is in accord with a process in which bond breaking is primarily important in the formation of the transition state, indicating a dissociative mechanism, presumably Id, for the Co(iii) complexes. Table 4.5  Effect  of overall charge on a complex on the rate of its aquation.a kaq/s−1 at 25 °C M = Co(iii) b

trans

cis

Complex 2+

[MCl(NH3)5] [MCl2(NH3)4]+ [MCl(en)2(NH3)]2+ [MCl2(en)(NH3)2]+ [MCl2(en)2]+

M = Cr(iii) b

cis

−6

1.7 × 10 — 3.7 × 10−7 — 2.4 × 10−4

trans −6

1.9 × 10−3 2.9 × 10−7 2.3 × 10−5 3.1 × 10−5

9.3 × 10 — — — 3.3 × 10−4

4.9 × 10−5 — — 3.9 × 10−5

a

 ata from ref. 3a, Table 3.13, and ref. 3b; see also ref. 17. D The faster rate of the cis isomer is due to the π-­donor character of Cl− (see Figure 4.3).

b

Table 4.6  Relative  effect of overall charge on rate of aquation of analogous complexes of Co(iii) and Cr(iii). Co cis

Rate ratio  MCl2  NH3 4 

 kaq

 MCl  NH3 5 

 kaq





 MCl  NH3 5 

2

 kaq

trans 3

cis

trans

—

1.1 × 10

—

5.3

141

182

35.5

4.2

2

 MCl2  en 2 

 kaq

Cr

Chapter 4

200

4.1.3  Steric and Structural Effects of Spectator Ligands Evidence favouring a dissociative mechanism is the observation that for several complexes of the type trans-­[CoCl2(A–A)2]+, where A–A is ethylenediamine or a C-­alkylethylenediamine, the rate of aquation increases with increasing steric crowding due to substitution in the ethylenediamine molecule,19a as indicated by the data in Table 4.7 for the following reaction:   



trans-­[CoCl2(A−A)2]+ + H2O → [CoCl(A−A)2(OH)2]2+ + Cl−

(4.1)

  

Increasing chelation also slows the rate of aquation in a progressive manner19b in the sequences cis-­[CoCl2(NH3)4]+ > cis-­[CoCl2(en)2]+ > cis-­ [CoCl2(trien)]+ and trans-­[CoCl2(NH3)4]+ > trans-­[CoCl2(en)(NH3)2]+ > trans-­ [CoCl2(en)2]+ and there is a similar trend for the chloropentaamine also (Table 4.8). These results are also compatible with a dissociative mechanism. Rate data for aquation of the three geometrical isomers I, II and III of [CoCl(dien)(tmd)]2+ and of the corresponding isomers of [CoCl(dien)(en)]2+ have been reported.20 In each case, the complex having tmd, which forms a six-­membered ring, aquates much faster than the corresponding complex having en, which forms a five-­membered ring. At 25 °C, the ktmd/ken ratio is 16.8, 31.7 and 2.3 for the complexes I, II and III, respectively.

Thus, whereas the increase in rate is ca. 32-­fold on changing from en (H2NCH2CH2NH2) to tmd (H2NCH2CH2CH2NH2) in the Co(iii) complex II, Table 4.7  Rate  constants illustrating the steric effect of spectator ligands. A–A

105k/s−1 at pH 1 and 25 °C

H2NCH2CH2NH2 H2NCH2CH2NHMe H2NCH2CH2NHEt H2NCH2CH2NH(n-­C3H7) H2NCH2CH(Me)NH2 d,l-­H2NCH(Me)CH(Me)NH2 meso-­H2NCH(Me)CH(Me)NH2 H2NCH2C(Me)2NH2 H2NC(Me)2C(Me)2NH2

3.2 1.7 6.0 12.0 6.2 15.0 42.0 22.0 Instantaneous (very fast) reaction (i.e. k very large)

Ligand Replacement Reactions of Octahedral Complexes

201

Table 4.8  Rate  constants for aquation of some chloropentaamine and dichloro-

tetraamine complexes of Co(iii) at 25 °C (pH 1) for the dichloro complexes and at 35 °C for the chloro complexes19b.

Complex cis-­[CoCl2(en)2]+ cis-­[CoCl2(trien)]+ trans-­[CoCl2(NH3)4]+ trans-­[CoCl2(en)(NH3)2]+ trans-­[CoCl2(en)2]+

104kaq/s−1 2.5 1.5 18 2.3 0.31

Complex

106kaq/s−1

[CoCl(NH3)5]2+ cis-­[CoCl(en)2(NH3)]2+ cis-­[CoCl(trien)(NH3)]2+ [CoCl(dien)(en)]2+ [CoCl(tetraen)]2+

6.7 1.4 0.67 0.52 0.25

the rates are almost identical for the analogous Cr(iii) complexes. The rate differences for the cobalt complexes have been explained in terms of a dissociative mechanism with a trigonal bipyramidal (tbp) intermediate; considerable stretching of the diamine ring is necessary to form the tbp intermediate, which is therefore formed more readily by the relatively more flexible system having a six-­membered bidentate chelate ring (formed by tmd), compared with the five-­membered ring system (formed by en), which restricts the distortion required to form the trigonal bipyramid (IV and V). The near identity of the rates for the two Cr(iii) complexes is in keeping with an associative character in these cases. The kinetics of aquation, mercury(ii)-­assisted aquation and base hydrolysis of some configurational isomers of chloro(diethylenetriamine)bis(monoamine)cobalt(iii) complexes have also been reported.21

The aquation rate constants of several Co(iii) complexes of the type [CoCl2(N4)]+ and [CoCl(X)(N4)]+, where N4 is a macrocyclic ligand such as 1–4, have been reported. The rate is much faster for the tet-­b complex than in the case of its analogous cyclam complex (Table 4.9).22

Chapter 4

202

Table 4.9  Aquation  of cobalt(iii) complexes of some macrocyclic ligands. Complexa

kaq (for Cl− loss, aquation)/s−1 at 25 °C

trans-­[CoCl2(cyclam)]+b trans-­[CoCl2(tet-­b)]+c trans-­[CoCl(NO2)(cyclam)]+ trans-­[CoCl(NO2)(tet-­b)]+ trans-­[CoCl(NCS)(tet-­b)]+ trans-­[CoCl(NCS)(cyclam)]+ trans-­[CoCl(CN)(cyclam)]+ trans-­[CoCl(CN)(tet-­b)]+ trans-­[CoCl(cyclam)(NH3)]2+

1.1 × 10−6 9.3 × 10−4 4.3 × 10−5 4.1 × 10−2 7.0 × 10−7 1.1 × 10−9 4.8 × 10−7 3.4 × 10−4 7.3 × 10−8

a

cyclam = 1,4,8,11-­tetraazacyclotetradecane (1); tet-­b = d,l-­1,4,8,11-­tetraaza-­5,5,7,12,12, 14-­hexamethylcyclotetradecane (2) (the corresponding meso isomer is known as tet-­a). b For aquation of the dibromo complex the kaq value is 2.2 × 10−5 s−1. c For aquation of the dibromo complex the kaq value is 3.8 × 10−2 s−1.

Table 4.10  Aquation  of cobalt(iii) complexes of macrocyclic ligands having unsaturation in the macrocyclic ring.

N4

trans-­[CoX2(N4)]+

kaq (loss of X−)/s−1 at 25 °C

3

X = Cl X = Br X = Cl X = Br

1.6 × 10−2 5.1 × 10−2 2.1 × 10−3 0.12

4

Acceleration of the reaction by the steric effect of the methyl groups of tet-­b suggests a dissociative mechanism (see Chapter 1, Table 1.1). However, for the analogous complexes of the macrocyclic ligands 3 and 4, having unsaturation in the macrocyclic ring, the rates are much faster than for the tet-­b complexes as seen from the kaq values given in Table 4.10.22 It is also of significance that whereas in the cases of the cyclam and tet-­b complexes the bromo complexes react faster than the chloro complexes, the reverse is true for the complexes of the unsaturated macrocyclic ligands 3 and 4. This suggests that cobalt in these complexes has become “soft” (“hard” in the cases of cyclam and tet-­b complexes). The unsaturated macrocyclic ligands effectively reduce the positive charge on Co sufficiently to cause acceleration of the aquation in a dissociative mechanism. For the complexes of the cyclic tetraimine 4 the lability is comparable to those for Co(iii) porphyrins.23 Similar studies have been reported on complexes of some saturated N4 macrocyclic ligands with larger ring sizes.24 At 25 °C, kaq of trans-­[CoCl2(cyclam)]+ is lower than that of trans-­[CoCl2(en)2]+ (see Table 4.8) by a factor of 102; the kaq value of trans-­[CoCl2(cyclam)]+ at 25 °C is lower than that of cis-­[CoCl2(trien)]+ at 25 °C by a factor of 102 due to the chelate effect; the kaq value of trans-­[CoCl (cyclam)(NH3)]2+ at 25 °C is lower than that of cis-­[CoCl(trien)(NH3)]2+ at 35 °C by a factor of 10. All these data indicate a decrease in kaq on replacing an open-­chain quadridentate N4 ligand by a macrocyclic N4 ligand. This is the macrocyclic effect on reaction rate.

Ligand Replacement Reactions of Octahedral Complexes 2+

203 2+

On changing from [MCl(NH3)5] to [MCl(MeNH2)5] , the rate of aquation of the Co(iii) complex (at 50 °C) increases from 3.6 × 10−5 to 6.5 × 10−4 s−1, but that of the Cr(iii) complex decreases from 1.8 × 10−4 to 1.4 × 10−5 s−1, in agreement with dissociative activation for Co(iii) and associative activation for Cr(iii).25 For the aquation of [CoX(L)](LL)2]n+ and [CoX(LL)(LLL)]n+ (L, LL and LLL are uni-­, bi-­ and tridentate ligands, respectively), where X is the leaving group, the plot of log kaq versus δ (N–Co–N), which is the deformation frequency in the IR spectra of these complexes, is linear with a negative slope. This is believed to reflect the ease with which these complexes can distort via a deformational vibrational mode to a trigonal bipyramid structure, and hence considered as evidence for dissociative activation.26 Characterization of a five-­coordinate intermediate provides evidence for a D mechanism in the reaction of an octahedral complex (coordination number = 6). One such piece of evidence has come from scavenging experiments in methanol–water mixed solvents in the solvolysis of [CrX(OH2)5]2+ complexes. Irrespective of the nature of X and the catalysing metal ion, the product ratio R = [Cr(OH2)5(MeOH)]3+/Cr(iii)total is almost the same for any particular solvent composition (Table 4.11).27 However, the unequivocal nature of this evidence has been questioned.28 Further evidence was provided by Posey and Taube in connection with the aquation of [CoX(NH3)5]2+ complexes.29 They observed that in the formation of [Co(NH3)5(OH2)]3+ from [CoX(NH3)5]2+ induced by metal ions in water enriched in 18O, the aqua product had the same 16O : 18O ratio when Hg2+ was used to withdraw X− (Cl−, Br− or I−), showing that a common intermediate results in this case, as is possible for a simple dissociation mechanism. However, when Ag+ or Tl+ was used to assist in the removal of X−, the 16O/18O ratio varied detectably with the nature of X− and the same was also true for spontaneous (uncatalysed) aquation. Since only the more reactive Hg2+generates a common intermediate, it was concluded that aquation of such complexes does not occur by a simple dissociation (D) mechanism and that solvent participation is also of importance in the transition state, the extent of which depends on the nature of X−, which presumably remains weakly bonded in the transition state, hence indicating an Id mechanism. Table 4.11  Catalysed  aquation of some complexes of chromium(iii) in mixed solvents.

R at 25 °C in different solvent compositions (H2O–MeOH)

Catalysing metal ion 28% MeOH

Complex 2+

[CrI(OH2)5] [CrI(OH2)5]2+ [CrI(OH2)5]2+ [CrCl(OH2)5]2+

3+

Tl Hg2+ Ag+ Hg2+

0.20 0.20 0.21 0.19

46% MeOH

64% MeOH

87% MeOH

0.32 0.33 0.32 0.34

0.46 0.47 0.47 0.47

0.74 0.72 0.74 0.72

Chapter 4

204 −

Again, for complexes of the above type, where X was acetate or substituted acetate ligands, where the basicity of the ligand differed by a factor of ca. 104, suggesting a comparable difference in the strength of the Co–X bond, the rates differed only by a factor of ca. 10 (Table 4.12), which is not compatible with a purely dissociation mechanism, where rupture of the Co–X bond should only be of importance in the transition state.30 Also included in Table 4.12 are the rate constants for base hydrolysis, which are dependent on the OH− ion concentration, and Basolo et al.30 showed that the plot of log kOH versus log Ka (Ka for HX) gives a straight line, indicating that the weaker base is also dislodged at a faster rate (as for aquation), but the spread of kOH values is much greater (a factor of ca. 100 as against ca. 10 for aquation). This is reasonable on the basis of an associative mechanism, since with decreasing basicity of the ligand (i.e. decreasing electron-­donating character of R3CCO2−, a minimum for F3CCO2− due to the strong electron-­withdrawing character of fluorine in the CF3 group), the attack by an incoming anionic nucleophile such as OH− will be facilitated more than that by a polar non-­ionic nucleophile, such as H2O:

For the aquation of trans-­[CoX(en)2(NH3)]2+, the activation energy increases in the sequence Cl− < Br− < NO3− (for X−), and the rate increases in the same sequence due to the increasingly favourable entropy of activation. To account for this, a solvent-­assisted dissociation (SAD) mechanism was suggested.31 Interaction of the solvent by hydrogen bonding with the outgoing group in the transition state (see Figure 4.1) decreases strongly in the sequence Cl− > Br− > NO3−, and this lowers the activation energy of dissociation and more markedly the entropy of activation in the reverse order. The order Cl− > Br− follows from the fact that in protic solvents chloride is more strongly solvated than bromide, while the difference between chloride and nitrate is due to the fact that in the latter the oxygen atoms of the nitrate remain hydrogen bonded to the >NH of the amine ligands (see Figure 4.2) and hence Table 4.12  Aquation  and base hydrolysis rate constants of some [CoX(NH3)5]2+

complexes (to convert all the k values into s−1 it is necessary to divide the k values in min−1 by 60).30

X−

Ka for HX

kH2 O /min−1 at 70 °C

kOH /M−1 min−1 at 25 °C

CF3CO2− CCl3CO2− CHCl2CO2− CH2ClCO2− CH3CO2− CH3CH2CO2− (CH3)2CHCO2− (CH3)3CCO2−

5 × 10−1 2 × 10−1 5 × 10−2 1.4 × 10−3 1.8 × 10−5 1.5 × 10−5 1.5 × 10−5 1.0 × 10−5

3.3 × 10−3 3.2 × 10−3 9.6 × 10−4 3.5 × 10−4 4.9 × 10−4 1.9 × 10−4 1.6 × 10−4 2.6 × 10−4

4.4 4.3 1.6 2.5 × 10−1 4.2 × 10−2 2.7 × 10−2 3.4 × 10−2 1.8 × 10−2

Ligand Replacement Reactions of Octahedral Complexes

205

Figure 4.1  Solvent-­  assisted dissociation in the aquation of trans-­[CoCl(en)2(NH3)]2+.

Figure 4.2  Intramolecular  hydrogen bonding in trans-­[Co(NO3)(en)2(NH3)]2+. interaction with solvent is insignificant until complete dissociation of the nitrate from the complex occurs. The value of K (equilibrium constant) for the aquation increases in the same order of the ligands and the higher value for the nitrato complex is due to greater solvation of the nitrate ion, once it has been released from the complex. The case of [CoX(NH3)5]2+ is closely parallel. The basic feature of the SAD mechanism is that the dissociation of the Co–X bond is assisted by the solvent, which is hydrogen bonded to the departing ligand, and in the transition state, as the Co–X bond is elongated and finally breaks off, H2O slips into the position vacated by the departed ligand. Deuterium substitution, both in the complex and in the solvent, slows the rate of aquation of [CoCl(NH3)5]2+ and the results are said to be in agreement with the SAD mechanism.32 However, the factor by which the rate is decreased on changing from H2O to D2O is virtually the same for cis-­and trans-­[CoCl(L)(en)2]2+ (L = NO2−, Cl−), i.e. it is independent of the position of L and also its nature, which, as we shall soon see, may determine the mechanism. Hence the rate reductions are not diagnostic of the mechanism; however, in any case, the reduction in rate indicates solvent participation in some fashion. Further support for this SAD mechanism comes from observations33 on the effect of D2O on the rate of solvolysis in [CrX(NH3)5]2+ (see Table 4.13). For such complexes the variation in the rate of hydrolysis is primarily due to the ∆S‡ term, since ∆H‡ values are almost the same for X− = C1−, Br− and I−. This means that the energy expended in bond rupture must be at least partially offset by a compensating solvation of the incipient halide ion, i.e. the solvent may be considered as competing with the metal ion for the halide. On changing from H2O to D2O, there is a slight increase in ∆H‡ whereas there is a much greater effect on ∆S‡. Chloride, the most solvated halide in the series Cl−, Br− and I−, is released more slowly in D2O compared with H2O than is iodide, which is

Chapter 4

206

Table 4.13  Solvolysis  of [CrX(NH3)5] at 25 °C. 2+

33

∆H‡/kcal mol−1 X Cl Br I

kH2 O /min−1a −4

5.6 × 10 6.2 × 10−3 6.0 × 10−2

∆S‡/eub

H2O

D 2O

H2O

D 2O

kD2 O /kH2 O

22.4 21.5 21.4

23.0 22.9 23.0

−9.0 −6.1 −1.9

−8.2 −2.3 4.2

0.60 0.95 1.51

 ivide by 60 to obtain values in s−1. D 1 eu = 1 cal K−1 mol−1.

a b

the least solvated halide.34 Since ∆S‡ for the reaction parallels the entropy of solvation of the halide being released, it is evident that solvation of the leaving halide is important in the transition state, suggesting the SAD mechanism. Observations on the acceleration of the rate of aquation of [CrX(NH3)5]2+ due to ion-­pair formation35 by various anions (halides, acetate, benzoate, succinate, oxalate, sulfate, etc.), polyvalent anions producing a more marked acceleration than monovalent anions, and dependence of the effect of such ions on the total ionic strength (an increase in ionic strength retards the reaction) led Jones et al.33 to conclude there is a dissociative activation, viz. a solvent-­assisted dissociation of the ion pair formed between the cationic complex and the accelerating anion is involved in the rate-­determining step. As a result of an increase in the electron density in Cr(iii) in the ion pair, dissociation of the Cr–X bond is facilitated, while simultaneously the hydrogen-­ bonded H2O slips into the position vacated by X−. Such ion-­pair formation is expected to retard the rate of an associative process.

4.1.4  Electronic Effects of Spectator (Non-­leaving) Ligands Apart from steric effects, the non-­leaving ligands can influence the rate of aquation by electronic effects also. Data on the aquation of the cis and trans isomers of [CoCl(L)(en)2]n+ complexes (see Table 4.14) are in accord with the dissociative mechanism.36,37 The results in Table 4.14 indicate that electron donors facilitate a dissociative reaction, whereas electron-­withdrawing groups facilitate an associative reaction (as discussed below):

Also, π-­donor ligands such as Cl− and OH− accelerate the loss of Cl− in the cis position compared with that in the trans position; with H2O (a weak π-­donor) and NH3 (neither a π-­donor nor π-­acceptor) the cis and trans isomers aquate at comparable rates, but with a strong π-­withdrawing ligand such as CN− the trans isomer reacts much faster than the cis isomer. All of these observations can be nicely accounted for on the basis of a dissociative mechanism.38 Thus, when the leaving group (Cl−) leaves the coordination sphere of Co(iii) it leaves

Ligand Replacement Reactions of Octahedral Complexes

207

Table 4.14  Rate  constants (k) for aquation reactions of cis-­ and trans-­[CoCl(L) (en)2] and accompanying stereochemical changes in the product [Co(L) (en)2(OH2)] at 25 °C.36,37 cis-­[CoCl(L)(en)2]

trans-­[CoCl(L)(en)2]

[Co(L)(en)2(OH2)] L H2O OH Cl Br NCS N3 NH3 NO2 CN

4

−1

10 kaq/s

0.016 120 2.4 1.4 0.11 0.033 0.005 0.018 0.0062

[Co(L)(en)2(OH2)] 4

−1

% cis

% trans

10 kaq/s

% cis

% trans

— 84 76 >95 100 86 ∼100 ∼100 ∼100

— 16 24 SCN− > N3− > Br− > Cl−), a changeover from predominantly D-­IP to D-­CB occurring between SCN− (N-­bonded) and N3−. The pseudo-­first-­order rate constant for the base hydrolysis of trans-­ [RhCl2(en)2]+ and the bromo and iodo analogues varies with [OH−] as follows: kobs = k1 + k2[OH−] Values of k1 and k2 and the corresponding activation parameters and also for the independently determined kaq {rate constant for aquation of trans-­ [RhX2(en)2]+}166 are given in Table 4.38.

Chapter 4

244

Table 4.38  Base  hydrolysis of trans-­[RhX2(en)2] . +

Kinetic parameter 6

−1

10 k1/s ∆H‡ (k1)/kJ mol−1 ∆H‡ (kaq)/kJ mol−1 ∆S‡ (k1)/kJ−1 mol−1 ∆S‡ (kaq)/kJ−1 mol−1 106k2/dm3 mol−1 s−1 ∆H‡ (k2)/kJ mol−1 ∆S‡ (k2)/kJ−1 mol−1

X = Cl

X = Br a

5.33 ± 0.21 109 ± 2.6 103 ± 1 −21.5 ± 7.7 −38 ± 4 3.85 ± 0.43a 146.7 ± 5.9 +87.9 ± 17.3

X=I b

17.9 ± 0.6 107.9 ± 1.4 105 ± 1 −13.4 ± 4 −21 ± 4 16.8 ± 0.9b 140 ± 2.7 84.8 ± 7.7

384 ± 6c 102 ± 0.5 105 ± 1 −5.6 ± 1.6 +4 ± 4 181 ± 11c 130.5 ± 3 +73.9 ± 9

a

 t 60.4 °C. A At 60.6 °C. At 59.8 °C.

b c

The near identity of the ∆H‡ and roughly of the ∆S‡ values for k1 and kaq indicate that the k1 path is identical with that of aquation. The k2 path makes a contribution to kobs that is less than that of k1 (k2/k1 ≈ 0.5–0.9 dm3 mol−1). The reactions lead to extensive stereochemical rearrangement (100% trans → cis for X = C1, Br; ca. 50% for X = I) and are likely to proceed by a sort of SN1CB mechanism as values of ∆H‡ and ∆S‡ (corresponding to k2) lie at the upper end of an isokinetic plot that includes data for trans-­[Rh(OH)X(en)2]+; reactions further down the isokinetic plot are postulated to have a mechanism closer to the SN1IP end of the SN1CB–SN1IP mechanistic profile.167 The base hydrolysis of [Rh(NCS)(NH3)5]2+has also been studied.167 On consideration of LFAE, both the displacement mechanism involving a trapezoidal octahedral intermediate (LFAE = 3.63Dq) and the dissociation mechanism (SN1IP or SN1CB) involving a square-­pyramidal intermediate (LFAE = 4Dq) appear energetically feasible for the OH-­dependent base hydrolysis process. For a dissociative process, ∆H‡ values may be expected to follow the sequence of Dq values of the [RhX(NH3)5]2+complexes. However, ∆H‡ decreases linearly with increase in Dq in the sequence I− < Br− < Cl−, whereas for the azido and isothiocyanato complexes ∆H‡ increases with increase in Dq in the sequence N3− < NCS−. Hence the trend in ∆H‡ disfavours a simple dissociation mechanism. However, the ∆S‡ value decreases in order of increasing solvation of X− in protic solvents, SCN− < I− < N3− < Br− < Cl−, and this supports a solvent-­assisted dissociation mechanism for reactions of the ion pair [RhX(NH3)5]2+·OH− in which rupture of the Rh–X bond is assisted by solvation of the X− being released. Base hydrolysis of [M(NCS)(NH3)5]2+ (M = Co, Rh, Cr) in water–ethanol and water–acetone mixtures has been investigated153 and the results suggested significant ion-­pair formation. The dependence of the rate constant on the dielectric constant of the medium and other well-­known empirical solvent parameters, such as the Grunwald–Winstein parameter, Y, the Kosower parameter, Z, and the Dimroth parameter, ET, which express the solvating and ionizing power of the solvent, indicates significant associative character in all cases. However, there appears to be more associative character in Co(iii), somewhat less in Rh(iii) and even less in Cr(iii), and this is in agreement

Ligand Replacement Reactions of Octahedral Complexes

245

with earlier observations in aqueous solutions discussed above. The linear dependence of the activation parameters, ∆H‡ and ∆S‡, on different solvent parameters has been interpreted as being due to a solvent-­assisted dissociation of the M–NCS bond with synchronous bond formation to the metal by the entering nucleophile. ∆H‡ decreases and the rate increases with increase in concentration of ethanol in water–ethanol mixtures, despite the fact that ∆S‡ becomes more negative with increasing ethanol concentration. This seems to rule out H2O as the entering nucleophile in the rate-­determining step; instead, OH− appears likely, although on electrostatic considerations attack by H2O was considered more probable earlier,167 as in a reaction through a trapezoidal octahedral intermediate the X− and the entering nucleophile (H2O or OH−) will be in close proximity. The view that OH− is the attacking nucleophile is also in agreement with large positive slopes for the log k versus 1/κ plots (where κ is the dielectric constant for the solvent mixture) for all three complexes.153 Base hydrolysis of [Co(S2O3)(NH3)5]+ in ethanol–water and acetone–water follows both alkali-­dependent and alkali-­independent paths.168 Values of the rate constants for each of these paths and their corresponding activation parameters are markedly dependent on solvent polarity; their correlation with different solvent parameters suggests associative activation, possibly Ia (cf. conclusions from earlier observations in aqueous solution).81,82 An isokinetic relationship was observed168 between the ∆H‡ and corresponding ∆S‡ values for the different paths, for base hydrolysis in water and also in mixed solvents, and also for replacement of S2O32− in the complex with Cl− and NH3 in aqueous solutions,81 suggesting that similar mechanisms are operative. A significant contribution in the field was the proposal for a concentrated E2 mechanism169 for the base hydrolysis of cis-­[CoX2(cyclam)]+ (X = Cl or Br; cyclam = 1,4,8,11-­tetraazacyclotetradecane), in which cleavages of the N–H and Co–X bonds occur synchronously, to give a five-­coordinate intermediate without the intervention of the six-­coordinate conjugate base, and this mechanism is consistent with general base catalysis as follows:

The ratio kOHBr/kOHCl = 9 at 25 °C, showing a considerable dependence of the rate of base hydrolysis on the nature of the two analogous leaving groups, which is inconsistent with an SN1CB mechanism, but consistent with an E2 mechanism. However, there has also been criticism of this mechanism.137b For the base hydrolysis of [Cr(ONO)(NH3)5]2+, evidence was presented153 that the rate is independent of [OH−] and the reaction occurs by a pseudo-­ substitution process, involving rupture of the O–N and not the Cr–O bond:

Chapter 4

246

but may also occur as follows:

Incidentally, aquations of the nitrito complexes [Cr(ONO)(NH3)5]2+ and [Cr(ONO)(OH2)5]2+ are unexpectedly fast and take place by O–N bond fission and not CrIII–O bond rupture.170 Anionic complexes such as [Co(CN)5X]3− (X− = Cl−, Br−, I−) (Adamson and Basolo171), NCS−, N3− (Haim and Wilmarth172), [Co(CN)5(S2O3)]4− (Banerjea and Das Gupta173), [Fe(CN)5(SO3)]5− and [Fe(CN)5(NH3)]3− (Le Gross174) all undergo base hydrolysis by an OH−-­independent path. An SN1 dissociation mechanism was suggested in the case of [Co(CN)5(S2O3)]4− on the basis of experimental evidence.173 The same is possibly also true in the other cases above. The overall negative charge on the complex is obviously expected to favour M–X bond dissociation. Activation volumes, ∆V‡, for the anation of [Co(CN)5(OH2)]2− and aquation of [Co(CN)5X]3− (X = Cl, Br or I) are all positive (ca. +8–14 cm3 mol−1 at 40 °C), which is in keeping with a dissociative mechanism for ligand replacement reactions of such complexes having overall negative charge.175

4.3  Ligand Replacement Reactions of [M(CO)6] It was mentioned in Chapter 3 (Section 3.6) that for the ligand replacement reactions of M(CO)6 (M = Cr, Mo and W), the reactivity order is Cr < Mo > W, for which an explanation was offered. Basolo and co-­workers176 studied CO substitution in the 17e V(CO)6 with phosphines (PR3) and phosphites (P(OR)3): V(CO)6 + L → [V(CO)5L] + CO The rate law is second order, first order with respect to each of the reactants, and there is a perceptible dependence of the rate on the nature of L. The results suggest an associative process. The relatively low ∆H‡ and perceptibly negative ∆S‡ are also consistent with an associative process (Table 4.39). The lability of V(CO)6 is remarkable compared with that of 18e carbonyls. For example, the associative replacement of CO with PBun3 is ca. 1010 times faster in V(CO)6 than that for the second-­order (interchange) pathway in Cr(CO)6, and V(CO)6− is essentially inert to thermal ligand substitution. Hard bases also lead to associative ligand replacement177 in V(CO)6, but the product V(CO)5L cannot be characterized owing to its fast disproportionation into [V(L)6][V(CO)6]2 and V(CO)6. However, the kinetic data suggest that replacement of a CO with L is the rate-­determining step in these cases also. The observed nucleophilicity order (in dichloromethane) is L = py > Et3N > MeCN > MeOH > acetone > THF > 2,5-­Me2THF > DMF > MeNO2 > Et2O.

Ligand Replacement Reactions of Octahedral Complexes

247

Table 4.39  Kinetic  parameters for the reaction of V(CO)6 with L in hexane. L

k/M−1 s−1a

∆H‡/kcal mol−1

∆S‡/cal K−1 mol−1

PMe3 PBun3 PMePh2 P(OPri)3 P(OMe)3 PPh3 PPri3 AsPh3

132 50.2 3.99 0.94 0.70 0.25 0.11 0.018

7.6 ± 0.7 7.6 ± 0.4 8.9 ± 0.3 — 10.9 ± 0.2 10 ± 0.4 — —

−23 ± 3 −25 ± 2 −26 ± 1 — −23 ± 1 −28 ± 2 — —

a

k values at 25 °C.176b

Table 4.40  Activation  volumes for reactions of some hexacarbonyls. M(CO)6

PR3

Kinetic order

Solvent

∆V‡/cm3 mol−1

Cr(CO)6 Mo(CO)6 W(CO)6

PPh3 PPh3 PBu3

First First Second

Cyclohexane Isooctane Cyclohexane

+15 ± 1 +10 ± 1 −10 ± 1

The activation volumes ∆V‡ for replacement of a CO with PR3 in M(CO)6 in non-­polar inert solvents have been reported178 (Table 4.40). Owing to the nature of the solvents used the electrostriction effect will be insignificant, hence the sign and magnitude of the ∆V‡ values indicate (see Chapter 2, Section 2.4.2) an associative reaction for W(CO)6 but a dissociative reaction for the Cr and Mo hexacarbonyls. For the reaction Cr(CO)6 + N3− → [Cr(CO)5(NCO)]− + N2 the ∆V‡ value is ∼0. This was interpreted in terms of the transition-­state structure shown below to which bond making and bond breaking contribute almost equally:178

4.4  Reactions of s-­ and p-­Block Metals Beryllium forms tetrahedral complexes of Be2+ by reactions that are discussed in Chapter 3. The formation of complexes by several s-­and p-­block metal ions are discussed in Chapter 5. Isotopic exchange and racemization studies of tris-­β-­diketonate complexes of Al3+, Ga3+ and In3+ and also of Si(iv) and Ge(iv) have been reviewed.179 From studies on the ligand exchange of Ga(acac)3 with labelled Hacac in THF + H2O, the following rate law was reported: Rate = [Ga(acac)3](k1 + k2[H2O] + k3[Hacac])

248

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A dissociative mechanism involving an intermediate containing a unidentate ligand was proposed.180 Comparison with similar exchange reactions of other actylacetonato complexes led to the following reactivity order: Ti(iv) > In(iii) > Ga(iii) > Al(iii) > Be(ii) > Pd(ii) > Co(iii) > Si(iv) > Ge(iv) > Cr(iii). This sequence should be interpreted with caution because the coordination number is not the same in all cases, and it is not possible to interpret the exchange reaction by a single mechanism in all cases, but the lability order Al(iii) < Ga(iii) < In(iii) suggests a dissociative mechanism in these instances. The lability order Cr(iii) < Co(iii) is also surprising.

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254

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158. G. C. Lalor and T. Carrington, J. Chem. Soc. A, 1969, 2509. 159. (a) T. P. Jones and J. K. Phillips, J. Chem. Soc. A, 1971, 1881; (b) T. P. Jones and J. K. Phillips, J. Chem. Soc. A, 1968, 674. 160. (a) M. A. Levine, T. P. Jones, W. E. Harris and W. J. Wallace, J. Am. Chem. Soc., 1961, 83, 2453; (b) D. L. Gay and G. C. Lalor, J. Chem. Soc. A, 1966, 1179. 161. A. W. Adanson and R. G. Wilkins, J. Am. Chem. Soc., 1954, 76, 3379. 162. E. Jorgensen and J. Bjerrum, Acta Chem. Scand., 1958, 12, 1047. 163. (a) G. Guastalla and T. W. Swaddle, J. Chem. Soc., Chem. Commun., 1973, 61; (b) E. Zinato, C. Furlani, G. Lanna and P. Riccieri, Inorg. Chem., 1972, 11, 1746; (c) D. A. House, Acta Chem. Scand., 1972, 26, 2847. 164. C. Chatterjee and P. Chaudhuri, Indian J. Chem., 1971, 9, 1132. 165. S. C. Chan, J. Chem. Soc., 1965, 3207, and references cited therein. 166. A. J. Poe and C. Vuik, Can. J. Chem., 1975, 53, 1842; J. Chem. Soc., Dalton, 1976, 661. 167. C. Chatterjee, P. Chaudhuri and D. Banerjea, Indian J. Chem., 1970, 8, 1123. 168. M. N. Bishnu, R. N. Banerjee and D. Banerjea, Indian J. Chem., 1983, 22A, 948. 169. R. W. Hay and P. R. Norman, J. Chem. Soc., Chem. Commun., 1980, 734. 170. T. C. Mattts and P. Moore, J. Chem. Soc. A, 1960, 219, 1977. 171. A. W. Adamson and F. Basolo, Acta Chem. Scand., 1955, 9, 1261. 172. A. Haim and W. K. Wilmarth, Inorg. Chem., 1962, 1, 573. 173. D. Banerjea and T. P. Das Gupta, J. Inorg. Nucl. Chem., 1967, 29, 1021. 174. J. Le Gross, Compt. Rend., 1956, 242, 1605. 175. D. A. Palmer and H. Kelm, Z. Anorg. Allg. Chem., 1979, 450, 50. 176. (a) Q. Shi, T. G. Richmond, W. C. Trogler and F. Basolo, J. Am. Chem. Soc., 1982, 104, 4032; (b) Q. Z. Shi, T. G. Richmond, W. C. Trogler and F. Basolo, J. Am. Chem. Soc., 1984, 106, 71. 177. T. G. Richmond, Q. I. Shi, W. C. Trogler and F. Basolo, (a) J. Chem. Soc., Chem. Commun., 1983, 650; (b) J. Am. Chem. Soc., 1984, 106, 76. 178. R. Brower and T.-­S. Chen, Inorg. Chem., 1973, 12, 2198. 179. K. Saito, Pure Appl. Chem., 1974, 38, 325. 180. C. Chatterjee, K. Matsuzawa, H. Kido and K. Saito, Bull. Chem. Soc. Jpn., 1974, 47, 2809. 181. E. Grunwald and S. Winstein, J. Am. Chem. Soc., 1948, 70, 846.

Chapter 5

Catalysed Reactions and Formation Reactions 5.1  Electrophilic and Nucleophilic Catalysis Electrophilic and nucleophilic catalysis is well known1 and is interpreted by the following general scheme:   



kA k   SA  S  A   products 

(5.1)

  

where S = substrate complex and A = electrophile or nucleophile. Based on this,   



kobs = kKA[A]/(1 + KA[A])

(5.2a)

  

If KA is small enough such that KA[A] 10.5 (see Chapter 4, Section 4.2).2b Acid-­ catalysed aquation of nitro complexes at high acid concentrations have been reported.3 Acid-­catalysed aquation of [Cr(NH3)5(N3)]2+ has been studied.4a The reaction proceeds by concurrent acid-­independent and acid-­dependent paths:   



kobs = k0 + kH[H+]

(5.4)

  

where k0 and kH are the specific rate constants for the acid-­independent and acid-­dependent paths, corresponding to processes involving rupture of the Cr–N and Cr–N3H bonds, respectively. However, in the absence of acid no reaction is observed owing to the overwhelming back-­reaction; but in presence of acid the back-­reaction is prevented by protonation of the free N3−. Hence k0 corresponds to an acid-­assisted process whereas kH corresponds to an acid-­catalysed process. This is the most common type of behaviour that has been observed in many such reactions. The magnitude of the ΔS‡ value (−4.9 eu) for the acid-­independent path for the azidopentaammine complex of Cr(iii) suggests4a an associative mechanism with significant bond formation by the entering nucleophile H2O, since it is known that fixing a molecule of H2O in the transition state should lead to an entropy loss of 3–6 eu (see Chapter 2). On the other hand, if rupture of the Cr(iii)–N3 bond is primarily important in the transition state, we should expect a positive value for ΔS‡ owing to the gain in the rotational entropy of the unbound azide (+11 eu) minus the entropy loss due to solvation of N3− (−7.5 eu in aqueous solution). For the acid-­catalysed path, if we assume a similar mechanism involving the conjugate acid, we should expect a positive ΔS‡ corresponding to kH due to the contribution from the desolvation of the

Catalysed Reactions and Formation Reactions

257

proton in the conjugate acid formation pre-­equilibrium. However, even on simple electrostatic considerations, it should be relatively much easier energetically for the Cr(iii)–N3H bond to dissociate compared with the Cr(iii)–N3 bond, hence it is likely that transformation of the conjugate acid involves significant dissociation of the Cr(iii)–N3H bond, concurrent with bond formation by the incoming H2O, as in an Ia process. On the basis of this mechanism, the positive contribution to ΔS‡ due to desolvation of H+aq may be more than compensated by the negative contributions from the solvation of the released HN3, and the fixing of the water molecule. Hence the observed ΔS‡ value of −5.8 eu in this case is in agreement with an associative interchange process for transformation of the protonated complex. Aquation of trans-­[CoF2(en)2]+ proceeds similarly by concurrent acid-­ independent and acid-­dependent paths.4b However, aquation of cis-­ [Cr(CN)2(en)2]+ occurs exclusively by the acid-­dependent path4c with   

  

kobs = kKH[H+] = kH[H+]

(5.5)

However aquation of [Cr(CN)(NH3)5]2+ in an acidic medium occurs by both acid-­independent and acid-­dependent paths, the latter being dominant.4d It is worth mentioning that for aquation of complexes of the type [Cr(OH2)5L]3+ the k0 path may involve loss of LH+ (rather than L), if the pKa of LH+ is greater than 2,4e forming a transient intermediate [Cr(OH)(OH2)4]2+, which is transformed into the product [Cr(OH2)6]3+ by reacting with H3O+. Aquations of [Cr(ONO)(NH3)5]2+ (ref. 4f) and [Cr(ONO)(OH2)5]2+ (ref. 4g) are also acid catalysed, but the reactions occur by N–O (not Cr–O) bond rupture (pseudo-­substitution process, see Chapter 1) and in the latter case there is an additional path that is dependent on [H+]2 Work on the stepwise dissociation (X− loss) of [Cr(atda)X2]2− {atda = o-­C6H4(CO2)[N(CH2CO2)2]3−; X = NCS, N3} in acidic solution were carried out by Banerjea and co-­workers. In both cases k0 and kH[H+] paths were observed. In the case of the azido complex the intermediate could be isolated and its dissociation studied. Based on spectroscopic data, the ligand field activation energy values for dissociative and associative reactions were estimated and on consideration of experimental ΔH‡ values for k0 and kH plausible mechanisms were proposed (unpublished, see ref. 59b, pp. 112–114). Aquation (dissociation) of metal complexes containing a chelating ligand, such as ethylenediamine, bipyridine, etc., having no free basic site in the chelated ligand, are also acid catalysed. This is easily understood in terms of the general scheme for the dissociation of a metal chelate shown in Figure 5.1. Transformation of the intermediate I, in which the bidentate ligand is bound unidentately due to one-­ended dissociation of the chelate ring, to the product is facilitated both by electrophiles (E+) and by nucleophiles (Nu−), which prevent back-­reaction by combining with the free end of the ligand or by occupying the vacated coordination position of the metal, respectively. Further, electron withdrawal by E+ weakens the M–AA bond and electron donation by Nu− causes a similar effect by decreasing the electron

258

Chapter 5

Figure 5.1  General  mechanism of dissociation of a metal chelate. E+ = electrophile (Lewis acid such as H+ or a metal ion); Nu− = nucleophile (Lewis base).

affinity of the metal. Such a one-­ended dissociation, although easily possible with 2,2′-­bipyridine, is not favourable in the case of 1,10-­phenanthroline, where because of three fused rings the structure is rigid and no free rotation about the C–C axis joining the two pyridine rings (as is possible in the case of 2,2′-­bipyridine) is possible here. Hence, whereas the dissociation of [Fe(bipy)3]2+ occurs partly by an acid-­catalysed path,5 this path is absent6 in the case of [Fe(phen)3]2+. However, ligands such as biguanide, oxalate, malonate, acetylacetonate, etc., which have free basic sites even in the chelated form, are able to add H+ (and also Mn+), prior to chelate ring opening, and this facilitates the chelate ring opening and final loss of the ligand as follows, taking the case of oxalate bound to M3+ as an example:

Catalysed Reactions and Formation Reactions

259

The possibility of protonation of the chelated biguanide prior to opening of the chelate ring leads to some apparently paradoxical results. Thus, for both Cr(iii) and Co(iii) and for any particular series (tris or bis) of complexes of the biguanides, the rate of dissociation in acidic solution decreases in the following sequence of decreasing basicity of the ligands, although the stability of the complexes decreases in the same sequence: biguanide > N1-­hexylbiguanide > N1-­phenylbiguanide.7 This is because protonation of the chelated ligand, which facilitates the cleavage of the metal–ligand chelate ring, is expected to be less favoured in the same sequence as the decreasing basicity of the free ligands. Again, although ethylenediamine forms much less (thermodynamically) stable complexes with Cr(iii) and Co(iii) compared with biguanide, the ethylenediamine complexes are much more inert (kinetically) than the corresponding biguanide complexes, because protonation (which facilitates metal–ligand chelate ring opening) cannot occur prior to opening of the chelate ring in the case of ethylenediamine complexes, owing to the absence of any free basic group in the chelated ethylenediamine. Acid-­catalysed dissociation of complexes of en involves an uncatalysed chelate ring opening pre-­equilibrium followed by protonation of the free –NH2 (cf. Figure 5.1). Thus, whereas for [Cr(bigH)3]3+ the ΔH‡ value is 10.1 kcal mol−1, the corresponding value for [Cr(en)3]3+ is 24.3 kcal mol−1 and [Co(bigH)3]3+ (bigH = biguanide) suffers dissociation in an acidic medium at a measurable rate at room temperature (ca. 30 °C), whereas [Co(en)3]3+ remains completely unchanged in 1 M HClO4, even for months at 30 °C.7 Dissociation of the bis-­complexes of di(2-­pyridyl)amine formed by Cu(ii), Ni(ii) and Co(ii) in aqueous acidic media proceeds in two steps through formation of the mono-­complexes, and for each of the steps the observed rate constant shows an acid-­independent and an acid-­dependent path for which a plausible reaction scheme has been proposed, which involves chelate ring opening as the rate-­determining step for both paths.8a Dissociation of [M(trien)]2+ (M = Cu, Ni; trien = triethylenetetramine) has been reported.8b There is evidence9 that in aqueous solution the Ni(ii) complex is six-­coordinated cis-­[Ni(OH2)2(trien)]2+; [Cu(trien)2+] is likely to be similar. A reaction scheme has been proposed for their dissociation that is consistent with the rate law derived from experimental results on the dissociation of [Cu(trien)2+] and the kobs value under pseudo-­first-­order conditions:   

  

kobs = k0 + kH′[H+] + k″H[H+]2

(5.6)

where k0, kH′ and kH″ are composite constants, being products of some equilibrium constants and a rate constant. The case for [Ni(trien)2+] is similar but here kH″ = 0. At 25 °C (I = 0.1 M) the values of kobs (experimental) in 0.05 M HClO4 are 228 s−1 (Cu) and 3.0 s−1 (Ni). The reported values of k0, kH′ and kH″ at 25 °C (I = 0.1 M) and corresponding ΔH‡ and ΔS‡ values are given in Table 5.1.

Chapter 5

260

Table 5.1  Dissociation  of [M(trien)(OH2)2] (M = Cu, Ni). 2+

Cu −1

ΔH‡/kJ mol−1 ΔS‡/J K−1mol−1

−1

−1

Ni −2

−1

−1

k0/s

kH′/M s

kH″/M s

k0/s

3.0 69.3 −4.6

2.57 × 102 22.1 −125.7

8.3 × 104 16.4 −96.5

1.20 39.5 −112.5

kH′/M−1 s−1 37 30.8 −112.6

For both complexes the k values are low, in addition to the kCu/kNi ratio, 2.5 for k0 and 6.95 for kH′; for [M(NH3)(OH2)5]2+ (M = Cu, Ni) at 25 °C the ratio is ca. 3.5 × 103 and for [M(en)(OH2)4]2+ at 25 °C the ratio (corresponding to k0)10 is ca. 7.7 × 102. The enhanced chelate effect of quadridentate trien causes a dramatic decrease in individual rate constants of both the Cu(ii) and Ni(ii) complexes and this is more pronounced in the case of Cu(ii) than Ni(ii), causing decrease in the rate ratios. A large volume of work has been carried out by Banerjea and co-­workers on the acid-­catalysed dissociation of metal complexes, such as [Co(NO2) (NH3)5]2+,11 several oxalato,12,13 malonato13,14 and acetylacetonato15 complexes of Cr(iii), several biguanide and N1-­substituted biguanide complexes of Cu(ii),16 Ni(ii),17,18 Cr(iii),7 Co(iii)7 and Rh(iii),19 bis(biguanide)oxalato complex of Cr(iii),20 [Co(ox)(phen)2]+,21a and [Co(mal)(phen)2]+,21b ethylenedibiguanide complexes of Cu(ii) and Ni(ii)17a and m-­phenylenedibiguanide complex of Cu(ii),17b M(baen) [M = Cu, Ni; baen = bis(acetylacetonato)ethylenediamine(2–)],22a Cu(ii) complexes of N-­salicylideneglycinate and N-­salicyliden eglycylglycinate17b and some complexes of U(vi).22b Such kinetic information has also made possible the synthesis of some complexes in the senior author's laboratory, such as [PdCl2(bigH)],24 [Cr(ox) (bigH)2]2ox·2H2O and [Cr(ox)(bigH)2]ClO4.20 Redox decomposition of 1,2,6-­[CoNO2)3L3] (L3 = 3NH3, dien)25 and [Co (acac)3]26 in solution are also acid catalysed; but acid retards (inverse dependence of rate on [H+], negative catalysis) the redox decomposition of [Ag(endibigH2)]3+.27 Base hydrolysis of [Cr(big)3] formed from [Cr(bigH)3]3+ in alkaline solution shows7 an inverse dependence on [OH−]. Acid (H+)-­catalysed dissociation of [Co(NO2)(NH3)5]2+ in strongly acidic media forms [Co(NH3)5(OH2)]3+,28 but in HOAc–NaOAc buffer slow H+-­catalysed redox decomposition takes place.11 In HOAc–NaOAc buffer, [Cr(NH3)5(ONO)]2+ forms [Cr(OAc)(NH3)5]2+ by an H+-­catalysed path.29 Dissociation of a chelated bidentate ligand is kinetically a two-­step process, viz., chelate ring opening then dissociation of the unidentately bound bidentate ligand, and for elucidation of the mechanism it is necessary to know which of these steps is rate determining, and this may sometimes be possible by a suitable approach as in the following example. The bis(malonato) complex [Co(mal)2(en)]− aquates in acidic solution forming [Co(mal)(en)(OH2)2]+ exclusively by an acid-­catalysed path, then the

Catalysed Reactions and Formation Reactions

261

loss of the malonate from the diaqua product occurs in concurrent acid-­ independent and acid-­dependent paths.30 The ΔH‡ versus ΔS‡ isokinetic plot for the acid-­catalysed dissociation of L from several Co(iii) complexes, [Co(L)(NH3)5]n+, where L is a monocarboxylate (n = 2) or a unidentatedly bonded dicarboxylate (n = 1), is linear and data for [Rh(ox)3]3−, [Co(ox)(phen)2]+ and [Co(mal)(phen)2]+ also fit this line (Figure 5.2). This suggests that in the dissociation of the latter three complexes the chelate ring opening is fast followed by rate-­determining dissociation of the now unidentatedly bonded dicarboxylate. However, the corresponding isokinetic plot for the acid-­catalysed dissociation (loss of dicarboxylate) of [Co(mal)2(en)]− and [Co(mal)(en)(OH2)2]+ and also of [Co(ox)3]3−, [Co(mal)3]3− and [Co(ox)2(OH2)2]− is another straight line (of much larger slope) (Figure 5.2), hence it has been reasonably concluded that in these cases the chelate ring opening is the rate-­determining step followed by a fast dissociation of the unidentendately bonded dicarboxylate.30 In the dissociation of [M(Me6tren)(OH2)]2+ [trigonal bipyramidal (tbp) structure] (M = Cu, Ni, Co) in acidic solution, only the first M–NMe2 bond dissociation is rate determining, with all the subsequent steps leading to loss of Me6tren being fast.31 At 25 °C, the observed lability order is Cu2+ > Ni2+ > Co2+.

Figure 5.2  Isokinetic  plots for the acid-­catalysed dissociation path (see text):30

(1) [Co(ox)(phen)2]+; (2) [Rh(ox)3]3−; (3) [Co(ox)(NH3)5]+; (4) [Co(phthalate)(NH3)5]+; (5) [Co(succinate)(NH3)5]+; (6) [Co(o-­MebzH)(NH3)5]+; (7) [Co(mal)(NH3)5]+; (8) [Co(salH)(NH3)5]+; (9) [Co(glyH)(NH3)5]3+; (10) [Co(phen)2]+; (11) [Co(mal)3]3−; (12a) [Co(ox)3]3−; (12b) [Co(ox)3]3−; (13) [Co(mal)2en]−; (14) [Co(mal)(en)(OH2)2]+; (15) [Co(ox)2(OH2)2]−; (16) [Co(O2CCH3)(NH3)5]2+; (17) [Co(NH3)5(O2CCH2Cl)]2+; (18) [Co(O2CCHCl2) (NH3)5]2+; (19) [Co(O2CCMe3)(NH3)5]2+.

Chapter 5

262 2+

The dissociation of [M(bipy)3] [M = Fe(ii), Ni(ii)] shows an acid dependence, as indicated in Scheme 5.1, taking [Fe(bipy)3]2+ as an example.5 Applying the usual steady-­state approximation for the concentration of the species I with unidentately bonded bipyridine, the following expression is obtained for kobs at any acid concentration, which explains the observed acid dependence of the rate of dissociation on acid concentration:   



kobs = {(k3 + KHk4[H+])/(k2 + k3 + KHk4[H+])}k1

(5.7)

  

At high acid concentration a limiting rate is observed, since k2 and k3 will be negligible compared with KHk4[H+] and hence under this condition kobs ≈ k1. At [H+] ≈ 0, we obtain kobs ≈ k1k3/(k2 + k3) = k1′ (say). The value of k1′/k1 = k3/ (k2 + k3) is ca. 0.15 (25 °C) and hence, at very low acidity, each time a single

Scheme 5.1

Catalysed Reactions and Formation Reactions

263

Fe–N bond breaks the bond reforms approximately 85% of the time, and the second bond breaks (leading to complete dissociation) only about 15% of the time. The behaviour of [Ni(bipy)3]2+ is similar. However, the rigidity of the fused-­ring system of 1,10-­phenanthroline prevents similar one-­ended dissociation, hence the dissociation of the tris-­phenanthroline chelates of these metals show no such acid dependence.6 An insight into the mechanism of transformation of the conjugate acid (protonated form of the complex) into the product (i.e. the role of water in the transition state for the reaction in aqueous acidic solution, which indicates whether the reaction is SN1CA or SN2CA), has been obtained from various evidence in some cases; the observed dependence of the rate on H+ ion concentration and various acidity functions has been particularly useful. The Hammett acidity function,32 H0, expresses the ability of the system to protonate a base B:   

B + H+ ⇌ BH+



(5.8)

  

By definition,   



H0   log aH  log  f BH / f B  +

+

 pK BH+ +log  [B]/[BH+ ] 

(5.9) a where K BH+ is the acid dissociation constant of BH , H+ denotes the activity of H+, f BH+ and fB are the activity coefficients of BH+ and B, respectively, and [ ] denotes molar concentration. f 1, hence under this condition In very dilute aqueous solution, f BH + B   

  



  

+

H0   log aH+  pH

(5.10)

Similarly, the acidity functions H+ and H– can be defined with respect to cationic and anionic bases, respectively, and these are all parallel functions.33 According to Hammett and Zucker,34 if the aquation of a substrate S occurs through the formation of its conjugate acid, SH+:   

  

S + H+ ⇌ SH+

(5.11)

followed by reaction of SH+ with water to form the product, then in the case of a dissociative (SN1) transformation of SH+, the plot of log k versus –H0 should be linear with slope = 1. On the other hand, if SH+ reacts by an associative (SN2) process (with H2O acting as the nucleophile), then the plot of log k versus log [H+] should be linear with slope = 1. This enables us to distinguish between SN1CA and SN2CA mechanisms. Another criterion was suggested by Bunnett,35 which involves plotting (log k + H0) versus log aH2 O (aH2 O being the activity of H2O in the solution) and the slope ω of the linear plot is diagnostic of the mechanism [see (I) in Table 5.2]. Later, Bunnett and Olsen36 suggested plotting (log k + H0) versus (log [H+] + H0) for weakly basic substrates and the slope ϕ of the linear plot is diagnostic of the mechanism [see (II) in Table 5.2]. For strongly basic substrates, the authors suggested plotting log k versus (log [H+] + H0).

Chapter 5

264

Table 5.2  Values  of ω and ϕ in assignment of mechanism (see text for further detail). (I) ω value

Mechanism

≤0 ∼+2 >3.3

SN1CA (i.e. D-­CA) SN2CA(H2O) (i.e. A-­CA) (involving attack by H2O) SN2CA(H3O+) (i.e. A-­CA) (involving attack by H3O+)

(II) ϕ value

Mechanism

Ni2+ > Co2+. This behaviour is in contrast to what is observed in six-­coordinate systems (Cu2+ > Co2+ > Ni2+) and the order of ΔH‡ (values in kJ mol−1) is Cu2+ (81.1) > Co2+ (61.9) > Ni2+ (58.9). If it is assumed that

274

Chapter 5

dissociation of the first amine group proceeds via an approximately tetrahedral transition state, the approximate crystal field activation energy values are 0.10Dq, 2.70Dq and 5.30Dq for high-­spin cobalt(ii), nickel(ii) and copper(ii), respectively. The uncertainties in the calculations make these values somewhat speculative, but it is noteworthy that the highest crystal field activation energy is that of copper(ii), which also exhibits the largest ΔH‡ value. Dissociation of the nickel(ii) and copper(ii) complexes of the macrocyclic ligand [16]-­ane-­N4 (1,3,6,9,11,14-­hexaazacyclohexadecane) in acidic media was reported by Banerjea and co-­workers.23a Whereas dissociation of the copper(ii) complex occurs through concurrent acid-­independent and acid-­ dependent paths, that of the nickel(ii) complex is independent of acid, even in 1 M HClO4. Structural models indicate that just a slight distortion of these square-­planar complexes would cause the ligand to bind in a tridentate manner, giving rise to a trigonal bipyramidal complex with two other ligands in apical positions (these apical ligands may be donor solvent molecules, e.g. H2O). The ligand field stabilization energies (see Chapter 1, Table 1.8) indicate that for a strong-­field ligand such as the one under consideration, d8 nickel(ii) will have little tendency to change from square-­planar to any other geometry. The conditions are more favourable, however, for d9 copper(ii) to assume trigonal bipyramidal geometry (see Chapters 2 and 3). In such a structure, the HN< group, which is freed from metal binding, can accept a proton in a fast pre-­equilibrium which accounts for the observed acid dependence. Since the nickel(ii) complex does not show acid dependence, even at 1 M acid concentration, it appears that the two unbound HN< groups of the macrocycle in the square-­planar complex do not have sufficient residual basicity to add protons. This is possibly due to the involvement of the lone pairs of the N of these two HN< groups in a hydrogen bonding interaction with a neighbouring metal-­bound HN< group forming an HN⋯HN–M hydrogen bond, which appears probable from examination of a structural model (see the structures shown).

Catalysed Reactions and Formation Reactions

275

−1

4

At 50 °C (I = 1 M) the 10 k0 values (s ) are 2.2 (Cu) and 0.8 (Ni), and the 104kH value for the Cu(ii) complex is 2.0 M−1 s−1. For the Ni(ii) complex the ΔH‡ value (kJ mol−1) is 100.4 ± 4.2 (k0) and the corresponding value for the Cu(ii) complex is 94.5 ± 5.5. The ΔS‡ values (J K−1 mol−1) (k0) are −17.8 ± 0.8 (Ni) and −25.3 ± 2.2 (Cu), and the corresponding values for the kH path (Cu) are 58 ± 2.5 and -70 ± 2.5 respectively. The ionic strength has hardly any effect on k0 but kH increases appreciably with increase in ionic strength. The ratio k0Cu/k0Ni is 2.75 (50 °C) and 3.25 (60 °C); such low values have been observed in two other cases of square-­planar complexes of quadridentate ligands.22,23b However, even in the case of [M(trien)(OH2)2]2+ the ratio is similar due to a pronounced chelate effect,8 but in the case of [M(endibigH2)]2+ {where, endibigH2 = ethylenedibiguanide} a significantly higher value has been observed.17a Hay et al.46 reported similar studies on the Cu(ii) complex of 1,5,8,12-­tetraa zacyclooctadecane ([18]-­ane-­N4) and of the Ni(ii) complex of 1,5,8,12-­tetraaza cycloheptadecane ([17]-­ane-­N4). Dissociation of the Cu(ii) complex displayed saturation kinetics beyond 0.4 M HClO4, the kobs value being independent of [HClO4] at concentrations ≥0.4 M. However, the Ni(ii) complex did not display this behaviour, and kobs showed a first-­order dependence on [HClO4] (up to 0.5 M used in the investigation). The attainment of a limiting value of kobs at 0.4 M HClO4 in the case of the Cu(ii) complex is due to following reaction:   



kH

H2 O,k CuL2   H #CuLH3   Cu2aq  LH

(5.21)

  

The estimated value of KH is 64 M−1 (at 25 °C). Hence at 0.4 M HClO4 the [CuL]2+ will be virtually completely transformed to [CuLH]3+, which accounts for the observed saturation kinetics, as in the case of Ni(en)2+ aq mentioned earlier [see Chapter 1, eqn (1.54)]. The behaviour of the Ni(ii) complex suggested that no significant amount of [NiLH]3+ was formed even in 0.5 M HClO4. At 25 °C, dissociation of the Ni(ii) complex has been reported to be 1010-­fold faster than that of [Ni(cyclam)]2+. Acid-­catalysed dissociation of the Cu(ii) and Ni(ii) complexes of the quadridentate macrocyclic ligand 1,5,9,13-­tetraaza-­2,4,4,10,12,12-­hexamethylcy clohexadecane-­1,9-­diene in aqueous HClO4 solution was reported by Banerjea and co-­workers.23b Both complexes dissociate fairly slowly in acidic media ([H+] = 0.1–1.0 M; I = 1.0 M) with the experimental pseudo-­first-­order rate constant, kobs, for the nickel(ii) complex showing acid dependence as follows:   



  

kobs = k0 + kH[H+]

(5.22)

The acid-­independent path (k0) is, however, absent in the case of the copper(ii) complex, which reacts exclusively by the kH path. It should be noted that contrary to the case of the Ni(ii) complex of [16]-­ane-­N4 reported earlier,23a which reacts exclusively by an acid-­independent path, the Ni(ii) complex mentioned above reacts by an acid-­dependent path also. In the cases of these complexes, the protonation occurs at an unsaturated site (C=N bond) in the macrocyclic ring forming the conjugate acid as shown in Scheme 5.7.

Chapter 5

276

Scheme 5.7 At 60 °C (I = 1 M) the kH values for both complexes are ∼10−4 M−1 s−1; kHCu/kHNi = 3.87, comparable to that for [M[16]-­ane-­N4)]2+.23a The difference in rates for the Cu(ii) and Ni(ii) complexes is primarily due to the ΔH‡ values, and the ΔS‡ values are highly negative due to solvation of the M2+ and LH+ (protonated ligand) in the dissociation process. An increase in ionic strength has no effect on k0 whereas kH increases perceptibly.

5.1.2  Electrophilic Catalysis by Metal Ions Metal ions, like other electrophiles, e.g. H+ and NO+, also catalyse reactions of metal complexes. Catalysis by electrophiles, especially labile metal ions, is analogous to acid catalysis and can occur in nearly all cases where acid catalysis is observed. A major difference between catalysis by protons and by metal ions lies in differences in their binding abilities; the HSAB (hard and soft acids and bases) principle works here and can be used as a general guide in correlating metal ion-­assisted and metal ion-­catalysed processes.47a,b A proton can bind readily to class a or “hard” substrates (e.g. F, O, N donors) and because of its small size can penetrate donor sites which large metal ions find sterically hindered. Metal ion catalysts fall into two broad categories: (1) class a or “hard” metal ions (e.g. Be2+, Al3+, Fe3+, Th4+), which bind most readily to N, O and F donor sites and hence catalyse the removal of such ligands; and (2) class b or “soft” metal ions (e.g. Hg2+, Ag+, Tl+, Tl3+), which can catalyse reactions that protons do not accelerate (e.g. aquation of Cl−, Br− or I−, and other class b donor ligands, Hg2+ being most effective). Soft metal ions such as Hg2+ are also effective for pseudo-­halide complexes. Thus, aquation of [CrX(OH2)5]2+ (X− = N3−. SCN−)48b,c are catalysed by Hg2+. Aquation of [CoF2(en)2]+ is catalysed by Be2+ < Al3+ ≈ Sc3+ < Th4+; the rate constant increases with increase in charge on the catalysing metal ion.47c The same trend has been reported47c for the aquation of [CrF4en]−. For the

Catalysed Reactions and Formation Reactions 3+

2+

277 3+

Al -­catalysed aquation of [CoF(NH3)5] , kobs = kAl[A1 ]; at 25 °C and I = 2 M, kAl = 1.6 × 10−3 M−1 s−1, and below pH 3 no dependence of the rate on pH was observed.47d However, the Al3+-­catalysed aquation of [CrF(NH3)5]2+ shows a rate retardation with decrease in pH, which is due to protonation of the bound F− ligand, as is evident from the following expression47e for kobs:   

kobs = k[Al3+]/(1 + KH[H+])



(5.23)

  

At 35 °C and I = 2 M, pKH = 3.0 and kAl = 7.5 × 10−4 M−1 s−1 (at pH 3). The aquation of [Cr(O2CMe)(OH2)5]2+ shows normal acid catalysis (k0 and 2+ kH paths, see Section 5.1.1), but aquation catalysed by Cr2+ aq and V aq has also 47f 2+ been observed. For the Cr aq-­catalysed reaction the rate law is −d[CrIII]/dt = (k1 + k2/[H+])[Cr2+][CrIII] and a similar rate law has been observed at high concentrations of V2+ aq. The path showing inverse first order in [H+] is due to the involvement of [Cr(O2CMe(OH)(OH2)4]+, the OH− bridging the M2+ aq in the activated complex. In reactions catalysed by Hg(ii) in the presence of Cl−, catalysis by HgCl+ and even HgCl2 in addition to Hg2+ is often observed.48a For the aquation (Br− loss) of [CoBr(CN)5]3− catalysed by Hg(ii) halides, the observed rate law leads to49   

kobs = k1 + k2[HgX+] + k3[HgX2]



(5.24)

  

in which k2 is ca. 103k3, with practically no catalytic effect of [HgX3]−. In the presence of Hg2+, [Co(SCN)(NH3)5]2+ suffers concurrent aquation and linkage isomerization to [Co(NCS)(NH3)5]2+ through formation of the adduct [(H3N)5Co–SCN–Hg]4+ (K = 1.2 M−1 at 25 °C); both reactions have nearly the same rate (k = 0.23 s−1 at 25 °C).50a Studies on the two-­step Hg2+-­catalysed aquation (SCN− loss) of [Co(NCS)2 (dmgH)2]− have been reported,50b with values of kinetic parameters as given in Table 5.9. For the uncatalysed aquation in acidic solution, the values for the first step are kaq (25 °C) = 3.8 × 10−8 s−1 estimated using ΔH‡ = 132 kJ mol−1 and ΔS‡ = +54 J K−1 mol−1. This is followed by a much slower reaction. The acceleration by Hg2+ is due to considerable lowering of the enthalpy of activation. Based on all the evidence, a mechanism involving electrophilic attack by Hg2+ on the

Table 5.9  Kinetic  parameters for mercury(ii)-­catalysed dissociation of a bisthiocyanato complex of Co(iii).

Parameter 2+

−1

−1

kHg /M s (at 25 °C) ΔH‡/kJ mol−1 ΔS‡/J K−1 mol−1

First step

Second step

2.85 ± 0.05 89 ± 2 +60 ± 5

1.48 ± 0.06 63 ± 3 −30 ± 7

Chapter 5

278 −

leaving SCN (i.e. SE2 mechanism) has been proposed without involvement of an intermediate CoIII–NCS–HgII of sufficiently long lifetime. The relative effectiveness of metal ions, for example Ag+, Hg2+ or Tl3+, in promoting aquation of chloride ligand in complexes of Cr(III) can be correlated with the hardness or softness of these ions.50c Linkage isomerization of cis-­ [Cr[CN)2(en)2]+ to cis-­[Cr(NC)2(en)2]+ is also catalysed by Ag+.4c For Hg(ii)-­catalysed aquation of [MX(NH3)5]2+, the reported ΔV‡ values in 0.31 M [H+]51a are given in Table 5.10. From the similarity of the ΔV‡ values, a dissociative mechanism following rapid M–X–Hg bridge formation was proposed. In the Hg2+-­catalysed aquation of [MCl(NH3)5]2+, the effectiveness of Hg2+ increases in the sequence Ru(iii) < Rh(iii) < Co(iii).51b The aquation52 of [IrCl6]3− is retarded by Ga3+, Fe3+, Ce3+ and Eu3+ owing to the formation of less reactive ion pairs and the effects of the added cations correlate with their ionic radii. In contrast, the aquation of [IrCl6]3− is considerably accelerated53 by In3+ [the rate is doubled on addition of only 10−5 M In3+, KE (see eqn (5.27) ≈ 100 at 20 °C] and also by Hg2+. Aquation of the extremely inert [OsF6]2− is catalysed by Zr(iv). Surprisingly, the rate constants for Zr(iv)-­catalysed aquations of [OsF6]2−, [ReF6]2−, [PtF6]2− and [PF6]2− are almost the same under identical conditions. Comparison with the rate constant for depolymerization of polynuclear Zr(iv) species in aqueous solution shows that this is the rate-­determining step that generates the catalytically active Zr(iv) species.54 Hg2+-­ and NO+-­induced aquations of [CoC1(NH3) (tren)]2+ and [CoN3(NH3)(tren)]2+ have been reported.55 Solvolysis of azido complexes catalysed by HNO2 have been studied.56 For [CoN3(NH3)5]2+ in nitrate medium (up to 1 M), the established rate-­law is   

−d[complex]/dt = (k1 + k2[NO3−])[H+][HNO2][complex]



(5.25)

  

This rate law can be accommodated by several mechanisms, including one that involves a five-­coordinate intermediate for the NO3−-­independent path. The authors reviewed the problems associated with interpreting the results of competition experiments in relation to Hg(ii) catalysis and nitrous acid–azide reactions, and ideas for resolving some of the ambiguities were presented. A combination of rate-­law and competition evidence still supports the five-­coordinate Co(iii) complexes as intermediates in many of these catalysed aquations. A general rate-­law equation for catalysis by a single electrophile E (E = H+ or Mm+) is   

Table 5.10  Values  of activation volume for mercury(ii)-­catalysed aquation of some halidopentaammine complexes MX(NH3)5.

‡

3

−1

ΔV /cm mol Temperature/°C

M = Co(iii)

M = Co(iii)

M = Rh(ii)

M = Cr(iii)

X = Cl

X = Br

X = Cl

X = Cl

−1.7 ± 1.0 15

+0.8 ± 0.5 15

+0.7 ± 0.4 25

−1.0 ± 0.4 15

Catalysed Reactions and Formation Reactions



−d[complex]/dt = {k0 + k′EKE[E]/(1 + KE[E])}[complex]

279

(5.26)

  

where k0 is the rate constant for the spontaneous (uncatalysed) aquation, kE′ is the rate constant for aquation of the complex–E adduct and KE is the equilibrium constant for the adduct formation (binding of E to the complex). Hence kE′ × KE = kE, where kE is the rate constant for the catalysed path. Depending on the two limiting situations corresponding to KE[E] being (a) very much larger than 1 or (b) very much smaller than 1, eqn (5.26) leads to the observed rate constant as   



(a) kobs ≈ k0 + kE′

(5.27)

(b) kobs ≈ k0 +kE′KE[E] = k0 + kE[E]

(5.28)

  

  

In agreement with the view that the electrophile forms an adduct with the complex which then undergoes the reaction, it has been observed that for such catalysed aquation of oxalato complexes the plot of log kE versus log KE(ox) is linear, where kE is the observed rate constant for the catalysed path and KE(ox) is the formation constant for the monooxalato complex of the electrophile, E (a metal ion). This and similar observations have been reported by the senior author and others for complexes of many other such ligands. Catalysis of the aquation of cis-­[Cr(C2O4)2(OH2)2]− into [Cr(C2O4)(OH2)4]+ by several metal ions of the first transition (3d) series in acidic (HClO4) media has been reported.57a The results showed that log kM (kM being the rate constant for the metal ion-­catalysed path) varies linearly with log KMox (where KMox is the formation constant of the monooxalato complex of M2+); the reactivity sequence is Cu2+ > Ni2+ > Co2+ > Mn2+. This suggests the operation of the mechanism shown in Scheme 5.8. In a similar study57b of metal ion-­assisted dissociation of [Cr(ox)3]3− and [Cr(ox)(OH2)4]+ and Fe3+-­assisted dissociation of cis-­[Cr(ox)2(OH2)2]−, the experimental results showed that   



kobs = k + kM[Mm+]

(5.29)

k = k0 + kH[H+]

(5.30)

  

where   

  

Scheme 5.8

Chapter 5

280

Mm+ aq, the kM 2+ +

3+

2+

For different values decrease in the order Fe >> Cu > Ni2+ > 2+ 2+ Zn > Co > Mn (H has a position between Co2+ and Mn2+). This order reflects the affinity of binding of Mm+ to oxalate, indicating the formation of a reactive intermediate as in the case of the bisoxalato complex, since a plot of log kM versus log KMox is linear. kFe/kH at 25 °C is 3500 for [Cr(ox)3]3−, 429 for cis-­[Cr(ox)2(OH2)2]− and 137 for [Cr(ox)(OH2)4]+. Even for the cationic [Cr(ox) (OH2)4]+ complex Fe3+ is a far superior catalyst to H+, as would be expected from the above mechanism. For the kFe path the plot of ΔH‡ versus ΔS‡ is linear for the mono-­, bis-­and tris-oxalato complexes and so also is log kMtris versus log kMbis, which suggests similar mechanisms. For any of the substrates the ΔH‡ values for kM are nearly identical and the variation in kM values is essentially due to variations in the ΔS‡ values; the plot of ΔS‡ versus log kM is linear. Direct evidence for the formation of a [Cr(ox)3]3−–Cu2+ adduct has been furnished.57c Dissociation of cis-­[Cr(mal)2(OH2)2]− in acidic medium is catalysed by M2+ aq ions (M = Co, Ni, Cu, Zn). The reaction occurs in two concurrent paths, H+ catalysed and M2+ catalysed:58a   



kobs = kH[H+]+kM[M2+]

(5.31)

  

For different M2+, the kM values parallel the values of the formation constants of the [M(mal)] complexes (mal = malonate). Similar studies on [Cr(mal)3]3− have been reported.58b Several such studies in the senior author's laboratory have established the generality of this mechanism for H+-­and Mm+-­catalysed dissociation of metal chelates.59 The metal ion-­catalysed dissociation of tris(biguanide) complexes of Co(iii) and Cr(iii) has also been investigated.60 It was observed that in these cases the value of log kM increases linearly with (Z – S)/r2 of the catalysing metal ion M2+ (where M = Mn, Co, Ni, Cu, Zn, Z is the atomic number of M, S is Slater's screening constant and r is the radius of M2+), which is a measure of the complexing ability of M2+.61 This also suggests the formation of an intermediate of the following type:

Similarly, for the metal ion-­catalysed dissociation of [Co(C2O4)(NH3)5]+, Dash and Nanda62a,b observed the following order of catalytic activity of M2+ and M3+ ions, which parallels the thermodynamic stabilities of the intermediate dinuclear species having CoIII(µ:η2-­ox)Mm+:

Catalysed Reactions and Formation Reactions

281

Mn(ii) < Co(ii) < Ni(ii) < Cu(ii) < Zn(ii); and In(iii) < Al(iii) < Ga(iii) < Fe(iii) In a similar study of the pentaamminesalicylatocobalt(iii) ion, the observed rate constant was found to vary with the concentrations of the catalysing metal ion and H+ as follows:62c

  



kobs = k0 + kH[H+] + (kMKM[Mn+]/[H+])(1 + KM[Mn+]/[H+])

(5.32)

  

This is analogous to eqn (5.26). From similar studies on [Co(O2CR)(NH3)5]2+, the following order of catalytic activity was reported62d,63,64 Fe3+ > Al3+ > In3+ > Cu2+ > Zn2+ The metal exchange in metal complexes of multidentate ligands, such as EDTA, occurs by a similar mechanism: NiY 2  + *Ni2+ NiY*Ni Ni2+ +*NiY 2 

The dissociation of NiY2− is likewise catalysed by Cu2+, Zn2+, etc.:65 NiY 2  +M 2+ NiY  M Ni2+ +MY 2 

In both cases, the following type of dinuclear intermediate is formed:

In this, each metal ion remains bonded to an iminodiacetate grouping, hence a plot of the logarithm of the rate constant versus the logarithm of the formation constant of the iminodiacetate complex of the catalysing metal ion is linear, which is indeed evidence for the aforesaid mechanism.66 The reaction [CoCl(edta)]2− → [Co(edta)]− + Cl− is catalysed by metal ions (Sc3+, In3+, La3+, Tl3+, Hg2+, etc.), Hg2+ being most effective and Tl3+ is similar. Rate constants for the catalysed reaction correlate with a function incorporating the formation constants of complexes of M with Cl− and carboxylate. Rudakov and Kozhevnikov67a reported a correlation between the rate constant for the metal ion-­catalysed solvolysis of tert-­butyl halides and the stability constants of the corresponding metal–halide complexes MX(m−1)+. They also showed67b that a similar correlation applies, albeit rather approximately, to metal ion catalysis of the aquation of halido complexes of cobalt(iii), chromium(iii) and rhodium(iii). The catalysts mentioned include not only Ag+, Hg2+ and Tl3+, but also species such as HgCl+ and TlC12+. The correlation can be improved by making allowance for the different coulombic repulsions in

Chapter 5

282

systems of differently charged products. If the rate constant for catalysed aquation is kM and that for the uncatalysed aquation is k0, the stability constant of the halido complex, MX(m−1)+, of the catalysing metal is KMX, the product of the charges on the reactants is ZAZB and C is a coulombic interaction constant, then the correlation conforms to the equation   



log(kM/k0) = −0.7 + 0.84(logKMX) − CZAZB

(5.33)

  

This equation and that for the catalysed aquation of tert-­butyl halides are very similar, hence the catalytic activities of the metal ions in the solvolysis of inorganic and organic halides must be similar. A few other examples of metal ion-­catalysed reactions of metal complexes are the isomerization of trans-­[Cr(ox)2(OH2)2]− to the cis isomer,68a racemization of [Cr(ox)3]3− by M2+ ions68b and racemization of [Cr(ox)2phen]− by Cu2+ ions.68c But racemization of [Co(bigH)3]3+ has been reported to be retarded by cations.276 For metal ion catalysis of outer-­sphere electron transfer reactions, see Chapter 7.

5.1.3  Nucleophilic Catalysis Catalysis by anions (nucleophiles) is also known for various types of reactions of metal complexes. In ligand replacement reactions this is generally observed where the metal ion is coordinatively unsaturated and can add the nucleophile to increase its coordination number, as in square-­planar complexes which can add up to two such nucleophiles. Thus, the dissociation of [M(ox)2]2− and several other complexes of Pd(ii) and Pt(ii) are catalysed by both H+ and Cl−/Br− (see Chapter 3); intermediates (adducts) such as [M(ox) (oxH)Cl]2− and [M(oxH)2Cl2]2− are involved (where oxH represents chelated oxalate protonated at the free carbonyl oxygen and not the unidentate oxH−). Such studies have been reported69a for [Pt(ox)2]2− in aqueous HX having excess X− (X = Cl, Br) forming [PtX4]2−. The experimental results led to kobs = (k1 + k2[H+])[[X−] Similarly, the reaction of [Pd(bigH)2]2+ (bigH = biguanide) in aqueous HCl solution in the presence of excess Cl− occurs in two steps, forming [PdCl2(bigH)] and [PdCl4]2−. For each of the two steps the experimental results led to (see Chapter 3)69b kobs = k1[H+][Cl−] + k2[H+]2[Cl−]2 Similar observations have been reported in the reaction of [Pt(ox)2]2− with SCN−,70 and of [Pd(ox)2]2− and [Pd(mal)2]2− with Cl− in aqueous HCl solution.71 However, since in all such reactions the X− remains bonded to the metal in the final product, these are more appropriately considered as acid catalyzed ligand replacement of bidentate ligands with X− ions. Catalysis by Cl− and Br−, in addition to catalysis by H+, has been reported72a for the aquation of [M(NH3)5(ONO)]2+ (M = Co, Cr), where the reaction proceeds through reactive intermediates of the following type (X = Cl, Br):

Catalysed Reactions and Formation Reactions

283

and this splits (by O–N bond fission, rate-­determining step) to eliminate XNO (which is rapidly hydrolysed, forming HNO2 + HX), leaving (NH3)5M−OH2+, which immediately protonates to form [M(NH3)5(OH2)]3+. Aquation of [CoF(NH3)5]2+ by Al3+ has been reported to be accelerated by the anions SO42−, ox2− and mal2− forming both the aquapetaammine and anionopentaammine complexes of cobalt(III) as products.72b Outer sphere association involving the cationic Co(III) complex-bivalent anion-Al(III) has been proposed. Aquation of [CrX(OH2)5]2+ (X = CI, Br) is also specifically catalysed by HNO2, due to the ready formation of [Cr(ONO)X(OH2)4]+ as intermediate by direct attack of HNO2 on coordinated H2O without Cr–OH2 bond fission.73 This is supported by the observation that aquation of [CrBr(NH3)5]2+ is not accelerated by HNO2. Similar catalysed aquation has been reported for cis-­ and trans-­ [CrCl2(OH2)4]+ and cis-­[CrCl(NH3)4(OH2)]2+, but no such catalysis of the trans-­ [CrCl(NH3)4(OH2)]2+ due to the absence of an H2O cis to the Cl ligand needed for formation of the following type of intermediate:73

Oxygen exchange between [Cr(NH3)(OH2)]3+ and solvent (H2O) is catalysed by CO2 through the formation of an intermediate bicarbonato complex by a pseudo-­ substitution process that requires no Cr–OH2 bond fission and the experimental observations (at pH 1.5–5) are in conformity with the following rate-­law [M = (H3N)5Cr3+]:74a Rate = k1[M–OH23+] + (k2 + k3[CO2])[M–OH2+] At 25 °C and I = 0.1 M (NaClO4), k1 = (6.3 ± 0.5) × 10−5 s−1, k2 = (17 ± 0.4) × 10−5 s−1 and k3 = 7 M−1 s−1. This rate law suggests three concurrent reactions, viz. spontaneous (uncatalysed) reactions of M–OH23+ and M–OH2+ and CO2-­catalysed reaction of M–OH2+. For the catalysed exchange, the following mechanism has been proposed:

Chapter 5

284

The labelled CO2 (*OCO) results from rapid exchange of oxygen between the dissolved CO2 with H2O* in the aqueous solution containing H2O*. This mechanism of exchange of oxygen between M–OH2+ and H2O* involves insertion of the CO2 (labelled with *O) into the O–H bond of M–O–H forming the labelled bicarbonato complex, which by oxygen scrambling through a seven-­ coordinate transition state leads to the oxygen exchange. However, no such CO2-­catalysed oxygen exchange occurs with the analogous complex of Co(iii). This is because, unlike in the case of Cr(iii), which has the affinity for seven coordination [hence ligand replacement reactions of Cr(iii) are generally associative, see Chapter 4], Co(iii) lacks the tendency for seven coordination and its reactions are dissociative (see Chapter 4). Aquation and anation reactions of Cr(iii) are also catalysed by CO2.74a In a similar manner, exchange of H2O of [Cr(OH2)6]3+ with that of solvent (H2O*) is strongly catalysed by SO2−3 (kex is increased by a factor of 2.5 × 104 at 25 °C (I = 0.65 M) through the formation of a bisulfito complex.74b Oxo anions such as sulfite, nitrate and acetate labilize the X− in [CrX(OH2)5]2+.74c Linkage isomerization of cis-­[Cr(CN)2(en)2] to cis-­[Cr(NC)2(en)2] catalysed by Ag+ has been reported.75a In a solution having a fairly high concentration of HClO4, cis-[Cr(en)2F2]+ undergoes aquation, leading to replacement of one F− by H2O, which is acid catalyzed due to protonation of a F− ligand in the complex. In the presence of a fairly high concentration of F−, the reaction is further accelerated due to the formation of an intermediate complex having a CrIII−F−H...F− bond which forms the product by loss of HF2−.75b Anation of cis-[Cr(en)2(OH2)2]3+ by oxalate has been reported to be catalysed by nitrate ions.75c Some other examples of nucleophilic catalysis are the effect of OH− ion on the rate of dissociation of [Ni(bipy)3]2+,76 racemization of [Co(edta)]−,77 dissociation of [Cr(bigH)3]3+,7 and the reaction of edta4− with several complexes of Cr(iii) in alkaline media.78 In many such cases the OH− accelerates the reaction, but in some cases it retards the reaction (negative catalysis).75a Halide anation reactions of Pt(iv) complexes proceed via halide-­assisted paths.79a Thus, for the bromide anation reaction of cis-­[PtBrCl4(OH2)]− carried out at 25 °C, the results led to the following rate law (k = 0.43 M−2 s−1): −d[complex]/dt = k[Br−]2[complex] and for chloride anation at 50 °C the rate law is (k1 = 2.7 M−2 s−1; k2 = 2.95 × 10−3 M−2 s−1):

−d[complex]/dt = (k1[Br−][Cl−] + k2[Cl−]2)[complex]

Base hydrolysis of trans-­[PtBr2(CN)4]2− in the presence of Br− occurs by two concurrent paths:79b Rate = (k1[OH−] + k2[Br−])[PtIV complex] At 25 °C (I = 1 M KNO3) the following values were reported for this reaction, which is reversible (equilibrium constant {Keq}):

Catalysed Reactions and Formation Reactions 2

−1

−1

‡

285 −1

‡

−1

−1

k1 = 1.15 × 10 M s ; ΔH1 = 26.4 kcal mol ; ΔS1 = +38.3 cal K mol

k2 = 4.17 × 10−3 M−1 s−1; ΔH2‡ = −13.7 kcal mol−1; ΔS2‡ = −24.4 cal K−1 mol−1



Keq = 6.98 × 104 Neither the conjugate base mechanism nor the normal ion-­pair mechanism (see Chapter 4) is possible in this system. The authors suggested that the highly positive value of the entropy of activation for the k1 path is consistent with an Id mechanism. The k2 path has been considered as a reaction occurring through an activated complex in which the Br− is associated with a Br− ligand, and this is expected to increase the electron density on the Pt, facilitating dissociation of the Pt–Br bond with a fast entry of OH− into the vacated coordination position, and this, according to the authors, is consistent with the fairly high negative value of the entropy of activation for this path. Unlike in most cases of reactions of complexes of Pt(iv), no catalytic effect of [Pt(CN)4]2− was observed. The presence of several multivalent anions such as sulfate, oxalate and malonate has been reported80 to accelerate the Al3+-­assisted aquation of [CoF(NH3)5]2+. Nucleophilic catalysis of electron transfer reactions is known. Reduction of [Co(NH3)6]3+ by V2+aq takes places by concurrent anion-­dependent (F− < SO42− < Cl− < I−) and non-­dependent paths:276a Rate = (k1 + k2[Xn−1])[V2+aq][CoIII complex] Similar observations have been reported276b for the dinuclear complex [(NH3)5Co–(µ-­NH2)–Co(NH3)5]5+. Since for any catalysing anion the ratio k2/k1 has nearly the same value for both complexes of Co(iii), which have an overall charge of 3+ and 5+, respectively, association of the catalysing anion with V2+ was proposed for the anion-­dependent path. A similar rate law has been reported276c for the reduction of [Fe(phen)3]3+ by Fe2+ aq and it is reasonable that association of the catalysing anion with the Fe2+ is responsible for the anion-­dependent path; however, the authors proposed association of the anion with the Fe(iii) complex by nucleophilic attack on C of a phen ligand.

5.1.4  E  lectron Transfer Mechanism of Reactions of Metal Complexes As a rule, a ligand replacement reaction of a metal complex is catalysed by a compound of the metal in another oxidation state81–96 due to operation of an inner-­sphere electron transfer mechanism.94 Thus, Co(ii) catalyses ligand exchange and substitution reactions83,84 of trans-­İCoX(dmgH)2L]. Examples of Cr2+-­catalysed reactions of Cr(iii) complexes are well known;85–92 some specific examples are aquation of [CrX(OH2)5]2+ (X = Br, I),85 aquation of [CrF(NCS)(OH2)4]2+,86 ‡ aquation and ligand replacement (I− by F−, Cl−, Br−) reactions of [CrI(OH2)5]2+,87 aquation of [Cr(N3)(OH2)5]2+,89a ‡

I n this system the reaction proceeds to ca. 80% completion, forming almost exclusively [CrF(OH2)5]2+. Based on the observed rate law, equilibrium formation of an unproductive (for product formation) intermediate CrIII–NCS–CrII followed by its transformation into a fluoride-­ bridged species which forms the product in the rate-­determining step has been proposed.

Chapter 5

286 3− 89b

2+ 90

[Cr(CN)6] , [Cr(CN)(OH2)5] , and cis-­ and trans-­[Cr(NH3)4(OH2)2]3+.91 − Ligand (Br ) substitution of [CrBr(OH2)5]2+ by Cl− and F− is also catalysed by + Cr2+ aq. Formation of an intermediate [Cr(BrX)(OH2)5] has been postulated from − − the following rate law for substitution of Br by Cl (for substitution by F− the rate law is more complex since most of the uncomplexed F− is present in the acidic aqueous solution as HF):92 Rate = kCl[Cl−][Cr2+][CrBr(OH2)2+ 5 ] Aquation of [Fe(CN)6]3− is catalysed by [Fe(CN)6]4− and also acid catalysed (no reaction at pH 7).93 Ligand replacement and ligand exchange reactions of complexes of Pt(iv) are catalysed by Pt(ii) and several examples are known.94,95 Similarly, Pd(ii) catalyses both the forward and reverse reactions of the following equilibrium:96 [PdCI2(en)2]2+ + 2Br− ⇌ trans-­[PdBr2(en)2] + 2CI− Interestingly, anation of trans-­[PtCl(CN)4(OH2)]− by Br− involves reductive elimination followed by oxidative addition:97,98 trans-­[PtCl(CN)4(OH2)]− + Br− → [Pt(CN)4]2− + ClBr + H2O [Pt(CN)4]2− + ClBr → trans-­[PtBr(Cl)(CN)4]2− [Co(phen)3]2+ catalyses the racemization of [Co(phen)3]3+ by an outer-­ sphere electron transfer mechanism.99a Reactions of the complexes [CoX (NH3)5](3−x)+ (Xx− = F−, Cl−, Br−, N3−, NCS−, NO3−, MeCO2−, SO42−, S2O32−, etc.) with CN− are catalysed by Co2+ (see Chapter 7).99b

5.2  Formation Reactions A reaction that involves replacement of a metal-­bound solvent molecule by another ligand is commonly referred to as formation, and is called anation if the entering ligand is an anion. Such reactions can be represented by a general mechanistic scheme (Figure 5.3) in which a solvent molecule (H2O in the given scheme) in the inner coordination sphere is replaced with an entering monodentate ligand X. Substitution at a labile metal ion is generally discussed in terms of the mechanism represented by route B in Figure 5.3, as first proposed by Eigen.100 For oppositely charged substrate and entering ligand, the outer-­sphere complex is an ion pair, but in favourable cases fairly strong outer-­sphere association may result between a cationic complex and an uncharged ligand, or even between an anionic substrate and an anionic ligand due to hydrogen bonding,101 stacking interactions,102 etc. Route A (D mechanism) is considered most appropriate for highly charged anionic complexes of metals.103 Formation of the outer-­sphere complex is, of course, diffusion controlled, setting an upper limit for the rate of formation in solution.

Catalysed Reactions and Formation Reactions

287

Figure 5.3  General  scheme for the formation of metal complexes in which (2) and

(7) are outer-­sphere complexes; L = 5H2O. A: route (1) → (6) → (4) represents a D process. B: route (1) → (2) → (7) → (4) also represents a D process, but involving outer-­sphere association in a pre-­equilibrium and hence D-­OS, but may be designated D-­IP when the outer-­sphere complex is an ion pair. Route (1) → (2) → (3) → (4) represents concerted interchange processes Id or Ia depending on the relative importance of M–OH2 bond dissociation and M–X bond making in the formation of the transition state. If the outer-­sphere complex is an ion pair, the process is designated I-­IP. D: route (1) → (2) → (5) → (4) represents an associative A process with outer-­sphere association in a pre-­equilibrium which may be designated A-­IP if the outer-­sphere complex is an ion pair.

Metals can be classified according to the relative rates of steps (1) → (2) and (2) → (4). Group A contains metal ions that have a comparatively low charge density at their surface, e.g. the alkali metals (including Li+) and the heavier alkaline earths (Ca2+, Sr2+ and Ba2+). The water molecules are held comparatively weakly, with the result that steps (1) → (2) and (2) → (4) have comparable rates. Group B contains metals of intermediate surface charge density, in which the water molecules are held moderately strongly, so that the second step is now slower than the first. This group includes many bivalent metals, such as Mg2+, Fe2+, Co2+ and Ni2+. Group C contains metal ions of sufficiently high surface charge density that the water molecules are very strongly held and hydrolysis (i.e. splitting of an H+ from a bound water molecule) can occur before water loss. This group includes small, highly charged ions such as Be2+, Al3+ and Fe3+ (probably), and tetravalent ions. It is the least well understood group since the overall mechanisms may involve a number of non-­separable steps. Substitution at the hydrolysed ion M(OH)(n−1)+ is faster than at the non-­hydrolysed ion Mn+ [in fact M(OH)(n−1)+ is usually in group B], so anything that promotes hydrolysis will also promote substitution. In any event, for group C the overall rate constant is largely influenced by the nature of the incoming ligand, for example, its basicity. This influence is much less apparent with metals of groups A and B, except insofar as the formation of second and subsequent bonds with chelating ligands may become rate determining, Most group A metal ions have the d0 electronic configuration of the noble gases, and the interaction of the metal with a ligand is primarily electrostatic in nature. Because the charge density (charge:radius) is small, the water molecules are weakly held and the rate of their loss is comparable to the

Chapter 5

288 9

10

−1

diffusion-­controlled value of around 10 –10 s . The complex formation rate constant for an ion of this group is therefore nearly equal to the maximum possible value, unless chelation steps are important, and the stability of the complex is reflected in the dissociation rate constant. With group B metal ions, step (2) → (4) is considerably (i.e. >10 times) slower than step (1) → (2), so a fast pre-­equilibrium step is followed by the rate-­limiting loss of a water molecule from the inner coordination sphere. To a first approximation (provided that the concentration of outer-­sphere complex is small compared with that of the uncomplexed metal ion), the observed overall formation rate constant kr is given by the equation   



kf = Koskwl

(5.34)

  

where Kos is the equilibrium constant for the formation of the outer-­sphere complex and kwl is the rate constant for the water loss process On this basis, kwl is essentially the rate constant for water exchange at the metal ion, and this can frequently be measured independently by NMR spectroscopy. Unfortunately, it has not yet been possible to measure Kos directly in such reacting systems, so it has not been possible to confirm the mechanism unequivocally. However, the values of kf do appear to be more or less typical of the metal ion and similar in magnitude to that of the water exchange rate constant, kex. Some of the log kex values in Figure 5.4 are lower than the values reported subsequently (see Chapter 1, Figure 1.3), viz. 8.9 for Cr(ii), 9.3 for Cu(ii) and 5.25 for Ti(iii), but these do not alter the trend.

Figure 5.4  Rate  constants for water exchange of di-­ and trivalent hexaaqua metal ionsofvariousd-­electronconfigurations(dn).104Asimilartrendisobservedforkf(M−1 s−1) for the formation of [M(OH2)5L]x+ from [M(OH2)6]n+.

Catalysed Reactions and Formation Reactions

289

It is found that the forward rate constant is approximately inversely proportional to the ionic radius for metals of a given charge type, but the presence of partially filled d-­orbitals in the metal ion can cause considerable complication. This effect of d-­electron configuration on formation (kf ) and water exchange (kex) rate constants for various di-­and trivalent metal ions is shown in Figure 5.4. The general trend of rates is reminiscent of the ligand field stabilization diagrams and, indeed, general correlations have been made between substitution rates and ligand field activation energies, although the detailed picture has not been satisfactorily explained. Companion105 examined the possibility of improving the LFAE calculations by the use of many-­electron methods and improved ligand-­field parameters. The increases in rate constant on going from V2+ to Cr2+ (d3 to d4) and from Ni2+ to Cu2+ (d8 to d9) have been rationalized in terms of the Jahn–Teller effect. Thus, in octahedral Cr2+ and Cu2+ there are two axial water molecules that are less firmly held than the four equatorial water molecules, and the incoming ligand is logically assumed to replace the former. The consequence of this is that Cr2+ and Cu2+ behave more like metals of group A than group B. A mechanism of the type shown in routes B, C or D (Figure 5.3) will, under certain conditions, show a first-­order dependence on the concentration of the entering ligand, X, irrespective of whether transformation of the outer-­ sphere complex into the product (inner-­sphere complex) is associative or dissociative. Thus, if the outer-­sphere complex (OSC) is formed in a fast equilibrium, followed by its transformation into the product in a rate-­determining step, the formation may be represented by   

K OS



kos M  OH2  X #M  OH2 X   M  X  H2 O (OSC)

(5.35)

  

Under the condition TX >> TM (TX and TM being the total concentrations of X and M in the system, respectively), this leads to   



kobs = kosKosTX/(1 + KosTX)

(5.36a)

  

It follows that if KosTX > 1, and under this condition kobs will reach a limiting value and become independent of TX:   

  

kobs ≈ kos

(5.37)

Chapter 5

290 106

Experimental evidence is available to indicate that outer-­sphere complexation of an aqua complex entity does not markedly alter the value of its water exchange rate constant, kex. Hence the experimental value of kos in either case will be nearly the same as kex or even less for a dissociative transformation of the outer-­sphere complex. A lower value is often seen as X has to enter into an unfavourable competition with H2O (present in extremely high concentration) for the vacant coordination site of M. However, the value of kobs will never exceed kex for a dissociative process in the transformation of the outer-­sphere complex, as in a D or Id mechanism. For associative transformations, significantly higher kos values are generally observed. Moreover, since for the large number of different uncharged ligands (or those having the same magnitude of charge), the values of Kos will be nearly the same, the kf value will be more or less insensitive to the nucleophilicity of different entering ligands (unchanged or bearing the same charge) for a dissociative process. However, for an associative process, as in an A or Ia mechanism, there will be considerable variation in kf, with the nucleophilicity of the entering ligand, and in this case for some good nucleophiles Kos may be higher than kex. These criteria have been widely used to distinguish between associative and dissociative formation processes. Thus, the limiting rate constants for the anation of [Co(NH3)5(OH2)]3+ by Xn− the values at 45 °C (kex included for comparison) are given in Table 5.11. Hence klim is fairly insensitive to the nature of the entering ligand. However, compared with the case of [Co(NH3)5(OH2)3+], the anation reactions of aquacobalamin with Yn− show that the rate is far less sensitive towards the entering nucleophiles (see Chapter 4); even [Co(CN)5(OH2)]2− shows much greater discrimination107 towards Yn− and similarly for many other such complexes of Co(iii).108 Thus, the ratio k N3 k Br  is 5.3 for [Co(CN)5(OH2)]2− but 1.2 for aquacobalamin. It is reasonably tempting to consider the behaviour of aquacobalamin as corresponding to an almost D process (rather than Id) as in the following proposed reaction scheme: k1

Co(L)OH2 # Co(L)  H2 O k 1

k  Co(L)(X) Co(L)  X  2

This leads to the formation rate constant kf = k1k2/k−1 [see eqn (4.5) in Chapter 4] and rate of product (P) formation d[P]/dt ≈ kf[Co(L)(OH2)][Y] under the condition k2[Y] > k−1, if follows from eqn 4.3 in Chapter 4) that Rate = d[P]/dt ≈ k1[Co(L)(OH2)] Table 5.11  Rate  constants of reactions of aquapentaamminecobalt(iii). Xn−

klim/s−1

Xn−

klim/s−1

H2O N− 3 SO24−

1 × 10−4 (kex) 1 × 10−4 2.4 × 10−5

Cl− NCS−

2.1 × 10−5 1.6 × 10−5

Catalysed Reactions and Formation Reactions

291

i.e. k1 = kf. This limiting value of kf, attained at a high concentration of Y, is independent of the nature of Y. However, for the other Co(iii) complexes, which presumably react through a precursor outer-­sphere complex in an Id process, the corresponding limiting value is kos [see eqn (5.13)] and somewhat greater discrimination is observed towards entering Y; although this limiting k may approach kex, it can never exceed kex, as is to be expected. Substitution of the axial H2O in the cobaloxime trans-­[Co(dmgH)2Me(OH2)] is also very rapid (as for aquacobalamin) and the k value shows little dependence on the nature of the entering ligand, pointing to a dissociative mechanism as for aquacobalamin.109 [Cr(OH2)6]3+ has been found to react with a variety of anions, viz. Cl−, Br−, I−, NO3− and SCN− (both as N-­bonded and S-­bonded), of strong acids at widely different rates [kan25 °C = (0.08–73) × 10−8 M−1 s−1],110 indicating significant incoming ligand participation in the transition state; the slope (less than 1) of the LEER plot (see Chapter 2) of ∆G (free energy of reaction) versus ∆G‡ (free-­energy of activation) also suggests some associative character (Ia mechanism). In contrast, the range of rates for reaction of [Cr(OH)(OH2)5]2+ with the same series of ligands is small [kan25 °C = (0.26–8.5) × 10−5 M−1 s−1],110a which suggests a predominantly dissociative mechanism (Id). It should be noted that a dissociative mechanism is more likely in the presence of a π-­donor ligand such as OH− cis to the leaving group (see Chapter 4, Figure 4.3). Rate constants for the formation of [Cr(NH3)5X]2+ from [Cr(NH3)5(OH2)]3+ and a variety of univalent anions, X−, also vary by less than an order of magnitude, and the observed rates are only slightly less than the rate of water exchange, kex, of [Cr(NH3)5(OH2)]3+. Thus, with NCS− and Cl− as entering ligands, the ratio of the formation rates, kNCS/kCl, at 25 °C is 60 for [Cr(OH2)6]3+ and only 6 for [Cr(NH3)5(OH2)]3+. Interestingly, for [Co(OH2)6]3+ also this ratio is ≥ 43, which indicates that an Ia mechanism may be operative here also, despite a general preference for dissociative reactions encountered with cobalt(iii) complexes. Since NH3 is a better σ-­donor than H2O, it has been proposed110 that the tendency is towards a dissociative mechanism with increasing σ-­donor character of the non-­leaving ligands. Most of the later work111 supports an Ia mechanism for reactions of [Cr(OH2)6]3+ and an Id mechanism for reactions of [Cr(OH2)5(OH)]2+ and [Cr(NH3)5(OH2)]3+. Application112 of LFER has indeed also been fruitful to lend further support to the mechanisms in such cases (see Chapter 4, Figure 4.6). Lo and Swaddle113 reported that the formation of [Cr(DMF)5X]2+ from [Cr(DMF)6]3+ shows a strong dependence of the rate on the entering ligand (X−). Based on the ∆H‡ and ∆S‡ values reported by them, the kf/Kos = kos values are given in Table 5.12. The behaviour of [Cr(DMF)6]3+ is therefore similar to that of [Cr(OH2)6]3+ and the mechanism is predominantly associative. Table 5.12  k os values for the formation of [Cr(DMF)5X]2+ from [Cr(DMF)6]3+. X− 108kos/s−1 at 25 °C

Cl−

Br−

NCS−

N3−

7.1

0.41

66.2

1580

Chapter 5

292

Scheme 5.9 The reactions of [Co(NH3)5(OH2)]3+ with various ligands (X) are fairly insensitive to the nature of X (X = Cl−, Br−, NO3−, NCS−, H2PO4−, NH3),114 with rate constants in the range 1.3 × 10−6–2.5 × 10−6 M−1 s−1 (at 25 °C), as compared with the water exchange rate constant of 6.6 × 10−6 (value of kex/55.5, expressed in M−1 s−1). Melson and Wilkins115 showed that the second-­order rate constants for the formation of mono-­complexes of Ni(ii) from [Ni(OH2)6]2+ with NH3, N2H4, py, phen, bipy, terpy, etc., are in all the range 2 × 103–3 × 103 M−1 s−1 (at 25 °C), which also suggests a dissociative mechanism as for the reactions of [Co(NH3)5(OH2)]3+. A temperature jump method was used to obtain equilibrium constants and rate constants for formation of mono-­ and bischloro complexes of Fe3+ aq from [Fe(OH2)6]3+. For the mono species, the interchange process involving the ion pair [Fe(OH2)6]3+·Cl− (Kos = 1.1 M−1) is characterized by kf = 21.8 M−1 s−1 and kd = 37 s−1 at 25 °C (I = 2.6 M). The equilibrium formation constant (kf/kd) is 5.9 M−1. For the bis-­species the calculated value of kf is 324 M−1 s−1 from the measured116 kd = 185 s−1 and K2 = 1.75 M−1. A high-­pressure laser temperature jump technique was used to investigate the reaction between [Fe(OH2)6]3+ and Br− at pressures up to 2.76 kbar. The reaction involves both the hexaaqua ion and its conjugate base (see Scheme 5.9). For the formation reaction ∆V‡ (for k1) = −19 ± 4 cm3 mol−1 at I = 2.0 M and 25 °C, and an Ia mechanism was accordingly proposed. For the overall reaction involving [Fe(OH2)6]3+ and Br−, ∆V‡ (for k1K1) = −8 ± 4 cm3 mol−1. Previously reported values are for Cl− −4.5 ± 1.1 cm3 mol−1 and for NCS− an average value of 0 ± 5 cm3 mol−1. The large variation in the rate constants (k1K1) for reactions of [Fe(OH2)6]3+ with these three ions (at 25 °C k1K1 = 122, 4.8 and 1.6 for NCS−, Cl− and Br− ions respectively) is as expected for an Ia mechanism. In contrast, the reactions involving the hydroxo species [Fe(OH2)5(OH)]2+ are believed to proceed by an Id mechanism117.§ The equilibria corresponding to Ka, Kb and Kc involve the loss of an H+ from an aqua ligand bound to Fe3+. This is a general behaviour pattern (compare the cases of [Cr(OH2)6]3+ and [Cr(OH2)5(OH)]2+) due to the π-­donor character of the OH group strongly labilizing the ligand H2O in a cis position (see Chapter 4). In keeping with this, the observed rate constants, kex, for the §

Different values have been reported by Grant and Jordan118 (see Chapter 4, Table 4.18).

Catalysed Reactions and Formation Reactions

293

3+

2+

17

water exchange reactions of Fe and Fe(OH) studied by O NMR spectroscopy are 1.6 × 102 and 1.2 × 105 s−1, respectively, at 25 °C.118 This subtle difference in the mechanism is also revealed by the activation volumes, ∆V‡, for these water exchange reactions, the values being −5.4 and 2+ 119 +7.0 cm mol−1 for Fe3+ For the water exchange aq and Fe(OH) aq, respectively. 3+ ‡ 3 of Cr aq also ∆V is significantly negative (−9.6 cm mol−1),120 as expected for an associative process, but +2.7 cm3 mol−1 for Cr(OH)2+ aq, indicating a dissociative process. The effect of varying ionic strength and the anions NO3−, ClO4−, SO42−and − Cl on the rate of formation of Fe(NCS)2+ aq ion was investigated by a stopped-­flow method121 and the catalysed formation of Fe(NCS)2+ aq was investigated in the presence of aliphatic carboxylic acids (HL = RCO2H; R = H, Me or Pr) by the temperature jump technique.122 The latter reactions involve the fast formation of Fe(OH)(HL)(NCS)aq+ ions, followed by a slower loss of HL from this first-­ formed species. A mechanism analogous to that proposed by Rorabacher (ICB effect, see Figure 5.6) has been suggested. A stopped-­flow method was used to study the reactions of the [Fe(OH2)6]3+ ion with the α-­hydroxycarboxylic acids glycolic, dl-­lactic, dl-­malic and benzilic acids. The main pathway involves reactions of [Fe(OH2)5(OH)]2+ ion with HL, and an Id mechanism was proposed.123 A similar study involved reactions with acetohydroxamic acid, hydroxyproline hydroxamic acid, tryptophan hydroxamic acid and histidine hydroxamic acid.124 The second-­order rate constants at 25 °C are in the range (1.14–5.90) × 103 M−1 s−1. A spin change is involved in the formation of low-­spin bis(imidazole) complexes of the type trans-­[Fe(L)X2]+ (L = protoporphyrin-­IX dimethyl ester, X = imidazole or N-­propyl-­ or N-­methylimidazole) from reactions of high-­spin [Fe(L)Cl] with X.125 High-­spin six-­coordinate complexes of the type [Fe(L)CIX] are involved, and slightly different mechanisms have been proposed for the reactions with the unsubstituted and substituted imidazoles studied. Imidazole reacts more rapidly and a mechanism was proposed involving assisted loss of Cl− through hydrogen bonding with imidazole (kobs = Kk1[X]2/(1 + K[X]): k

[Fe(L)CI]  X#[Fe(L)CI(X)] k1 [Fe(L)Cl(X)]  X  [Fe(L)X 2 ]+  Cl  (assisted)

For the reaction with N-­methylimidazole, the proposed mechanism, after the rapid first step (equilibrium constant K), is as follows {kobs = k2K[X]/(1 + K[X])}: k1

[Fe(L)Cl(X)]# [Fe(L)X]  Cl  (unassisted) k2

k3 [Fe(L)X]+  Cl   X   [Fe(L)X 2 ]  Cl 

For the reaction with N-­methylimidazole, k1 = 8.1 ± 0.7 s−1 and K = 2.6 ± 0.3 M−1 at 25 °C. The high-­spin six-­coordinate intermediate was not detected

Chapter 5

294 125b

spectrally when X = imidazole and was detected only when X = N-­ propylimidazole or N-­methylimidazole.125a For the first-­stage anation of trans-­[M(L)(OH2)2]3− [M = Co, Rh, Cr; L = meso-­ tetrakis(p-­sulfonatophenyl)porphyrinate ion] by NCS−, the activation parameter values corresponding to kan have been reported (Table 5.13).126 Whereas the values of ∆S‡ would suggest a change in mechanism from dissociative for Co(iii) to associative for Rh(iii) and Cr(iii), the values of ∆V‡ suggest that all three react dissociatively. Hence a D mechanism has been proposed for Co(iii) and an Id mechanism for Rh(iii) and Cr(iii). For labile hexaaqua metal ions [M(OH2)6]2+ of the first transition (3d) series, the rate of the reaction [M(OH2)6]2+ + L → [M(OH2)5L]2+ + H2O for various ligands, particularly monodentate ligands, is nearly the same as the rate of water exchange of these aqua metal ions (suggesting a dissociative mechanism), and the rate follows the sequence (the order of k values in s−1 at 25 °C is given in parentheses)104 Cu  II   Cr  II   Mn  II   Zn  II   Fe  II   Co  II   Ni  II   V  II 

  10  9

  10  7

  10    10    10    10  6

5

4

2

In keeping with the predictions of the ligand field theory (see Chapter 1), [V(OH2)6]2+(d3) and [Ni(OH2)6]2+(d8) are much less labile than the others. Rorabacher and co-­workers127a summarized the considerable problems associated with the measurement of the water exchange kinetic parameters for [Cu(OH2)6]2+ and used data from results of the reaction of this ion with NH3 forming the monoammine complex [Cu(NH3)(OH2)5]2+ to estimate those parameters. For the formation and dissociation of the monoammine complex, the kinetic parameters are kf (25 °C) = 2.3 × 108 dm3 mol−1 s−1, ∆H‡ = 18.8 kJ mol−1, ΔS‡f = −20.9 J K−1 mol−1, kd (25 °C) = 2.0 × 104 s−1, ΔHd‡ = 39.75 kJ mol−1 and ΔS‡d = −29.3 J K−1 mol−1. On the basis of these data and evidence that the reaction proceeds by a dissociative mechanism, they reported the following as the “best” kinetic parameters for the water exchange reaction of [Cu(OH2)6]2+ with the solvent (H2O*) in solution: kex (25 °C) = 2.0 × 109 s−1, ΔH‡ex = 18.8 kJ mol−1 and ΔS‡ex = −4.2 J K−1 mol−1. Table 5.13  Values  of kinetic parameters for anation of some metal porphyrinate complexes.

M3+

T/K

kan/dm3 mol−1 s−1 ∆H‡/kcal mol−1 ∆S‡/cal K−1 mol−1 ∆V‡/cm3 mol−1

Co Rh Cr

293.2 278.2 288.2

103 1.21 × 10−2 9.47 × 10−4a

18.4 16.5 16.8

+14.4 −10.3 −12.8

+15.4 +8.8 +7.4

The rate constant for dissociation is 6.38 × 10−4 s−1, ∆H‡ = 15.7 kcal mol−1, ∆S‡ = −18.5 cal K−1 mol−1, ∆V‡ = +8.2 cm3 mol−1.

a

Catalysed Reactions and Formation Reactions 9

2+

295 4

2+

The high lability of d [Cu(OH2)6] {and also of d [Cr(OH2)6] } is due to the Jahn–Teller effect, which causes axial elongation, leading to a structure in which a trans pair of the two H2O ligands are rather weakly bonded and are very labile, permitting fast replacement of one of these H2O by the incoming ligand L. This is followed by rapid inversion in which the axial bonds in L–Cu–OH2 undergo contraction with concurrent elongation of a trans pair of equatorial H2O–Cu–OH2 bonds, thereby defining a new equatorial plane (see Figure 5.5) containing the donor atom of the ligand L, and the resultant structure will also be labile like the hexaaqua ion. If this was not the case, the second bond formation in the case of a bidentate ligand (or polydentate ligand) would require the loss of an equatorial water molecule, and this process is most likely to be slow (estimated value ca. 104 s−1), but that is contrary to the facts. In agreement with the above explanation regarding the lability of [Cu(OH2)6]2+, it has been reported that ligand replacement reactions of complexes of Cu(ii) are relatively slow in the absence of Jahn–Teller distortion, as in the case of the trigonal bipyramidal complex [Cu(tren)(OH2)2]2+. This reacts with pyridine (py) and imidazole (im) (replacement of H2O by these ligands) with rate constants kf (at 25 °C) of 1.6 × 105 and ca. 2 × 105 dm3 mol−1 s−1, respectively. Their dissociation rate constants kd (at 25 °C) are 1.7 × 103 and ca. 1.5 × 102 s−1, respectively.128

Figure 5.5  Schematic  representation of the dissociative mechanism for ligand substitution of [Cu(OH2)6]2+ with subsequent Jahn–Teller inversion. The large circle represents the first layer of solvent molecules surrounding the inner coordination sphere.127b

Chapter 5

296 2+

In [Cu(L)(OH2)2] , where L = N(CH2CH2NMe2)3 = Me6tren, for replacement of H2O with NCO−, Cl− or Br− the k values are in the range 1–60 s−1 at 10–35 °C; the process is believed to be Id.129 Kinetic studies in aqueous solution of the formation of 1 : 1 complexes of Cu2+ with imidazole and ammonia have been reported.130 At 25 °C, the reported kf values (dm3 mol−1 s−1) are 5.7 × 108 (NH3) and 2.0 × 108 (im), which are typical for reactions of Cu2+ aq with simple neutral ligands. The kf values for the corresponding reactions with en and bipy are 4 × 109 and 4 × 107 dm3 mol−1 s−1, respectively. The high value in the case of en is because the second amino group of en can assist the first water loss by the internal conjugate base mechanism (see Figure 5.6), whereas the rate for bipy is rather low, presumably because of the steric effect associated with the rotation about the C–C bond linking the two rings in bipy. There is no correlation between the kf values and basicities of the Data given in four neutral ligands im, NH3, en and bipy. Data given in Table 5.14131 for the formation reaction (L of denticity x)  M  OH2 6 

2

kf  L    M  OH2 6  x (L)

2

 xH2 O

where M = Cu, Ni, in aqueous solution at 25 °C show that the log(kCuf/kNif ) values in many cases are in the range 4.5–5.0, which has been considered “normal”. Since these M2+ ions, bearing identical charge, are of nearly the same size (their ionic radii for coordination number six are 0.83 Å for Ni2+ and 0.87 Å for Cu2+), it is reasonable to assume that electrostatic and steric factors will be almost the same in both cases, leading to very similar values of Kos for their outer-­sphere association with any ligand. Hence, k″f being equal to Ni KOS .kOS [see eqn (5.36b)], the ratio of kCu f /k f is also practically the same as the corresponding ratio for kos ≈ kex. This may be interpreted as strong evidence that the mechanisms are similar for these [M(OH2)6]2+, i.e. dissociative, as is known from various evidence in the case of [Ni(OH2)6]2+. Values of log(kCu/kNi) outside the above range must be rationalized by invoking a special mechanism for either the Ni(ii) complex or the Cu(ii) complex (or both). Thus, the ratio for the ethylenediamine rate constants is slightly low (but borderline). The anomalously large rate constant for the reaction of this ligand with Ni2+ has already been attributed to the operation of an internal conjugate base (ICB) mechanism (see later), and it may be argued127 that the same explanation holds with copper, but that in this case the diffusion-­controlled limit (estimated to be ca. 6 × 109 M−1 s−1) for a neutral ligand) has been reached before the full enhancement in kf has been attained. Consistent with this interpretation is the observation that for the ligand Ia the log(kCu/kNi) ratio is quite large. The ligand is less basic than en (its pKa being 9.1 compared with 10 for en) and therefore the rate enhancement by the ICB is diminished; thus the dif2+ . In this case, it was confusion limit is not exceeded in its reaction with Cuaq cluded that the second bond formation is rate determining.127 In the cases of the protonated ligands it is necessary for a proton to be lost for formation of the first chelate ring. The low KCu/kNi ratios observed in their cases is probably due to a reduced rate constant for Cu2+ resulting from a shift in the rate-­ determining step to the proton loss preceding the second bond formation.

Catalysed Reactions and Formation Reactions −1

297

−1

Table 5.14  Rate  constants (M s ) for formation of mono-­complexes of Ni(ii) and Cu(ii) from [M(OH2)6]2+ (data from ref. 131 and 132).a

Ligand H2O NH3 en enH+ (Ia) (Ia)H+ trienH22+ trienH33+ tetrenH22+ tetrenH33+ TEA (Ib)c (Ib)H+ (Ic)H+ imid phen phenH+ bipy bipyH+ (II)H+ (II)H22+ ida2− Hida− (III)H+ terpyH+ terpyH22+

Log kNi b

4.5 3.7 5.6 2.3 2.6 −0.7 1.9 ∼1.0 3.0 0.5 2.5 2.4 0.9 0.4 3.8 3.5 0.5 3.3 1.4 3.2 ∼1.0 4.9 −1.1 1.5 2.0 ∼–0.3

Log kCu b

∼9.2 8.3 9.6 5.2 8.2 3.0 6.9 4.2 7.6 5.0 7.5 7.5 4.7 4.7 8.8 7.8 k−1 (as is to be expected with a ligand forming a stable chelate), then eqn (5.40b) leads to (as in the case of a monodentate ligand)   



kf ≈ Kosk1

(5.40c)

  

If, however, k2 is not much different from k−1, then obviously kf will be less than Kosk1, and this appears to be true for the reactions of [Ni(OH2)6]2+ with malonate and other dicarboxylate ligands136 and several other cases.137 Low formation rates of chelates because of frequent first-­bond formation and rupture before the closure of the ring is called sterically controlled substitution. However, aliphatic diamines such as ethylenediamine (en) react much more rapidly (see the data in Table 5.14) and a similar effect but less so is seen with N,N-­diMeen, N,N′-­diMeen and triMeen138 and also tetraMeen.133 Rorabacher and co-­workers139 explained this by an internal conjugate base (ICB) mechanism as illustrated in Figure 5.6 with the [Ni(OH2)6]2+–en system as an example. The hydrogen-­bonded species (b) formed from species (a) will be in equilibrium with the internal conjugate base (b′):

Chapter 5

304

Figure 5.6  Operation  of internal conjugate base (ICB) mechanism in reaction of [Ni(OH2)6]2+ with ethylenediamine.

In the internal conjugate base (ICB) there will be considerable labilization of the bound H2O in the cis position that makes the unidentate complexation by the bidentate ligand, ethylenediamine [step (b) to (c)], proceed much faster than the complexation by a unidentate133 monoamine such as NH3. Once the ethylenediamine binds unidentately, ring closure is also quite fast owing to the proximity of the unbound NH2 of the ethylenediamine to the replaceable aqua ligand in the cis position and labilization of the cis-­H2O (see Chapter 4, Figure 4.3) by the bound NH2 group (step k″). Use of the methyl-­substituted ethylenediamines enabled Turan138 to obtain information on the relative importance of the steric and ICB effects in complex formation by Ni(ii) hexaaqua ion and a linear relationship was observed between the ICB effect and the pKa of the protonated ligand. However, Jordan140 critically appraised the evidence for the ICB mechanism and in particular the reason for invoking it in the reaction of Ni2+ aq with ethylenediamine and other polyamines. He concluded that at pH < 6.8

Catalysed Reactions and Formation Reactions

305

the published data can be accounted for on the basis of “normal” reaction rates (including the well-­established labilizing effect of a coordinated amino group in the intermediate having unidentately bound en) and requires no special explanation such as the ICB mechanism. However, for [Ni(OH2)6]2+ kfen  102 kfNH 3, hence the rate for unidentate binding of en must be faster than that for Ni2+–NH3 formation, which can be accounted for by the ICB mechanism. It should further be noted that hydrogen bonding interaction will strengthen the outer-­sphere association of the Ni2+ aq and en, leading to a significantly higher value of Kos, and hence a higher value of kf. 141 In the reaction of Ni2+ first bond formation is rate aq with dicarboxylates determining, but for reaction with amino acids142 ring closure (by N of NH2) is rate determining, and this is also true for reactions of Co2+ aq with malonate, malate and glycolate143 and (in contrast to corresponding reaction in water) for reaction of Ni2+ aq with malonic acid and its methyl and n-­butyl derivatives and also cyclopropane-­1,1-­dicarboxylic acid in dioxane–water (25 : 75 v/v).144 For reaction of Ni2+ aq with glycolate and lactate in aqueous medium, ring closure has been reported as the rate-­determining step,145 but for succinate the ∆S‡ value suggests that Ni2+–OH2 bond dissociation is the rate-­determining step, which is the normal behaviour of labile aqua metal ions reacting with monodentate ligands.146 For the formation of [Co(Cys)3]3− from [Co(Cys)2(OH2)2]− in alkaline aqueous medium, the second step (ring closure) is rate determining with k1/k2 ≈ 102 (30–50 °C; I = 1 M; pH 11.75), where k1 and k2 are the observed rate constants for the two steps under the experimental conditions.147 Anations of [Ti(OH2)6]2+ and of [V(OH2)6]3+ (see Table 5.18) are associative (Ia), much more so for Ti(iii) with a significant dependence of kf on the nature of the entering ligand Xn−, and this is also true in several other cases such as for Rh(iii) and Ir(iii) (see Table 4.17 in Chapter 4). Complex formation reactions of Mo(iii) are also associative.151b For the anation reaction of [M(OH2)6]3+ with Cl− at 25 °C, the kan values (M−1 s−1) are Table 5.18  Anation  rate constants for [M(OH2)6]3+ (M = Ti,148 V149). k



f [M(H2 O)6 ]3+  X n    [M(H2 O)5 X](3 n )  H2 O

[Ti(OH2)6]3+–Xn−(13 °C)

[V(OH2)6]3+–X−b (25 °C)

Xn−

kf/M−1 s−1

X−

kf/M−1 s−1

H2O NCS− ClCH2CO−2 CH3CO2− HC2O4−

8.6 × 103(kex/55.5)a 1 × 104 2.1 × 105 1.8 × 105 3.9 × 105 (10 °C)

NCS− N3− Cl− Br− HC2O4−

1.1 × 102 9 × 102 ≤3 ≤10 1.3 × 103

a

 5.5 is the molar concentration of H2O in dilute aqueous solution; kex/55.5 gives the value in 5 M−1  s−1 for comparison with the other values; kex = 4.77 × 105 s−1.150 Further evidence for associative reactions of [V(OH2)6]3+ has been furnished by Perlmutter-­ Hayman and Tapuhi.151a

b

Chapter 5

306 149

149

−3 149

149

in the sequence Fe (9.4) > Co (3) ≈ V (3) >> Mo (4.6 × 10 ) > Ru (∼10−6)151c ≈ Rh (5.5 × 10−6)151d > Cr (3 × 10−7).149 On consideration of LFAE (see Chapter 1), octahedral Mo(iii) is expected to be more inert than Cr(iii). Hence the observed much higher kan value for [Mo(OH2)6]3+ is probably due to catalysis by traces of [Mo(OH2)6]2+ present as an impurity. In contrast to [Fe(OH2)6]3+ (3d5) and [Co(OH2)6]3+ (3d6), [Ru(OH2)6]3+ (4d5) and [Rh(OH2)6]3+ (4d6) are very inert. Similarly, compared with [M(OH2)6]2+ of the 3dn series, [Ru(OH2)6]2+ (4d6) is very inert. Thus, at 25 °C, the rate constant values (M−1 s−1) for the formation of 1 : 1 complexes by [M(OH2)6]2+ with unidentate ligands are ∼106 (Fe), ∼105 (Co), ∼103 (Ni), ∼108 (Cr), ∼10−3 (Ru) (see ref. 149, data in Table 1.3). In a complex [M(L)(OH2)x], L influences the lability of the M–OH2, as the data in Table 5.19 indicate. The kex values in Table 5.19 indicate that σ-­donor amines (N-­donors) enhance the lability of the M–OH2, and this is a generally observed behaviour;158 the effect increases with increase in the number of donor N in the inner coordination sphere. Bipyridine (bipy) and terpyridine (trpy), being both σ-­donors and π-­acceptors, have hardly any effect. Obviously in these cases the effect of σ donation is more or less compensated by π back-­donation. The high lability of [Co(NCS)2(OH2)2] and [Co(NCS)3(OH2)]− is due to their tetrahedral structure. EDTA causes considerable labilization of H2O in [Cr(EDTA)(OH2)]− for replacement by X− [at 25 °C, pH ≈ 5, kf = 5.2 M−1 s−1 (X− = OAc−) and 102 M−1 s−1 (X− = N3−)];159 [Cr(NH3)5(OH2)]3+ reacts significantly more slowly.160,161 Cayley and Margerum102a reported results (Table 5.20) for the following type of formation reactions: k

f [Ni(A)(OH2 ) x ]2   L   [Ni(A)(L)(OH2 ) x  n ]  nH2O

Table 5.19  Water  exchange rate constants (kex) at 25 °C. Complex

kex/s−1a

Complex

kex/s−1

[Ni(OH2)6]2+ [Ni(OH2)5(NH3)]2+ [Ni(OH2)4(NH3)2]2+ [Ni(OH2)3(NH3)3]2+ [Ni(OH2)4(en)]2+ [Ni(OH2)2(en)2]2+ [Ni(OH2)2(trien)]2+ [Ni(OH2)4(bipy)]2+ [Ni(OH2)2(bipy)2]2+ [Ni(OH2)3(terpy)]2+ [Ni(OH2)2([12] aneN4)]2+

3.6 × 104 2.5 × 105 6.1 × 105 2.5 × 106 4.4 × 105 5.4 × 105 5.7 × 105d 4.9 × 104 6.6 × 104 5.2 × 104e 2.1 × 107f

[Co(OH2)6]2+ [Co(OH2)5(NH3)]2+ [Co(OH2)4(NH3)2]2+ [Co(OH2)4(mal)] [Co(OH2)5(NCS)]+ [Co(OH2)2(NCS)2] [Co(OH2)(NCS)3]−

2.2 × 105 1.6 × 107 6.5 × 107 2.2 × 107 9.5 × 106 ∼3 × 108 >5 × 108

a

Ref. 152 (except where indicated otherwise). Ref. 153. c Ref. 154. d Ref. 155. e Ref. 156. f Ref. 157. b

Catalysed Reactions and Formation Reactions

307

Table 5.20  Rate  constants of formation reactions illustrating the effect of stacking interactions.

Substrate complex

L

kf/dm3 mol−1  s−1 (25 °C)

[Ni(bipy)(OH2)4]2+ [Ni(phen)(OH2)4]2+

terpy NH3 bipy phen terpy bipy terpy

5.6 × 104 1.5 × 103 3.7 × 104 3.0 × 103 1.0 × 105 2.0 × 104 2.8 × 105

[Ni(terpy)(OH2)3]2+

Figure 5.7  Stacking  interaction of bipy with [Ni(phen))(OH2)4]2+. The results of reactions with [Ni(phen)(OH2)4]2+ indicate that while both NH3 and phen react at comparable rates, both bipy and terpy react much faster, especially terpy {see also the case of [Ni(terpy)(OH2)3]2+}. The authors explained this as being due to higher values of Kos [kf = Koskos, see eqn (5.36b)] for the bipy and even more for terpy, resulting from what they mentioned as a stacking interaction.102a In the reaction of bipy, the incoming bipy orients with one of its rings parallel to that of the bound phen or terpy rings, thereby permitting interaction between the π-­systems of this bipy ring and those of the bound phen or terpy, leading to a stacking interaction (see Figure 5.7) in which the other bipy ring due to permissible rotation about the C–C bond axis ( joining the two rings) orients appropriately (in a plane vertical to that of the other ring) such that its ring N is more favourably placed to bind to a coordination site of Ni vacated by loss of an equatorial H2O in a dissociative process that is rate determining. It is obvious that concurrent with this bond formation there will be a shifting of the bipy rings, making the bipy free of the stacking interaction, and in this situation the N of the freed ring of the

Chapter 5

308

bipy being in proximity to an axial aqua ligand will displace this and bind to the Ni, leading to fast ring closure. In the case of terpy as incoming ligand, it is possible that two of its rings are involved in the stacking interaction (hence stronger than in the case of bipy), leading to a higher value of Kos and hence a higher value of kf. A similar explanation holds good for the much higher rate of reaction of [Ni(terpy)(OH2)3]2+ with terpy than with bipy. Such a stacking is obviously not possible in the case of phen, having three fused rings, which does not permit rotation of one of the rings with respect to the others. Further evidence was presented by Renfrew et al.102b for enhancement of the rate of ternary complex formation in reactions of octahedral [NiIIL(OH2)x] with L′ (where both L and L′ are aromatic) due to large outer-­sphere association constants attributable to a stacking interaction between the incoming and the bound ligands. The labilization of an aqua ligand in [M(L)(OH2)x]m+ due to L is related to the electron-­donating ability of L (as measured by the Edwards nucleophilicity, En) as follows:162 M log(kML ex /k ex) = γEn M m+ where kML ex and k ex are the water exchange rate constants for [M(L)(OH2)x] n+ and [M(OH2)6] , respectively, and γ is a proportionality constant that has been estimated for Fe(iii), Cr(iii), Co(ii), Ni(ii) and (VO)2+, and is related to the “softness” parameter δ:163

γ = −5.5δ + 5.5 Values of δ have been evaluated for 26 metal ions; hence kML ex values may be predicted knowing kM ex. The effect of strong σ-­donor ligands in increasing the lability of M–OH2 is seen in the rate constant data on the entry of ox2− in the Cr(iii) complexes (Table 5.21) (see ref. 59b, p. 102). At 25 °C, the kox values for [Cr(OH2)6]3+, [Cr(NH3)5(OH2)]3+ and cis-­[Cr(bigH)2(OH2)2]3+ are in the ratio 1 : 15.5 : 165, respectively. In the reaction of a square-­planar complex of Ni(ii) of a quadridentate (2N, 2O) ligand with bidentate heterocyclic amines, forming six-­coordinate complexes, the LFER plot (log k versus log K) is a straight line with a slope significantly less than 1; this suggests an associative mechanism for the formation of these ternary complexes.164 Table 5.21  Effect  of bound σ-­donor ligands on formation rate constants. Substrate complex

105kox/dm3 mol−1 s−1 (25 °C)

[Cr(OH2)6]3+ [Cr(ox)(OH2)4]+ cis-­[Cr(ox)2(OH2)2]− [Cr(NH3)5(OH2)]3+ cis-­[Cr(bigH)2(OH2)2]3+a

4 53 14 62 660

a

bigH = H2NC(NH)–NH–C(NH)NH2  (biguanide).

Catalysed Reactions and Formation Reactions 165

309 166

The formation of the monooxalato and monomalonato complexes of Cr(iii) from [Cr(OH2)6]3+ and the protonated forms of these ligands in acidic aqueous solution occurs by a mechanism involving outer-­sphere association (due to ion-­pairing and hydrogen-­bonding interactions) between the reacting species and its transformation into the product by an essentially dissociative process, in which loss of a water molecule from the hexaaquachromium(iii) ion is only important in the transition state, with no significant bond formation by the incoming ligand. The same is true for the formation of the bis-­ complexes of these ligands from the corresponding mono-­complexes. The ∆H‡ values (24–29 kcal mol−1) for these reactions are comparable to the ∆H‡ value (26.1 kcal mol−1) for water exchange of [Cr(OH2)6]3+. A much lower ∆H‡ value (12.4 kcal mol−1) for the corresponding reaction of [Cr(OH2)6]3+ with glycine, however, suggests significant associative character of the process.167 The formation of [CrX(NH3)5]2+ complexes (X− = chloride or thiocyanate)168 and azide169 have been studied and the results are consistent with an SN1-­IP (D-­IP) mechanism. [Cr(OH)(OH2)5]2+ also reacts with N3− by a similar process. In a study of the corresponding reaction of [Cr(NH3)5(OH2)]3+ with acetate in HOAc–NaOAc buffer, both SN1 and SN1-­IP parallel reaction paths have been established.160 The formation of [Co(N3)(NH3)5]2+ from [Co(NH3)5(OH2)]3+ and azide [in HN3– N3− aqueous buffer, as for the Cr(iii) reaction] also proceeds by an SN1-­IP mechanism.161 A similar mechanism has also been suggested for the formation of [Co(glycine)(NH3)5]3+ from [Co(NH3)5(OH2)]3+ and glycine in aqueous solution.170 In the formation of [Cr(ox)2(AA)]− from cis-­[Cr(ox)2(OH2)2]− (ox = oxalate, AA = phen or bipy), the results are consistent with a dissociative mechanism for the reagent-­independent path and an associative mechanism for the reagent-­dependent path.171 In the formation of [Cr(ox)(bigH)2]+ from cis-­[Cr(bigH)2(OH2)2]3+, an SN1-­IP mechanism has been proposed.172 Similar formation of [Cr(bigH)2(AA)]3+ (AA = phen or bipy) also involves dissociative transformation of the outer-­sphere complexes.173 Other formation reactions studied by Banarjea's group include the formation of Cr(iii)–EDTA complexes from several chelate complexes of Cr(iii) (with oxalate, acetylacetonate and biguanide) and EDTA in alkaline174 media and the formation of an 175 Rh(iii)–EDTA complex from RhCl3− all of which 6 and EDTA in HCl media, have interesting mechanistic features. These reactions occur in several steps, of which only one is rate determining. It is worth noting in this connection that rates of formation reactions are often pH dependent, owing to an acid–base equilibrium preceding the rate-­ determining step in the complex formation. The acid–base equilibrium may involve the substrate or the entering ligand, or both (see below), depending on the system and reaction conditions. Anation of trans-­[Cr(L)(OH2)2]3+ (L = tet-­a, tet-­b) by SCN− exhibits the rate law kobs = (k1 + k2Ka/[H+])/(1 + Ka[H+]) corresponding to parallel paths involving [Cr(L)(OH2)2]3+ and [Cr(L)(OH) (OH2)]2+.176 Aquation of the complex having L = tet-­b is significantly faster than that of the complex having L = tet-­a.

Chapter 5

310 m+

x−

In the anation of [M(OH2)6] by the anion L of a weak acid {with pKa values such that [LH](x−1)− and [M(OH2)5(OH)](m−1)+ are both present in significant amounts in the solution under the experimental conditions}, one cannot distinguish between the following proton-­ambiguous paths 1 and 2:

The formation of [M(OH2)5L](m−x)+ may involve only reaction of [M(OH2)6]m+ with Lx− or of [M(OH)(OH2)5](m−1)+ with HL(x−1)−, and in either case the observed rate constant (when TM > TFe >> THL, and where [FeL2+] is negligible compared with TFe) is   



kobs = (k1[H+] + k2Kh)(QTFe + Kh + [H+])/(Kh + [H+])Q

  

Scheme 5.10

(5.42a)

Chapter 5

312

which on rearrangement leads to   

  

kobs(Kh + [H+])Q/(QTFe + Kh + [H+]) = k1[H+] + k2Kh

(5.42b)

A plot of the left-­hand side of eqn (5.42b) versus [H+] yielded a good straight line in each case, from the slopes and intercepts of which the k1 and k2 values at the experimental temperature and ionic strength could be evaluated from a knowledge of the Kb value (evaluated using literature data188 for ∆H‡ and ∆S‡). From the k1 and k2 values thus determined, the reverse rate constants k−1 (=Q/k1) and k−2 (=k2Kb/Q) were calculated. All these constants were determined at three temperatures (25, 35 and 45 °C), permitting the evaluation of the corresponding activation parameters ∆H‡ and ∆S‡. For both k1 and k2, a plot of ∆H‡ versus ∆S‡ is linear and this isokinetic trend suggests a similar mechanism. The ∆H‡ values for k2 are comparable to ∆H‡ for kexH2 O of [Fe(OH)(OH2)5]2+, which suggests an Id mechanism, but the ∆H‡ values for k1 are perceptibly lower than ∆H‡ for kexH2 O of [Fe(OH2)6]3+, hence the mechanism is Ia. The difference in mechanism is in agreement with the results of complex formation by other ligands.189,190 The ∆S‡ values are significantly negative for the k1 and k2 paths, which agrees with rate-­determining unidentate binding followed by a fast chelation step, as is the case for complexation of cis-­[Cr(ox)2(OH2)2]− by 1,10-­phenanthroline and 2,2′-­dipyridyl.171 Similar behaviour has been observed in the complexations of Fe(iii) with malono-­and succinodihydroxamic acids.191 In the pH range 2.7–3.9, benzohydroxamic acids LH react with [Fe(nta) (OH2)], forming [Fe(nta)(L)]−. The equilibrium constant for the formation of this complex (evaluated by spectrophotometry)187 {as in the case of formation of [Fe(L)(OH2)4]2+ from [Fe(OH2)6]3+ and LH} is 0.3 (at 30 °C, I = 1 M); the corresponding value for [Fe(L)(OH2)4]2+ is 133. Results of kinetic studies have shown187 that at 25 °C, kf for the formation of [Fe(nta)(L)]− is ca. four times that for the formation of [Fe(L)(OH2)4]2+, while kd (for the dissociation of L−) of [Fe(nta)(L)]− is ca. 5000 times that of [Fe(L)(OH2)4]2+, accounting for the considerably lower value for the equilibrium constant. Such destabilization of a ternary complex always occurs when both ligands are strong σ-­donors. However, in the case of a ternary complex in which one ligand is a strong σ-­donor whereas the other has π-­withdrawing character, stabilization of the ternary complex occurs, as in the case of [Ni(nta)(im)(OH2)]−.192 In this case, kinetic studies have shown that at 25 °C kf increases by a factor of 11.6 but k0 increases by a factor of 11.1, causing a marginal increase in stability of the ternary complex.193 The ternary complex [Ni(nta)(im)(OH2)]− has higher stability than [Ni(nta) (NH3)(OH2)]− owing to the higher formation rate for the former compared with the latter.193 The higher formation rate is due to a higher value for outer-­ sphere association.194 At 25 °C and I = 1 M, kfim / kfNH3 is 12.4, but only 1.25 for the reaction of [Ni(OH2)6]2+, whereas kdim / kdNH3 is 4.1. Another case of proton ambiguity (some other cases have already been mentioned) is in the anation of [Cr(NH3)5(OH2)]3+ with azide in HN3–N3− buffer,

Catalysed Reactions and Formation Reactions 169,195

313 169

reported by two groups. Whereas Banerjea and Sarkar were able to interpret the results quantitatively, assuming concurrent reactions of N3− with [Cr(NH3)5(OH2)]3+ and [Cr(OH)(NH3)5]2+, the other group196 interpreted the results on the basis of reactions of N3− and HN3 with [Cr(NH3)5(OH2)]3+. Their argument that under the pH conditions used in anation (pH 3.4–4.3 by Banerjea and Sarkar but 2.5–4.5 by the other group) the [Cr(OH)(NH3)5]2+ concentration is negligible, whereas that of HN3 is considerable, is not tenable since the pKa value of HN3 (4.11 at 50 °C)192 is comparable to that of [Cr(NH3)5(OH2)]3+ (4.24 at 50 °C),190 hence at pH ≈ 4 ± 0.5 the [Cr(OH)(NH3)5]2+ concentration will also be significant. Both mechanisms lead to the relation kobs/[N3−] = (A + B[H+])/(K + [H+]) where A and B are rate constants and K is the Ka of [Cr(NH3)5(OH2)]3+ in the mechanism proposed by Banerjea and Sarkar and Ka of HN3 in the mechanism proposed by the other group. The ambiguity is clearly due to the closeness of the pKa values of the substrate and the protonated ligand and is not easy to resolve. In this connection, the work of Chaudhury and Siddhanta197 is worth mentioning. They studied the reaction of cis-­[Co(en)2(OH2)2]3+ with benzoic and substituted benzoic acids in ethanol–water mixtures at pH 4.4–5.0 and concluded that an associative (A) mechanism (designated SN2) is operative on the basis of the first-­order dependence of the rate on the benzoate concentration and increase in rate with decrease in dielectric constant of the medium. They used dubious arguments to rule out a mechanism involving the formation of an ion pair as a precursor, which would also allow for the observed first-­order dependence of the rate on the benzoate concentration even if the transformation of the ion pair is actually dissociative (Id) as is indeed expected for any such cobalt(iii) complex.198,199 Also, the authors did not consider the possibility of a mechanism of the type shown in eqn (5.43), which would lead to the same results, despite the dissociative character of the rate-­determining step (k1). k1

k2 M(OH2 )3 # M   H2 O;M3  L   M  L2  k1

where M(OH2)3+ represents cis-­[Co(en)2(OH2)2]3+ and L− is the benzoate ion. Under the experimental conditions where TFe >> TL and [H+] >> [Fe3+ aq] (T is the total analytical concentration of the species given as a subscript and [FeL2+ aq] represents the equilibrium concentration of the species formed), the above reaction scheme leads to the following: k−1 >> k2[L−1],  kobs ≈ k1k2[L−]/k−1 = k[L−] kobs = k1k2[L−]/(k−1 + k2[L−]) This accounts for the observed first-­order dependence on benzoate concentration. Also, since the dielectric constant of the medium would influence k1 and k−1 in an opposite manner, the change in the gross rate constant k with

Chapter 5

314

dielectric constant would be chiefly due to an increase in k2 (a step involving reaction between oppositely charged ions) with decrease in dielectric constant. A change in the nature of L−1 will affect k2, and hence also k. The observed ∆H‡ value (109.8 kJ mol−1) for the reaction with benzoate compares favourably with the corresponding value for the reaction of [Co(NH3)5(OH2)]3+ with salicylate(1–) (∆H‡ = 129 kJ mol−1)200 and with glycine (∆H‡ = 115.9 kJ mol−1),170 where reactions occur by Id processes involving outer-­sphere precursor complexes. Hence the conclusion that the mechanism is SN2 appears untenable. The reaction of [Co(NH3)5(OH2)]3+ with various amino acids has also been interpreted on the basis of an Id mechanism involving an outer-­sphere complex.201 Complexation of U(vi) with salicylate and some substituted salicylates has been studied at pH 7.0–8.5 by stopped-­flow spectrophotometry.201 The mechanism involves binding of UO2(OH)+ species through the carboxylate of the salicylate, forming an inner-­sphere complex in a fast pre-­equilibrium followed by a slower ring closing (rate determining) through the phenolic group of the salicylate. Results of similar studies have been reported on U(vi) complexes with chromotropic acid and chromotrope 2B,202b and of Mo(vi) (at pH 7.5–8.5) and V(v) (at pH 2.5–3.5) with benzo-­, o-­methylbenzo-­ and o-­hydroxybenzohydroxamic acids203a and of B(iii) with alizarin-­S and quinalizarin203b in aqueous solution. Kinetics of the formation of the tris-­complex of Fe2+ with methyl 2-­pyridyl ketoxime in solution and its isomerization ( fac → mer), its acid-­catalysed dissociation and its oxidation to the Fe(iii) complex by S2O82− have been reported.204 An interesting situation in a complex formation process may arise in a complex mechanism showing a decrease in rate with increase in concentration of the entering ligand. This has been observed205 in complex formation by M2+ (M2+ = Co2+, Cu2+) with the quadridentate ligand N,N′-­p-­chlorophen yloxalylbis(thiosemicarbazides) LH2 (2N, 2S donor)205b in DMF medium in the presence of acetate. Here the results are consistent with the following scheme (S = DMF): K OAc

MS62   OAc # MS5 (OAc)  S k1

MS5  OAc # MS4 (OAc)  S k1

k 1

MS6 2  # MS52   S k1

Catalysed Reactions and Formation Reactions

315

k 2

Fast MS52   LH2 #MS5 (LH2 )2   ML  5S  2H k 2

Under the condition TL >> TM, the above scheme leads to   

kobs 

 k1 k2TL  k1 k2  K OAcTOAc k1  k2TL



k1 k2 TL  k 1 k 2 k 1  k2 TL

(5.44)

  

If k−1 >> TL and kʹ−1 >> kʹ2TL, then   

  

k k k k  kobs  1 2 K OAcTOAc  1 2 k 1  k1

   TL  k2 K OAc  k2 

(5.45)

This suggests an increase in rate with increase in concentration of the entering ligand (i.e. normal behaviour), actually observed for M2+ = Ni2+. However, if k2TL >> k−1 and kʹ2TL >> kʹ−1, then   



kobs

k1 k2 K OAcTOAc k 1 k 2  k2 k2  k1 K OAcTOAc  K1  TL

(5.46)

  

which suggests a decrease in rate with increase in ligand concentration, which has actually been observed for M2+ = Co2+ and Cu2+. Graphical evolutions have indicated k−2 = 0 and kʹ'−2 = 0 for the Ni2+–LH2 system, hence in this case eqn (5.46) leads to

  



 k k kobs   1 2  k 1

k1 k2    K T  TL OAc OAc  k  1  

= (k1TOAc + kʹ1)TL

(5.47)

  

However, if k−1k−2KOAcTOAc/k2 α-­alanine > l-­phenylalanine > l-­valine > l-­methionine > β-­alanine > sarcosine

This is not quite the sequence of pKa values of the amino acids. However, the spread of kf is fairly small; at 35 °C and I = 0.1 M at pH 6.0, for A = nta3− the kf values range from 1.3 M−1 s−1 for sarcosine to 3.9 M−1 s−1 for glycine, and for A = ada3− the corresponding values are 0.97 and 2.3 M−1 s−1, respectively. It is significant that the size of the chelate ring formed by the bidentate L− and bulky substituent on the reacting NH2 group have a noticeable effect on the kf value. The mechanism proposed (shown in Figure 5.8 with glycine as an example)208 involves fast unidentate chelation by the zwitterion form of the

Figure 5.8  Plausible  mechanism of the following type of reaction of amino acids with nickel(ii) complexes of nta3− and ada3−:

Catalysed Reactions and Formation Reactions

317

ligand, followed by its fast deprotonation and rate-­determining ring closure. For this particular system, it may not be unreasonable to assume that the K′ value will be comparable to the equilibrium constant for the following reaction: K

Ni  nta   CH3 COO # Ni  nta   OOCCH3 

2

At 25 °C and I = 0.1 M, K = 2.34.192 We may also reasonably assume Kʹa = Ka for glycine (Ka = 2.45 × 10−10 at 25 °C, I = 0.1 M). With these values, using eqn (5.50a) and (5.50b) we obtain k = 6 × 104 s−1, K″ = 2.9 × 104. This k value compares well with the reported kf value (4.1 × 104 M−1 s−1) for the reaction of [Ni(OH2)6)]2+ with glycinate ion,220 and also agrees with the fact that both [Ni(OH2)6)]2+ and [Ni(nta)(OH2)2]− react with NH3 at comparable rates: at 25 °C and I = 0.1 M, the 10−3kf values are 3.9 and 4.6, respectively.193 k1

2  NiA   LH# Ni  A  L   H+

The hydrogen-­bonded species (d) in the scheme in Figure 5.8 is essentially an internal conjugate base and is likely to be unreactive; a ring closure step involving the species (c) is rate determining. Similar rate-­determining ring closure has been proposed for the reactions of [Ni(OH2)2(trien)]2+ with glycine,210 sarcosine211 and ethylenediamine,212 and of [Ni(OH2)2(trien)]2+ and [Ni(OH2)2(tren)]2+ with pada.212 The scheme in Figure 5.8 leads, when TL >> TNi, to   



1/kobs = [H+]/TLkK ʹKʹa + (1 + Kʺ)/k

(5.49)

  

For the formation of [Ni(nta)(gly)]2−, the following values were evaluated (at 25 °C and I = 0.1 M) graphically using eqn (5.49):209   



kKʹKʹa = 3.41 × 105  s−1

(5.50a)

k/(1 + Kʺ) = 2.07  s−1

(5.50b)

  

and   

  

Complexation of [Co(nta)(OH2)2]− with NCS−, bipy and phen forming 1 : 1 : 1 ternary complexes has been investigated,213 and also similar reactions of [Ni(nta)(OH2)2]− with bipy and phen and reactions of [Co(AA)(H2O)4]2+ (AA = bipy, phen) with ntaH2−.213 Rate studies have also been reported for the stepwise replacement of biguanide in [Cu(bigH)2]2+ by bipy and phen (see Chapter 3). In the 1 : 1 complex formation between adenosine monophosphate and M2+ (M = Co214 or Ni215), ring closure is also the slowest step. The reactions involve stepwise coordination of the phosphate group followed by that of adenine ring nitrogen.

Chapter 5

318 216

Margerum and Rosent reported the results of kinetic studies on the formation of six-­coordinate Ni(ii) complexes of the type Ni(L)(NH3) from Ni(L) and NH3 and the reverse (dissociation) reaction (Table 5.22). The results in Table 5.22 indicate that the electron-­donating or -­withdrawing ability of the ligand L is more decisive than the overall charge, n, on the complex for its liability to inner-­sphere water replacement. Coordination of en, gly− and dien bring about significant increases in kf, the tridentate dien showing the largest effect. The influence of L no doubt arises from its ability to weaken Ni–OH2 bonds through increased M–L σ-­bonding. Both en, which coordinates through two amine nitrogens, and glycinate, bonding through one amine nitrogen and one carboxylate oxygen, produce about the same effect. However, ida2−, which binds through two carboxylate groups and one amine nitrogen, has a much smaller effect than dien on kf. A further increase in the denticity of the polyamine as in trien and tetren causes a marked decrease in kf. Thus, provided that the bound L occupies only two or three coordination positions, its ability to labilize the ligand being replaced dominates. On the other hand, if the ligand forms two or more chelate rings and is relatively inflexible, its inability to accommodate the structural changes that are part of the ligand substitution process becomes important. This effect is further illustrated in the relatively low rate of reaction of [Ni(edta)]2− with NH3, and a similar though lesser effect observed with [Ni(nta)(OH2)2]− compared with [Ni(dien)(OH2)3]2+. The unsaturated diamines phen and 5,6-­dimethylphen show anomalously low kf values for complexes of such bidentate ligands reacting with NH3. Metal–ligand π-­bonding, which is a feature of these complexes, removes Table 5.22  Rate  data (at 25 °C)216 for the systems. n

kf

n

 Ni  L   OH2  x   NH3 #  Ni  L   NH3   OH2  x 1   H2 O kd La

n

10−3kf/M−1 s−1 kd/s−1

K (=kf/kd)/M−1

H2O en dien trien tetren (4-­coordinate) gly− α,β-­Diaminopropionate(1–) ida2− nta3− N-­Hydroxyethylenediaminetriacet ate(3–) edta4− phen 5,6-­Dimethyl-­1,10-­phenanthroline

+2 +2 +2 +2 +2 +1 +1 0 −1 −1

2.8 12 46 2.1 0.08 14 2.7 2.5 4.6 0.29

2.6 66 641 86 1.5 61 36 65 26 8.3

1050 174 72 25 53 225 75 38 177 35

−2 +2 +2

0.43 1.5 2.2

20 3.8 3.9

21 398 555

a

 ien = diethylenetriamine; trien = triethylenetetramine; tetren = tetraethylenepentamine; all d other abbreviations are well known.

Catalysed Reactions and Formation Reactions

319

negative charge density from the metal ion to the ligand, and thus compensates for the labilizing effect of σ-­bonding in which negative charge is transferred from ligand to metal. Pasternack et al.217a determined, by the temperature jump method, the rates of formation and dissociation of ternary complexes in the reaction of Cu(bipy)2+ with B [B = ethylenediamine, glycinate(1–), α-­alaninatc(1–), β-­alaninate(1–)]:   

2+

k12

n+

 Cu  bipy  B   Cu  bipy    B# k21 

  

(5.51)

and compared the results with those for the binary system   

k12

 Cu2+  B# CuBn +



(5.52)

k21

  

Similar studies have been reported for the formation of some other ternary complexes of Cu(ii),217b and for the [Cu(bipy)(acac)] system (acac = acetylacetonate).218 The results of these investigations, summarized in Table 5.23 are significant and indicate that the dissociation rates are largely responsible for the relative stabilities of the different systems. For [Co(bipy)(gly)]+ also, the comparatively high stability219 compared with that of [Co(gly)2] can be attributed to a lower dissociation rate for the former (kd = 53 s−1 at 25 °C) than the latter (330 s−1), the formation rate constants being similar (kf = 1.6 × 106 and 2.0 × 106 M−1  s−1, respectively).220 For the reversible formation of the ternary complexes [Cu(bipy)L)]+ from Cu(bipy)2+ and L− (L− = (gly−, α-­ala− and acac−), the formation and dissociation rates are comparable to those for the corresponding formation and dissociation rates for CuL+. However, for [Cu(bipy)en]2+, the rate of dissociation of en (k21 = 1.4 s−1) is much closer to the rate of dissociation of en from [Cu(en)2]2+ (1.5 s−1) than that of Cu(en)2+ (0.1 s−1). This difference in the behaviour of the ternary complex formed by the uncharged ethylenediamine ligand compared with that of the monoanionic gly−, α-­ala− or acac− shows that once the second ligand is attached to Cu(ii), the influence of the charge of the complex on the dissociation rate becomes important.220,221 The rate constant for Table 5.23  Rates  of formation and dissociation of some binary, CuL, and ternary,

Cu(bipy)L, complexes of copper(ii) and their relative stabilities, at 25 °C and I = 0.1 M217a,218.

L

Δlog Ka

k1/M−1 s−1a

k2/s−1b

k12/M−1 s−1c

k21/s−1c

k12/k1

k21/k2

en gly− α-­ala− β-­ala− acac−

−1.29 −0.35 −0.26 −0.59 +0.32

3.8 × 109 4.0 × 109 1.3 × 109 2.0 × 108 9.0 × 108

0.1 22 12 11 3.2

2.0 × 109 1.6 × 109 1.0 × 109 3.4 × 108 1.1 × 109

1.4 19 10 110 1.9

0.5 0.4 0.8 1.7 1.2

14 0.9 0.8 10 0.6

a

Ref. 219. Ref. 51. c Ref. 50. b

Chapter 5

320 −

−

the addition of β-­ala (compared with α-­ala ) to [Cu(bipy)] is lower (β, 3.4 × 108 M−1 s−1; α, 1.0 × 109 M−1 s−1). This is also observed for the corresponding reactions of Cu2+ aq ion and is a manifestation of the slow rate of closure of a six-­membered ring, which becomes rate determining. However, β-­ala− dissociates at a much faster rate from [Cu(bipy)(β-­ala)]+ {kd, 110 s−1; only 10 s−1 for [Cu(bipy)(α-­ala)]+}. This may be the combined effect of a bulky bipy and a six-­membered ring, inducing sufficient strain in the Cu(ii)–ligand bonds, leading to rapid loss of the least strongly bound ligand. For Co(gly)+ also,220 the formation rate is similar (kf = 1.5 × 106 M−1 s−1) and its dissociation rate (kd = 34 s−1) is more comparable to that for [Co(bipy)(gly)]+. After correcting for the usual statistical effects and further correcting the forward rates for the effect of charge on the substrate on the formation of the precursor outer-­ sphere complex, the rate constants for water loss from [Co(OH2)6]2+, [Co(gly) (OH2)4]+ and [Co(bipy)(OH2)4]2+, calculated on the basis of the Eigen–Wilkins mechanism, are 8 × 105, 5 × 106 and 2 × 106 s−1, respectively.220 Since charge type does not correlate well with water labilization,206 the lower labilizing effect of bipyridine relative to glycinate is a manifestation of lower reactivity due to π-­t ype back-­bonding from metal to a coordinated unsaturated ligand. Hoffmann and Yeager222 examined the effect of various ligands coordinated to nickel(ii) on the rates of formation and dissociation of the corresponding malonato complexes. Ring closure and ring opening, respectively, appear to be the rate-­determining processes with this six-­membered chelate ring system. The rates were found to increase steadily with increase in the number of coordinated aliphatic amines; both Ni(pn)2+ (pn = 1,3-­propylenediamine) and Ni(trien)2+ react faster than Ni(dien)2+. In contrast, bipy and phen have little effect on the rates, but acac−, which is both a σ-­and a π-­donor, increases the rates. Hague and co-­workers223–226 studied the rates of formation of ternary complexes of Mg(ii), Mn(ii), Co(ii) and Zn(ii) (at 16 °C, I = 0.1 M) to elucidate the behaviour of these metal ions in enzyme systems which require them. Several systems were studied in which the entering ligands were the same and the bound spectator (inert) ligands were varied to determine the effect of previously bound ligands, particularly those having a negative charge, on the reactivity of the metal ion towards an entering negatively charged ligand. Of further interest was the testing of the hypothesis based on experimental observations that Mg2+ and Mn2+ behave similarly in biological systems, but Ca2+ behaves differently; all these have their origin in the substitution and dissociation rates of these metal ions. In a series of experiments, the reaction of 8-­hydroxyquinoline (oxnH) with − 2− Mn2+ = uranil-­N,N-­diacetate), Mn(ATP)2− (S = adeaq, Mn(nta) , Mn(uda) (uda nosine-­5 triphosphate) and Mn(P3O10)3− were studied.224 Two parallel paths involving oxnH and oxn− were observed. The value of k2 for [Mn(OH2)6]2+ is 1 × 108 M−1 s−1, which is the normal value for a rate-­determining water loss (dissociative process). The presence of a negatively charged polydentate ligand causes a large decrease in k2, particularly with the highly charged ATP4− and P3O105−; for Mn(P3O10)3−, the value of k2 is 7.9 × 105 M−1 s−1. Several factors 2+

Catalysed Reactions and Formation Reactions

321

are responsible for this decrease. One is a statistical effect that arises from the fact that the number of water molecules that are available for replacement decreases as the inert ligand (L) occupies more coordination positions. Another is the coulombic repulsion between the negatively charged complex and the approaching anionic ligand, oxn−. However, neither of these is sufficient to account for the total observed decrease, which is very large. It appears that the detailed stereochemistry of the complex influences Mn(ii) rates in a manner analogous to that known for Ni(ii) reactions.216 The rates of decomposition of these ternary Mn(ii) complexes lie in the relatively narrow range 35–210 s−1 [Mn(nta)(oxn)2− 90, Mn(uda)(oxn)− 210, Mn(ATP)(oxn)3− 35 and Mn(P3O10)(oxn)4− 80 s−1, all at 16 °C, I = 0.1 M], and are similar to the rate constant for the dissociation of the binary complex, for which the kd value is 140 s−1: kd [Mn(oxn)(OH2 )4 ]  2H2 O   [Mn(OH2 )6 ]2   oxn 

For the reactions of the protonated ligand oxnH, the k1 values lie in the range 105–106 M−1 s−1. Hague and Zetter224 concluded that solvent loss from the inner coordination sphere of Mn(ii) is not rate determining, but the slow step involves deprotonation of oxnH after its unidentate binding to Mn(ii) in a fast pre-­equilibrium followed by a much faster ring closure. Similar conclusions were drawn from the results obtained in studies on Cu(ii) complexes.217

where species XIV represents Hoxn bound unidentately and the final product species XV has chelated oxn−. Based on this, the formation rate constant kf is equal to Kk; a reasonable K value is ∼103 M−1, which with a normal value of k (∼102 s−1) for proton transfer to solvent followed by a fast ring closure would account for the observed kf value of ∼105 M−1 s−1. Similar studies (at 25 °C, I = 0.1 M) were reported by Hague and Eigen223 on the rates of formation of LMg(oxn) ternary complexes. However, for reactions of oxn− with Mn2+ aq and 2 8 −1 −1 5 Mn Mg Mg2+ aq ions, the ratio of kf is ∼10 (kf = 1.1 × 10 M s at 16 °C; kf = 3.8 × 10 −1 −1 M s at 25 °C), and this arises primarily from the difference in the rates of water loss from the inner coordination sphere of these metal ions. However, the effect of a coordinated ligand on the rate (kf ) is much smaller in the case of Mg(ii) than it is with Mn(ii), as the data in Table 5.24 indicate for the formation of LM(oxn) from LM and oxn−. As for Mn(ii), the dissociation rates for the Mg(ii) complexes are little altered by ternary complex formation, the kd values at 25 °C, I = 0.1 M being 7, 41, 54 and 6 s−1 for [Mg(oxn)]+, [Mg(uda)(oxn)]−, [Mg(ATP)(oxn)]3− and [Mg(P3O10) (oxn)]4−, respectively. The rates of the corresponding Ca2+ complexes are too fast,224 with kf >2 × 108 M−1 s−1 and kd >105 s−1. Hence Ca2+ is kinetically considerably different from Mg(ii) and Mn(ii), which become rather similar in the highly chelated environment of a highly charged anionic ligand.

Chapter 5

322

Table 5.24  Rates  of formation of some Mn(ii) and Mg(ii) complexes of oxinate, LM(oxn).

kf/M−1 s−1a L

Mn(ii)224

Mg(ii)223

H2O uda2− ATP4− P3O105−

1.1 × 108 3.6 × 107 1.0 × 106 7.9 × 105

3.8 × 105 8.0 × 104 9.6 × 104 4.7 × 104

a

At 16 °C I = 0.1 M for Mn(ii) and at 25 °C I = 0.1 M for Mg(ii).

It has been proposed223,224 that if the catalytic pathway of a metalloenzyme reaction includes a conformational change, which can take place only when the metal is bound, then the dissociation rate constant for the metal ion– enzyme–substrate ternary complex has to be less than the turnover number (usually 102–103 s−1)¶ for the enzyme-­catalysed reaction. Both Mg(ii) and Mn(ii) would satisfy such a requirement, but Ca(ii) would not be suitable; on the other hand, Ca(ii) would be more suitable than Mg(ii) or Mn(ii), if no conformational change occurs. It is known that Mn(ii) and Mg(ii) can interchange as an enzyme activator in several cases. Cayley and Hague225 studied the formation of several ternary complexes of Zn(ii) of the type [Zn(L)(pada)] [pada = pyridine-­2-­azo-­p-­dimethylaniline; L = dien, trien, ida2−, ethylenediamine-­N,N′-­diacetate (edda2−), cysteinate (cys2−), P3O5−10] at 25 °C, I = 0.3 M. The formation rates for all these ternary complexes were kf = (0.9–15) × 106 M−1 s−1, which were similar to those of the binary complex, [Zn(pada)]2+, formed from Zn2+aq ion and pada (kf = 6.8 × 106 M−1 s−1). However, the dissociation rates for these ternary complexes [kd = (0.6–14) × 104 s−1] are much higher than that of [Zn(pada)]2+ (kd = 2.7 × 102 s−1). Cobb and Haugue226 studied the formation of [CoII(L)(pada)] complexes, analogous to some of the Zn(ii) complexes mentioned above. A behaviour parallel to that of the Zn(ii) series was observed, the chief difference being due to the difference in the water exchange rates of the two metal ions (Table 5.25). Since for the Zn(ii) complexes the LFSE is zero, the influence ¶

The metalloenzyme (EM)-­catalysed reaction of a substrate (S) forming the product (P) can be expressed asfollows:

Expressing the concentration of total EM as [EM]0: d[P]/dt = k3[EM]0[S]/{(k2 + k3)/k1 + [S]} = k3[EM]0[S]/(K + [S]) where K = (k2 + k3)/k1. In the presence of excess S, it may be that K ​1.0 × 105 5.6 × 103

of the LFSE on the rates of the Co(ii) complexes do not appear to be very significant. [Cu(bigH)2]2+ [bigH = H2N–C(=NH)–NH–C(=NH)–NH2] is a typical square-­ planar complex, which reacts with bipy and phen forming [Cu(L)2]2+ (L = bipy, phen) through formation of intermediate [Cu(bigH)(L)]2+. These reactions (each of the steps) have been studied (see Chapter 3, Section 3.5.7), as also the formation of [Cu(bipy)(phen)]2+ in the reaction of [Cu(bigH) (bipy)]2+ with phen and of [Cu(bigH)(phen)]2+ with bipy. For each of these reactions, kf = k0 + k1[L], where k0 and k1 correspond to entering ligand (L)-­ independent and entering ligand-­dependent paths, respectively. At 25 °C and I = 0.1 M, the k0 values are in the range 0.3–7 s−1 and the k1 values are in the range 0.3 × 103–1.7 × 103 M−1 s−1. Results have shown that compared with bipy, the greater labilization of bigH bound to Cu(ii) is caused by phen (see Chapter 3, Section 3.5). For the following reaction:   

  

[Cu(bipy)2]2+ + pada ⇌ [Cu(bipy)(pada)]2+ + bipy

(5.53)

kf = 107 M−1 s−1.226 For the corresponding reaction of [Cu(en)2]2+ and [Cu(gly)2], the kf values are 1.3 × 105 and 5 × 105 M−1 s−1, respectively. The high value in the case of [Cu(bipy)2]2+ has been attributed to the fact that [Cu(bipy)2]2+ in solution is not square planar, but owing to steric interference the bipyridyl molecules occupy non-­planar octahedral edges with two water molecules occupying the remaining cis positions. Presumably, a similar non-­square-­planar geometry holds for [Cu(bipy)(pada)]2+ also. Attack of pada at either of the sites occupied by water molecules followed by expulsion of a bipyridine molecule would allow ligand replacement to occur without requiring extensive ligand reorganization, which is necessary with square-­planar complexes.

Chapter 5

324

In NH4NO3–NH4OH-­buffered 10% v/v dioxane–water medium (pH 7.0– 8.5), the pseudo-­first-­order rate constant for the formation of the complexes M(baen), i.e. ML (M = Cu2+, Ni2+) conforms to the equation 1/kobs = 1/k + 1/(kKosTL) where TL is the total ligand concentration in the solution, Kos is the equilibrium constant for the formation of an intermediate outer-­sphere complex and k is the rate constant for the formation of the complex ML from the intermediate.22a Under the experimental conditions, the free ligand (pKa >14) exists virtually exclusively in the undissociated form (baen H2 or LH2) which is present mostly as a ketoamine in the internally hydrogen-­bonded state [baen H2 = bis(acetylacetone)ethylenediamine]. Although the observed formation rate ratio kCa/kNi is of the order of 105, as expected for systems having “normal” behaviour (loc. cit.), the individual rate constants are very low (at 25 °C, kCu = 50 s−1 and i = 4.7 × 10−4 s−1) owing to the highly negative ∆S‡ values (−84.2 ± 3.3 J K−1 M−1 for CuL and −105.8 ± 4.1 J K−1 M−1 for NiL); the much slower rate of formation of the nickel(ii) complex is due to the higher ∆H‡ value (41.2 ± 1.0 kJ M−1 for CuL and 78.2 ± 1.2 kJ M−1 for NiL) and more negative ∆S‡ value compared with that of CuL. The Kos values [10.8 M−1 for Cu(ii) and 10.9 M−1 for Ni(ii) at 25 °C] are much higher than expected for simple outer-­sphere association between [M(OH2)6]2+ and LH2 and may be due to hydrogen bonding interaction. For M2+ incorporation in meso-­tetra(p-­sulfonatophenyl)porphyrin (TPPS), the rates are significantly slower, by several orders of magnitude, than the corresponding kex values for the [M(OH2)6]2+ ions; at 25 °C the kf/kex values227 are 5 × 10−10 (Mn2+), 5.3 × 10−8 (Fe2+), 8.9 × 10−8 (Co2+), 3.1 × 10−7 (Ni2+) and 3.7 × 10−9 (Cu2+); the kf values at 25 °C are 7.45 (Cu2+), 1.7(Fe2+), 1.6(Co2+), 0.016(Mn2+) and 0.01(Ni2+) dm3 mol−1 s−1. In the reaction of [Co(OH2)6]2+ with dimethylglyoxime (dmgH2)228a at pH ≈ 4.4–5.0 in 20% v/v ethanol–water solution, a strong signal appears at 460 nm on the stopped-­flow time scale due to the formation of a brown species that decays slowly owing to oxidation of the cobalt(ii) complex into the yellow cobalt(iii) species, which is trans-­[Co(dmgH)2(OH2)2]+ since its absorption spectrum is identical with that of the completely aquated trans-­ [Co(dmgH)2(NCS)2]−.228b Formation of the cobalt(ii) complex exhibits a complex order in dmgH2, at fixed pH, and retardation of the rate due to a decrease in pH in conformity with the following mechanism:   



  

k

Co2   dmgH2 #[Co(dmgH)]  H

(5.54a)

k [Co(dmgH)]  dmgH2  [Co(dmgH)2 ]  H

(5.54b)

and   

  

When TdmgH2 >> TCo, this scheme leads to the relationship   



  

 kobs kK  TdmgH2  /  [H ]  K TdmgH2  2

(5.54c)

Catalysed Reactions and Formation Reactions

325

which leads to TdmgH2 / k 1 / k  [H ] / kKTdmgH2 obs

and this allows the evaluation of k and K: at 25 °C (I = 0.1 M NaClO4), K = 2.96 × 10−4 and k = 2 × 104 s−1 (∆H‡ = 50.7 kJ mol−1 and ∆S‡ = 6.0 J K−1 mol−1).228a In the reaction of [Cu(OH2)6]2+ with dmgH2 in solution leading to a product of 1 : 2 (Cu : dmgH2) stoichiometry, stopped-­flow measurements at a wavelength corresponding to the absorbance maximum of the product, where the reactants have almost no absorbance, indicated instantaneous formation of a highly coloured species (too fast to measure on the stopped-­flow time scale), which then decays to the product at a measurable rate independent of the concentration of dmgH2 and the pH (4.0–5.2). The decay is presumably due to an intramolecular electron transfer from the ligand system to the copper(ii) of the initial product, forming a copper(i) complex, [Cu(dmgH)2], in which the (dmgH)−2 moiety is a univalent radical ion with delocalized charge.228a 2+ In the reaction of Feaq in the presence of hydroxylamine [to keep the iron in the iron(ii) state] at pH 5.5 (NaOAc–HOAc buffer) in 20% v/v ethanol–water medium (I = 0.1 M, adjusted with NaClO4) with an excess of dimethylglyoxime, formation of a coloured iron(ii)–dmgH complex (λmax 460 nm) occurs in two consecutive steps.228a The first step exhibits a complex order in dmgH2 whereas the second slow step is first order in dmgH2 at any fixed pH. The fast step is presumably similar to that of the reaction of Co2+ aq with dmgH2, leading to [Fe(dmgH)2], followed by a slow step, presumably following the reaction   



  

kslow [Fe(dmgH)2 ]  dmgH2   [Fe(dmgH)3 ]  H+

(5.55)

A detailed mechanism was discussed and values of the specific rate constants were reported with corresponding ∆H‡ and ∆S‡ values from studies at pH 5.4–5.9 and different concentrations of dmgH2 in solution. The ∆V‡ values for complex formation by Fe3+ with NCS− and other ligands in H2O, DMF and DMSO suggest a gradual change in mechanism from Ia to Id as the sizes of the coordinated solvent molecules and the entering ligand increase.229 Rate constants for the formation of 1 : 1 Ln3+–tartrate complexes have been reported230a for Ln = La, Nd, Sm, Gd, Dy, Ho, Tm and Lu. The kf values at 25 °C range from 5.6 × 108 M−1 s−1 for Sm3+ to 1.9 × 108 M−1 s−1 for La3+. The kf/Kos value was found to be about two orders of magnitude lower than kex, which according to the authors is because ring closure is the rate-­determining step. The kd values were in the range from 10.8 × 103 s−1 (La3+) to 1.5 × 103 s−1 (Tm3+). Complex formation of Ce(iv) with 8-­hydroxyquinoline and derivatives (S) and a subsequent slow intramolecular redox process leading to the formation of Ce(iii) and oxidation products of S have been reported.230b Another class of reactions of metal complexes are metal exchange and ligand exchange reactions, a few cases of which have already been mentioned, and solvent exchange of solvated metal ions, which was mentioned fairly elaborately in Chapter 4. A few more typical examples of these are considered below.

Chapter 5

326

The exchange of Ln (Ln = La, Pr, Sm, Gd, Er) with Ce in [Ce(edta)]− was studied at pH 6–7; reaction occurs through formation of the bridged species [Ce(edta)Ln]2+ formed by coordination of Ln3+ with one of the CO2− groups of the coordinated edta4−.231 For the exchange of Ln3+ (Ln = Ce, Eu, Tb, Tm) with the respective [Ln(dtpa)]− in acetate buffer at pH 4.2–6.8, the following rate law was reported (M = Ln3+):232 3+

3+

Rate = (k1[M] + k2[H+][M] + k3[H+] + k4[H+]2 + k5[H+]3)[M(dtpa)]2− For water exchange of [Ln(OH2)8]3+ (Ln = Pr–Yb), kex ≈ 107–108 s−1 at 25 °C.233a Ligand exchange in [Ln(nta)2]3− with labelled free nta3− in solution was studied by 1H NMR spectroscopy as a function of pH for Ln = La, Ce, Pr, Nd, Sm, Eu, Tm, Yb, Lu and also Y. The kinetic path values of k−3 are nearly constant for all the Ln, being 2 × 10−3–6 × 10−3 M−1 s−1, which are about 104 times smaller than k−1. An Id mechanism was proposed for the pathway for k−1. The pathway for (b) involves progressive chelation of the entering ligand and concurrent dechelation of the leaving ligand (an interchange process), and the pathway for (c) may either involve an Id mechanism or a slow H+ migration from the partially bonded Hnta2−.233b k1

(a) M  nta 23 # M  nta   nta  k1 k2

(b) M  nta 2  *nta # M  nta  *nta   nta3 3

3

k2

k3

(c) M  nta 23  H #M  nta   Hnta 2  k3

For ligand exchange of [M(acac)3] in acetylacetone, kex decreases in the sequence M = Co > Cr > Ru > Rh and the activation parameters (see Table 5.26) suggest a dissociative mechanism for [Co(acac)3] and an associative mechanism for the others.234a A similar reactivity order (Co > Cr > Ru ≈ Rh) was reported for the racemization of these M(acac)3 in chlorobenzene.234b The rate law for fluoride exchange of [TaF6]− in HF is235 Rate = (k1 + k2[F−])[TaF−6] Table 5.26  Activation  parameters for ligand exchange of [M(acac)3] at 25 °C. M

∆H‡/kJ mol−1

∆S‡/J K−1 mol−1

Co Cr Ru Rh

152.2 120.4 114.5 118.3

80.3 −20.3 −61.0 −79.4

Table 5.27  Kinetic  parameters for fluoride exchange in hexafluorotantalate(v). Parameter k (25 °C) ∆H‡/kcal mol−1 ∆S‡/cal K−1 mol−1

k1 path

k2 path 2

4.9 × 10 s 1.8 ± 0.3 −40 ± 4

−1

2.9 × 104 M−1 s−1 2.4 ± 0.5 −31 ± 4

Catalysed Reactions and Formation Reactions

327

Such a two-­term rate law is normal for square-­planar complexes (see Chapter 3) but is unusual for octahedral complexes. Values of the rate constants and the reported activation parameters are given in Table 5.27. The high negative value of ∆S‡ for k2 is consistent with an associative mechanism, obviously involving TaF27− as the species in the transition state. However, the large negative ∆S‡ value for k1 is surprising for a dissociative process. Obviously, the k1 path cannot be a simple dissociative process; the authors suggested an HF-­dependent associative mechanism. The ∆H‡ values for both k1 and k2 are remarkably lower. Complex formation by s-­and p-­block elements have also been reported. A few cases involving Be2+ and B(iii), which form tetrahedral complexes, were mentioned in Chapter 3. The cases of some of the post-­transition elements, particularly Zn, were mentioned earlier in this chapter and also in Chapters 2 1 and 4. Complex formation by M3+ aq (M = Al, Ga, In) with SO 4− has kf values of 3 4 5 −1 −1 1.2 × 10 (Al), 2.1 × 10 (Ga) and 2.6 × 10 M s (In) at 25 °C.236 Based on the 2 Eigen–Wilkins mechanism,100 kf/Kos = kos. For M3+ aq·SO −4 the theoretically cal3 −1 culated value of Kos (see Chapter 2) is 1.7 × 10 M . The kex values are 16 (Al), 7.6 × 102 (Ga) and 4.0 × 104 s−1 (In). Hence the kos values (at 25 °C) are 0.7 (Al), 12.4 (Ga) and 153 s−1 (In), which are very much lower than the corresponding kex values. This is surely due to the weak nucleophilicity of SO24−, which has to compete unfavourably with H2O, which is present at an enormously high concentration. In the reaction of aluminium ion in aqueous solution (pH ≈ 2.6–3.2) with ferron (8-­hydroxy-­7-­iodoquinoline-­5-­sulfonic acid) forming a 1 : 1 coloured complex, studied by the stopped-­flow method, the major pathway involves Al(OH)+2aq; assuming a normal Id mechanism237 (Eigen–Wilkins model), the water exchange of this species is (4.0 ± 0.4)×107 s−1. For such non-­transition metal ions, the reactions are predominantly dissociative, as indicated by the significantly positive ∆V‡ values for ligand exchange in [AlL6]3+ and [GaL6]3+ complexes (L = DMSO or DMF) in CD3NO2 solution.238 Factors that affect ligand exchange in such non-­transition metal ion complexes have been discussed.239 Studies on the kinetics of complex formation by Al(iii) with chromotropic acid and 3-­hydroxypicolinic acid using the stopped-­flow method have been reported.240 In the former case, the reaction appeared to involve Al(OH)2+ and the undissociated ligand (kf = 3.8 × 102 M−1 s−1 at 25 °C). In the latter, there is evidence for participation of Al3+ and Al(OH)2+. The aqueous Al(iii) propionate system was studied by the pressure jump technique.241 Two relaxations, attributed to the formation of the mono-­and bis-­complexes, were observed. Formation of the mono-­complex of Ga(iii) with 5-­nitrosalicylate in aqueous solution involves two pairs of “proton-­ambiguous” paths.242 Attributing the observed rate constant exclusively to the two reactions involving the monoprotonated ligand with Ga3+ and [Ga(OH)]2+ led to upper limits for the rate constants of 4.1 × 102 and 6.4 × 103 M−1 s−1, respectively. Comparison with other results led to the suggestion that the reactions involve an associative mechanism. Kinetics of complex formation by Ga3+ aq with 8-­hydroxyquinoline243 and 4-­(2-­pyridylazo)resorcinol244 have been reported; the reaction involves Ga(OH)2+ aq as the active species.

Chapter 5

328 245a

Kinetic data have been reported for complex formation by Mg2+ aq with 3 inorganic phosphates, HP2O7− and HP3O410− , and also nucleotide phosphates, ADP3−, CDP3−, ATP4− and CTP4−, by temperature jump spectrophotometry. The aim of this investigation was to determine the effects of the phosphate backbone and the two different ring systems (purine and pyrimidine) on the complexation mechanism. In the case of the inorganic phosphates, the experimental results were consistent with the following mechanism (for clarity, charges have been omitted for all species except H+): k1

kf

KM

H2 L#HL  H+ ;HL  M#MLH#ML  H+ K in

kd

+

HIn#H  In



All proton transfer reactions including that of the pH indicator HIn are assumed (reasonably) as rapid pre-­equilibria. A similar mechanism was 2+ 245b reported earlier for the reactions of Co2+ aq and Ni aq with these phosphates. The results were analysed in terms of the usual mechanism for complex formation by labile metal ions, the rate-­determining step being the dissociation of an Mg − OH2 bond. For the Mg2+–nucleotide phosphate systems the experimental results were consistent with the following mechanistic scheme in which 1 : 1 complex formation was coupled with the formation of a 2 : 1 complex: ka

kf

kf

kd

kd

HL#H+  L; M  L#ML; M  ML#M2 L Kin

HIn#H+  In 

The rate-­determining step in the formation of both the 1 : 1 and 2 : 1 complexes appears to be water loss from the inner hydration shell of the metal ion, and the adenine and cytosine nucleotides of a given charge type exhibit the same kinetic behaviour. Values of the evaluated rate constants are given in Table 5.28.

Table 5.28  Rate  constants for the formation and dissociation of complexes in the Mg2+–phosphate system [data at 15 °C, I = 1 M (KNO3)].245a.

Phosphate

kf/M−1 s−1

kd/s−1

kʹf/M−1 s−1

kʹd/s−1

HP2O73− HP3O104− ADP3− CDP3− ATP4− CTP4−

3.85 × 106 8.5 × 106 3.8 × 106 3.8 × 106 8.7 × 107 8.7 × 106

2.55 × 103 8.5 × 102 2.3 × 103 2.3 × 103 7.8 × 102 7.8 × 102

1.0 × 105 1.0 × 105 6.0 × 105 6.0 × 105

1.0 × 104 2.0 × 104 1.0 × 104 1.0 × 104

Catalysed Reactions and Formation Reactions

329

Values of the water-­exchange rate constant kex calculated from k1 and the computed outer-­sphere formation constant (modified to take into account the known tendency of K+, the cation of the supporting electrolyte, to complex with polyphosphates) are in good agreement with the experimentally determined245c value of ∼105 s−1. A shock-­wave apparatus was used246 to follow the reaction of Mg2+ with malonate and tartrate in water. The results were analysed in terms of the following three-­step process. K OS

k1

k2

k1

k2

2 Mg 2+ aq  L # Mg  OH 2 L #Mg  L# MgL 

It was suggested that in both cases the ring-­closure step (k2) is significantly slower than that involving dissociation of the singly bound intermediate Mg−L (k−1). The evaluated values of k2 for malonate and tartrate at 20 °C are calculated as 1.6 × 105 and 1.8 × 105 s−1, respectively, and that of k−1 is 6 × 105 s−1 for both (assuming k1 to be 3 × 105 s−1for both); values of k−2 are estimated to be 3 × 104 and 8 × 103 s−1. Kinetic studies of the complexation of Zn2+ with glycine have been reported.247 An ultrasonic relaxation method was used to study the kinetics of formation of lead acetate in aqueous solution. The formation (from Pb2+ aq and CH3CO−2) and dissociation rate constants are 7.5 × 109 M−1 s−1 and 5.7 × 107 3− s−1, respectively.248 Studies by NMR spectroscopy249 of the Pb2+ system aq–nta 6 11 −1 −1 yielded 4.4 × 10 and 1.3 × 105 (1.6 ± 0.5)×104 (1.8 ± 0.4)×104 (4.7 ± 0.6)×104 >1.1 × 105 (2.5 ± 0.3)×104

This enables the rate constant for the exchange process to be determined. The rate constant for the complexation reaction is estimated to be 6 × 107 mol−1 s−1 at 25 °C, with an activation energy of 6.5 kcal mol−1. NMR spectroscopy has also been used252 to measure the kinetics of the binding of five ionophores to K+ in 80 : 20 v/v methanol–chloroform. The results suggest that the rate of decomplexation at the water/membrane interface is not generally the rate-­determining step in the transport reaction sequence. Similar results have been reported for the reaction between Na+ and several macrocycles (biological carriers) in methanol.253 The kinetics of the binding of dibenzo-­30-­crown-­10 with univalent cations in methanol have been studied (data reported in Table 5.29).254 The formation rate constants (kf ) follow the order K+ < Rb+ ≈ Cs+ ≈ Tl+. Interestingly, K+ has the highest hydration energy among these ions. The high rate constants suggest that a stepwise mechanism is involved.

Kinetic studies have been reported for the formation of the Na+ complexes of dibenzo-­18-­crown-­6 (DBC) in DMF, methanol and dimethoxyethane255 of dicyclohexyl-­18-­crown-­6 (DCC)255 and valinomycin256 in methanol, and of compound XVI257 in ethylenediamine. In all three studies, use was made of the difference between the nuclear relaxation rates of 23Na in the free and complexed forms of the metal; this is larger than that observed with the 1H signal.

Catalysed Reactions and Formation Reactions

331

Table 5.30  Rate  and equilibrium parameters for the formation of alkali metal cryptates in methanol at 25 °C.266

+

M

+

Li

Na+ K+ Rb+

a

Cryptate

   

(2,2,1) (2B ,2,2) (2,2,1) (2B ,2,2) (2,2,1) (2B ,2,2) (2,2,1) (2B ,2,2)

10−7kf/ M−1 s−1

ΔH‡f/kJ ΔS‡f/J K−1 mol−1 mol−1 kd/s−1

ΔH‡d/kJ mol−1

ΔS‡f/J K−1 mol−1 Log K

1.88

13.3

−61

78.4

23.8

−129

5.38

3.3 8.74

— 15.3

— −42

2.1 × 105 0.0196

— 64.6

— −61

2.19 9.65

8.78 33.6

— 10.0

— −48

2.78 0.969

55.1 70.0

−52 −10

7.50 8.54

25.7 30.2

— −0.1

— −83

0.158 60.0

76.8 56.3

−3 −22

9.21 6.7

31.5

—

—

20.4

70.1

+7

7.19

a

The subscript B indicates that a benzene ring has been added between the two oxygen atoms in one of the bridges.

With valinomycin, the dissociation rate constant (25 °C) was in the range (1–3) × 105 s−1 and the activation energy was 39.75 kJ mol−1. Similar high activation energies were found for the dissociation of Na+ from the other ligands (ranging from 34.7 kJ mol−1 for DCC in methanol to 52.7 kJ mol−1 for DBC in DMF) and are presumably the energies associated with the configurational changes that occur during decomplexation. Various aspects of ligand exchange and replacement reactions of labile transition metal complexes have been reviewed.258–261 Studies in various non-­ aqueous solvents (S) and also mixed (H2O + S) solvents provide useful information on such reaction mechanisms.262–264 Rate and equilibrium data for the formation of complexes of alkali metal ions with the cryptands XVII in MeOH have been reported (Table 5.30).265

Chapter 5

332

As reported in work on this type of system, the variation in equilibrium constant is reflected primarily as a variation in the dissociation rate constant kd, the formation rate constant kf remaining one to two orders of magnitude below the diffusion-­controlled value (which, in these cases, is around 3.5 × 109 dm3 mol−1 s−1). The introduction of a benzene ring into (2,2,2) influences both kf and kd for the three heavier members of the series but only kf for Li+ (Table 5.30). K+ fits most effectively into the cavity of (2B,2,2), whereas with the smaller cryptand (2,2,1) the preference is for Na+. The kinetics of dissociation of the thallium cryptates (2,2,2) Tl+ and (2,2,1) + Tl in water and 90 : 10 methanol–water were compared267 with the previously determined values for the potassium complexes; both direct and acid-­ catalysed dissociation rate parameters are reported. More such studies have been reported on the formation and dissociation (spontaneous and acid-­catalysed) of such complexes (Tables 5.31 and 5.32).268 Steric hindrance due to benzene ring in the cryptand (2B,2,2) see (XVII) increases kd but decreases kf (see Table 5.31); kd increases with increase in the donor capacity of the solvent (see Table 5.32). We shall now briefly discuss, with a few typical examples, ligand replacement processes involving replacement of unidentate ligands, other than solvent, with polydentate ligands, of replacement of polydentate ligands with unidentate ligands (other than solvent) and of a polydentate ligand with another such ligand. Kinetic studies on the formation of [RhCl(edta)]2− in the reaction of [RhCl6]3− with edta in aqueous HCl–NaCl solution have been reported269 From the observed dependence of rate on HCl and edta concentration, it follows that   

kobs = kʹ[H2edta2−] + kʺ[H3edta−]



(5.56)

  

Table 5.31  Rate  constants for formation and uncatalysed and acid-­catalysed dissociation of some cryptates of alkali metals at 25 °C268a–c.

Cryptate

Solventa

kf/M−1 s−1

kd/s−1

kd(H)/M−1 s−1

[Na(2B,2B,2)]+ [Na(2,2,2)]+ [K(2,2,2)]+ [Na(2B,2,2)]+ [K(2B,2,2)]+ [K(2B,2B,2)]+ [Rb(2,2,2)]+ [Rb(2B,2,2)]+ [Rb(2B,2B,2)]+ [Na(2B,2B,2)]+ [K(2B,2B,2)]+ [Rb(2B,2B,2)]+

A A A A A A A A A B B B

≤109 — 4.5 × 108 — 5.8 × 107 2.9 × 107 1.8 × 108 1.3 × 108 8.0 × 107 4.9 × 107 1.5 × 108 1.1 × 108

≤0.2 — 3 × 10−3 — 5.7 × 10−3 2.0 × 10−2 0.17 3.32 18.8 1.2 0.27 1.3 × 102

5 × 102 3.3 × 104 15 6.7 × 103 17 1.4 24 1.3 × 102 ≤5 — — —

a

Solvent: A, propylene carbonate; B, methanol.

Catalysed Reactions and Formation Reactions

333

−1

Table 5.32  Rates  of dissociation (s ) of some metal cryptates in various solvents at 25 °C268d.

Solvent Cryptate

EtOH

DMSO −4

+

6.0 × 10 0.71 — ∼13 2.6 × 10−3 0.14 11 ∼2 × 103 — 0.3 4.1 × 10−3 9.2 × 10−2 —

[Li(2,1,1)] [Na(2,1,1)]+ [Ca(2,1,1)]2+ [Li(2,2,1)]+ [Na(2,2,1)]+ [K(2,2,1)]+ [Rb(2,2,1)]+ [Cs(2,2,1)]+ [Ca(2,2,1)]2+ [Na(2,2,2)]+ [K(2,2,2)]+ [Rb(2,2,2)]+ [Ca(2,2,2)]2+

N-­Methylpropionamide

DMF −2

2.1 × 10 ∼5 — — 0.75 — — — — — 2.7 — —

−2

1.4 × 10 — ∼0.2 — 0.25 ∼2.6 — — 8 × 10−4 — 0.4 — 4.4 × 10−2

4.8 × 10−3 0.47 — — 0.17 1.35 — — — 5.7 0.13 0.5 —

Table 5.33  Kinetic  parameters for the formation of the Rh(iii)–edta complex (see text).

Ionic strength

102k1K1/M−1 s−1

102k2K2/M−1 s−1

1 2 ΔH‡/kcal mol−1 ΔS‡/cal K−1 mol−1

8.4 2.83 21.2 −5.8

1.0 0.43 9.2 −43

Parallel reactions involving transformation of the outer-­sphere species [RhCl6]3−.H2edta2− and [RhCl6]3−.H3edta− has been proposed. This suggests k′ = k1K1 and k″ = k2K2, where K1 and K2 are the outer-­sphere formation constants for the two outer-­sphere complexes. Hydrogen bonding interaction is obviously of importance in the outer-­sphere association between the anionic species. Expressing [H2edta2−] and [H3edta−] in terms of the total edta and total HCl concentrations, it can be shown that   

  

kobs = k1K1([edta]T) − [HCl]T) + k2K2[HCl]T

(5.57)

Using this equation, the values of the gross constants were evaluated graphically. Values of these constants at two ionic strengths at 90 °C and ∆H‡ and ∆S‡ values are reported in Table 5.33. The relative magnitudes of the activation parameters are qualitatively in agreement with a sort of dissociative mechanism for the k1 path and an associative mechanism for the k2 path. An increase in ionic strength causes a significant decrease in these gross rate constants, which is as expected. The kinetics of reactions of nickel(ii) aminocarboxylates with CN− and of the reverse reaction can be accommodated by the following scheme (charges

Chapter 5

334

on the complexes having L have been omitted for clarity; the charge on NiL will vary with that of L):270 NiL + CN− ⇌ NiL(CN)  Fast NiL(CN) + CN− ⇌ NiL(CN)2  Fast NiL(CN)2 + CN− ⇌ NiL(CN)3 Rate-­determining NiL(CN)3 + CN− ⇌ NiL(CN)24− + L  Fast This scheme applies to complexes of a variety of monoaminocarboxylate ligands L, including ida, nta, edta, etc., and this is different from the corresponding reaction scheme for reaction of [Ni(OH2)6]2+ or [Ni(trien)]2+ with CN−, where four CN− per Ni(ii) are required in the transition state of the rate-­ determining step. The complexes [Fe(L−L)3]2+ (L–L = phen) or substituted phen271 or a Schiff base such as (Ph)(py)C = N−C6H4R (N, N donor)272 react with CN− to form [Fe(L−L)2(CN)2] initially; further reaction forming [Fe(CN)6]4− through [Fe(L−L)(CN)4]2− is a very slow process. Formation of the Schiff base complexes [Fe(L−L)2(CN)2] follows a simple second-­order rate law, but for the analogous phen complexes the rate law is Rate = (k1 + k2[CN−]) [Fe(phen)2+ 3 ] The k1 term has been assigned to rate-­determining dissociation of [Fe(phen)3]2+ to [Fe(phen)2]2+ and the k2 term to bimolecular nucleophilic attack of Fe(ii) by CN−. For the complexes of phen and of the Schiff base, the rate of reaction with CN− increases greatly with increase in mole fraction of MeOH and EtOH for the reactions carried out in H2O–MeOH and H2O– EtOH mixed solvents. Reactions of [Ni(S−S)2]2− complexes (S–S = a dithiolene ligand) mentioned in Chapter 3 are also examples of this class of reactions. Kinetic studies on the following reaction systems have been reported: [Co(mnt)2(L−L)]n− + PR3 ⇌ [Co(mnt)2(PR3)]− + L−L where L–L is en, bipy, phen (n = 1) or mnt2− (n = 3) and mnt2− is maleonitriledithiolate. Based on the experimental observations, Scheme 5.11 was proposed for the forward and reverse reactions.273 For reactions involving replacement of a polydentate ligand with another polydentate ligand, various plausible schemes have been proposed (see Scheme 5.12) based on detailed studies in several such systems.274,275 The reactions of [Ni(S−S)2]2− with another similar S–S′ ligand mentioned in Chapter 3 are also in this category. The reaction scheme (d) is a modification of (c) which can apply when two polydentate (e.g. terdentate) ligands are replaced with one polydentate (e.g. quinque-­or sexidentate) ligand.

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335

Scheme 5.11

Scheme 5.12

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127. (a) L. S. W. L. Sokol, T. D. Fink and D. B. Rorabacher, Inorg. Chem., 1980, 19, 1263; (b) D. N. Hague, Labile Metal Complexes, in Inorganic Reaction Mechanisms, Specialist Periodic Reports, The Chemical Society, London, 1976, vol. 4, Part II, ch. 4. 128. G. R. Cayley, D. Cross and P. F. Knowles, J. Chem. Soc., Chem. Commun., 1976, 837. 129. J. H. Coates, P. R. Collins and S. F. Lincoln, Aust. J. Chem., 1980, 33, 1381. 130. H. Dibler and Ph. Rosen, Ber. Bunsenges. Phys. Chem., 1972, 76, 1031. 131. D. B. Moss, C. T. Lin and D. B. Rorabacher, J. Am. Chem. Soc., 1973, 95, 5179. 132. T. S. Roche and R. G. Wilkins, J. Am. Chem. Soc., 1974, 96, 5082. 133. R. G. Wilkins, Acc. Chem. Res., 1970, 3, 408. 134. (a) R. Bustorf and T. A. Kaden, Helv. Chim. Acta, 1974, 57, 1035; (b) T. Hertill and T. A. Kaden, Helv. Chim. Acta, 1974, 57, 1328; (c) T. A. Kaden, Helv. Chim. Acta, 1971, 54, 2307; (d) D. K. Cabbiness and D. W. Margerum, J. Am. Chem. Soc., 1969, 91, 6540; (e) D. K. Cabbiness and D. W. Margerum, J. Am. Chem. Soc., 1970, 92, 2151; (f) M. Kodama and E. Kimura, J. Chem. Soc. Chem. Commun., 1975, 326, 891; (g) C. T. Lin, D. B. Rorabacher, G. R. Cayley and D. W. Margerum, Inorg. Chem., 1975, 14, 919; (h) D. H. Busch, Acc. Chem. Res., 1978, 11, 392; (i) A. P. Leugger, L. Hretli and T. A. Kaden, Helv. Chim. Acta, 1978, 61, 2296; ( j) T. J. Riedo and T. H. Kaden, Helv. Chim. Acta, 1979, 62, 1089; (k) L. L. Diaddario, L. L. Zimmer, T. E. Jones, L. S. W. L. Sokol, R. B. Cruz, E. L. Yee, L. A. Ochrymowycz and D. B. Rorabacher, J. Am. Chem. Soc., 1979, 101, 3511. 135. D. Banerjea and S. Bandynopadhyay, J. Indian Chem. Soc., 1996, 73, 153 and references cited therein. 136. H. Hoffman, Ber. Bunsenges. Phys. Chem., 1969, 73, 432. 137. (a) D. B. Rorabacher, T. S. Turan, J. A. Defever and W. G. Nickels, Inorg. Chem., 1969, 8, 1498; (b) R. G. Perarson and O. P. Anderson, Inorg. Chem., 1970, 9, 39; (c) T. A. Kaden, Helv. Chim. Acta, 1970, 53, 617. 138. T. S. Turin, Inorg. Chem., 1974, 13, 1584. 139. (a) D. B. Rorabacher, Inorg. Chem., 1966, 5, 1891; (b) R. W. Taylor, H. K. Stepien and D. B. Rorabacher, Inorg. Chem., 1974, 13, 1282. 140. R. B. Jordan, Inorg. Chem., 1976, 15, 748. 141. G. Calvaruso, F. P. Cavasino and E. Di Dio, J. Chem. Soc., Dalton, 1972, 2632. 142. J. E. Letter and R. B. Jordan, J. Am. Chem. Soc., 1975, 97, 2381. 143. S. Harada, M. Kawasawa and T. Yasunaga, Bull. Chem. Soc. Jpn., 1980, 53, 2074. 144. G. Arcoleo, F. P. Cavasino, E. Di Dio and C. Sbriziolo, J. Chem. Soc. Dalton Trans., 1980, 41. 145. T. Inoue, K. Sugahara, K. Kojima and R. Shimozawa, Inorg. Chem., 1983, 22, 3977. 146. T. Inoue, K. Kojima and R. Shimozawa, Inorg. Chem., 1983, 22, 3972.

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147. R. Majumdar, PhD thesis, Chapter 2, Calcutta University, 1995. 148. (a) H. Dibler, Z. Phys. Chem., 1969, 68, 64; (b) P. Chaudhuri and H. Dibler, J. Chem. Soc., Dalton, 1977, 596. 149. D. W. Margerum, G. R. Cayley, D. C. Weatherburn and G. K. Pagenkopf, Kinetics and Mechanisms of Complex Formation and Ligand Exchange, in Coordination Chemistry, ACS Monograph 174, ed. A. E. Martell, Washington, D.C., 1978, vol. 2; data in Table 1–4. 150. K. Thompson and A. G. Sykes, Inorg. Chem., 1979, 18, 2025. 151. (a) B. Perlmutter-­Hayman and E. Tapuhi, Inorg. Chem., 1979, 18, 2872; (b) Y. Sasaki and A. G. Sykes, J. Chem. Soc., Chem. Commun., 1973, 767; J. Less Common Metals, 1974, 36, 125. J. Chem. Soc, Dalton Trans., 1978, 1468; (c) T. W. Kallen and J. E. Earley, Inorg. Chem., 1971, 10, 1149; (d) M. J. Pavelich and G. M. Harris, Inorg. Chem., 1973, 12, 423. 152. (a) A. G. Desai, H. W. Dodgen and J. P. Hunt, J. Am. Chem. Soc., 1969, 91, 5001; 1970, 92, 798; (b) M. Grant, H. W. Dodgen and J. P. Hunt, J. Am. Chem. Soc., 1970, 92, 2321. 153. (a) R. Murray, S. F. Lincoln, H. H. Glaeser, H. W. Dodgen and J. P. Hunt, Inorg. Chem., 1969, 8, 554; (b) P. E. Hoggard, H. W. Dodgen and J. P. Hunt, Inorg. Chem., 1971, 10, 959. 154. H. A. Zeltmann and L. O. Morgan, Inorg. Chem., 1970, 9, 2522. 155. D. P. Rablen, H. W. Dodgen and J. P. Hunt, Inorg. Chem., 1976, 15, 931. 156. D. Rablen and G. Gordon, Inorg. Chem., 1969, 8, 395. 157. J. H. Coates, D. A. Hadi, S. F. Lincoln, H. W. Dodgen and J. P. Hunt, Inorg. Chem., 1981, 20, 707. 158. J. P. Hunt, Coord. Chem. Rev., 1971, 7, 1. 159. Y. Sulfab, R. S. Taylor and A. G. Sykes, Inorg. Chem., 1976, 15, 2388. 160. D. Banerjea and S. Sarkar, Z. Anorg. Allg. Chem., 1970, 373, 73. 161. T. W. Swaddle and G. Guastalla, Inorg. Chem., 1969, 8, 1604. 162. M. Tanaka and S. Yamada, J. Chem. Soc. Chem. Commun., 1976, 178. 163. S. Yamada and M. Tanaka, J. Inorg. Nucl. Chem., 1975, 37, 587. 164. M. Cusumano, J. Chem. Soc., Dalton Trans., 1980, 2456. 165. D. Banerjea and S. Dutta Chaudhuri, J. Inorg. Nucl. Chem., 1970, 32, 1617. 166. D. Banerjea and C. Chatterjee, J. Inorg. Nucl. Chem., 1969, 31, 3845. 167. D. Banerjea and S. Dutta Chaudhurim, J. Inorg. Nucl. Chem., 1968, 30, 871. 168. N. V. Duffy and J. E. Earley, J. Am. Chem. Soc., 1967, 89, 272. 169. D. Banerjea and S. Sarkar, Z. Anorg. Allg. Chem., 1972, 393, 301. 170. D. Banerjea and J. Roy, Z. Anorg. Allg. Chem., 1973, 400, 89. 171. D. Banerjea and J. Roy, Z. Anorg. Allg. Chem., 1973, 399, 115. 172. D. Banerjea and S. Sengupta, Z. Anorg. Allg. Chem., 1973, 397, 2. 173. S. Sen Gupta and D. Banerjea, Z. Anorg. Allg. Chem., 1976, 424, 93. 174. (a) D. Banerjea and P. Chaudhuri, J. Inorg. Nucl. Chem., 1968, 30, 3259; (b) D. Banerjea and P. Chaudhuri, J. Inorg. Nucl. Chem., 1970, 32, 2697; (c) D. Banerjea and P. Chaudhuri, Z. Anorg. Allg. Chem., 1970, 372, 268. 175. D. Banerjea and B. Chattopadhyay, Indian J. Chem., 1970, 8, 993.

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176. E. Madej, A. Katafias, P. Kita and O. Monsted, Eur. J. Inorg. Chem., 2006, 5098. 177. B. Perlmutter-­Hayman, F. Secco, E. Tapuhi and M. Venturini, J. Chem. Soc., Dalton Trans., 1980, 1124. 178. (a) J. N. Mandal and G. S. De, Indian J. Chem., 1979, 17A, 254; (b) Stability Constants of Metal Ion Complexes, Special Publication No. 17, complied by L. G. Sillen and A. E. Martell, The Chemical Society, London, 1964; Table 2. 179. C. Chatterjee and T. H. Kaden, Helv. Chim. Acta, 1975, 58, 1881. 180. S. Gouger and J. Stuehr, Inorg. Chem., 1974, 13, 379. 181. E. Mentasti, E. Pelizzetti, F. Secco and M. Venturini, Inorg. Chem., 1979, 18, 2007. 182. E. Mentasti, F. Secco and M. Venturi, Inorg. Chem., 1980, 19, 3258. 183. D. R. Prasad, T. Ramasami, D. Ramaswamy and M. Santappa, Inorg. Chem., 1980, 19, 3181. 184. K. R. Ashley, J. G. Leopoldt and V. K. Joshi, Inorg. Chem., 1980, 19, 1608. 185. R. Majumdar, PhD thesis, Chapter 3, Calcutta Univertsity, 1995. 186. D. Banerjea and B. Chakravarti, Inorg. Chim. Acta, 1996, 240, 117. 187. P. K. Das, S. G. Bhattacharyya, R. N. Banerjee and D. Banerjea, J. Coord. Chem., 1989, 19, 311; Proc. 25th Int. Confc. Coord., Chem., Nanjing (China), 1987, p. 272. 188. M. Milburn, J. Am. Chem. Soc., 1957, 79, 537. 189. B. B. Hasinoff, Can. J. Chem., 1979, 57, 77. 190. T. W. Swaddle and A. E. Merbach, Inorg. Chem., 1981, 20, 4212. 191. D. Banerjea and S. Bandyopadhyay, J. Indian Chem. Soc., 1994, 71, 483. 192. D. Banerjea, T. A. Kaden and H. Sigel, Inorg. Chem., 1981, 20, 2586. 193. D. Banerjea, T. A. Kaden and H. Sigel, Inorg. Chim. Acta, 1981, 56, L53. 194. D. Banerjea, Science Review (Indian Science News Association, Kolkata), 1984, 2, 43. 195. (a) R. W. Renfrew, P. Osvath and D. C. Weatherburn, Aust. J. Chem., 1980, 33, 45; (b) G. R. Cayley and D. W. Margerum, J. Chem. Soc., Chem. Commun., 1974, 1002. 196. S. Castillo-­Blum and A. G. Sykes, Inorg. Chem., 1984, 23, 1049. 197. C. Chaudhury and S. K. Siddhanta, Indian J. Chem., 1982, 21A, 1079. 198. D. A. Palmer and H. Kelm, Coord. Chem. Rev., 1981, 36, 89. 199. R. W. Hay, Coord. Chem. Rev., 1982, 41, 191. 200. B. Mohanty and A. C. Dash, Indian J. Chem., 1979, 17A, 296. 201. T. Murakami, K. Ogino, H. Kobayashi, H. Yamazaki and K. Saito, Bull. Chem. Soc. Jpn., 1971, 44, 120. 202. (a) S. Gangopadhyay, R. N. Banerjee and D. Banerjea, Transition Met. Chem., 1985, 10, 325; Proc. 23rd Int. Confc. Coord. Chem., Boulder, Colorado (USA), 1984, p. 198; (b) S. Basak and D. Banerjea, Transition Met. Chem., 1992, 17, 426. 203. (a) K. Basu and D. Banerjea, J. Indian Chem. Soc., 1993, 70, 893; (b) A. Bhowmick and D. Banerjea, Indian J. Chem., 1990, 29A, 1208. 204. T. Pramanik and D. Banerjea, Indian J. Chem., 1994, 33A, 826.

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205. (a) S. Datta Ray, PhD thesis, Chapter 2, Calcutta University, 1995; D. Banerjea, Proc. 34th Int. Confc. Coord. Chem., Edinburgh (UK), 2000, p. 1092; (b) S. S. Singh and S. S. Tiwari, J. Indian Chem. Soc., 1986, 63, 517. 206. M. Cusumano, J. Chem. Soc., Dalton Trans., 1980, 2456. 207. A. Das, S. Ganmgopadhyay and D. Banerjea, Transition Met. Chem., 1986, 11, 259. 208. A. Pandit, A. Das and D. Banerjra, Transition Met. Chem., 1988, 13, 459. 209. J. C. Cassalt and R. G. Wilkins, J. Am. Chem. Soc., 1968, 90, 6045. 210. R. K. Steinhaus and B. I. Lee, Inorg. Chem., 1982, 21, 1829. 211. R. K. Steinhaus and L. H. Kolopajlo, Inorg. Chem., 1985, 24, 1839. 212. M. A. Cobb and D. N. Hague, J. Chem. Soc., Faraday Trans. 1, 1072, 68, 932. 213. P. Chattopadhyay and D. Banerjea, Polyhedron, 1994, 13, 1981. 214. R. S. Taylor and H. Dibler, Bioinorg. Chem., 1976, 6, 257. 215. A. Peguy and H. Dibler, J. Phys. Chem., 1977, 81, 1335. 216. D. W. Margerum and H. M. Rosen, J. Am. Chem. Soc., 1967, 89, 1088. 217. (a) R. F. Pasternack, P. R. Huber, U. M. Huber and H. Sigel, Inorg. Chem., 1972, 11, 276; (b) V. S. Sharma and D. L. Leussing, Inorg. Chem., 1972, 11, 138, 1955. 218. M. Munakata and K. Yamada, Bull. Chem. Soc. Jpn., 1978, 51, 3500. 219. (a) R. Griesser and H. Sigel, Inorg. Chem., 1970, 9, 1238; (b) R. Griesser and H. Sigel, Inorg. Chem., 1971, 10, 2229. 220. R. F. Pasternack, P. R. Huber, U. M. Huber and H. Sigel, Inorg. Chem., 1972, 11, 420. 221. R. F. Pasternack and H. Sigel, J. Am. Chem. Soc., 1970, 92, 1646. 222. H. Hoffmann and E. Yeager, Ber. Bunsenges. Phys. Chem., 1970, 74, 641. 223. D. N. Hague and M. Eigen, Trans. Faraday Soc., 1966, 62, 1236. 224. D. N. Hague and M. S. Zetter, Trans. Faraday Soc., 1970, 66, 1176. 225. G. R. Cayley and D. N. Hague, Trans. Faraday Soc., 1971, 67, 786. 226. M. A. Cobb and D. N. Hague, J. Chem. Soc., Chem. Commun., 1971, 20. 227. D. Banerjea and S. Bandyopadhyay, J. Indian Chem. Soc., 1996, 73, 153. 228. (a) S. Gangopadhyay, A. Das and D. Banerjea, J. Indian Chem. Soc., 1989, 66, 517; (b) S. Das, R. N. Banerjee, S. G. Bhattacharyay and D. Banerjea, Transition Met. Chem., 1983, 8, 241. 229. K. Ishihara, S. Funbahashi and M. Tanaka, Inorg. Chem., 1983, 22, 3589. 230. (a) S. S. Yun and J. L. Bear, J. Inorg. Nucl. Chem., 1976, 38, 1041; (b) A. K. Pandit, A. Das and D. Banerjea, Transition Met. Chem., 1991, 16, 324. 231. E. Brucher and I. Banyai, Acta Chim. Acad. Sci. Hungary, 1977, 95, 147. 232. A. J. Mahmood, Chem. Abstr., 1979, 99, 93.061. 233. (a) D. T. Richens, Chem. Rev., 2005, 105, 1961, data in Table 15 and citations therein; (b) Y. Gefeller and A. E. Merbach, Inorg. Chim. Acta, 1978, 29, 217. 234. (a) H. Kido, Bull. Chem. Soc. Jpn., 1980, 53, 82; (b) R. C. Fay, A. Y. Girgis and U. Mabunde, J. Am. Chem. Soc., 1970, 92, 2056. 235. N. A. Matwiyoff, L. B. Asprey and W. E. Wageman, Inorg. Chem., 1970, 9, 2014.

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236. J. Micali and J. Stvehr, J. Am. Chem. Soc., 1968, 90, 6967. 237. J. H. Fergusson, K. Kustin and A. Phipps, Inorg. Chim. Acta, 1980, 43, 49. 238. C. Ammann, P. Moore, A. E. Merbach and C. M. McAterr, Helv. Chim. Acta, 1980, 63, 268. 239. D. L. Pisaniello and S. F. Lincoln, J. Chem. Soc., Dalton Trans., 1980, 699. 240. C. Baiocchi and E. Mentasti, Ann. Chim., 1981, 71, 631; Chem. Abstr., 1982, 96, 41636n. 241. M. Hiraishi, J. Sci. Hiroshima Univ., Ser. A-­2, 1980, 44, 311; Chem. Abstr., 1981, 94, 72234t. 242. B. Perlmutter-­Hayman, F. Secco, E. Tapuhi and M. Venturini, J. Chem. Soc., Dalton Trans., 1980, 11244. 243. T. Biver, L. Ghezzi, V. Malvaldi, F. Secco, M. Tinc and M. Venturini, J. Phys. Chem. B, 2009, 113, 1598. 244. T. Biver, G. Boggioni, F. Secco and M. Venturini, Langmuir, 2008, 24, 36. 245. (a) C. M. Frey, J. L. Banyasz and J. E. Stuehr, J. Am. Chem. Soc., 1972, 94, 9198; (b) G. Hammes and M. Morell, J. Am. Chem. Soc., 1964, 86, 1497; (c) M. Eigen and K. Tamm, Z. Elektrochem., 1962, 66, 93, 197. 246. G. Platz and H. Hoffmannm, Ber. Bunsenges. Phys. Chem., 1972, 76, 491. 247. (a) J. A. Miceli and J. E. Stuehr, Inorg. Chem., 1972, 11, 2763; (b) W. M. Grant, J. Chem. Soc., Faraday Trans. 1, 1973, 69, 560. 248. T. Yasunaga and S. Harada, Bull. Chem. Soc. Jpn., 1971, 44, 848. 249. D. L. Robenstein, J. Am. Chem. Soc., 1971, 93, 2869. 250. D. L. Robenstein and R. L. Kula, J. Am. Chem. Soc., 1969, 91, 2492. 251. E. Shchori, J. Jagur-­Grodzinski, Z. Luz and M. Shporer, J. Am. Chem. Soc., 1971, 93, 7133. 252. D. H. Haynes, FEBS Lett., 1972, 20, 221. 253. D. Baneerjea, Coordination Chemistry, Asian Books, New Delhi, 3rd. edn (2nd reprint), 2016, ch. 7, p. 599. 254. R. Winkler, Struct. Bonding, 1972, 10, 1. 255. P. B. Chock, Proc. Natl. Acad. Sci. U. S. A., 1972, 69, 1939. 256. E. Shchori, J. Jagur-­Grodzinski and M. Shporer, J. Am. Chem. Soc., 1973, 95, 7842. 257. M. Shporer, H. Zemel and Z. Luz, FEBS Lett., 1974, 40, 357. 258. M. Ceraso and J. L. Dye, J. Am. Chem. Soc., 1973, 95, 4432. 259. R. G. Wilkins, Acc. Chem. Res., 1970, 3, 408. 260. K. Kustin and J. Swinehart, Prog. Inorg. Chem., 1970, 13, 107. 261. J. P. Hunt, Coord. Chem. Rev., 1971, 7, 1. 262. D. N. Hague, Labile Metal Complexes, in Inorganic Reaction Mechanisms, Vol. 2, Specialist Peeriodical Reports, The Chemical Society, London, 1972, Part II, ch. 4. 263. W. J. MacKellar and D. B. Rorabacher, J. Am. Chem. Soc., 1971, 93, 4379. 264. Z. Luz and D. B. Rorabacher, J. Chem. Phys., 1964, 40, 1066. 265. (a) H. P. Bennetto and E. F. Caldin, Chem. Commun., 1969, 599; (b) H. P. Bennetto, J. Chem. Soc. A, 1971, 2211; (c) H. P. Bennetto and E. F. Caldin, J. Chem. Soc. A, 1971, 2190, 2198, 2207.

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266. (a) B. G. Cox, I. Schneider and H. Schneider, Ber. Bunsenges. Phys. Chem., 1980, 84, 470; (b) B. G. Cox, D. Knop and H. Schneider, J. Phys. Chem., 1980, 84, 320. 267. R. Gresser, D. W. Boyd, A. M. Albrecht-­Gary and J. P. Schwing, J. Am. Chem. Soc., 1980, 102, 651. 268. (a) B. G. Cox, J. Garcia-­Rosas and H. Schneider, Ber. Bunsenges. Phys. Chem., 1982, 86, 293; (b) B. G. Cox, J. Garcia-­Rosas and H. Schneider, J. Phys. Chem., 1980, 84, 3178; (c) B. G. Cox, I. Schneider and H. Schneider, Inorg. Chim. Acta, 1981, 49, 153; (d) B. G. Cox, J. Gracia-­Rosas and H. Schneider, J. Am. Chem. Soc., 1981, 103, 1054. 269. D. Banerjea and B. Chattopadhyay, Indian J. Chem., 1970, 8, 993. 270. L. C. Coombs, D. W. Margerum and P. C. Nigam, Inorg. Chem., 1970, 9, 2081. 271. J. Burgess, Inorg. Chim. Acta, 1971, 5, 133. 272. J. Burgess, G. E. Ellis, D. J. Evans, A. Porter, R. Wane and R. D. Wyvill, J. Chem. Soc. A, 1971, 44. 273. (a) D. A. Sweigart and D. G. DeWit, Inorg. Chem., 1970, 9, 1582; (b) D. G. DeWit, M. J. Hynes and D. A. Sweigart, Inorg. Chem., 1971, 10, 196. 274. (a) D. Banerjea and P. Chaudhuri, J. Inorg. Nucl. Chem., 1968, 30, 3259; 1970, 32, 2697; Z. Anorg. Allgm. Chem., 1970, 372, 268; (b) R. M. Countryman and H. M. N. H. Irving, J. Inorg. Nucl. Chem., 1971, 33, 1819; (c) M. Kodama, Y. Fujii and T. Ueda, Bull. Chem. Soc. Jpn., 1970, 43, 2085; (d) S. Funahashi and M. Tanaka, Inorg. Chem., 1970, 9, 2092. 275. (a) M. Kodama, S. Karasawa and T. Watanabe, Bull. Chem. Soc. Jpn., 1971, 44, 1815; (b) S. Funahashi, S. Yamada and M. Tanaka, Inorg. Chem., 1971, 10, 257; (c) P. E. Reinbold and K. H. Pearson, Inorg. Chem., 1970, 10, 2325; (d) J. D. Carr and D. R. Baker, Inorg. Chem., 1971, 10, 2249; (e) S. Funahashi, M. Tabata and M. Tanaka, Bull. Chem. Soc. Jpn., 1971, 44, 1586. 276. P. Ray and N. K. Dutt, J. Indian Chem. Soc., 1943, 20, 81.

Chapter 6

Isomerization, Optical Inversion and Racemization Reactions 6.1  Linkage Isomerization Nitrito complexes of Co(iii) having a Co–ONO linkage were first described by Jorgensen,1 who also observed that they rearrange to the corresponding nitro isomers having a Co–NO2 linkage. Usually the nitrito form is less stable than the nitro form, hence the former isomerizes to the latter not only in solution but even (although more slowly) in the solid state (more readily on heating). Under a pressure of 20 kbar, violet [Ni(ONO)2(en)2] changes to red [Ni(NO)2(en)2] at 126 °C. However, the conversion of nitro to nitrito occurs on heating at atmospheric pressure.2 An abundance of experimental evidence now exists in support of nitrito → nitro isomerization in complexes of CoIII and others such as PtIV, RhIII and IrIII, e.g. [M(ONO)(NH3)5]n+ to [M(NO2)(NH3)5]n+. The UV and visible spectra of the two isomers differ. The d–d absorption band for M–ONO is at longer wavelength than that of M–NO2 since the Dq value of the former is lower than that of the latter and the two classes generally have different colours, e.g. [Co(ONO)(NH3)5]2+ is red and [Co(NO2)(NH3)5]2+ is yellow. Examples are known in which both –ONO and –NO2 are present in the same complex,3 as in [Co(ONO)(NO2)(en)2]NO3. The two types of bonding are readily distinguishable by their IR spectra, the best diagnostic regions being 1060 cm−1, where a strong band appears due to M–ONO stretching vibrations, and a band at 820 cm−1 due to M–NO2 deformation. Adell4 carried out extensive investigations on the nitrito to nitro isomerization, viz. of [Co(ONO)(NH3)5]2+ to [Co(NO2)(NH3)5]2+ and of cis-­ and   Mechanisms of Reactions of Metal Complexes in Solution By Debabrata Banerjea and M. K. Bharty © Debabrata Banerjea and M. K. Bharty 2023 Published by the Royal Society of Chemistry, www.rsc.org

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348 +

+

trans-­[Co(ONO)2(en)2] to [Co(NO2)2(en)2] in the solid state. Basolo and Hammaker5 studied the nitrito to nitro isomerization of [M(ONO)(NH3)5]Cl2 (M = Rh, Ir) and of [Pt(ONO)(NH3)5]Cl3, and similar studies on cis-­[Co(NO2) (ONO)(en)2]+ were reported by Pearson et al.3 The isomerization rate was determined spectrophotometrically. The rates were first order in all the different cases (different metals and different complexes of the same metal) and the rates were very similar, which rules out the possibility of an intermolecular mechanism. Studies were made both in the solid state and in aqueous solution, and the reaction was much faster in solution. Since the solid salts also isomerize, it was concluded that the change occurs by an intramolecular process. Pearson et al.3 suggested that the reaction in solution also occurs by an intramolecular process, because the rates of isomerization in solutions containing no excess nitrite ion are far too great to be compatible with an intermolecular process involving dissociation of the bound ONO− and its re-­entry as NO2−, and further no uptake of labelled nitrite added to the solution takes place during isomerization. This has been proved by use of 18O as a tracer.6 It was found that neither the oxygen attached to cobalt and nitrogen nor that attached only to nitrogen exchanges with the solvent or with added nitrite ion in solution in the process of isomerization. The same was also observed with cis-­[Co(NO2) (ONO)(en)2]+. Furthermore, this optically active cation mutarotates, with no racemization, at the same rate as it isomerizes to the dinitro complex.7 An X-­ray study of thermally induced nitrito to nitro isomerization and the photochemically induced nitro to nitrito isomerization of CoIII complexes showed that both occur intramolecularly by rotation of the NO2 group in its own plane, probably via a seven-­coordinate intermediate.8 Similarly, the base-­catalysed nitrito to nitro isomerization of [M(NH3)5(ONO)]2+ (M = Co, Rh, Ir) is intramolecular and occurs without 18O exchange of the coordinated ONO− with H218O, 18OH− or free N18O2− in solution.9 These results prove that the isomerization occurs via an intramolecular process and for this an SNi (substitution, nucleophilic internal displacement) mechanism involving an intermediate in which the nitrito ligand is bonded to the M by the N and one of the O, shown in eqn (6.1), was suggested.6,7   

  



(6.1)

However, an elegant 17O NMR study using specifically labelled [Co(17ONO) (NH3)5]2+ and [Co(ON17O)(NH3)5]2+ established that intramolecular O to O exchange in the nitrito ligand occurs at a rate comparable to that of the spontaneous M–ONO to M–NO2 isomerization.10 The intermediate is therefore likely to be a π-­bonded species as shown, which can account for all the experimental observations on isomerization and oxygen exchange (scrambling) mentioned above.

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349

Scheme 6.1  An intramolecular mechanism that accounts for all the observations,10 with k1/k2 = 1.2, is shown in Scheme 6.1. A similar intramolecular mechanism operates for the linkage isomerization11 of [M(ONO)(NH3)5]2+ to [M(NO2)(NH3)5]2+ (M = RhIII, IrIII) and of [Co(SCN)(NH3)5]2+ to [Co(NCS)(NH3)5]2+.11 In keeping with this, the values of the activation volumes ΔV‡ are all negative for these systems [−6.7, −7.4 and −5.9 cm3 mol−1 for the Co(iii), Rh(iii) and Ir(iii) nitrito complexes, respectively,12 and −5.3 cm3 mol−1 for the Co(iii)–isothiocyanato complex].13 Since the molar volumes for the above nitrito– and nitro–cobalt(iii) complexes are 82.3 and 69.4 cm3 mol−1, respectively, an intramolecular mechanism, in which there is effectively both O and N bonding of the NO2− to the Co(iii) in the transition state, is expected to lead to a negative ΔV‡, as observed. In contrast, for isomerization of [(Et4dien)Pd–XCN]+ to [(Et4dien)Pd–NCX]+ (X = S or Se), the kinetic parameters are identical with those for solvolysis and also for the replacement of XCN− with Br−. This suggests an intermolecular (dissociation) mechanism [eqn (6.2) and (6.3)].14,15 If the reaction is carried out in the presence of added Br− in the solution then the Br− also competes [eqn (6.4)], hence kisomer ≈ kBr ≈ kd.   



 d Et dien  Pd  XCN   Solvent    Et 4 dien  Pd   Solvent    4   

Et dien  Pd   Solvent    4

Slow k

2

  

 Et 4 dien  Pd   Solvent  

 XCN 

(6.2)



Fast  XCN    Et 4 dien  Pd  NCX   Solvent

(6.3)

  



2

2



Fast  Br    Et 4 dien  Pd  Br   Solvent

(6.4)

This is further supported by the fact that the N-­bonded isomer reacts with Br− at a slower rate than does the S-­bonded isomer. Since, unlike NO2−, the group SCN− is linear and a prohibitive amount of energy may be required to bend it, as is necessary for the operation of the SNi mechanism, the intermolecular mechanism is more likely in the thiocyanato to isothiocyanato isomerization. For the isomerization of [(H2O)5Cr–SCN]2+ to [(H2O)5Cr–NCS]2+ also, the rate of isomerization is equal to the rate of aquation and the reaction must be by an intermolecular mechanism believed to involve an ion-­ pair intermediate, in which the SCN− after its dissociation from the metal remains closely associated with the Cr(iii) species of reduced coordination

Chapter 6

350

Scheme 6.2

Scheme 6.3

Scheme 6.4 number, for its easy re-­entry to give the N-­bonded isomer (Scheme 6.2).16 This assumes that dissociation (a) of the ion pair [(H2O)5Cr]3+·[SCN]− is relatively slow compared with its rearrangement (b), to give finally the N-­bonded isomer, or its reaction with water (c). However, an intramolecular mechanism has been proposed based on subsequent studies on the linkage isomerization of [(H2O)5Cr–SCN]2+,17a [(H2O)4Cr(SCN)(NCS)]+,17b and [(NH3)5Co(SCN)]2+.18a In agreement with this, no exchange with N14CS− in solution was observed in the case of [(H2O)5Cr–SCN]2+.17a A mechanism in which SCN− rotates has been suggested (Scheme 6.3).17,18a This involves synchronous bond fission immediately and new bond formation. The novel rearrangement shown in Scheme 6.4 is akin to a linkage isomerization process and occurs intramoleularly.18b The following is an interesting linkage isomerization that has been studied in the solid state:18c K[(NC)5Cr–CN–Fe] → K[(NC)5Cr–NC–Fe] At 50–70 °C the rate constant is of the order of 10−4 s−1.

6.2  Geometrical Isomerization This is known in both square-­planar and octahedral complexes. Geometrical isomers of inert complexes, such as the square-­planar complexes of Pt(ii) and octahedral complexes of Pt(iv), Co(iii), Cr(iii), Rh(iii), Ir(iii), etc., are

Isomerization, Optical Inversion and Racemization Reactions

351

well known. Our present state of knowledge on the kinetics and mechanisms of the interconversion of geometrical isomers of square-­planar and octahedral complexes is discussed in the following sections.

6.2.1  Square-­planar Complexes Thermodynamic data on the cis ⇌ trans equilibria in benzene solutions of [PtX2(ER3)] (E = P, As; R = Me, Et, Prn, Bun, n-­pentyl; X = Cl, I) have been reported. Both cis-­ and trans-­[PtCl2(PEt3)2] are stable in benzene solution at room temperature, but if a trace amount of PEt3 is added to the solution of either isomer, isomerization to a cis ⇌ trans equilibrium mixture is completed within 30 min.19 Similarly, [PtCl2(SEt2)2] isomerizes in solution in the presence of Et2S, which acts as a catalyst.20 These observations, that nucleophiles catalyse isomerization, suggest the operation of an associative mechanism in which isomerization proceeds through the formation of a five-­coordinate intermediate, resulting from the addition of the free ligand which acts as a catalyst. Similar results have also been reported for palladium(ii) complexes21 of the type [PdX2L2]. As a result of all such studies, the general mechanism that has been proposed is shown in Figure 6.1, where L may be the same as A or a different ligand. Each of the three pathways (I, II and III) occurs under various conditions, and the conditions favouring each were delineated by Redfield and Nelson.22 In some investigations21a of catalysed isomerization, the lack of interchange of A and L was taken to imply that any mechanistically important pentacoordinate species cannot have a regular geometry, but must be distorted in such a way that A and L can never become equivalent with respect to the central metal ion in the five-­ coordinate intermediate. More information was provided by Favez et al.,23 who studied the cis ⇌ trans isomerization of [PtX2(PR3)2] (X = halide) in the presence of PR3 in CH2Cl2 using 31P NMR spectroscopy to characterize reaction intermediates. With this and other spectroscopic techniques, they identified four-­and five-­ coordinate species, such as [PtX(PR3)3]+, [PtI2(PMe3)3] and [PtX(PR3)4]+, in solution. Their final conclusion was that isomerization takes place by a double displacement mechanism:   

  

  

Fast  ptX  PR 3 3   X  cis‐ ptX 2  PR 3 3   PR 3 

(6.5)

  trans‐ptX 2  PR 3    PR 3  ptX  PR 3 2   X     2

(6.6)

Slow

The Cl−-­catalysed isomerization of the Pt(ii) complexes may take place through five-­coordinate intermediates interchanging by a pseudo-­rotation process (Figure 6.2).24 However, the isomerization process shown in Scheme 6.5 follows a dissociative pathway.25 Another example of a dissociative mechanism of reaction of a square-­planar Pt(ii) complex is uncatalysed cis ⇌ trans isomerization of cis-­[Pt(o-­tolyl)Cl (PEt3)2] in methanol and ethanol.26 Whereas the substitution of the coordinated chloride occurs with the usual two-­term rate law [see Chapter 3, eqn (3.2)], the

Chapter 6

352

Figure 6.1  Mechanism  of nucleophile-­catalysed cis ⇌ trans isomerization of square-­planar complexes.

isomerization takes place at a rate that is dependent only on the concentration of the starting isomer, with a first-­order rate constant that is nearly two orders of magnitude smaller than k1 for the substitution reactions of the cis complex. The isomerization is sensitive to mass-­law retardation by chloride ions. The entropy change that accompanies the formation of the transition state during isomerization has a large positive value (+21 cal K−1 mol−1 in methanol). The authors suggested a dissociative mechanism in which the rate-­determining step is the cleavage of the Pt–Cl bond to form a three-­coordinate intermediate, which is consistent with the experimental data. However, it provides only a very minor path for substitution in this complex.

Isomerization, Optical Inversion and Racemization Reactions

353

Figure 6.2  Mechanism  of chloride-­catalysed isomerization of [Pt(N–O)(L)Cl)] type

complexes (N–O = glycinate, sarcosinate or N,N′-­dimethylglycinate; L = DMSO). Exchange of Cl− takes place during isomerization.

Scheme 6.5 The isomerization of cis-­[PdX2(am)2], prepared by the reaction of an amine (am) with [PdX2(PhSCH2CH2SPh)], is catalysed by the excess of amine present in the reaction mixture.27 The rate of isomerization obeys the rate law Rate = k[cis-­isomer][am] and for the isomerization process the following sequence of reactions has been proposed:

  



Fast cis‐ pdX 2  am 2   am   pdX  am 3  X  trans‐ pdX 2  am 2   am

  

(6.7)

6.2.2  Octahedral Complexes Interconversions of the geometrical isomers of octahedral complexes of cobalt(iii) are known from early days.28 Thus, Jorgensen29 recognized the cis ⇌ trans interconversion in [CoCl2(en)2]+. Werner30 reported that prolonged boiling or evaporation to dryness of cis-­[Co(NO2)2(en)2]NO3 and [Co(NO2)(Cl) (en)2]Cl yields the corresponding trans isomers. Similar observations were reported later by Delepine31 in the case of cis-­K3[M(C2O4)2Cl2] (M = RhIII, IrIII). Subsequent studies on a few typical systems established a mechanism for

Chapter 6

354

the isomerization process. Most of the studies were carried out on complexes of the type [M(A–A)2(X)2] and [M(A–A)2(X)(Y)], which show cis ⇌ trans isomerization and also optical isomerization (Δ ⇌ Λ) of cis isomers.32 The general mechanisms for such isomerization and racemization reactions are shown in Figure 6.3. This involves either a dissociation process with dissociation of an X (or Y) to generate the intermediate (Figure 6.3a), an intramolecular process involving one-­ended dissociation of a chelate (Figure 6.3b), or a twist mechanism without any metal–ligand bond rupture (Figure 6.3c), as in the racemization of some [M(A–A)3] (see Figure 6.8), or even an associative catalysed process (Figure 6.3d). For a number of optically active ions of the type cis-­[M(en)2(X)(Y)]+ (M = Co, Cr) there is an initial change in optical rotation (a mutarotation process) having a rate comparable to that for aquation [eqn (6.8)]. The change in optical rotation is due to the difference in molar rotation of two complex species. The aqua complex which results as above then racemizes to the racemic product

Figure 6.3  Possible  mechanisms for trans–cis interconversion in complexes of the type [M(A–A)2X2].32

Isomerization, Optical Inversion and Racemization Reactions

355 −

(Δ + Λ in a 1 : 1 molar ratio, hence zero rotation) without loss of X . Hence such aqua complexes are, in general, most suitable to study the relationship of isomerization, racemization and substitution (using H2O* exchange).33   



cis-­Λ-­[M(A−A)2(X)(Y)]+ + H2O → cis-­Λ-­[M(A−A)2(X)(OH2)]2+ + Y−

  

(6.8)

A cis ⇌ trans isomerization that has been extensively studied is the interconversion of the green praseo {trans-­[CoCl2(en)2]+} and violet violeo {cis-­[CoCl2(en)2]+} ions. If an aqueous solution of the green praseo salt is concentrated on a steam bath, the crystals of the violet violeo salt separate out from the solution. This violet (cis form) salt can be transformed into the green (trans form) salt by evaporation of its hydrochloric acid solution. From a study of the isomerization process in the presence of labelled Cl− ion (containing 36Cl), it has been shown that isomerization is accompanied by a completely random distribution of 36C1−, by its exchange with ligand Cl− bound in the complex.34 It was also established that there was no direct replacement of chloride bound in the complex with chloride in solution. Hence isomerization is presumably associated with the known reversible aquation of the complex. The particular isomer that separates from solution is largely determined by the relative solubilities of the isomeric salts. The less soluble cis-­[CoCl2(en)2]Cl is obtained from aqueous solution, whereas from hydrochloric acid solution trans-­ [CoCl2(en)2]C1·HCl separates out, being much less soluble. On the other hand, an aqueous solution of either cis-­ or trans-­[CoCl2(en)2]NO3 yields the less soluble trans isomer upon concentration. Ettle and Johnson34 pointed out that cis ⇌ trans interconversion in [CoCl2(en)2]+ takes place by the mechanism shown in Scheme 6.6. Both cis-­and trans-­[CoCl(en)2(OH2)]2+ isomerize to the cis ⇌ trans equilibrium mixture at rates comparable to their respective rates of formation from cis-­or trans-­[CoCl2(en)2]+.35 It is also known that when [MX2(A–A)2] is formed again from [MX(A–A)2] intermediate of tbp structure by the entry of the X, it will form either the cis or the trans isomer depending on the site of entry (Figure 6.4). Assuming that all three sites of attack are equally accessible, the entry of X will lead to cis and trans isomers in a 2 : 1 ratio. However, the reaction will proceed until the equilibrium concentrations of cis and trans isomers are reached. Aquation of cis-­[CoCl2(en)2]+ produces 100% cis-­[CoCl(en)2(OH2)]2+ as the initial product; the aquation of the trans isomer produces directly a mixture of

Scheme 6.6

356

Chapter 6

Figure 6.4  Dissociation  mechanism for cis ⇌ trans interconversion of [MX2(A–A)2] type complexes; ∠2.3, ∠2,4 and ∠3,4 indicate the site of entry of X. Entry between 2 and 3 will lead to the trans isomer whereas the other two will lead to the cis isomer.

∼35% cis-­ and ∼65% trans-­[CoCl(en)2(OH2)]2+, hence the sequence of changes indicated earlier has to be modified. Also, the isomerization at steam bath temperature (>80 °C) must certainly involve the formation of diaqua complex, due to further aquation of [CoCl(en)2(OH2)]2+. A dissociation mechanism involving a trigonal bipyramidal intermediate, which readily provides a path for isomerization, explains all the observed facts (Figure 6.4). Here X may be any ligand, such as Cl− or H2O for the Co(iii) systems under discussion. Further insight into the mechanism of cis ⇌ trans interconversion in cis-­ [CoCl(en)2(OH2)]2+ was obtained from ΔV‡ for the cis to trans and trans to cis conversions, viz. +8.0 and +5.1 cm3 mol−1, respectively.33b,c Hence the molar volume of the activated complex for the reactions is 5–8 cm3 mol−1 larger than that of either of the isomers, which rules out a twist mechanism; bond cleavage of the Co–OH2 bond forming a five-­coordinate intermediate [CoCl(en)2]2+ of trigonal bipyramidal structure (which can form both cis and trans isomers only on re-­entry of the H2O) (Figure 6.4) is in agreement with the observed ΔV‡ values and the mechanism seems likely to be an Id process. From a comparison of k(cis ⇌ trans) and k(racemization) values for several [Co(X)(en)2(OH2)]n+ complexes (Table 6.1),33 it is seen that for X = OH−, C1−, Br−, N3− and probably SCN− and H2O (but not NH3) the loss of optical activity of cis-­Λ-­(+)-­[Co(X)(en)2(OH2)]n+ arises mainly from cis → trans conversion rather than true racemization cis-­Λ ⇌ cis-­Δ forming Δ and Λ in a 1 : 1 molar ratio in the equilibrated product in solution. The exchange process shown in eqn (6.9) is stereoretentive and has ΔV‡ value of +5.9 cm3 mol−1; hence an Id mechanism with little bond formation by the entering H2O was suggested for the exchange involving essentially a square-­pyramidal intermediate.36,37 It is possible that for the cis–trans interconversion of a chloroaqua complex, a square-­pyramidal intermediate is generated initially, but this must change to a trigonal bipyramidal structure to permit isomerization. A similar mechanism has been suggested for a number of complexes of this type, [CoX(en)2(OH2)]n+.   

trans-­[Co(en)2(∗OH2)2]3+ + 2H2O ⇄ trans-­[Co(en)2(OH2)2]3+ + 2H2O∗   

(6.9)

Isomerization, Optical Inversion and Racemization Reactions

357

−1

Table 6.1  Rate  constants (k/s ) for isomerization, racemization and H2O* exchange in complexes of Co(iii), [Co(X)(en)2(OH2)]n+, in solution at 25 °C.33

X

105k s−1 (cis → trans)

105k s−1 (trans → cis)

105k s−1 (racemization)

105k s−1 (exchange)

OH− Br− Cl− N3− NCS− H2O NH3

220 5.1 2.0 5.2 0.014 ∼0.012 Ni2+ > Zn2+ > Co2+ > Cd2+ > Ca2+ > Sr2+ > Ba2+ > Mg2+, which, with the exception of Mg2+, is the order of the formation constants of the monooxalato complexes of the catalysing metal ions.47 The results were interpreted in terms of a one-­ended dissociation mechanism, in which the catalysing cation attacks the chelated oxalate to form a five-­coordinate intermediate, with only one end of oxalate bound to Cr(iii) (III). Subsequently, the released end of the oxalate becomes re-­bound to chromium, yielding the isomerized product with simultaneous release of Maq2+ ion.

Evidence has been presented48 that the dissociative trans ⇌ cis isomerization of trans-­[Cr(mal)2(OH2)2]− (mal = malonate) involves dissociation of an aqua ligand, but the corresponding reaction of the trans-­[Cr(ox)2(OH2)2]− (ox = oxalate) involves one-­ended dissociation of a Cr–oxalate chelate ring;

Chapter 6

360 −1

the observed ΔV values are +8.9 and −16.6 cm mol , respectively, in the two cases. The negative value for the oxalato system results from considerable solvation of the free –COO− end resulting from the oxalate chelate ring opening. However, for the malonato complex, having a more stable six-­membered chelate ring, the five-­coordinate intermediate is generated by dissociation of an aqua ligand, which leads to the normal expected positive value for ΔV‡. For the trans ⇌ cis isomerization of [Cr(A–A)2(OH2)2]− (A–A = oxalate, malonate) in mixed solvents H2O–S (S = MeOH, EtOH, dioxane, acetone), the rate is S dependent for the oxalato complex and for each S decreases with increasing concentration of S,49 but for the malonato complex the rate is not S dependent but increases with decrease in the concentration of H2O.50 The observed difference also suggests a difference in mechanism. Both oxalate ring opening and twist mechanisms were proposed for the oxalato complex, whereas a dissociation mechanism involving dissociation of an aqua ligand was proposed for the malonato complex. For the trans ⇌ cis isomerization of [Co(O2CMe)(en)2(OH2)]2+, the value of ΔV‡ is +7.9 cm3 mol−1 (in 0.05 M HClO4) and this positive value is in keeping with a dissociative reaction.51 The mer ⇌ fac isomerization of the complexes [Co(dien)(L)(OH2)2]3+ (L = H2O, NH3, etc.) takes place by a dissociative mechanism involving a trigonal bipyramidal intermediate, resulting from loss of an aqua ligand.52 ‡

3

6.3  Other Types of Structural Isomerization Examples of such isomerizations are square-­planar ⇌ tetrahedral, square-­ planar ⇌ square-­pyramidal, square-­planar ⇌ octahedral, etc., and also conformational isomerization such as δ ⇌ λ as in [Fe(CN)4(L–L)]− (L–L = en, pn, etc.), where the isomerism is due to the conformational change of the chelated diamine. The inherent lability of the square-­planar ⇌ tetrahedral equilibria, which mostly involve Ni(ii), requires that NMR line broadening53 or photochemical perturbation methods54 be used {as in the case of [NiCl2(dpp)] (dpp = 1,3-­bis(diphenylphosphino)propane}, mentioned under photochemical reactions in Chapter 9, Section 9.2. For the square-­planar ⇌ tetrahedral interconversion in [NiX2(PR3)2], the first-­order rate constants are 105–106 s−1 at 25 °C (in CDCl3 or CD2C12).54 In the square-­planar ⇌ square-­pyramidal equilibria, of course, there is a change in coordination number. The kinetics associated with the equilibria in eqn (6.11) (L = an N2S2 quadridentate macrocyclic ligand, H2O is in the axial position in the square-­planar five-­coordinate species) were followed in a manner similar to that for the square-­planar ∏ tetrahedral equilibria in [NiCl2(dpp)] using a concentration jump induced photochemically.55 The values of k1 and k–1 at 20 °C (in aqueous solution) are 4.5 × 106 and 1.5 × 107 s−1, respectively.   

  



(6.11)

Isomerization, Optical Inversion and Racemization Reactions

361

The square-­planar ⇌ octahedral equilibria can be represented in a general way as in eqn (6.12), where L4 is generally a quadridentate macrocyclic ligand or two bidentate ligands and S is a unidentate ligand, such as H2O.   

  

k1

k2

k1

k2

ML 4  S # M(L)4 S; ML 4 (S)  S # ML 4 (S)2

(6.12)

The equilibria are invariably labile (see, however, ref. 56) and as the intermediate ML4(S) is usually present at a very low concentration compared with the other species, the perturbation method (T-­jump),57 ultrasonic absorption58 or photoexcitation59 gives only one relaxation and hence only an incomplete description of the system. Microwave T-­jump and NMR studies rule out four-­ coordinate ⇌ five-­coordinate interconversion as the rate-­determining step in the system in chlorobenzene [see eqn (6.13), where A is a tetradentate ligand and L is a substituted pyridine].   



[Ni(A)] + L ⇌ [Ni(A)(L)];  [Ni(A)(L)] + L ⇌ [Ni(A)(L)2]

(6.13)

  

Study of the exchange of ML4(S)2 with S by NMR line broadening allows the evaluation of k−2 and further information from relaxation data allows the evaluation of k1, and hence the ratio k−1/k2, as was done for [Ni(12-­aneN4)]2+ changing to [Ni(12-­aneN4)(OH2)2]2+ through [Ni(12-­aneN4)(OH2)]2+studied by 17O NMR and T-­jump methods.57 The values obtained for this system in aqueous solution at 25 °C (I, 3 M LiClO4) were k−2 = 4.2 × 107 s−1, k1 = 5.8 × 103 M−1 s−1 and k−1/k2 = 0.016.57 In this system, the first step is rate determining. For some earlier work on [Ni(L)]2+ (low-­spin square-­planar) ⇌ Ni(L)(OH2)2]2+ (high-­spin octahedral) equilibria (L = a tetradentate polyamine), the original literature can be consulted.60 Octahedral ⇌ tetrahedral equilibria in Co(ii) complexes in non-­aqueous solvents are well known. Kinetic data on such systems are too few, but those available from T-­jump studies in pyridine solution were interpreted as shown in Scheme 6.8.61 Spin equilibria (L.S ⇌ H.S.) in octahedral complexes have been studied in a few cases. An example is [FeII(L–L)3]2+, where L–L = 2-­(2′-­pyridyl)imidazole, studied in acetone solution, for which the following values were reported:62 k1 = 5 × 106 s−1 (∆V‡, +5.5 cm3 mol−1) and k−1 = 1 × 107 s−1 (∆V‡, −5 cm3 mol−1). Similarly for [FeIII(sal2trien)]+ (in aqueous solution),63 k1 = 6.1 × 107 s−1 and k−1 = 1.3 × 108 s−1. Some rate data are available on conformational isomerization, δ ⇌ λ, in [Fe(CN)4(1,2-­diamine)]− complexes in CD3OD + DCl at 25 °C.64 The

Scheme 6.8

Chapter 6

362

Table 6.2  Kinetic  parameters for conformational isomerism (see text). 1,2-­Diamine

k1/s−1

∆H‡/kJ mol−1

∆S‡/J K−1 mol−1

en meso-­bn (2R,3S) cis-­chxn (1R,2S)

3 × 108 2 × 107 8 × 104

25 30 43

0 −3 −8

Scheme 6.9 interconversion slows and ∆H‡ increases as the bulkiness of the carbon substitution increases (Table 6.2). It is believed that the interconversion proceeds through an envelope conformation (Scheme 6.9).64–66

6.4  Optical Inversion It is often true that the configuration of an optical isomer of a species is retained when it is converted to another species by a ligand substitution process. Thus, (+)-­[CoCl2(en)2]+ can be transformed to (+)-­[CoCl(en)2(OH2)]2+, (+)-­[CoCl(NCS)(en)2]+ and (+)-­[Co(ox)(en)2]+. The (+) in these cases refers to dextrorotatory and similarly (−) refers to laevorotatory with respect to the Na D-­line (589 nm). However, a change in configuration in such substitutions sometimes takes place and this phenomenon is called optical inversion, the first example of which was reported by Bailar and Auten,67 who denoted it “Walden inversion” in inorganic complexes (by analogy with similar examples in organic compounds). but many prefer to call it Bailar inversion. Their system is shown in Scheme 6.10. More systematic and extensive investigations of this system68,69 led to the conclusion that aquation, rather than the carbonate used, is an important factor, as shown in Scheme 6.11. Investigations by another group70 indicated that inversion occurs as a consequence of base hydrolysis of d-­[CoCl2(en)2]+ in the presence of Ag+ ion (Scheme 6.12). The authors proposed that inversion occurs due to a trans-­ attack displacement process involving both Ag+ and OH− (see Figure 6.6).70 However, appropriate dissociative mechanisms are equally feasible. Base hydrolysis of d-­[CoCl2(en)2]+ ion accompanied by inversion can take place in the absence of Ag+ ion in concentrated (OH− ≥0.25 M) but not in dilute (OH− 10 have been followed by stopped-­flow spectrophotometric method.102

Inversion occurs through a copper–hydroxo species (Scheme 6.18) which leads to Rate = kobs[Cu(tet-b)(Blue)]total[OH−], where kobs = kKOH[OH−]/(1 + KOH[OH−]); reported values are KOH = 51.6 M−1, k1 = 5.5 s−1, I = 5.0 M at 25 °C.102 A similar kinetic pattern has been observed for the blue → red change for [Cu(tet-­b)]2+ (tet-­b = (racemic)-­5,5,7,12,12,14-­hexamethylcyclam) (Scheme 6.18), but the reactions are much slower as two inverting nitrogens and less favourable ring conformational changes are involved and the reported values at 25 °C, I = 0.1 M are KOH = 5.06 × 102 M−1 and k = 6.2 × 10−3 s−1.103a Similar data for several other systems having an anion in place of OH− have also been reported.103b

References 1. S. M. Jorgensen, Z. Anorg. Chem., 1893, 5, 169. 2. R. Ferrararo and L. Fabbrizzi, Inorg. Chim. Acta, 1978, 26, L15.

Isomerization, Optical Inversion and Racemization Reactions

381

3. R. G. Pearson, P. M. Hemry, J. G. Bergmann and F. Basolo, J. Am. Chem. Soc., 1954, 76, 5920. 4. (a) B. Adell, Z. Anorg. Allg. Chem., 1952, 271, 49; 1955, 279, 220; (b) B. Adell, Acta Chem. Scand., 1952, 5, 54, 941. 5. F. Basolo and G. Hammaker, Inorg. Chem., 1962, 1, 1. 6. R. Murmann and H. Taube, J. Am. Chem. Soc., 1956, 78, 4886. 7. R. K. Murmann, J. Am. Chem. Soc., 1955, 77, 4190. 8. I. Grenthe and E. Nordin, Inorg. Chem., 1979, 18, 1109, 1869. 9. W. G. Jackson, G. A. Lawrance, P. A. Lay and A. M. Sargeson, Inorg. Chem., 1980, 19, 904. 10. W. G. Jackson, G. A. Lawrance, P. A. Lay and A. M. Sargeson, J. Chem. Soc., Chem. Commun., 1982, 70. 11. D. A. Buckingham, I. I. Creaser and A. M. Sargesonm, Inorg. Chem., 1970, 9, 655. 12. M. Mares, D. A. Palmer and H. Kelm, Inorg. Chim. Acta, 1978, 27, 153. 13. D. A. Palmer, R. van Eldik and H. Kelm, Inorg. Chim. Acta, 1978, 30, 83. 14. F. Basolo, W. H. Baddley and K. L. Weidenbaum, J. Am. Chem. Soc., 1966, 88, 1576. 15. M. Cusumano, G. Guglelmo and P. Marricchi, Inorg. Chim. Acta, 1978, 30, 29. 16. A. Haim and N. Sutin, J. Am. Chem. Soc., 1965, 87, 4210; 1966, 88, 434. 17. (a) M. Orhanovic and N. Sutin, J. Am. Chem. Soc., 1968, 90, 4286; (b) L. D. Brown and D. P. Pennington, Inorg. Chem., 1971, 10, 2117. 18. (a) D. A. Buckingham, I. I. Creaser and A. M. Sargeson, Inorg. Chem., 1970, 9, 655; (b) J. N. Armor and H. Taube, J. Am. Chem. Soc., 1970, 92, 2560; (c) J. E. House Jr and J. C. Bailar Jr, Inorg. Chem., 1969, 8, 672. 19. J. Chatt and R. G. Wilkins, J. Chem. Soc., 1952, 273, 4300; 1953, 70; 1956, 525. 20. H. D. K. Drew and G. H. Wyatt, J. Chem. Soc., 1934, 56. 21. (a) P. Haake and R. M. Pfeiffer, J. Chem. Soc., Chem. Commun., 1969, 1330; (b) P. Haake and R. M. Pfeiffer, J. Am. Chem. Soc., 1970, 92, 4996; (c) H. K. J. Powell and D. G. Cooper, J. Chem. Soc., Chem. Commun., 1974, 749; (d) D. A. Redfield, L. W. Cary and J. H. Nelson, Inorg. Chem., 1975, 14, 50. 22. D. A. Redfield and J. H. Nelson, J. Am. Chem. Soc., 1974, 96, 6219. 23. R. Favez, R. Roulet, A. N. Pinkerton and D. Schwarzenbach, Inorg. Chem., 1980, 19, 1356. 24. L. F. Erikson, T. A. Ferrett and L. F. Buhse, Inorg. Chem., 1983, 22, 1461. 25. G. W. V. Cave, F. P. Fanizzi, R. J. Deeth, W. Errington and J. P. Rourke, Organometallics, 2000, 9, 1355. 26. G. Faraone, V. Ricevutto, R. Romeo and M. Trozzi, J. Chem. Soc. A, 1971, 1877. 27. L. Cattalini and M. M. Martelli, J. Am. Chem. Soc., 1969, 91, 312. 28. N. Serpone and D. G. Bickley, Kinetics and Mechanisms of Isomerization and Racemization Process of Six-­coordinate Chelate Complexes, in Inorganic Reaction Mechanisms, Part II, ed. J. O. Edwards, Progress in InorganicChemistry, ed. S. J. Lippard, vol. 17, Interscience, New York, 1972, pp. 391–566.

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29. S. M. Jorgensen, J. Prakt. Chem., 1889, 39, 16; 1890, 41, 449. 30. A. Werner, Ber., 1911, 44, 2452, 3278. 31. M. Delepine, Compt. Rend., 1922, 175, 1409. 32. G. A. Lawrence and D. R. Stranks, Acc. Chem. Res., 1979, 12, 403. 33. (a) M. L. Tobe, Studies on Chemical Structure and Reactivity, ed. J. H. Ridd, Methuen, London, 1966; (b) W. G. Jackson and A. M. Sargeson, Inorg. Chem., 1978, 17, 1348; (c) W. G. Jackson, Inorg. Chim. Acta, 1987, 126, 147. 34. G. W. Ettle and C. H. Johnson, J. Chem. Soc., 1939, 1490. 35. M. E. Baldwin, S. C. Chan and M. L. Tobe, J. Chem. Soc., 1961, 4637. 36. Y. Kitamura, S. Nariyuki and K. Yoshitani, Inorg. Chem., 1985, 24, 3021. 37. S. B. Tong, H. R. Krasuse and T. W. Swaddle, Inorg. Chem., 1976, 15, 2643. 38. (a) D. D. Brown and C. K. Ingold, J. Chem. Soc., 1953, 2680; (b) B. Bosnich, C. K. Ingold and M. L. Tobe, J. Chem. Soc., 1965, 4074; (c) D. D. Brown and R. S. Nyholm, J. Chem. Soc., 1953, 2696. 39. F. Basolo, J. Am. Chem. Soc., 1950, 72, 4393. 40. J. Bjerrum and S. E. Rasmussen, Acta Chem. Scand., 1952, 6, 1265. 41. W. Krause and H. Taube, J. Am. Chem. Soc., 1961, 83, 1280. 42. D. C. Olson and C. S. Garner, Inorg. Chem., 1963, 2, 415. 43. D. R. Stranks and N. Vanderhock, Inorg. Chem., 1976, 15, 2639. 44. D. F. Martin and M. L Tobe, J. Chem. Soc., 1962, 1388. 45. A. Werner, Ann. Phys., 1914, 406, 261. 46. (a) R. E. Hamm, et al., J. Am. Chem. Soc., 1953, 75, 609; 1955, 77, 2083; 1961, 83, 340; (b) G. L. Welch and R. E. Hamm, Inorg. Chem., 1963, 2, 295. 47. (a) K. R. Ashley and R. E. Hamm, Inorg. Chem., 1965, 4, 1129; (b) H. Kelm, H. Stieger and G. M. Harris, Z. Phys. Chem., 1969, 67, 98. 48. M. Casula, G. Illuminati and G. Ortaggi, Inorg. Chem., 1972, 11, 1062. 49. K. Uchida and Y. Takinami, Bull. Chem. Soc. Jpn., 1981, 54, 2298. 50. K. Uchida and Y. Takinami, Bull. Chem. Soc. Jpn., 1980, 53, 3522. 51. G. A. Lawrence and S. Suvachittanont, J. Coord. Chem., 1979, 9, 13. 52. D. A. House, Inorg. Chim. Acta, 1978, 30, 281. 53. (a) L. H. Pignolate, W. de W. Horrocks and R. H. Holm, J. Am. Chem. Soc., 1970, 92, 1855; (b) G. N. La Mar and E. O. Sherman, J. Am. Chem. Soc., 1970, 92, 2691; (c) L. H. Pignolate and G. N. La Mar, Dynamics of Intramolecular Rearrangements in NMR of Paramagnetic Molecules, ed. G. N. La Mar, W. D. Horrocks and R. H. Holm, Academic Press, New York, 1974, ch. 8; (d) R. H. Holm, Stereochemically Non-­rigid Metal Chelate Complexes in Dynamic Nuclear Magnetic Resonance Spectroscopy, ed. L. M. Jackson and F. A. Cotton, Academic Press, New York, 1974, ch. 9. 54. J. L. McGarvey and J. Wilson, J. Am. Chem. Soc., 1975, 97, 2531. 55. L. Campbell, J. J. McGarvey and N. S. Samman, Inorg. Chem., 1978, 17, 3378. 56. A. M. Martin, K. J. Grant and E. J. Billo, Inorg. Chem., 1986, 25, 4904. 57. (a) J. H. Coates, D. A. Hadi, S. F. Lincoln, H. W. Dodgen and J. P. Hunt, Inorg. Chem., 1981, 20, 707; (b) M. Cusumano, J. Chem. Soc., Dalton

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Trans., 1976, 2133, 2137; (c) I. Sachinidis and M. W. Grant, J. Chem. Soc., Chem. Commun., 1978, 157. 58. (a) A. F. Godfrey and J. K. Beattie, Inorg. Chem., 1983, 22, 3794; (b) J. K. Beattie, M. T. Kelso, W. E. Moody and P. A. Tregloan, Inorg. Chem., 1985, 24, 415. 59. K. Ivin, R. Jamison and J. J. McGarvey, J. Am. Chem. Soc., 1972, 94, 1763. 60. (a) D. H. Busch, Acc. Chem. Res., 1978, 11, 392; (b) L. Fabbrizzi, M. Micheloni and P. Paoletti, Inorg. Chem., 1980, 19, 535; (c) L. Fabbrizzi, J. Chem. Soc., Dalton Trans., 1979, 1857; (d) K. J. Ivin, R. Jamison and J. J. McGarvey, J. Am. Chem. Soc., 1972, 94, 1763; (e) C. Creutz and N. Sutin, J. Am. Chem. Soc., 1993, 95, 7177; (f) R. G. Wilkins, R. Yelin, D. W. Margerum and D. C. Weatherburn, J. Am. Chem. Soc., 1969, 91, 4326; (g) B. Bosnich, R. D. Gillard, E. D. McKenzie and G. A. Webb, J. Chem. Soc. A, 1966, 1331. 61. R. D. Farina and J. H. Swinehart, Inorg. Chem., 1972, 11, 645. 62. J. DiBenedetto, V. Arkle, H. A. Goodwin and P. C. Ford, Inorg. Chem., 1985, 24, 455. 63. (a) I. Lawthers and J. J. McGarvey, J. Am. Chem. Soc., 1984, 106, 4280; (b) R. A. Binstead, J. K. Beattie, T. G. Dewey and D. H. Turner, J. Am. Chem. Soc., 1980, 102, 6442. 64. Y. Kuroda, N. Tanaka, M. Goto and T. Sakai, Inorg. Chem., 1989, 28, 997. 65. C. J. Hawkins and J. A. Palmer, Coord. Chem. Rev., 1982, 44, 1. 66. C. J. Hawkins, R. M. Peachey and C. L. Szoredi, Aust. J. Chem., 1978, 31, 973. 67. J. C. Bailar Jr and R. W. Auten, J. Am. Chem. Soc., 1934, 56, 774. 68. J. C. Bailar Jr, F. G. Jonelis and E. H. Hoffman, J. Am. Chem. Soc., 1936, 58, 2224. 69. J. C. Bailar Jr and D. F. Peppard, J. Am. Chem. Soc., 1940, 62, 820. 70. F. P. Dwyer, A. M. Sargeson and I. K. Rid, J. Am. Chem. Soc., 1963, 85, 1215. 71. L. J. Boucher, E. Kyuno and J. C. Bailar Jr, J. Am. Chem. Soc., 1964, 86, 3656; 1965, 87, 4458. 72. (a) F. Basolo and R. G. Pearson, Mechanisms of Inorganic Reactions: A Study of Metal Complexes in Solution, John Wiley, New York, 2nd edn, 1967, pp. 268–270; (b) F. Basolo, Adv. Chem. Ser., 1967, 62, 408. 73. (a) J. C. BailarJr, J. Hoslam and E. M. Jones, J. Am. Chem. Soc., 1936, 58, 2226; (b) R. D. Archer and J. C. Bailar Jr, J. Am. Chem. Soc., 1961, 83, 812. 74. (a) L. J. Boucher, E. Kyuno and J. C. Bailar Jr, J. Am. Chem. Soc., 1965, 87, 4458; (b) E. Kyono and J. C. Bailar Jr, J. Am. Chem. Soc., 1966, 88, 1125; (c) E. Kyuno and J. C. Bailar Jr, J. Am. Chem. Soc., 1966, 88, 1120. 75. (a) R. G. Wilkins and M. J. G. Williams, J. Chem. Soc., 1957, 1763; (b) G. A. Lawrence and D. R. Stranks, Inorg. Chem., 1978, 17, 1804; (c) A. Yamigashi, Inorg. Chem., 1986, 25, 55; (d) F. Basolo, J. C. Hayes and H. M. Neumann, J. Am. Chem. Soc., 1954, 76, 3807; (e) E. L. Blinn and R. G. Wilkins, Inorg. Chem., 1976, 15, 2952. 76. (a) F. Basolo, J. C. Hayes and H. M. Neumann, J. Am. Chem. Soc., 1953, 75, 5102; (b) Stability Constants of Metal-­Ion Complexes, compiled by L. G. Sillen and A. E. Martell, Special Publication No. 25, The Chemical Society, London, 1971, Table 906.

384

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77. (a) A. Werner, Justus Liebigs Ann. Chem., 1912, 386, 1; (b) A. Werner, Ber. Dtsch. Chem. Ges., 1912, 45, 1228, 3061. 78. (a) F. A. Long, J. Am. Chem. Soc., 1939, 61, 570; 1941, 63, 1353; (b) C. A. Bunton, J. H. Carter, D. R. Llewellyn, C. O'Conner, A. L. Odell and S. Y. Yih, J. Chem. Soc., 1964, 4615, 4622, 4627; (c) J. A. Broomhead, J. Chem. Soc. A, 1971, 645; (d) L. Damrduer and R. M. Milburn, J. Am. Chem. Soc., 1968, 90, 3884; 1971, 93, 6481. 79. R. Charonnat, Annu. Chim., 1931, 16, 202. 80. C. H. Johnson and A. M. Mead, Trans. Faraday Soc., 1935, 31, 1621. 81. (a) D. W. Cooke, Y. A. Im and D. H. Busch, Inorg. Chem., 1962, 1, 13; (b) D. Sen and W. C. Fernelius, J. Inorg. Nucl. Chem., 1959, 10, 269. 82. P. Ray and N. K. Dutt, J. Indian Chem. Soc., 1943, 20, 81. 83. J. C. Bailar Jr, J. Inorg. Nucl. Chem., 1958, 8, 165. 84. A. Rodger and B. F. G. Johnson, Inorg. Chem., 1988, 27, 3061. 85. J. C. Bailar Jr, J. Inorg. Nucl. Chem., 1980, 42, 201. 86. C. S. Springer Jr and R. E. Sievers, Inorg. Chem., 1967, 6, 852. 87. (a) D. Banerjea and J. C. Bailar Jr, Transition Met. Chem., 1985, 10, 331; (b) D. Banerjea, Transition Met. Chem., 1988, 13, 160. 88. A. Y. Girgis and R. C. Fay, J. Am. Chem. Soc., 1970, 92, 7061. 89. G. A. Lawrence and D. R. Stranks, Inorg. Chem., 1977, 16, 929 and references therein. 90. R. C. Fay and T. S. Piper, J. Am. Chem. Soc., 1962, 84, 2303; 1963, 85, 500. 91. G. Gordon and R. H. Holm, J. Am. Chem. Soc., 1970, 92, 5319. 92. R. H. Holm, D. H. Gerlach, J. G. Gordon II and M. G. McNamee, J. Am. Chem. Soc., 1968, 90, 4184. 93. G. H. Searle, R. F. Keene and S. F. Lincoln, Inorg. Chem., 1978, 17, 2362. 94. (a) H. Kawaguchi, T. Ama and T. Yasui, Bull. Chem. Soc. Jpn., 1984, 57, 2422; (b) H. Kawaguchi, M. Shimizu, T. Ama and T. Yasui, Bull. Chem. Soc. Jpn., 1989, 62, 753. 95. See H. Bruner, Optical Activity, in The Organometallic Chemistry of Iron, ed. E. A. Koerner Von Gustorf, F.-­W. Grevels and I. Fischler, Academic Press, New York, 1978, p. 299. 96. H. Brunner, Adv. Organomet. Chem., 1980, 18, 151. 97. L. E. Nivorozhkin, A. L. Nivorozhkin, M. S. Korobov, L. E. Konstantinovsky and V. I. Minkin, Polyhedron, 1985, 4, 1701 and references therein. 98. (a) E. J. Billo, Inorg. Chem., 1981, 20, 4019; (b) R. W. Hay and M. Akbar Ali, Inorg. Chim. Acta, 1985, 103, 23; (c) A. M. Martin, K. J. Grant and E. J. Billo, Inorg. Chem., 1986, 25, 4904; (d) P. J. Connolloy and E. J. Billo, Inorg. Chem., 1987, 26, 3224. 99. E. K. Barefield, A. Bianchi, E. J. Billo, P. L. Connolly, P. Paoleti, J. S. Summers and D. G. Van Derveer, Inorg. Chem., 1986, 25, 4197. 100. (a) R. W. Hay and M. P. Pujari, J. Chem. Soc., Dalton Trans., 1986, 1455; (b) R. W. Hay and M. Govan, Transition Met. Chem., 1992, 17, 119.

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101. (a) L. Anichini, L. Fabbrizzi, P. Paoletti and R. W. Clay, Inorg. Chim. Acta, 1977, 24, L21; (b) E. J. Billo, Inorg. Chem., 1984, 23, 236. 102. C.-­S. Lee and C.-­S. Chung, Inorg. Chem., 1984, 23, 639. 103. (a) B.-­F. Liang, D. W. Margerum and C.-­S. Chung, Inorg. Chem., 1979, 18, 2001;(b)C.-­S.Lee,G.T.WangandC.-­S.Chung,J.Chem.Soc.,DaltonTrans.,1984, 109.

Chapter 7

Electron Transfer Reactions 7.1  Introduction Chemists are quite familiar with electron transfer reactions. Many of the commonly studied fast reactions also belong to this category. Broadly speaking, electron transfer reactions are of two types: (i) those with net chemical change [i.e. oxidation–reduction (redox) reactions] [eqn (7.1)] and (ii) those with no net chemical change (i.e. electron exchange between two different oxidation states of an element) [eqn (7.2)].   



[Fe(OH2)6]2+ + CeIV(aq) → [Fe(OH2)6]3+ + CeIII (aq)

(7.1)

[∗Fe(CN)6]4− + [Fe(CN)6]3− ⇌ [∗Fe(CN)6]3− + [Fe(CN)6]4−

(7.2)

  

  

Redox reactions are of great importance in chemistry and biochemistry. Many of the classical analytical methods are based on very rapid stoichiometric redox reactions, and redox reactions are also widely used in preparative chemistry. The role of some transition metal ions in vivo as metalloenzymes depends on their ability to participate selectively in electron transfer reactions. However, the rates of redox reactions even in analogous systems often do not show the same trend as that of the E° values. Thus, for the [M(OH2)6]3+– [M(OH2)6]2+ couple the values of the reduction potential at 25 °C are −0.41 V for M = Cr and −0.26 V for M = V, yet the rate constant for the reduction of [Co(NH3)6]3+ by Craq2+ at 25 °C is 8 × 10−5 M−1 s−1 and that for reduction by Vaq2+ is 3.7 × 10−3 M−1 s−1, i.e. faster by a factor of 102, despite the fact that in terms of the E° values Craq2+ is a more powerful reducing agent than Vaq2+. Also, Euaq2+, despite being comparable to Craq2+ in terms of E° value (−0.43 V) reduces [Co(NH3)6]3+ with a rate constant of 2 × 10−2 M−1 s−1 at 25 °C.   Mechanisms of Reactions of Metal Complexes in Solution By Debabrata Banerjea and M. K. Bharty © Debabrata Banerjea and M. K. Bharty 2023 Published by the Royal Society of Chemistry, www.rsc.org

386

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387

1

Taube and co-­workers suggested classifying the electron transfer reactions into two groups, depending on their general mechanisms: (i) those involving an outer-­sphere activated complex and (ii) those involving an inner-­sphere or bridged activated complex. The outer-­sphere mechanism will operate in systems where both the oxidizing and reducing agents are inert to substitution and the electron transfer rate is more rapid than the rate of substitution. On the other hand, the inner-­sphere mechanism is generally observed in systems where at least one of the reactants is labile to substitution and a suitable bridging group is available to bind as a bridging group to two centres, between which electron transfer takes place, in the activated complex. Developments in the field are covered in two books2 and several reviews.3,4 Some typical rate data on electron transfer reactions are given in Table 7.1. The k values refer to the observed rate law: Rate = k[oxidant][reductant]

Table 7.1  Rate  constants for some electron exchange reactions. System

Temperature/°C

k1/M−1 s−1

[VO(OH)]2+–[VO(OH)]+a [VO]2+–[VO(OH)]2+ [V(OH2)6]3+–[V(OH2)6]2+a [Cr(OH2)6]3+–[Cr(OH2)6]2+a MnO4−–MnO42− [Mn(OH2)6]3+–[Mn(OH2)6]2+a [Fe(phen)3]3+–[Fe(phen)3]2+ [Fe(CN)6]3−–[Fe(CN)6]4− [Fe(OH2)6]3+–[Fe(OH2)6]2+a [Fe(edta)]−–[Fe(edta)]2− [Co(phen)3]3+–[Co(phen)3]2+ [Co(OH2)6]3+–[Co(OH2)6]2+ [Co(NH3)6]3+–[Co(NH3)6]2+ [Co(en)3]3+–[Co(en)3]2+ [Co(sep)3]3+–[Co(sep)3]2+c [Co(ox)3]3−–[Co(ox)3]4− [Co(edta)]−–[Co(edta)]2− [Co(NH3)6]3+–[Cr(bipy)3]2+ [Co(NH3)5(OH2)]3+–[Cr(bipy)3]2+ [Mo(CN)8]3−–[Mo(CN)8]4−a [Ru(NH3)6]3+–[Ru(NH3)6]2+a [Ru(bipy)3]3+–[Ru(bipy)3]2+ [Os(bipy)3]3+–[Os(bipy)3]2+ [Os(phen)3]3+–[Os(phen)3]2+a [IrCl6]2−–[IrCl6]3−a Cuaq2+–Cuaq+a

25 0 25 25 0 25 0 0 25 0 0 0 25 25 25 25 25 4 4 25 25 25 0 25 25 25

1 × 10−2 0.29 1 × 10−2 2 × 10−5 7 × 102 3.2 × 10−4 >105 3.5 × 102 1.1 >103 1.1 0.75b 1.7 × 10−7 2 × 105 5.0 9 × 10−7 4 × 10−7 7.1 6.5 × 102 3.2 × 104 4 × 103 4 × 108 >105 3.2 × 108 2 × 105 1 × 10−5

a

 ef. 5; others are from various literature sources. R 5 M−1 s−1 at 25 °C.6 c sep = sepulchrate. b

388

Chapter 7

7.2  Outer-­sphere Mechanism In this mechanism, electron transfer takes place from one species to the other without these being chemically bonded at any stage in the course of the reaction; in the case of metal complexes as reactants, coordination spheres of the metals remain intact. Electron exchange between [Co(NH3)6]2+ and [Co(NH3)6]3+ (which can be studied by isotopic labelling of the Co in either of the oxidation states) or the reduction of [Co(NH3)6]3+ by [Cr(OH2)6]2+ takes place by an outer-­sphere mechanism. This mechanism always operates in the reaction between two metal complexes both of which are substitution inert and also when no good bridging group is present, as this prevents the possibility of binding the two centres through a bridging group as is essential for the inner-­sphere mechanism to operate (see Section 7.3). Further examples of this are the electron exchange observed in the following systems: MnO4−–MnO42−, IrCl62−–IrCl63−, Fe(CN)63−–Fe(CN)64−, Fe(phen)33+–Fe(phen)32+, Os(bipy)33+–Os(bipy)32+, etc. In cases such as [Co(NH3)6]3+–[Co(OH2)6]2+, the Co(ii) species is labile to substitution, but the lack of bridging properties of the ligand (NH3) and other considerations make it likely that this system also belongs to this group. Although CN− has good bridging properties, the substitution-­inert nature of [Fe(CN)6]3− and [Fe(CN)6]4− discourages the formation of an intermediate [(NC)5Fe–CN–Fe(CN)5]6− needed for operation of the inner-­sphere mechanism. In an outer-­sphere process, electron transfer takes place through the barrier due to the ligands in the coordination spheres of the metals and the outer solvation shells of the complexes. Considerable insight into such an electron transfer process in solution is provided by the electron tunnelling theory developed by Weiss, Marcus, Eyring and others.7 The possibility of an electron leaking through a potential barrier that would be classically impenetrable is a well-­known quantum-­mechanical phenomenon. The result is that the electron can transfer at distances considerably greater than would correspond to actual collision of the reactants.

7.3  T  he Marcus Equation: Marcus Cross-­relation and Its Applications7,8 Based on a sound theoretical approach, Marcus evaluated the different components that contribute to the overall ∆G‡ for an outer-­sphere reaction, and derived an equation for calculating theoretically the rate constant for an outer-­sphere reaction from the exchange rate constants and equilibrium constant for the overall reaction, which as usual can be calculated from the potential data. Marcus's work was recognized by the award of the Nobel Prize in Chemistry in 1992. The Marcus equation (Marcus cross-­relation)7c for the rate constant of an outer-­sphere electron transfer reaction of the type in eqn (7.3) is shown in eqn (7.4), where log f12 is expressed by eqn (7.5).†   

 or the reaction in eqn (7.3), if n electrons are transferred then log K12 = (n/0.0591)  E 1  E 2  at F 25 °C, where E 1 and E 2 are the standard reduction potentials of the two couples.

†

Electron Transfer Reactions



389

Ox1 + Red2 ⇌ Red1 + Ox2

(7.3)

k12 = (k11k22K12f12)½

(7.4)

log f12 = (log K12)2/4log(k11k22/Z2)

(7.5)

  

  

  

where Z is the number of collisions per second between particles in solution (∼1011 M−1 s−1 at 25 °C), k11 and k22 are the electron exchange rate constants for the Ox1–Red1 and Ox2–Red2 couples and k12 and K12 are the rate and equilibrium constants, respectively, for the redox process of eqn (7.3). The Marcus equation shows that the rates of redox reactions depend on an intrinsic factor (through k11 and k22) and a thermodynamic factor (through K12). This equation is another example of a linear free energy relationship (LFER). The cross-­relation [eqn (7.4)] is based on the following derived expression in terms of the free energies of activation and free energy of reaction [eqn (7.6)]:

  



‡ ‡ ΔG‡12 = 0.5(ΔG11 + ΔG22 + ΔG12 − 2.303RTlog f12)

(7.6)

  

when the oxidizing powers of Ox1 and Ox2 are comparable, i.e. E 1  E 2, then K12 ≈ 1 and according to eqn (7.5) log f12 ≈ 0, hence f12 ≈ 1, and in that case ΔG‡12 is expressed as in eqn (7.7):   



ΔG‡12 ≈ 0.5(ΔG‡11 + ΔG‡22 + ΔG12)

(7.7)

  

Hence, for an electron transfer from a series of analogous metal complexes as reductants to a particular oxidant, or from a particular reductant to a series of analogous oxidants, the plot of ΔG‡ versus ΔG (and hence also of log k versus log K for the redox reaction) should be linear with a slope equal to 0.5. This is illustrated in Figure 7.1 for the oxidation of a series of Fe(ii) and Ru(ii) complexes by Ce(iv) in dilute sulfuric acid.8a–c Later work indicated that for some inner-­sphere electron transfer reactions also a type of Marcus relationship is applicable.8d,e The usefulness of the Marcus equation can be illustrated by calculating the rate constant for the following outer-­sphere redox reaction:8g,h   



[Mo(CN)8]3− + [Fe(CN)6]4− ⇌ [Mo(CN)6]3−

(7.8)

  

For this system, the different terms of eqn (7.4) have the following values at 25 °C: k11 = 3 × 104 M−1 s−1, k22 = 7.4 × 102 M−1 s−1, K12 = 1.0 × 102 and f12 calculated using eqn (7.5) is 0.85. The factor f12 is a correction factor for the difference in free energies of the two reactants and is often close to unity, as in this case, substituting all the appropriate values in eqn (7.4), the calculated value of k12 is 4 × 104 M−1 s−1, which compares well with the experimental value of 3 × 104 M−1 s−1. It is worth noting that this redox reaction occurs much faster than exchange of CN− with either of the reactant complexes. Hence electron transfer in this process proceeds without any M–CN bond breaking; this is

390

Chapter 7

Figure 7.1  LFER  illustrating the Marcus cross-­relation in the oxidation of [M(phen)3]2+, [Mbipy)3]2+ and [M(terpy)2]2+ (M = Fe, Ru) and derivatives with Ce(iv) in dilute sulfuric acid.8a–c ●, Fe(ii) complexes; ▲, Ru(ii) complexes; (3, 15), phen; (2, 12), 5-­Me-­phen; (9), 4,7-­Me2-­phen; (1, 11), 5,6-­Me2-­phen; (4), Ph-­phen; (5), 5-­Cl-­phen; (8), 5-­NO2-­phen; (7), 3-­SO3H-­phen; (6), 5-­SO3H-­phen; (14), bipy; (10), 5,5-­Me2-­bipy; (13), terpy.

evidence that outer-­sphere reactions occur without any metal–ligand bond cleavage. The data in Table 7.2 for a number of other outer-­sphere cross-­reactions show fairly reasonable agreement between the calculated and experimental values. Hence one can have great confidence in the Marcus relation; a considerable disagreement between calculated and experimental values can be taken as evidence that a mechanism other than the outer-­sphere mechanism is operational. A modification of the Marcus equation has been proposed for exothermic and endothermic atom transfer reactions.9 An outer-­sphere electron transfer may be represented as in eqn (7.9). First the oxidant (O) and reductant (R) come together to form a precursor complex (a). This is followed by activation of the precursor complex (b), which includes reorganization of solvent molecules and changes in metal–ligand bond lengths that must occur before electron transfer can take place. The final step is dissociation of the ion pair (successor complex) (c) into the product species.   

O  R  (OR)  (OR)*  (O R + )  O + R + (7.9) (a) (b) (c) We can clarify the activation and electron transfer steps by taking the example of electron exchange between the hexaaqua ions of Fe(iii) and Fe(ii)

Electron Transfer Reactions

391

Table 7.2  Calculated  and observed rate constants for outer-­sphere cross-­reactions.8f k12/M−1 s−1 Log K12

Reaction system (Red + Ox) 2+

3+

[Ru(NH3)6] + [Ru(NH3)5(py)] [Ru(NH3)6]2+ + [Co(phen)3]3+ [V(OH2)6]2+ + [Co(en)3]3+ [Mo(CN)6]4− + [IrCl6]2− [Fe(CN)6]4− + [IrCl6]2− [Fe(CN)6]4− + [MnO4]− [Co(terpy)2]2+ + [Co(bipy)3]3+

4.40 5.42 0.25 2.18 4.08 3.40 0.55

Observed 6

1.4 × 10 1.5 × 104 5.8 × 10−4 1.9 × 106 3.8 × 105 1.7 × 105 64

Calculated 4 × 106 1 × 105 7 × 10−4 8 × 105 1 × 106 6 × 104 39

(studied with iron in one of the oxidation states labelled with a radioactive isotope of Fe). The free energy of activation, ∆G‡, for this reaction is 7.9 kcal mol−1. One may wonder why it is not zero since the reaction involves no net chemical change (∆G = 0) and the products are exactly the same as the reactants. In order for the electron transfer to take place, the energies of the participating electron orbitals must be the same, as required by the Franck– Condon principle. In this particular reaction an electron is transferred from a t2g orbital of Fe(ii) (HS) to a t2g orbital of Fe(iii) (HS). The Fe–O bond lengths in these high-­spin (HS) outer-­orbital aqua complexes of Fe(ii) and Fe(iii) are unequal (2.21 Å = r2 and 2.05 Å = r3 in the Fe(ii) and Fe(iii), respectively) and this shows that the energies of the orbitals are not identical. If the electron transfer could take place without an input of energy, the products would be hexaaquairon(iii) with an Fe–O bond length as in the hexaaquairon(ii) species and hexaaquairon(ii) with a bond length as in the hexaaquairon(iii) species; both of these product species could then relax to their ground states with the release of energy. This would clearly violate the first law of thermodynamics, since the production of energy out of nothing is impossible. Hence there ought to be an input of energy for the electron transfer to take place. The actual process occurs with shortening of the Fe–O bonds (r2) in the Fe(ii) species and elongation of the Fe–O bonds (r3) in the Fe(iii) species until the Fe–O bond length is the same [r* = (2r2r3)/(r2 + r3)] in both species6 and the participating orbitals are of nearly the same energy, to permit electron transfer. If (r2 − r*) and (r* − r3) do not exceed the amplitudes of the metal–ligand vibrations, there will be no activation energy requirement for the exchange process. For the [Fe(OH2)6]n+ (n = 2, 3) system, |rn − r*| values are sufficiently large to require activation and hence a slow rate {the [Co(OH2)6]n+ (n = 2, 3) system is similar}.10 For the MnO4−–MnO42− system,11 however, the |rn − r*| values just exceed the amplitude and the calculated activation is only 1.6 kcal mol−1, hence kex = 7 × 102 M−1 s−1 at 0 °C. This is in accord with the Frank–Condon principle, which states that electron rearrangement/transfer is very much faster than nuclear motion. As far as electrons are concerned, nuclear motions are “frozen” during the time period required for electron transfer. This is due to the great difference in the masses of an electron and a nucleus. Electron exchange between hexaaqua species of Fe(ii) and Fe(iii),

Chapter 7

392

both of which are high-­spin outer-­orbital complexes, is quite slow (kex ≈ 1.1 M−1 s−1 at 25 °C), but in the case of hexacyano complexes, both of which are low-­spin (LS) inner-­orbital complexes, the exchange is quite fast (kex ≈ 102 M−1 s−1 at 0 °C). Considerably less reorganization of Fe–CN bond lengths is needed for the low-­spin hexacyano complexes, hence the much faster rate. The same is also true for the [Fe(phen)3]2+–[Fe(phen)3]3+ system (kex > 105 M−1 s−1 at 0 °C), where both species are strong-­field complexes. This is how Libby accounted for the difference in the rates of electron transfer in these and other similar systems.12a The free energy of activation may be expressed as the sum of three terms:   

  

ΔG‡ = ΔG‡c + ΔG‡l + ΔG‡o

(7.10)

where ΔG‡c is the energy needed to bring the oxidant and reductant into a configuration in the transition state in which they are separated by the required distance (for charged reactants this includes work to overcome Columbic repulsion), ΔG‡l is the energy required for bond compression and stretching in the inner-­sphere of the two complex species to achieve orbitals of same energy and ΔGo‡ is the energy needed for solvent reorganization outside the inner coordination sphere, i.e. in outer spheres of two reactant species. The importance of the magnitude of bond distortion in controlling the rate of electron exchange is revealed by comparing the [M(NH3)6]2+–[M(NH3)6]3+ couples of Co and Ru ions. In case of the cobalt complexes, the Co–N bond length is 2.115 Å for Co(ii) and 1.936 Å for Co(iii). This rather large difference requires considerable elongation and compression of the Co–N bonds in the hexaammines to permit electron exchange [transfer from Co(ii) to Co(iii)], and consequently the rate is slow (kex = 1.7 × 10−7 M−1 s−1 at 25 °C). In the case of the Ru species the difference in the Ru–N bond lengths for Ru(ii) and Ru(iii) oxidation states is only ca. 0.04 Å (2.144–2.104 Å) and for this system the rate of electron exchange is much faster (kex = 4 × 103 M−1 s−1 at 25 °C). These cobalt and ruthenium systems are of course not entirely analogous, since in the case of the cobalt system Co(ii) is high spin whereas Co(iii) is low-­spin, but in the case of the Ru system both oxidation states are low spin, hence in these hexaammine complexes the M–N bond lengths are nearly the same as both are of inner-­orbital type. Thus, the change in spin state accompanying electron transfer indirectly becomes a contributing factor for the slow rate of electron transfer. Even without a change in spin state, if the electron transfer is accompanied by a change in bonding from inner-­orbital to outer-­orbital type, this also causes a slow rate of electron transfer, as such a change in bonding requires considerable adjustment of the metal–ligand bond lengths. This explains why outer-­sphere reduction of [Co(NH3)6]3+ by [Cr(OH2)6]2+ is much slower than reduction by [V(OH2)6]2+. As a result of electron transfer to Co(iii), the Cr(ii) species (t2g4) changes to Cr(iii) species (t2g3) with an accompanying change in bonding from outer-­orbital to inner-­ orbital type. However, for V(ii) to V(iii) the configuration changes from t2g3 to t2g2, in both of which the bonding is of the inner-­orbital type. It should be

Electron Transfer Reactions

393

remembered that the metal–ligand bond length is much more susceptible to the oxidation state of the metal in the case of outer-­orbital complexes (as for the hexaaqua complexes) than for the inner-­orbital complexes. This explains why the V–O bond length does not change drastically as one goes from the V(ii) to the V(iii) state. The electron exchange in the case of the Co(ii)–Co(iii) system is slow for hexaammine and trisethylenediamine complexes, and also for trisoxalato complexes of these ions, but the rates are much faster for the hexaaqua and also for trisphenanthroline complexes (Table 7.1). A similar difference is observed in the Ru(ii)–Ru(iii) systems also, where the electron exchange in [Ru(bipy)3]2+/3+ is 105 times faster than in [Ru(NH3)6]2+/3+ at 25 °C (Table 7.1). Remembering that the difference in energy between the HS and LS configurations for both d7 and d6 metal ions in an octahedral field is a function of Dq (see Chapter 1) and considering the fact that LS [Co(OH2)6]3+ is more stable than the HS form by only ca. 2000 cm−1, i.e. ∼6 kcal mol−1, it is possible to excite the low-­spin hexaaquacobalt(iii) to its HS state by expenditure of this small amount of energy and thus allow electron transfer to take place between the HS hexaaquacobalt(ii) and the corresponding Co(iii) species in the excited HS state having the comparable configurations t25eg2 and t24eg2, respectively. For a strong-­field ligand such as phen, the HS [Co(phen)3]2+ can be excited by spending a small amount of energy to the LS state (t26eg1) to permit electron exchange with the LS [Co(phen)3]3+ (t2g6). However, in the cases of ligands such as NH3, en and ox2−, a lot more energy will be needed to excite either LS Co(iii) to HS Co(iii) or HS Co(ii) to LS Co(ii). This may be understood by referring to energy diagrams appropriate for d6 and d7, as shown in Figure 7.2, which depicts variations in energies of HS and LS states with Dq. Hence electron exchange in the complexes of these ligands is very slow. Although reacting species are in no way chemically bonded in the course of the reaction involving an outer-­sphere mechanism, the orbitals of the reactants which extend in space overlap at the closeness of the reactants in the precursor complex, and this permits the electron flow from the reductant to the oxidant by what is referred to as “electron tunnelling”. In an outer-­sphere mechanism the inner coordination spheres of the oxidant and reductant remain intact (without a bridging ligand being commonly shared by both as is true in an inner-­sphere mechanism discussed in the next section). However, in the case of an outer-­sphere reaction, ion association through a counter ion is known to accelerate the reaction. Thus, in the electron exchange in the MnO42−–MnO4− system the rate varies markedly in the following order of the counter cation:13a,b Cs+ > K+ > Na+ > Li+. This specific cation effect is due to an MnO42−⋯M+⋯MnO4− type of ion aggregate formation which facilitates electron transfer.13c Similar results have been reported13c,d in the [Fe(CN)6]3−–[Fe(CN)6]4− system: Cs+ > Rb+ > K+ > NH4+ > Na+ > Li+ and Sr2+ > Ca2+ > Mg2+. Also, in the reaction between aqua and hydroxo-­aqua types of complexes there can be a specific interaction between the oxidant and reductant centres through hydrogen bonding, OH− being more effective than H2O in this regard (see Chapter 5, Section 5.2). Banerjee and co-­workers112a–d

Chapter 7

394

Figure 7.2  Qualitative  representation of the change in energy with Dq showing the

changeover from HS to LS state for octahedral complexes of d4, d6 or d7 ion with increasing Dq. In the region A′–A″ near the crossover point (A) there will be a small energy change in HS–LS inter-­conversion.

reported the catalytic activity order K+ > Na+ > Li+ in several redox reactions. Oxidation of [Fe(CN)6]4− by peroxodiphosphate in acidic medium was reported to be catalysed by [Cu(edta)]2− and based on experimental observations the following reaction scheme was proposed:113 k1

[HFe(CN)6 ]3  [Cu(edta)6 ]2  # [HFe(CN)6 ]2   [Cu(edta)6 ]3  k2

3 2[Cu(edta)]  H2 P2 O8 k  2[Cu(edta)]2   2HPO42 

3

2

The reaction is retarded by [Fe(CN)6]3− as expected on the basis of this proposed scheme. Although the reaction has also been reported to be catalysed by Feaq2+ and Feaq3+, Cuaq2+ retards the reaction, which is astonishing.

7.4  Inner-­sphere Mechanism This was first suggested by Taube and co-­workers1,4 based on experimental evidence in several systems. The reaction shown in eqn (7.11), studied by Taube's group,1a,b is the first established example of an inner-­sphere redox reaction in acidic aqueous solution.‡   

[CoCl(NH3)5]2+ + [Cr(OH2)6]2+ + 5H+ + 5H2O → [Co(OH2)6]2+ + 5NH4+ + [CrCl(OH2)5]2+ (7.11)

  

In an inner-­sphere mechanism, the activated complex for electron transfer contains a bridging ligand, which simultaneously occupies a position in the ‡

 wing to the lability of [Cr(OH2)6]2+ to substitution, it is possible that the bridged complex O [(H2O)5CrIII]–(OH)–CrII(OH2)5]4+ is involved.

Electron Transfer Reactions

395

first coordination sphere of each of the two metal ions between which electron transfer takes place. The formation of this involves the substitution of a ligand in the labile partner by another one in the other partner, which is capable of bridge formation. Some common bridging ligands are H2O, OH−, O2−, O22−, F−, Cl−, Br−, I−, N3−, SCN−, CN−, RCO2−, etc. Almost 200 different bridges were examined in the 1960s.14a Evidence for the formation of a ligand-­bridged intermediate is usually indirect, although very persuasive. In rare cases, the lifetime of the bridged intermediate allows direct observation of its formation in the reaction system. Thus, for the reaction of cis-­[RuCl2(NH3)4]+ with [Cr(OH2)6]2+ forming cis-­[RuCl(NH3)4(OH2)]+ and [CrCl(OH2)5]2+, the participation of the bridged intermediate [(NH3)4ClRu–Cl–Cr(OH2)5]3+ is indicated by spectral observations, which also suggests that the formation of this intermediate and the electron transfer are relatively rapid, so that this intermediate spends most of its lifetime in the RuII–Cl–CrIII form, while the dissociation of this form of the intermediate is rate determining.14b The bridging ligand may simply facilitate electron transfer, but in a number of cases there is also a transfer of the bridging ligand (atom or group transfer) from the oxidant to the reductant, causing a concurrent net electron transfer as well. The bridging mechanism is evidently not effective unless ligand replacement in at least one of the partners is faster than the rate of electron transfer. It should be mentioned that with H2O or OH− as bridging group, the bridge is often formed by hydrogen bonding without any prior substitution in the coordination spheres of the reactants, hence electron transfer may be viewed as an outer-­sphere process in the following cases:

Bridging by OH− is more effective and for the exchange in the Craq3+–Craq2+ system the rate is ca. 105-­fold faster with [Cr(OH)(OH2)5]2+ (kex = 0.7 M−1 s−1 at 25 °C) than with [Cr(OH2)6]3+ (kex ≈ 2 × 10−5 M−1 s−1 at 25 °C). Similar observations have been reported in the Fe(iii)–Fe(ii) and Tl(iii)–Tl(i) systems. The role of a good bridging ligand in accelerating an electron transfer reaction is also illustrated by the fact that reduction of Fe(iii) by Npaq3+ is ca. 100 times faster for FeClaq2+ than Feaq3+.15 Many such examples are known.

7.4.1  Atom (or Group) Transfer Processes Evidence for such processes comes from several sources. In the oxidation–reduction of oxo anions, it is possible to show oxygen atom transfer by labelling with 18 O, provided that it can be demonstrated that rapid exchange with solvent oxygen does not occur under the experimental conditions. In this way, atom transfer has been demonstrated for the oxidation of SO32− with ClO−, ClO2−, ClO3− and BrO3− and also in the oxidation of NO2− with HClO [eqn (7.12)].16   

  



(7.12)

Chapter 7

396 2+

By the ingenious selection of [Cr(OH2)6] as a reducing agent, Taube and co-­workers1 were able to demonstrate the transfer of a large number of univalent atoms and groups. The situation here is ideal for testing if atom/group transfer takes place or not, since Cr(ii) is labile whereas the product Cr(iii) is inert to substitution. Thus, in the reaction of an oxidant M–L with Craq2+, if L is found in the Cr(iii) product, the L must have been transferred during the electron transfer process itself, since any subsequent reaction of Cr(iii) with L released from the reduction product of M–L would in general be slow and this can be ascertained independently in most cases. Thus, [CoCl(NH3)5]2+ reacts with [Cr(OH2)6]2+ at a fast rate, leading to the formation of [Co(OH2)6]2+, [CrCl(OH2)5]2+ and 5NH4. If an outer-­sphere mechanism operates, then one expects the formation of [Cr(OH2)6]3+ as the Cr(iii) product, which may undergo anation by Cl−, forming [CrCl(OH2)5]2+. However, the observed rate constant17a,b for the reduction at 25 °C is 6 × 105 M−1 s−1, whereas the rate for anation of [Cr(OH2)6]3+ by Cl− is exceedingly slow (kan = 2.9 × 10−8 M−1 s−1 at 25 °C).17c Hence, during the time for complete reduction there will be hardly any formation of the [CrCl(OH2)5]2+, which, however, is formed in this reaction as the Cr(iii) product almost exclusively. This indicates that the reduction takes place by a different mechanism in which transfer of Cl to the Cr(iii) centre is synchronous with the electron transfer, and a mechanism was proposed to account for this that involves the formation of a bridged intermediate [eqn (7.13)].1a,b   

k1

[CoCl(NH3 )5 ]2   [Cr(OH2 )6 ]2  #[(NH3 )5 Co – Cl – Cr(OH2 )5 ]4   H2 O k1

  

(7.13)

Because of the extremely labile character of [Cr(OH2)6]2+ to substitution, this reaction proceeds very fast. Electron transfer then occurs from Cr(ii) to Co(iii) in this precursor complex (which is the rate-­determining step) whereby the Co(iii) centre becomes Co(ii) and the Cr(ii) becomes Cr(iii). Since the Co(ii)–Cl bond is expected to be very labile whereas the Cr(iii)–Cl bond is known to be very inert, decomposition of this successor complex into the products takes place by a fast change involving dissociation at the Co(ii)– Cl bond, followed by other fast changes forming the final products (Scheme 7.1). On carrying out the reaction in the presence of added 36Cl− in solution, no uptake of 36Cl− into the product [CrCl(OH2)5]2+ was observed, supporting the mechanism illustrated in Scheme 7.1. A similar mechanism operates in many other such Co(iii)–Cr(ii) systems where transfer of a bridging group takes place simultaneously with electron transfer1c involving various bridging groups such as H2O, OH−, F−, Cl−, Br−, I−, SCN−, CN−, SO42−, PO43−, etc. (OH− transfer has been demonstrated by 18O labelling).18 From observations on the rates of reduction of several Co(iii) complexes of the type [CoL(NH3)5]m+ with [Cr(OH2)6]2+, it is seen that for inner-­sphere reduction of Co(iii) complexes the rate increases dramatically with increasing bridging character of the L. In the case of the hexaammine complex, where an inner-­sphere mechanism is not possible owing to the inability of the bound

Electron Transfer Reactions

397

Scheme 7.1 NH3 to act as a bridging ligand, the rate constant for reduction at 25 °C is 8 × 10−5 M−1 s−1.19 This rate increases to ca. 0.1 M−1 s−1, i.e. by a factor of ca. 1.25 × 103, when L = H2O20 and,compared with the aqua complex, the hydroxo complex (L = OH−) reacts 1.5 × 107 times faster (k = 1.5 × 106 M−1 s−1).20 For L = NCS− (Co–N bond), bridging can occur only through the free S of the SCN−. The S being a “soft” ligand is not preferred by the “hard” Cr(ii), hence bridging is not very effective and as such k = 19 M−1 s−1 (kNCS),21 but for the analogous S-­bonded (Co–SCN) complex bridging occurs through the “hard” N-­bound SCN (Co– SCN–Cr), which is very effective, hence in this case the k value22 is kNCS = 1.9 × 105 M−1 s−1, i.e. kSCN/kNSC = 1.9 × 105 M−1 s−1/19 M−1 s−1 = 104.23 For L = N3− also k = 3 × 105 M−1 s−1.21 In some other cases,23,24 the ratio kSCN/kNSC is ≥ 104. Thus, electron exchange in the [Cr(OH2)6]2+–[Cr(SCN)(OH2)5]2+ system proceeds ca. 2.9 × 105 times faster (k = 40 M−1 s−1 at 25 °C) than in the [Cr(OH2)6]2+–[Cr(NCS) (OH2)5]2+ system (k = 1.4 × 10−4 M−1 s−1 at 25 °C). Atom transfer has also been established in several other systems,1d such as [CrX(OH2)5]2+–[Craq]2+ and [CrX(NH3)5]2+–[Craq]2+. Exchange of ligands in [Cr(OH2)6]3+ with the solvent is slow, but is catalysed by Craq2+. Here also a group transfer with an electron transfer mechanism is involved. For [*Cr(OH2)6]2+ + [CrCl(OH2)5]2+ ⇌ [*CrCl(OH2)5]2+ + [Cr(OH2)6]2+ kex = 9 M−1 s−1 at 0 °C. A similar mechanism is involved in the dissolution of the insoluble anhydrous CrCl3 in water in the presence of traces of Craq2+. The Craq2+-­catalysed complexation of Craq3+ with various ligands also occurs by an inner-­sphere electron transfer process:25   

Craq 2+  L# CrL aq 

2

 Craq 3+# Craq  L  Craq 

5

  

This reaction scheme leads to Rate = k[Craq2+][Craq3+][L] where k is a composite rate constant.

3

  CrL aq   CrL aq 2+  (7.14)

Chapter 7

398

In the Fe(iii)–Cr(ii) reaction in the presence of a halide, the product is Fe(ii) and [Cr(halide)(OH2)5]2+, indicating atom transfer. The reactions of [CrF(OH2)5]2+ and cis-­[CrF2(OH2)4]2+ with Craq2+ are interesting. Both transfer one F− predominantly and the rate parameters are also very similar, suggesting that double bridges are not particularly favoured, even if conditions are suitable for double bridge formation. Often groups not serving as bridging groups find their way into the Cr(iii) products. Thus, in the reaction of [CrCl(NH3)5]2+ with [Cr(OH2)6]2+ in the presence of P2O74−, both Cl− and P2O74− are found in the Cr(iii) product. The rates are also changed by such non-­ bridging ligands. The following type of interaction has been suggested in the transition state: [(NH3)5Co]3+⋯Cl−⋯[Cr(OH2)4]2+⋯P2O4−7 Owing to the high lability of Eu3+,26 inner-­sphere reduction by Euaq2+ does not occur with transfer of the bridging ligand to the Eu3+ product,27 and in such reductions evidence for an inner-­sphere path is indirect. In all the examples cited above, the kCl step (see Scheme 7.1) is rate determining, hence kobs = KkCl [where K = k1/k−1, being the equilibrium constant for precursor complex formation, eqn (7.13)]. Reactions in which the precursor complexes are formed readily (fast process), followed by their transformation to successor complexes, which decompose to form products in a rate-­determining step, are believed to be the most common reaction pattern for electron transfer reactions. However, the inner-­sphere reduction of trans-­[CoCl(en)2(OH2)]2+ by Feaq2+ proceeds by substitution on the labile Fe2+, forming CoIII–Cl–FeII bridged species (precursor complex) followed by rate-­ determining electron transfer forming the bridged CoII–Cl–FeIII (successor complex), which undergoes fast CoII–Cl bond rupture forming FeCl2+ and Co2+ as products.28,29 For a system reacting by such a scheme, if the intermediate formed in a fast pre-­equilibrium is sufficiently stable and transforms to products in a slow (rate-­determining) step, then it has sometimes been possible to detect kinetically a sufficiently stable binuclear precursor complex, as in the reaction of [Co(nta)(NH3)5] with Feaq2+, where the pseudo-­first-­order rate constant attains a limiting value at high [Fe2+], which is in conformity with the reaction scheme shown in eqn (7.15).30   

K

k [Co(nta)(NH3 )5 ]  Feaq 2  #[Co(nta)(NH3 )5 Feaq ]2    Fe(nta)aq  Co aq 2 

(7.15)

  

6

−1

−2

−1

At 25 °C (I = 1 M), K = 1.1 × 10 M and k = 9.4 × 10 s . Examples are known where the successor complex breaks down without any transfer of the bridging ligand, as shown in eqn (7.16).31   

  

k [Co(edta)]2   [Fe(CN)6 ]3#[(edta)Co  NC  Fe(CN)5 ]5  [Co(edta)]

[Fe(CN)6 ]4 

(7.16)

Based on various evidence, arguments have been put forth31 that in the intermediate the metals are present as Co(iii) and Fe(ii) (hence diamagnetic).

Electron Transfer Reactions

399

Therefore, in this system the precursor complex is formed in a fast step followed by its fast transformation into the successor complex that forms the products in a rate-­determining slower step. At 25 °C (I = 0.66 M; pH 5) the values of the equilibrium constant (K) for formation of the precursor complex and rate constant (k) for transformation of the successor complex into the products are K = 670 M−1 and k = 6 × 10−3 s−1. Another similar example is the following overall reaction:32   

[Co(CN)5]3− + OH− + IrCl62− → [Co(CN)5(OH)]3− + IrCl63−



(7.17)

  

The reaction occurs through the formation of a precursor complex [Cl5IrIV– Cl–CoII(CN)5]5−, which transforms into the successor complex [Cl5IrIII–Cl– CoIII(CN)5]5−. The latter undergoes CoIII–Cl bond rupture {as this is less inert than IrIII–Cl; compare the aquation rates of [MCl(NH3)5]2+: M = Co, kaq = 1.7 × 10−6 s−1 at 25 °C;64 M = Ir, kaq = 6.3 × 10−7 s−1 at 95 °C 2} with concurrent capture of OH− from the solution by [Co(CN)5]2−, forming the final product [Co(OH) (CN)5]3−. For this system at 25 °C (I = 0.1 M; [OH−], 1 × 10−3 M), the reported values are k = 4.0 s−1, ΔG‡ = 15.4 kcal mol−1 and ΔS‡ = −4.2 cal K−1 mol−1. However, in this system, concurrent outer-­sphere reaction could not be ruled out. Systems are known, however, where precursor complex formation (k1 step, Scheme 7.1) is rate determining. In such cases, the reaction rate is controlled by the rate of substitution into the coordination sphere of the labile reactant. This is the situation in the reduction of some Co(iii) complexes by [V(OH2)6]2+, hence the rate is almost insensitive to the nature of the bridging group (see the data in Table 7.3). The kinetic parameters given in Table 7.3 closely match those for ligand substitution (replacement of H2O) in [V(OH2)6]2+ (see the data in Table 7.4). Another example of this scheme is in several reductions by Craq2+ having rate constants in the range 106–107 M−1 s−1 (see ref. 3, Table 8.5), for which Halpern (quoted by Taube and co-­workers33) suggested a rate-­limiting substitution on the Craq2+. Orhanovic and Sutin34 reported that Craq2+ reductions of the complexes [FeX(OH2)5]2+ (X = Cl, SCN, N3) have mostly the same rate; at 25 °C, I = 1 M, the 10−7k values are 2.8, 2.9 and 1.0 M−1 s−1, respectively. Table 7.3  Rate  parameters for reduction of some Co(iii) complexes by [V(OH2)6]2+ at 25 °C.a

Complex (oxidant)

Kobs/M−1 s−1

ΔH‡/kcal mol−1

ΔS‡/cal K−1 mol−1

[CoCl(NH3)5]2+ [CoBr(NH3)5]2+ [Co(N3)(NH3)5]2+ [Co(oxH)(NH3)5]2+ cis-­[Co(N3)(NH3)(en)2]2+ cis-­[Co(N3)(en)2(OH2)]2+ trans-­[Co(N3)2(en)2]+ trans-­[Co(N3)(en)2(OH2)]2+

10 25 13 12.5 10.3 16.6 26.6 18.1

— — 11.7 12.2 12.6 12.1 12.2 11.0

— — −14.0 −12.9 −12.0 −12.0 −11.0 −16.0

a

Literature data from various sources (see ref. 3, Table 8.6).

Chapter 7

400

Table 7.4  Kinetic  parameters for ligand substitution in [V(OH2)6] .

2+ 3

Reaction of [V(OH2)6]2+

Rate constant (25 °C)

ΔH‡/kcal mol−1

ΔS‡/cal K−1 mol−1

Water exchange Anation by NCS−

1 × 102 s−1 (kex)a 28 M−1 s−1

16.4 13.5

−5.5 −6.9

k ex/55.5 = 1.8 M−1 s−1, 55.5 M−1 s−1 being the molar concentration of H2O in dilute aqueous solution; this value is more appropriate for comparison with the k value (in M−1 s−1) for anation by NCS−.

a

Hence formation of the precursor complex FeIII–X–CrII is rate determining and this is obviously controlled by the rate of dissociation of a CrII–OH2 bond in [Cr(OH2)6]2+ for which the kex/55.5 value is ∼107 M−1 s−1 (see Chapter 1, Figure 1.3). A third type of inner-­sphere mechanism is where break-­up of the successor complex is rate determining (k2 step, Scheme 7.1). This is generally observed when both the metal centres in the successor complex are substitution inert, as in the reduction of [RuCl(NH3)5]2+ with [Cr(OH2)6]2+, where the Ru(ii) and Cr(iii) in the successor complex [(H3N)5RuII–Cl–CrIII(OH2)5]4+ are both inert to substitution, hence break-­up of the complex is slow and rate determining. As the Cr(iii) product is CrCl2+, break-­up occurs by Ru–Cl bond rupture. The presence of a suitable bridging ligand in the inert oxidant does not automatically guarantee an inner-­sphere reduction by a labile reductant. Even in such a system the outer-­sphere mechanism may operate if electron transfer by the outer-­sphere path can proceed much faster than substitution on the labile metal centre. This is the situation in the [RuBr(NH3)5]2+ and [V(OH2)6]2+ system, where the second-­order rate constant for the reduction is 5.1 × 103 M−1 s−1 at 25 °C, which is very much faster than the anation rate for [V(OH2)6]2+ mentioned above. With a multi-­atomic bridging group, the precursor complex formation may be due to either remote attack or adjacent attack.35 Thus, with [Co(SCN)(NH3)5]2+, adjacent attack on the S gives a rate constant of 8 × 104 M−1 s−1, while remote attack on N gives a rate constant of 1.9 × 105 M−1 s−1 for reductions by [Cr(OH2)6]2+. As a result of these two concurrent paths, 29% of the Cr(iii) product is Cr–SCN (green) and 71% is Cr–NCS; on standing the Cr–SCN slowly rearranges into a more stable Cr–NCS (purple) species.36a In the case of N-­bonded [Co(NCS)(NH3)5]2+, the Cr(iii) product of reduction by [Cr(OH2)6]2+ is 100% [Cr(SCN)(OH2)5]2+, indicating reaction by only remote attack36a on the S of the N-­bonded SCN−. It is generally found21 that reduction of M–SCN2+ by Feaq2+ or Craq2+ is ∼104 times faster than that of M–NCS2+ [M = Co(NH3)5 or Cr(OH2)5]. However, the ratio is much less for reduction by Vaq2+, which supports the view that in this case reduction is substitution controlled.24 Based on X-­ray crystallography, [Co(S2O3)(NH3)5]+ has S-­bonded S2O32−. Its reduction by Craq2+ yields a green Cr(iii) complex having an O-­bonded S2O32−, supporting remote attack.36a Reduction of [Co{NH2CH2C(O)O}(NH3)5]2+ by Craq2+ forms [Cr{O(O)CCH2NH2}(NH3)5]2+, also indicating remote attack.36c

Electron Transfer Reactions

401 III

n+

−

In the reaction of [Co (X)(NH3)5] with CN , two types of behaviour have been observed. When X is a ligand having good bridging properties (such as Cl−, Br−, I−, N3−, NCS−, OH−, S2O32−, etc.), the product is [CoIII(CN)5X]n−, whereas in the case of a ligand having poor or no bridging properties (such as X = NH3, MeCO2−, CO32−, SO42−, PO43−, etc.) the product is [Co(CN)6]3−; when X is F− or NO3−, both types of products are obtained. Ray37 made use of this reaction for the synthesis of [Co(S2O3) (CN)5]4− (isolated as the sodium salt). It has been shown that these reactions are catalysed by Co(ii); the product [CoIII(CN)5X] is obtained due to operation of an inner-­sphere group transfer–electron transfer mechanism, whereas [Co(CN)6]3− is formed due to operation of an outer-­sphere electron transfer mechanism.38 In the presence of excess CN−, Co(ii) exists in solution essentially as [Co(CN)5]3− with traces of [Co(CN)6]4− in equilibrium. The inner-­sphere (IS) reactions (I) involve participation of [Co(CN)5]3−, whereas outer-­sphere (OS) reactions (II) involve participation of [Co(CN)6]4− (Scheme 7.2). Evidently, the rate laws for outer-­sphere and inner-­sphere processes would take the forms expressed in eqn (7.18) and (7.19), respectively.   



Rate = kos[(H3N)5CoIIIX][CoII][CN−]

(7.18)

Rate = kis[(H3N)5CoIIIX][CoII]

(7.19)

  

  

These rate laws have been verified experimentally. These are apparently ligand replacement but are not so in the true sense. Transfer of a bridging ligand from one metal centre to another in an electron transfer reaction is authentic evidence for an inner-­sphere mechanism. However, transfer of a bridging ligand is not an essential prerequisite for an inner-­sphere process. Electron (le−) transfer from Craq2+ to trans-­[(H3N)5Cr–(OH)–CrCl(NH3)4]4+ takes place somewhat faster than to trans-­[CrCl(NH3)4(OH2)]2+ at 25 °C (k = 3.0 and 1.1 M−1 s−1, respectively).39 The reaction products indicate exclusive transfer to the [Cr(NH3)4] centre in the former. The bridging group may serve in various ways to facilitate the electron transfer. A negatively charged ligand would serve to bring two metal ions

Scheme 7.2

402

Chapter 7

closer, reducing their electrostatic repulsion. Furthermore, if the ligand has mobile electrons (e.g. due to π-­bonding), then it may serve as a good “conductor” for the movement of electrons. The mechanism involving the formation of a bridged intermediate accounts for the fact that the electron exchange reaction between [Cr(OH2)6]2+ and [Cr(OH2)6]3+, studied by radioactive isotope labelling, is extremely slow, whereas the exchange between [Cr(OH2)6]2+ and [CrCl(OH2)5]2+ is very rapid and atom transfer occurs in this case.1b,40 In the case of Pt(iv)–Pt(ii) exchange involving a Cl bridge (see Chapter 3), an atom transfer would not be sufficient to bring about a net two-­electron transfer. Conversely, in the reaction shown in eqn (7.20) there are two bridging groups which are transferred, but only one electron is transferred.41   

  

cis-­[Cr(N3)2(OH2)4]+ + *Cr2+aq ⇌ Cr2+aq + cis-­[*Cr(N3)2(OH2)4]+

(7.20)

Haim42 postulated a double-­bridged intermediate to explain the absence of an H+-­catalysed path in the reduction of cis-­[Co(N3)2(en)2]+ by Feaq2+, whereas reduction of the trans isomer proceeds via an H+-­catalysed path. Further work showed43 that cis-­[Cr(N3)2(OH2)4]+ is a partial product in the reduction of cis-­[Co(N3)2(en)2]+ and cis-­[Co(N3)2(NH3)4]+ by Craq2+. Hence the reduction occurs partly through the formation of a double -­bridged intermediate in these cases. A similar situation44 arises in the reduction of the bisformato complex cis-­[Co(O2CH)2(en)2]+ by Craq2+. Huchital45 provided evidence for double-­bridged paths in the reactions of Craq2+ with CrIII-­oxalato complexes, [Cr(ox)3]3− and cis-­and trans-­[Cr(ox)2(OH2)2]−. All these results indicate that electron transfer need not necessarily take place by atom transfer. The bridge merely facilitates electron transfer by providing a suitable reaction path; whether ligand transfer also takes place simultaneously or not depends on the nature of the bonds of the bridging ligand holding the two metal centres, which results from the electron transfer process. Thus, in the reaction of [CoCl(NH3)5]2+ with Craq2+, the bridged intermediate formed is CoIII–Cl–CrII and then, as a result of electron transfer from CrII to CoIII, the former becomes CrIII and the latter CoII. As the CoII–Cl bond is labile, whereas the CrIII–Cl bond is inert, the bridged complex dissociates at the CoII–Cl bond, leading to the formation of CrIII–Cl species and CoII. However, in the reaction of IrCl62− with Craq2+ at 0 °C, 71% of the reaction is by the outer-­sphere path forming [Cr(OH2)6]3+ and IrCl63−, whereas 29% of the reaction is by the inner-­sphere path through the bridged intermediate IrIV–Cl–CrII in which, as a result of electron transfer from CrII to IrIV, the former becomes CrIII and the latter IrIII. However, the IrIII–Cl bond is also inert like the CrIII–Cl bond, so that the bridged complex dissociates 39% at the CrIII–Cl bond, forming [Cr(OH2)6]3+ and IrCl63−, and 61% at the IrIII–Cl bond, forming [IrCl5(OH2)]2− and [CrCl(OH2)5]2+.4e Reaction of [IrBr6]2−–Craq2+ also proceeds by similar inner-­and outer-­sphere paths; the former produces the successor complex [Br5IrIII–Br–CrIII(OH2)5], which undergoes transformation at measurable rates in parallel paths, breaking the Ir–Br bond in one case and the Cr–Br bond in the other, thereby forming a mixture of [IrBr6]3−,

Electron Transfer Reactions 2−

403 3+

2+

46

[IrBr5(OH2)] , [Cr(OH2)6] and [CrBr(OH2)5] as products. However, the reaction of [IrBr6]2−with [Co(CN)5]3− proceeds 100% by the inner-­sphere path, and the successor complex breaks at Ir–Br and Co–Br bonds.46 There are other known examples of systems where reactions occur by both inner-­ and outer-­sphere paths concurrently, as in the reduction of IrCl62− by Coaq2+.47 Another example is the reduction of VO2+ by V2+.48 Reduction of [Co(edta)]−, [Co(Hedta)(H2O)] and [Co(Hedta)Cl]− by Craq2+ proceeds by an inner-­sphere process,49a but reduction by Tiaq3+ by an outer-­sphere process.49b Reduction of [CoX(NH3)5]2+ (where X is a good bridging ligand such as F− or OH−) by Tiaq3+ involves an inner-­sphere process.49c Reduction of RuIII by TiIII is catalysed by CrIII.49d Non-­bridging ligands also influence the rate of inner-­sphere electron transfer,114a as is evident from the data in Table 7.5 on the reduction (inner-­ sphere, Cl− bridging) of complexes of Co(iii) by Feaq2+ in aqueous solution. Extensive studies in this field were made by Benson and Haim114b and some of the rate constants reported by them are given in Table 7.5 (entries 1, 3–5, 9, 11 and 12), along with some data from other sources. The data qualitatively show that k increases with decrease in the ligand field effect (f value)114c of the non-­bridging ligand and the authors offered a plausible reason for this. However, in subsequent work, Bifano and Linck114d reported that in an analogous pair of complexes (see entries 1 and 2), on changing from NH3 (f = 1.25) to py (f = 1.23) the ligand field strength decreases by only ca. 1.6% but (at 25 °C) k increases by a factor of ca. 44, whereas on changing from Cl− (f = 0.78) to Br− (f = 0.72) the ligand field strength decreases by ca. 7.7% but k increases (at 25 °C) by a factor of only 1.1 (see the data for entries 9 and 10). The large increase in rate on changing from NH3 to py was according to them due to Table 7.5  Inner-­  sphere reduction of Co(iii) complexes (X− is a bridging ligand) by hexaaquairon(ii).a

log(k/M−1 s−1)b Entry 1 2 3 4 5 6 7 8 9 10 11 12 a

Complex 2+

cis-­[CoX(en)2(NH3)] cis-­[CoX(en)2(py)]2+ cis-­[CoCl(en)2(OH2)]2+ trans-­[CoX(en)2(OH2)]2+ trans-­[CoX(en)2(NH3)]2+ [CoX(NH3)5]2+ cis-­[CoCl(NH3)4(OH2)]2+ trans-­[CoX(NH3)4(OH2)]2+ trans-­[CoX2(en)2]+ trans-­[CoX(Br)(en)2]+ cis-­[CoX2(en)2]2+ trans-­[CoX2(NH3)4]+

X = Cl

X = Br

X = N3

−4.74 (5.36) −3.10 (5.97) −3.34 (5.62) −0.62 −4.18 −2.85 (5.78) −1.46 ∼1.0 −1.50 −1.44 −2.80 +0.34

−5.21 — −3.55 −1.03 — −3.05 — — −1.74 — — —

−3.64 — −2.16 −0.097 — −2.06 −0.44 +1.38 — — — —

 ef. 3, pp. 315 and 337. R k values at 25 °C; I = 1 M, ClO4−. Values in parentheses are for reduction by hexaaquachromium(ii) ion.

b

404

Chapter 7

the much weaker σ-­donor character of py; an increase in σ-­donor character will reduce the electron affinity of the Co(iii), causing a reduction in rate. The same explanation is tenable in the other cases also. On changing from 2en (en, f = 1.28) to 4NH3 (NH3, f = 1.25) the rate increases by a factor of ca. 78 (data for entries 1 and 6). In the complexes [CoCl(en)2L]2+, on changing from L = NH3 to H2O (f = 1.00) in the cis series (entries 1 and 3) the k value increases by a factor of 25.6, whereas in the trans series (entries 4 and 5) the increase is by a factor of 3.6 × 103. It has further been shown that in the [CoXL5]2+ (X = Cl, Br) series of complexes a change in the non-­bridging ligand makes an identical change in ∆G‡ (and hence in k) in both systems (X = Cl, Br).114e Hence the effect of non-­ bridging ligands is independent of the bridging ligand (X−). The data in Table 7.5 further show that for the reduction of analogous complexes of Co(iii), the rate decreases in the order N3− > Cl− > Br− of the bridging ligand; but the variation is not uniform and depends perceptibly on the non-­bridging ligand; however, the relative effect depends on the non-­bridging ligand, and also on the reductant, as seen from the data on the reduction of cis-­[CoCl(en)2L]2+ by hexaaquachromium(ii) ion (see ref. 3, p. 315) and comparison with the corresponding values for reduction by hexaaquairon(ii) ion (Table 7.5). Hence for the reduction of cis-­[CoCl(en)2L]2+ by Cr(ii) and Fe(ii) the ratio kCr/kFe is of the order of 108 for L = H2O and py, but 1010 for L = NH3. For Cr(ii) reduction of the same Co(iii) complexes, log(k/M−1 s−1) values at 25 °C, I = 1 M for different L are H2O 5.62, py 5.97 and NH3 5.36. Hence the order of rates is NH3 < H2O < py, On the basis of the σ-­effect viewpoint, the position of H2O is anomalous, since the basicity order (as seen from the pKa values of their conjugate acids)114f is NH3 (9.5) > py (5.3) > H2O (−1.7), and this is also expected to be their order of decreasing σ-­bonding ability. This can be accounted for by considering the possibility of π back-­bonding in the case of py. Evidence for such π-­bonding is the values of the formation constants (K) of the mono complexes of Ni(ii) and Co(ii) with NH3 and py in aqueous solution:114g at 25 °C, the log K values are for Ni(ii) 2.8 (NH3) and 1.8 (py) and for Co(ii) 2.1 (NH3) and 1.4 (py). These data show that Kpy/KNH3 = 0.1 (Ni) and 0.2 (Co), although the ratio KHpy/KNHH3 is 6.3 × 10−5; KH is 1/Ka, which is a measure of the proton affinity (basicity). The back-­bonding leads to a bond order of 2 for the M–N(py) bond. This accounts for the much higher stability of the py complexes than is to be expected. The data in Table 7.5 further show that for an analogous pair of complexes, reduction of the trans isomer is faster than that of the cis isomer, and replacement of 4NH3 by 2en also causes a decrease in rate as en is a better σ-­donor than NH3. For reduction of the complexes cis-­[CoCl(en)2L]2+ by Craq2+ and Feaq2+ the ratio kCr/kFe is of the order of 109 for L = H2O and py, but 1010 for L = NH3. For the reduction of cis-­[CoX(en)2L]2+ by Feaq2+ the ratio kH2O/kNH3 is 25.6 (Cl−), 45.9 (Br−) and 30 (N3−). Cannon and Earley115 reported the effect of non-­bridging ligands in the CrII–CrIII exchange system. Sutter and Hunt50 reported asymmetric induction in the oxidation of [Cr(phen)3]2+ with d-­[Co(phen)3]3+, ∼90% of the product being l-­[Cr(phen)3]3+.

Electron Transfer Reactions

405

This result implies an alignment in the transition state which is sensitive to the lack of spherical symmetry in the reactants. Doyle and Sykes51 noted that values of the ratio of the anion-­catalysed and uncatalysed rate constants, k2/k1 (F− > SO42− > Cl− > I−), in the reduction of [(H3N)5Co(µ-­NH2)Co(NH3)5]5+ by Vaq2+ are nearly the same as the ratios reported earlier for the reduction of [Co(NH3)6]3+.52 According to them, the results imply that the anion must be associated with the V2+ centre since preferential affinity of the dinuclear Co(iii) complex bearing a charge of 5+ over the mononuclear Co(iii) bearing a charge of 3+ would be predicted, and that the anion Xn− is remote from the Co(iii) ion. The anion thus serves a more important role than just charge neutralization.52 It has been argued that outer-­sphere reactions of complexes of the type [CoClL5]n+ that show similar relative rates when L is changed, as do inner-­sphere reactions, presumably have a geometrical structure independent of the nature of L.53 In the oxidation of [Co(edta)]2− by IrCl62−, both inner-­and outer-­sphere pathways have been reported. The outer-­sphere path is enhanced in the presence of added [Co(en)3]3+ due to ion-­pair formation {[Co(edta)]2−·[Co(en)3]3+}.54 The Creutz–Taube ion, [(H3N)5Ru–L–Ru(NH3)5]5+ (L = pyrazine), displays redox properties yielding both 4+ and 6+ species.55 Much interest has been focused on the extent to which the pyrazine ring bridging the two Ru centres facilitates electron transfer. A variety of spectroscopic studies support the view that low-­energy electron tunnelling across the bridge delocalizes the charge, making the 5+ ion symmetrical. Studies on many such similar dinuclear complexes such as [(bipy)2ClRu–L–RuCl(bipy)2]3+ (L = pyrazine) as model systems for inner-­sphere precursor complexes allowed the evaluation of Frank–Condon barriers, solvation energies, etc.56 In this bridged complex, the rate constant for electron exchange is 3 × 109 s−1, whereas for the analogous complex with L = 4,4′-­bipyridine the rate constant is 1 × 108 s−1. In such complexes with various bridging ligands (L), the length of the bridge is less important than its nature in influencing the rate of electron transfer.4c The subtle question of the intimate mechanism of electron transfer by the inner-­sphere path, i.e. a detailed idea of how electron density is shifted from the reductant to the oxidant in the bridged intermediate, now remains to be answered. Some progress has been made in this direction by the use of various organic ligands as bridging groups. Basically two types of intimate mechanism have been considered: (i) a chemical mechanism, in which an electron is transferred to the bridging group, thus reducing it to a radical anion, whereupon an electron hopping process eventually carries the electron to the oxidant metal ion, and (ii) a tunnelling mechanism, whereby the electron simply passes from the reductant to the oxidant by quantum mechanical tunnelling through the barrier constituted by the bridging ligand. In using organic bridging groups to investigate the nature of the intimate mechanism, the problem arises of distinguishing between adjacent and remote attack by the reductant on the potential bridging group. In the case of benzoate ion as a bridging group, attack must be on the coordinated carboxyl

Chapter 7

406 18

group and there is evidence (using O as a tracer) to show that this actually occurs on the free carboxyl oxygen [remote attack, eqn (7.21)].57a   



[CrIII–OC(O18)–Ph] + Cr2+aq → Cr2+aq + [CrIII–18OC(O)–Ph]

(7.21)

  

Another example of a remote attack is provided by the following Craq2+ reduction of a cobalt(iii) complex where the Cr(iii) product suggests remote attack on the carbonyl oxygen:

  

(7.22)

   

One might argue that the initial product contains Cr(iii) bonded to N of the pyridine ring (I), which then isomerizes by the action of excess Cr2+ that binds to the carbonyl oxygen; electron transfer from the Cr(ii) to the Cr(iii) forms the O-­bonded Cr(iii) product (II) with loss of the N-­bonded Cr(ii) (Scheme 7.3). However, from actual experiments it is known that this isomerization is very slow and the equilibrium is more in favour of the compound I (Scheme 7.3), since the N of pyridine is a better donor (ligand) than the carbonyl O of the amide (–CONH2) group. Hence the Cr(iii) product formed in the reaction in eqn (7.22) established remote attack. In keeping with this is also the observation that the reaction of [Co(NH3)5(py)]3+ with [Cr(OH2)6]2+ proceeds much more slowly and exclusively by an outer-­sphere mechanism, and the Cr(iii) product is [Cr(OH2)6]3+. Hence the metal-­bound N of pyridine does not have bridging ability. A remote attack in the case of an ambidentate bridging ligand, as expected, leads to linkage isomerization of the ligand in the product complex as in the following examples (k values at 25 °C):58   

  

[Co(CN)(NH3 )5 ]2   [Co(CN)5 ]3 k [Co(NC)(CN)5 ]3  Co2   



k  2.9  102 M 1 s 1

Scheme 7.3

(7.23)

Electron Transfer Reactions

407

[Co(NO2 )(NH3 )5 ]2   [Co(CN)5 ]3 k [Co(ONO)(CN)5 ]3  Co2      

k  3.4  104 M 1 s 1



(7.24)

An indication that a chemical mechanism can be operative in remote attack is afforded by the rate data given in Table 7.6.59 The first four pairs of reactions in Table 7.6 are inner-­sphere in nature, and presumably involve electron tunnelling through the bridging ligand as usual. The results indicate that tunnelling to CoIII is characteristically ∼103–106 times faster than to CrIII. However, the much smaller value of RCo/RCr in the last example in Table 7.6 implies that the electron transfer occurs in a somewhat different manner. It seems plausible that the bridged ligand accepts the electron from the reductant and is thereby converted to a reduced form (a radical ion for which there is convincing experimental evidence),60 this process being rate determining; this is followed by a relatively faster transfer of the electron from the reduced bridged ligand to the M(iii), hence the observed rate is far less sensitive to the nature of M(iii). Unlike in the case of reduction by Cr(ii), which yields substitution-­inert Cr(iii) complexes, the reactions of V(ii) and Eu(ii) yield substitution-­labile products on oxidation. Hence the evidence for atom transfer that has been applied as a criterion of mechanism is not applicable in these cases. Results of investigations by Candlin and co-­workers,21,61 Endicott and Taube,62 Espenson and co-­workers23,63 and Sykes and co-­workers64 have shown that the rate of reduction of [Co(NH3)5L]2+ by Vaq2+ or Euaq2+ is much less sensitive to the nature of L (unlike in the case of reduction by Craq2+). In this respect, the behaviour of [Co(NH3)5L]2+ is similar to that of the reduction by [Cr(bipy)3]2+ (Table 7.7), which surely proceeds by an outer-­sphere mechanism. Hence an outer-­sphere mechanism is implied in reductions by Vaq2+ or Euaq2+. It is worth noting that the ratio of the rate constants for outer-­sphere reduction of the azido and N-­bonded thiocyanato complexes is ca. 3.7 for reduction by [Cr(bipy)3]2+ but the same ratio is ca. 1.6 × 104 for inner-­sphere reduction by Craq2+. Indeed, this criterion has often been applied as evidence for the mechanism.61 The trend in reactivity in reductions of metal complexes at a dropping mercury electrode with the ligand field parameter 10Dq was first recognized by Vlcek.65 Candlin et al.21 showed that the same trend is operative in the Table 7.6  Rate  constants for reductions of analogous complexes of Co(iii) and Cr(iii).

Reactants (M = Co, Cr) 2+

[MF(NH3)5] –Cr2+aq [MCl(NH3)5]2+–Cr2+aq [M(OH)(NH3)5]2+–Cr2+aq [M(OOCMe)(NH3)5]2+–Cr2+aq

[M{–NC5H4(CONH2-­4)}(NH3)5]2+–Cr2+aqa a

Rate ratio (RCo/RCr) ∼106 ∼103 ∼106 ∼103 ∼10

NC5H4(CONH2-­4) is isonicotinic acid amide with the N of the pyridine ring bonded to the M.

Chapter 7

408

Table 7.7  Second-­  order rate constants for the reduction of Co(iii) complexes by various reductants at 25 °C in aqueous solution.

Rate constants for reduction by M2+/M−1 s−1 Oxidant

Craq2+

Vaq2+

Euaq2+

[Cr(bipy)3]2+

[CoF(NH3)5]2+ [CoCl(NH3)5]2+ [Co(N3)(NH3)5]2+ [Co(ONO2)(NH3)5]2+ [Co(NCS)(NH3)5]2+ [Co(SCN)(NH3)5]2+ [Co(O2CMe)(NH3)5]2+

2.5 × 105 6 × 105 3 × 105 ∼90 19 1.9 × 105 0.35

2.6 10 13 — 0.3 30 1.2

2.6 × 104 3.9 × 102 1.9 × 102 ∼1 × 102 ∼0.7 0.12 0.18

1.8 × 103 8 × 105 4.1 × 104 — 1.1 × 104 — 1.2 × 103

outer-­sphere reduction of Co(iii) complexes by Vaq2+ and [Cr(bipy)3]2+. An indication of such a trend in the reduction of [CoCN)5X]3− was also apparent from the pulse radiolysis data of Baxendale et al.66a and results reported by Venerable66b on the outer-­sphere reduction of [Co(CN)5X]3− (X− = CN−, H−, OH−, NCS−, I−) along with similar data on such complexes with X− = NO2−, N3− and Cl− reported earlier,66c which showed a correlation between the rate of reduction of [Co(CN)5X]3− by hydrated electrons (eaq−) with the ligand field parameter 10Dq for the complexes; the second-­order rate constant for reduction by Vaq2+ and [Cr(bipy)3]2+ was found to increase monotonically with decrease in 10Dq (k2 increases with the wavelength for the d–d transition, t2g → eg*). Chen and Gould67a reported rates of reduction of a series of Co(iii) complexes of the type [Co(NH3)5R]3+ (where R represents a variety of organic ligands, including py) by Vaq2+ for comparison of the earlier reported values for reduction by Craq2+ and noted that when R = py (also NH3) and some other cases where R is not a bridging ligand the reactions are obviously outer-­ sphere in nature and the following linear relationship holds: log kV = 1.1log kCr + 1.85 where kV and kCr are the specific second-­order rate constants for reduction by these Maq2+ ions. The observed slope of ∼1 is in agreement with Marcus's theory,67b which predicts a slope of l for outer-­sphere reactions. However, in all other cases where reactions are by an inner-­sphere mechanism, they surprisingly observed that in these cases also a linear free energy relationship holds but the slope is much less than 1: log kV = 0.4log kCr + 0.22 A slope of >1 has also been observed in some other inner-­sphere systems [see eqn (7.25) and (7.26)]. Hence a slope of nearly 1 for such log–log plots has often been used as a criterion for an outer-­sphere reaction. Gould and co-­workers68 reported the reduction of a series of Co(iii) complexes by [Ru(NH3)6]2+, which offers no sites for inner-­sphere attack and hence are undoubtedly outer-­sphere in nature, and compared the results with those of the reductions of these Co(iii) complexes by Maq2+ ions (M = V,

Electron Transfer Reactions

409

3+

Eu, Cr) and Uaq . They found that the second-­order rate constants for reduction of a common oxidant decrease in the order kU > kRu > kV > kEu > kCr and the following relationships hold: log kRu = 1.05log kCr + 0.48 log kRu = 1.05log kEu + 0.96 log kRu = 1.05log kCr + 2.3 log kU = 1.08log kRu + 1.76 log kU = 1.14log kCr + 4.15 However, the reductions of 16 pentaammine complexes of Co(iii) by Craq2+ and Euaq2+, eight of which are structurally constrained to outer-­sphere pathways and the others are known to be mainly inner-­sphere, all fit the following linear correlations with nearly unit slope: log kRu = 0.93log kCr + 0.97 log kU = 1.23log kCr + 3.05 Hence Gould68c warned against the indiscriminate use of this criterion (nearly unit slope) for the assignment of mechanisms. The reduction of platinum(ii) halidoammine complexes by [Ru(NH3)6]2+ and Vaq2+ proceeds by two successive one-­electron transfers. The first step is rate determining, and for this the following correlation was observed:68c log kV = 0.89log kRu + 1.68 These log–log plots represent a good way of storing large amounts of data.68a The intrinsic electron transfer rate constant ket of a redox reaction: k0

ket M  III   N  II # M  III   N  II     M  II   N  III   Products

was used to define the selectivity (s) of a reductant69b using ket values for the reduction of [Co(NH3)6]3+ (A) and [Co(NH3)5(OH2)]3+ (B) by a reductant. Since both oxidants are of similar size and have the same charge, the K0 values will be similar. The selectivity of the reductant was defined as s = log ketA − log ketB If K0 is very small then ket cannot be evaluated; in such a case, the experimentally evaluated k (= K0ket) values are used in place of ket to express s, as the K0 values are similar. The following correlation was found:  s 1.3ER  2.3 where ER is the reduction potential of the reductant. Banerjea and co-­workers70 studied the reduction of [Co(acac)3], [Co(ox)3]3− and [Co(Hedta)(OH2)] by reductants R = hydrazine, hydroxylamine and formic acid in acidic media, for which

Chapter 7

410 +

+

kobs = kH[H ] + (kR + kH·R[H ])[R] where kH, kR and kH·R are the specific rate constants for the acid-­dependent, reductant-­dependent and both acid-­and reductant-­dependent paths, respectively. For the reduction of [Co(acac)3] and [Co(ox)3]3− by these reductants, a linear relationship was observed for the kR and kH·R paths [eqn (7.25) and (7.26)].   

3



3 log kRCo(acac)  2.11  1.62log kRCo(ox )3



Co(ox )3 3 log kRCo(acac)  5.67  2.72log kH·R

  

(7.25)

3

  

(7.26)

12b

On the basis of Marcus's theory, a slope of 1.0 is expected for an outer-­ sphere process and this was indeed observed in many systems by Gould and co-­workers (see earlier), who also observed a slope very different from 1.0 for inner-­sphere processes. Hence these reductions seem to occur by a sort of inner-­sphere process. It is significant that the relative k values (at 50 °C) given in Table 7.8 have been observed. The results in Table 7.8 show that despite the near identity of their reduction potentials, [Co(ox)3]3− (E° = +0.57 V) is reduced much faster than [Co(Hedta)(OH2)] (E° = +0.60 V) and N2H5+. Taking this into consideration and the evidence for the inner-­sphere nature of reductions, it has been proposed that the intermediate in the reduction process results from a proton bridging the Co(iii) centre and the reductant resulting from overlap of the vacant 1s orbital of H+ with a filled t2g (dxy) orbital of Co(iii) and an appropriate unshared filled orbital (sp3) of nitrogen (in the case of N2H5+ and NH3OH+) or oxygen (in the case of HCOOH), as shown in the structure III for [Co(Hedta) (OH2)] and N2H5+ as an example.70d Formation of such an H+-­bridged intermediate will obviously be more favourable with anionic [Co(ox)3]3− than the uncharged [Co(Hedta)(OH2)], accounting for much faster reductions of the former.

Table 7.8  Rate  constants for the reduction of some Co(iii) complexes (see text). Rate constant +

kRN2H5

[Co(ox)3]3−

[Co(Hedta)(OH2)]

[Co(acac)3]

1

13.3

25.3

1

19.4

150.0

kRNH3OH



1

1.7

6.8

3 OH kHNH R



1

N 2 H5+

kHR

N 2 H5+

kR

kRNH3OH kHNHR3OH

428.0

1.26

9.92

4.72

4.00

3.11

1.40



+

kHN2RH5

25



Electron Transfer Reactions

411

Scheme 7.4

An inverse dependence of rate on hydrogen ion concentration was observed in the case of the intramolecular redox decomposition of the Ag(iii) complex of ethylenedibiguanide in acidic media, which suggests that the conjugate base form of the complex is the reactive species.71 However, for reduction of the tetranuclear [Mn4(µ-­O)6(bipy)6]4+ by ascorbic acid in aqueous solution (pH 1.5–6.0), the protonated species [Mn4(µ-­O)5(µ-­OH)(bipy)6]5+ is kinetically more reactive.72 Gould and co-­workers5 reported kinetic studies on the reduction of [Ag(endibigH2)]3+ (endibigH2 = ethylenedibiguanide) by Maq2+ ions (M = V, Cr, Fe, Eu) and also [Ru(NH3)6]2+ and suggested that the reduction proceeds through the formation of Ag(ii) species in the rate-­determining step followed by a faster reduction to Ag(i). The observed rate constants (M−1 s−1 at 20.5 °C, I = 1 M) are 0.67 (Feaq2+), 3.4 ± 0.2 × 105 (Euaq2+), 1.07 ± 0.05 × 106 (Vaq2+), >5 × 107 (Craq2+) and 15 ± 1 ([Ru(NH3)6]2+). All these reactions, despite being first order in the reductant, consume two equivalents of 1e− reductant. This is in agreement with the proposed reaction sequence shown in Scheme 7.4. The authors also studied the reduction of the silver(iii) complex by Tiaq3+, Cuaq+ and [Fe(Me2phen)3]2+ and reported rate constants (M−1 s−1 at 20.5 °C, I = 1 M) of 50 ± 2, > 6 × 102 and 4 × 102, respectively. In these cases, only one equivalent of the reductant is consumed in the formation of Ag+ product. Hence in these cases, the [Ag(endibigH2)]2+ species formed in the rate-­determining step suffers spontaneous decomposition [which is obviously much faster than the rate of reduction of the Ag(ii) by these reductants] forming Ag+ and oxidation products of the ligand (not identified). According to Gould and co-­workers, all of the reductions of [Ag(endibigH2)]3+ are presumably outer-­sphere in nature.

7.5  Comproportionation Comproportionation is also an electron transfer reaction. Rate constants for reactions of the type given in eqn (7.27) show a linear dependence on −∆G°, varied by substitution in bipy and terpy ligands.73   

[RuIV(O)(bipy)(terpy)]2+ + [RuII(bipy)(terpy)OH2]2+ → 2[RuIII(bipy)(terpy)OH]2+  (7.27)   

Chapter 7

412

7.6  Mixed Outer-­ and Inner-­sphere Reactions There are a limited number of systems where the energetics of outer-­and inner-­ sphere reactions are comparable, and both provide concurrent paths for the overall reaction. Determination of the rates of the overall reaction shown in eqn (7.28) and of the isotopic exchange between the Fe(ii) and Fe(iii) species [eqn (7.29)] allows the relative importance of outer-­ and inner-­sphere paths to be assessed; the latter contributes ca. 65% to the Fe2+ + FeCl2+ reaction at 25 °C.74 See also IrX62−–[Cr(OH2)6]2+ (X = Cl, Br) systems (Section 7.4.1).   



Fe2+ + FeCl2+ → Fe3+ + FeCl+ → Fe2+ + Cl−

(7.28)

*Fe2+ + FeCl2+ ⇌ *FeCl2+ + Fe2+

(7.29)

  

  

7.7  E  stimation of Redox Rate Constants for   Inner-­sphere Reactions There is evidence that a type of Marcus cross-­relationship may be applied to inner-­sphere redox reactions.8d,e The rates of the inner-­sphere reactions shown in eqn (7.30) and (7.31) are related as shown in eqn (7.32), where k12 and k13 are the rate constants and K12 and K13 are the equilibrium constants for the reactions shown in eqn (7.30) and (7.31) and k22 and k33 are the rate constants for self-­exchange in the Cr2+–CrCl2+ and Fe2+–FeCl2+ systems, respectively (for f12 and f13, see Section 7.3). The validity of this relation has been shown by the observed and calculated k values (see Table 7.9), which show reasonably good agreement.   



[CoCl(NH3)5]2+ + Cr2+ aq → Products; k12K12f12

(7.30)

[CoCl(NH3)5]2+ + Fe2+ aq → Products; k13K13f13

(7.31)

k12/k13 = (k22K12f12/k33K13f13)½ ≈ k22K12/k33K13)½

(7.32)

  

  

  

Table 7.9  Calculated  and observed rate constants (M−1 s−1) for Craq2+ reductions of Co(iii) complexes using eqn (7.32) and experimental data for reductions by Feaq2+.8c log k

Oxidant 2+

[CoCl(NH3)5] [CoF(NH3)5]2+ cis-­[CoCl(NCS)(en)2]+ trans-­[CoCl(NCS)(en)2]+

Observed

Calculated

6.4 5.9 6.3 6.4

6.6 5.5 5.7 5.6

Electron Transfer Reactions

413

The same approach directed to the following three inner-­sphere reactions leads to k2 ≈ (k1k3)½, which is of course not valid, as is evident from the experimental values (at 25 °C) of k1, k2 and k: [Cr(NCS)]2   Cr 2   Cr 2   [Cr(SCN)]2  : k 1 1.4  104 M 1 S1 [*Cr(N 3 )]2   Cr 2   *Cr 2   [Cr(N 3 )]2  : k 2  6.1M 1 S1 [Cr(SCN)]2   Cr 2   Cr 2   [Cr(NCS)]2  : k 3  40M 1 S1

However, taking into account the presence of precursor complexes, a better understanding is possible.75 Precursor complex formation is also a feature of many redox reactions, e.g. oxidations by CrVI in acidic solution.76 It is of interest that for the outer-­sphere reduction of a series of Co(iii) complexes cis-­and trans-­[CoCl(en)2(L)] (L = H2O, NH3, py, Cl−) by [Ru(NH3)6]2+ and their inner-­sphere reduction by [Fe(OH2)6]2+ (involving bridging Cl−), the ratio R = kRu/kFe is fairly constant (log R values at 25 °C are in the range 5.4–5.9).77

7.8  E  lectron Transfer Reactions in Heterogeneous Systems Electroreductions at a cathode are examples of this class. Polarography affords a means of determining the rates of such processes. For an irreversible diffusion-­controlled process A + ne− → product, the half-­wave potential, E½, is related to the rate constant for electron transfer by the following expression:   



E½  E  

2.303 RT  nF

  

log 0.886kE  (t / D )½

(7.33)

where kE° is the rate constant at the standard potential E° of the system, α is the transfer coefficient and the other terms have their usual significance (see Chapter 2).78 For the purpose of comparison of a number of analogous systems, it is best to consider the rate constant values at a rather arbitrarily chosen reference potential Er, using eqn (7.33) in the following form:   



  

E½ Er 

2.303RT  nF

log 0.886kr (t / D )½



(7.34)

A comparison of the polarographic and spectroscopic behaviours of a number of substitution-­inert complexes of Co(iii), Cr(iii) and Rh(iii) is interesting. In such complexes, the lowest unoccupied orbital is an antibonding eg orbital of very low electron affinity, so that electron transfer from the electrode to this orbital is impossible. A change of configuration to a suitable excited state is necessary to precede the electron transfer, and the energy of the excited state may be regarded as a function of the difference between the ground and excited states of the M(iii). Hence the energy of the transition state depends on the ligand field strength of the complex, i.e. the energy

Chapter 7

414

difference (10Dq) between the t2g and eg levels. The transition state need not necessarily correspond to a spectroscopic excited state, since the electrode field is a contributing factor in its formation. If Dq is large, the species may not be reduced, as is observed in the case of [Co(CN)6]3−. On the other hand, [Co(NH3)6]3+, having a much lower Dq, shows a reduction wave. Replacement of one of the ligands with another ligand L will split the t2g and eg levels to give a smaller difference between the ground and excited states, and this makes the complexes [CoIIIL(NH3)5]n+ more easily reducible and E½ shifts to a less negative value than for [Co(NH3)6]3+. The decrease in the energy difference between the levels is sufficient to reduce considerably the activation energy for the rearrangement process, so that even for complexes of the type [CoIII(CN)5L]n−, reduction waves are observed. In agreement with these ideas, Vlcek65 observed that the greater the separation of the parent ligand (L) and the other ligands in the spectrochemical series, the greater is the shift (towards less negative values) in E½. The plot of E½ (for the process CoIII → CoII) versus the difference, ∆ν1, in the wavenumber of the long-­ wavelength d–d absorption band (first band) of the complexes CoIIIA6 and CoIIIA5B was found to be linear for complexes of series A = NH3 and B = H2O, NO3−, NO2−, F−, Cl−, Br−, NCS−, A = NH3 and B = C2O42−, SO42−, S2O32−, and also A = CN− and B = H2O, Cl−, Br−, I−. The activation energy of the aquation of these complexes versus ν1 also follows the linear trend. On the basis of this observation, a model of the electrode process was developed. In the course of the electrode process, a change in the electronic configuration of the complex takes place that is proportional to tetragonal splitting caused by ligand B. The ligand B also influences the ligand trans to itself in the complex, this being more influenced by the electrode field. Similar results were also obtained for the ammine complexes of Cr(iii) and Rh(iii). Using polarography, Banerjea and Banerjee79 determined E½, kr and the corresponding activation energy, Ea, for the single-­step diffusion-­controlled irreversible 2e reduction of [Cu(en)2]2+ and [Cu(gly)2] {these are likely to be present in solution as tetragonally distorted [CuL2(OH2)2] (L = en or gly)} at a dropping mercury electrode. The results (Table 7.10) indicate that the complex [Cu(gly)2], having a much lower Dq value, is reduced more readily than [Cu(en)2]2+. Similar studies on a number of Ni(ii)80a and Co(ii)80b complexes (Tables 7.11 and 7.12) have been reported by Banerjea and co-­workers. These complexes undergo a single-­step, irreversible and diffusion-­controlled 2e reduction. Table 7.10  Polarographic  data for [Cu(en)2]2+ and [Cu(gly)2] complexes. Complexa

Wavenumber of principal ligand field bandb/cm−1

E½/V vs. SCE, 10−2k/cm s−1, at −0.7 at 25 °C) V vs. SCE, at 25 °C) Ea/kcal mol−1

[Cu(en)2]2+ [Cu(gly)2]

18 200 15 800

0.368 0.230

a

I n 0.1 M KNO3 solution; glyH = glycine. This is related to the Dq value of the complex.

b

0.32 7.68

14.2 11.7

Electron Transfer Reactions

415

Table 7.11  Characteristic  features of the electroreduction of some nickel(ii) complexes at 25 °C.

Complex

10Dq/cm−1

−E½/V vs. SCE

kb/cm s−1

Eab/kcal mol−1

[Ni(OH2)6]2+ [Ni(NH3)4(OH2)2]2+ [Ni(en)3]2+ [Ni(gly)2(OH2)2] [Ni(Hedta)(OH2)]−

8500 10 000 11 200 10 000 10 000

0.997 1.030 1.345 1.163 1.535

1.80 × 10−3 3.5 × 10−4 1.6 × 10−6 1.2 × 10−6 6.13 × 10−9

19.9 20.7 26.9 14.6 15.2

a

a

e n = ethylenediamine; glyH = glycine. k and Ea values at −1.0 V vs. SCE in 0.1 M KNO3 solution.

b

Table 7.12  Electroreduction  of some cobalt(ii) complexes in 1 M KCl solution, 25 °C.

Complex

10Dq/cm−1

−E½/V vs. SCE

ka/cm s−1

Eaa/kcal mol−1

[Co(OH2)6]2+ [Co(NH3)4(OH2)2]2+ [Co(en)3]2+

9300 10 500 11 500

1.183 1.246 1.353

33.41 7.29 0.21

26.9 31.2 33.7

k and Ea values at −1.5 V vs. SCE.

a

From a comparison of k and E½ values, it is seen that for Ni(ii) complexes of different charge types (viz. 2+, 0 and 1−) having the same Dq value, the ease of reduction decreases in the sequence [Ni(NH3)4(OH2)2]2+ > [Ni(gly)2(OH2)2] > [Ni(Hedta)(OH2)]−, i.e. in the order of decreasing electron affinity of Ni(ii) in the complexes. Again, in the case of cationic complexes having a charge of 2+, the ease of reduction decreases in the sequence [Ni(OH2)6]2+ > [Ni(NH3)4(OH2)2]2+ > [Ni(en)3]2+ with increase in Dq value in the same sequence, i.e. k decreases with increase in Dq (Table 7.11). A plausible explanation for these observations has been offered as follows.80b For the six-­coordinate complexes of M(ii), the d electrons are accommodated in different orbitals as (dxz)2(dyz)2(dxy)n < ( σ *z2 )1( σ *x2  y2)1 [n = 1 for Co(ii) and 2 for Ni(ii)]. At the electrode surface, if the complex is oriented with its XY-­plane perpendicular to the electrode surface, then the σ *x2  y2 orbital would be destabilized by the electrostatic field of the electrode. The same also holds good for the dxy orbital, which will be destabilized almost to the same extent and the electron distribution will change to (dxz)2(dyz)2 < ( σ *z2 )n < (dxy)1 < ( σ *x2  y2)1; the difference in energy between the dxy and σ *x2  y2 will be related to Dq. The dxy orbital is suitably oriented to accept electrons from the electrode surface to cause the reduction. However, for this process to occur, the dxy orbital has to be vacated by transferring the electron to be paired in the σ *x2  y2 orbital. The energy required for the process would naturally contribute to the activation energy, thus accounting for a relatively slow rate of electron transfer, which gives rise to an irreversible wave, and the rate will decrease in the sequence of increasing Dq. The same explanation can account for the observed behaviour of tetragonally distorted six-­coordinate complexes of d9 Cu2+, although a somewhat

Chapter 7

416

Table 7.13  Electroreduction  of some Ni(ii) and Cu(ii) complexes of 1,3,6,9,11,14-­h exaazacyclohexadecane.

kb/cm s−1

Eab/kcal mol−1

Complex

−E½/V vs. SCE

[NiL]2+ [CuL]2+ 1st wave 2nd wave

1.350

1.65c

17.78

0.624 1.352

31.22d 0.465c

26.34 12.14

a

a

 = 1,3,6,9,11,14-­hexaazacyclohexadecane. L k and Ea values in 0.1 M KCl solution at 25 °C. c At a potential of −1.6 V vs. SCE. d At a potential of −1.0 V vs. SCE. b

different explanation was offered earlier.79 However, since in the case of d9 Cu(ii) complexes the dxy orbital will be doubly occupied, one electron has to be promoted to 4s also to permit the 2e reduction in a single step. The square-­planar complexes of Cu(ii) and Ni(ii) formed by the macrocyclic ligand 1,3,6,9,11,14-­hexaazacyclohexadecane are reduced irreversibly at a dropping mercury electrode.81 The Ni(ii) complex gives a single diffusion-­ controlled wave due to the irreversible 2e reduction process, but the Cu(ii) complex gives two well-­separated irreversible diffusion-­controlled waves corresponding to two consecutive 1e transfer processes (Table 7.13). Both of these complexes are reduced much faster than the complexes of these M2+ with gly−, en, etc. However, it is not clear why these square-­planar complexes are reduced irreversibly whereas the low-­spin square-­planar [Ni(biguanide)2]2+ is reduced reversibly.80a

7.9  Solvated Electrons82 Blue solutions of alkali metals in liquid ammonia have been known since the very early days83 and the presence of solvated (ammoniated) electrons in this solution was postulated by Kraus83b from observations of the electrical conductance of such solutions. Gibson and Argo84 observed that the blue colour of solutions of alkali and alkaline earth metals in different amines (MeNH2, EtNH2, etc.) has a similar absorption spectrum, indicating that the same species is responsible for the blue colour, and suggested this to be the solvated electron. Since 1958, chemists realized that the primary reducing species formed in water irradiated with ionizing radiation (such as X-­rays and γ-­rays) is the hydrated electron (eaq−) identified by Hart and Boag,85 and not the hydrogen atom: h H2 O   Haq    OH  eaq  This is the most useful method of preparation of eaq− and is applicable over the entire pH range. The same species is also formed in some photochemical processes. Thus, solutions containing reducing ions such as I− or [Fe(CN)6]4− on irradiation with UV light (253.7 nm) yield hydrated electrons; similarly, a metal ion M2+, which can be a first-­row transition metal ion, in solution also

Electron Transfer Reactions

417 −

•

(z+1)+

facilitates the generation of eaq as it can capture the OH, forming M + OH−. The OH− combines with the H+, forming H2O. Hence the net reaction is as shown in eqn (7.35).   



  

h Maq Z +    Maq( Z  1)   eaq  H2 O



(7.35) 86

Thus, radiolysis of a solution containing Cu(ii) generates Cu(iii). Solvated electrons may also be generated in other polar solvents, such as alcohols and aliphatic amines. The various processes that produce solvated electrons are the following:    1. Solvation of radiolytically generated secondary electrons. 2. Photochemical excitation leading to photolysis. 3. Dissolution of alkali metals in liquid ammonia, when solvated electrons are formed in equilibrium with solvated M+ ions in the medium. Similar behaviour has also been observed in ethylenediamine. 4. Interaction of OH− ions with atomic H in solution of pH > 12 produces hydrated electrons [eqn (7.36)] with almost 100% efficiency.87 For this process, the rate constant k is 2.2 × 107 M−1 s−1 at 25 °C.88 This reaction was corroborated spectrophotometrically using flash photolysis (radiolysis). The reaction is reversible89 and for the back-­reaction k is 2.3 × 1010 M−1 s−1.   



  

k OH  H   H2 O  eaq 

(7.36)

   The hydrated electron has an absorption band in the red region (λmax = 720 nm; εM = 15 800 M−1 cm−1). At 25 °C the diffusion coefficient of eaq− is 4.7 × 10−5 cm2 s−1 and it is therefore less mobile than H+ (9.5 × 10−5 cm2 s−1) in aqueous solution. The fact that the hydrated electrons were not detected earlier is because of their evanescent and elusive nature (see the data below). eaq− + eaq− → H2 + 2OHaq− (k = 1 × 1010 M−1 s−1 at 25 °C, pH ≈ 11) eaq− + H2O → H + OH− (k = 16 M−1 s−1 at 25 °C, pH = 8.4) eaq− + H3Oaq+ → H + H2O (k = 2 × 1010 M−1 s−1 at 25 °C, pH = 2−4) Thus, in acidic solutions the reducing species that react with other solutes present in the solution are effectively the same, whether produced directly or by the reaction in eqn (7.37).

  



H2 + OH → H + H2O

(7.37)

  

Chemical means of differentiating the two reducing species, viz. H atom and eaq−, are known. N2O reacts with eaq− and with H atoms to produce N2 at rates that differ by a factor of about 104. It should be noted, however, that H does not always act as a reducing agent; in strongly acidic solution it oxidizes

Chapter 7

418 2+

−

3+

Fe to Fe [eqn (7.38)]. N2O and O2 are scavengers for eaq [eqn (7.39) and (7.40)].90,91   



H + Haq+ + Feaq2+ → H2+ Feaq3+

(7.38)



N2O + eaq− → N2 + OH + OH−

(7.39)

O2 + eaq− → O−2

(7.40)

  

  

  

Baxendale et al.92a studied the rate of reduction of 3d-­block aqua metal ions, M(OH2)6]2+, with eaq− and found the following order of reactivity (k values at 25 °C): Mn2+ (k = 7.7 × 107 M−1 s−1) < Fe2+(k = 3.5 × 108 M−1 s−1) < Zn2+ (k = 1.5 × 109 M−1 s−1) < Co2+ (k = 1.2 × 1010 M−1 s−1) < Ni2+(k = 2.2 × 1010 M−1 s−1) < Cu2+(k = 3.0 × 1010 M−1 s−1) < Cr2+(k = 4.2 × 1010 M−1 s−1). This is not the order of electron affinity of M2+, which follows the sequence Mn2+ < Fe2+ < Cr2+ < Co2+ < Zn2+ < Ni2+ < Cu2+. The reported93 values of 10−10k (M−1 s−1) for reduction of several Maq3+ by eaq− {Euaq3+ 6.1, Ybaq3+ 4.3, Smaq3+ 2.5, Tmaq3+ 0.3, Craq3+ 6, Coaq3+ >10 and92 [Co(NH3)6]3 + 9} conform to the following equation:   



  

10 10 k  8.0  3.5EM 3  M2

(7.41)

These results have been of great significance in understanding the mechanism of electron transfer to a metal centre for its reduction to a lower oxidation state. The estimated value94 of E° for   



eaq− + Haq+ → ½H2 (g)

(7.42)

  

is +2.70 V. Some representative values of 10−10k (M−1 s−1) at 25 °C for M3+ ions and their complexes are as follows: Maq3+: Al, 0.20; La, 0.034; Pr, 0.029; Nd, 0.059; Gd, 0.055; Sm, 2.5; Yb, 4.3; Tm, 1.3; Eu, 6.1; Cr, 6.0. [M(NH3)6]3+: Co, 9.0; Ru, 7.4; Rh, 7.9; Os, 7.2; Ir, 1.3. [Cr(en)3]3+, 5.3; [Cr(ox)3]3−, 1.8; [Cr(CN)6]3−, 1.5; [Cr(edta)]−, 2.6; [Co(bipy)3]3+, 8.3; [Co(ox)3]3−, 1.2; [Co(NO2)6]3−, 5.8; [Co(CN)6]3−, 0.27; [CoCl(NH3)5]2+, 5.4; cis-­[CrCl2(en)2]+, 7.1; trans-­[CoCl2(en)2]+, 3.2; cis-­[Co(NCS)2(en)2]+, 6.9; cis-­[Cr(NCS)2(en)2]+, 4.2; cis-­ [Cr(ox)2(OH2)2]−, 1.3; [Cr(OH)4(OH2)2]−, 0.02. The hydrated electron reacts with Co(iii) complexes of the type [Co(CN)5X]3− − (X = CN−, H−, NO2−, NCS−, OH−, N3−, Cl− and I−), the rate constant k2 increases monotonically with decreasing 10Dq and the mechanism involves an outer-­ sphere process;66b k2 corresponds to the rate law −d[complex]/dt = k2[complex][eaq−] where complex = [Co(CN)5X]3−. The k2 values at 20 °C and pH 7–13 are given in Table 7.14. Based on consideration of the structure of the hydrated electron having a high binding energy of 1.7 eV, Venerable66b put forth arguments that inner-­ sphere reduction by eaq− is unlikely and that outer-­sphere reaction is feasible.

Electron Transfer Reactions

419

Table 7.14  k 2 values for some [Co(CN)5X] complexes at 20 °C and pH 7–13. 3−

X−

k2/M−1 s−1

CN− H− NO2− NCS− OH− N3− Cl− I−

5.4 × 109 6.7 × 109 8.0 × 109 1.6 × 109 1.2 × 1011 1.3 × 1010 1.8 × 1010 2.1 × 1010

On consideration of the observed effect of ligands on the rates of reduction by eaq− of a number of complexes, Anbar and Hart93b concluded that the rates increase in the following sequence of the ligands, which is the same as for outer-­sphere reductions and reductions at the dropping mercury electrode of Co(iii) complexes of the type [CoX(NH3)5](3−n)+: OH− < CN− < NH3 < H2O < F− < Cl− < Br− < I− The reactions are ∼20% slower in D2O. In a hydrated electron the immediate environment of e− has four H2O arranged tetrahedrally around the e− having a strong electrostatic attraction with the H atoms of highly polar H2O molecules bearing a δ+ positive charge. This accounts for the high binding energy of 1.7 eV for eaq−.66b

7.10  Oxidative Addition Reactions95 Oxidative addition and its reverse, reductive elimination, are two important classes of reactions in organometallic chemistry that are involved in many homogeneous processes in solution such as the Monsanto process and alkene hydrogenation using Wilkinson's catalyst. It is believed that oxidative addition reactions are also involved in heterogeneous catalysis such as hydrogenation catalysed by platinum metal. In oxidative addition, a low-­valent transition metal complex reacts with a molecule XY to yield a product in which both the oxidation number and coordination number of the metal are increased [a typical example is given in eqn (7.43)].   

(7.43)    

The d8, four-­coordinate, 16e starting complex is converted into a d6, six-­coordinate, 18e product by addition of Cl2; the oxidation number of Ir increases from I to III. Consistent with this change, νC=O increases from 1967

Chapter 7

420 −1

to 2075 cm . The Cl–Cl bond is broken with concomitant formation of two Ir–Cl bonds. Oxidative addition products of carbonyl complexes exhibit increased values of νC=O compared with the starting complexes. Various XY molecules react in this fashion (Table 7.15). However, cases are known where XY adds (bond formation with M) without X–Y bond rupture; hence this is non-­oxidative addition, examples being the formation of W(CO)2 (PiPr3)2(H2)96 and Ir(X(CO)(O2)(PPh3)2.97 The high C–H bond energy (∼470 kJ mol−1) and low polarity make C–H bonds relatively unreactive in general. Oxidative addition of C–H forms an M–C bond and opens the way to further reaction of organic groups on the metal (insertion, for example). Addition to a metal is said to “activate” the C–H bond. Oxidative addition of H2 affords a way of activating the fairly strong H–H bond (430 kJ mol−1) for reaction by first coordinating it to a metal as two hydrido ligands that could then undergo further reaction. The first step in oxidative addition of H2 might be regarded as coordination of the intact molecule in an η2 fashion.96 That the formation of η2-­H2 complexes represents the first step in H–H bond breaking is seen by the fact that they are sometimes in equilibrium in solution with dihydrido complexes, as shown by variable-­ temperature NMR spectroscopy of [W(CO)3(PiPr3)2(H2)], which undergoes ∼15% conversion into [W(H)2(CO)3(PiPr3)2] in solution in the equilibrium state. A water molecule (average bond energy 493 kJ mol−1) is activated by oxidative addition to [Ir(PMe3)4]+ forming cis-­[Ir(H)(OH)(PMe3)4]+; the Ir(i) is converted to Ir(iii) in this process.98a Other Group 8, 9 and 10 metal complexes having a d8 configuration also undergo oxidative addition; the tendency to become oxidized to d6 increases on going down a triad or to the left as shown in Scheme 7.5.98b There are known cases where early transition metals undergo oxidative addition, as in the following example (d6 W0 → d4 WII):   



[W(CO)4(bipy)] + SnCl4 → [WCl(SnCl3)(CO)2(bipy)] + 2CO

(7.44)

  

Table 7.15  Species  that can be added oxidatively.

Atoms that become separated

Atoms that remain attached

X–X: H2, Cl2, Br2, I2, RSSR, etc. C–C: Ph3C–CPh3, (CN)2, C6H5CN, MeC(CN)3 H–X: H–X (X = Cl, Br, I, CN, ClO4, RS, etc.), HC≡CPh, R3SiH, etc. C–X: MeI, PhI, CH2Cl2, CCl4, AcCl, PhCH2COCl, etc. M–X: HgCl2, MeHgCl, R3SiCl, RGeCl3, R3SnCl, Ph2BX, (PPh3)AuCl, etc. Ionic: PhN2+, BH4−, Ph3C+, BF4−

O2; NO; SO2; CF2=CF2, RC≡CR'; RNCO; RNCS; RN=C=NR′ RCON3 R2C=C=O, CS2, (CF3)2CO, etc.

Electron Transfer Reactions

421

Scheme 7.5

Scheme 7.6 Some reactions of the main-­group elements can be viewed as oxidative additions, e.g. the addition of F2 to BrF3 and PF3 yielding BrF5 and PF5, respectively.    (i) One-­electron oxidative addition [Co(CN)5]3− undergoes free-­radical oxidative addition in one-­electron steps 7 (d ↓d6) (Scheme 7.6). Some dinuclear complexes undergo oxidative addition at two centres (d8↓d7) [eqn (7.45)].   



[Rh{µ2-­CN(CH2)3NC}4Rh]2+ + 4Br2 → [BrRh{µ2-­CN(CH2)3NC}4RhBr][Br3]2 (7.45)

  

(ii) Addition of oxygen – an intermediate case Some molecules XY undergo oxidative addition without cleavage of the X–Y bond. A good example is the addition of O2 to trans-­[IrX(CO)(PPh3)2] forming [IrX(CO)(PPh3)2(η2-­O2)] (IV).99 This has a six-­coordinate IrIII or IrII containing η2-­O2− depending on the nature of X. As usual, structural data can help make the distinction. For X = Cl, the O–O distance is 1.30 Å, which is comparable to that in the superoxide ion O2− (1.28 Å).100 This indicates transfer of electron density from the metal into the π* orbitals of O2. When X = I the O–O distance is 1.51 Å, which is very near to that in organic peroxides RO–OR (1.49 Å).101 Obviously, as the metal complex becomes a better base, an entire pair of electrons is donated to dioxygen, giving an IrIII complex. These results show that all degrees of electron donation are possible.

  

Chapter 7

422

Table 7.16  ν C=O for some trans-­[IrCl(XY)(CO)(PPh3)2] complexes . a

XY

νC=O/cm−1

XY

νC=O/cm−1

HCl I2 Br2 ICl Cl2

2046 2067 2072 2074 2075

O2 NCCH=CHCN MeI F2C=CF2 (NC)2C=C(CN)2

2015 2029 2047 2049 2054

In free CO νC=O = 2143 cm−1, but in trans-­[IrCl(CO)(PPh3)2] νC=O = 1967 cm−1.

a

Table 7.16 presents several adducts of trans-­[IrCl(CO)(PPh3)2] with neutral molecules in which electron density is transferred from the metal, as seen from the νC=O values.95d Owing to oxidative addition, the oxidation state of Ir tends to increase from +1 to +3, thereby weakening the Ir → CO π-­bond, consequently strengthening the C–O bond and hence causing an increase in νC=O. The rate of O2 addition is also dependent on X; the rate increases in the order I > Br > Cl with increase in ∆H‡ in the reverse order.102 Because all X–Y bonds are not broken in the adduct, we should formally label these as complexes of IrI. The actual extent of electron transfer can be determined only from structural and spectroscopic data. The lowering of νC=O indicates increased metal to CO back-­bonding leading to increased electron density in the π* orbital of CO that lowers the C–O bond order. In contrast, νC=O for Ag(CO)+ is actually higher (2204 cm−1),103 indicating strengthening of the C–O bond (increase in bond order) on coordination to Ag+; consistent with this, the C–O distance in Ag(CO)+ is 1.077 Å compared with 1.128 Å in free CO.

7.10.1  Mechanisms of Oxidative Addition95d There are four distinct mechanisms for such processes, which are briefly discussed below.

7.10.1.1 Concerted Pathway This operates in cases of non-­polar reagents A–B such as H2, R–H, having a C–H bond and also those with an Si–H bond. These lack π-­bonds and hence a three-­centred σ complex is formed (Scheme 7.7, step a), followed by intramolecular A–B bond cleavage (Scheme 7.7, step b) as a result of back-­donation from M to the σ* orbital of A–B, to form the oxidized complex in which the

Scheme 7.7

Electron Transfer Reactions

423

resulting fragments will be mutually cis (but subsequent isomerization to the trans form may occur in favourable cases). Sometimes the intermediate is stable and step b does not occur. The reaction of H2 with Vaska's complex, trans-­[IrCl(CO)(PPh3)2] (16e species), forming the 18e species [IrCl(H)2(CO)(PPh3)2] having the two H− ligands in mutually cis positions is a typical example of this mechanism:103

Many C–H bond activation reactions also follow a concerted mechanism involving the formation of an intermediate having agostic interactions. The C–H approaches the M with the H pointing towards the M and then the C–H bond pivots around the H to bring the C closer to the MN in a side-­on arrangement, followed by C–H bond cleavage forming H–M–C.104 The addition occurs with retention of stereochemistry at carbon.

7.10.1.2 Bimolecular Associative (SN2) Pathway In the reaction of a square-­planar complex ML4 with an alkyl halide such as MeI [eqn (7.46)], nucleophilic attack by the metal centre at the less electronegative atom in the reagent, R–X, leads to cleavage of the R–X bond forming [M–R]+ species and X−; this is followed by rapid coordination of the X− to the metal centre, obviously in the vacant coordination site trans to R. This mechanism is generally assumed in the addition of polar and electrophilic reactants such as alkyl halides (RX), benzyl and acyl halides and halogens (X2). In some cases, the product of the first step is stable [eqn (7.47)].   

(7.46)    



[IrCO(CpL] + MeI → [IrI(Me)(CO)(Cp)L]

(7.47)

  

This is better considered as electrophilic attack on metal rather than oxidative addition. The more nucleophilic the metal, the greater is its reactivity in SN2 addition, as illustrated by the following reactivity order for Ni(0) complexes: [Ni(PR3)4] > [Ni(PAr3)4] > [Ni(PR3)2(alkene)] > [Ni(PAr3)2(alkene)] > [Ni(cod)2] (R = alkyl, Ar = aryl). Steric hindrance at carbon slows the reaction and hence

Chapter 7

424 i

the reactivity order: MeI > EtI > PrI. A better leaving group X at carbon accelerates the reaction and hence the reactivity order: ROSO2(C6H4Me) > RI > RBr > RCl.

7.10.1.3 Ionic Mechanism This is similar to the SN2-­t ype mechanism mentioned above and involves the stepwise addition of two distinct ionic ligand fragments of the reagent. This mechanism operates with reagents that are dissociated in solution (prior to their interaction with a metal centre), as in the oxidative addition of HCl which remains ionized in solution as H+ and Cl− [eqn (7.48)].   



 [PtH(PPh3 )3 ]  Cl   [Pt(PPh3 )4 ]  H  Cl    [PtH(PPh3 )2 ] -PPh3 -PPh3 18e,d10 ,Pt 0 Tetrahedral

16e,d8 ,Pt II Square-planar

16e,d8 ,Pt II Square-planar

(7.48)

  

The following overall reaction also occurs in two steps as in eqn (7.48): PtIICl(H)(PR3)2 + H+ + Cl− → PtIVCl2(H)2(PR3)2 Hence the rate law for the first step of such reactions is generally Rate = k[complex][H+]

7.10.1.4 Radical Mechanism In addition to undergoing SN2-­t ype of oxidative addition, alkyl halides and similar reagents can add to a metal centre via a radical mechanism,105 some examples of which are given in Scheme 7.8. Processes in oxidative additions of a few classes of complexes with XY are shown in Schemes 7.9 and 7.10.

7.10.2  Five-­coordinate Eighteen-­electron Substrates Lower-­valent Group 8, 9 and 10 metals tend to form five-­coordinate complexes that conform to the 18e rule (coordinatively saturated). Hence their oxidative additions may involve SN2 displacement, which increases the coordination

Scheme 7.8

Electron Transfer Reactions

425

Scheme 7.9

Scheme 7.10 by one and the oxidation state by two units (Scheme 7.9, paths a and e). The reaction [eqn (7.49)] involves SN2 attack by the metal on XY = Br2, to afford [ML5Br]+ and Y− = Br−; the former may or may not undergo subsequent attack by Y− to displace a ligand (in either case, oxidative addition takes place). The intermediate [OsBr(CO)3(PPh3)2]Br can be isolated; on heating this to 60 °C, Br− displaces CO to form cis-­[OsBr2(CO)2(PPh3)2].   



[Os(CO)3(PPh3)2] + Br2 → cis-­[OsBr2(CO)2(PPh3)2] + CO

(7.49)

  

SN2 displacement by carbonylate anions is common and is expected to result in inversion of configuration at C when XY is an optically active halide such as Ph(Me)HC*Cl. Relative nucleophilicities of several carbonylate anions are given in Table 7.17.105 [FeCp(CO)2]− (and other carbonylate anions) may also react with H+ to produce hydrides such as [FeCp(H)(CO)2]; because we conventionally consider the H ligand as hydride (H−) and alkyls as R−, the oxidation number of the metal is increased by two in both cases. Alternatively, prior ligand loss (Scheme 7.9, path b) may occur to give a species that could go on to react via path c + f or path d (Scheme 7.9). A probable

Chapter 7

426

Table 7.17  Relative  nucleophilicities of some carbonylate anions.

106

Anion

Relative nucleophilicity −

[FeCp(CO)2] [RuCp(CO)2]− [NiCp(CO)]− [Re(CO)5]− [WCp(CO)3]− [M(CO)5]− [MoCp(CO)3]− [CrCp(CO)3]− [Co(CO)4]−

7 × 107 7.5 × 106 5.5 × 106 2.5 × 104 5 × 102 77 67 4 1

case of path b + d is the reaction shown in eqn (7.50), which occurs only under conditions that are sufficiently vigorous to promote ligand dissociation, in contrast to H2 addition to 16e square-­planar complexes, which takes place at room temperature and atmospheric pressure.   



  

[Os(CO)5  H2  [OsH2  CO)4   CO 100  C,80 atm

(7.50)

7.10.3  Four-­coordinate Sixteen-­electron Substrates Sixteen-­electron square-­planar complexes (which may be intermediates produced from ligand dissociation from ML5) can undergo both SN2 (path c) and concerted (path d) additions. Addition of Mel to Vaska's compound is an SN2 process characterized by a large negative value of ∆S‡ in non-­polar solvents; ∆S‡ becomes less negative in polar solvents, thereby increasing the rate. The rates of SN2 additions are promoted by “polar” substrates such as acyl halides, alkyl halides, hydrogen halides, halogens and pseudohalogens which contain good leaving groups. Concerted additions to square-­planar complexes involve less polar X–Y and include addition of H2 and R3SiH to [IrX(CO)X(Ph2PCH2CH2PPh2)]. Pearson107 pointed out that reactions proceeding with reasonably low activation energies involve electron flow between orbitals on each reactant with the same symmetry properties. For oxidative additions, electrons must flow from a filled metal orbital into an antibonding orbital of X–Y; this allows the X–Y bond to be broken and new bonds to the metal to be formed. For oxidative of addition of X–Y to square-­planar complexes, ML4, both cis and trans additions have been observed.

7.10.4  Four-­coordinate Eighteen-­electron Substrates With d10 four-­coordinate complexes (18e), ligand loss must also occur on oxidative addition as the product formula shows. This can be prior to (Scheme 7.10, path g, etc.) or concurrent with oxidative addition (Scheme 7.10, path h). For example, solutions of Pt(PPh3)4 contain Pt(PPh3)3 as the major species in

Electron Transfer Reactions

427

addition to some Pt(PPh3)2. It is likely that oxidative addition of RBr proceeds (via path g) according to eqn (7.51).   



[Pt(PPh3)3] + RBr → trans-­[PtR(Br)(PPh3)2] + PPh3

(7.51)

  

On the other hand, [Ni{P(OEt)3}4] is barely dissociated and may undergo oxidative addition (via path h, Scheme 7.10) with loss of two P(OEt)3 ligands, unless kinetically significant quantities of the unsaturated species too small to detect are present. The reactivity order for d10 complexes in general parallels ease of oxidation by two units: Pt0 ≈ Pd0 ≈ Ni0 > AuI > CuI ≫ AgI.

7.11  Reductive Elimination95d Many species can undergo reductive elimination, the reverse of oxidative addition, on heating and even by photolysis in some cases. Reductive elimination decreases both the coordination number and the oxidation state of the metal. Sometimes the molecule eliminated is different from the one added oxidatively. For example, the PdIV species resulting from oxidative addition given in eqn (7.52) is not isolable, but its presence as an intermediate explains the observed products. A requirement for reductive elimination seems to be that eliminated groups be cis oriented. This is evident from the product distributions for some gold complexes (Scheme 7.11).108   

(7.52)    

The reaction is retarded by added phosphine, which suggests that reductive elimination proceeds with ligand loss. This is reasonable because elimination increases the electron density on the metal. In contrast, some reductive

Scheme 7.11

Chapter 7

428

eliminations are promoted by added ligands. For example, cis-­[PtPh2(PPh3)2] eliminates biphenyl at a lower temperature in the presence of added phosphine. Presumably, a trigonal bipyramidal intermediate that is formed places the phenyl groups in a more favourable geometry for elimination. A different role for an added ligand seems plausible for the reductive elimination of alkane from [NiR2(bipy)], where dissociation of one end of the chelate is not favoured. Added olefins promote this reaction; the more electron-­withdrawing the olefin, the faster is the elimination of R2. This can be understood if the function of the olefin is to coordinate to Ni and drain off excess electron density, thereby stabilizing a transition state in which added electron density is placed on the metal. For trans-­[PdMe2(PPh3)2], elimination of ethane (Me–Me) involves an intermediate resulting from loss of a PPh3 ligand forming the trigonal (Y-­shaped) planar [PdMe2(PPh3)]. Even the cis complex, cis-­[PdMe2(PPh3)2], undergoes ethane elimination ∼100 times faster than in the case of the analogous chelating diphosphine complex [PdMe2(dppe)] (dppe = Ph2PCH2CH2PPh2). In the case of the trans-­spanning diphosphine complex trans-­[PdMe2(trans-­ diphos)], no ethane elimination occurs under conditions where the analogous cis-­[PdMe2(dppe)] undergoes the ethane elimination fairly easily. In its reaction with CD3I, [PdMe2(trans-­diphos)] eliminates D3C–Me forming [PdMe(I)(trans-­diphos)] through intermediate formation of ­[PdMe2(CD3) (trans-­diphos)]+ having CD3 bonded axially in this intermediate of square-­ pyramidal structure and hence cis to both Me groups, which are mutually trans. The mechanism is therefore associative. However, in cases of octahedral complexes, the analogous elimination reactions occur via the dissociative route (Scheme 7.12). Elimination reactions of dialkyl compounds of d8 metals have been well studied. There are significant variations in the thermal stability of [MR2L2] depending on M and L. Good donors such as bipy or PR3 stabilize the dialkyl compounds, and the PtII compounds are more stable than the PdII analogues.109 In the case of [PtIMe3(diphos)], both MeI and ethane are eliminated through parallel paths, as shown in Scheme 7.13.110 Reductive elimination involving two metal centres is also known, as in the following example:111–115   



7[Co(H)(CO)4 + [Co(Ac)(CO)4] → [Co2(CO)8] + MeCH= O

  

Scheme 7.12

(7.53)

Electron Transfer Reactions

429

Scheme 7.13

References 1. (a) H. Taube, H. Myers and R. L. Rich, J. Am. Chem. Soc., 1953, 75, 4118; (b) H. Taube and H. Myers, J. Am. Chem. Soc., 1954, 76, 2103; (c) H. Taube, J. Am. Chem. Soc., 1955, 77, 4481; (d) A. E. Ogard and H. Taube, J. Am. Chem. Soc., 1958, 80, 1084. 2. (a) H. Taube, Electron Transfer Reactions of Complex Ions in Solution, Academic Press, New York, 1970; (b) R. D. Cannon, Electron Transfer Reactions, Butterworths, London, 1980. 3. R. G. Linck, Rates and Mechanisms of Oxidation–Reduction Reactions, in MTP International Review of Science, ed. M. L. Tobe, Butterworths, London, Inorganic Chemistry Series One, 1972, ch. 8, vol. 9. 4. (a) A. Haim, Acc. Chem. Res., 1975, 8, 264; (b) N. Sutin, Acc. Chem. Res., 1982, 15, 275; (c) C. Creutz, Prog. Inorg. Chem., 1983, 30, 1; (d) J. F. Endicott, K. Kumar, T. Ramasami and P. Rotzinger, Prog. Inorg. Chem., 1983, 30, 141; (e) A. Haim, Prog. Inorg. Chem., 1983, 30, 273; (f) T. J. Meyer, Prog. Inorg. Chem., 1983, 30, 389; (g) N. Sutin, Prog. Inorg. Chem., 1983, 30, 441; (h) R. A. Marcus, Angew. Chem., Int. Ed. Engl., 1993, 32, 1111. 5. S. P. Ghosh, M. C. Ghosh and E. S. Gould, Inorg. Chim. Acta, 1994, 225, 83. 6. J. F. Endicott, B. Durham and K. Kumar, Inorg. Chem., 1982, 21, 2437. 7. (a) J. Weiss, Proc. R. Soc. London, Ser. A, 1954, 222, 128; (b) R. J. Marcus, B. J. Zwolinski and H. Eyring, J. Phys. Chem., 1954, 58, 432; (c) R. A. Marcus, J. Chem. Phys., 1956, 24, 966, 970; 1957, 26, 867; Annu. Rev. Phys. Chem., 1964, 15, 155.; (d) T. W. Newton, J. Chem. Educ., 1968, 45, 571; (e) N. Sutin and C. Creutz, J. Chem. Educ., 1983, 60, 809. 8. (a) M. H. Ford-­Smith and N. Sutin, J. Am. Chem. Soc., 1961, 83, 1830; (b) R. J. Companion, N. Purdie and N. Sutin, Inorg. Chem., 1964, 3, 1091; (c) J. D. Miller and R. H. Prince, J. Chem. Soc. A, 1966, 1370; (d) N. Sutin, Annu. Rev. Phys. Chem., 1966, 17, 119; (e) A. Haim and N. Sutin, J. Am. Chem. Soc., 1966, 88, 434; (f) R. A. Marcus and N. Sutin, Biochim. Biophys. Acta, 1985, 811, 265; (g) R. J. Companion, N. Purdie and N. Sutin,

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Chapter 7

Inorg. Chem., 1964, 3, 1091; (h) M. Chou, C. Creutz and N. Sutin, J. Am. Chem. Soc., 1977, 99, 5613; 5623. 9. P. Blowers and R. I. Masel, J. Phys. Chem. A, 1999, 103, 7047. 10. B. S. Brunschwig, J. Logan, M. D. Newton and N. Sutin, J. Am. Chem. Soc., 1980, 102, 5798. 11. T. K. Sham and B. S. Brunschwig, J. Am. Chem. Soc., 1981, 103, 1590. 12. (a) W. J. Libby, J, Phys. Chem., 1952, 56, 863; (b) R. A. Marcus, Discuss. Faraday Soc., 1960, 29, 129; Annu. Rev. Phys. Chem., 1964, 15, 155; cf., N. S. Hush, Z. Electrochem., 1957, 61, 734; J. Chem. Phys., 1958, 28, 962; Trans. Faraday Soc., 1961, 57, 557. 13. (a) J. C. Sheppard and A. C. Wahl, J. Am. Chem. Soc., 1957, 79, 1020; (b) L. Gjertsen and A. C. Wahl, J. Am. Chem. Soc., 1959, 81, 1572; (c) A. C. Wahl, Z. Elektrochem., 1960, 60, 90; (d) M. Shopover, G. Ron, A. Loewenstein and G. Navon, Inorg. Chem., 1965, 4, 362. 14. (a) H. Taube and E. S. Gould, Acc. Chem. Res., 1964, 2, 321; (b) W. G. Movius and T. G. Linck, J. Am. Chem. Soc., 1969, 91, 5934. 15. T. W. Newton, G. E. McCrary and W. G. Clark, J. Phys. Chem., 1968, 72, 4333. 16. H. Taube, Rec. Chem. Prog., 1956, 17, 25. 17. (a) M. C. Moore and R. N. Keller, Inorg. Chem., 1971, 19, 747; (b) N. Sutin, Chem. Br., 1972, 148; (c) J. H. Espenson, Inorg. Chem., 1969, 8, 1554. 18. H. Taube, J. Am. Chem. Soc., 1960, 82, 526. 19. A. Zwickel and H. Taube, J. Am. Chem. Soc., 1961, 83, 793. 20. T. L. Toppen and R. G. Linck, Inorg. Chem., 1971, 10, 2635. 21. P. Candlin, J. Halpern and D. L. Trimm, J. Am. Chem. Soc., 1964, 86, 1019. 22. C. Shea and A. Haim, J. Am. Chem. Soc., 1971, 93, 3055. 23. J. H. Espenson, Inorg. Chem., 1965, 4, 121. 24. D. P. Fay and N. Sutin, Inorg. Chem., 1970, 9, 1291. 25. (a) D. E. Pennington and A. Haim, Inorg. Chem., 1967, 6, 2138; (b) R. D. Cannon and J. E. Earley, J. Chem. Soc. A, 1968, 1102. 26. M. Eigen, Rev. Pure Appl. Chem., 1963, 6, 105. 27. D. W. Carlyle and J. H. Espenson, J. Am. Chem. Soc., 1969, 91, 599. 28. N. Sutin, Acc. Chem. Res., 1968, 1, 225. 29. A. Haim and N. Sutin, J. Am. Chem. Soc., 1966, 88, 5343. 30. R. D. Canon and J. Gardiner, J. Am. Chem. Soc., 1970, 92, 3800; R. G. Linck, Rates and Mechanisms of Oxidation–Reduction Reactions, in MTP International Review of Science, ed. M. L. Tobe, Butterworths, London, Inorganic Chemistry Series One, 1972, ch. 8, vol. 9, p. 313. 31. (a) A. W. Adamson and E. Conick, Inorg. Chem., 1963, 2, 129; (b) D. H. Huchital and R. G. Wilkins, Inorg. Chem., 1967, 6, 1022. 32. B. Grossman and A. Haim, J. Am. Chem. Soc., 1970, 92, 4835. 33. M. Green, K. Schug and H. Taube, Inorg. Chem., 1965, 4, 1184. 34. M. Orhanovic and N. Sutin, J. Am. Chem. Soc., 1968, 90, 4286. 35. H. Taube, Adv. Chem. Ser., 1965, 14, 107. 36. (a) C. Shea and A. Haim, J. Am. Chem. Soc., 1971, 93, 3055; (b) R. J. Restivo, G. Ferguson and R. J. Balahura, Inorg. Chem., 1977, 16, 167; (c) C. S. Glennon, J. D. Edwards and A. G. Sykes, Inorg. Chem., 1978, 17, 1654.

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37. P. Ray, J. Indian Chem. Soc., 1927, 4, 327. 38. (a) J. P. Candlin, J. Halpern and S. Nakamura, J. Am. Chem. Soc., 1963, 85, 2517; (b) J. Halpern and S. Nakamura, Proc. 8th Int. Confc. Coord., Chem., Vienna, 1964, p. 271; (c) C. K. Jorgensen, Inorg. Chem., 1964, 3, 1201; (d) J. L. Burmeister, Coord. Chem. Rev., 1966, 1, 205. 39. D. W. Hoppenjans, J. B. Hunt and L. Penzhorn, Inorg. Chem., 1968, 7, 1467. 40. H. Taube and E. L. King, J. Am. Chem. Soc., 1964, 76, 4053. 41. R. Smellgrove and E. L. King, J. Am. Chem. Soc., 1962, 84, 4609. 42. A. Haim, J. Am. Chem. Soc., 1963, 85, 1016. 43. A. Haim, J. Am. Chem. Soc., 1966, 88, 2324. 44. J. R. Ward and A. Haim, J. Am. Chem. Soc., 1970, 92, 475. 45. D. H. Huchital, Inorg. Chem., 1970, 9, 486. 46. W. S. Melvin and A. Haim, Inorg. Chem., 1977, 16, 2016. 47. A. G. Sykes and R. N. V. Thorneley, J. Chem. Soc. A, 1970, 232. 48. T. W. Newton and F. S. Baker, Inorg. Chem., 1964, 3, 569. 49. (a) H. Ogino, E. Kikkawa, M. Shimura and N. Tanaka, J. Chem. Soc., Dalton Trans., 1981, 894; (b) R. Marcec and M. Oranpvic, Inorg. Chem., 1978, 17, 3672; (c) S. Zakir-­Ali, P. Chalilkayil and J. E. Earley, Inorg. Chim. Acta, 1981, 48, 57; (d) R. A. Kelley and J. E. Earley, Inorg. Chim. Acta, 1983, 76, L167. 50. J. H. Sutter and J. B. Hunt, J. Am. Chem. Soc., 1969, 91, 3107. 51. J. Doyle and A. G. Sykes, J. Chem. Soc. A, 1967, 795. 52. P. H. Dodel and H. Taube, Z. Phys. Chem., 1955, 44, 92. 53. R. G. Linck, Inorg. Chem., 1968, 7, 2394. 54. D. Olmstead, X. Hua, P. Osvath and A. G. Lappin, Dalton Trans., 2010, 39, 1375. 55. C. Creutz and H. Taube, J. Am. Chem. Soc., 1969, 91, 3988; 1973, 95, 1086. 56. T. J. Meyer, Acc. Chem. Res., 1978, 11, 94. 57. (a) H. Taube and C. S. Gould, Acc. Chem. Res., 1969, 2, 321; (b) F. Nordmeyer and H. Taube, J. Am. Chem. Soc., 1868, 90, 1162. 58. J. Halpern and S. Nakamura, J. Am. Chem. Soc., 1965, 87, 3002. 59. (a) R. G. Gaunder and H. Taube, Inorg. Chem., 1970, 9, 2627; (b) J. A. Stritar and H. Taube, Inorg. Chem., 1969, 8, 2281. 60. (a) J. E. French and H. Taube, J. Am. Chem. Soc., 1969, 91, 695; (b) E. S. Gould, J. Am. Chem. Soc., 1972, 94, 4360; (c) M. Z. Hoffman and M. Simic, J. Am. Chem. Soc., 1972, 94, 1757. 61. J. P. Candlin and J. Halpern, Inorg. Chem., 1965, 4, 766. 62. (a) J. F. Endicott and H. Taube, J. Am. Chem. Soc., 1964, 86, 1686; (b) J. F. Endicott and H. Taube, Inorg. Chem., 1965, 4, 437. 63. (a) K. M. Davies and J. H. Espenson, J. Am. Chem. Soc., 1969, 91, 3093; (b) J. H. Espenson, Acc. Chem. Res., 1970, 3, 347. 64. (a) A. G. Sykes and M. Green, J. Chem. Soc. A, 1971, 3221; (b) M. Green, R. S. Taylor and A. G. Sykes, J. Chem. Soc. A, 1971, 509. 65. A. A. Vlcek, Discuss. Faraday Soc., 1958, 26, 164.

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66. (a) J. H. Baxendale, E. M. Fielden and J. P. Kune, Proc. R. Soc. London, Ser. A, 1965, 286, 320; (b) G. D. Venerable, Inorg. Chem., 1971, 10, 1999; (c) M. Anbar and P. Neta, Int. J. Appl. Radiat. Isot., 1967, 18, 493. 67. (a) J. C. Chen and E. S. Gould, J. Am. Chem. Soc., 1973, 95, 5539; (b) R. A. Marcus, J. Phys. Chem., 1963, 67, 353. 68. (a) E. R. F. Fan and E. S. Gould, Inorg. Chem., 1974, 13, 2647; (b) M. K. Loar, M. A. Sens, G. W. Loar and E. S. Gould, Inorg. Chem., 1978, 17, 330; (c) E. S. Gould, Inorg. Chem., 1979, 18, 900 and references cited therein. 69. (a) C. S. Glennon, T. D. Hand and A. G. Sykes, J. Chem. Soc., Dalton Trans., 1980, 19; (b) H. M. Huck and K. Weigharch, Angew. Chem., Int. Ed., 1980, 19, 558. 70. (a) A. K. Basak and D. Banerjea, Z. Anorg. Allg. Chem., 1976, 425, 277; (b) A. K. Basak and D. Banerjea, J. Indian Chem. Soc., 1977, 54, 1015; (c) A. K. Basak and D. Banerjea, Indian J. Chem., Sect. A., 1977, 17, 250; (d) A. K. Banerjee, A. K. Basak and D. Banerjea, Indian J. Chem., 1979, 18A, 332. 71. R. N. Banerjee and D. Banerjea, Indian J. Chem., 1979, 17A, 246. 72. M. Chakraborty, N. J. Singh, P. C. Mandal, S. Das and S. Mukhopadhyay, J. Phys. Chem. A, 2011, 115, 4882. 73. B. T. Farrer and H. H. Thorp, Inorg. Chem., 1999, 38, 2497. 74. R. J. Campion, T. J. Conocchioli and N. Sutin, J. Am. Chem. Soc., 1964, 86, 4591. 75. N. Sutin, Annu. Rev. Nucl. Sci., 1962, 12, 285. 76. J. K. Beattic, Chromium(vi) Oxidation of Inorganic Substrates, in Inorganic Reaction Mechanisms, Part II, ed. J. O. Edwards, of the series Progress in Inorganic Chemistry, ed. S. J. Lippard, Interscience, New York,1972, vol. 17, p. 93. 77. R. C. Patel and J. F. Endicott, J. Am. Chem. Soc., 1968, 90, 6364. 78. (a) J. Koutecky, Collect. Czech. Chem. Commun., 1953, 18, 597; 1955, 20, 980; (b) K. B. Oldham and E. P. Parry, Anal. Chem., 1968, 40, 65. 79. D. Banerjea and P. Banerjee, J. Indian Chem. Soc., 1972, 49, 1321. 80. (a) S. Sen Gupta and D. Banerjea, J. Indian Chem. Soc., 1976, 53, 741; (b) A. K. Banerjee and D. Banerjea, Indian J. Chem., 1983, 22A, 937. 81. M. Upadhyay, S. Gangopadhyay and D. Banerjea, J. Indian Chem. Soc., 1985, 62, 1029. 82. (a) Solvated Electron, in Adv. Chem. Ser., ed. E. J. Hart, ACS, Washington D.C., 1965, vol. 50; (b) F. Basolo and R. G. Pearson, Mechanisms of Inorganic Reactions: A Study of Metal Complexes in Solution, John Wiley, New York, 2nd edn, 1967, pp. 515–519; (c) D. C. Walker, Q. Rev., Chem. Soc., 1967, 21, 79; (d) E. Hart, Surv. Prog. Chem., 1969, 5, 129; (e) J. M. Herbert and M. P. Coons, Annu. Rev. Phys. Chem., 2017, 68, 447. 83. (a) W. Weyl, Ann. Phys. Chem., 1864, 197, 601; (b) C. A. Kraus, J. Am. Chem. Soc., 1907, 29, 557; (c) R. Hoffmann, Angew. Chem., Int. Ed., 2009, 48, 8198. 84. G. E. Gibson and W. L. Argo, J. Am. Chem. Soc., 1918, 40, 1327. 85. (a) E. J. Hart and J. W. Boag, J. Am. Chem. Soc., 1962, 84, 4090; (b) E. J. Hart, Science, 1964, 146, 19.

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86. D. Myerstein, Inorg. Chem., 1971, 10, 639. 87. J. Jortnen and J. Rabani, J. Am. Chem. Soc., 1961, 83, 3868. 88. M. S. Matheson and J. Rabani, J. Phys. Chem., 1965, 69, 1324. 89. S. Gordon, E. J. Hash, M. S. Matheson, J. Rabani and J. K. Thomas, Discuss. Faraday Soc., 1968, 36, 193. 90. E. Janata and R. Schuter, J. Phys. Chem., 1982, 86, 2078. 91. S. E. Forest, M. J. Stimson and J. D. Simon, J. Phys. Chem. B, 1999, 103, 3963. 92. (a) J. H. Baxendale, E. M. Fielden, C. Capellos, J. M. Francis, J. V. Davies, M. Ebert, C. W. Gilbert, J. P. Keene, E. J. Land, A. J. Swallow and J. M. Nosworthy, Nature, 1964, 201, 468; (b) M. Anbar and E. J. Hart, J. Phys. Chem., 1965, 69, 973. 93. (a) J. K. Thomas, S. Gordon and E. J. Hart, J. Phys. Chem., 1964, 68, 1524; (b) M. Anbar and E. J. Hart, J. Phys. Chem., 1965, 69, 973. 94. (a) J. H. Blaxendale, Radiat. Res. Suppl., A, 1964, 4, 139; (b) E. J. Hart, S. Gordon and E. M. Fielden, J. Phys. Chem., 1966, 70, 150. 95. (a) J. P. Collman and W. R. Roper, Adv. Organomet. Chem., 1968, 7, 53; (b) A. J. Deeming, Oxidative Addition, in Reaction Mechanisms in Inorganic Chemistry, ed. M. L. Tobe, Butterworths, London, Inorganic Chemistry Series One, 1972, ch. 4, vol. 9; (c) J. K. Stille, in The Chemistry of the Metal-­Carbon Bond, ed. F. R. Hartley and S. Patai, Wiley, Chichester, 1985, vol. 2, p. 625; (d) R. H. Crabtree, The Organometallic Chemistry of the Transition Metals, Wiley, Hoboken, NJ (USA), 5th edn, 2009. 96. (a) G. J. Kubas, R. R. Ryan, B. I. Swanson, P. I. Vergamini and H. J. Wasserman, J. Am. Chem. Soc., 1984, 106, 452; (b) G. J. Kubas, Acc. Chem. Res., 1988, 21, 120; (c) R. H. Crabtree, Acc. Chem. Res., 1990, 23, 95. 97. (a) D. Milstein, J. C. Calabrese and I. D. Williams, J. Am. Chem. Soc., 1986, 108, 6387; (b) J. P. Collman and W. R. Roper, Adv. Organomet. Chem., 1968, 7, 54. 98. A. Vaska, Acc. Chem. Res., 1976, 9, 175. 99. R. S. J. LaPlaca and J. A. Ibers, J. Am. Chem. Soc., 1965, 87, 2581. 100. D. C. Hrncin, R. D. Rogers and J. L. Atwood, J. Am. Chem. Soc., 1981, 103, 4277. 101. P. B. Chock and J. Halpern, J. Am. Chem. Soc., 1966, 88, 3511; see also L. Vaska, Science, 1971, 174, 587. 102. P. K. Hurlburt, O. P. Anderson and S. H. Strauss, J. Am. Chem. Soc., 1991, 113, 6277. 103. (a) J. Curtis and R. Eisenberg, J. Am. Chem. Soc., 1985, 107, 3148; (b) S. J. P. Pool, M. A. Klingshirn, R. D. Rogers and K. H. Shaughnessy, J. Organomet. Chem., 2003, 690, 3522. 104. (a) T. R. Cundari, J. Am. Chem. Soc., 1994, 116, 340; (b) M. Brookhart, M. L. H. Green and G. Perkin, Proc. Natl. Acad. Sci. U. S. A., 2007, 104, 6908; (c) W. Scherer, A. C. Dunbar, J. E. Berquera-­Lozada, D. Schmitz, G. Erekerling, D. Kratzert, D. Stalke, A. Lanza, P. Macchi, N. R. M. Casati, J. Ebad-­Allah and C. Kuntscher, Angew. Chem., Int. Ed., 2015, 54, 2505. 105. C. D. Hoff, Coord. Chem. Rev., 2000, 206, 451. 106. R. B. King, Acc. Chem. Res., 1970, 3, 417.

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107. R. G. Pearson, Symmetry Rules for Chemical Reactions, Wiley, New York, 1976. 108. S. Komiya, T. Albright, R. Hoffman and J. K. Kochi, J. Am. Chem. Soc., 1976, 98, 7255. 109. J. M. Brown and N. A. Cooley, Chem. Rev., 1988, 88, 1031. 110. K. I. Goldberg, J. Y. Yan and E. L. Winter, J. Am. Chem. Soc., 1994, 116, 1573. 111. B. Douglas, D. McDaniel and J. Alexander, Concepts and Models of Inorganic Chemistry, John Wiley, New York, 3rd edn, 1993, p. 676. 112. (a) S. K. Saha, M. C. Ghosh and P. Banerjee, J. Chem. Soc., Dalton Trans., 1986, 1301; Inorg. Chim. Acta, 1987, 126, 29; J. Chem. Kinetics, 1988, 20, 699; (b) M. Ali, S. K. Saha and P. Banerjee, J. Chem. Soc., Dalton Trans., 1990, 187; (c) M. Gupta, S. K. Saha and P. Banerjee, Bull. Chem. Soc. Jpn., 1990, 63, 609; (d) S. Gangopadhyay, M. Ali and P. Banerjee, J. Chem. Soc., Perkin Trans. 2, 1992, 781. 113. L. M. Bharadwaj, D. N. Sharma and Y. K. Gupta, J. Chem. Soc., Dalton Trans., 1980, 1526. 114. (a) J. E. Earley, Prog. Inorg. Chem., 1970, 13, 243; (b) P. Benson and A. Haim, J. Am. Chem. Soc., 1965, 87, 3826; (c) C. K. Jorgensen, Modern Aspects of Ligand Field Theory, Elsevier, New York, 1971, ch. 26; (d) C. Bifano and R. G. Linck, J. Am. Chem. Soc., 1967, 89, 3945; (e) R. G. Linck, Inorg. Chem., 1970, 9, 2529; (f) J. O. Edwards, Inorganic Reaction Mechanisms, Benjamin, New York, 1964; Table 4.1; (g) Stability Constants:, Special Publication No. 17 (1964), No. 25 (1971), compiled by L. G. Sillen and A. E. Martell, The Chemical Society. London; Tables 27 and 275. 115. R. D. Cannon and J. E. Earley, J. Chem. Soc. A, 1968, 1102.

Chapter 8

Activation of Molecules by Coordination and Reactivity of Coordinated Ligands 8.1  Introduction The electron distribution in ligands is changed when they bind to metals, and this changes their reactivity. Hence coordinated ligands are often activated to attack by nucleophiles or electrophiles, which does not occur in the free molecules. Often a pronounced effect is observed on binding of a metal to an organic substrate that greatly influences the chemical reactivity of the latter in a variety of ways. For example, it is remarkable that K[PtCl3(CH2=CH2)] is resistant to the oxidizing action of an acidified solution of KMnO4 in the cold, although both Pt(ii) compounds and CH2=CH2 are readily oxidized under these conditions. The chemical reactivity of coordinated ligands is a field with potential and tremendous technological implications. Commercial applications of this concept are to be found in the polymerization (Ziegler‒Natta), oxidation (Wacker) and hydroformylation reactions of olefins. Metal complexes are also well known activators of O2, N2, H2, etc.

8.2  Activation of Some Diatomic Molecules 8.2.1  Activation of Dihydrogen by Coordination Hydrogen, H2 (dihydrogen), is not a highly reactive molecule for reduction and/or hydrogenation, presumably owing to its high bond energy (434 kJ mol−1). Thus, it fails to reduce an acidified solution of MnO4− unless traces of   Mechanisms of Reactions of Metal Complexes in Solution By Debabrata Banerjea and M. K. Bharty © Debabrata Banerjea and M. K. Bharty 2023 Published by the Royal Society of Chemistry, www.rsc.org

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436 2+

a metal salt such as of Cu are present to catalyse the reaction. Similarly, for hydrogenation a catalyst is required. A number of examples of the activation of molecular hydrogen by metal complexes in solution are known; most active are some transition metal compounds. One of the earliest studies of such a reaction was reported by Calvin,1 who showed that in quinoline solution Cu(i) salts of organic acids catalyse the reduction of Cu(ii) salts and of benzoquinone by hydrogen. Detailed studies2 further showed that the reaction is homogeneous, independent of the concentration of the oxidant, and follows the rate law given by eqn (8.1). A plausible explanation is the formation of a Cu(ii) hydride as shown in eqn (8.2). For the Cu(OAc)-­catalysed process, Ea for H2 splitting is 13 kcal mol−1 (13.7 kcal mol−1 for D2, which therefore reacts more slowly).2d   

Rate = k[Cu(i)salt]2[H2]

(8.1)

  

Cu 2Cu Cu  H2#Cu   H2  # Cu   H2 Cu # 2 CuH   4 Cu   2 H  (8.2) Cu   

In the absence of an oxidizing agent, the system catalyses the ortho‒para-­ hydrogen conversion,3a and also the exchange of deuterium with a hydrogen donor in solution,3b,c which demonstrates the reversibility of the H2 splitting reaction step.4 Similarly, Cu2+, Ag+ and Hg2+ in aqueous solution are known to catalyse the reduction of a number of oxidizing agents [such as Cr(vi)] by hydrogen.3a,5 In the case of Cu2+ and Ag+, hydride species such as CuH+ and AgH are believed to be formed: Cu2+ +H2 → CuH+ + H+; Ag+ + H2 → AgH + H+ whereas in the case of Hg2+, the rate-­determining step is believed to be Hg2+ + H2 → Hg + 2H+ The Hg formed then reacts rapidly with the oxidizing agent to form the product and regenerating the Hg2+ catalyst.6 Compounds of Pd, Pt, Ru, Rh, Ir, etc., are very effective hydrogen activating catalysts. From a detailed study of H2 activation by Cuaq2+ in aqueous HClO4 solution, the following mechanism was proposed, which explains all the observed facts including a change of order from 1 to 2 with respect to Cu2+ as the concentration of H+ increases (Scheme 8.1). Applying the steady-­state principle to the reactive, the rate law shown in eqn (8.3) was derived.

Scheme 8.1

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437

The rate decreases with increase in acid concentration, as expected.   

Rate = k1k3[H2][Cu2+]2/(k3[Cu2+] + k2[H+])

(8.3)

  

Based on the derived rate law in eqn (8.3) we obtain    (a) at low [H+], Rate = k1[Cu2+][H2]; (b) at high [H+], Rate = (k1k3/k2)[Cu2+]2[H2]/[H+].    The rate decreases with increase in acid concentration as expected. For reduction of Cr(vi) by H2 catalysed by Ag+, there is both a heterolytic and a homolytic activation path [eqn (8.4a and b)].5a   

Ag+ + H2 ⇄ AgH + H+

(8.4a)

  



2Ag+ + H2 → 2AgH+

(8.4b)

  

The catalytic action of all of these compounds is due to their ability to activate molecular hydrogen, either by homolytic splitting [H2 → 2H; ∆H(g) = 434 kJ mol−1) or by heterolytic splitting (H2 → H+ + H−; ∆H(g) = 1676 kJ mol−1]. The heterolytic splitting (unfavourable in the gaseous state) is energetically favoured in solvents of high polarity such as water, more particularly when H+ or H− or both are fixed. Thus, even a strong base such as OH− or NH2− (in liquid ammonia) can facilitate heterolytic cleavage of the H2 molecule7 by removing the H+, while AgF, which provides both the Lewis acid (Ag+) and the Lewis base (F−) needed to fix H− and H+, respectively, is very effective in activating hydrogen.8a In the absence of any H+ binder, solvation of H+, which is highly exothermic (ΔHaq = −1150 kJ mol−1),8b is a contributing factor. Investigations9 of some H2 activating systems in D2O have shown that equivalent amounts of hydride and deuteride are produced, indicating heterolytic cleavage of H2 (Scheme 8.2). A case of homolytic fission, which occurs under suitable conditions when both the H atoms can be picked up, as with Cu(i), was mentioned above, and another well-­characterized example is the reaction of H2 with [CoII(CN)5]3− in aqueous solution [eqn (8.5)].10   



2[CoII(CN)5]3- + H2 ⇌ 2[HCoIII(CN)5]3-

  

Scheme 8.2

(8.5)

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The equilibrium constant is 1.6 × 10 M at 25 °C, ∆H = −11.2 kcal mol−1, ∆S = −15 cal K−1 mol−1 and the Co–H bond energy is 57 kcal mol−1. The observed rate law for this process is11 11

5

−1

Rate = k{[Co(CN)5]3−}2[H2] The anion [Co(CN)5H]3− has been isolated as sodium and caesium salts. In all such cases the Mx+H formed is actually M(x+1)+H− due to an internal electron transfer. [Co(CN)5]3− is a homogeneous catalyst for the hydrogenation of a variety of substrates, including conjugated diolefins, such as butadiene and styrene. [IrCl(CO)(PPh3)2] (Vaska's compound) in benzene solution also brings about homolytic splitting of H2 [eqn (8.6)].12 [IrCl(CO)(PPh3)2] and its rhodium analogue effect the catalytic hydrogenation of ethylene, propylene and acetylene.   



[IrICl(CO)(PPh3)2] + H2 ⇌ [IrIIICl(H)2(CO)(PPh3)2]

(8.6)

  

Many transition metal complexes having suitable ligands bound to the metal activate H2 by similar homolytic splitting. This can be understood by referring to Figure 8.1, in which the M–H2 bonding scheme is shown. The σ-­bond between the metal and H2 is a µ-­bond,13a similar to that of CH2=CH2 in metal‒alkene complexes such as the Zeise's salt, and results from overlap of the bonding σ MO of H2, which is filled, with an empty σ hybrid bonding orbital of the metal atom. In addition, there is back-­bonding (π type) involving a filled metal orbital (dπ) with the vacant σ antibonding MO of H2 (Dewar‒Chatt‒Duncanson model).13b,c As a result, if the metal is strongly π donating, it may put enough electron density into the antibonding MO of the H2 such that it is equal to the electron density in the σ-­bonding MO of the H2. In such a situation, the H–H bond will cease to exist, i.e. resulting in H–H bond fission (homolytic) with each of the H atom binding to the metal atom by an electron pair σ-­bond, one electron being contributed by each of the bonded atoms, viz. the metal (M) atom and the H atom. This is analogous to the combination of Cl with M forming M–Cl, which effectively raises the oxidation state of the M by one unit (+1) for each M–H bond formed while the oxidation state of H is reduced to −1. Hence the H in the M–H bond may be viewed as a hydride (H−). For the process H + e−→ H−, ∆H ≈ −72.4 kJ mol−1.

Figure 8.1  Bonding  (σ and π) in M–H2.

Activation of Molecules by Coordination and Reactivity of Coordinated Ligands

439

Scheme 8.3 In HCl solution, Ru(iii) chloride catalyses the oxidation of H2 by Ru(iv) and Fe(iii), for which a heterolytic splitting mechanism has been proposed (Scheme 8.3).14 In dilute HCl solution, Ru(ii) catalytically causes hydrogenation of olefinic compounds such as maleic and fumaric acids.15 Ruthenium(ii)‒phosphine‒halide complexes are good hydrogenation catalysts.16 1-­Heptene and 1-­hexyne are hydrogenated rapidly at 25 °C and 1 atm, using [RuCl2(PPh3)4] as a catalyst in ethanol‒benzene (1 : 1) solution; the intermediate [RuH(Cl) (PPh3)3] has been isolated from concentrated solution and is the most active catalyst. The catalyst is very specific for reducing 1-­alkenes rapidly but internal alkenes very slowly. Another good catalyst for homogeneous hydrogenation is Wilkinson's catalyst, [RhCl(PPh3)3],17a which hydrogenates 1-­hexyne and 1-­hexene at room temperature and atmospheric pressure; an intermediate [Rh(H)2Cl (PPh3)2(S)] is believed to be formed, where S is a molecule of the solvent. The formation of metal hydride species during such hydrogenation reactions has been detected by NMR spectroscopy from the appearance of NMR peaks characteristic of M–H bonds in a solution of the catalyst treated with hydrogen. On addition of alkene, the characteristic NMR peaks disappear owing to transfer of hydrogen from the metal hydride species to the alkene. This catalyst is selective for 1-­alkenes (i.e. terminal olefins) rather than 2-­alkenes. It is also a useful catalyst for the oxo process (a hydroformylation reaction used to convert an alkene to an aldehyde). [RhH(PPh3)3] is more active than [RhCl(PPh3)3] for hydrogenation. A plausible mechanism of the hydrogenation of an alkene (H2C=CHR) using Wilkinson's catalyst is shown in Scheme 8.4.4b,17b–d Hydrogenation of H2C=CH2, however, takes place very slowly owing to the formation of reasonably stable[RhCl(η2-­C2H4)(PPh3)2]. Migration of H and its bonding with an alkene carbon atom constitute a process that is notionally insertion of the alkene into an Rh–H bond. In all the bisphosphine species in the reactions in Scheme 8.4 the phosphines may be in cis positions, for which evidence has been furnished.17e Several other good hydrogenation catalysts are known, such as [Rh(H)(CO)(PPh3)3], [Rh(diene)(PPh3)2]+, [Ru(H)Cl(PPh3)3] and [IrCl(CO)(PPh3)2] (Vaska's complex, which is also an efficient catalyst for oxidation by O2, see Section 8.2.2). Homogeneous hydrogenation by Wilkinson's catalyst and similar complexes and asymmetric homogeneous catalytic hydrogenation have been reviewed.17f Solutions of RhCl3 and [RhCl3(py)3] in ethanol have also been reported to hydrogenate simple alkenes. A mixture of H2PtCl6 and SnCl2 (1 : 5 molar ratio)

Chapter 8

440

Scheme 8.4 in the presence of HCl is an efficient hydrogenation catalyst, which hydrogenates acetylene and ethylene at 20 °C and 1 atm pressure in methanol solution. It is thought that an anionic Pt(ii) complex having SnCl3− and H− ligands is the active species. In fact, complexes of the type [Pt(SnCl3)5]3−, [HPt(SnCl3)4]3−, [HPt(SnCl3)2(PEt3)2]−, etc., have been isolated, the first compound being converted into the second compound by hydrogen gas under slight pressure.18 A wide variety of soluble Ziegler-­like systems catalyse the hydrogenation of alkenes at ambient temperature and moderate pressure of hydrogen.19 The type of system used consists of a metal acetylacetonate such as [Cr(acac)3], [Fe(acac)3] or [Co(acac)3], or a metal alkoxide such as Ti(OiPr)4 or V(OEt)3, or cyclopentadienyl complexes such as Cp2TiCl2, Cp2ZrCl2, CpFeCl(CO), CpCo(CO)2, etc., together with an excess of an aluminium alkyl such as AltBu3 in heptane or toluene. Most alkenes are hydrogenated by such catalyst systems.

Activation of Molecules by Coordination and Reactivity of Coordinated Ligands

441

Scheme 8.5 A variety of microorganisms have the ability to activate hydrogen for exchange with water, ortho‒para conversion in hydrogen and reduction reactions. This ability has been ascribed to an enzyme (or system of enzymes) named hydrogenase. Stephenson and Stickland20a first recognized hydrogen activation by microorganisms. Studies by various workers suggest that these enzymes contain an active M(ii) site (M = Fe, Ni) and sometimes molybdenum may also be an active component. On the basis of the HD production in H2–D2O exchange under the influence of such enzymes, the familiar heterolytic splitting of H2 was postulated by Rittenberg and Krasna20b (Scheme 8.5). In the presence of a substrate, S, hydride transfer and subsequent protonation yield the reduction product [eqn (8.7)].   



  

k3 Fast EH  S   E  SH ;SH  H  SH2

(8.7)

Tamuja and Miller21 studied the D2/HD product ratio as a function of pH and enzyme concentration and modified the mechanism, in that EH− may also exchange with the solvent. Rittenberg22 suggested that a basic site in the enzyme picks up the proton, thus favouring the reaction shown in eqn (8.8). Such a model readily explains the observed maximum activity at an optimum pH, since at low pH the basic site may be neutralized, whereas at high pH the Fe(ii) may have bound OH−. Gest and co-­workers,23 however, proposed homolytic splitting of hydrogen and fixation of the H atom at two Fe(ii) centres followed by internal electron transfer leading to FeIII–H having bound H−.   





(8.8)

  

8.2.2  Activation of Dioxygen by Coordination The standard reduction potential for the overall reaction [eqn (8.9)] is +1.23 V (pH ≈ 0), +0.80 V (pH = 7.5). Hence the overall process is thermodynamically favoured. However, as this reaction takes place in several steps, each involving the addition of one electron, and since the reduction potential for the very first step [eqn (8.10)] is −0.2 to −0.5 V (depending on the medium and pH), this step is thermodynamically unfavourable. This is why O2 as such fails to act as a reactive oxidizing agent in the absence of a suitable catalyst.   



O2 + 4H+ + 4e− → 2H2O

(8.9)

O2 + e− → O− 2

(8.10)

  

  

Chapter 8

442

Many metal complexes (particularly those of the transition metals) have affinity for binding oxygen in one of the ways shown in Figure 8.2. The nature of these bondings and also the catalytic and biological roles of M–O2 complexes have been the subject of several publications and reviews.24,25 The side-­on bridging mode (d) is rare; one reported example is a peroxodicopper(ii) complex with an unusually long O–O bond having a [µ-­η2 : η2(O2)]2− bridge, which has been found to be in equilibrium with the bis-­µ-­oxo-­diCu(iii) isomer.26 A special issue of Accounts of Chemical Research has been published in the year 2007 which includes several articles on O2 activation by metalloenzymes and models.27,28 Based on various evidence several bonding modes (a)–(d) are known (see Figure 8.2), in (a) the O2 is bound as O2− with the oxidation state of the M raised by one unit, in (b) depending on the nature of other ligands attached to the M the bonding may be expected to be of a type between the following two extremes:    (i) metal‒alkene-­like binding (Figure 8.3a); (ii) O22− bound to the M.    In case (ii), the oxidation state of the metal is raised by two units. The electron transfer in the M → O2 π interaction may be sufficient to transform the bound O2 to O2− and even O22−. In the O2 adduct of Vaska's compound [eqn (8.11)] (Figure 8.3b), the O–O distance is 1.30 Å, which in the analogous complex having I in place of Cl is 1.53 Å, and this effect of an Ir–I bond weakening the O–O bond of O2 bound to Ir is expected. One needs to compare these values with the O–O distances in O2− (1.28 Å) and O22− (1.48 Å). Hence

Figure 8.2  Bonding  modes of dioxygen.

Figure 8.3  (a)  Metal‒alkene-­like binding in M–O2 and (b) structure of [IrCl(CO)(O2) (PPh3)2]; Cl, CO, O2 and Ir are coplanar.

Activation of Molecules by Coordination and Reactivity of Coordinated Ligands

443

the latter (iodo compound) surely represents case (ii) above, and the former would correspond formally to that of O2− bound to Ir (with its oxidation state increased by one unit). However, the diamagnetism of the complex would correspond to IrIII(O22−) and it liberates H2O2 on acidification. In any case, the Ir–O2 bond is not like an M–alkene bond. The O2 binding is reversible29 for the chloro compound, but irreversible for the iodo compound.   



trans-­[IrCl(CO)(PPh3)2] + O2 → [IrCl(CO)(O2)(PPh3)2]

(8.11)

  

X-­ray determination of the structure of the O2 adduct of Vaska's compound, [IrCl(CO)(O2)(PPh3)2] (Figure 8.3b), shows that the O–O bond axis is perpendicular to the P–Ir–P trans pair of axial bonds; the midpoint of the O–O bond, CO and Cl occupy the three corners of a trigonal plane around Ir, which is almost at its centre.30 Considering Ir(iii)–O22− bonding, the overall geometry is roughly octahedral for Ir. Many Co(ii) complexes of quadridentate (2N,2O) Schiff base ligands, such as those resulting from condensation of two moles of salicylaldehyde with one mole of ethylenediamine, or two moles of acetylacetone and its derivatives with one mole of ethylenediamine, form O2 adducts particularly if a suitable base, B, is also present bonded to the Co(ii) (in a trans position to O2) in the adduct.31 The reaction (O2 binding) is mostly reversible and experimental evidence indicates Co(iii)–O2− binding. The third type of binding (as in Figure 8.2c) is also observed with many complexes of Co(ii), such as [Co(terpy)(AA)(OH2)]2+ (AA = bipy, phen), which reversibly add O2 [eqn (8.12)].   

(8.12)    

A similar reaction has also been reported for [Co(trien)(OH2)2]2+. Tetramesitylporphyrin–RhII, (TMP)RhII, and its hydride complex, (TMP)Rh–H, react with O2 to form the superoxo [(TMP)RhIII–O2], µ-­peroxo [(TMP)RhIII–O–O– RhIII(TMP)] and hydroperoxo [((TMP)RhIII(OOH)], complexes.32 Since the oxygen adducts in such complexes exist as either O2− or O22−, these are readily available for oxidation, as the thermodynamically hindered step (O2 + e− → O2−) has been overcome by binding to the metal. In the case of [CoIIL4(B)(O2)], the transformation into [CoIIIL4(B)(O2−)] and the role of the base B bonded trans to O2 can be easily explained on the basis of the MO diagram for the adduct (Figure 8.4). Considering the ligand field effect, if the adduct base B is a strong-­field ligand it will considerably raise the energy of the unpaired electron of Co(ii) in the strong-­field square-­planar

444

Chapter 8

Figure 8.4  Energy  level diagram (qualitative) for the d orbitals of Co(ii) in (a) square-­ planar and (b) square-­pyramidal fields in CoL4 and CoL4(B), respectively; (c) MO diagram (qualitative) for the Co–O2− bond in CoL4(B)(O2) involving d orbitals of Co(ii) in CoL4(B) with the antibonding π MOs of O2. Only 7e of Co(ii) are shown in (a) and (b); the electron distribution due to Co–O2 bond formation is shown in (c) and antibonding pi MOs of O2 shown in (d).

CoL4 when this forms square-­pyramidal CoL4(B) (Figure 8.4b). MO formation by O2 with appropriate orbitals of cobalt in CoL4(B) (Figure 8.4c) will bring the unpaired electron of the metal to an antibonding π MO, which is essentially localized on the O2, being close in energy to the π* O2 orbitals. This is equivalent to transfer of an electron from the Co(ii) to O2, thereby forming Co(iii) and O2−, respectively. The affinity of [Co(acacen)(B)] for O2 increases with increase in the basicity of B.31 ESR spectroscopic studies of [Co(salen)(py)(O2)] labelled with 17O have shown that the unpaired electron resides nearly 100% in an MO almost localized on the O2.33 The O–O distance in this compound (from crystallographic data) is 1.26 Å (compared with 1.28 Å in the O2− in KO2 and 1.21 Å in free O2). Equilibrium studies on O2 absorption by a number of Co(ii) complexes of the type CoL4(B), L4 being a variety of quadridentate Schiff bases (2N,2O donors) and various B, viz. pyridine and substituted pyridines, substituted imidazoles, t BuNH2, etc., have shown that the plot of log Keq for O2 absorption (reversible) versus E½ for the polarographic reduction of the complex [E½ is related to the redox potential of the Co(iii)‒Co(ii) couple] is linear;31c log Keq increases with increasing negative value of E½, i.e. increase in difficulty of reduction of Co(iii) to Co(ii) [hence increasing stability of the Co(iii) state in the complex]. This is in accord with the fact that the O2 absorption leads to the formation of a CoIII– O2− bond. Examples of O2 activation by complexes of nickel28 and by low-­valent metal complexes of Cr, Co and Rh34 have been reported.

Activation of Molecules by Coordination and Reactivity of Coordinated Ligands

445

8.2.3  Activation of Dinitrogen by Coordination35 The discovery of stable complexes of N2 opened up the possibility of nitrogen fixation with such complexes as intermediates.36 An important development in this area was the discovery of some phosphine complexes of Mo and W in zero oxidation state also having bound N2, which readily give NH3 in fairly good yield in acidic media (Scheme 8.6).37 Similar results were obtained with [M(N2)2(dppe)2] (M = Mo or W; dppe = Ph2PCH2CH2PPh2)†.38 Both reactions take place at room temperature and normal atmospheric pressure. The reduction could be effected by Na–Hg to make the complex containing N2. With the complexes of PR3 the yield of NH3 was ca. 0.7 NH3 per molybdenum atom, but with the corresponding tungsten compound a better yield of ca. 1.7 NH3 per W atom and little N2H4 were obtained; the rest was liberated as N2 in both cases. The acid used was a dilute solution of H2SO4 in methanol in the NH3 generation step. The overall reaction is expressed by eqn (8.13).   

N2 + M(0) + 8H+ → 2NH4+ + M(vi)



(8.13)

  

By judicious choice of the reagents and solvent, Chatt and co-­workers were able to isolate and fully characterize several intermediate species having MNNH, MNNH2, MNNH3+, MN, MNH and MNH2 from the hydrazine-­ and ammonia-­producing reactions, and also to observe their interconversion and degradation in reaction solutions by 15N NMR spectroscopy and other techniques.37–39 This led to the proposal of a cycle of reduction of N2 to NH3 at a single metal site as shown in Scheme 8.7, which was denoted the Chatt cycle,37,40 and the suggestion that it could be operative in nitrogenase if a single metal site is involved, which of course is still not known.41 Another proposal for N2 reduction at a mononuclear Mo centre is the Schrock cycle (Scheme 8.8).42 The electrons shown in these proposed schemes are provided by the lower-­ valent Mo and as a result the Mo(0) ultimately becomes Mo(vi). In some such

Scheme 8.6

†



Chapter 8

446

Scheme 8.7

Scheme 8.8 reaction systems N2H4 is also formed with NH3. This is likely to be due to the following transformations:   

 (8.14)   

Schrauzer et al.43 reported that a solution of molybdate and thiol ligands such as l-­(+)-­cysteine slowly reduces N2 to NH3 in the presence of NaBH4; the presence of ATP significantly stimulates the reduction. Based on such findings, it was proposed earlier that reduction of N2 to NH3 by nitrogenase enzyme occurs at Mo in its active site. However, the evidence is not in conformity with this suggestion.44 Leigh45 proposed a cycle for N2 reduction to NH3 at the Fe(ii) centre of a bis(diphosphine) complex of Fe(ii), with the Fe cycling between +2 and

Activation of Molecules by Coordination and Reactivity of Coordinated Ligands

447

zero oxidation states. Various workers have proposed Chatt-­t ype cycles for the reduction of N2 on Fen+ centres of synthetic models having Fe in +1 and +2 oxidation states that forms both N2H4 and NH3 as the reduction products, with the Fe cycling between Fen+/Fe(n+1)+/Fe(n+2)+/Fe(n+3)+ states.46 Based on knowledge from N2 reduction using metal complexes, an electrochemically driven cycle for the reduction of N2 to NH3 was developed.47 Incidentally, van Tamelen et al.48 reported that titanium(iv) isopropoxide, Ti(OiPr)4, in the presence of a strong reducing agent such as sodium naphthalide, Na+(C10H8)−, reduces N2 to NH3 under ambient conditions (Scheme 8.9). The reduction can also be effected electrochemically; some hydrazine is formed along with ammonia. An earlier and indeed the first such observation was by Vol'pin and Shur,49 who showed that N2 can be reduced by strong reducing agents, such as RMgX, LiA1H4, LiR, AlR3 and Mg + MgI2, in the presence of compounds of various transition metals (Ti, V, Cr, Fe, Mo, etc.) in aprotic media. Generally complex nitrides of the metals were produced which formed ammonia on treatment with acid. Reduction of N2 in protic media was observed for the first time by Shilov in 1970; since then several systems [having Ti(ii), Mo(iii), V(ii), Nb(iii), Ta(iii) and an appropriate reducing agent] have been found that were capable of reducing N2 to NH3, N2H4 or NH3 + N2H4 in water or methanol.50 Shilov also discussed all earlier work on mono-­ and dinuclear transition metal complexes having N2 as a ligand that are capable of releasing some part or nearly all of the bound N2 as NH3 and/or N2H4 on treatment with acid (with reducing agents such as Na–Hg or Zn–Hg in some cases); a few examples of such dinuclear complexes are given in eqn (8.15)‒(8.19). HCl was used in all cases except those with Zn–Hg or Na–Hg, where lutidine·HCl or lutidine· HOSO2CF3 was used as the proton source.   



[{ZrCp∗2(N2)2}2(µ-­N2)] → N2H4(33%) + N2(67 %)

(8.15)

[{V(1,2-­Me2N−C6H4)2py}2(µ-­N2)] → N2H4(33 %)

(8.16)

  



Scheme 8.9

Chapter 8

448   



M  S2 CNEt 2    μ  N 2    N 2 H4  100%  3 2    M  Nbor Ta 

  

(8.17)

MoCP *2 Me3   μ  N 2    NH3  16%    2



  



  

 withZn-Hg, NH  32 - 36%

(8.18)

3

 WCP *2 Me3   μ  N 2    NH3  17%  2    withZn-Hg, NH3  38% 

(8.19)

Activation of N2 by a heterodinuclear Nb–Mo complex (Scheme 8.10) and P4 by an Nb complex (Scheme 8.11) has been reported.51 Conversion of N2 to RCN using an Mo(iii) complex has also been reported.52 The molybdenum complex of a triamidoxime ligand converts N2 to NH3 at room temperature and atmospheric pressure, where the Mo shuttles between +6 and +3 states.42 Activation/fixation of N2 by complexes of scandium,53a yttrium,53b samarium,54g titanium,54a zirconium,54b,c hafnium,55a,56 vandium,54d niobium,54e tantalum54f and molybdenum57,58 have been reported. Activation of N2 by soluble metal complexes and base metal atoms55,59 and activation of N2 by Group 4 metal complexes60 have been reviewed, and also catalytic reduction of N2 by d-­block metal complexes.61

Scheme 8.10  Conversion  of N2 to RCN (OTf = F3CSO3−).

Activation of Molecules by Coordination and Reactivity of Coordinated Ligands

449

Scheme 8.11  Conversion  of P4 to a phosphide complex.

A molybdenum complex of a PNP “pincer” ligand binds N2 to form a dinuclear complex (1), which, in the presence of a proton source (2,6-­lutidinium trifluoromethanesulfonate) and an electron source (CoCp2), catalyses the conversion of N2 to NH3 very efficiently under ambient conditions.62 Another similar efficient catalyst is a dinuclear (N2-­bridged) Mo complex of an unsymmetrical PNP ligand having a P(adamantyl)2 in place of one PtBu2 on each of the PNP ligands in 1 has been reported.63 The first N2 complex reported64 was [Ru(NH3)5(N2)]2+, but this shows no activation of N2 and no reduction of its N2 ligand to NH3 or N2H4 has been achieved so far.

8.3  Reactivity of Coordinated Ligands Reactions of metal complexes of N2 mentioned in the preceding section are examples of this type.

8.3.1  Reaction of Metal-­bound CO Ligand Carbon monoxide bound to a metal of low π basicity is very sensitive to nucleophilic attack.65 A well-­known example of this is the water-­gas shift reaction (CO + H2O ⇌ CO2 + H2) in which CO bound to a metal in a complex is activated to undergo nucleophilic attack. The proposed cycles for the Fe(CO)5-­ and Pt(PiPr3)3-­catalysed processes are shown in Schemes 8.12 and 8.13.

Chapter 8

450

Scheme 8.12

Scheme 8.13 Another example of nucleophilic attack on a CO ligand in a metal complex is shown in eqn (8.20).   



(8.20)

  

Because of the positive charge, the CO ligand in [Mn(CO)6]+ is more sensitive to nucleophilic attack [eqn (8.21)] than in the uncharged Mo(CO)6.

  



(8.21)

  

Nucleophilic attack by MeOH instead of H2O can give an ester [M(CO2R)Ln] [eqn (8.22)], which is stable as it does not have β-­H.

  

MeOH,Et 3 N CO –Et 3 N ·HCl PtCl2  PR 3 2   PtCl  CO   PR 3 2  Cl  PtCl  CO2 Me   PR 3 2 



(8.22)

  

8.3.2  Reactions of Coordinated CO2 and SO2 Eqn (8.23) shows a typical example of a reaction of coordinated CO2 forming a metal ketene complex (Wittig reaction).66 Insertion of CO2 into M–C and M–N bonds may also involve initial coordination of CO2 [eqn (8.24) and (8.25)].   

(8.23)    

Activation of Molecules by Coordination and Reactivity of Coordinated Ligands



451

CO2 ,150° C ZnEt 2   Zn  O2 CEt 2 

(8.24)

CO2 ,20° C Ti  NMe2 4   Ti  OCONMe2 4

(8.25)

  

  

Coordinated SO2 undergoes insertion into an M–C bond as shown in eqn (8.26).   



[PtPh(Cl)(PEt3)2(SO2)] → trans-­[Pt(SO2Ph)(Cl)(PEt3)2]

(8.26)

  

8.3.3  Reactions of Coordinated NO in Nitrosyl Complexes Various reactions of coordinated NO in nitrosyl complexes are known, as illustrated with the following examples.

8.3.3.1 Nucleophilic Addition This reaction is the basis of the nitroprusside test for HS− (and hence of H2S) and RS− in alkaline media that leads to the formation of a red‒violet product [eqn (8.27)]. The reaction with EX− = HO− proceeds further with 2 reacting with HO− to form [(NC)5Fe–NO2]4−, and both steps are reversible so that on acidification [(NC)5Fe–NO]2− is re-­formed.   

(8.27)



  

Other examples of this class of reactions are given in eqn (8.28) and (8.29).    PhNH2 Ru  NO   Cl   bipy 2   RuN  O   NH2Ph   Cl   bipy 2  2

  

  Ru=NPh  Cl   bipy 2   H2 O

2

(8.28)

2







(8.29)

8.3.3.2 Electrophilic Addition The product 4 formed in the reaction shown in eqn (8.30) is the first example of a complex with an HNO ligand (which itself is a fugitive species).67   





(8.30)

Chapter 8

452

8.3.3.3 Oxygenation/Oxidation Some nitrosyl complexes add O2 [eqn (8.31)].   

2[LxM(NO)] + O2 → 2[LxM(NO2)]



(8.31)

8.3.3.4 Reduction Reversible 1e− reduction [eqn (8.32)] and non-­reversible reduction [eqn (8.33)] are known. The product in eqn (8.33) has a deprotonated hydroxylamine, H2NO, ligand.68 The following are a few typical examples:   

  

 e

-

 Et 2 NCS2 2 Mo  No 2  #  Et 2 NCS2 2 Mo  No 2  e

  



  

(8.32a)

2

(8.32b) 

Zn,H RuCl  py 4  NO     RuCl  py 4  NH3    H2 O 

(8.32c)

THF 2  S4  Mo  NO 2  + N 2 H4    2  S4  Mo  H2 NO  NO    N 2



S4  a tetrathiolate

 2−

II

(8.33)

3−

Two-­electron reduction of [Fe(CN)5(NO)] forms [Fe (CN)5(HNO)] containing a novel HNO ligand that is stable at pH 6, but on increasing the pH it forms [FeII(CN)5(NO)]4−.69

8.3.3.5 Reaction with Alkenes This is a synthetic route for amines [eqn (8.34)].   



(8.34) 70a,b

8.3.4  Reactions of Some Coordinated Organic Ligands

Reactivity to nucleophilic attack on ligands is increased by (a) an overall positive charge on the complex, (b) the presence of other π-­acid ligands, (c) coordinative saturation of the metal (which prevents attack on the metal) and (d) high nucleophilicity of the attacking reagent. Ligand displacement competes, and it is not always possible to predict which will occur. Nucleophilic addition to an η4-­diene is shown in eqn (8.35).   

(8.35)   



Activation of Molecules by Coordination and Reactivity of Coordinated Ligands

453

Cyclopentadienyl anion is the most common η ligand, and is usually inert to nucleophilic attack except when the complex is cationic and the attack is by a strong nucleophile [eqn (8.36)]. 5

  

(8.36)



  

Cyclohexadienyl complexes are more reactive to nucleophilic attack [eqn (8.37)].   

(8.37)





  

An example of nucleophilic attack on a coordinated η6-­arene is shown in eqn (8.38).   

(8.38)

   

The neutral product of two hydride additions could be expected to be [Ru(η6-­C6H7)2], since attack at an even ligand is expected to be favoured. This expected pattern is observed in the analogous iron complex. As usual with other π-­donor ligands, the free molecules are unreactive towards nucleophiles as their electron-­rich multiple bonds favour reaction with electrophiles. Factors that promote attack by electrophiles include (a) zero or negative charge on the complex, (b) low oxidation state of the metal and (c) late transition metals (having more d electrons).70c Because transition metals generally contain some non-­bonding electrons, the metal may be the initial site of electrophilic attack followed usually by transfer to a π-­donor ligand. For example, ferrocene can be protonated by

Chapter 8

454 +

strong acids to form [Cp2FeH] . Protonation of a hexamethylbenzene complex converts a “triene” ligand into a dienyl ligand [eqn (8.39)].   



(8.39)

  

Enyl complexes generally cannot be protonated. An exception is the 20e Ni(ii) complex, which is converted to an 18e species by protonation [eqn (8.40)].   

(8.40)





  

The Ph3C+ cation is commonly employed to abstract H−. Hydride abstraction at a CH2 group makes a π orbital free on C and converts polyene ligands into polyenyls containing one more C. Thus, η4-­cyclohexadiene is converted to η5-­ cyclohexadienyl and η6-­cycloheptatriene is converted to η7-­cycloheptatrienyl. Also, η6-­cycloocatriene is converted to η7-­cyclooctatrienyl [eqn (8.41)].   

(8.41)

   

Activation of Molecules by Coordination and Reactivity of Coordinated Ligands

455

71

Cobalt(ii) complexes have been used as catalysts in the synthesis and controlled hydrolysis of peptides. Both [CoCl(H2NCH2COOEt)(trien)]2+ and [Co(H2NCH2COOEt)(trien)]3+ react rapidly and almost quantitatively with glycine ester giving [Co(H2NCH2CONHCH2COOEt)(trien)]3+, hence the same product is obtained from cis-­[CoCl2(trien)]+ and glycine ester, glyOEt, but no peptide formation occurs in the absence of the cobalt(iii) complex. The reaction can be used to link glycine to a variety of amino acids and peptide esters, and compounds such as [Co(trien)(glyglyNH2)]3+, [Co(trien)(glyalaOEt)]3+ and [Co(trien)(glyglyglyOEt)]3+ (trien = triethylenetetramine, gly = glycine unit, ala = alanine unit) can be prepared from cis-­[Co(trien)(glyOEt)]2+ and an appropriate amino acid, amino acid amide or peptide ester. Hydrolysis (in alkaline solution) of the peptide chelate leads to hydrolytic elimination of the peptide chain, leaving the metal-­bound terminal amino acid coordinated to the metal. [Co(OH)(OH2)(trien)]2+ reacts with peptides, amino acid esters and amino acid amides, hydrolysing them and leaving the N-­terminal amino acid residue coordinated to the metal. The metal chelates formed by acetylacetonate (acac−) represent an interesting field of reactivity of coordinated ligands. Extensive investigations have been made on the reactions of M(acac)3 [M = Cr(iii), Co(ii), Rh(iii)], which have pseudo-­aromatic character because of electron delocalization in the M–acac chelate rings.72 Electrophilic substitution (as in benzene) takes place easily at the carbon atom of CH of the chelate ring, which leads to replacement of the hydrogen attached to this carbon with an electrophile, E+. These reactions include halogenation, thiocyanation, acylation, formylation, chloromethylation and aminomethylation and proceed by the plausible mechanism shown in eqn (8.42).   



(8.42)

  

Substitution may be effected in one, two or all three chelate rings, except in acylation and formylation, which do not yield the tris-­substituted chelates. In support of the aforementioned mechanism is the observation that when an optically active (optical isomer) chelate is used, the substitution takes place without racemization, which indicates that no opening and closing of the chelate rings take place in the process. Similar studies were also reported on metal chelates of 8-­hydroxyquinolines.73 Many applications of the activity of coordinated ligands have been made use of in template reactions for the syntheses of multidentate macrocyclic ligands, comparable to those which occur in Nature (such as vitamin B12, haemoglobin, haemocyanin, chlorophyll, etc.) and which are difficult to synthesize by conventional methods where polymerization may take place or by-­products of undesired nature may result. It is now possible to use metal

456

Chapter 8

complexes as templates to control the steric course of multistep reactions that produce macrocyclic ligands. Developments in the field were reviewed by Busch and co-­workers.74 Martell and Calvin75 reported that the formation of a quadridentate Schiff base ligand (baenH2) from acetylacetone and ethylenediamine is catalysed by Cu(ii) ions and the reaction probably involves attack on the Cu(ii)‒acetylacetonate complex [eqn (8.43)].   





(8.43)

  

When anhydrous [M(en)3]2+ [M = Ni(ii), Cu(ii)] complexes are refluxed with anhydrous acetone (in the presence of anhydrous CaSO4 to remove water formed in the reaction), or heated at 110 °C in a sealed tube, or irradiated under UV light, compound 5 is formed as a result of controlled condensation reactions leading to elimination of water and H+.76

Similar compounds can be obtained using α-­ and β-­hydroxyketones instead of acetone. Reduction of these compounds by catalytic hydrogenation or electrochemical methods gives saturated products, from which the metal can be removed electrochemically or chemically, such as by treatment with cyanide solution, to release the quadridentate macrocyclic amine 6.

Activation of Molecules by Coordination and Reactivity of Coordinated Ligands

457

This has gone a long way to providing a clue as to how the very complex metal porphyrins present in biological systems are formed. Metal ion-­ and metal complex-­catalysed reactions of substrates are well known and many of these mimic metalloenzyme-­catalysed reactions of substrates in biological systems that proceed through the formation of an intermediate having the substrate bonded to the metal. Some typical examples are given below.

8.3.5  Catalysed Oxidation of Coordinated Ligands Oxidation of ascorbic acid by molecular oxygen is catalysed by Cu2+ (and also by Fe3+) and involves the formation of an O2–CuII–ascorbate mixed-­ligand complex (and similarly for the Fe3+ case),77 as indicated in Scheme 8.14. The oxidation (in the presence of Cu2+ or Fe3+) is first order in the substrate, the O2 and the metal ion, which agrees with the proposed scheme. In the presence of a Cu(ii) chelate, however, a different mechanism operates,78 since the rate becomes independent of the concentration of O2 employed, and O2 is reduced to H2O instead of H2O2. The mechanism that appears to operate here is shown in Scheme 8.15. The mechanism of oxidation by the enzyme ascorbic acid oxidase (ascorbate oxidase) appears similar,79 involving two successive one-­electron steps, with the formation of a free radical intermediate (which was detected by ESR

Scheme 8.14  Mechanism  of Cu2+-­catalysed oxidation of ascorbate ion.

Chapter 8

458

Scheme 8.15  Mechanism  of oxidation of ascorbate catalysed by a Cu(ii) chelate, [CuIIL] (L = trien), an ascorbate oxidase model.

Scheme 8.16  Mechanism  of oxidation of catechol catalysed by an Mn(ii) chelate, a tyrosinase model; a similar mechanism occurs with other M(ii) species (see text).

spectroscopy),80 followed by re-­oxidation of the CuI‒enzyme to CuII‒enzyme by O2, which is reduced to H2O. Oxidation of catechol by O2 catalysed by Mn(ii) chelates is believed to occur according to the proposed general mechanism shown in Scheme 8.16.81 The order of catalytic activity of the metal ions investigated is Mn(ii) > Co(ii) > Fe(ii) > Cu(ii) > Ni(ii).

Activation of Molecules by Coordination and Reactivity of Coordinated Ligands

459

The enzymatic oxidation of catechols to quinones by O2 with the enzyme tyrosinase is similar. The only difference is that the stable form of the enzyme contains Cu(i) in the absence of O2 and substrate, and O2 is reduced to H2O rather than H2O2. The reaction is actually a two-­electron oxidation. It is believed that the H2O2 remains bound to the metal centre after dissociation of the quinone and oxidizes a second molecule of the substrate. Oxidation of α-­amino acids to α-­keto acids and NH3 by O2 in the presence of pyridoxal and transition metal ions represents a model for the enzymatic oxidation brought about by amine oxidase. Such reactions of model compounds have been investigated by several groups.82 They involve the formation of a Schiff base complex as the first step, as shown in Scheme 8.17. For several bivalent metal ions of the first transition series, the relative rate of catalysis was as follows:83 Mn(ii) > Co(ii) > Cu(ii) >> Ni(ii). The catalytic activity of Ni(ii) was found to be very slight, almost indistinguishable from the spontaneous rate. A similar mechanism may apply to amine and amino acid oxidase enzymes, which require Cu(ii) and pyridoxal phosphate for activity (in pyridoxal phosphate the –CH2OH group of pyridoxal is replaced with –CH2OPO3H). Many metal chelates catalyse the decomposition of H2O2 into H2O and O2 and thus represent models of the enzymes called catalases. For full activity, the complex metal ion must have at least two coordination positions occupied by easily replaceable H2O molecules.84 Thus, in case of Cu(ii), [Cu(en)(OH2)2]2+ has high activity, that of [Cu(dien)(OH2)]2+ is much less and [Cu(trien)]2+ shows no activity at all, but [Mn(trien)(OH2)2]2+ and [Fe(trien) (OH2)2]3+ are active.85 The difference is due to the two replaceable H2O bound to these Mn+ ions, because of the preference of Mn(ii) and Fe(iii) for six coordination, which permits binding of two HOO− to form the reaction

Scheme 8.17

Chapter 8

460

Scheme 8.18  Mechanism  of decomposition of H2O2 catalysed by Mn(trien) (OH2)2]2+ (catalase model).

intermediate (Scheme 8.18). On the basis of the observation that the reaction is second order in peroxide concentration when Mn2+ is the catalyst, Hamilton86 proposed the mechanism shown in Scheme 8.18. The proposed mechanism visualizes electron transfer thorough the Mn2+ from one bound HO2− to the other forming the products: HO2− + HO2− → 2(OH)− + O2, and this presumably occurs through a transient Mn4+ intermediate [eqn (8.44)].   



Mn2   OOH 2  O   O2 2     Mn   O   O2  (8.44)  Mn 4     H2 O 2

  

The Mn(ii) complex [MnII2(H2dapsox)(MeOH)(OH2)][ClO4]2·H2O [H2dapsox = 2,6-­diacetylpyridine bis(semioxamazide)], which rapidly catalyses the disproportionation of superoxide (hence a superoxide dismutase mimic), has been synthesized.87 Peroxidases are enzymes that catalyse the oxidation of many organic compounds, such as aromatic amines and phenols, by H2O2.88 The Fe(iii)– H2O2–catechol ternary system serves as a model for peroxidase activity.89 The proposed mechanism is shown in Scheme 8.19. Similarly, RCH2OH can be oxidized by H2O2 to RCHO (overall reaction: RCH2OH + H2O2 → RCHO + 2H2O).

8.3.6  Amino Acid Ester Hydrolysis90 These reactions are also catalysed by metal ions through the formation of metal ion–ester complexes. For hydrolysis of ethyl glycinate (EG), the highest rate is observed for Cuaq2+ ion and the rate decreases on coordination

Activation of Molecules by Coordination and Reactivity of Coordinated Ligands

461

Scheme 8.19  Mechanism  of Fe(iii)-­catalysed oxidation of catechol by H2O2 (a model for peroxidase activity).

of Cu(ii) to another ligand. The observed rate of base hydrolysis is kobs = k0 + k1[OH−]. Denoting k1(complex)/k1(free metal ion) = k′, the observed values of k′ for [Cu(EG)(OH2)4]2+, [Cu(ida)(EG)(OH2)] and [Cu(nta)(EG)]− are 1.2 × 105, 2.9 × 104 and 1.2 × 102, respectively, under comparable conditions.91a Comparison of the hydrolysis rates of the ethyl esters of α-­ and β-­alanine by [Cu(nta)(OH2)2]− shows that the k′ values are 2.8 × 102 and 30, respectively; hence interaction of the metal ion with the carbonyl oxygen of the β-­alanine ester 7 is much less than for the α-­alanine ester 8, owing to the weaker coordination of carbonyl oxygen in a six-­membered ring.

Cobalt(iii) complexes also catalyse the hydrolysis of amino acid esters,91b through the formation of ternary complexes as in the Cu(ii) examples mentioned above. Studies by the pH-­stat method have been reported for the acid and base hydrolysis of a series of α-­amino acid esters chelated to Pd(ii) in [Pd(en)(H2NCHR'CO2R)]2+; compared with the free esters, base hydrolysis of the chelated esters is 104‒105 times faster.92a Similarly, an amino acid ester chelated to (bipy)Pd2+ hydrolyses much faster than the free ester.92b

8.3.7  Decarboxylation of β-­Keto Acids A typical example of an enzyme-­catalysed reaction is decarboxylation of β-­keto acids, a reaction of biological importance. Westheimer and co-­workers93,94 obtained evidence that the decarboxylation of α,α-­dimethyloxaloacetic acid, and presumably also of other β-­keto acids, proceeds by a mechanism

462

Chapter 8

involving the dianion of the acid, as illustrated using oxaloacetate as an example (Scheme 8.20),95 where M2+ is either an aqua metal ion or the metal ion complexed with the enzyme protein. A similar mechanism operates for other β-­keto acid anions. At 25 °C, the Cu2+-­catalysed96 decarboxylation of the dianion of oxaloacetic acid, −O2C–COCH2CO2−, occurs ca. 104 times faster than the spontaneous97 (uncatalysed) decarboxylation. However, compared with the spontaneous reaction, the metalloenzyme-­catalysed reaction is ca. 109 times faster. Also, it is found that in the simple metal ion-­catalysed reaction there is a close relationship between the catalytic activity of the metal ion and the formation constant of the metal complex; in fact, the catalytic activity follows (more or less) the Irving and Williams98 natural stability sequence for complex formation, viz. Mn(ii) < Fe(ii) < Co(ii) < Ni(ii) < Cu(ii) > Zn(ii). Munakata et al.99 observed that the catalytic activities of metal ions in the decarboxylation of oxaloacetic acid are related to the electronegativities of the metals. Gelles and Salama determined the rate constants (kMA)100 for the decarboxylation of the 1 : 1 complexes (MA) of the dianion (A2−) of oxaloacetic acid with a number of metal ions, and also the formation constants101 of these complexes. These values are given in Table 8.1, together with the rates for uncatalysed (spontaneous)96 decarboxylation of oxaloacetic acid and its anions. It should

Scheme 8.20  Mechanism  of decarboxylation of oxaloacetate catalysed by a metal

ion, M2+ [H2O ligands attached to M(ii) are not shown]. The enolic complex (b) is reported to be inactive.95 For the 3-­oxoglutaric acid‒Cu2+ system, the pKa value for the equilibrium a ⇌ b + H+ is 3.81 at 25 °C (I = 0.6 M); a = keto form and b = enol form (inactive) of the metal complex.

Activation of Molecules by Coordination and Reactivity of Coordinated Ligands

463

be noted that the copper(ii) complex is about 10 times more reactive than the manganese(ii) complex. All these results clearly show that additional factors are operative in the natural enzymatic reactions, where manganese is the enzyme activator for such decarboxylating enzymes. The metal ion-­catalysed and enzyme-­catalysed decarboxylations also differ in some finer mechanistic details. Thus, the Mn(ii)-­catalysed decarboxylation of oxaloacetate dianion shows a 6% carbon isotope effect (C−12k/C−13k = 1.06, shown using HO2CCOCH213CO2H),94 indicating that the cleavage of the carbon–carbon bond occurs in the rate-­determining step (i.e. k is rate determining; see Scheme 8.20). The enzymatic reaction with Mn(ii)-­activated enzyme does not show the carbon isotope effect, but the reaction is slower in D2O than in H2O, while the rate of the Mn(ii)-­catalysed reaction is unaffected by the solvent change (hence in the enzymatic reaction k′ is the rate determining step). For metal ion-­catalysed decarboxylation of oxaloacetic acid (H2A) (Table 8.1), the plot of log kMA versus log KMA is not linear, and in fact it does not even form a smooth curve. However, the plot of log kMA versus log KMox (where KMox is the formation constant of the monooxalato complex of the metal ion M2+) is linear. This linear free energy relationship (LFER) lends support to Hammond's postulate102 that the transition state of the decarboxylation should closely resemble the intermediate a in Scheme 8.20. A similar result was noted in the sequence of transition metal ion-­catalysed decarboxylation of 3-­oxoglutaric acid (acetonedicarboxylic acid, HO2CCH2COCH2CO2H), where the catalytic activities of the metal ions follow the sequence of the thermodynamic stabilities of the corresponding malonates.103 These observed free energy relationships are closely comparable to what has been observed in the metal ion-­catalysed dissociation of the metal chelates of oxalate104 and biguanide105 ligands. Rund and co-­workers106 studied the kinetics of the decarboxylation of α,α-­ dimethyl oxaloacetate catalysed by Mn(ii) and Ni(ii) in aqueous solution,

Table 8.1  Decarboxylation  rate constants97,100 and formation constants101 of 1 : 1

metal oxaloacetates (MA) and decarboxylation rate constants of the metal-­free species97 (37 °C, I = 0.1 M).

Species H2A HA− A2− CaA MnA CoA NiA CuA ZnA a

k = kspon or kMA.

104k/s−1a

Log KMA

0.058 2.57 0.7 24 65 240 230 660 310

— — — 2.6 2.8 3.1 3.5 4.9 3.2

Chapter 8

464

in the presence and absence of dioxane (Table 8.2) and of various chelating ligands, in an attempt to set up enzyme models for understanding enzymatic actions. The observed first-­order rate constants were found, except at fairly high dioxane concentrations, to be proportional to the catalyst concentration, and increased by as much as three orders of magnitude (103) [both Mn(ii)-­ and Ni(ii)-­catalysed rates were enhanced to the same extent] on addition of dioxane owing to an increased association constant (K) for the catalyst‒substrate complex due to lowering of the dielectric constant (D) of the medium. Certain complex ions of Mn(ii) and Ni(ii) also catalyse decarboxylation, provided that the complexes have some easily replaceable aqua ligands bound to the metal. Complexes such as [M(phen)3]2+ or [M(edta)]2−, in which all six coordination positions of the metal ion are occupied by the bound ligands, were found to be inactive. This shows that the β-­keto acid binds to the metal bound to another ligand, forming a ternary complex, prior to the decarboxylation; hence in the presence of replaceable aqua ligands in the complexes, with anionic oxygen donor ligands, such as 8-­hydroxyquinoline-­5-­sulfonate, the catalytic activity was lower than that of the uncomplexed metal ions. Other ligands, such as 1,10-­phenanthroline, enhanced the catalytic activity of Mn(ii) greatly, but hardly in the case of Ni(ii) (Table 8.2). A similar effect was also observed107 in the case of the dianion of oxaloacetic acid. 2,2′-­Bipyridine has a similar effect in greatly enhancing the catalytic activity of Mn(ii) (ca. 10-­fold) but that of Cu(ii) and Ni(ii) only slightly (ca. twofold); significantly, acetate ion has little effect (see the data in Table 8.3).108,109 Transition metal ions also catalyse the hydrolysis of amino acid esters.90 In the catalysed hydrolysis of methyl glycinate, Cu2+ is 104 times more active than Cu(gly)+ for the Cu2+-­catalysed reaction and the rate is 108 times that for the uncatalysed reaction.110 In the absence of metal ions, 2,2′-­bipyridine and 1,10-­phenanthroline ligands have no effect on the rate of decarboxylation, indicating the involvement of a ternary complex of the metal ion as an intermediate; in fact, a 1 : 1 : 1 ternary complex of Cu(ii) with 5-­nitro-­1,10-­phenanthroline and oxaloacetate was isolated.111 These observed results on the preferential activation of Mn(ii) qualitatively resemble the behaviour of the manganese-­ activated enzyme oxaloacetate decarboxylase, and it was found to be due to Table 8.2  Decarboxylation  of α,α-­dimethyl oxaloacetate at 25 °C. Catalyst

Mediuma

Mn2+ Mn(phen)2+ Ni2+ Ni(phen)2+ Ni2+ Ni(phen)2+

H2O (D = 78) H2O (D = 78) H2O (D = 78) H2O (D = 78) H2O + dioxane (D = 60) H2O + dioxane (D = 60)

a

 = dielectric constant of the solvent. D Kk = kcat.

b

103k/s−1

K/M−1

103Kk/M−1 s−1b

0.11 2.5 2.07 8.33 3.5 6.33

61 11 95 20.3 1.9 × 103 6.6 × 102

6.7 27.5 196.7 169 6650 4178

Activation of Molecules by Coordination and Reactivity of Coordinated Ligands

465

Table 8.3  Catalysed  decarboxylation of 3-­oxoglutaric acid (AH2) at 25 °C a

[pH = 4.66, I = 0.1 M (NaOAc)].108

Catalyst

kcat/M−1 s−1

Cu2+ Cu(OAc)+b Cu(bipy)2+ Ni2+ Ni(bipy)2+ Mn2+ Mn(bipy)2+

0.41 0.51 0.79 5.5 × 10−3 9.0 × 10−3 2.0 × 10−4 2.0 × 10−3 2.0 × 10−4c

a

HO2 CCH2 COCH2CO2H   MeCOCH2CO2H   MeCOMe .  CO CO ,slow 2

b

109

2 not catalysed by M n+

Values reported by Prue; Cu(OAc)2 is inactive.  his is kspon for the spontaneous decarboxylation of AH2 (in strongly acidic solution, [H+] >0.2 M), T i.e. in the absence of a metal catalyst; for the monoanion, AH−, kspon = 7.5 × 10−4 s−1 and for the dianion, A2−, kspon = 4.7 × 10−5 s−1.

c

an effect on the specific rate of decarboxylation of the active complex. Rund and Plane106a also found that lowering the dielectric constant of the solvent greatly enhances K, the catalyst‒substrate association, but has little effect on k, its rate of decarboxylation (Table 8.2). Conversely, the effect of coordination by 1,10-­phenanthroline has a small effect on K (it decreases slightly as expected on statistical considerations), but in the case of Mn(ii) it greatly increases the rate constant k (by a factor of ca. 23), whereas in the case of Ni(ii) k increases by a factor of only ca. 4 (Table 8.2). The fact that k for Ni(ii) is not so strongly affected in these simple model systems suggests that the specific enhancement of the activity of Mn(ii) by the enzyme protein is presumably not just the result of a lowered dielectric constant (in the vicinity of the binding site), which would cause stronger binding of the substrate, but to a specific coordination of the metal by the enzyme, such as to greatly speed up the decarboxylation of the bound substrate. Edta4− destroys the catalytic activity of Ni(ii) by binding to all the coordination positions, thus preventing direct binding to the substrate, which is apparently necessary for the reaction. Similar results were found for high concentrations of other ligands, such as 1,10-­phenanthroline, due to the formation of [Ni(phen)3]2+ with no vacant coordination site. Less obvious is the low reactivity of 8-­hydroxyquinoline-­5-­sulfonate complexes. A possible explanation is that this negatively charged ligand lowers the effective charge of the metal ion, and it has been found previously that ions bearing a 1+ charge do not have catalytic activity.112 Some literature data are summarized in Table 8.4,113 which clearly demonstrate the relative effect of K and k on kcat for enhancement of the activity of different metal ions by bipy or phen. The effect of 1,10-­phenanthroline or 2,2′-­bipyridine in enhancing the catalytic activity presumably arises from the back-­bonding (π, M → L), leading

Chapter 8

466

Table 8.4  Formation  constants and rate constants for decarboxylation of binary a

(MA) and ternary (MAB) complexes of 3-­oxoglutaric acid at 25 °C; I = 0.6 M.113

Catalyst Mn2+ Mn(bipy)2+ Co2+ Co(bipy)2+ Co(phen)2+ Ni2+ Ni(bipy)2+ Ni(phen)2+ Cu2+ Cu(bipy)2+ Cu(phen)2+ Zn2+ Zn(bipy)2+ Zn(phen)2+

K/M−1

105k/s−1

104Kk/M−1 s−1b

8.1 15 20 20 25 25 20 20 150 360 270 10 8 6

2.0 7.0 6.7 15.5 20.0 9.7 13.2 16.3 200 267 170 8.5 26.7 51.7

1.6 10.5 13.4 31.0 50.0 24.2 26.4 36.6 3000 9612 4590 8.5 21.3 31.0

a

HO2 CCH2COCH2CO2H   MeCOCH2 CO2H.  CO2 Kk = kcat.

b

to a higher effective charge on the metal. The greater enhancement of Mn(ii) than of Ni(ii) could then be accounted for in terms of the smaller nuclear charge of Mn(ii), which allows greater (more favourable) back-­donation. It also might be that there is a finite amount of π-­bonding from any coordinated oxygen base to metal. As the aqua ligands are substituted by phen, π-­bonding (L→M) from the coordinated oxygens of the substrate increases, thereby polarizing the C–C bond in

(as in the case of oxaloacetate), facilitating decarboxylation. Whatever the explanation might be, the results are most suggestive of a type of binding between metal ions and the enzyme, similar to that between metal ions and phen or bipy, perhaps through appropriately located imidazole rings in the enzyme molecule. A much lowered (microscopic) dielectric constant near the binding site may be an additional factor enhancing the catalytic process as a result of stronger association with the substrate. Decarboxylation of a number of β-­oxodicarboxylic acids, such as oxaloacetic acid (HO2CCOCH2CO2H) and oxalosuccinic acid [HO2CCOCH(CO2H) CH2CO2H], is known to be catalysed by enzymes activated by a variety of M2+ and M3+; Mn2+ is five times more effective than Ni2+.114

Activation of Molecules by Coordination and Reactivity of Coordinated Ligands

467

Scheme 8.21

8.3.8  Wacker Process115 This was first reported by Smidt et al.115a and was exploited commercially for the production of acetaldehyde by Wacker Chemie AG in Germany.115b Basically three sets of reactions as shown in Scheme 8.21 are involved in the process [eqn (8.45)].   

  

  

1 PdCl24  CuCl2 catalyst CH CH2  O2    MeCHO 2 2

(8.45a)

d[CH3CHO] [CH2  CH2 ][PdCl24 ] k dt [H ][Cl  ]2

(8.45b)

The mechanism of the first step set of changes has attracted much attention.116 Kinetic studies in aqueous solution led to the rate law in eqn (8.45b), which shows a first-­order dependence on CH2=CH2 and PdCl42−, an inverse dependence on H+ and an inverse square dependence on Cl−. The rate is slowed when D2O is used as the solvent, but no deuterium appears in the product. This indicates that all four H atoms of the CH3CHO originated from the CH2=CH2. Only a very small isotope effect is seen when C2D2 is substituted for CH2=CH2, suggesting that no C–H bonds are broken in the rate-­determining step. On the basis of this information and other experimental observations, several proposals for the sequence of changes shown in Scheme 8.22 seem appropriate. The transformation c→d in this scheme involves nucleophilic attack and migration of OH− to the µ-­bonded CH2=CH2. In a similar manner, propylene (MeCH=CH2) forms acetone.

8.4  Insertion Reactions117 These are non-­oxidative addition reactions in which a molecule inserts into a metal‒ligand bond, thereby transforming the ligand into another ligand. Insertions into M–C and M–H bonds are especially common. Both 1,1 [eqn (8.46)] and 1,2 [eqn (8.47)] insertion reactions are possible, whereas 1,3 [eqn (8.48)] and 1,4 [eqn (8.49)] insertions are rare.   

 (8.46)   

M–L + X–Y → M–X–Y–L  (e.g. M–H + F2C=CF2 → M–CF2–CF2H)   

(8.47)

Chapter 8

468

Scheme 8.22

M–H + N2CHR → M–N=N–CH2R

(8.48)

M–H + H2C=CH–CH=CH2 → M–CH2–CH=CH–CH3

(8.49)

  

  

The term “insertion” describes only the result of the reaction, it has no mechanistic significance. Many insertion reactions are reversible, and the reverse reaction is called extrusion or elimination. Table 8.5 shows representative insertion reactions involving transition metals. For thermodynamic reasons, insertion of CO generally takes place into M–R and not into M–H bonds. In contrast, alkene insertion is common for M–H and much less common for M–R bonds. Polymerization of an alkene involves repeated alkene insertion into an M–R bond. Thermodynamics favours the reaction with M–R and hence its comparatively rare occurrence must be due to kinetic factors being unfavourable. Comparison of the barriers of insertion of CH2=CH2 into the M–R bond in [(Cp*){P(OMe)3}MR(CH2=CH2)] (where R = H or Et, M = Rh or Co) showed118 that the reaction involving M‒H has a 25‒42 kJ mol−1 lower barrier. This corresponds to a migratory aptitude ratio kH/kR of 106‒108. Reactions involving M–H bonds are almost always kinetically more facile.

Activation of Molecules by Coordination and Reactivity of Coordinated Ligands

469

Table 8.5  Some  representative common insertion reactions. Inserted molecule

Bond

Insertion product

CO

M–CR3 M–OH M–NR2 M–H M–R M–NR2 M–OH, M–OR M–M M–H M–H M–H M–H M–R M–M M–H M–R′ M–η3-­C3H5 M–H M–H M–R M–M M–η3-­C3H5 M–Me M–H M–CR3 M–CR3 M–Cl

MC(O)CR3 MC(O)OH MC(O)NR2 MO2CH MC(O)OR MOC(O)NR2 MOCO2H(R) MSC(S)M MS2CH and MSC(S)H M–C2H5 M–CF2CF2H MC(R)=CH(R′) cis or trans M(η3-­allyl) MSn(Cl)2M MCH=NR MCR'=NR MC(=NR)(CH2CH=CH2) M(RNCHS) M(RNCHNR) MS(R(O)2 or MOS(OR) MOS(O)M MS(O)2CH2CH=CH2 MOS(O)2Me MOOH MOOCR3, MOCR3 MSCR3 MCH2Cl

CO2

CS2 CH2=CH2 CF2=CF2 RC≡CR' CH2=C=CH2 SnCl2 RNC RNCS RN=CNR SO2 SO3 O2 S8 CH2N2

Although, in principle, insertion reactions are reversible, for thermodynamic reasons generally reaction in only one of the two possible directions is observed. Thus, SO2 inserts into M–R bonds to give alkyl sulfinate complexes [eqn (8.50)], which rarely eliminate SO2. Conversely, diazoarene complexes readily eliminate N2 [eqn (8.51)] but N2 insertion into an M–aryl bond has not been observed so far.   



M–R + SO2 → M–SO2R

(8.50)

M–N=N–Ar → M–Ar + N2

(8.51)

  

  

One way to view insertion reactions is to consider that the ligand X migrates with its M–X bonding electrons to attack the antibonding π-­orbital of the A = B ligand. In this intramolecular attack, nucleophilic attack on A = B, the migrating group R retains its stereochemistry.

Chapter 8

470

8.4.1  Insertion of CO Carbon monoxide shows a strong tendency to insert into metal‒alkyl bonds, forming metal‒acyl compounds. Most of these reactions proceed according to eqn (8.52) and in general follow the mechanism shown in eqn (8.53).119   





(8.52)

  

 (8.53)   

The rate law for the mechanistic scheme in eqn (8.53) is   

Rate = ‒d[substrate]/dt = (k1k2[L]/k−1 + k2[L])[substrate]   

(8.54)

There are three possible situations, all of which have been observed in real cases, as follows:    (a) If k−1 > k2[L], eqn (8.53) reduces to Rate = (k1k2/k–1)[L][substrate].    In this case, the intermediate almost always goes back to the substrate and the overall rate is governed by the second step (attack by L) and the overall kinetics are those of a second-­order process.    (c) If k–1 ≈ k2[L], then the situation is much more complicated and eqn (8.53) can be rearranged to the following form:   Rate = kobs[substrate] (where kobs = k1k2[L]/(k–1 + k2[L]).

  



  

In this case, the intermediate is trapped by L, at a rate that is comparable to that of the reverse migration. The expression in (b) can be rearranged to the form in eqn (8.55) and by plotting 1/kobs versus 1/[L] we can evaluate k1 from the intercept (= 1/k1) and then the ratio k–1/k2 from the slope (= k–1/k1k2). 1/kobs = (k–1/k1k2[L]) + 1/k1

(8.55)

When the incoming ligand is labelled CO (13CO), the product contains only one labelled CO, which is cis to the newly formed MeCO group. This shows

Activation of Molecules by Coordination and Reactivity of Coordinated Ligands

471

that the Me group migrates to a CO in the cis position in the complex, rather than free CO attacking the Mn–Me bond. Any stereochemistry at the alkyl C is retained in the insertion, which is consistent with the mechanistic scheme shown in eqn (8.52). The labelled CO can be located in the product by NMR and IR spectroscopy. The substitution process of the reaction [eqn (8.56)] has features that are more complex than in a conventional ligand replacement process. There is evidence that such reactions occur through an acyl intermediate having a vacant coordination site [eqn (8.53)] or a weakly bonded solvent molecule (S) as shown in 7.   



[MeMn(CO)5] + PPh3 → [MeMn(CO)4(PPh3)] + CO

(8.56)

  

This arises from cleavage of the Mn–Me bond and its replacement by an MeC(O)–Mn bond. The substitution reaction is completed by binding of an entering ligand in the vacant coordination site or by replacement of the weakly bonded S by a more strongly binding ligand [such as PPh3 in eqn (8.52)] and restoration of the Mn–Me bond with loss of CO from the M–C(=O)–Me. When this entering ligand is CO, the net reaction is “insertion” of CO into the Mn–Me bond [eqn (8.57)].   





(8.57)

  

Based on experimental results, the most attractive hypothesis for acyl intermediate formation is methyl migration in which the Me group engages in nucleophilic attack (indicated by -­-­-­-­in 8) on the C atom of a neighbouring CO ligand (which is cis to the Me), for which evidence is available.121

Chapter 8

472 ‡

This interpretation is consistent with a strongly negative ∆S value (−21.1 cal K−1 mol−1) for the reaction; this negative value indicates incorporation of an additional ligand into the activated complex. Moreover, electron-­ withdrawing groups substituted into CH3 dramatically slow the reaction by reducing the nucleophilicity of the CH3. Thus, [(O2NCH2)Mn(CO)5] is much less reactive than [MeMn(CO)5)]. By studying the reverse of the reaction shown in eqn (8.57), i.e. α-­ elimination of CO from the 13C-­labelled acyl complex [(Me13CO)Mn(CO)5], it has been shown that the labelled C ends up in a CO cis to the Me (Scheme 8.23). Incidentally, the 13C-­labelled acyl complex can be prepared easily by reacting [Mn(CO)5]− with Me13C(O)Cl. It is a general strategy to study a reverse reaction to obtain information about the forward reaction, since forward and reverse reactions of a thermal process must follow the same path according to the microscopic reversibility principle. In this particular example, the labelled C of the acyl group ends up as CO cis to Me, hence the CO group to which the Me migrates in the forward reaction must also be cis to the Me. The situation is favourable in this case as subsequent scrambling of the COs is not fast. Hence we know that the Me and CO must be mutually cis for the insertion process, but it is still not certain if the migration is of Me to CO or vice versa. It is possible to use reversibility arguments to show that it is the Me, and not CO, that migrates. This has been inferred by studying the elimination of labelled CO that is cis to the acyl, MeCO, group in [(MeCO)Mn(13CO)(CO)4]. If the CO of the acyl group migrates during elimination, then the product will have the Me in the same position as that of MeCO in the starting complex and hence remain cis to the labelled CO, and no trans product can be formed. If, however, the Me migrates, then it will end up in both cis and trans positions of the labelled CO, and this has actually been observed, as is evident from Scheme 8.24. Path c in Scheme 8.24 involves the loss of labelled *CO followed by migration of Me or CO of Me–CO, which will lead to the identical product MeMn(CO)5. The actual distribution of products is shown in Scheme 8.25, indicating that some reaction involves loss of labelled CO also.

Scheme 8.23

Activation of Molecules by Coordination and Reactivity of Coordinated Ligands

473

Scheme 8.24

Scheme 8.25

The observation mentioned above implies that the reaction involves the loss of a CO (including the *CO) from a cis position to the COMe followed by migration of the Me of COMe to the position vacated the loss of CO. As there are two CO groups that are cis to both *CO and COMe, the cis isomer (having Me cis to *CO) constitutes 50% of the product. Hence in the insertion reaction also the Me migrates to a CO in a cis position and not the CO.121 The labelled complex for the above study was prepared by a photolytic process and characterized by its IR spectrum. NMR spectroscopy is also useful for locating the labelled CO. The conclusion drawn for this Mn complex is not necessarily true for all other similar systems, since an acyl group has occasionally been found at the site originally occupied by the alkyl group.

474

Chapter 8

The product distribution follows from rational considerations. The four CO groups adjacent (i.e. in a cis position) to MeCO are equally likely to be lost. If the CO trans to the labelled *CO is lost, then Me migration results in the formation of a complex in which 13CO and Me are trans to each other; loss of a CO that is cis to both 13CO and MeCO followed by migration of Me results in a product complex in which 13CO and Me are mutually cis. Finally, loss of 13 CO results in the formation of a product complex having no 13CO. These considerations suggest that in the case of Me migration the products should be in the ratio 1 : 2 : 1, as observed (Scheme 8.25). Hydroformylation (oxo process)4b,122 is used on an industrial scale for the production of aldehydes from which other useful products (mostly alcohols) are made. Basically, the reaction is as shown in eqn (8.58), yielding n-­ and iso-­products in a ca. 3 : 1 ratio. The reaction presumably occurs according to Scheme 8.26.   



 (8.58)

  

The steps involve migration and insertion of a cobalt-­bound H into the Co–alkene bond, then migration and insertion of a cobalt-­bound Co into the Co–C2H5 bond that results from the previous migration‒insertion and rearrangement step, and finally insertion of H from Co into the Co–CO bond of the Co–CO(alkyl) and elimination of alkyl‒CHO. In addition to the formation

Scheme 8.26  R  = H, alkyl.

Activation of Molecules by Coordination and Reactivity of Coordinated Ligands

475

of a mixture of straight-­(n-­) and branched-­(iso-­) chain products, some loss of alkene by hydrogenation to RCH2CH3 and of RCHO to RCH2OH also occurs. The rhodium carbonyl cluster [Rh6(CO)16] behaves similarly. In 1968, Wilkinson and co-­workers123 discovered that some rhodium catalysts are very effective in the homogeneous hydroformylation of alkenes. One such effective catalyst is [RhH(CO)(PPh3)3]. The presence of the bulky triphenylphosphine greatly favours anti-­Markovnikov addition124 (Scheme 8.27), leading to higher yields of straight-­chain aldehydes of greater use. [Rh6(CO)16] produces a mixture of straight-­and branched-­chain aldehydes, and Co2(CO)8 gives similar results, with considerable hydrogenation of aldehydes to alcohols. The branched-­chain products are formed by Markovnikov addition.124 The use of [RhH(CO)(PPh3)3] in the presence of excess PPh3 leads to a process that could be operated below 100 °C and a pressure of a few atmospheres with no loss of alkene and of the aldehyde through hydrogenation, and giving a yield of the straight-­chain aldehyde of over 95%. Even with [Rh6(CO)16] the yield of the straight-­chain aldehyde is ca. 40%. Thus, at 50 °C and a pressure of 300 psi using CO and H2 in a 1 : 1 ratio, the yields of aldehydes formed from 1-­hexene in a 22 h reaction time using [Rh6(CO)16] as a catalyst were heptanal 43.6%, 2-­methylhexanal 41.1% and 1-­ethylpentanal 12.3%. The possible mechanisms of the hydroformylation process, and also of isomerization and hydrogenation, using such catalysts have been discussed.124 In hydroformylation, the first step is the formation of a conventional alkene addition compound, its transformation into a metal‒alkyl

Scheme 8.27  Schematic  representation of the hydroformylation of 1-­alkenes using [RhH(CO)L3] (where L = PPh3) as catalyst.

Chapter 8

476

Scheme 8.28 species by transfer of metal-­bound hydrogen to the double bond and rearrangement, followed by insertion of CO between the M–alkyl bond and then addition of hydrogen (from M–H) to the M–CO‒alkyl at the M–CO bond and splitting off of the alkyl‒CHO (Scheme 8.26). Insertion of CO into CH3OH in the presence of a rhodium catalyst, [Rh(CO)2I2]−, with CH3I as a co-­catalyst is an important method (Monsanto process) for producing acetic acid on a commercial scale. The process (Scheme 8.28) operates at 150‒200 °C and 30‒60 bar pressure.125 Similar reactions occur in the carbonylation of other alcohols and of C2H4; in the latter case an intermediate, [Rh{C(O)Et}(CO)I3]−, is formed [eqn (8.59)].126   



  

HI  C H  EtC  O  Rh  CO  I3  Rh  CO 2 I2   2

4





(8.59)

The similar Ir-­catalysed MeOH → MeCO2H process has certain advantages and is the basis of the Cativa process. Cobalt-­and nickel-­catalysed processes for the carbonylation of methanol to acetic acid have also been studied. The active catalysts in these processes are [Rh(CO)2I2]−, [CoH(CO)4] and [Ni(CO)4].124,127a Another useful catalyst for the conversion of methanol to acetic acid is [RhCl(CO)2(Ph2PCH2CH2P(O)Ph2)].127b [Ni(CO)4] and [Co2(CO)8] are good catalysts for the following type of reaction:   



catalyst ROH  CO  C2 H4   CH2  CHCO2 R

(8.60)

  

As with [Rh(CO)2I2]− catalyst, the steps involve oxidative addition (a), insertion (b) and reductive elimination (c), with the rhodium shuttling between +l and +3 oxidation states. The latest developments in the carbonylation of methanol to acetic acid have been reviewed.128

8.4.2  Insertion of Sulfur Dioxide Coordination of SO2 is sometimes followed by intramolecular insertion of SO2 into an M–C σ-­bond of an M–alkyl/aryl compound in a cis position [eqns (8.61) and (8.62)].129   

trans-­[PtPh(Cl)(PEt3)2] + SO2 → PtPh(Cl)(PEt3)2(SO2)] → trans-­[Pt(SO2Ph)(Cl) (PEt3)2] (8.61)   

AuMe3(PMe3)+SO2 → AuMe3(PMe3)(SO2) → AuMe2(SO2Me)(PMe3)  (both cis and trans isomers) (8.62)   

Activation of Molecules by Coordination and Reactivity of Coordinated Ligands

477

Intermolecular insertion of SO2 is well known, and it may or may not involve prior formation of an SO2 adduct as a transient intermediate. The overall process is shown in eqn (8.63), where R is an alkyl or aryl group or a similar σ-­bonded carbon group and the product M–SO2R can adopt four different types of structures as shown in Figure 8.5.   



M–R + SO2 → M–SO2R

(8.63)  

However, the S-­sulfinate is commonly formed as a stable product.130 Thus, on refluxing [CpFeMe(CO)2] in liquid SO2, [CpFe(SO2Me)(CO)2] (S-­sulfinate) is formed, but O-­sulfinates have been isolated for some metals. The other isomers are also known, which (formed initially) may also be transformed to the most stable S-­sulfinate.131 However, for substrates that conform to the 18e rule, the path through O,O′-­sulfinate formation may be ruled out since the formation of two M–O bonds would lead to violation of the 18e rule in the O,O′-­sulfinate.132 The S-­sulfinates have νS‒O in the 1250–1000 and 1100‒1000 cm−1 regions.130 However, the O,O′-­sulfinates and O-­sulfinates are difficult to distinguish by νS‒O as they appear in the ranges 1085–1050 and 1000‒820 cm−1 or lower. In contrast to CO (mentioned in the preceding section), SO2 inserts directly into M–C bonds. Kinetic studies have shown that SO2 behaves as a Lewis acid attacking (SE2) the alkyl ligand rather than the metal (Scheme 8.29), so that more electron-­donating alkyl groups react faster. As indicated in Scheme 8.29, backside attack leads to inversion130 of configuration at C.

Figure 8.5  Sulfinates  formed by SO2 insertion.

Scheme 8.29  Mechanism  of SO2 insertion.130

Chapter 8

478

Also, in contrast to CO, extrusion of SO2 occurs much less readily and hence its insertion is seldom reversible.

8.4.3  Insertion of Carbon Dioxide Carbon dioxide undergoes a variety of insertion reactions (Table 8.6) fairly easily; the reaction possibly occurs via initial complexation forming an M– CO2 linkage.133 Insertions into M–H, M–R, M–OH, M–OR, M–O and M–NR2 are well known.134 Reaction of CO2 (COS and CS2 are similar) with [W(CO)5(OH)]− forms the bicarbonato complex [W(CO)5{OC(O)OH}]− (COS and CS2 form the mono-­ and bis-­thiocarbonato complexes, respectively).133 However, in the cases of complexes of Ti, Zr, etc., which have a high affinity for oxygen (i.e. oxophilic), deoxygenation of CO2 to CO takes place, and in some other cases a disproportionation reaction takes place [eqn (8.64)].

  



Ru  CO 4 

2

 2 CO2 

  

1 2 Ru2  CO 8   CO  CO32  2

(8.64)

Carbon dioxide insertion into an M–H bond is probably involved in the catalytic reduction of CO2 with H2 forming HCO2H (formic acid). Although this is an uphill task thermodynamically (∆G = +33.5 kJ mol−1), the reaction Table 8.6  Insertion  of CO2 across different substrates. Substrate undergoing CO2 insertion

Product(s) of insertion

M–H

M–O–C(O)H and

M–R

M–O–C(O)R and

M–OH M–OR M–O–M

M–O–C(O)–OH M–O–C(O)–OR

M–NR2

M–O–C(O)NR2 and

Activation of Molecules by Coordination and Reactivity of Coordinated Ligands

479

becomes favourable under pressure and in the presence of a base to deprotonate the formic acid formed. An efficient catalyst reported is [Cp*IrCl(AA)]+ [AA = 4,4′-­bis(hydroxo)bipyridine], which gives over 40 000 turnovers per hour at 120 °C and a pressure of 60 atm.135 Reaction of an aqua complex with CO2 to form a bicarbonate complex also apparently involves insertion of CO2 across an M–OH bond. By isotope labelling of oxygen of the aqua ligand, it was demonstrated that no breaking of the M–O bond occurs during the reaction [eqn (8.65)].   



[(H3N)5Co–*OH2]3+ + CO2 → [(H3N)5Co–*O–C(=O)OH]2+ + H+ (8.65)

  

The reaction involves insertion into the O–H bond (not the Co–O bond) by electrophilic attack on the O of M–OH2 by the C of CO2. Reactions of all M–OH (and also M–OR) may be similar. The dinuclear Cu(ii) complex [Cu{HB(3,5-­ i Pr2pz)3}2(µ-­OH)2] reacts with CO2 to form the bridged carbonato complex 9, which has potential for the fixation of CO2.136

Insertion of CO2 into an M–C bond is also well known; the reaction with organolithium compounds occurs vigorously [eqn (8.66)]. Similarly, Grignard reagents also undergo CO2 insertion reactions. CO2 insertion reactions with some other MR and MR2 also occur readily. Thus, BeR2 forms Be(O2CR)2, but ZnR2 are less reactive [eqn (8.67)]. CdEt2 and HgEt2 are even less reactive than ZnEt2.   

RLi + CO2 → R–C(=O)–OLi   



  

150 C   Zn  O2 CEt 2  Et 2 Zn  2CO2 

(8.66)

(8.67)

Insertions into M–NR2 bonds (also B–N and Si–N bonds) are well known [e.g. eqns (8.68) and (8.69)].   



25° C PhB  NHEt 2  CO2   PhBOC  O  NHEt 2

(8.68)



25° C in NHEt 2 Me3Si  NEt 2   CO2    Me3SiOC  O   NEt 2 

(8.69)

  

  

Like CO2, cumulenes (X=C=Y), such as C3O2, PhNCO, PhNCS and Ph2C=C=O, can also insert across M–X bonds [eqns (8.70) and (8.71)]. These reactions are facilitated if X is a good donor; thus, for a series of complexes of W–X

Chapter 8

480

(X = amide, alkoxide or alkyl) the reactivity of RNCS was found to decrease in the order W–N > W–O > W–C.137   





(8.70)

  





(8.71)

  

8.4.4  Insertion of Carbon Disulfide Insertion of CS2 is known for all the elements that undergo CO2 insertion. Like CO2, CS2 undergoes insertion reactions with M–H, M–R, etc., bonds [eqn (8.72)].138a Other insertions include into M–N bonds [M = Sb(iii), Zr(iv), Nb(v), Ta(v), etc.]. Reaction of Au(iii) chloride with CS2 leads to a novel insertion into an Au–Cl bond, forming an orange complex, [AuCl2(η2-­S2CCl)].138b   



M–H + CS2 → MS2CH or MSC(S)H

(8.72)

  

Insertion of CS2 into the Ru–H bond in [RuH(Cl)(CO)(PPh3)2 (4-­vinylpyridine)] leads to the formation of [RuCl(η2-­S2CH)(CO)(PPh3)2]·THF in THF medium as an orange complex having a dithioformate ligand.139 Reaction of a metallathioborane of Rh with CS2 under reflux yielded ca. 37% of a product having the HCS2 as a bridging ligand between Rh and B of the cluster.140

8.4.5  Insertion of Olefins (Alkenes) Olefins can insert into M–C and M–H bonds. When the metal has sufficient d electrons for back-­bonding, the olefin is often coordinated before insertion. This may require ligand dissociation to vacate a coordination position [eqns (8.73) and (8.74)].   



Cp2NbVH3 + H2C=CH2 → [Cp2NbIII(H)(η2−H2C=CH2)] + H2

(8.73)

  





(8.74)

  

Insertion converts a hydrido‒ or alkyl–olefin complex into an alkyl complex by addition of H or C to the β-­carbon of the olefin. Note that the

Activation of Molecules by Coordination and Reactivity of Coordinated Ligands

481

oxidation state of the metal does not change. This insertion is related to attack on coordinated olefins by nucleophiles, except that here the hydride or alkyl nucleophile is pre-­coordinated to the metal. In contrast to CO, olefins often undergo multiple insertions, providing a means for their polymerization. However, the final product in the above reaction resists further insertion. The reverse of olefin insertion into an M–H bond is the β-­elimination reaction of metal‒alkyl complexes, which can occur only if a vacant orbital and a vacant coordination position are available and β-­H is present (Scheme 8.30).‡ If these conditions are fulfilled, a facile route exists for the decomposition of metal‒alkyl complexes. The initial product is a hydrido–olefin complex, which may then dissociate, forming an olefin. A study141 of the thermal decomposition of [CpFe(n-­C4H9)(CO)(PPh3)] revealed that the reaction was severely retarded by the addition of PPh3. This suggests that dissociation of PPh3 to create a vacant coordination position of Fe is a necessary intermediate step in the decomposition via β-­elimination for this 18e complex. The overall process is shown in eqn (8.75).   





(8.75)

  

Fundamental features of olefin insertion were shown in an investigation of some scandium complexes [eqn (8.76)].142 [Cp*2ScR] is a coordinatively

Scheme 8.30  Olefin  insertion and the reverse process, β-­elimination.

‡

α-­Elimination is also known, as in the following case:

The abs­ence of β-­H and the presence of bulky ligands apparently favour intramolecular abstraction of an α-­hydrogen and elimination of CMe4 from an M(CH2CMe3)4Cl or M(CH2CMe3)5 intermediate. However, these are much rarer than β-­elimination and are most often seen with early transition metals. α-­Eliminations are promoted by steric crowding around the metal, addition of phosphines (which may abstract the α-­H or increase steric crowding) and lack of β-­H.

Chapter 8

482 0

unsaturated d , 14e compound stabilized by bulky electron-­donating Cp* ligands. Such compounds represent ideal models for studying insertion and β-­elimination reactions because no complications arise from the need to dissociate other ligands to provide vacant coordination sites, or from the kinetics of olefin coordination because this does not happen for d0 complexes, which are incapable of back donation.   



Cp*2SCR + H2C=CH2 → Cp*2ScCH2CH2R  (R = H, Me, Et, npr) (8.76)

  

The rate of ethylene insertion into [Cp*2ScR] decreases in the order R = H >> CH2(CH2)nCH3 (n > 2) > nPr > Me > Et. This order can be understood in terms of an acyclic, planar, four-­centre transition state involving incipient formation of a bond between the metal and an olefinic carbon. This creates a partial positive charge on the β-­carbon, and H or alkyl C migrates as an anion to the β-­carbon. Metal and R add in a cis fashion to the olefin [eqn (8.77)].

  





(8.77)

  

H inserts fastest because its non-­directional s orbital can overlap more effectively with the β-­C orbital than can alkyl sp3 hybrids. Among alkyls, the rate decreases with increasing strength of the M–alkyl bond, except for R = Et. The ground state of [Cp*2ScEt] is apparently stabilized by agostic interaction of the unsaturated metal with the electrons in a C–H bond, which is impossible for larger alkyls owing to steric hindrance of the Cp* rings and for Me because of its size and geometry. The rates of β-­H elimination from a series of substituted complexes [Cp*2Sc(CH2CH2C6H4X-­4)] showed that the rate decreased with increase in the electron-­withdrawing power of X, indicating the development of positive charge on the β-­carbon in the transition state. Polymerization of ethylene catalysed heterogeneously by TiIII salts and aluminium alkyls (Ziegler‒Natta catalysts)143 also involves olefin insertion.

8.4.6  Olefin (Alkene) Polymerization144 Breslow and Newburg145a reported that a fresh mixture of Cp2TiCl2 and Et2AlCl is a highly active homogeneous catalyst for the polymerization of ethylene, being as active as the usual Ziegler-­t ype heterogeneous catalysts,145b obtained from TiCl4 and Al(alkyl)3, but better than the latter in some respects. During 1958–1960, much valuable work with such catalysts was carried out by Natta and co-­workers. From studies made with such homogeneous catalysts, it was

Activation of Molecules by Coordination and Reactivity of Coordinated Ligands

483

Scheme 8.31  The  process continues until suitably terminated. possible to elucidate the mechanism of the catalytic action of all such bimetallic catalyst systems, including the Zeigler-­t ype heterogeneous catalysts. The mechanism of the polymerization (Scheme 8.31) is believed to involve insertion of an alkene into an M–alkyl bond as shown for Ti–Et generated by the action of AlEt3 on TiCl4 (in the usual Ziegler‒Natta type of catalyst) or of Et2AlCl on Cp2TiCl2. Metallocene complexes of Ti, Zr and Hf are useful catalysts for alkene polymerization reactions,146 and so also are lanthanide metallocenes.147 The mechanism shown in Scheme 8.31 is the Cossee‒Arlman mechanism.148 It involves η2 attachment of ethylene at a vacant site on a Ti atom on the surface of the catalyst, followed by insertion of the ethylene between Ti and C in the Ti–Et bond. This extends the carbon chain from two to four atoms, leaving a vacant site on Ti. The process is repeated and the carbon chain grows in length until this process is terminated by a hydride transfer to the metal [eqn (8.78)]. A similar reaction occurs with other alkenes RCH=CH2 such as propylene (CH3CH=CH2). The steric hindrance inherent in the catalyst surface coordination site ensures the formation of stereoregular polymers, since when an alkene RCH=CH2 binds to the surface coordination site the group R, which is bulkier than H, always points away from the catalyst surface and hence when the molecule migrates and is inserted into the Ti–C bond it always has the same orientation. This is called cis-­insertion of the alkene, and explains why the polymers produced are stereoregular (isotactic polymer). In the catalysed polymerization, the chain will continue to grow until termination occurs by a reverse of insertion reaction involving hydride transfer to the metal. The hydride, of course, becomes replaced by, say, Cl−, regenerating the catalyst.   

 (8.78)   

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484

8.4.7  Acetylene (Alkyne) Polymerization A breakthrough in this field was the report that Ni compounds are efficient catalysts for the cyclotrimerization of acetylene to benzene.149 As a result of subsequent developments, it is now known that alkynes can be selectively dimerized, cyclotrimerized and even oligomerized (forming higher polymers) by a variety of d-­block metal compounds and their complexes and also by complex compounds of Group 3 metals (Sc, Y, including lanthanides such as La and Ce). Thus, [Cp*RuH3(L)] (L = PMe3, PCy3, PPh3),150 an Hf–carboranyl complex,151 some tris-­pyrazolylborate complexes of Ru152 and [Cp*2M (SiMe3)2] (M = Y, La, Ce)153 are efficient in the dimerization of terminal alkynes, RC≡CH. In such dimerizations, open-­chain isomeric products (a), (b) and (c) are formed (Figure 8.6). NMR studies indicated153a that M–acetylides [Cp*2M‒C≡CR]n, generated by σ-­bond metathesis of the precursor complex with the CH bond of the alkyne, are the active species in the catalytic cycle. The postulated reaction mechanism is shown in Scheme 8.32. A similar mechanism has been proposed for

Figure 8.6  Open-­  chain products formed in the dimerization of terminal alkynes

(RC≡CH): (a) and (b) are cis-­and trans-­isomers, respectively, of disubstituted 1-­buten-­3-­ynes and (c) is 2,4-­disubstituted 1-­buten-­3-­ynes.

Scheme 8.32

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485

the catalytic open-­chain dimerization and trimerization of RC≡CH by a zirconocene catalyst,153b where the active species is Cp*2Zr–C≡CR. Cyclotrimerization of alkynes to arenes is catalysed by NbCl5, TaCl5, MoCl6, WCl6, [CpCo(η4-­C6Me6)] and [CpRh(cod)],154 and also by Mo2Cl6(tht)3 (tht = tetrahydrothiophene).155 The dinuclear Mo compound was found to be very active in the cyclotrimerization (and also oligomerization) of a variety of alkynes. Mononuclear MoCl3(tht)3 shows similar behaviour, leading to the postulate that the active species in the case of the dinuclear complex is also a mononuclear species generated from the dinuclear species, viz. MoCl3 (tht)2(L); species with L=PhC≡CR and R = Me, Et have been isolated.155 The complex with R = Et cyclotrimerizes PhC≡CH and PhCMe, but polymerizes EtC≡CMe. The trimerization occurs in a stepwise manner through intermediates. The mechanism proposed for the cyclotrimerization of HC≡CH by a fac-­[Ir(triphos)]+ catalyst is shown in Scheme 8.33.156 Some of the intermediates were isolated and characterized by X-­ray diffraction analysis. The dinuclear Hf‒dicarbollide‒methyl complex [(Cp*)(η5-­C2B9H11) Hf(µ-­η2 : η3-­C2B9H11)Hf(Cp*)Me2] catalyses the regioselective dimerization of terminal alkynes RC≡CH (R = Me, nPr, tBu) to 2,4-­disubstituted 1-­buten-­3-­ynes selectively, rather than to trimers and higher oligomers due to a “self-­ correcting” mechanism.151 Lanthanide and Group 3 carbyls [(Cp*)2LnCH(SiMe3)2] (Ln = Y, La, Ce) are active catalyst precursors for the oligomerization of terminal alkynes RC≡CH (R = alkyl, aryl, SiMe3).153 Regioselectivity and the extent of oligomerization depend strongly on the Ln used and also the alkyne substituent R. Use of the Y complex leads to selective dimerization to 2,4-­disubstituted 1-­buten-­3-­ynes, but mixtures of 1,4-­and 2,4-­disubstituted 1-­buten-­3-­ynes are formed when R = Ph and SiMe3. Reactions using the [Cp*2M(SiMe3)2] complexes produce, in addition to dimers, higher oligomers (trimers, tetramers) of various types (allenes and dienes). Cp*2Sc‒R (R = H, Me, n-­C5H11) are effective in the

Scheme 8.33

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Scheme 8.34 polymerization of alkynes through generation of Cp*2Sc‒C≡CH, which is the active species.157 The proposed catalytic cycle for the dimerization of HC≡CH is shown in Scheme 8.34. Rhodium(i) tris(pyrazolyl)borate complexes (TpR2)Rh(cod) (R = Me, iPr, Ph) serve as efficient catalysts for the highly stereoregular polymerization of phenylacetylene derivatives, p-­YC6H4(C≡CH) [Y = H, Me, Cl, CN, NO2, MeCO2, C(O)Me], to form polyphenylacetylenes having a head-­to-­tail, cis-­ transoidal structure. The catalyst activity is strongly affected by the substituents R in the 3-­and 5-­positions of the pyrazolyl fragment; the more sterically demanding R groups lead to higher catalytic activity.158 In the polymerization of PhC≡CH with an Rh(i) pyrazolate catalyst the products obtained159 are PhC≡C–CH=CHPh (52%), PhC≡C–C(Ph)=CH2 (31%), 1,2,5-­C6H3Ph3 (14%) and 1,3,5-­C6H3Ph3 (3%). The complexes Cp*2M–R (M = Sc,160 Y161) catalyse the head-­to-­tail dimerization of propyne to 2-­methyl-­1-­penten-­3-­yne. A nickelocene-­based alkyne polymerization catalyst is prepared by reacting NiCp2 with LiR.162 The polymerization mechanisms shown in Schemes 8.31‒8.33 have been proposed in investigations with some specific catalysts, but these are presumably applicable in cases of other metal-­based catalysts for alkyne polymerization. RuH2(PR3)4 (R = Et or nBu) are active catalysts for the co-­dimerization of alkynes with 1,3-­butadiene forming trans-­RC≡C–CH=CH–CH2Me (R = nPr, n Bu, tBu, n-­C6H13, Ph).163 Alkynes can be co-­trimerized with nitriles to form pyridines; the reaction is catalysed by [CpCo(H2C=CH2)2] and other such labile complexes.154a The catalyst system Cp2ZrEt2 + NiCl2(PPh3)2 has been used for the co-­trimerization of two different alkynes with PhCN, forming pyridine derivatives.154b

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In some alkyne polymerizations, the isolable product contains the polymeric alkyne bonded to the catalysing metal as a ligand. Thus, in the reaction of Fe(CO)5 with PhC≡CPh eight different products are formed,164 one of which is (η4-­cyclo-­C4Ph4)Fe(CO)3. In the reaction of ZrCl4 with an alkyne, (η6-­arene)Zr(iv) is formed.165 Polymerizations of alkynes with d-­block metal-­ based catalysts are of use in organic syntheses.166

8.4.8  A  symmetric Synthesis Catalysed by Coordination Compounds167 There has been considerable interest in the stereospecific synthesis of organic compounds using optically active coordination compounds. Chiral catalysts are needed for the production of drugs, pesticides, pheromones and fragrances that are chiral but where only one of the enantiomers is the species with desired properties. A significant advance was the discovery that asymmetric epoxidation can be achieved with complexes of Mn(iii) containing chiral chelating ligands, as in the following example where alkenes are converted to epoxides [eqn (8.79)] [note that Mn(iii) has a square-­pyramidal geometry with the Cl in an axial position].   



(8.79)

  

Some cis-­substituted alkenes can be converted to chiral epoxides having more than 90% of an enantiomer by this process.168 A chiral and water-­ soluble Mn(iii) complex of a porphyrin derivative was prepared that catalyses enantioselective epoxidation and C–H hydroxylation.169 Hydrogenation of double bonds using an Rh(i) diphosphine complex as catalyst may be carried out stereospecifically using a chiral diphosphine such as diop or binap (Scheme 8.35). The chirality of the diphosphine makes two transition states possible, leading to two enantiomeric products (diastereomers), and therefore subject to differences in equilibrium concentrations, activation energies and reaction rates; one enantiomeric product may thus form with the exclusion of the other. The (diop)Rh(i)-­catalysed enantioselective hydrogenation is used commercially to prepare S-­(l)-­DOPA [i.e. S-­(l)-­3,4-­ dihydroxyphenylalanine], a drug used for the treatment of Parkinson's disease, and aspartame [i.e. S-­(l)-­phenylalanyl-­S-­(l)-­aspartic acid], used as an artificial sweetener, The (binap)Ru(ii) catalyst may become even more useful.170 The work on the design and development of catalysts for asymmetric hydrogenation and epoxidation of alkenes was recognized by the award of the Nobel Prize in Chemistry to W. S. Knowles, R. Noyori and K. B. Sharpless in 2001.171

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Scheme 8.35  Some  chiral diphosphine ligands: (a) R,R-­(−)-­diop; (b) R-­(+)-­binap; (c) S-­(−)-­binap.

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89. G. A. Hamilton and J. P. Friedman, J. Am. Chem. Soc., 1963, 85, 1008; G. A. Hamilton, J. P. Friedman and P. M. Campbell, J. Am. Chem. Soc., 1966, 88, 5266; G. A. Hamilton, J. W. Hanifin Jr and J. P. Friedman, J. Am. Chem. Soc., 1966, 88, 5269. 90. H. Kroll, J. Am. Chem. Soc., 1952, 74, 2036. 91. (a) D. Hopgood and R. J. Angelici, J. Am. Chem. Soc., 1968, 90, 2508, 2511; (b) P. A. Sutton and D. A. Buckingham, Acc. Chem. Res., 1987, 20, 357. 92. (a) R. W. Hay and P. Banerjee, J. Chem. Soc., Dalton Trans., 1981, 352; (b) R. W. Hay and A. K. Basak, J. Chem. Soc., Dalton Trans., 1982, 1819. 93. R. Steinberger and F. H. Westheimer, J. Am. Chem. Soc., 1951, 73, 429. 94. S. G. Seltzer, A. Hamilton and F. H. Westheimer, J. Am. Chem. Soc., 1959, 81, 4018. 95. E. Gelles and R. W. Hay, J. Chem. Soc., 1958, 3673. 96. K. N. Leong and D. L. Leussing, Proc. 14th Int. Confc. Coord. Chem., Toronto, 1972, p. 608. 97. E. Gelles, J. Chem. Soc., 1956, 4736. 98. H. Irving and R. J. P. Williams, Nature, 1948, 162, 746. 99. M. Munakata, M. Matsui, M. Tabuschi and T. Shigematsu, Bull. Chem. Soc. Jpn., 1970, 43, 114. 100. E. Gelles and A. Salama, J. Chem. Soc., 1958, 3689. 101. E. Gelles and A. Salama, J. Chem. Soc., 1968, 3683. 102. G. S. Hammond, J. Am. Chem. Soc., 1955, 77, 334. 103. (a) J. E. Prue, J. Chem. Soc., 1952, 2331; (b) D. W. Larsen and M. W. Lister, Can. J. Chem., 1968, 46, 823. 104. (a) L. Goswami, S. Sarkar and D. Banerjea, Z. Anorg. Allg. Chem., 1977, 435, 301; (b) B. Chakravarti, R. N. Banerjee and D. Banerjea, Proc. 20th Int. Confc. Coord. Chem., Calcutta, 1979, p. 298. 105. R. N. Banerjee and D. Banerjea, J. Indian Chem. Soc., 1978, 55, 969. 106. (a) J. V. Rund and R. A. Plane, J. Am. Chem. Soc., 1964, 86, 367; (b) J. V. Rund and K. G. Claus, Inorg. Chem., 1968, 7, 860; 1969, 8, 59. 107. P. R. Bontchiev and V. Michaylova, J. Inorg. Nucl. Chem., 1967, 29, 2945. 108. R. W. Hay and K. G. Leong, J. Chem. Soc. A, 1971, 3639. 109. J. E. Prue, J. Chem. Soc., 1952, 2331. 110. (a) M. L. Bender and W. B. Turnquest, J. Am. Chem. Soc., 1957, 79, 1889; (b) M. L. Bender and R. J. Thomas, J. Am. Chem. Soc., 1961, 83, 4189. 111. Y. Yamane, M. Miyazaki, N. Yamaji and T. Ogashiwa, Chem. Pharm. Bull., 1971, 19, 2342. 112. H. A. Krebs, Biochem. J., 1942, 36, 303. 113. R. W. Hay, in Metal Ions in Biological Systems, ed. H. Sigel, Marcel Dekker, New York, 1976, vol. 5, p. 152. 114. (a) A. Kornberg, S. Ochoa and A. H. Mehler, J. Biol. Chem., 1948, 174, 159; (b) D. F. Herbert, in Methods of Enzymology, vol. 1, ed. S. P. Colowick and N. O. Kaplan, Academic Press, New York, 1995; p. 753; (c) J. F. Speck, J. Biol. Chem., 1949, 178, 315.

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Chapter 9

Photochemical Reactions of Metal Complexes 9.1  Introduction The behaviour of silver halides when illuminated with light, which is used in photography, is possibly the earliest observation of photochemistry, i.e. a light-­induced chemical reaction. Photochemistry is a fast developing area in chemistry and there is a profuse literature in the field.1–6 Energy is imparted to molecules when they are subjected to irradiation. The energy so imparted on irradiation in the UV–visible region (ca. 200–600 nm) amounts to ca. 150–50 kcal mol−1, which is a few times that imparted thermally at ordinary temperatures. Absorption of radiation by any species, including metal complexes, takes place at certain wavelengths which are characteristic of these species and depends on their electronic energy states (Table 9.1). The absorption bands (Table 9.1) for the simple anions, which are colourless, are due to charge-­transfer absorption in the UV region. The rather intense colours of CrO42− and MnO4− are also of the charge-­transfer type of absorption. For the complex ions of the transition metals, the absorption bands are due to d–d transitions, although for Ir(iii) these are in the UV region owing to the high value of 10Dq. In a transition metal complex, such as octahedral ML6, the d–d transition arises from promotion of an electron from the non-­bonding t2g to the antibonding eg* orbital and as a result there is weakening of the M–L bonds. Hence various reactions such as substitution (ligand replacement), isomerization and racemization processes occur easily with the species in the photoexcited state. The species in such an excited state may undergo several kinds of energy transfer, of which the most straightforward is luminescence, in which light energy is re-­emitted as the excited state returns to the ground state   Mechanisms of Reactions of Metal Complexes in Solution By Debabrata Banerjea and M. K. Bharty © Debabrata Banerjea and M. K. Bharty 2023 Published by the Royal Society of Chemistry, www.rsc.org

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Table 9.1  Absorption  of some species in aqueous solution. Species ν/cm−1

Species

ν/cm−1

Species

ν/cm−1

I− SCN− N3− S2O32– SO32– NO3−

[Cr(OH2)6]3+ [Cr(NH3)6]3+ [Cr(en)3]3+ [Cr(CN)6]3– [Co(OH2)6]3+ [Co(NH3)6]3+

17 400, 24 500 21 550, 28 500 21 600, 28 500 26 700, 32 200 16 600, 24 800 21 000, 29500

[Co(en)3]3+ [RhF6]3– [Rh(NH3)6]3+ [Ir(NH3)5(OH2)]3+ [Ir(NH3)6]3+ [IrCl(NH3)5]2+

21 400, 29 500 21 300, 27 800 34 100, 38 900 38 800, 47 000 39 800, 46 800 35 000, 44 100

44 200 47 000 42 500 46 000 50 000 47 600

Figure 9.1  Comparison  of photochemical and thermal activation of ground state

I forming products A, B and C; A′, B′ and C′ are the activated states in thermal reactions.

(a radiative process). Alternatively, energy in the excited state may be converted into vibrational energy and dissipated thermally as the species returns to the ground state (an internal conversion process). Of much greater interest is that a relatively long-­lived excited state may undergo a variety of reactions that may be different from those of a thermally activated state (Figure 9.1). Photochemically excited state II formed from the ground state I may luminesce with a rate constant kL, undergo intersystem crossing with a rate constant kIC to form an excited state III, or react with rate constants kA, kB, kC, etc., forming

Photochemical Reactions of Metal Complexes

499

the products A, B, C, etc. In Figure 9.1 it is also seen how products that are inaccessible owing to thermodynamically (A) or kinetically (C) unfavourable situations might result from photochemical reaction. The transformation I → A has a positive ∆G value, hence the change is thermodynamically unfavourable, whereas for the transformation I → C the activation energy is very high, hence the change is kinetically unfavourable. An example is that on irradiation of an aqueous solution of [Co(NH3)6]3+ at a lower wavelength ligand field band (29 500 cm−1, 338 nm) it readily changes to [Co(NH3)5(OH2)]3+, but an acidic solution of [Co(NH3)6]3+ can be boiled for hours without any change, despite favourable thermodynamics in this case, as the rate is extremely slow (the photochemically excited state is ∼1013 times more reactive than the ground state). Another complicating possibility is the so-­called intersystem crossing, in which a transition occurs to a different excited state III, which then undergoes any of the changes shown in Figure 9.1. For transition metal complexes, several types of photochemically excited states are feasible. A ligand field (LF) excited state results from a d–d transition which generally places electron density in an eg* (σ*) orbital, and this is obviously expected to cause M–L bond weakening and thus favour labilization of the ligand L, leading to its dissociation and replacement with another ligand L′ or re-­entry of L forming a geometrical or optical isomer (leading to racemization in the last case). By absorbing radiation corresponding to a charge-­transfer transition (M → L or L → M), a charge transfer (CT) excited state results. A third type is a transition localized on the ligand resulting in an intra-­ligand transition state, which may lead to ligand rearrangement or reaction of the ligand with other substrates. The other possibility is a charge transfer to solvent transition, leading to an oxidized metal ion and a solvated electron (e−solv), which may lead to electron transfer to the oxidized metal from another substrate and/or reduction of solvent or any other substrate by the e−solv. Irradiation of [Ru(bipy)3]2+ at 590 nm, or irradiation of other similar complexes of Ru(ii), produces a CT excited species {[Ru(bipy)3]2+}*, which is actually an oxidized ruthenium(iii) with an electron localized on bipy. This excited species is both a better oxidizing agent and a better reducing agent than [Ru(bipy)3]2+, as is evident from the reduction potential (E°) values (Table 9.2). {[Ru(bipy)3]2+}* is luminescent; quenching of the luminescence

Table 9.2  Reduction  potentials of ruthenium-­bipyridine complexes. E°/V at 25 °C

System −

{[Ru(bipy)3] }* + e   → [Ru(bipy)3] [Ru(bipy)3]3+ + e−  → {[Ru(bipy)3]2+}* [Ru(bipy)3]2+ + e−  → [Ru(bipy)3]+ [Ru(bipy)3]3+ e−  → [Ru(bipy)3]2+ 2+

+

+0.84 −0.84 −1.28 +1.26

Chapter 9

500

by another species Q (quencher) in solution is evidence of a reaction that could be reduction, oxidation or energy transfer depending on the nature of Q (Scheme 9.1). Energy transfer is involved in the role of [Ru(bipy)3]2+ as a sensitizer in the photoaquation of PtCl42−.7 These oxidation and reduction reactions may be used for the conversion and storage of light energy as chemical energy, since [Ru(bipy)3]3+ is an excellent oxidizing agent and [Ru(bipy)3]+ is an excellent reducing agent. [Ru(bipy)3]3+ oxidizes OH− to O2 and this is likely to be a step in the photolytic splitting of water into hydrogen and oxygen by solar energy in the presence of a similar Ru(ii) complex as the photocatalyst and a suitable quencher. However, dissociation of H2O into H2 + ½O2 by thermal excitation is not feasible under easily attainable conditions. The reaction Cl2 → 2Cl (∆H = 58 kcal mol−1) occurs to ∼1% at 975 °C (1 atm), but 40% dissociation of Cl2 takes place at room temperature on irradiation with UV light. A binuclear complex such as A5M–L–M′B5, where M and M′ may be two different metals in appropriate oxidation states or the same metal in different oxidation states, may absorb radiation, causing excitation due to electron transfer from one metal centre to the other in the species, which is known as an intervalence transfer (IT) process that is responsible for their usual intense colours. Redox reaction can be initiated in many such complexes by irradiating with radiation appropriate for the IT band of the species [eqn (9.1)].   

  

h [(H3 N)5 CoIII  NC  RuII (CN)6 ]   Co2   5NH3  [Ru(CN)6 ]3  (9.1)

It is a general and useful approximation to associate photosubstitution, photoisomerization and photoracemization reactions with LF excitation and photoredox reactions with CT excitation, but this rule is not an absolute one. A CT excitation may lead to ligand substitution by an indirect path (Scheme 9.2). As is evident from Scheme 9.2, the first step occurs exceedingly fast ( 1 suggests the involvement of an OH radical [eqn (9.4)]. However, the Φ for[Mo(CN)8]3− is even higher, suggesting a reaction taking place according to eqn (9.5). Prolonged photolysis of [M(CN)8]4− (M = Mo, W) forms a final blue species through initial formation of a red intermediate [eqn (9.6)], which reverts to the original species in the dark [eqn (9.7)]. In an isolated solid (as the potassium salt) of the blue species, an anion [M(CN)4(O)2]4− is present,14a,b but the species in solution is likely to be [M(CN)4(O)(OH)]3−. Evidence in support of the reaction shown in eqn (9.6) has been reported.14c The photochemistry of [M(CN)8]4− (M = Mo, W) has been reviewed.13c   



h [M(CN)8 ]3  H2 O  [M(CN)8 ]4   H  OH

(9.3)



[M(CN)8]3− + OH → [M(CN)8]4− + H+ + ½O2

(9.4)

[Mo(CN)8]3− + H2O + OH → [Mo(CN)8]4− + H+ + 2OH

(9.5)

  

  

  

h

[M(CN)8 ]4   2H2 O # [M(CN)7 (OH2 )]3  HCN  OH   



dark

red

[M(CN)7(OH2)]3− + OH− ⇌ [M(CN)4(O)(OH)]3− + 2HCN + CN−

  

(9.6) (9.7)

Photolysis of [Mn(C2O4)3]3− is consistent with the overall reaction shown in Scheme 9.4.15a Other [M(C2O4)3]3− (M = Fe, Co, etc.) complexes behave similarly.15,16 We shall now discuss in more detail the photochemistry of some six-­ coordinate complexes of Co(iii) and Cr(iii), and a few other six-­coordinate complexes that have been studied extensively. For both Co(iii) and Cr(iii), substitution, isomerization and racemization reactions are known to occur through photoexcitation of their complexes. In addition, in the case of Co(iii) and Cr(iii) complexes such as [Cr(bipy)3]3+, photoredox reactions are also known to occur. Photochemical data for a few representative complexes of Co(iii) are given in Table 9.3.17a,b

Scheme 9.4

Photochemical Reactions of Metal Complexes

503

Table 9.3  Photochemical  data for some Co(iii) complexes

17a,b

Complex ion [Co(NH3)6]3+ [CoX(NH3)5]2+ X = Cl X=I X = SCN (N-­bonded) X = Br X = NO2 (N-­bonded) X = N3 [Co(SO4)(NH3)5]+ [Co(CN)6]3− [CoX(CN)5]3− X = Cl X = Br X=I [Co(C2O4)3]3−

.

Irradiating radiation, λex/nm Products

Quantum yield, Φ/mol einstein−1

370

No reaction

—

370 550 370 550 370 550 370 550 370 370 550 370 370

[Co(NH3)5(OH2)]3+ [Co(NH3)5(OH2)]3+ Co(ii), I2 Co(ii), I2 [Co(ii)/Co(NH3)5(OH2)]3+ = 2.13 [Co(ii)/Co(NH3)5(OH2)]3+ = 0.243 [Co(ii)/Co(NH3)5(OH2)]3+ = 1.0 [Co(NH3)5(OH2)]3+ [Co(ii)/Co(NH3)5(OH2)]3+ = 1.85 Co(ii) Co(ii) No reaction [Co(CN)5(OH2)]2−

0.011 0.001 0.66 0.10 0.045 6.7 × 10−4 0.21 0.001 1.0 0.44 0.011 — 0.9

370 370 370 550 370 550

[Co(CN)5(OH2)]2− [Co(CN)5(OH2)]2− [Co(CN)5(OH2)]2− [Co(CN)5(OH2)]2− Co(ii), CO2 Co(ii), CO2

0.3 0.7 0.95 0.7 1.0 0.007

It can be seen that the nature of the products is a function of the wavelength (λex) of the irradiating radiation. Radiation of 550 nm, which causes LF excitation, produces more aquation than 370 nm radiation, which causes CT excitation and encourages redox reaction. However, the nature of the ligands, particularly the ease of oxidation of an attached ligand, is very important. Thus, [CoI(NH3)5]2+ undergoes 100% redox reaction, forming Co(ii) and I2, at both wavelengths mentioned above, but [CoCl(NH3)5]2+ undergoes 100% aquation of Cl− ligand (Cl− is replaced with H2O) at both wavelengths. Acidopentacyano complexes of Co(iii) differ from many other complexes of Co(iii) in giving only photoaquation upon ligand-­to-­metal charge transfer (LMCT) irradiation (see Table 9.3). For [CoX(CN)5]3− the quantum yields for aquation decrease in the order I− > CN− > Br− > Cl− for X−, which is the order of decreasing ease of oxidation of X− and not of M–X bond strengths (order of X− in the spectrochemical series). Hence the primary act involving the excited species is a homolytic M–X bond breaking, although the final product in each case corresponds to aquation and not a redox reaction. Because CN− is a very strong-­field ligand, the absorption in these cases in the near-­UV region (370 nm) causes LF excitation (due to a d–d-­t ype transition), and the same also holds good for [Co(CN)6]3−, which undergoes photoaquation. Adamson and Sporer17a accordingly proposed a three-­stage mechanism, the first stage being homolytic bond breaking but with the products remaining associated

Chapter 9

504

as in an outer-­sphere complex (and this change may involve one or more steps) [eqn (9.8)].   



III h    MII A 5  X  M XA 5  

(9.8)

  

The species MIIA5·X can react in one or two steps: it can either return to the ground state (back-­reaction to give [MIIIA5X]) or further react with water molecules in aqueous medium [eqn (9.9)].   



MIIA5 · X + H2O → MIIA5(OH2) · X

(9.9)

  

The outer-­sphere aqua MII species can further react in one of the two ways: if electron transfer from MII to X is energetically favourable, aquation will take place [eqn (9.10)]:   

  

MIIA5(OH2) · X → [MIIIA5(OH2)] + X−

(9.10)

but if X is not a good electron acceptor (oxidizing agent), the resulting products will correspond to a redox process [eqn (9.11)]. The free radical X may react in various ways. It may recombine with another X forming X2 (as for I), or else it may react as shown in eqn (9.12).17c   



MIIA5(OH2) · X → [MIIA5(OH2)] + X

(9.11)

[CoI(NH3)5]2+ + I → CoII + 5NH3 + I2

(9.12)

  

  

For the bromo complex Br2 is not generated, as free bromine oxidizes NH3 to N2 and becomes reduced to Br−. However, using the flash photolysis technique I and Br atoms were detected in the photolysis of the corresponding [CoX(NH3)5]2+ ions.18a The reverse of the reaction of eqn (9.8) accounts for the fact that quantum yields (Φ) are usually less than 1. The Φ values are temperature dependent, which can be explained in terms of different activation energies for the reactions of eqn (9.8)–(9.11). However, variations in the yields of several of the products with wavelength have been observed in some cases, which seems to suggest that different intermediates are involved in the photolytic aquation and redox reactions of these complexes. The use of very energetic UV radiation certainly causes a change in mechanism. Thus, both [Co(NH3)6]3+ and [Co(en)3]3+, which are unaffected by light in the near-­UV region (e.g. 370 nm), are decomposed by radiation of wavelength 254 nm, and similarly for several other complexes of Co(iii).18b,c Photolysis of cis-­[Co(NO2)2(en)2]+ in aqueous solution by irradiation in the LMCT band leads to the formation of Co(ii) and also the nitrito linkage isomer of the Co(iii) complex.18d Photolysis of [Co(acac)2(NH3)(N3)] in the 250–580 nm region leads to the formation of [Co(acac)2] and the azide radical.18d Irradiation of an aqueous solution of [(NC)5Co–NC–Co(NH3)5] in the 254–365 nm region (LF excitation) leads to photoaquation of the

Photochemical Reactions of Metal Complexes

505 2−

pentacyanocobaltate(iii) centre to give [Co(CN)5(OH2)] and [Co(CN)(NH3)5]2+ with Φ = 0.2–0.3. The wavelength of light used corresponds to a ligand field absorption band of the [CoIII(CN)5] chromophore. However, LF excitation of the same in the case of the linkage isomer [(NC)5Co–CN–Co(NH3)5] leads to very little reaction at that site. Since [Co(CN)6]3− is photoactive under the conditions, the result implies that coordination of the [CoIII(NH3)5]3+ moiety to the bridging CN group provides a new pathway for rapid deactivation of the LF excited states of the [Co(CN)6]3− in this binuclear complex.19a Photolysis of the S-­bonded isomer of the sulfinato complex [Co(SO2CH2 CH2NH2)(en)2]2+ forms the O-­bonded isomer.19b,c The latter is photoinert, but reverts thermally in the dark to the S-­bonded form with a rate constant of ∼10−5 s−1 at ca. 60 °C (t½ ≈ 600 h at room temperature). Photolysis of [Co{Me4[14]tetraeneN4}(OH2)2]3+ (Me4[14]tetraeneN4 = 1) complex in aqueous acidic medium at different wavelengths leads to reduction of Co(iii), but the macrocyclic ligand remains intact.19d

In the case of Cr(iii) complexes similar to those in Table 9.3, redox reactions do not occur; they undergo only substitution, isomerization and racemization reactions. Results of studies of some Cr(iii) complexes are given in Table 9.4.17a,20 In these cases the Φ values are almost insensitive to the wavelength used in the 300–700 nm region. The Φ values increase with increase in temperature owing to activation energies20a,21 as high as 14 kcal mol−1. Such a high activation energy for a photochemical reaction means that following photoexcitation, a process takes place that requires appreciable thermal energy for this to occur. Obviously the process is relatively slow, hence the reacting intermediate species must have a sufficient lifetime. The species must last a millisecond or so to account for the observed effect of temperature on Φ. This intermediate has been postulated20a,22 to be the doublet state 2Eg. To understand this, we need to consider the electronic energy states of a typical octahedral complex of d3 Cr(iii) taking into consideration the states that arise from the field-­free states 4F, 4P and 2G (this being the order of increasing energy) of a d3 ion in an octahedral field as follows: Free ion states

States in an octahedral field

4

4

F 4 P 2 G

A2g, 4T2g, 4T1g T1g 2 Eg, 2T1g, 2T2g, 2A1g 4

Chapter 9

506

Table 9.4  Photochemical  data for some Cr(iii) complexes

17c,20

Complex/system 3+

[Cr(OH2)6]

[Cr(NH3)6]3+ [Cr(NH3)5(OH2)]3+ [Cr(NH3)4(OH2)2]3+ [Cr(NH3)3(OH2)3]3+ [Cr(NH3)2(OH2)4]3+ [Cr(NCS)(NH3)5]2+ [Cr(NCS)(OH2)5]2+ [Cr(NH3)5(OH2)]3+ + SCN− [Cr(OH2)6]3+ + SCN− [Cr(OH2)6]3+ + Cl− (+)569-­[Cr(C2O4)3]3−

.

Irradiating radiation, λex/ nm

Reaction products

Quantum yield, Φ/mol einstein−1

254 540–730 254 320–600 320–600 320–700 320–700 320–700 360 560 400 575 560

O exchange using 18OH2 O exchange using 18OH2 [Cr(NH3)5(OH2)]3+ [Cr(NH3)5(OH2)]3+ [Cr(NH3)4(OH2)2]3+ [Cr(NH3)3(OH2)3]3+ [Cr(NH3)2(OH2)4]3+ [Cr(NH3)(OH2)5]3+ [Cr(NH3)5(OH2)]3+ [Cr(NH3)5(OH2)]3+ [Cr(OH2)6]3+ [Cr(OH2)6]3+ [Cr(NCS)(NH3)5]2+

0.03 0.02 0.49 0.32 0.25 0.16 0.014 0.0018 0.018 0.013 0.003 1 × 10−4 0.075

400 575 370 370

[Cr(NCS)(OH2)5]2+ [Cr(NCS)(OH2)5]2+ [CrCl(OH2)5]2+ Racemization (no decomposition or aquation)

0.0024 0.002 0.006 0.045

Figure 9.2  Ordering  of energy sates of d3 in an octahedral field. Energy separations of the states are only qualitative [2Eg = (2Eg, 2T1g), see text]. According to convention, the lowest (ground) state is considered to have zero energy.

The ordering of these octahedral field states in the sequence of energy is shown in Figure 9.2. The ground electronic state in octahedral d3 is 4A2g(F), corresponding to the electronic configuration t32g in which each of the t2g orbitals is singly occupied, with all the three electrons having parallel spins. There are three other quartet states of higher energy (excited states), viz. 4T2g(F), 4T1g(F) and 4T1g(P);

Photochemical Reactions of Metal Complexes

507 2

1

the first two correspond to the excited state configuration t 2g, eg (all three electrons singly occupying an orbital, two in t2g orbitals and one in an eg orbital) with their spins parallel as in the case of the t32g ground state mentioned above; the third state [highest quartet state, 4T1g(P)] corresponds to the excited configuration t12g, eg2 (all three electrons singly occupying an orbital with their spins parallel). The absorption spectra of Cr(iii) complexes in the visible region are due to the following two spin-­allowed transitions (in order of increasing energy): A2g(F) → 4T2g(F), 4A2g(F) → 4T1g(F)

4

The third band due to 4A2g(F) → 4T1g(P) in most cases appears in the UV region, where it is obscured by strong CT bands. The doublet “free ion” state 2 G [configuration t32g with one of the t2g orbitals occupied by a pair of electrons with opposite spins (L.S. state) with another t2g orbital being singly occupied by an electron] splits in an octahedral ligand field into four states of which the first two are almost identical in energy, viz. 2Eg(G) and 2T1g(G), and these two together are referred to as the 2Eg state; these are followed by higher energy states 2T2g(G) and 2A1g(G) in that order of increasing energy. In complexes of Cr(iii) of all the usual ligands, H2O, NH3, etc., the ordering of these energy levels in order of increasing energy is shown in Figure 9.2. The transitions 4A2g → 2Eg and 4A2g → 2T2g are very weak spin-­forbidden transitions and these weak absorptions are of not much significance in interpreting the absorption spectra of Cr(iii) complexes responsible for their colours. However, the excited doublet state 2Eg is surely of concern in the photochemistry of Cr(iii) complexes. Based on the usual interpretation of the relative rates of reactions of transition metal complexes (see Chapter 1), the excited t32g configuration with a vacant t2g orbital will allow facile reactions of this excited-­state species. From the widths of the absorption bands, Schläfer23 estimated the radiative lifetimes of these different excited states. The results indicate that the excited quartet states will revert to the ground state in 10−6–10−7 s, but the transition from 2Eg to 4A2g, which is spin forbidden, has a lifetime of 10−3–10−4 s. Collisional deactivation will cause the two higher quartet states to drop to the lowest excited quartet state 4T2g in 10−10–10−11 s, with no emission of radiation, The main path for transition from the 4T2g state to the ground state is also non-­radiative in nature as no fluorescence is observed. A radiationless transition from 4T2g to 2Eg also takes place in a time of 10−7–10−8 s.23 Hence photochemical excitation to any of the quartet states will cause some population of the 2Eg state. As this is rather long lived and is also expected to be labile, it is likely to be the intermediate for chemical reactions of Cr(iii) complexes on irradiation. As is evident from Table 9.4, the quantum yield (Φ) for [Cr(NH3)x(OH2)6−x]3+, falls steadily from [Cr(NH3)6]3+ to [Cr(OH2)6]3+. This decline is explained by the small energy difference (ca. 7 kcal mol−1)20b,c between the 2 Eg and 4T2g states in the [Cr(OH2)6]3+. In this case, thermal energy may cause a 2 Eg to 4T2g transition and then to the 4A2g state, thus causing a decrease in Φ.

508

Chapter 9

The larger separation (due to a higher ligand field strength) in [Cr(NH3)6]3+ would not allow this process. According to Wegner and Adamson,24 the excited state 4T2g of configuration t22geg1 will have Jahn–Teller distortion with a consequent elongation of two trans L–M–L bonds [as is known in cases of six-­coordinate complexes of Cu(ii) and Cr(ii)], and its lifetime could be long because its geometry is quite different from that of the undistorted ground state. Evidence reported25 in favour of the doublet state in the photochemical reaction of [Cr(en)3]3+ was later re-­examined and from the results it was concluded26 that involvement of the quartet state is more likely. For Λ-­[Cr(en)3]3+ in acidic aqueous solution also the lowest quartet state is photoreactive and the wavelength-­ independent Φ values for the three products in the 365–685 nm wavelength region, viz. Λ-­cis-­, ∆-­cis-­ and trans-­[Cr(en)2(enH)(OH2)]4+, are 0.10, 0.03 and 0.24, respectively.27 However, Kutal and co-­workers28 observed that irradiation at 669.2 nm populates the 2Eg state and enhances the reaction efficiency by ca. 50%; at this wavelength, there is little or no competitive absorption due to the quartet state. Very low Φ values (0.02–0.05) have been reported29 for photolysis of [Cr(NCS)4L]− (L = bipy, phen) by excitation with λex in the LF region, suggesting that the quartet states are of little significance in their photolysis. Based on available evidence, a mechanism involving the doublet 2Eg state seems likely in most photochemical reactions of Cr(iii) complexes. In the case of octahedral complexes of Cr(iii) of lower symmetry, such as in [CrX(NH3)5]2+ (X = CI, Br, I, SCN) and cis-­[Cr(en)2(OH)2]+, there is a difference between thermal and photolytic substitution, which for cis-­[Cr(en)2(OH)2]+ leads to different proportions of isomerization and aquation, and for [CrX(NH3)5]2+ to varying amounts of NH3 and X− aquation. Also, for many of these complexes, the photolytic products and Φ values are wavelength dependent.30 In such complexes of lower symmetry, the doublet and quartet states are closer in energy and their photochemical activity could derive from chemical activity of two excited states, either a doublet or a quartet or two quartet states. Irradiation of [CrX(NH3)5]2+ complexes at short wavelengths of ∼250 nm leads to an increase in X− labilization. Schläfer and co-­workers30,31 identified X− release with a CT state, and preferred the doublet state for NH3 release. In the photolysis of [Cr(en)3]3+ Schläfer32 identified a CT contribution. Chen and Porter33 identified photochemical activity of both a doublet and a quartet state in the photochemical reactions of reineckate ion, [Cr(NCS)4(NH3)2]−. Adamson et al.34 in studies on photosensitized reactions of [Cr(NCS)(NH3)5]2+ found support for their view35 that the doublet state leads only to SCN− aquation and the quartet state only to NH3 release. Adamson et al.1c,3,36 proposed the following empirical rules for predicting the results of photolysis in complexes of the type [CrX(NH3)5]2+:    1. Considering the fact that the six ligands lie in pairs at the ends of the three mutually perpendicular axes, the axis having the smallest average ligand field will be the one labilized, and the total quantum yield will be about that of a complex of Oh symmetry of the same average field.

Photochemical Reactions of Metal Complexes

509

2. If the labilized axis has two different ligands, then the ligand of greater ligand field strength will preferentially aquate. 3. The discrimination implied in Rules 1 and 2 will occur to a greater extent if the irradiation is in the first quartet band rather than the second.    However, these rules are not infallible.4 Thus, photolysis of [CrCl(NH3)5]2+ leads preferentially to cis-­[CrCl(NH3)4(OH2)]2+,30 which points to the labilized axis being a trans-­NH3–Cr–NH3 and not trans-­NH3–Cr–Cl as predicted by Rule 1. Photoaquation of [Cr(NH3)6]3+ and [CrX(NH3)5]2+ (X = Cl, Br, NCS) was investigated at high pressure under LF excitation. The ∆V‡ values (cm3 mol−1) obtained were −6.4 for NH3 aquation in the first complex and −13.0, −12.2 and −9.8 for the X− aquation in the second complexes for X = Cl, Br and NCS, respectively. The data suggested an Id mechanism.37a Ligand field excited photoaquation of cis-­[Cr(CN)2(NH3)4]+ in acidic aqueous solution involves mostly loss of NH3 and ca. 10% loss of CN−. The quantum yields are wavelength dependent (Φ ≈ 0.24–0.30 mol einstein−1 for NH3 loss and 0.010–0.022 mol einstein−1 for CN− loss). The product of NH3 release, [Cr(CN)2(NH3)3(OH2)]+, is a mixture of 1,2-­CN-­3-­H2O and 1,2-­CN-­6-­H2O isomers, the ratio of which varies from ca. 2 : 1 to 1 : 1 on moving the wavelength of the exciting radiation from the first to the second LF band.37b Photolytic aquation of cis-­ and trans-­[Cr(CN)2(en)2]+ was studied in acidic aqueous solutions at 12–14 °C.37c The trans isomer has essentially one reaction mode in which proton uptake is observed that is linked to one-­ended dissociation of a chelate ring and the proton becoming attached to the free NH2, and there is hardly any loss of CN−. However, the cis isomer undergoes a similar change and also perceptible loss of CN− (Table 9.5). Even on prolonged irradiation there is hardly any change in the case of trans-­ [Cr(CN)2(cyclam)]+ and three possible reasons for this unusual photoinactivity have been discussed.37d ESR and other evidence suggests the formation of CrV-­nitrido species in the UV photolysis of azido complexes of Cr(iii) such as [Cr(N3)L]n− (L = edta4−, nta3−, salen2−, etc.) [eqn (9.13) and (9.14)].37e Published reviews may be consulted for more details on the photochemistry of Cr(iii) complexes.38   



  

 [Cr III  N 3  L  OH2 ]n  h [Cr V  N  L   OH2 ]n   N 2

(9.13)

Table 9.5  Photoaquation  of cis and trans isomers of dicyanobis(ethylenediamine) chromium(iii).

Complex

λex/nm

H /mol einstein−1

CN /mol einstein−1

cis-­[Cr(CN)2(en)2]+

366 436 366 436

0.45 0.51 — 0.58

0.104 0.088 ∼0 ∼0

trans-­[Cr(CN)2(en)2]+





Chapter 9

510



h [Cr III (N 3 )2 L] n  1     [Cr V  N  L(OH2 )]n   N 2  N 3 H2 O

(9.14)

  

Available data on the photochemistry of Fe(ii) and Fe(iii) complexes (both H.S. and L.S. types) are well documented1c,e,4–6 so only selected aspects will be discussed here. A well-­studied system of Fe(ii) (L.S.) is [Fe(CN)6]4−, which exists in acidic aqueous solutions as [HFe(CN)6]3− (pH 3–4) and [H2Fe(CN)6]2− (pH 1). Two observed photochemical reactions of the species in these solutions are shown in eqn (9.15) and (9.16). Reaction (9.15) occurs with the highest efficiency in the CT bands, but does also occur in the LF band.39 Reaction (9.16) occurs over a wide range of wavelengths but with higher photochemical efficiency in the LF bands.40a The e−aq can be trapped using N2O as scavenger, generating an OH• radical that oxidizes [Fe(CN)6]4− to [Fe(CN)6]3− [eqn (9.17)]. The primary product in reaction (9.16) is thermally and photochemically unstable and forms various other products. The photochemistry of [Ru(CN)6]4− is similar40b to that of [Fe(CN)6]4−.   



  



h [Fe(CN)6 ]4   [Fe(CN)6 ]3  e aq  [Fe(CN)6 ]4 

Fe  CN 6 

4

3

h    Fe  CN 5  OH2    CN  H2 O



(9.15) (9.16)

  





(9.17)

  

Photochemical reactions of [Fe(CN)6]3− are quite complex and are summarized in Scheme 9.5.41 The visible spectra of [Fe(bipy)3]3+ and [Fe(phen)3]3+ are dominated by ligand → Fe CT absorption at ca. 600 nm. Following excitation of these in aqueous solution in this region their excited-­state species will oxidize water to OH radical42 or other reducing agents or other ground-­state complex ions.43

Scheme 9.5

Photochemical Reactions of Metal Complexes

511 2+

Photolysis of an acidic aqueous solution of [Fe(OH2)6] involves the primary process shown in eqn (9.18). This is followed by a set of complicated pH-­dependent processes. Excitation of an aqueous solution of [Fe(OH2)6]3+ and of [FeX(OH2)5]2+ (X = OH, CI, Br, SCN, N3) in the CT region (L → M) leads to oxidation of the coordinated ligands (H2O, X−) or of any reducing species present in solution (oxidation by OH or X radicals).   



h

[Fe(OH2 )6 ]2  /[Fe(OH2 )6 ]3  e aq

  

(9.18)

Photochemistry of [Fe(C2O4)3]3− mentioned earlier is of importance because of its role in actinometry in the CT region (λex < 550 nm).43,44 The Φ value is very low for the spin-­forbidden d–d band region (ca. 680 nm). Flash photolysis studies established the reaction sequence shown in Scheme 9.6.16,44a Results of investigations on the primary quantum yield at different wavelengths for the photodecomposition of [M(C2O4)3]3− (M = Mn, Fe, Co) have been reported.15 [Cr(C2O4)3]3− undergoes no photodissociation; however, the optical isomer of [Cr(C2O4)3]3− shows fairly efficient photoracemization;44a,45a photoracemization has also been observed for several other complex ions.44,45 Based on the general scheme for photodecomposition of [M(C2O4)3]3− [as shown earlier in this section for the Mn(iii) complex], the primary quantum yield is half that of the M(ii) complex formed. The primary quantum yield for photodecomposition of the tris-­oxalato complexes of Mn(iii), Fe(iii) and Co(iii) is nearly constant in each of the cases in the CT region and falls off in the d–d band region; the region of fall-­off of Φ matches the region of minimum absorption between the CT and d–d bands.1b Optical isomers of metal complexes for which photochemistry has been investigated44 include [M(C2O4)3]3− (M = Cr, Co, Rh), [Co(C2O4)(en)2]+, [Co(NO2)2(en)2]+, [Co(en)3]3+, [Co(edta)]−, etc. [Cr(C2O4)3]3− has been studied in detail because it undergoes no photochemical reaction other than racemization. Some of the results reported are given in Table 9.6, which show that neither the solvent nor the wavelength of light used has any significant effect on the quantum yield. Also, unlike thermal racemization, which shows acid catalysis, the quantum yield for photoracemization shows no detectable dependence on acidity. Experiments in D2O indicated a reduction in the thermal racemization rate, kH2 O / kD2 O = 1.26, and a comparable reduction in the quantum yield at 420 nm, H2 O /D2 O = 1.24. Similarly, 18O exchange experiments showed comparable ratios of krac/kexch = 2.6 and Φrac/Φexch = 2.3,

Scheme 9.6

Chapter 9

512

Table 9.6  Photoracemization 

45a

3−

of [Cr(C2O4)3] .

λ/nm

Solvent

Temperaturea/°C

I × 106/ einstein s−1

R × 107/M s−1

Φb

420 420 420 420 420 570 697

Water Water Water 30% ethanol 30% ethanol 30% ethanol 30% ethanol

1.0 8.0 15.0 15.0 4.0 4.0 4.0

1.85 1.53 1.98 1.09 0.39 3.17 0.27

1.50 1.64 1.85 3.65 0.35 2.58 0.18

0.081 0.106 0.093 0.043 0.089 0.082 0.068

a

The temperature dependence for thermal racemization in water is given by krac = 1.63 × 107exp(−15 000/RT) s−1 and in 30% ethanol–water by krac = 6.23 × 108exp(−17 600/RT) s−1, compared with Φ = 3.67exp(−2100/RT) and Φ = 25.8exp(−3200/RT), respectively. b The quantum yield, Φ, is the moles of complex racemized per einstein of light absorbed.

Table 9.7  Photochemical  decomposition and racemization of some metal complexes.45

Complex 3+

[Co(en)3] [Co(C2O4)(en)2]+ [Co(C2O4)2(en)]− [Co(edta)]− [Co(C2O4)3]3− cis-­[Co(NO2)2(en)2]+ cis-­[CoCl2(en)2]+ [Cr(C2O4)3]3− [Rh(C2O4)3]3−

Comments No photodecomposition No photodecomposition Photodecomposition Slight photodecomposition Photodecomposition No photodecomposition Photodecomposition No photodecomposition Photodecomposition

No photoracemization No photoracemization Photoracemization No photoracemization Slight photoracemization No photoracemization No photoracemizationa Photoracemization Photoracemization

a

 onsidering only aquation to cis-­[CoCl(en)2(OH2)]+, this occurs with retention as it does in the C thermal aquation.

respectively, for these two processes. These results indicated that the thermal and photochemical racemizations proceed by related mechanisms, presumably involving chelate ring opening and closing (see Chapter 6, Section 6.5), with the bond rupture being driven photochemically. It has been known from earlier reports that the principal thermal and photochemical reaction of [Co(C2O4)3]3− is oxidation–reduction to yield Co(ii) and CO2.46 However, for the optically active complex the rate of disappearance as deduced by following the rate of loss of optical activity had been found to be noticeably larger than that obtained from the decrease in concentration as measured directly by following the optical absorbance decrease. The results showed that about 15% of the time, loss of optical activity was through racemization. Photochemical observations made on the decomposition and racemization of several complexes are summarized in Table 9.7. For the Co(iii) complexes of the series [Co(en)3]3+ to [Co(C2O4)3]3−, it should be noted that the replacement of en by C2O42− labilizes the complex towards photo-­ (and

Photochemical Reactions of Metal Complexes

513

thermal) redox decomposition and also towards photo-­ (and thermal) racemization. This behaviour supports the view of Adamson17b that the photolability of Co(iii) complexes towards racemization (and aquation) appears to be closely related to the ease of oxidation of the acido groups and thus to the ease of photoredox decomposition. Similarly, the observation that there is no photodecomposition for [Cr(C2O4)3]3− and only a small amount for [Rh(C2O4)3]3− is understood in terms of these metals being more difficult to reduce to the divalent state than is Co(iii). In the photodecomposition of [Pt(ox)(PPh3)2], the primary reaction [eqn (9.19)] is followed by reactions leading to the formation of Pt0, [Pt2(PPh3)4] [but Pt(PPh3)4 in the presence of excess PPh3 in solution] and CO2.47   





 + C O  

h Pt  ox   PPh3 2    Pt  PPh3 2   ethanol 

_ *

*

2

4

(9.19)

  

The primary act in the photolysis of acidic aqueous solutions of Tl(iii) involves CT (L → M) excitation [eqn (9.20)] and is likely to be followed by the reaction shown in eqn (9.21).   

h

h



Tlaq 3 /Tl2   H  OH or Tl(OH)2 / Tl2   OH

(9.20)



Tl2+ + OH → Tl+ + H+ + ½O2

(9.21)

  

  

A photoinduced chain reaction between TlI and TlIII has been reported.48a FeII-­Sensitized photolysis of [Cu(C2O4)2]3− yields Cu and CO2.16 Extensive studies on the photochemistry of ceric ammonium nitrate solutions have been made by various groups.48a A postulated mechanism involves intramolecular electron transfer [eqn (9.22)] (CeIV nitrato species not specified). Dogliotti and Hayon48b concluded from flash photolytic studies that OH is first formed [eqn (9.23)].   



h CeIV (NO3 )  {CeIV ( NO3 )}*  CeIII  NO3

(9.22)

h CeIV (OH2 )   CeIII  H OH

(9.23)

  



  

48c,d

However, Martin and co-­workers favoured the original postulates of NO3 radical formation and also photolysis of HNO3 in the solution [eqn (9.24)]. The reaction CeIII + NO3 → CeIV(NO3−) also takes place. Dogliotti and Hayon48b also detected HSO4 radicals in the flash photolysis of Ce(iv) sulfate solutions.

  



  

h

HNO3 /NO2  OH; OH  HNO3  NO3  H2 O

(9.24)

Photolysis of trans-­[RhX2(NH3)4]+ (X− = halide) gives trans-­[RhX(NH3)4 (OH2)]2+.49 However, photoaquation of cis-­[RhX2(NH3)4]+ (X = Cl, Br) gives trans-­[RhX(NH3)4(OH2)]2+, whereas cis-­[RhX(NH3)4(OH2)]2+ forms the trans isomer on irradiation. These reactions occur through a square-­ pyramidal intermediate having X− in an apical position.50 Photoaquation51

Chapter 9

514 2+

of cis-­ and trans-­[RhCl(NH3)4(OH2)] gives trans-­[Rh(NH3)4(OH2)2]3+, but trans-­[RhCl(NH3)4(OH)]+ undergoes concurrent aquation and isomerization on irradiation, forming cis-­[Rh(NH3)4(OH)2]+. Based on these and earlier results, the preference for an apical position in the square-­pyramidal intermediate in these photochemical reactions52 is OH− < NH3 < H2O < Cl− < Br− < I−. Both cis-­ and trans-­[Rh(NH3)4(OH)2]+ are inert to photoisomerization, because the OH− ligand, being in the extreme left of the above series, having the weakest preference for the apical position is not photolabile.53 Irradiation of cis-­ and trans-­[RhBr(X)(en)2]n+ (X = Br−, H2O, NH3) in their ligand field bands gives trans-­[RhBr(en)2(OH2)]2+ in all cases except for cis-­ [RhBr(en)2(NH3)]2+, which forms cis-­[Rh(en)2(NH3)(OH2)]3+, and similarly for trans-­[RhBr(en)2(NH3)]2+. These results can be interpreted by a two-­stage mechanism with a square-­pyramidal intermediate, but with the postulate that the transformation (a) → (b) is difficult, whereas (c) → (d) is favourable, and this rationalizes the observed mixture of products in the photoaquation54 of cis-­[RhBr(en)2(NH3)]2+.

Similar observations have been reported in the photolysis of cis-­and trans-­ [RhX(I)(en)2]n+ (X = I−, NH3, H2O)55 and of cis-­[RhCl(en)2(NH3)]2+.56 Similar arguments can be put forth for the products of photolysis of cyano-­Co(iii) complexes.57 Extensive studies have been made on the photochemical behaviour of [IrX(NH3)5]n+ (X = Cl−, Br−, I−, H2O, NH3, MeCN, PhCN), trans-­[IrI2(NH3)4]+ and trans-­[IrI(NH3)4(OH2)]2+. Photolysis at ligand field frequencies generally resulted in slow photoaquation with 100% retention of stereochemistry, where relevant. However, photolysis at CT bands gives mainly photoaquation with a small redox change. These results broadly parallel those for the analogous Rh(iii) complexes [but there is no redox contribution in the case of Rh(iii)].58 As with Rh(iii), trans-­[IrX(OH)(en)2]+ (X = Cl, Br, I) photoaquates with concurrent photoisomerization, giving cis-­[Ir(OH)2(en)2]+.59 Photosubstitution in Rh(iii) complexes has been discussed in relation to Co(iii) and Ir(iii),60 and also Ru(ii) and Ru(iii).61 Through photochemical reactions, UO22+ effects the oxidation of a wide range of organic and inorganic substrates, which have been well covered in reviews.2,4 There has been wide use of the UO22+–oxalate system in actinometry. The photochemical yield is measured by loss of C2O42− and the products formed, mainly CO2, CO and HC(O)OH, vary in their proportions with pH, but are independent of λex and temperature. Low pH favours CO formation. Heidt et al.62 carried out thorough studies of the pH dependence of the CO

Photochemical Reactions of Metal Complexes

515

Scheme 9.7 yield and correlated increased CO formation with C–O bond breaking by H+ in the energy-­rich complex (Scheme 9.7). In MeCN solution at room temperature, the complex [NiCl2(dpp)] (dpp = Ph2PCH2CH2CH2PPh2) exists as a mixture of planar and tetrahedral complexes with an equilibrium constant [planar]/[tetrahedral] = 1.33. Laser irradiation of a 1 mM solution of the complex in MeCN at 1060 nm causes complete conversion of the tetrahedral to the planar form. In the dark, the original equilibrium is re-­established and the rate of conversion of the planar to the tetrahedral isomer, followed either at 470 nm (decrease in planar concentration) or at 380 nm (increase in tetrahedral concentration), showed a relaxation time τ of 0.95 µs at 24 °C, leading to a k1 + k−1 value corresponding to k1 planar # tetrahedral k1

The value so obtained at 24 °C (k1 + k−1 = 1/τ) was 1.05 × 106 s−1. Since K = k−1/k1 = 1.33, the k1 and k−1 values obtained were 4.5 × 105 and 6.0 × 105 s−1, respectively.63 Carassiti and co-­workers64 investigated the photolytic aquation of PtCl42− [eqn (9.25)]. They also suggested reasonably that the higher value of the quantum yield (Φ = 0.65) in the CT region is due to this CT state, which extensively aquates before it deactivates to a lower LF state. The near constancy of Φ (0.13–0.14) at 313, 404 and 472 nm was attributed to the two excited singlet states resulting from the first two LF excitations and the triplet state from the LF excitation at 472 nm all undergoing radiationless transition (no luminescence observed), crossing to the lowest component of 3Eg triplet state, which corresponds to the electronic configuration ( d z2)1(dxz,dyz)4(dxy)2( d x2  y2 )1, which would be vulnerable to nucleophilic attack in the axial position by H2O. However, Sastri and Langford65 suggested that not all the excited species reach the 3Eg state, as this could give Φ close to 1. *Br− exchange with PtBr42− is also light sensitive and involves a [PtBr3(OH2)]− intermediate generated photochemically.66 Photosubstitution reactions of [PtBr(Et4dien)]+ were investigated in aqueous solution by irradiation at 313 mm. The aqua complex is the product in acidic or neutral solutions even in the presence (0.01 M) of free Br− or Cl−. In alkaline medium (pH 12) the product is the hydroxo complex. In both acidic and alkaline solutions, the reported quantum yield is 0.004.67   



  

h

PtCl 42   H2 O/[PtCl3 ( OH2 )]  Cl 

(9.25)

Chapter 9

516

An interesting photochemical reaction of Pt(iv) is the nitrito → nitro linkage isomerization of [Pt(ONO)(NH3)5]3+.68 Quantum yields for the photochemical exchange [eqn (9.26)] and substitution reactions (*Cl− denotes isotopically labelled Cl− such as β-­active 36Cl−) [eqn (9.27)] are of the order of 103 and increase with increase in temperature, pointing to chain mechanisms.69,70   

  



  

h PtCl62   *Cl   [PtCl5* Cl]2   Cl 

(9.26)

h PtCl62   X   [PtCl5 X]2   Cl 

(9.27)

where X− = Br− or I−. Rich and Taube69 postulated the formation of a Cl atom and a Pt(iii) species: h PtCl62    PtCl52   Cl

in the photoirradiation step of Cl− exchange with PtCl62−. This is followed by a chain reaction that accounts for very high Φ value for the photoexchange [eqn (9.28) and (9.29)].   



PtCl52− + *Cl− → [Pt*Cl5]2− + Cl−

(9.28)

[Pt*Cl5]2− + [PtCl6]2− → [Pt*Cl6]2− + [PtCl5]2−

(9.29)

  

  

In the absence of exchangeable anions photoaquation occurs [eqn (9.30)].   

  

h PtCl62   H2 O  [PtCl5 ( OH2 )]  Cl 

(9.30)

This process is reversed thermally by the addition of Cl− ions.71,72 Similar behaviour was reported for the photoaquation of [PtBr6]2−, which gives a wavelength-­independent quantum yield in acidic solution (Φ = 0.4 at 313, 365, 433 and 530 nm), leading to the conclusion that radiationless transitions of all the excited states populate a triplet LF state which aquates.73 In the photochemical Br− exchange the Φ value is ∼102, leading to the postulation of a radical chain reaction as for [PtCl6]2−,18a,73,74 involving PtBr52− or PtBr42− [eqn (9.31) and (9.32)]; the latter (PtBr42−) can also act as a chain carrier [eqn (9.33) and (9.34)].18a,64a,74c   

 PtBr62  h [PtBr5 ]2   Br

(9.31)



h PtBr62   [PtBr4 ]2   Br2

(9.32)



[PtBr4]2− + *Br− → [Pt*Br4]2− + Br−

(9.33)

[Pt*Br4]2− + PtBr62− → [Pt*Br6]2− + [PtBr4]2−

(9.34)

     

  

  

A transient Br2− species was detected in the photoaquation of trans-­ [PtBr2(NH3)4]2+.2 Photochemical reactions of PtI62− are complicated by fast

Photochemical Reactions of Metal Complexes

517 71

thermal processes, but aquation is accelerated and a long-­lived intermediate was identified in flash photolysis.18a As for the thermal reactions,75a photochemical reactions of Pt(iv) are also sensitive to trace amounts of Pt(ii). Photoisomerization of square-­planar complexes of Pt(ii) were first reported over a century ago.76 Photochemical conversion of cis-­and trans-­[PtCl2(PEt3)2] into a cis ⇌ trans equilibrium mixture is rapid,75b but the thermal reaction is exceedingly slow unless catalysed by free PEt3 (see Chapter 6).77 The excitation occurs by radiation in the wavelength range of d–d bands (Φ ≈ 0.01).75b,c A later investigation78 of cis ⇌ trans isomerization in [PtCl2(PEt3)2] reported Φ values of 0.09 and 1.0 for the forward and reverse reactions, respectively. Both cis-­and trans-­[PtCl2(py)2] undergo photoisomerization and photodissociation at 313 nm with Φisom ≈ 0.04 and Φdiss ≈ 0.025 for the cis isomer, but for the trans isomer Φ ≈ 10−3.79a The cis → trans conversion of Pt(gly)2 occurs on irradiation in the d–d band region, but the trans isomer is not affected. For the isomerization of cis-­Pt(gly)2, the Φ values are 0.12 (254 nm) and 0.13 (313 nm).64a,79b,c The process is intramolecular, proceeding via a twist mechanism (as no exchange occurs with 14C-­labelled gly− in the solution), whereas the corresponding thermal isomerization is intermolecular and requires the presence of free gly− in solution and hence exchange occurs with the 14C-­labelled gly− in the solution. Irradiation of trans-­[M(gly)2] (M = Pd, Pt) in the CT region leads to decompositions.64a,79b Fairly convincing evidence was presented for destabilization of the square-­ planar geometry79c,80 in favour of a distorted tetrahedral structure, due to promotion of a non-­bonding d electron of Pt(ii) into the d x2  y2 orbital, and the same process was suggested for [Ni(CN)4]2−.81 It is known that some of the d–d bands in square-­planar complexes of Pt(ii) are due to such singlet– triplet transitions, which explains why many such photochemical reactions occur on irradiation in the d–d region, but there is still some confusion as to the exact assignment.82 The proposed activated state would be vulnerable to both isomerization and substitution; the latter is more probable for complexes of monodentate ligands, whereas for those of chelating ligands intramolecular isomerization will be favoured.4b The photochemical steady state is solvent dependent; more polar solvents favour the polar cis isomers, as emphasized by studies on [PtX2(PPh3)2] isomers.83 Work has been reported on the photochemistry of metallocene and metal–benzene complexes,84 including substitution in the ferrocene ring.85

References 1. (a) A. W. Adamson and P. D. Fleischauer, Concepts of Inorganic Photochemistry, Wiley, New York, 1975; (b) F. Basolo and R. G. Pearson, Mechanisms of Inorganic Reactions: A Study of Metal Complexes in Solution, John Wiley, New York, 2nd edn, 1967, pp. 332–334, 546-­548 and 654-­665; (c) A. W. Adamson, W. L. Waltz, E. Zinato, D. W. Watts, P. D. Fleischauer and R. D. Lindholm, Chem. Rev., 1968, 68, 541; (d) C. G. Garcia, J. E. de Lima and N. Y. M. Iha, Coord. Chem. Rev., 2000, 196, 219; (e) J. Chen and W. R. Browne, Coord. Chem. Rev., 2018, 374, 15.

518

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2. P. Chini, Inorg. Chim. Acta, Rev., 1968, 2, 31. 3. A. W. Adamson, Coord. Chem. Rev., 1968, 3, 169. 4. (a) E. Rabinowich and R. L. Belford, Spectroscopy and Photochemistry of Uranyl Compounds, Monograph on Nuclear Energy, Chemistry Division McMillan Co., New York, 1964, vol. 1; (b) V. Balzani and V. Carassitim, Photochemistry of Coordination Compounds, Academic Press, New York, 1970. 5. M. A. Jamieson, N. Serpone and M. Z. Hoffman, Coord. Chem. Rev., 1981, 39, 121. 6. R. B. Bucat and D. W. Watts, Inorganic Photochemistry, in Inorganic Chemistry Series One, M. T. P. International Review of Science, ed. M. L. Tobe, Butterworths, London, 1972, ch. 5, vol. 9 and references cited therein. 7. J. N. Demas and A. W. Adamson, J. Am. Chem. Soc., 1971, 93, 1800. 8. (a) E. O. Fischer, H. P. Koglert and P. Kuzel, Chem. Ber., 1960, 93, 3006; (b) W. Strohmeier and K. Gerlack, Z. Naturforsch., 1960, 16b, 413; (c) W. Strohmeier and D. V. Hobe, Chem. Ber., 1961, 94, 761; (d) W. Strohmeier, D. von Hobe, G. Schönauer and H. Laporte, Z. Naturforsch., 1962, 17b, 502; (e) W. Strohmeier, Angew. Chem., Int. Ed. Engl. 1964, 3, 370. 9. (a) G. R. Dobson, I. W. Stolz and R. K. Sheline, Adv. Inorg. Radiochem., ed. H. J. Emeleus and A. G. Sharpe, Academic Press, New York, 1966, vol. 8, pp. 1–89; (b) I. W. Stolz, G. R. Dobson and R. K. Sheline, J. Am. Chem. Soc., 1962, 84, 3589; J. Am. Chem. Soc., 1963, 85, 1013; (c) A. J. Rest and J. J. Turner, J. Chem. Soc., Chem. Commun., 1969, 1026. 10. (a) J. R. Moss and W. A. G. Graham, J. Chem. Soc., Chem. Commun., 1970, 835; (b) E. L. Wehry, Q. Rev., Chem. Soc., 1967, 213; (c) M. Wrighton, Chem. Rev., 1974, 74, 401; (d) Inorganic Photochemistry, in Advances in Inorganic Chemistry, ed. R. van Eldik and G. Stochel, Academic Press, London, 2011, vol. 63. 11. (a) A. J. Rest, J. Chem. Soc., Chem. Commun., 1970, 345; (b) A. J. Rest and J. J. Turner, J. Chem. Soc., Chem. Commun., 1970, 375. 12. I. J. Chang and D. C. Nocera, Inorg. Chem., 1985, 24, 4309. 13. (a) V. Balzani and V. Carassiti, Ann. Chim., 1961, 51(508), 533; 1962, 52, 432; (b) V. Balzani, H. Bertoluzza and V. Carassiti, Bull. Soc. Chim. Belg., 1962, 71, 821; (c) A. Samotus and J. Szklawzewicz, Coord. Chem. Rev., 1999, 125, 63. 14. (a) S. J. Lippard, H. Nozaki and B. J. Russ, J. Chem. Soc., Chem. Commun., 1967, 118; (b) S. J. Lippard and B. J. Russ, Inorg. Chem., 1967, 6, 1943; (c) V. Balzani, M. M. Manfrin and L. Moggi, Inorg. Chem., 1969, 8, 47. 15. (a) G. B. Porter, J. G. W. Doering and S. Karanka, J. Am. Chem. Soc., 1962, 84, 4027; (b) T. B. Copestake and N. Uri, Proc. R. Soc. A, 1955, 228, 252. 16. J. A. van den Berg and P. Schneider, J. S. Afr. Chem. Inst., 1968, 21, 113. 17. (a) A. W. Adamson and A. H. Sporer, J. Am. Chem. Soc., 1958, 80, 3865; (b) A. W. Adamson, Discuss. Faraday Soc., 1960, 29, 163; (c) A. Haim and H. Taube, J. Am. Chem. Soc., 1963, 85, 595.

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18. (a) S. A. Penkett and A. W. Adamson, J. Am. Chem. Soc., 1965, 87, 2514; (b) J. F. Endicott and M. Z. Hoffan, J. Am. Chem. Soc., 1964, 87, 3348; (c) V. Balzani, V. Carassiti, L. Moggiand and N. Sabbatini, Inorg. Chem., 1965, 4, 1247; (d) P. Natarajan and A. Radhakrishnan, J. Chem. Soc., Dalton Trans., 1982, 2403; (e) J. L. Reed, Inorg. Chem., 1981, 20, 2590; Inorg. Chem., 1982, 21, 2829. 19. (a) M. Nishizawa and P. C. Ford, Inorg. Chem., 1981, 20, 2016; (b) H. Macke, V. Houlding and A. W. Adamson, J. Am. Chem. Soc., 1980, 102, 6888; (c) V. H. Houlding, H. Macke and A. W. Adamson, Inorg. Chem., 1981, 20, 4279; (d) T. G. Hoon and C. Y. Mok, Inorg. Chim. Acta, 1982, 62, 231. 20. (a) R. A. Plane and J. P. Hunt, J. Am. Chem. Soc., 1957, 79, 3343; (b) M. R. Edelson and R. A. Plane, J. Phys. Chem., 1959, 63, 327; (c) M. R. Edelson and R. A. Plane, Inorg. Chem., 1964, 3, 231. 21. A. W. Adamson, J. Inorg. Nucl. Chem., 1960, 13, 275. 22. (a) H. L. Schlӓfer, Z. Phys. Chem., 1957, 11, 5; (b) H. L. Schlӓfer, Z. Elektrochem., 1960, 64, 886. 23. H. L. Schlӓfer, Z. Phys. Chem., 1965, 69, 2201. 24. E. E. Wegner and A. W. Adamson, J. Am. Chem. Soc., 1966, 88, 394. 25. R. Fukuda, R. T. Walters, H. Macke and A. W. Adamson, J. Phys. Chem., 1979, 83, 2097. 26. A. D. Kirk, J. Phys. Chem., 1981, 85, 3205. 27. M. C. Cimolino and R. G. Linck, Inorg. Chem., 1981, 20, 3499. 28. X. Yang, C. A. Surton and C. Kutal, Inorg. Chem., 1982, 21, 2893. 29. B. Olarte, H. Krentzun and C. Bifano, Chem. Abstr., 1981, 95, 159743c. 30. H. F. Wasgestian and H. L. Schlӓfer, Z. Phys.-­Chem., 1968, 57, 282; 1968, 62, 127. 31. P. Riccieri and H. L. Schlӓfer, Inorg. Chem., 1970, 9, 727. 32. H. L. Schlӓfer, Z. Phys.-­Chem., 1969, 65, 107. 33. S. N. Chen and G. B. Porter, Chem. Phys. Lett., 1969, 6, 41. 34. A. W. Adamson, J. E. Martin and F. D. Camessot, J. Am. Chem. Soc., 1969, 91, 7530. 35. E. Zinato, R. D. Lindholm and A. W. Adamson, J. Am. Chem. Soc., 1969, 91, 1076. 36. (a) A. W. Adamson, J. Phys. Chem., 1967, 71, 798; (b) A. W. Adamson, Rec. Chem. Prog., 1968, 29, 191; (c) A. W. Adamson, Pure Appl. Chem., 1969, 20, 25. 37. (a) K. Angerman, R. van Eldik, H. Kelm and F. Wasgestian, Inorg. Chem., 1981, 20, 955; (b) P. Riccieri and E. Zinato, Inorg. Chem., 1983, 22, 2305; (c) A. D. Kirk and G. B. Porter, Inorg. Chem., 1980, 19, 445; (d) N. A. P. Kane-­Maguire, W. S. Krippen and P. K. Miller, Inorg. Chem., 1983, 22, 696; (e) S. I. Arshankov and A. L. Poznyak, Z. Anorg. Allg. Chem., 1981, 481, 201; (f) S. I. Arshankov and A. L. Poznyak, Russ. J. Inorg. Chem., 1981, 26, 850. 38. (a) M. A. Jamieson, N. Serpone and M. Z. Hoffman, Coord. Chem. Rev., 1981, 39, 121; (b) A. D. Kirk, Coord. Chem. Rev., 1981, 39, 225. 39. G. Stein, Adv. Chem., 1965, 50, 238.

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Photochemical Reactions of Metal Complexes

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68. V. Balzani, N. Sabbatini and V. Carassiti, Progress in Coordination ­Chemistry, ed. M. Cais, Elsevier, Amsterdam, 1968, p. 80. 69. R. L. Rich and H. Taube, J. Am. Chem. Soc., 1954, 76, 2608. 70. R. Dreyer and K. Konig, Z. Chem., 1966, 6, 271. 71. E. Blasius, W. Prteez and R. Schmitt, J. Inorg. Nucl. Chem., 1961, 19, 115. 72. R. Dreyer, I. Dreyer and D. Rettig, Z. Phys. Chem., 1963, 224, 199. 73. V. Balzani, M. F. Manfrin and L. Moggi, Inorg. Chem., 1967, 6, 354. 74. (a) V. Balzani and V. Carassiti, J. Phys. Chem., 1968, 72, 383; (b) G. Schmidt and W. Z. Herr, Z. Naturforsch., 1961, 16a, 748; (c) A. W. Adamson and A. H. Sporer, J. Am. Chem. Soc., 1958, 80, 3865. 75. (a) F. Basolo and R. G. Pearson, Mechanisms of Inorganic Reactions: A Study of Metal Complexes in Solution, John Wiley, New York, 2nd edn, 1967, pp. 237–238; (b) P. Haake and T. A. Hylton, J. Am. Chem. Soc., 1962, 84, 3774; (c) B. L. Shaw and P. R. Brookes, J. Chem. Soc., Chem. Commun., 1968, 919. 76. L. Ramberg, Chem. Ber., 1910, 43, 580. 77. J. Chatt and R. G. Wilkins, J. Chem. Soc., 1952, 273. 78. S. H. Goh and C. Y. Mok, J. Inorg. Nucl. Chem., 1977, 39, 531. 79. (a) L. Moggi, C. Varani, N. Sabbatini and V. Balzani, Mol. Photochem., 1971, 3, 141; (b) V. Balzani, V. Carassiti, F. Scandola and L. Moggi, Inorg. Chem., 1965, 4, 1243; (c) F. Scandola, O. Traverso, V. Balzani, G. L. Zucchini and V. Carassiti, Inorg. Chim. Acta, 1967, 1, 76. 80. D. S. Martin Jr, M. A. Tucker and A. J. Kassman, Inorg. Chem., 1965, 4, 1682; 1966, 5, 1298. 81. C. J. Ballhausen, N. Bjerrum, R. Dingle, K. Eriks and C. R. Hare, Inorg. Chem., 1965, 4, 514. 82. (a) H. B. Gray and C. J. Ballhausen, J. Am. Chem. Soc., 1963, 85, 260; (b) D. S. Martin Jr and C. R. Lenhardt, Inorg. Chem., 1964, 3, 1368. 83. S. H. Mastin and P. Haake, J. Chem. Soc., Chem. Commun., 1970, 202. 84. R. B. Bucat and D. W. Watts, Inorganic Photochemistry, in Inorganic Chemistry Series One, M. T. P. International Review of Science, ed. M. L. Tobe, Butterworths, London, 1972, ch. 5, vol. 9, p. 182 and references cited therein. 85. H. C. H. A. van Riel, Tetrahedron Lett., 1969, 3085.

Subject Index acetylene (alkyne) polymerization, 484–487 acid-­catalysed reactions, 22 acid catalysis, 256–276 acid hydrolysis, 3 activated complex, 80 activation of molecules coordinated ligands, reactivity of amino acid ester hydrolysis, 460–461 catalysed oxidation, 457–460 coordinated CO2 and SO2, 450–451 coordinated organic ligands, 452–457 β-­keto acids, decarboxylation of, 461–467 metal-­bound CO ligand, 449–450 NO in nitrosyl complexes, 451–452 Wacker process, 467 diatomic molecules dihydrogen by coordination, 435–441 dinitrogen by coordination, 445–449 dioxygen by coordination, 441–444 insertion reactions, 467–469 acetylene (alkyne) polymerization, 484–487

carbon dioxide, 478–480 carbon disulfide, 480 of CO, 470–476 coordination compounds, asymmetric synthesis catalysed by, 487–488 olefin (alkene) polymerization, 482–483 olefins (alkenes), 480–481 sulfur dioxide, 476–478 amino acid ester hydrolysis, 460–461 anation reaction, 3 angular overlap model (AOM), 116 aquapentaamminecobalt(iii), rate constants of reactions of, 290 aquation, 3 associative mechanism, 9 base hydrolysis, 3, 46, 231–246 Berry pseudo-­rotation, 118 bimolecular associative (SN2) pathway, 423–424 biphasic irreversible system, 77–79 π-­bonding theory, 109, 122 of the trans effect, 121–125 Brønsted relationship, 91 carbon dioxide, 478–480 carbon disulfide, 480 Cardwell's mechanism, 118 catalases, 459 catalysed oxidation, 457–460

522

Subject Index

centre of gravity rule, 31, 34 Chan and Wong's mechanism, Pt(ii) complexes, 120 cis-­labilizing effect, 165 cis-­labilizing series, 124–125 classical bond-­weakening concept, 109 cluster formation reactions, 4 CO, 470–476 Co(iii) complexes, 503 collision theory of reaction rates, 80 competition methods, 79 concurrent irreversible reactions, 76 conformational isomerization, 3, 360–362 constant-­flow method, 67 coordinated CO2 and SO2, 450–451 coordinated organic ligands, 452–457 Cr(iii) complexes, 506 crystal field stabilization energy (CFSE), 35, 39. See also ligand field stabilization energy (LFSE) crystal field theory, 30, 35 Debye–Hückel interionic potential, 14 Debye–Hückel theory, 87, 88, 89 Dewar, Chatt and Duncanson (DCD) model, 438 for metal‒olefin bonding, 121 dihydrogen by coordination, 435–441 dinitrogen by coordination, 445–449 dioxygen by coordination, 441–444 direct chemical analysis, 59–60 dissociative mechanism, 8–9 effective mass, 98 Eigen–Wilkins mechanism, 13, 327 electrochemical methods, 69–70 electrometric methods, 63–66 electron-­donating ability, 308

523

electron paramagnetic resonance (EPR) methods, 71–72 electron spin resonance (ESR), 72 electron-­transfer mechanism, 5, 285–286 electron transfer reactions, 3, 66 atom (or group) transfer processes, 395–411 comproportionation, 409 in heterogeneous systems, 413–416 inner-­sphere mechanism, 394–395 inner-­sphere reactions, redox rate constants for, 412–413 Marcus equation, 388–394 mixed outer-­and inner-­sphere reactions, 412 outer-­sphere mechanism, 388 oxidative addition reactions, 419–422 bimolecular associative (SN2) pathway, 423–424 concerted pathway, 422–423 five-­coordinate eighteen-­ electron substrates, 424–426 four-­coordinate eighteen-­ electron substrates, 426–427 four-­coordinate sixteen-­ electron substrates, 426 ionic mechanism, 424 mechanisms of, 422–424 radical mechanism, 424 rate constants for, 387 reductive elimination, 427–429 solvated electrons, 416–419 electron tunnelling theory, 388 electrophilic catalysis. See also nucleophilic catalysis acid catalysis, 256–276 by metal ions, 276–282

524

electrophilic substitutions, 3 electrostatic crystal field theory, 30 energy profiles for reaction, 16–17 enthalpy of activation, 81 enthalpy of transition, 82 entropy of activation, 81 EPR spectroscopy, 75–76 experimental evidence for mechanisms, 16–18 external pressure effect, 84–86 first-­order reaction, 6 five-­coordinate eighteen-­electron substrates, 424–426 flash photolysis, 71 flow methods, 67–69 fluoride exchange, 326 force constant, 98 formation reactions, 3, 286 aquapentaamminecobalt(iii), rate constants of reactions of, 290 general mechanistic scheme, 286–287 hydroxycuprate(ii) species, 302 ligand substitution, dissociative mechanism for, 295 low-­spin bis(imidazole) complexes, 293 macrocyclic ligands, 297 metal porphyrinate complexes, 294 outer-­sphere complex (OSC), 289 rate constants for water exchange, 288 second-­order formation rate constant, 289 four-­coordinate complexes, structural changes in, 377 four-­coordinate eighteen-­electron substrates, 424–425 four-­coordinate sixteen-­electron substrates, 424 free-­energy of activation, 81 Fuoss–Eigen equation, 14, 15

Subject Index

geometrical isomerization, 350–351 octahedral complexes, 353–360 square-­planar complexes, 351–353 geometric isomerism, 20 Grinberg's theories, 113 of the trans effect, 107, 108 group replacement factor (GRF), 211 Grunwald–Winstein equation, 224 Hammett relationship, 94–96 hydroxycuprate(ii) species, 302 Ilkovic equation, 65 inner-­sphere mechanism, 392–393 insertion reactions, 4 internal conjugate base (ICB) mechanism, 303, 304 internal conversion process, 498 intersystem crossing, 499 intimate mechanism, 10, 16 iodide ion, polarization of, 108, 109 ionic strength effect, 87–89 isokinetic trend, 82 isotope effects, 97–99 isotropic tracers, 67 β-­keto acids, decarboxylation of, 461–467 kinetic features of the reactions, 50 kinetic macrocyclic effect, 302 kinetic trans effect (KTE), 110, 111 kinetic wave, 65 lability of complexes, 26–29 of coordination compounds, 28 ion size and charge, 28 lanthanides, 40 non-­transition elements, 40 lanthanides, 40 Lewis acids and bases, 3 ligand exchange, 326

Subject Index

ligand field (LF) excited state, 499 ligand field stabilization energy (LFSE), 35, 36, 37, 38, 39, 47, 48 ligand field theory, 29, 30–52, 35, 38, 40, 41, 46, 48, 49, 50 ligand replacement reactions, 8–16 activation parameters, 218–223 charge on reaction rate, 198–199 group replacement factor (GRF), 211 leaving ligand, 194–198 of [M(CO)6], 246–247 process, 207–211 rate of replacement, comparison of, 226–228 reaction mechanism, 218–223 of s-­and p-­block metals, 247–248 solvent effect, 223–226 spectator (non-­leaving) ligands, electronic effects of, 206–207 spectator ligands, steric and structural effects of, 200–206 stoichiometric mechanisms, 228–231 ligand substitution, dissociative mechanism for, 295 linear free energy relationship (LFER), 97, 389 linkage isomerization, 347–350 linkage isomers, 3 low-­spin bis(imidazole) complexes, 293 Marcus equation, 388–394 Marcus's theory, 408, 500 metal-­bound CO ligand, 449–450 metal–ligand bond axes, 30 metal–ligand bond distances, 33 metal porphyrinate complexes, 294 microscopic reversibility, 2

525

migration–insertion reactions, 4 molar absorbance, 60 molecularity, 5–8 monophasic reversible reactions, 76–77 NMR spectroscopy, 72–75. See also nuclear magnetic resonance (NMR) methods NO in nitrosyl complexes, 451–452 nuclear magnetic resonance (NMR) methods, 71–72 nucleophilic catalysis, 282–285 nucleophilic constant, 91, 92 nucleophilicity, 89–90 values, 94 nucleophilic scale, 91 nucleophilic substitutions, 3 octahedral complexes, 31, 353–360 octahedral wedge structure, 36 olefin (alkene) polymerization, 482–483 olefins (alkenes), 480–481 one-­electron oxidants, 18 optical density, 60 optical inversion, 362–364 optical isomerism, 20 in tetrahedral complexes, 377 optical isomerization (racemization), 364–365 intermolecular mechanism, 365–367 intramolecular mechanism, 367–372 unsymmetrical chelating ligands, tris chelates of, 372–375 order of a reaction, 5–8 order of reaction, 21–26 outer-­sphere complex (OSC), 13, 289 outer-­sphere mechanism, 388 oxidative addition reactions, 4

526

photoaquation of cis and trans isomers, 509 photochemical decomposition and racemization, 512 photochemical reactions aqueous solution, 498 charge transfer (CT) excited state, 499 Co(iii) complexes, 503 Cr(iii) complexes, 506 internal conversion process, 498 intersystem crossing, 499 intra-­ligand transition state, 499 ligand field (LF) excited state, 499 vibrational energy, 498 photochemical (light-­induced) reactions, 4 photolysis, 508 photoracemization, 4 [Cr(C2O4)3]3−, 512 photoredox reactions, 4 photosubstitution, 4 planar complexes, configurational changes in, 378–379 Planck's constant, 81 polarimetric methods, 66 polarization theory, 109 polarographic method, 65 primary isotope effect, 98 primary kinetic isotope effect, 98 prompt reactions, 501 proton magnetic resonance (PMR) techniques, 71 pseudo-­first-­order rate constant, 12, 14, 23 pseudo-­first-­order reaction, 21 pseudo-­substitution process, 4, 41 pseudo-­zeroth-­order rate, 6, 7 Pt–X bond lengths, 113 pulse radiolysis, 76 rate-­determining step, 2, 9, 11 rates of reactions, factor affecting

Subject Index

external pressure effect, 84–86 ionic strength effect, 87–89 solvent influence, 89 temperature effect, 80–84 reaction mechanisms, 2, 21–26 reactions involving structural changes, 3 reactions of coordinated ligands, 3 reactive intermediates, 18–20 redox reaction, 3 reductive elimination, 4 relative nucleophilicities, 90–94 relaxation methods, 70–71 rhombic twist, 369 ruthenium-­bipyridine complexes, 499 Schrödinger equation, 98 secondary isotope effect, 98 second-­order formation rate constant, 289 second-­order reaction, 6, 7 sigma trans effect, 109 solvent influence, 89 solvolysis, 3 spectator ligands, 8, 25 spectator (non-­leaving) ligands, electronic effects of, 206–207 spectator ligands, steric and structural effects of, 200–206 spectrophotometric methods, 60–63 square-­planar complexes, 17, 351–353 π-­bonding theory of the trans effect, 121–125 cis effect in terms of discrimination, 165–167 cis-­labilizing effect, 165 complexes of platinum(ii), 105–107 coordination number five, 177–182 of copper(ii), 167–172 energy profile for reactions of, 154–155

Subject Index

gold(iii) complexes, 161–163 higher coordination number, 182–183 mechanism of reaction, 125–154 of nickel(ii), 167–172 palladium(ii) complexes, 156–161 platinum(iv) complexes, trans effect in, 155 tetrahedral complexes, reactions of, 172–177 trans effect in terms of discrimination, 164–165 trans effect theories, 107–121 stacking interactions, 307 static trans effect (STE), 110 steady-­state principle, 11 sterically controlled substitution, 303 stoichiometric mechanisms, 10, 16, 228–231 stoichiometry, 17 stopped-­flow method, 68 substrate constant, 91 sulfur dioxide, 476–478 Syrkin's resonance theory, 112, 113 Taft relationship, 96–97 temperature effect, 80–84

527

terdentate ligands of the type M(L′) (L″), 376 tetragonal configuration, 32 tetragonal twist, 369 tetrahedral complex, 31 thermodynamic effect, 118 thermodynamic sense, 26 thermodynamic trans effect, 118 third-­order reaction, 7 π-­trans effects of ligands, 110 σ-­trans effects of ligands, 110 transition state theory, 80 trans-­labilizing series, 124–125 ultra-­fast T-­jump method, 71 ultrasonic absorption, 76 UV-­visible spectroscopic methods, 16 valence bond theory, 28, 29, 40, 41 vibrational energy, 498 volume of activation, 85 Wacker process, 467 water exchange rates, 26, 306 X-­ray crystallographic data, 113 X-­ray photoelectron spectra, 116 zero-­point vibrational energies, 98 zeroth-­order reaction, 6 zwitterion, 316