ORGANIC ION RADICALS Chemistry and Applications
Zory V. Todres Columbus, Ohio, U.S.A.
Marcel Dekker, Inc. TM
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ORGANIC ION RADICALS Chemistry and Applications
Zory V. Todres Columbus, Ohio, U.S.A.
Marcel Dekker, Inc. TM
Copyright 2003 by Marcel Dekker, Inc. All Rights Reserved.
New York • Basel
ISBN: 0-8247-0810-5 This book is printed on acid-free paper. Headquarters Marcel Dekker, Inc. 270 Madison Avenue, New York, NY 10016 tel: 212-696-9000; fax: 212-685-4540 Eastern Hemisphere Distribution Marcel Dekker AG Hutgasse 4, Postfach 812, CH-4001 Basel, Switzerland tel: 41-61-260-6300; fax: 41-61-260-6333 World Wide Web http://www.dekker.com The publisher offers discounts on this book when ordered in bulk quantities. For more information, write to Special Sales/Professional Marketing at the headquarters address above. Copyright © 2003 by Marcel Dekker, Inc. All Rights Reserved. Neither this book nor any part may be reproduced or transmitted in any form or by any means, electronic or mechanical, including photocopying, microfilming, and recording, or by any information storage and retrieval system, without permission in writing from the publisher. Current printing (last digit): 10 9 8 7 6 5 4 3 2 1 PRINTED IN THE UNITED STATES OF AMERICA
Copyright 2003 by Marcel Dekker, Inc. All Rights Reserved.
To my wife, Irina The cloudless beauty of her heart, the profundity of her mind, and the depth of her feelings have always provided reliable support for me. For my children, Vladimir and Ellen, their mother represents a supreme example.
Copyright 2003 by Marcel Dekker, Inc. All Rights Reserved.
Foreword
This volume presents an overview of the widely varied subject matter that is relevant to understanding the chemistry of organic ion radicals. Detailed consideration is given to all aspects of ion radical solution chemistry, from discussing how ion radical structures are described theoretically to presenting the reagents commonly used to generate them, methods for their characterization, and synthetic methods based on ion radicals. Less usual topics are covered in depth in chapters on how to identify and study reactions proceeding through ion radicals, and how to optimize yields for such reactions. A chapter is devoted to solvent effects, which are especially important for such reactions. Examination of the practical applications of ion radicals includes coverage of their roles in biological systems as well as in material chemistry, ranging from optoelectronics, organic metals, and magnets to lubricants and the manufacture of paper. This book gathers into a single place the widely scattered material needed to consider ion radicals from an organic chemist’s point of view. A highly useful aspect of this book is its incorporation of relevant studies from the Russian literature, a significant amount of it originating from Todres’ own group in Moscow. Little of this material has previously received proper emphasis, or even citation, in the English-language literature. Stephen F. Nelsen, Ph.D. Professor Department of Chemistry University of Wisconsin Madison, Wisconsin
Copyright 2003 by Marcel Dekker, Inc. All Rights Reserved.
Preface
Contemporary organic chemistry puts great emphasis on investigations of the structure and reactivity of intermediate species, originating in the pathway from the starting compounds to the end products. Knowledge of properties of the intermediate species and penetration into the mechanism of reactions open new routes to increase the rate of formation and the yield of the desired final products. Until recently, chemists focused their attention on radicals or charged species of the cationic/anionic type. Species of an intermediate nature combining ionic and radical properties—ion radicals—remained outside the scope of their investigations. Perfected instrumental techniques markedly advanced fine experiments. As a result, the species, which were little (if at all) known to the chemists of previous decades, now came to the forefront. The behavior of organic ion radicals has become an area of current interest. Ion radicals arise by one-electron oxidation or one-electron reduction of organic compounds in isolated redox processes and as intermediates along the pathways of reactions. A conversion of an organic molecule into an ion radical brings about a significant change in its electron structure and a corresponding alteration in its reactivity. This conversion allows necessary products to be obtained under mild conditions with high yields and improved selectivity of transformation. In addition, there are several reactions that can proceed only by the ion radical mechanism and lead to products otherwise inaccessible. The theme of the book is the formation, transformation, and application of ion radicals in typical conditions of organic synthesis. The book presents an overview of organic ion radical reactions and explains the principles of ion radical organic chemistry. Methods of determining ion-radical mechanisms and controlling ion radical reactions are also reviewed. When applicable, issues relating to ecology and biology are addressed. The inorganic participants in the ion radical organic reactions are also considered. One of the chapters gives representative synthetic procedures and considers the background of related synthetic approaches.
Copyright 2003 by Marcel Dekker, Inc. All Rights Reserved.
The book also provides a review of current practical applications as well as an outlook on those predicted to be important in the near future. The reader will learn of the progress that has been made in technical developments utilizing the organic ion radicals. Electronic and optoelectronic devices, organic magnets and conductors, lubricants, and other applications are considered. The behavior of organic ion radicals is still not completely understood. Thus, new interpretations of scientific data appear frequently in the literature. I have attempted to synthesize the ideas from various references that complement one another, although the connections among them may not be immediately obvious. (An Author Index is included to help readers find such connections in the book.) Science is a collective affair, and its main task is to produce trustworthy and general knowledge. My apologies are due to any authors who have contributed significantly to the development of this vast field but, for various reasons, have not been cited. The many contributors who are cited certainly do not reflect my preferences; their publications have been selected as illustrative examples that may allow the reader to follow evolution of the corresponding topics. Every new branch of science passes through several stages in its progress, including the latent phase and the phase of increased interest. These initial phases have apparently already passed in organic ion radical chemistry; they spawned decades of heated debates. In recent years, however, the boil has cooled to a simmer. The development of organic ion radical chemistry, especially with respect to its practical applications, has allowed chemists to nail down ion radicals more confidently. Having become a regular division of general knowledge, organic chemistry of ion radicals is now entering the third stage of development, namely, putting the ideas elaborated into general operation, including them in the common property of organic chemistry. It is now necessary to generalize the obtained data and to treat them comprehensively. Grafting the new branch to the organic chemistry tree is the aim of this book. I have worked in the field of organic ion radicals and their applications for several decades and have become more and more fascinated by the beauty of this area and the diversity it presents. Understanding the role of ion radicals is as difficult as it is interesting. I hope that this attempt to graft this branch to the organic chemistry tree will be useful both for further advancing basic research and for facilitating new practical applications. During my entire working life, I, like other researchers, have felt the pressure of the scientific community’s appraisal. Criticism is crucial! The writing of this book was aided by discussions with my colleagues and friends. I am indebted to all of them for their corrections and polemics. At the same time, their support was a great incentive. The scientific atmosphere that surrounded me during my tenure at the Research and Educational Institute of the Cleveland Clinic Foundation, the time when I was writing the book, also stimulated my creative work. The book is meant for researchers and technologists who are carrying out syntheses and studying principles governing the choice of optimal organic reaction conditions. It will be useful for physical organic chemists, ecologists, biologists, and specialists in electronics of organic materials, as well as professors, researchers, and students. As for students, I have assumed that the reader is not acquainted with the field itself but possesses background knowledge in chemistry, both general and organic, a knowledge usually provided during undergraduate studies in most, if not in all, countries. Zory V. Todres
Copyright 2003 by Marcel Dekker, Inc. All Rights Reserved.
Contents
Foreword Preface Chapter 1. Nature of Organic Ion Radicals and Their Ground-State Electron Structure 1.1 Introduction 1.2 Unusual Features 1.3 Acid–Base Properties of Organic Ion Radicals 1.4 Metallocomplex Ion Radicals 1.5 Organic Ion Radicals with Several Unpaired Electrons and/or Charges 1.6 Polymeric Ion Radicals 1.7 Inorganic Ion Radicals in Reactions with Organic Substrates 1.8 Conclusion References Chapter 2. Formation of Organic Ion Radicals 2.1 Introduction 2.2 Chemical Methods of Organic Ion Radical Preparation 2.3 Equilibria in Liquid-Phase Electron-Transfer Reactions 2.4 Electrochemical Methods Versus Chemical Methods 2.5 Formation of Organic Ion Radicals in Living Organisms 2.6 Organic Ion Radicals in Solid Phases 2.7 Isotope-Containing Organic Compounds as Ion Radical Precursors 2.8 Conclusion References
Copyright 2003 by Marcel Dekker, Inc. All Rights Reserved.
Chapter 3. Basic Principles of Organic Ion Radical Reactivity 3.1 Introduction 3.2 Principle of the “Detained” Electron That Controls Ion Radical Reactivity 3.3 Principle of the “Released” Electron That Controls Ion Radical Reactivity 3.4 Behavior of Organic Ion Radicals in Living Organisms 3.5 Conclusion References Chapter 4. How to Discern Ion Radical Mechanisms of Organic Reactions 4.1 Introduction 4.2 Why Does a Reaction Choose the Ion Radical Mechanism? 4.3 Chemical Approaches to the Identification of Ion Radical Reactions 4.4 Physical Approaches to the Identification of Ion Radical Reactions 4.5 Examples of Complex Approaches to the Discernment of the Ion Radical Mechanism of Particular Reactions 4.6 Conclusion References Chapter 5. How to Optimize Organic Ion Radical Reactions 5.1 Introduction 5.2 Physical Effects 5.3 Effect of Chemical Additives 5.4 Solvent Role 5.5 Salt Effects 5.6 Conclusion References Chapter 6. Organic Ion Radicals in Synthesis 6.1 Introduction 6.2 Reductive Reactions 6.3 Ion Radical Polymerization 6.4 Cyclization 6.5 Ring Opening 6.6 Fragmentation 6.7 Bond Formation 6.8 Conclusion References Chapter 7. Practical Applications of Organic Ion Radicals 7.1 Introduction 7.2 Organic Ion Radicals in Optoelectronics 7.3 Organic Metals 7.4 Organic Magnets 7.5 Organic Lubricants 7.6 Ion Radical Routes to Lignin Treatment 7.7 Conclusion References
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Chapter 8. Concluding Remarks 8.1 Introduction 8.2 SRN1 Reaction 8.3 Stereochemical Aspects of Ion Radical Reactivity 8.4 Conclusion References
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1 Nature of Organic Ion Radicals and Their Ground-State Electron Structure
1.1 INTRODUCTION Organic chemistry represents an extensive body of facts, from which the contemporary doctrine of reactivity is built. The most important basis of this doctrine is the idea of intermediate species that arise along the way from starting material to final product. Depending on the nature of the chemical transformation, cations, anions, and radicals are created during an intermediate stage. These species form mainly as a result of bond rupture. Bond rupture may proceed heterolytically or homolytically (Scheme 1-1): R X ← R–X → R X
or
.
.
R–X → R X
Ions or radicals formed from a substrate further react, with other ions or radicals acting as reactants. Such changes in chemical bonds can be accompanied by a one-electron shift. The concept of the one-electron shift (Pross 1985) can be illustrated by Scheme 1-2 for nucleophilic substitution: Nu R↑↓Z → Nu↑↓R Z The species Nu↑↓R of Scheme 1-2 is not a radical pair; it is a covalent molecule of the product resulting from the SN2 reaction. The process of transfer of the R group to a Nu reactant proceeds in synchronicity with a one-electron shift and R–Z bond disruption. At . . that time, two radical particles, Nu and R (formed in the course of reaction) remain immediately close and therefore unite rapidly. A one-electron shift may or may not lead to the formation of radical particles. There are many reactions that consist not of a one-electron shift, but of a one-electron transfer. The initial results of such one-electron transfers involves the formation of ion radicals. . This book concentrates on species of the type (RX) , i.e., on cation and anion radicals. These species are formed during reaction, when an organic molecule either loses one electron from the action of an electron acceptor or acquires one electron from the action of
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an electron donor, Scheme 1-3: .
R–X e → (R–X)
or
.
R–X e → (R–X)
Ion radicals differ from their starting molecules only in a change of the total number of electrons; no bond rupture or bond formation occurs. As will be seen from the following chapters, after ion radical formation, cleavage and association reactions often occur. Geometry changes upon electron loss or gain can take place. Reactions with the participation of ion radicals bring their own specific opportunities. The concept of molecular orbital helps explain the electron structure of ion radicals. When one electron abandons the highest occupied molecular orbital (HOMO), a cation radical is formed. If one electron is introduced externally, it takes up the lowest unoccupied molecular orbital (LUMO), and the molecule becomes an anion radical. These ion radicals have a dual character. They contain an unpaired electron and are, therefore, close to radicals. At the same time, they bear a charge and, naturally, are close to ions. Being radicals, the ion radicals are ready to react with radicals. Like all other radicals, they can dismutate and recombine. Being ions, the ion radicals are able to react with particles of the opposite charge and are prone to form ionic aggregates. In contrast with radicals, the ion radicals are especially sensitive to medium effects. As long as an unpaired electron of an ion radical occupies an orbital covering all atoms of a molecule, a definite distribution of spin density occurs between individual atoms. The distribution determines the activity of one or another position of an ion radical species. From the point of view of organic synthesis, such properties of ion radicals as stability, resistance to active medium components, capacity to disintegrate in the needed direction, and the possibility of participating in electron exchange are especially important. All these properties become understandable (or predictable in cases of unknown examples) from the organic ion radical electron structure. Therefore, our account will be based on the analysis of connections between the structure of ion radicals and their reactivity or physical properties. This chapter concerns the peculiarities of conjugation in aromatic ion radicals, electron structures, and acid–base properties of ion radicals that have originated from molecules of different chemical classes. 1.2 UNUSUAL FEATURES 1.2.1 Substituent Effects The aim of this section is to show that substituent effects for organic ion radicals are quite different from those of their parent, neutral molecules. Amino and nitro compounds are good examples. N,N-Dimethylaniline is a molecule with a lone electron pair on the nitrogen atom. Of course, there is a strong interaction between this pair and the -electron system of the ben. zene ring. We often mark the symbol of the cation radical, i.e. , on the nitrogen atom. However, according to the ab initio Hartree–Fock molecular orbital calculations (Zhang, R., et al. 2000), this nitrogen atom is in fact negatively charged (0.708), and the positive charge is distributed on the carbon atoms, especially on the two methyl groups (0.482 on each). Influenced by the positive-charge delocalization along the cation radical, the benzene ring becomes an electron-deficient unit with a positive charge of 0.744. Summation yields the total charge of 1,000 for the N,N-dimethylamine cation radical. As seen, conventional ideas may not be applicable to the chemistry of ion radicals.
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TABLE 1-1 Nitrogen HFC Constants (aN ) from Experimental ESR Spectra of Nitro Compounds Compounds
Constant aN, mT
Nitroalkanes
2.4–2.5
Nitrobenzene 2-Chloronitrobenzene 2,6-Dichloronitrobenzene Nitrodurene
1.0 0.9 1.4 2.0
Reference Stone & Maki (1962), McKinney & Geske (1967) Geske & Maki (1960) Starichenko et al. (2000) Starichenko et al. (2000) Geske & Ragle (1961)
In anion radicals of nitro compounds, an unpaired electron is localized on the nitro group, and this localization depends on the nature of the core molecule bearing this nitro substituent. The value of the hyperfine coupling (HFC) constant in the ESR spectrum reflects the extent of localization of the unpaired electron; aN values of several nitro compounds are given in Table 1-1. Let us compare HFC data from Table 1-1. Aliphatic nitro compounds produce anion radicals, in which an unpaired electron spends its time on the nitro group completely. In the nitrobenzene anion radical, an unpaired electron is partially delocalized over the aromatic ring due to conjugation, and the nitrogen hyperfine coupling constant decreases by one half as compared with the aliphatic counterparts. Diminution of the -conjugation in the PhNO2 system as a result of the nitro group distortion intensifies the localization of the unpaired electron on the nitro group. In the nitrobenzene anion radical, however, an unpaired electron is not evenly spread between the nitro group and the benzene ring. The calculation (Stone & Maki 1962) of spin distribution in the nitrobenzene anion radical showed that the nitro group retains about 0.65–0.70 of the unit spin density (i.e., an unpaired electron resides mainly on the nitro group). These calculations are based upon HFC constants aN and aH taken from the nitrobenzene anion radical ESR spectrum. The same result arises from the unpaired electron distribution by the use of the MO Hueckel approximation: 0.31 of the unit spin density over the phenyl nucleus and 0.69 on the nitro group (Todres 1981). Of course, it is the entire molecule that receives the electron upon reduction. However, the nitro group is the part where the excess electron spends the majority of its time. Consideration of quantum chemical features of the nitrobenzene anion radical is of particular interest. The model for the calculation includes a combination of fragment orbitals for Ph and NO2, and the results are represented in Scheme 1-4. The left part of the scheme . refers to the neutral PhNO2, and the right part refers to the anion radical, PhNO 2 (Todres 1981). Some changes in all the orbital energies accompany the placing of an electron on the lowest unoccupied molecular orbital. According to the calculations, relative energy gaps remain unchanged for the orbitals in the nitrobenzene anion radical compared with those of the parent nitrobenzene. For the sake of graphic clearness, Scheme 1-4 disregards the difference mentioned, keeping the main feature of the equality in the energy gaps. The nitro group in the parent nitrobenzene evidently acts as -acceptor, which pulls the electron density out of the aromatic ring. An unpaired electron will obviously occupy the first vacant -orbital of the nitro fragment (i.e., the lowest-energy-fragment orbital). Interaction between occupied and vacant orbitals is the most favorable. In the nitrobenzene
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SCHEME 1-4 .
anion radical, the one-electron populated fragment orbital of NO 2 will pump spin density in the ring. Such interaction should be very advantageous due to the level proximity of the . lowest vacant ring orbital and the highest occupied orbital of NO 2 . Therefore, the nitro group can, in fact, act as -donor in the nitrobenzene anion radical. Such prediction is not self-evident, since the nitro group in neutral aromatic nitro compounds is recognized as a strong -acceptor and, in principle, even as a reservoir for four to six additional electrons. Comparing half-wave potentials of reversible one-electron reduction of meta-dinitrobenzene and other benzene derivatives, one can determine the Hammett constant for the nitro . group in the nitrobenzene anion radical. When the NO2 group is transformed into the NO 2 group, a change in both the sign and the value of the correlation constant is observed (Todres, Pozdeeva et al. 1972a,b). A formal comparison of the Hammett constants for the NO2, . . NO 2 , and NH2 groups shows NO 2 to be close to NH2 in terms of donating ability: . m(NO2) 0.71, m(NO 2 ) 0.17, m(NH2) 0.16. Having captured the single electron, the nitro group then acts as a negatively charged substituent. Similarly, the stable anion radical resulting from aryl diazocyanides . [(ArNBNCN)] contains a substituent [(NBNCN) ] that interacts with the aryl ring as a donor (Kachkurova et al. 1987). Using other nitro derivatives of an aromatic heterocyclic . series, the generality and statistical relevance of the observed m(NO 2 ) constant were established (Todres, Zhdanov et al. 1968; Todres, Pozdeeva et al. 1972a). The sign and absolute magnitude of the Hammett constant are invariant regardless of which cation (K, Na, or Alk4N) in the anion radical salts of nitro compounds was studied. Such invariance is caused by the linear dependence between electrochemical reduction potentials of substituted nitrobenzenes and the contribution of the lowest vacant *-orbital of the nitro group to the -orbital of this anion radical, which is occupied by the single electron (Koptyug et al. 1988). Experiments on the reactivity of the anion radicals under consideration are set out later. The ability of nitrobenzene anion radicals to undergo coupling with benzenediazo cations has been studied (Todres, Hovsepyan et al. 1988). This reaction is known to proceed for aromatic compounds having donor-type substituents (NH2, OH). Aromatic compounds containing only the nitro group do not participate in azo-coupling. It is also worth noting that benzenediazo cations are strong electron acceptors. For instance, the interaction between benzene- or substituted benzene-diazonium fluoroborates and the sodium salt of the naphthalene anion radical results in electron transfer only (Singh et al. 1977). The products were naphthalene (from its anion radical) and benzene or its derivatives (from ben-
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zene- or substituted benzene-diazonium fluoroborates). The potassium nitrobenzene anion radical appears to react with diazonium cations according to the electron transfer scheme, too (as the naphthalene anion radical does). Products of azo-coupling were not found (Todres, Hovsepyan et al. 1988). The potassium salt of the phthalodinitrile (ortho-dicyanobenzene) anion radical also reacts with an electrophile according to the electron transfer scheme. If the electrophile is tert-butyl halide, the reaction proceeds via the mechanism, including at the first-stage dissociative electron transfer from the anion radical to alkyl halide, followed by recombination of the generated tertiary butyl radical with another molecule of the phthalodinitrile anion radical. The product mixture resulting in the reaction includes 4-tert-butyl-1,2-dicyanobenzene, 2-tert-bytylbenzonitrile, and 2,5-di(tert-butyl)benzonitrile (Panteleeva and co-authors 1998). An attempt to detain an unpaired electron was made by means of the second nitro group in the anion radical of nitrobenzene (Todres, Hovsepyan et al. 1988). The potassium salt of the anion radical of ortho-dinitrobenzene did yield an azo-coupled product according to Scheme 1-5 (the nitrogen oxide evolved was detected). The reaction leads to a para-substituted product, entirely in accordance with the calculated distribution of spin density in the anion radical of ortho-dinitrobenzene (Todres 1990). It was established, by means of labeled-atom experiments and analysis of the gas produced, that azo-coupling is accompanied by conversion of one of the nitro groups into the hydroxy group and liberation of nitric oxide. In other words, the initial radical product of azo-coupling is stabilizing by elimination of the small nitrogen monoxide radical to give the stable nonradical final product (Todres, Hovsepyan et al. 1988), Scheme 1-5.
SCHEME 1-5
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During transformation of the parent molecule to the corresponding anion radical, changes in substituent effects appear to be possible not only for the nitro group, but also for other substituents. We have just observed the opportunity to use the nitro group as a donor (not an acceptor) in the anion radicals of aromatic nitro compounds. In the case of AlkO and AlkS substituents, we have a chance to encounter the donor-to-acceptor transformation of the thioalkyl group after one-electron capture by thioalkylbenzenes (Bernardi et al. 1979). Both groups, AlkO and AlkS, are commonly known as electron donors. However, in the anion radical form, these groups exert nonidentical effects. The methoxy group keeps its donor properties, while the methylthio group appears to be an acceptor. This is evident from a comparison of the ESR spectra of the nitrobenzene anion radical and its derivatives, in particular the MeO- and MeS-substituted ones. The introduction of substituents into nitrobenzene in general affects the value of aN arising from the splitting of an unpaired electron by the nitrogen atom in the anion radical. If the group introduced is a donor, the aN(NO2) value increases; if it is an acceptor, the aN(NO2) value is reduced. As follows from such a comparison of aN constants, the MeO and EtO groups act similarly to the Me and S groups (donors). At the same time, the MeS and EtS groups act similarly to CN, SO2Me, and SO2Et groups (acceptors) (Ioffe et al. 1970, Alberti, Martelli, & Pedulli 1977, Alberti, Guerra et al. 1979, Bernardi et al. 1979). The sharp contrast between the electronic effects exerted by the oxyalkyl and thioalkyl groups in anion radicals was explained by means of group orbital-energy diagrams. The usual mechanism involving n,-conjugation requires the MeO or MeS group to be situated in the same plane as the aromatic ring of the parent (neutral) molecules. According to calculations (Bernardi et al. 1979), “the most stable conformation is the planar” for the anion radical of anisol. In the case of the anion radical of thioanisol, however, “the preferred conformation is orthogonal.” The planar conformation is stabilized by the usual n,-conjugation between the ring and oxygen or sulfur. Such n,-conjugation is impossible in the orthogonal arrangement, and only the -electrons of the sulfur or oxygen appear to be involved. Only the -orbitals of these atoms are symmetrically available for overlapping with the aromatic -orbitals when fragments of the molecule are oriented perpendicularly. However, interaction between the -electrons of the nucleus and the vacant -orbitals of the substituent is also possible in principle, because this interaction is symmetrically allowed. In practice, , and *, interactions are not too important in the case of uncharged molecules, since the gap between the aromatic -orbitals and /*-orbitals of the substituents is too wide. This is obvious from the left part of Scheme 1-6. Conversion of a neutral molecule into an anion radical leads to occupation of the lowest-energy vacant orbital. The latter is the -orbital of the benzene ring in both anisole and thioanisole. Charge transfer is possible only by means of an interaction between vacant and occupied orbitals and only if an energy gap between them is not too wide. As the *-orbitals of the anisole MeO group are too far away from the ring -orbital occupied by the single electron, the conjugation conditions in the anion radical compared with the neutral molecule remain unchanged. This is evident from the right part of Scheme 1-6. The thioanisole MeS group differs from the anisole MeO group in the fact that the * is at a lower energy level (Alberti, Guerra et al. 1979). In this case, population of the lowest vacant aromatic -orbital by a single electron changes the conjugation conditions. Namely, the *, interaction becomes more favorable than the n, interaction, because the energy gap between the *- and the -orbitals is narrower. In other words, conditions created in the anion radical promote charge transfer from the ring to the substituent rather than from the substituent to the ring, as is the case in the neutral molecule. That is why the
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SCHEME 1-6
orthogonal conformation is stabilized instead of the planar one. The conversion of thioanisole into the anion radical causes the change in the orientation of the MeS group relative to the aromatic ring plane. This is depicted on the right part of Scheme 1-6. Once again, not the energy levels but the relative energy gaps remain unchanged for the anion radicals as compared to the parent neutral molecules. Change in the substituent nature after transformation of neutral aromatics into the corresponding ion radicals may be used intuitively for preparation of some unusual derivatives. One may transform an organic molecule into its anion radical, change the substituent effect, perform the desired substitution, and, after that, take a surplus electron off by means of a soft oxidant and eventually obtain the desired unusual derivative in its stable form. Studies along this line are intriguing in the cases of both anion and cation radicals. 1.2.2 Connection Between Ion Radical Reactivity and Electron Structure of Ion Radical Products The reaction of aryl and hetaryl halides with the nitrile-stabilized carbanions (RC–CN) leads to derivatives of the type ArCH(R)CN. Sometimes, however, dimeric products of the type ArCH(R)CH(R)Ar are formed (Moon et al. 1983). As observed, 1-naphthyl, 2pyridyl, and 2-quinolyl halides give the nitrile-substituted products, while phenyl halides, as a rule, form dimers. The reason consists of the manner of a surplus electron localization in the anion radical that arises upon replacing halogen with the nitrile-containing carbanion. If the resultant anion radical contains an unpaired electron within LUMO covering mainly the aromatic ring, such an anion radical is stable, with no inclination to split up. It is oxidized by the initial substrate and gives the final product in the neutral form, Scheme 1-7: .
[Ar] CH(R)CN e → ArCH(R)CN If the anion radical formed acquires an unpaired electron on the CN group orbital, this group easily splits off in the form of the cyanide ion. So the dimer is formed as the final
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product, Scheme 1-8: .
2PhCH(R)[CN] → 2CN PhCH(R)CH(R)Ph One-electron reduction of organyl halides often results in the spontaneous elimination of halide and the formation of organyl radicals according to Scheme 1-9: .
.
RX e → RX → R X The organyl radicals resulting in this cleavage can combine with the nucleophile anion, Scheme 1-10: .
.
R Y → RY
The anion radical of this substituted product initiates a chain reaction network, Scheme 1-11: .
.
RY RX → RY RX
According to Saveant (1994), an important contribution to the overall efficiency of . this substitution reaction is given by the step in which the RY anion radical is formed. In this step, an intramolecular electron-transfer/bond-forming process occurs when the nu. cleophile Y attacking the radical R begins to form the new species, characterized by an elongated two-center three-electron CY bond. An unpaired electron in this anion radical is at first allocated on a “low-energy” C-Y* molecular orbital. With the progress of the formation of the C–Y bond, the energy of the * molecular orbital increases sharply until a changeover occurs. If R is Ar, the * molecular orbital of the molecule becomes the LUMO. An internal transfer of the odd electron to the LUMO then takes place. So it follows that the substitution under consideration will be easier when the energy of the * . molecular orbital available in the ArY species is lower. Papers by Rossi et al. (1994), Galli with co-workers (1995), and Borosky et al. (2000) have again underlined the rule: The lower the energy of the LUMO of the RY (ArY) species, the easier (faster) the reaction . . between R (Ar ) and Y. It is worth noting, however, that primary halide-containing anion radicals may be somewhat stable if an aromatic molecule has another electron acceptor group as a substituent— such as the nitro, cyano (Lawless et al. 1969), carbonyl (Bartak et al. 1973), or pyridinyl group (Neta & Behar 1981). In these cases, dehalogenation reactions proceed as intramolecular . . electron transfers from the groups NO through the conjugated system to the car2 , CN bon–halogen fragment orbital. After that, the halide ion is eliminated. The splitting rate depends on the nature of the halogen (I Br CI) and on the position of the halogen with respect to another substituent (ortho para meta) (Alwair & Grimshaw 1973; Neta & Behar 1981; Behar & Neta 1981; Galli 1988). The cleavage proceeds more easily at those positions that bear maximal spin density. Change of the nitro group to the nitrile or carbomethyl group leads to some facilitation of halogen elimination: A greater portion of spin density reaches the carbon–halogen orbital and the rate of dehalogenation increases. For instance, the anion radical of 4-fluoronitrobenzene is characterized by the aF HFC constant of 0.855 mT and high stability (Starichenko et al. 1981). In contrast, the anion radical of 4-fluorobenzonitrile has a significantly larger aF HFC constant of 2,296 mT and does readily dissociate into the benzonitrile -radical and the fluoride ion (Buick et al. 1969). It should be emphasized that the cause of halide mobility in aromatic anion radicals is quite opposite to that in heterolytic aromatic substitution at the carbon–halogen bond. In anion radicals, the carbon–halogen bond is enriched with electron density and, after halide
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ion expulsion, an aromatic -radical is formed. In neutral molecules, the carbon–halogen bond conjugated with an acceptor group becomes poor with respect to its electron density; a nucleophile attacks a carbon atom bearing a partial positive charge. Some kind of -binding was established between the nitro group and chlorine through the benzene ring in 4-nitrochlorobenzene (Geer & Byker 1982). As a result, the inductive effect of chlorine becomes suppressed in the neutral molecule. In the anion radical, LUMO populated by one electron comes into operation. The HOMO role turns out to be insignificant. In anion radicals, this orbital can cause only a slight disturbance. The negative charge to a significant degree moves into the benzene ring, and this movement is enforced at the expense of the chlorine-inductive effect. The carbon–chlorine bond is enriched with an electron. Eventually, Cl leaves the anion radical species. This event is quite simple, and its simplicity is based on the -electron character of HOMO and LUMO. However, there are some cases when an unpaired electron is localized not on the -orbital but on the -orbital of an anion radical. Of course, in such cases a simple molecular-orbital consideration based on the approach does not coincide with experimental data. Chlorobenzothiadiazole may serve as a representative example (Gul’maliev et al. 1975). Although the thiadiazole ring is a weaker acceptor than the nitro group, the elimination of the chloride ion from the 5-chlorobenzothiadiazole anion radical does not take place (Solodovnikov & Todres 1968). At the same time, the anion radical of 7-chloroquinoline loses the chlorine anion (Fujinaga et al. 1968). Notably, 7-chloroquinoline is very close to 5-chlorobenzothiadiazole in terms of structure and electrophilicity of the heterocycle. To explain this difference, calculations are needed that can clearly take into account the -electron framework of the molecules compared. Solvation of intermediate states on the way to a final product should be involved in the calculations as well (Parker 1981). The alkyl halide anion radicals do not have -systems entirely. Nevertheless, they are able to exist in solutions. The potential barrier for the C–Cl cleavage is estimated to be ca. 70 kJmol1 (Abeywickrema & Della 1981; Eberson 1982). The carbon–halogen bond may capture one electron directly (Casado et al. 1987; Boorshtein & Gherman 1988). It is interesting to compare SCl and SCN in relation to NO2 as the reference group. Aryl sulfenyl chlorides and thiocyanates were subjected to two independent model reductive cleavage reactions by treatment with (a) cyclooctatetraene dipotassium (C8H8K2) in tetrahydrofurane (THF) or (b) HSiCl3 R3N (R alkyl) in benzene (Todres & Avagyan 1972, 1978). As shown, aromatic sulfenyl chlorides under conditions (a) and (b) produce disulfides or thiols; the presence of the nitro group in the ring does not affect the reaction. Aryl thiocyanates without the nitro group behave in a similar way. However, aryl thiocyanates containing the nitro group in the ring are converted into anion radicals, with the SCN remaining unchanged. This pathway is represented by Scheme 1-12: .
O2NC6H4SCN 1⁄2C8H8K2 → 1⁄2C8H8 NCSC6H4NO2 K Splitting of the SCN group is not observed and, after the one-electron oxidation, the . initial NCSC6H4NO 2 anion radical produces NCSC6H4NO2. The recoveries are close to quantitative; disulfides and thiols are not observed. The thiocyanate group (SCN) thus competes less successfully with the nitro group (NO2) for the extra electron than the sulfenyl chloride group (SCl). The conclusion just outlined was entirely confirmed by quantum chemical calculations. The results of the calculations are shown on the Table 1-2 (LCAO MO CNDO/2 approach, Todres, Tomilin, & Stankevich 1982). As seen in the table, the SCl group charge depends slightly on whether or not the NO2 group is present in the benzene ring. In the case
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TABLE 1-2 Effective Charges (qi ) on R and SX Groups in p-RC6H4SX Anion Radicalsa R
SX
qR
qsx
H NO2 H NO2
SCI SCI SCN SCN
0.05 0.20 0.06 0.48
0.78 0.76 0.54 0.27
a
The rest of the charge, up to unity, is in the benzene rings.
of thiocyanate anion radicals, the charge on the SCN group is diminished by 50% if the NO2 group is present in the molecule. Thus, SCl and SCN have different electron-attraction properties. That conclusion was not predictable a priori. Until recently, the extent of polarization in the SCl group has been considered to be comparable to that in the SCN group, according to the S –X scheme. For instance, Kharash et al. (1953), have pointed out that nitroaryl thiocyanates as well as nitroaryl sulfenyl chlorides, when dissolved in concentrated sulfuric acid, are converted into the same nitroaryl sulfenium ions, O2NArS. However, the preceding findings indicate otherwise. Other pertinent examples include splitting of the anion radicals from p-nitrophenyl metyl sulfone and p-cyanophenyl methyl sulfone (Pilard et al. 2001). The nitrophenyl species undergoes the preferential cleavage of the Ar–S bond, whereas the cyanophenyl species expels both CN and CH3SO2 groups in the two parallel cleavage reactions. All the examples show that foreseeing a splitting direction and extending it from one parent compound to another is risky, especially in the organic chemistry of radical ions. 1.2.3 Bridge Effect Peculiarities Let us consider ion radicals of paracyclophanes. As the basis for our consideration, we will choose the following species, depicted in Scheme 1-13 (a)–(e). (a) The anion radical of pseudogeminal-[2.2]paracyclophane-4,7,12,15-tetrone, in which the 1,4-benzoquinone units lie one underneath the other (b) The anion radical of syn-[2.2](1,4)naphthalenophane-4,7,14,17-tetrone, in which there is the same spatial situation (c) The anion radical of anti-[2.2](1,4)naphtalenophane-4,7,12,15-tetrone, in which the naphthoquinone units are further apart (d) The cation radical of syn-[2.2](1,4)naphthalenophane-4,714,17-tetramethoxy derivative, which is close to case (b) (e) The cation radical of anti-[2.2](1,4)naphthalenophane-4,7,14,17-tetramethoxy derivative, which resembles the case (c) structure As seen, cyclophane structures (b)–(e) have the unique feature that the through-bond distance within the paracyclophane fragment is held constant while the spatial distance between the ion radicalized and neutral moieties is changed. So the relative importance of through-bond and through-space mechanisms for electron transfer can be learned directly from experimental data on these molecules. All the depicted compounds (Scheme 1-13) were studied in electrochemical reduction [cases (a) to (c)] or oxidation [cases (d) and (e)]; two one-electron peak potentials were re-
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vealed, with differences significantly higher than 20 mV (Wartini, Valenzuela et al. 1998; Wartini, Staab et al. 1998). A large difference between the first two reductions or oxidation potentials is indicative of the delocalization of the first (unpaired) electron (Rak & Miller 1992). In other words, the two electrochemically active fragments can accept or lose a single electron, the second one-electron transfer being markedly hampered. In their turns, ESR and ENDOR spectra of the anion radicals under investigation gave evidence of delocalization of an unpaired electron over the whole molecule in each case. Because of the close spatial contact of the quinone units [0.31 nm between the centers of the 1,4-benzoquinone rings, Scheme 1-13(a)], one may suppose that the unpaired electron simply jumps over this narrow gap. If so, the whole-molecule delocalization would be impossible in the case of the mutual anti-arranged 1,4-naphthoquinone units [see structure (c) in Scheme 1-13]. However, this anti-arranged anion radical shows the full spin electron delocalization. Consequently, ,-conjugation is realized in the anion radicals of the paracyclophanes considered. In the same way, the displacement of the unpaired electron over the whole molecules was observed for (d) and (e) cation radicals from Scheme 1-13, in which 1,4-dimethoxynaphthalene units are syn- or anti-annelated to [2.2]paracyclophane. Taken together, the experimental results considered provide direct evidence for the through-bond mechanism of electron transfer in these paracyclophane systems. In a recent study, the electron transfer between 1,4-dimethoxybenzene and 7,7-dicyanobenzoquinone methide moieties in synor anti-cyclophane systems reached the same conclusion: The through-bond mechanism
SCHEME 1-13
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can remain the dominant reaction pathway at short donor–acceptor distances as well (Pullen et al. 1997). Other examples of the specified bridge effect deal with anion radicals of aryl derivatives of the tri-coordinated boron or tri- and four-coordinated phosphorus. In the tris(pentafluorophenyl)boran anion radical, spin density is effectively transferred from the boron p-orbital to an antibonding -molecular orbital of the phenyl rings (Kwaan et al. 2001). Studies of the phosphorus-containing aromatic anion radicals have been more informative. As known, phosphorus interrupts conjugation between aryl fragments in the corresponding neutral compounds. In contrast to the neutral molecules, the phosphorus atom does transmit conjugation in Ar3P and Ar3P(O) anion radicals. At least formally, the P atom appears to be a bridge, not a barrier. Spin delocalization takes place along the whole molecule in each case. Perhaps, the phosphorus unfilled p- or d-orbitals take part in this transmission effect (Il’yasov et al. 1980). The cation radicals of triaryl phosphine and trialkyl phosphites have also been probed to hyperconjugation; photoelectron, ESR spectra, and reactions with aromatic radicals were studied. As it turned out, the cation radicals of Ar3P are characterized by strong hypercon. jugation (Egorochkin and others 2001). The cation radical [P(OMe/OEt)3] contains an unpaired electron predominantly on the phosphorus atom, which facilitates the radical . coupling of this cation radical with the Ar radical from aryl diazonium salts. In [PhP(OMe/ . . . OEt)2] , and [Ph2P(OMe/OEt)] , and [Ph3P] cation radicals, an unpaired electron is increasingly shifted from the phosphorus atom to the phenyl ring(s). This reduces the spin density at the central phosphorus atoms, making the reaction of the mentioned cation radi. cals with Ar slower and eventually preventing it altogether (Yasui, Fujii et al. 1994; Yasui, Shioji, Ohno 1994). Similarly to the cation radicals, the phosphoranyl radicals with and without the aryl ligand(s) exhibit small and large values of phosphorus HFC constants, respectively, in the ESR spectra (Boekenstein et al. 1974; Davies, Parrot, & Roberts 1974; Davies, Griller, & Roberts 1976). This means that the unpaired electron is located mainly on the aryl ligands of aryl phosphoranyl radicals or entirely on the central phosphorus in the case of alkyl phosphoranyl radicals. Let us direct our attention now to the PBC bond in phosphaalkene ion radicals. The literature contains data on two such anion radicals in which a furan and a thiophene ring are bound to the carbon atom, and the 2,4,6-tri(tert-butyl)phenyl group is bound to the phosphorus atom. According to ESR spectra of the anion radicals, an unpaired electron is delocalized on a *-orbital built from the five-membered ring (furanyl or thienyl) and the PBC bond. The participation of the phosphaalkene moiety in this molecular orbital was estimated as about 60%. Although the cumbersome tris(tert-butyl)phenyl group is led out of the conjugation sterically, some moderate (but sufficient enough) transmission of the spin density does take place through the PBC bridge (Jouaiti et al. 1997). Scheme 1-14 depicts the structures under discussion. The same situation was revealed for the case of the anion radical of the phosphaalkene containing the phenyl ring linked to the carbon atom and the 2,4,6-tri(tertbutyl)phenyl group linked to the phosphorus atom of the PBC bond. The unpaired electron is delocalized in this anion radical on both the PBC bond and the phenyl ring (Geoffroy et al. 1992). In the benzene and furane derivatives containing two conjugated ArPBC bonds, an unpaired electron is delocalized on the whole bonds of the anion radicals, resulting in one-electron reduction. Neither localized structures nor diradical species were observed (Al Badri et al. 1999). However, if the carbon atom of this PBC bond is an integral part of the cyclopentadiene ring, the unpaired electron distribution proceeds ac-
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SCHEME 1-14
cording to a scheme of spin-charge scattering (Al Badri et al. 1997). Scheme 1-15 illustrates this special case. Hence, the possibility of acquiring aromaticity conferred by the presence of six -electrons in the five-carbon-membered ring considerably increases the electron affinity of this ring. As a consequence, one of the two -electrons of the PBC bond remains on the phosphorus atom, and another one combines with the excess electron to create the cyclopentadienyl -electron sextet. The situation is analogous to that in the diphenylfulvene anion radical, as analyzed in Chapter 3 (see Section 3.2.2). In the case of p-phosphaquinone, the 2,4,6-tri(tert-butyl)phenyl group is bound to the phosphorus atom, and the carbon atom of the PBC bridge is an integral part of the 3,5di(tert-butyl)-4-oxocyclohexa-2,5-dien-1-ylidene moiety. The corresponding anion radical (a structure of the p-benzosemiquinone type) is very similar to diarylphosphoranyl radicals, in the sense of the unpaired electron delocalization (Sasaki et al. 1999). In the phosphoric analog of biphenyl, namely, 4,4 ,5,5 -tetramethyl-2,2 -phosphinine, the corresponding anion radical was generated on a potassium mirror. ESR/ENDOR spectra and density functional theory calculations show that, in contrast with the neutral species, this anion radical is planar and that the unpaired electron is delocalized mainly on the phosphorus–carbon–carbon–phosphorus fragment of the two linked six-membered rings. The phosphorus p-orbitals play a large part in this delocalization (Choua et al. 2000). Diphosphaallene derivatives ArPBCBPAr are peculiar compounds because of the presence of the two orthogonal carbon–phosphorus double bonds. Those compounds were transformed into cation radicals upon electrochemical or chemical one-electron oxidation. As found, the unpaired electron is located on a molecular orbital constituted mainly of a p-orbital of each phosphorus atom and a p-orbital of the carbon atom (Chentit et al. 1997; Alberti, Benaglia et al. 1999). Upon electrochemical or chemical reduction, aromatic phosphaallene derivatives yield anion radicals. These species have two equivalent phosphorus
SCHEME 1-15
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nuclei. The unpaired electron oscillates between the two phosphorus atoms according to Scheme 1-16 (Sidorenkova et al. 1998, Alberti, Benaglia et al. 1999): .
.
ArPBC()–P( )Ar ↔ ArP( )–C()BPAr Consequently, the electron structures of the diphosphaallene ion radicals resemble those of the allene ion radicals, where the allenic fragment works as a bridge for conjugation (see Section 3.3.2). One quite unusual effect deserves to be described as applied to the CBC bridge connecting two phenyl rings in the 4-nitrostilbene anion radical. This bridge is good at transmission of conjugation in neutral stilbenes, but in stilbene anion radicals it can operate as a hollow on the conjugation route. At first glance, the distribution of unpaired electrons in the nitrostilbene anion radical has to be similar to that in the nitrobenzene anion radical. The consensus of opinion is that the LUMO of neutral aromatic nitro derivatives (the orbital that accommodates the introduced electron) is essentially an orbital of the “free” nitro group. The styryl fragment of neutral 4-nitrostilbene is conjugated with the nitro group and acts as a weak donor. This is indicated by values of the Hammett constants: (NO2C6H4–) is 0.23 and (–CHBCHPh) is 0.07. If the styryl substituent retained its donor nature in the anion radical state, an increase, not a decrease, in the value of the nitrogen HFC constant [a(N)] would have been observed. Experiments show that a(N) values for anion radicals of nitrostilbenes decrease (not increase) in comparison with the a(N) value for the anion radical of nitrobenzene (Todres 1992). Both “naked” anion radicals and anion radicals involved in complexes with the potassium cations obey such regularity. In cases of potassium complexes with THF as a solvent, a(N) 0.980 mT for PhNO2 anion radical and a(N) 0.890 mT for PhCBCHC6H4NO2-4 anion radical. In the presence of 18-crown-6-ether as a decomplexing agent in THF, a(N) 0.848 mT for PhNO2 anion radical and a(N) 0.680 mT for PhCHBCHC6H4NO2-2 anion radical. Reduction of nitrobenzene (Grant & Streitwieser 1978; Todres, Dyusengaliev et al. 1984) and 4-methoxynitrobenzene (Todres, Dyusengaliev et al. 1984) by uranium-, thorium-, and lanthanum-di(cyclooctatetraene) complexes leads to azo compounds. Scheme 117 illustrates these reductive reactions using uranium–(C8H8)2 complex as an example. Under the same conditions, 4-styryl nitrobenzene (4-nitrostilbene) undergoes cis-totrans isomerization only, with no changes in the nitro group (Todres, Dyusengalier et al. 1984, 1985), Scheme 1-18. Thus, it appears that the focal point of the reaction has transferred. The presence of a styryl (not a methoxyl) group protects the nitro group from reduction. For some reason or other, the styryl group causes a shift of excess electron density from the nitro to the ethylene fragment.
SCHEME 1-17
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SCHEME 1-18
This subtle difference between the anion radicals of nitrobenzene and nitrostilbene, observed experimentally, is well reproduced by quantum chemical calculations (Todres et al. 1984). Single-electron wave function analysis of the vacant orbitals in both molecules shows that one-electron reduction of cis-4-nitrostilbene must be accompanied by the predominant localization of an upaired electron in the region of the ethylene moiety. Participation of the nitro-group atomic orbitals appears to be insignificant. The nitro-group atomic coefficients in the molecular wave function for cis-4-nitrostilbene are half of those for nitrobenzene. The excess electron population (q) of the first vacant orbital for the nitro groups is 0.3832 for the nitrobenzene anion radical and 0.0764 for the nitrostilbene anion radical. The unpaired electron is localized largely on the ethylene fragment of the nitrostilbene skeleton (q 0.2629). Moreover, the first vacant level of the cis-4-nitrostilbene molecule has lower energy than that of the nitrobenzene molecule: 38 and 135 kJ, respectively. This means that 4-nitrostilbene is a more effective electron acceptor than nitrobenzene. This theoretical conclusion is verified by experiments. The charge-transfer N,Ndimethylaniline complexes formed by nitrobenzene or 4-nitrostilbene have stability constants of 0.085 Lmol1 and 0.296 Lmol1, respectively. Moreover, the formation of the charge-transfer complex between cis-4-nitrostilbene and N,N-dimethylaniline indeed results in cis-to-trans conversion (Dyusengaliev et al. 1995). This conversion proceeds slowly in the complex but runs rapidly via the nitrostilbene anion radical (Todres 1992). The cis–trans conversion of ion radicals will be considered in more details later (see Sections 3.2.5 and 7.2.1). It is interesting to compare the fate of the CBC bridge in the anion radicals of 4-nitrostilbene (see earlier) and 4-acetyl- , -diphenylstilbene (Wolf et al. 1996). When treated with potassium or sodium in THF and then with water, neutral 4-nitrostibene does not undergo a many-electron reduction of the nitro group or the CBC bridge. Under the same conditions, 4-acetyl- , -diphenylstilbene produces a pinacol, [Ph2CBC(Ph)C6H4C (OH)CH3]2. As calculations show, the carbonyl of the acetyl group in 4-acetyl- , diphenylstilbene is the site of significant reduction. The formal charge on the carbonyl carbon and oxygen atoms became significantly more negative upon addition of one electron, whereas the olefinic carbons become only slightly more negative (Wolf et al. 1996). It is worth pointing out that the phenylcarbonyl group is a stronger acceptor than the nitrophenyl group: (—COC6H5) 0.46 [ (—COCH3) 0.52], whereas (NO2C6H4–) 0.23 [ (—NO2) 0.78]. In addition, the phenylacetyl group in the molecule under consideration is conjugated only slightly, if at all, with the CBC bridge, because this molecule is propeller shaped (Hoekstra & Vos 1975). Although the acetyl group lies in the plane of the phenyl ring attached to it, it remains separated from the CBC bridge. In contrast, the nitro and ethylenic fragments in trans-4-nitrostilbene form the united conjugation system. Such conjugation is a necessary condition for the whole-contour delocalization of an unpaired electron in arylethylene anion radicals. Whether this condition is the only one or there is some interval of allowable strength for the acceptor is a question left to future experiments.
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1.3 ACID–BASE PROPERTIES OF ORGANIC ION RADICALS 1.3.1 Anion Radicals Let us compare anion radicals with dianions, which are definitely strong bases. For example, the cyclooctatetraene dianion (C8H 2 8 ) accepts protons even from such solvents as dimethylsulfoxide and N,N-dimethylformamide. The latter are traditionally qualified as aprotic solvents. In these solvents, the cyclooctatetraene dianion undergoes protonation, resulting in the formation of cyclooctatrienes (Allendoerfer & Riger 1965); see Scheme 1-19: C8H 82 2H → C8H10 Hence, such dianions with two introduced electrons are essentially the counterpart or aprotic equivalent, or base, of the corresponding C–H acid. 1.3.1.A Anion Radical Basicity Having one excess electron, anion radicals can be considered aprotic “half-equivalents” of the corresponding C–H acid. Reacting with protons, anion radicals display some dual behavior. As bases, they may add protons, get rid of their negative charges, and give radicals. This direction is represented by direction (a) in Scheme 1-20. As radicals with excess electrons, they may generate atomic hydrogens from protons and transform into parental uncharged compounds according to direction (b) in Scheme 1-20. Such dual anion radical reactivity towards protons depends on the difference between proton affinity to an electron and the first ionization potential of an anion radical. This difference may be not very strong. The fate of the competition between directions (a) and (b) in Scheme 1-20 also depends on the relative stability of the reaction products. It is reasonable to illustrate the duality with two extreme examples from real synthetic practice. Direction (a): Alkali metals transform saturated ketones into secondary alcohols. The reaction proceeds in the mixture of ethanol and liquid ammonia in the presence of ammonium chloride as a proton donor and follows Scheme 1-21 (Rautenstrauch et al. 1981).
SCHEME 1-20
SCHEME 1-21
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Controlled one-electron reductions transform 1,2,3,4-tetraphenyl-1,3-cyclopentadiene or 1,2,3,4,5-pentaphenyl-1,3-cyclopentadiene into the 1:2 mixtures of the dihydrogenated products and the corresponding cyclopentadienyl anions in the constant ratio 1:2 (Farnia et al. 1999). The anion radicals initially formed are protonated by the substrates themselves. The latter are thermodynamically very strong acids because of their strong tendency to aromatization. As to the cyclopentadiene anion radicals, they need two protons to give more or less stable cyclopentaenes. Scheme 1-22 represents the initial one-electron electrode reduction of C5HAr5, and Scheme 1-23 explains the ratio and the nature of the products obtained at the expense of the further reaction in the electrolytic pool. C5HAr5 e → (C5HAr5)
(1-22)
.
.
(C5HAr5) C5HAr5 → C5H2Ar 5 (C5Ar5) .
.
C5H2Ar 5 (C5HAr5) → C5H2Ar 5 C5HAr5 C5H2Ar 5 C5HAr5 → C5H3Ar5 (C5Ar5)
On the whole:
(1-23)
.
C5HAr5 2(C5HAr5) → C5H3Ar5 2(C5Ar5)
Hence, the case represents a typical example of the so-called “father–son” self-protonation process, provoked by the partial transformation of cyclopentadiene into the anion radical. Direction (b): The mixture of trichlorosilane and tributylamine in benzene reduces organic derivatives of valence 2 sulfur such as benzene sulfenates. However, nitrobenzene sulfenates remain intact in this system (Todres & Avagyan 1978), Scheme 1-24. Introduction of nitrobenzene sulfenates into the same mixture of trichlorosilane and tributylamine results in the evolution of hydrogen. As proven (Todres & Avagyan 1978), trichlorosilane with tributylamine yields the trichlorosilyl anion and tributylammonium cation. This stage starts the process involving one-electron transfer from the anion to a nitrobenzene sulfenate. At the time, the sulfenate produces the stable anion radical with the
SCHEME 1-24
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SCHEME 1-25
tributylammonium counterion. The latter anion radical gives off an unpaired electron to the proton from the counterion; see Scheme 1-25. Returning to direction (a), it is interesting to compare pKa values of aromatic anion . radicals after protonation, pKa(ArH 2), and pKa values of parental aromatics after protonation, pKa(ArH2). As seen from Table 1-3, if the anion radicals accept a proton, they hold it much more firmly than the parental neutral molecules (pKa values are positive). Contrary to the early indications (Kalsbek & Thorp 1994), the anion radical of C60 . fullerene is a very weak base. Acidity of C60H approaches that of dilute triflic acid; the . solutions in o-dichlorobenzene were compared. In dimethylsulfoxide, the pKa of C60H is estimated to be about 9, making it a slightly weaker acid than p-benzoic acid (see the review by Reed C., & Bolskar 2000). These data are consistent with the reports that the syn. thesis and ESR spectrum of C 60 are relatively insensitive to the presence of water. There . are aryl and methyl derivatives of C 60 that are stable and soluble in water (Sawamura et al. . 2000). The weak basicity of C 60 is due to its intrinsically high stability through delocaliza. tion of the negative charge toward the 50-electron system. When C60H comes up, it formally produces a carbon radical to the site of protonation, and the energetic cost of this localization is high. There is no electrochemical evidence for the reasonable expectation of . dimerization of C60H radicals (Cliffel & Bard 1994). The basicity of anion radicals therefore consists of proton landing. The proton landing assumes 1:1 stoichiometry with respect to an anion radical and a proton-donor
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.
molecule. For example, in the reaction of the naphthalene anion radical (C10H 8C 8 ) with methanol (Scheme 1-26, see below), this 1:1 stoichiometry should result in the formation of 50:50% mixture of naphthalene (C10H8) and dihydronaphthalene (C10H10). .
.
C10H 8 H → C10H 9 .
.
C10H 9 C10H 8 → C10H 9 C10H8 C10H 9 H → C10H10
On the whole:
.
2C10H8 2H → C10H8 C10H10
Surprisingly, Screttas and co-authors (1996) have found that the reaction of the lithium naphthalene anion radical with methanol in THF follows the 2:1 stoichiometry and leads to the C10H8–C10H10 mixture in a 95:5 ratio. The authors proposed an alternative, Scheme 1-27: .
.
C10H 8 Li MeOH → C10H 9 MeOLi . . C10H 9 C10H 8 Li → C10H 9 Li C10H8 C10H 9 Li → C10H8 LiH C10H 9 Li MeOH → MeOLi C10H10 LiH MeOH → MeOLi H2 The decisive point of the novel scheme is the amortization of C10H 9 Li by the elimination of the metal hydride. The authors admit that the weakness of their scheme is the lack of evidence for the formation of alkali metal hydride and for the formation of H2 from the (supposed) reaction between the protonating agent and the alkali metal hydride. However, the main sense of this scheme consists of its better agreement with the observed stoichiometry. As to the first step of Scheme 1-27 (the proton landing), it can certainly have a more intimate mechanism, say, electron transfer from the anion radical to the protonating
TABLE 1-3 Differences in pKa Values Between Aromatic Compounds and Their Anion Radicals ArH
pKa (ArH 2)
PKa (ArH 2)
pKa
Benzene NO2–Benzene Benzonitrile p-Me–Benzoate Benzaldehyde Fluorenone Benzophenone Acetophenone Benzamide Benzoate
23 11.3 10.4 7.8 7.1 6.6 6.2 6.2 1.9 4.2
13.5 3.2 7.3 5.5 8.4–10.5 6.3 9.2 9.9 7.7 12.0
36 14.5 17.7 13.3 15.5–17.6 12.9 15.4 16.1 9.6 7.8
Source: Bil’kis & Shteingarts 1987.
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SCHEME 1-29
agent, followed by hydrogen atom attack on the neutral naphthalene and production of dihydronaphthyl radical intermediate, according to Scheme 1-28: .
.
C10H 8 Li MeOH ⇔ [C10H 8 Li , MeOH] → .
.
[C10H8, LiMeO, H ] → MeOLi C10H 9 1.3.1.A Pathways of Hydrogen Detachment from Anion Radicals As a rule, the addition of an extra electron to a parent organic molecule leads to significant weakening of bonds in a forming anion radical, and thereby bond breaking is facilitated. According to Zhao, Y., & Bordwell (1996a, p. 2530), there are three feasible different . pathways of hydrogen detachment from anion radicals (AH ). These three pathways are compared in Scheme 1-29. Path (a) can be demonstrated with examples of anion radicals of amino and hydroxy derivatives of 2,1,3-benzothiadiazole (Asfandiarov et al. 1998) and . the azafullerene anion radical, C59HN (Keshavarz et al. 1996). Scheme 1-30 describes . one of these examples, hydrogen atom detachment from C59HN . .
.
C59HN → C59N H
Another example of the path (a) of Scheme 1-29 is the anion radical derived from fluorene. It undergoes a first-order decay to give the conjugate base (the fluorenide anion) and a hydrogen atom (Casson & Tabner 1969), according to Scheme 1-31. Pathway (a) of Scheme 1-29 has some kinetic preference, since it can be linked with the strongly exothermic dimerization of the hydrogen atoms formed by this pathway (Zhang & Bordwell 1992). As to case (b) of Scheme 1-29, hydride loss from organic anion radicals is generally not as favorable as the hydrogen atom loss because the solvation of the hydride ion and the organic anions is similar. Generally, path (a) is favored over path (b) in a wide set of organic anion radicals. Free energies of bond dissociation for the anion radicals to give a hydride ion and a radical by path (b) are the highest-energy pathway (Zhao, Bordwell 1996b, p. 6623).
SCHEME 1-31
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SCHEME 1-32
In Scheme 1-29, path (c) is sometimes a feasible process. Thus, for the anion radical derived from 4-nitrobenzyl cyanide, path (c) is favored over path (a) and path (b) by 84 and 150 kJmol1, respectively. Typical examples of anion radical deprotonation are the reactions in Scheme 1-32 (Zhao & Bordwell 1996a, p. 2530). . It is path (c) of Scheme 1-29 that describes the acidity of anion radicals, pKa(RH ). Table 1-4 compares the pKa values of nitro-substituted aromatic weak acids (RH) and their . anion radicals (RH ) in DMSO. Dianion radicals formed in path (c) of Scheme 1-29 are inherently unstable species because they bear a double negative charge as well as an odd electron. The nitro group can exert an unusually stabilizing effect since it can fasten both a negative charge and an added electron. As seen in Table 1-4, the anion radicals are less acidic than their parent comTABLE 1-4 Differences in pKa Values Between Aromatic Weak Acids .
(RH) and Their Anion Radicals (RH ) RH
pKa (RH)
pKa . (RH )
pKa
4-NO2C6H4OH 3-NO2C6H4OH 4-NO2C6H4SH 4-NO2C0H4COOH 4-NO2C6H4NHPh 4-NO2C6H4CH2CN 3-NO2C6H4CH2CN 2-Nitrofluorene
10.8 14.4 5.5 9.0 16.9 12.3 18.1 17.0
21.0 19.4 15.5 13.6 26.1 23.4 23.9 24.7
10.2 5.0 10.0 4.6 9.2 11.1 5.8 7.7
Source: Zhao & Bordwell 1996a, p. 2530.
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pounds. One can expect such weakening in acidity owing to the combined action of two factors: The negative charge retards the loss of a proton, and the formed dianion radical is inherently unstable. One very unusual case of prototropic isomerization was revealed for anion radicals of 1,4-dihydropyridine derivatives (Gavars et al. 1999). These anion radicals transform into 4,5-dihydropyridine analogs through proton detachment and addition. 1.3.2 Cation Radicals 1.3.2.A Cation Radical Acidity Deprotonation is a typical direction of cation radical reactivity. Cation radicals are usually strong H acids. (For example, the alkane cation radicals pass their protons on the alcohol molecules—Sviridenko et al. 2001). Bases that conjugate to these H acids are radicals, see Scheme 1-33: .
.
RH ⇔ H R
Scheme 1-34 displays two real cases of such reversible deprotonation (Neugebauer et al. 1972). As can be seen in the scheme, the cation radicals transform into radicals that are more or less stable and can be protonated reversibly. If the radicals formed are unstable, they perish before protonation. If the “initial” cation radicals have no hydrogen atoms, their stability appears to be higher. Deprotonation is typical for cation radicals that contain proton-active hydrogen atoms and form, after deprotonation, either quite stable or, the reverse, quite unstable radicals. Formation of radicals having a lower energy than that of the starting cation radicals is obviously favorable for their deprotonation. The cation radicals of toluene and other alkylbenzenes are illustrative examples. As shown (Sehested & Holcman 1978), these cation radicals lose protons even in very acidic aqueous solutions. The deprotonation rate does not, in general, depend on the medium acidity. In acetonitrile (AN), the toluene cation radical has high thermodynamic acidity, its pKa is between 9 and 13 (Nicholas &
SCHEME 1-34
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Arnold 1982). In the same solvent (AN), neutral toluene has pKa 45 (Breslow & Grant 1977). So, one-electron oxidation of toluene causes the acidity to grow by 60 pKa units! One-electron oxidation of toluene results in the formation of a cation radical in which the donor effect of the methyl group stabilizes the unit positive charge. Furthermore, the proton abstraction from this stabilized cation radical leads to the conjugate base, namely, the benzyl radical. This radical also belongs to the type. Hence, there is resonance stabilization in the benzyl radical. This stabilization is greater in the benzyl radical than in the cation radical of toluene. As a result, the proton expulsion appears to be a favorable reaction, and the acid–base equilibrium is shifted to the right. This is the main cause of the acidylation effects that the one-electron oxidation brings. It is interesting to compare the reactions in Scheme 1-35: .
.
(PhCH3) → H PhCH 2 . . (PhCH2CH2Ph) → H PhCH2CH Ph In dimethylsulfoxide, the two starting cation radicals of Scheme 1-35 have pKa values of 20 and 25, respectively (Bordwell & Cheng 1989). It is clear that both species give rise to the stabilized carboradicals after deprotonation. Electron-donating substituents increase the stability of the arene cation radical and render the odd-electron species less acidic; for example, the cation radical of hexamethylbenzene has a pKa value of only 2 in AN (Amatore & Kochi 1991). The cation radical of tris(bicyclopentyl)annelated benzene is not prone to proton loss, due entirely to the spin-charge location “more or less in the aromatic (nodal) phase” (Rathore & Lindeman et al. 1998), Scheme 1-36. The tetrahydropyrazine ring in the 5-methyluracil cation-radical is nonaromatic; spin delocalization in the corresponding methylene radical is impossible. Correspondingly, this cation radical does not expel the proton at all (Rhodes et al. 1988). Zhang, X.-M., & Bordwell (1994) determined N–H and O–H bond dissociation en. ergies of cation radicals, denoted BDE(AH ). They found that the energies are only slightly lower than those of their nitrogen or oxygen conjugate acids. However, when the . proton being detached is bonded to a carbon atom, the BDE(AH ) value is typically lower then the BDE(AH) of the corresponding neutral compound. Medium acidity is not essential for the deprotonation of the toluene cation radical, but medium basicity provokes the abstraction of the cation radical proton. Although this effect is obvious, special experiments were undertaken to define it more accurately (Bernstein 1992 and references therein). The studies show that at least three water molecules are re-
SCHEME 1-36
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quired to fulfill the deprotonation inside the cluster between the toluene cation radical and the water molecules. At the same time, only one molecule of ammonia is enough to achieve the deprotonation inside the corresponding cluster. Theoretical calculations for this and other cation radicals are in accord with the experimental results just mentioned (Camanyes et al. 1996). Cation radicals of weak acids may react either by heterolytic cleavage (loss of a proton to produce a radical) or by homolytic cleavage (loss of a hydrogen to form a carbocation). Alkane cation radicals in liquid hydrocarbon solution undergo ion–molecule reactions, such as proton transfer on a submillisecond scale (Werst et al. 1990). In a polar solvent, heterolytic cleavage leading to proton abstraction is usually facilitated because of the favorable solvation energy of the proton, and cation radicals are ordinarily much more acidic than the corresponding neutral compounds. Table 1-5 combines . acidity constants of organic compounds (AH) and their cation radicals (AH ) calculated for their solutions in dimethylsulfoxide (DMSO, a very polar solvent) at 25°C. The acidity of the phenol-family cation radicals depends on the stability of the corresponding phenoxyl radical, which formed after the proton abstraction according to Scheme 1-37: .
(ArOH) → H ArO
In aqueous solutions, the phenol cation radical has pKa 2 (Dixon & Murphy 1978; Holton & Murphy 1979). Surprisingly, the pKa value of the 2,4,6-triphenylphenol cation radical is equal to 5 only (Land et al. 1961; Land & Porter 1963). A natural question arises: Why does the stabilization of the phenoxyl radical with the shielding phenyl groups result in such a small increase in acidity? One can assume that the space shielding with the phenyl groups not only stabilizes this phenoxyl radical, but to some degree hinders proton removal. In a similar manner, the hydroxy group in 3,5-di(tert-butyl)-4-hydroxybenzylidene malodinitrile is localized in the benzene-ring plane and clamped between two methyl fragments of the tert-butyl group (for X-ray evidence, see Itoh et al. 1978). If there are two phenyl groups in the ortho positions of the phenol cation radical, the leaving hydroxylic proton experiences not only steric hindrance but also stoppage from complexation by the -system(s) of the phenyl ring(s). Although the size of the leaving proton is not significant, steric hindrance for deprotonation can be appreciable. Here’s an additional example: The cumene cation radical prefers to lose a proton from -carbon rather than -carbon. This is ascribed to the stereo electronic effect (Zhao, Ch.-X., et al. 1999). TABLE 1-5 Differences in pKa Values Between Aromatic Compounds and Their Cation Radicals AH
pKa (AH)
pKa (AH)
pKa
Toluene Fluorene Phenylacetonitrile Phenol Thiophenol
43.0 22.6 21.9 18.0 10.3
20.0 17.0 32.0 8.1 12.0
63 39.6 53.9 26.1 22.3
Source: Bordwell & Cheng 1989.
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SCHEME 1-38
Let us now compare the pKa values (in AN) of the cation radicals derived from aniline, N-methylaniline, and N-phenylaniline: 5.5, 4.2, and 1.8, respectively (Jonsson et al. 1996). An additional N-methyl substituent produces only a marginal effect on the pKa of the aniline cation radical. At the same time, the effect of a single N-phenyl substituent in aniline is considerable. The phenyl-group effect on the aniline cation radical acidylation is predictable. Deprotonation of the adenine cation radical is a convenient starting point for discussing the control of this reaction from the point of view of the relative stability of the radical formed. The proton loss is possible from the N(6)–H and N(9)–H bonds according to Scheme 1-38. It is obvious that one-electron oxidation has to touch the lone electron pair at the group NH2. Consequently, the N(6)–H deprotonation is expected. Nevertheless, the real process occurs at the expense of both N(6)–H and N(9)–H bonds (Dias & Vieira 1996). In the N(9)-deprotonated radical, the unpaired electron is delocalized over the entire purine system, incorporating four nitrogen and five carbon atoms; see Scheme 1-39.
SCHEME 1-39
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SCHEME 1-40
In the N(6)-deprotonated radical, an unpaired electron can be delocalized over only . three nitrogen atoms. These are one nitrogen of the group NH and two nitrogens of the sixmembered ring. An unpaired electron can also be delocalized over two carbon atoms (one by one of the six- and five-membered rings). This part of delocalization is depicted in Scheme 1-40. So the N(9) radical should be more stable than the N(6) one. That is why both radicals coexist in the system and both N(9) and N(6) deprotonations take place. A potential problem when discussing the pKa of alkylamine cation radicals is that the deprotonation can occur on both the nitrogen and the -carbon, as in Scheme 1-41: .
(Et2NH) ⇔ H Et2N
.
.
.
(Et2NH) ⇔ H EtNCH 2Me In the latter case, the aminomethylene radical is formed upon deprotonation of the cation radical. Unless proton equilibrium for one of these two radical types is much slower than for the other type, the radical that corresponds to the lowest pKa should be formed upon de. protonation of the cation radical (Et2NH) . This cation radical has pKa 5.3 in aqueous solutions at pH 3–9. Therefore, O2 saturation of the solutions does not affect the determined pKa (Jonsson et al. 1996). Since O2 reacts more rapidly with carbon-centered radicals than . with nitrogen-centered radicals, one can conclude that deprotonation of (Et2NH) takes place at the nitrogen rather than at the -carbon. All cation radicals can be categorized as H acids on the basis of the nature of their conjugate radicals. The benzene cation radical can be named a acid because it gives the corresponding -radical as a result of deprotonation. The toluene cation radical gives a benzyl radical of the type. Therefore, this cation-radical is a acid. The phenol cation radical should be named a hetero--acid (not a hetero--acid) because an unpaired electron in its conjugate base—the phenoxyl radical—is delocalized within the aromatic -system. Some amines can give cation radicals that transform into radicals having three electrons in
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the field of two nuclei according to Scheme 1-42: . . H → R12 N MC R12 N MCH2R2 ← 苶HR2 ←→ R12 N M CHR2 H
Because the three-electron bonded radicals are formed at the cost of the removal of the nitrogen p-electron, such cation radicals should considered p acids. Of course, the formation and behavior of these p acids have to be dependent on steric factors. Such dependence will be considered in Chapter 3 (Section 3.2.3). Works by Tomilin et al. (1996, 2000) and Bieti and co-authors (1998) should be mentioned as describing stereoelectronic requirements to formations and configurational equilibria of N-alkyl-substituted cation radicals. Deprotonation is essential in some cation radical reactions; the corresponding examples will be described in Chapter 6. Scheme 1-43 depicts a photoreaction between phenanthrene and triethylamine. This reaction includes photoinduced sequential electron-transfer, proton-transfer, and radical-recombination processes (Lawson et al. 1999).
SCHEME 1-43
Many authors note the fact that adsorption of alkenes and alkyl aromatics upon the alumosilicate surfaces leads to a marked increase in their proton acidity (e.g., Farne et al. 1972). Formation of cation radicals and their further deprotonation explain many features of the hydrocarbon catalytic transformations (Vishnetskaya et al. 1997). Scheme 1-44 (see below) represents this phenomenon. .
.
⬅SiO ArCH2R → ⬅SiO (ArCH2R) . . (ArCH2R) → H ArCH R
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Vishnetskaya and Romanovskii described the cation radical approach to such catalytic reactions in their reviews of 1993 and 1994. It is worth noting here that organic cation radicals are usually very sensitive to surrounding nucleophilic reagents. This often depresses their deprotonation. In channels of a redox-active catalyst (e.g., in zeolites), the cation radical deprotonation is not aggravated, because competing reactions with the outside reagents are precluded (Roth et al. 1997; Corma & Garcia 1998). Concerning catalytic transformation of aliphatic hydrocarbons, an important problem is what positions are active in the deprotonation. In other words, what is a radical generated by proton abstraction from the initial cation radical? The ESR spectra of alkane cation radicals show that the dominant hyperfine coupling is caused by hydrogen atoms of terminal methyl groups that lie in the plane of the carbon skeleton in its extended conformation. The proton loss from the cation radical is thought to involve the carbon–hydrogen bond with the highest unpaired electron density (Toriyama et al. 1982). Calculations show that spin enrichment of the carbon–hydrogen bond leads to its elongation. This results in a weakening of the carbon–hydrogen binding and facilitates the bond disruption (Shchapin & Chuvylkin 1996). Thus, the fully expanded cation radical of n-heptane undergoes a proton transfer according to Scheme 1-45: .
.
RH RH → R RH 2 Reaction 1-45 leads to selective formation of the 1-heptyl radical (Demeyer & Stienlet 1993). In the case of trans-octene-2, the cation radical deprotonation predominantly leads to the allyl radical formation (Fel’dman et al. 1993, 1996), Scheme 1-46: .
.
[CH3CHBCH(CH2)4CH3] → H (CH2MCHMCH) –(CH2)4CH3 Like any other rule, the rule of preferential deprotonation at the highest spin density position has its exceptions. For example, in the cation radical of silacyclohexane, the maximal-spin-density atoms are carbons bonded with silicon. Nevertheless, deprotonation (in Fresno matrices) occurs preferentially from the SiH2 (Komaguchi & Shiotani 1997). The authors pointed out, that the deprotonation energy from SiH2 in a silaalkane is a little bit smaller than that from MCH2 in an alkane. This suggests that formation of the silyl radical . MSiH might be energetically more favorable. The solvation energy and the energy of matrix arrangement can also be important factors to explain the observed high selective deprotonation. Very recently, some attempts were undertaken to uncover the intimate mechanism of the cation radical deprotonation. Thus, the reaction of 9-methyl-10-phenylanthracene cation radical with 2,6-lutidine was studied (Lu et al. 2001). The reaction takes place by a two-step mechanism that involves the intermediate formation of a cation radical/base complex prior to unimolecular proton transfer and separation of products. Based on the value of the kinetic isotope effect observed, it was concluded that extensive proton tunneling is involved in the proton-transfer reaction. The assumed structure of the intermediate complex involves -bonding between the unshared electron pair on nitrogen of the lutidine with the electron-deficient -system of the cation radical. 1.3.2.B Cation Radical Basicity At first glance, the question of cation radical basicity seems meaningless: Cation radicals are usually strong acids. Nevertheless, there are two relevant examples. Corrole and isocorrole (as the octaethyl derivatives) give cation radicals that are basic in nature. Both neu-
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tral compounds belong to the porphyrin family and are named contracted porphyrins. They are typified by their contracted macrocycle, one carbon less than porphyrins, but still contain four pyrrole rings (Vogel 1996). One of the pyrrole rings contains nitrogen without the N–H bond. As proven, this nitrogen accepts a proton when the free bases of corrole and isocorrole are oxidized to the cation radicals by iodine and silver perchlorate in methylene chloride (MC). Solutions of the protonated cation radical are stable for more than a week at 20°C (Endeward et al. 1998). Notably, the described protonation proceeds at the expense of MC. MC is known as an extremely weak acid. The starting neutral bases are not protonated in MC in the absence of the oxidants! Some cation radicals can appear as hydrogen acceptors. Thus, fullerene C60 is oxidized to the cation radical at preparative scale by means of photoinduced electron transfer. This cation radical reacts with various donors of atomic hydrogen (alcohols, aldehydes, ethers), yielding the fullerene 1,2-dihydroderivative. In the case of tert-butanol, propionic acid, and glycol, product formation is also initiated by H abstraction from the OH group. The reaction proceeds according to Scheme 1-47 (Siedschlag et al. 2000): .
C60 e → C60;
.
.
C60 RH → HMC 60 R ;
.
HMC 60 R e → HMC60MR
The naphthalene cation radical can react with the hydrogen atom. The 1-hydronaphthalene cation resulting from this reaction is more stable than the reactants. This process has no activation barrier. After its formation, the 1-hydronaphthalene cation can absorb a photon and lose one of the two hydrogen atoms on carbon 1, since these hydrogens are more weakly bound than the aromatic hydrogens. Because the C–D bond is slightly stronger than the C–H bond, the reaction mentioned here can lead to some deuterium enrichment (Bauschlicher 1998). On the other hand, some organic cation radicals can act as hydrogen atom donors. For example, one-electron oxidation of 1,3-cyclohexadiene (C6H8) by 3,4-di(tert-butyl)-4. chloro-1,2-benzoquinone leads to the formation of C6H 8 . In the presence of sulfur, dehydroaromatization of the cation radical takes place (Berberova et al. 1998), Scheme 1-48: .
.
.
.
C6H 2HS →HSSH; C6H 8 S → C6H 7 HS ; 7 C6H8 → C6H 7 C6H 8 ; . . H C6H6 ← C6H 7 → C6H6 C6H8 etc.
1.4 METALLOCOMPLEX ION RADICALS 1.4.1 Metallocomplex Anion Radicals Organic anion radicals are odd-electron species. When this odd-electron nature is retained after coordination with a metal or it emerges as a result of some redox transformation of an initial metallocomplex, the newly formed species may be unstable. As to the metallocomplex anion radical assembling, high basicity and, consequently, high nucleophilicity of the starting anion radical ligands may play the decisive role in coordination with electrophilic metal-containing reactants. In other words, the electrostatic stabilization of the forming metallocomplex may compensate for these unfavorable factors. Besides, relocation of an unpaired electron from the anion radical ligand to the central metal atom may also lead to stabilization of such complexes (see also Section 2.2.1). . For instance, the [(-C5Ph5)Co(CO)2] anion radical complex has the (n 1)py metal orbital populated with an unpaired electron, according to calculations (Connelly et al. 1986). On the other hand, reduction of (bpy)Cr(CO)4 (bpy 2,2 -bipyridine) to its anion radical is known to occur without any major change in its structure or composition. The
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.
ESR spectrum of the [(bpy)Cr(CO)4] ion radical resembles that of the uncoordinated . (bpy) anion radical, clearly indicating that the unpaired electron is localized on the (bpy) ligand. This conclusion is fully corroborated by the similarity between the electronic spec. . tra of [(bpy)Cr(CO)4] and (bpy) (Vlcek et al. 1998 and references therein). Obviously, the product of one-electron reduction of (bpy)Cr(CO)4 may best be formulated as a for. mally Cr(0) complex with an anion radical ligand, (bpy) Cr(CO)4. There are anion radical metallocomplexes with complete spin retention at the anion radical ligand (Glockle et al. 2001), as well as those having an unpaired electron on a metal atom entirely, or those that share an unpaired electron with all parts of the complex (Kaim 1987 and references therein). It is clear that the type of spin distribution depends on the participation of fragment orbitals in such a molecular orbital (MO), which is available to be populated with the incoming electron. In terms of electron transfer, organometallic chemistry distinguishes the electron-transfer pathway (if the lowest unoccupied orbital is populated) or the hole-transfer pathway (if the highest occupied orbital is fitted with an outgoing electron). For examples, see Andersson et al. (2000). The lower the MO energy, the more probable the spin density localization in the framework of this MO. At the same time, coordination with a metal often decreases the energy level of the organic ligand MO. This increases the stability of the ligand anion radical form. For example, 2,1,3-thiadiazole gives an unstable anion radical upon one-electron reduction. However, this stability sharply increases if the anion radical is coordinated with two W(CO)5 fragments (at two nitrogen atoms of the thiadiazole ring) (Bock et al. 1988). If the 2,1,3-thiadiazole ring is condensed with the benzene ring, the anion radical stability increases even without the metal coordination (Todres, Lyakhovetskii, & Kursanov 1969). Meanwhile, complexation with two M(CO)5 fragments (M Cr, Mo, or W) changes the ESR spectra of these anion radicals. This means the metal participates in spin delocalization (Bock et al. 1988). The central metal atom in complexes usually has a shell of the nearest noble gas, e.g., the shell of 18 electrons. Upon one-electron transfer, the electron number turns from an even into an odd one (e.g., 19 for anion radicals and 17 for cation radicals). The one-unit change of the electron amount leads to an increase in the complex reactivity. The ligandsubstitution reactions are markedly facilitated. Thus, one-electron reduction of (bpy)Cr(CO)4 results in facilitation of the CO–PPh3 replacement in an electron-transfer-catalyzed reaction (Miholova & Vlcek 1985). Sometimes, the catalytic cycles in Scheme 1-49 become available (Kochi 1986). One-electron reduction of the LM complex starts the reaction cycle in Scheme 1. 49. The LM anion radical that is formed reacts with a nucleophile (N) undergoing the L → N ligand change. In so doing, the complex does not lose the unpaired electron; the L ligand is displaced out of the coordination-sphere limits. The anion radical product . NM passes its unpaired electron to the initial uncharged complex LM. This results in . regeneration of the LM anion radical. The neutral product NM goes out of the catalytic cycle. Scheme 1-49 illustrates the substitution of triphenylphosphine for pyridine as a ligand in the manganese cyclopentadienyl carbonyl complex. The reaction is initiated electrochemically, and only one-electron consumption is sufficient to involve 290 molecules of the initial LM complex in this transformation. The reaction is completed after half an hour at room temperature. With no electron initiation, prolonged boiling in toluene is required (Kochi 1986).
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SCHEME 1-49
Ohst and Kochi (1986) traced changes in the electron structure that take place during substitution of the triethyl phosphine ligand for the carbonyl ligand in the iron– phenyl–phosphine–carbonyl complex; see Scheme 1-50. The initial anion-radical of Scheme 1-50 is formed from the diamagnetic ternary nuclear complex upon one-electron reduction. This anion radical undergoes spontaneous breaking of one of the phosphine–iron bonds. The further substitution restores the Fe–P bond, which has been opened previously. Such restoration makes the whole reaction macroscopically reversible (in the sense of the cluster-skeleton preservation).
SCHEME 1-50
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The main point of the Scheme 1-50 transformation consists of splitting of the bridge bond between the metal and the ligand (not the metal–metal bond). It is important because this type of splitting leads to formation of the odd-electron shell around one of the iron atoms. The iron with the odd-electron shell acquires enhanced reactivity and changes one of its ligands. Without one-electron transfer (in diamagnetic clusters), this transformation does not take place at all. Activation of an M–CO bond for nucleophilic substitution in anion radical metallocomplexes appears to be quite a general effect (Kaim 1987; Mao et al. 1989, 1992; Shut et al. 1995; Klein et al. 1996). Such activation seems to be the basis of metal-cluster catalytic activity. The iron–sulfur cluster (Bu4N)2Fe4S4(SPh)4 deserves to be mentioned here. The cluster is considered as a ferredoxin model (Inoue & Nagata 1986); it catalyzes an electron transfer from n-butyl lithium or phenyl lithium to S-phenyl thiobenzoate or phenylbenzoate (Inoue & Nagata 1986). Other examples concern the interaction between iron carbonyles and potassium . alkylthiolates that is accompanied by disproportionation. The anion radicals Fe2(CO) 8 , . . . Fe3(CO)11 , Fe4(CO)13 , and Fe(CO)2 are formed (Belousov et al. 1987). The interaction of iron carbonyls Fe(CO)5, Fe2(CO)9, and Fe3(CO)12 with (CH3)3NO occurs according to a one-electron redox–disproportionation scheme, giving rise to iron carbonyl anion radicals: . . . . Fe2(CO) 8 Fe3(CO)11 , Fe3(CO)12 , and Fe4(CO)13 (Belousov & Belousova 1999). Sometimes, the ligand substitution in metallocomplexes can also proceed without reduction. The complex Co(CO)3L2[L2 bis(diphenylphosphino)maleic anhydride] readily exchanges one of the carbonyl groups for trialkylphosphine. Kinetic studies showed that the substitution goes through the Co–CO bond disruption. The initial complex has tetragonal-pyramidal geometry and possesses one unfilled vacancy to coordinate an additional ligand. Studies of the complex Co(13CO)3L2 by the ESR method revealed that the excess electron presumably populates the antibonding orbital of the Co–CO fragment. This bond is just weakening. The substitution takes place according to Scheme 1-51: Co(CO)3L2 PPh3 → CO Co(CO)2L2PPh3 The initial 19-electron complex was classified as the 18-electron complex with the anion radical ligand (Mao et al. 1989). Insertion of new ligands into metallocomplex systems may proceed reversibly. Being reduced in the framework of the complex, these ligands lose the ability to be coordinated and leave the coordination sphere as products. One important example of such ligand sliding is the catalytic transformation of CO2 into CO. Rhenium, palladium, platinum, and nickel complexes were recommended to catalyze this process (Hawecker et al. 1986; Du Bois & Meidaner 1987). The Ni(II) complex with 1,4,8,11-tetraazacyclotetradecane is preferential (Beley and co-authors 1984). The preceding examples show that one-electron reduction of metalloorganic complexes or coordination between a metal and an anion radical ligand may expand an electron shell of the central metal atom. Sometimes, anion radical metallocomplexes contrast in this regard with the cation radical ones. Thus, some metalloporphyrins form cation radicals with charges and unpaired electrons on ligands (Shinomura et al. 1981) and anion radicals with charges and unpaired electrons on metals (Lexa et al. 1989). There is some lack in data on anion radical stabilization as a result of binding into a metallocomplex. One specific case was described for bis(trimethylsilyl) diacetylene (Kaim 1988). When treated with potassium, this compound gives tetrakis(trimethylsily)butatetraene, probably owing to some transformations of the unstable anion radical. This anion
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radical is stabilized in the presence of trimethyl aluminum, giving the anion radical metal. locomplex [(Me3Al)(Me3Si)CBCBCBC(SiMe3)(AlMe3)] K. 1.4.2 Metallocomplex Cation Radicals There is a large body of literature describing one-electron oxidation of metallocomplexes with organic ligands, metal–ligand interactions in cation radical products, and structural changes on the formation of cation radicals. Such abundance is caused by the chemical nature of metallocomplexes, in which the metallic center readily transforms into a higher state of oxidation. Kochi (1986) and Kaim (1987) have covered many of these problems in their reviews. Astruc (1995) has done one recent and very important generalization in the field. One of the most important intricacies in the redox chemistry of organyl metallocomplexes is that both a metal and a ligand are potential candidates for oxidation. One example can be taken from the porphyrin (P) family. It is possible to move between the two for. mulations. The porphyrine-oxidized manganese(III) bis-perchlorate, (P )MnIII(ClO4)2, gives the metal-oxidized manganese (IV) bis-chloride, (P)MnIV(Cl)2, when the two perchlorate ligands have been substituted with the two chloride ligands (Kaustov et al. 1997). The copper–porphyrin complex gives cation radicals with significant reactivity at the molecular periphery. This reactivity appears to be that of nucleophilic attack on this cation radical, which belongs to the type (Ehlinger & Scheidt 1999). A new bifunctional tetrathiafulvalene-type donor molecule (D––D) with a copper iodine bridge has recently been synthesized. Its cation radical salt, (D––D)BF 4 , has the hole localized on one of the D fragments, although copper is a metal of variable valence (Ramos et al. 1997). The porphyrin–cobalt complex gives rise to the cation radical with charge spin localization at the nitrogen atom of the porphyrin ring. The cation radical thus formed acquires enhanced reactivity and can add tolane (Kochi 1986), Scheme 1-52. The main point
SCHEME 1-52
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of the transformation in Scheme 1-52 is that the central cobalt atom eventually becomes an electron reservoir. Iron complexes with organic ligands are widely used as such reservoirs. For example, the stable 19-electron complex cyclopentadienyl benzene iron(I) is a clean electron reservoir that has a sufficiently negative oxidation potential and can be easily made, stored, and handled and can be weighed accurately (Alonso & Astruc 2000). Reversible one-electron oxidation of ferrocene and its derivatives to cation radicals (so-called “ferricenium” cations) is a well-known reaction. The cation radical center is localized at the iron atom. In contrast to this statement, the hole transfers though conjugated systems were proven for the bis(ferrocenyl) ethylene cation radical (Delgado-Pena et al. 1983) and the cation radical of bis(fulvaleneiron) (LeVanda et al. 1976). Scheme 1-53 depicts these structures. Extended Hueckel MO calculations (Kirchner et al. 1976) support the formulation that the unpaired electron is delocalized over both metallocene residues in the bis(fulvaleneiron) cation radical. Alternatively, very fast intraionic intervalence electron transfer may take place between the formal Fe(II) and Fe(III) atoms. In general, the ethylene bond in organic cation radicals is weakened, and the barrier to rotation becomes significantly less than that of the neutral ethylene derivative. This particular property of the ethylene bond in cation radicals has been used to probe for the mechanism of many reactions (Todres 1987). As revealed, cation radicals of ferrocenyl ethylenes do not undergo the cis-to-trans isomerization. Calculations show that the cation radical’s center is located exclusively at the iron atom, with no participation of the ethylene bond (Todres et al. 1992). Hence, oneelectron oxidation of this ferrocenyl ethylene occurs at the iron atom exclusively; anything else would be extremely unusual in this case. Thus, the recently described stable enol linked to a ferrocene redox center gives the cation radical upon one-electron oxidation. This species is better characterized as a ferricenium salt than as an enol cation radical (Schmittel & Langels 1998). Many hydroxyquinones can bind metal cation via chelate formation. Magnesium chelates of 5,8-dihydroxynaphtho-1,4-quinone and 1,4-dihydroxyanthra-9,10-quinone are characterized by large complex-formation constants. These constants become very small in cases of the corresponding semiquinones (anion radicals). This apparent paradox is explained by the much stronger intramolecular hydrogen bonding existing in the hydroxysemiquinones as compared to the neutral quinones (Alegria and co-workers 2000).
SCHEME 1-53
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Recently, variable-temperature absorption data were obtained for oxomolybdenum dithiolate complex [MoVO(qdt)2]1; here qdt is quinoxaline-2,3-dithiolate. The data clearly reveal the highly thermochromic nature of this compound. The authors (Helton and others 2001) explain the observed thermochromism through a thermally induced intramolecular electron transfer according to Scheme 1-54: .
[(qdt)MoVO(qdt)1 ⇔ [(qdt)MoIVO(qdt) ]1 The thermochromism in Scheme 1-54 represents valence tautomerism in molybdenum complexes and, at the same time, the inherent redox activity of a dithoilate ligand. This phenomenon may be significant with respect to the biological activity of various molybdenum enzymes. 1.5 ORGANIC ION RADICALS WITH SEVERAL UNPAIRED ELECTRONS AND/OR CHARGES In principle, an organic molecule can accept as many electron pairs as it has low-lying vacant orbitals. In the same way, high-lying occupied orbitals can release not a single, but several electrons. Such multielectron processes can result in the formation of poly(ion radicals). As seen later, the main topic of interest in poly(ion radicals) consists in their spin multiplicity. Therefore, it is meaningless to divide the material in the anion and cation radical parts. Let us therefore discuss first the reductive electron transfer. Anion radicals are the first products of reductive electron transfer. Mutual repulsion of the primary and subsequent excess electrons forms the basis of the impediment of polyelectron reduction at the initial step. However, an orbital, which already has one electron, can be populated by another electron if a proper reducer is chosen. This reducer must be able to overcome Coulombic repulsion. As a result, more or less stable dianions are formed. Like anion radicals, dianions can reversibly give excess electrons back. If skeleton rearrangement is absent, the starting, uncharged molecules are regenerated. In dianion diradicals, two vacant orbitals are populated, i.e., LUMO and the next one. This next orbital must lay low enough, as well. Such an orbital can be detected in molecules of unsaturated hydrocarbons of an extensive contour. Within this expanded contour, extensible electron delocalization somewhat decreases the repulsion energy. Several electronacceptor fragments (substituents with heteroatoms or with high conjugation ability) assist in formation of poly(anion radicals). Because poly(anion radical)s are polycharged species, they are very prone to association with counterions. The association decreases the total energy of the system, increases its stability, and facilitates its formation. The formation of poly(anion radical)s is usually successful if alkali metals are used as reducers in DME, THF, or MeTHF as solvents. Medium components, e.g., HMPA or cryptands, which weaken ion-pair interaction, work unfavorably. Even traces of proton admixtures evoke a transformation of reducing compounds into hydrogenated products instead of polycharged anions. As a rule, the higher the charge of a particle to be prepared, the lower must be the temperature and the longer the contact with a reducer. Working with such species requires an inert atmosphere, of course. Poly(ion radical)s excite today’s interest first of all because of their possible application in microelectronics. Therefore, the spin distribution is a very important point for these species. New organic triplet poly(ion radicals), with strong intramolecular ferromagnetic coupling and good thermal stability, may serve as building blocks for organic magnets.
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For example, sodium reduction of bis(dimesitylboron)-1,3-phenylene in 2-methyltetrahydrofurane-tetrahydrofurane (MeTHF-THF) mixture leads to the corresponding dianion diradical as a species of a triplet state (Rajca, Rajca, & Desai 1995). In the case of 1,8-diphenylnaphthalene, there is an interesting difference in the spin states of the anion radical on the one hand and of the trianion radical on the other. In the anion radical, the main contour for spin distribution is the naphthalene framework. In the trianion radical, the naphthalene system bears two electrons (nonassociated with any bond) while the third (unpaired) electron oscillates between both phenyl substituents at positions 1 and 8 (Gerson & Huber 1987). Teracyanoquinodimethide gives the anion radical and dianion. It is unable to form the trianion radical. However, with adhered -electron contour, the trianion radical does exist—such a form was described for tetracyano-2,7-pyrenoquinodimethide (Gerson, Heckendorn, & Cowan 1983). The anion radical of tetracyano-2,7-pyrenoquinodimethide contains “a half” an electron at each dicyanomethide group. The corresponding dianion holds “the whole” electron at the same sites, and the trianion radical has one electron at each C(CN)2 fragment and one electron even in the arene -system (Gerson & Huber 1987). This trianion radical was also described as an anion radical bearing two negatively charged dicyanomethide groups (Rieger et al. 1963). Several anion biradicals have recently been reported, composed of semiquinone and nitroxide functionalities. The ground-state triplet anion biradicals were derived from oneelectron reduction of open-shell acceptors, in which p-benzoquinone (BQ) is substituted by a nitronyl (Shultz & Farmer 1998) or nitronyl-nitroxide (NN) group (Kumai et al. 1994). Scheme 1-55 gives the BQNN example. . The anion diradical (BQNN) has two nondegenerated single occupied molecular orbitals (SOMOs). One is delocalized over the entire molecule and the other (SOMO ) is localized within the NN group. As a future cation diradical counterpart to this novel class of ground triplet species, the same Japanese crew has designed and prepared the following open-shell donors carrying the NN groups (Sakurai et al. 1996, 2000, Kumai et al. 1996). These donors are presented in Scheme 1-56. All mentioned donors were oxidized by excess I2 in THF or MeTHF. ESR spectroscopy at cryogenic temperatures (6–110 K) established the paramagnetic nature of the oxidized products. In order to rationalize the ground-state spin multiplicity, the MO calculations were performed on these paramagnetic species. The results strongly suggest that the species exist as the cation radicals with triplet ground states. There are two nondegenerated
SCHEME 1-55
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SCHEME 1-56
SOMO and SOMO , too. The delocalized spin on the donor site and the localized spin on the radical site in these cation diradicals are oriented in a parallel manner. As will be seen in Chapter 7 (Section 7.4), such a spin orientation results in ferromagnetic coupling (triplet interaction) between the spins. In contrast, an antiparallel type of mutual orientation of the spins results in antiferromagnetic coupling. The parallel type is crucial in the design of organic materials with magnetic properties. Another motif for preparing magnetic materials based on anion radicals is the socalled “metal radical” approach. The approach consists of designing an extended lattice containing paramagnetic metal ions whose magnetic interactions are determined by appropriate bridging radical ligands. This strategy has been applied successfully using as ligands the anion radicals of tetracyanoethylene (Brandon et al. 1998) or 1,3-bis(3,4-dihydroxy-5tert-butylphenyl)-5-tert-butylbenzene (the m-xylylene type of ligand; Caneschi and others 2001). Attempts to build di(anion radicals) with the participation of the fullerene C60 moiety deserve a special mention. Reduction of the benzoquinone-linked fullerene C60 first gives the monoanion containing the semiquinone radical and the C60 moiety. Further reduction of the monoanion produces the dianion having the semiquinone radical and the C60 anion radical. The ESR spectrum of this di(anion radical) clearly shows the triplet interaction between the semiquinone and C60 anion radical (Iyoda et al. 1996). In anion diradicals of fullerene C60 derivatives bearing a nitroxide group (namely, 2,2,6,6-tetramethylpiperidyl-1-oxyl), one electron is located in the fullerene moiety and experiences a strong exchange coupling with the nitroxide unpaired electron (Arena and co-authors 1997). At room temperature both triplet and singlet states of this anion diradical turn out to be mixed. Wienk and Janssen (1996) have designed N,N -bis[4-(diphenylamino)-phenyl]N,N -diphenyl-1,3-diaminobenzene, Scheme 1-57. This phenylaniline comprises alternating para- and meta-substituted phenylene (C6H4) rings to give the desired chemical stability from the p-diaminobenzene unit and ferromagnetic coupling from the m-diaminobenzene unit after two-electron oxidation. The N-phenylaniline oligomer depicted in Scheme
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SCHEME 1-57
1-57 was synthesized from 4-iodo-N,N-diphenylaminobenzene and N,N -diphenyl-1,3-diaminobenzene in diphenyl ether at 200°C, using butyllithium and copper(I) iodide as a catalyst. Oxidation of the product with more than two equivalents of thianthrenium perchlorate resulted in the formation of the dication diradical. This reaction proceeds in the trifluoroacetic acid (TFA) at room temperature. The product is stable at room temperature and has a triplet ground state, as evidenced from ESR spectroscopy. Using N,N -bis(p-anisyl)-1,3-diaminobenzene or N,N -bis( p-anisyl)-2,7-diaminonaphthalene as central units, Blackstock’s group prepared the analogous dication diradicals, which were found to be solution-stable triplets (Stickley & Blackstock 1994; Selby Blackstock 1999). Depending on the amount of thianthrenium perchlorate, N,N,N ,N ,N,N-hexakis (anisyl)-1,3,5-triaminobenzene gives its cation radical, dication diradical, and trication triradical as well (Stickley et al. 1997). These species are stable in methylene chloride at low temperatures (at 298 K they can exist for several days). Spin and charge are localized at each oxidized nitrogen atom. The dication diradical and trication triradical structures are ground-state triplet and quartet molecules (Sato et al. 1997). The latter two reactions demonstrate the possibility of using thianthrenium perchlorate in preparing magnet-active ion radicals with two or even three unpaired electrons and positive charges. (At this point one very important caution has to be stated: Thianthrenium perchlorate is a shock-sensitive explosive solid and should be handled with care.) From the materials just given above, one can conclude that mutual meta orientation (“meta through a benzene”) of the spin-bearing moieties is an indispensable condition for the existence of triplet states in aromatic di- or tri-(cation radicals). However, in fact, these systems have both singlet and triplet forms, and the questions are about what the difference is in the corresponding energy and which state is the more stable. Indeed, there is some experimental evidence of violation of the meta-topology rule. The following two cases are illustrative examples. 1. N,N,N ,N -Tetra(anisyl)-3,4 -diaminobiphenyl: Being oxidized with dichlorodicyanobenzoquinone in trifluoroacetic acid, this compound gives a dication di-
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SCHEME 1-58
radical according to Scheme 1-58. The di(cation radical) from Scheme 1-58 is stable in solution at room temperature. Studies by cyclic voltammetry and ESR spectroscopy led to the conclusion that the two centers are not wholly independent and that there is “leakage” of charge and spin between the two halves (Bushby et al. 1997a,b). Hence, the charge and spin distributions substantially overlap. This statement finds its confirmation in cases of other, similar systems (Karafiloglou & Launay 1998; Boman with co-authors 1999). 2. Phenothiazine derivatives: Comparison between the structures depicted in Scheme 1-59 can lead to a prediction that hexa(anisyl)-1,3,5-triaminobenzene is more similar to 10,10 -(1,3-phenylene)diphenothiazine than to 10,10 -(1,4phenylene)diphenothiazine. Nevertheless, upon oxidation in concentrated sulfuric acid, these two isomeric diphenothiazines give dication diradicals. As shown, m-phenylenediphenothaizine has a singlet state, and p-phenylenediphenothiazine has a triplet state (Okada et al. 1996; Baumgarten 1997). Because of geometry peculiarities, the resulting orbital energy splitting between the two SOMOs is larger in the meta isomer than in the para isomer.
SCHEME 1-59
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In these geometries, mixing with -orbitals of the central benzene ring destabilizes local cation radical orbitals (Fang et al. 1995). In particular, the energy difference between the two -orbitals with different symmetries (S and A) is larger in m-phenylene than in the p-phenylene framework simply because of the smaller number of -bonds in the m-phenylene structure than in the p-phenylene structure. Hence, the triplet state of the para-phenylene-bridged dication biradical is more stable than that of the meta-bridged one. According to quantum mechanical calculations performed at distortions close to 90°, para-derivatives may indeed form more stable triplet states than meta-phenylene-linked structures (Baumgarten 1997). If extended m-phenylene structures have several spin-bearing centers, their triplet coupling (at zero field) depends on spatial spin distribution. Namely, the coupling degree is inversely proportional to R3, where R is the distance between the “unpaired” electrons. Poly(arylmethyl) anion diradicals are an example of such dependence (Rajca & Rajca 1995). When the charge (electron pair) is localized at the terminal site, as in Scheme 1-60, ESR studies indicate ferromagnetic coupling between the remaining “unpaired” electrons. However, when the negative charge (electron pair) is localized at the center site, as in Scheme 1-61, antiferromagnetic coupling between the remaining “unpaired” electrons is observed. The distance between the two spin-bearing centers is much greater in the structure of Scheme 1-61 than in the structure of Scheme 1-60. Note that in this case a control of spin coupling via control of electron localization is attained. The straightforward control of electron localization was implemented via modification of the substituents at the center site (see Schemes 1-60 and 1-61). The mentioned anion diradicals were prepared by partial oxidation of relative polyanions with iodine (Rajca & Rajca 1995). When the carbanion-stabilizing 4-biphenylyl, in the place of 4-tertbutylphe-nyl, was employed as a substituent at the center site, the negative charge was confined to the center site and the spin density to the terminal (nonadjacent) site. Control of charge/spin localization at the specific sites leads to two different kinds of spin coupling: (1) When spin sites are adjacent, spin coupling between the “unpaired” electrons is ferromagnetic. (2) When spin sites are not adjacent, because the pathway between them includes a site with negative charge (electron pair), spin coupling between the “unpaired” electrons appears to be antiferromagnetic.
SCHEME 1-60
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SCHEME 1-61
Distribution of spin density in the framework of trianion radicals allows determining whether or not there is some conjugation between substituents at different sites. For example, reduction of di-nitro-spiro(indoline-benzopyran) with potassium tert-butoxide in dimethysulfoxide (DMSO) proceeds according to Scheme 1-62 (Alberti et al. 1995). As shown by means of electrochemical ESR spectroscopic methods and confirmed by ab initio calculations (Alberti et al. 1995), the first electron is accommodated in the more accessible nitrobenzopyran moiety. Next, the second electron fits into the nitroindoline moiety in virtually the same orbital that in the anion radical of mono-nitro-spiro(indolinebenzopyran), is occupied by the unpaired electron. The mono-nitro anion radical is also formed by the aforementioned reaction with tert-BuOK in DMSO, Scheme 1-63.
SCHEME 1-62
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SCHEME 1-63
As to the third step of the step-wise one-electron reduction of di-nitro-spiro(indolinebenzopyran), the third extra electron is again accommodated into the nitrobenzopyran unit without affecting the spin density distribution on the opposite fragment, which therefore remains essentially identical to that of the anion radical of mono-nitro-spiro(indoline-benzopyran). Hence, the spiro-nodal carbon atom prevents conjugate interactions between the indoline and benzopyran moieties. Scheme 1-64 points up one additional and important aspect to the problem of ferromagnetic coupling in dication diradicals. The stilbenoid compounds of Scheme 1-64 . . give N …/…N dication diradicals upon two consecutive one-electron oxidations (anodic or by means of NOPF6). The compound containing the methoxy group in the stilbenoid fragment has a half-life of up to two months, even under ambient conditions. This
SCHEME 1-64
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half-life is five times longer than that of the counterpart with no methoxy group in the stilbenoid moiety. However, introduction of the methoxy group results in torsional twisting on the substituted stilbenoid framework. This weakens ferromagnetic coupling in the CH3O-substituted dication diradical as compared with the H (stilbenoid-unsubstituted) analog (Michinobu et al. 2001). So ferromagnetic coupling needs not only the triplet orientation of the spins but also the spin-exchanging ability in the framework of a planar dication diradical. 1.6 POLYMERIC ION RADICALS Polymeric ion radicals are usually formed as a result of one-electron redox modifications of uncharged polymers containing electrochemically active groups. They attract enhanced attention in the sense of possible practical applications. Because polymeric ion radicals contain many spin-bearing groups, a similarity emerges between polymeric ion radicals and poly(ion radicals). However, there is some specificity peculiar just to polymers. For example, NOBF4 oxidation of the triphenylamine and diphenyltetraalkoxy fragments in a linear and networked polymer leads to the formation of cation radical centers (Bushby et al. 1997b). As a rule, the quantity of ion radical centers formed is low, not more than 10–20% of the theoretically possible values. This is a problem that is proving difficult to solve. A feasible explanation is that the effect is not electrostatic but steric. There is a certain difficulty in accommodating large counterions within rather rigid polymer networks of bent polymeric chains. Anion radicals engrafted to hyperbranched polymers (dendrimers) deserve special mention. Dendrimers are built up by the sequential synthesis of larger and larger generations, with branch points in each generation. The ground-state spin multiplicity and geometry of the reductively and oxidatively treated homo- or heteronuclear dendrimers were studied theoretically and deemed technically promising. In particular, poly(amidoamine) dendrimers were peripherally modified with diimide moieties; see the structure in Scheme 1-65. After reduction with dithionite, this dendrimer was cast into a film the electrical properties of which were isotropic. (This means that on the molecular and macroscopic levels there is a three-dimensional electron delocalization.) The conductivity was humidity dependent (water may take part in long-distance electron transfer). At 95% humidity, a mixed-valence film (0.55 electron per diimide) showed conduction at room temperature around 11 1cm1 (Miller & Mann 1996). As shown later, partially reduced mixed-valence materials are required for organic metals of high electrical conductivity. The convergent synthesis of a range of aryl ester dendrimers with peripheral tetrathiafulvalene units was also reported (Devonport et al. 1998). The dendrimers acquire some amount of the cation radical tetrathiafulvalene “tips” upon reaction with iodine in solutions. We have already pointed out that one of the main reasons for the modern interest in polymeric and oligomeric ion radicals is the design of organic compounds with magnetic properties (see Section 7.4). Oligomers and polymers allow, in principle, preparing organic high-spin molecules. A well-established strategy is to connect neutral or charged radicals as spin-carrying units via organic linkers. These units may act as ferromagnetic coupling parts and provide the desired alignment of electron spins. High-molecular ion radicals often have 1,3-phenylene (m-phenylene) as a coupling unit too (cf. Section 1.5). This unit couples ion radical centers at the 1 and 3 positions by an in-phase periodicity of the spin polarization. Cooperative interaction of a large number
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SCHEME 1-65
of unpaired electrons may eventually result in high-spin polymers (oligomers) and ultimately in ferromagnetic materials. A interesting theoretical strategy toward high-spin polyradicals/ion radicals was developed (Fukutome et al. 1987). This strategy aimed at the construction of ferromagnetic chains in which unpaired electrons can be introduced by reduction or oxidation of -conjugated segments. Those ion radicals were supposed to be positioned between the coupling units. In the language of solid-state physics, the ion radical segments are polarons.
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Thus, p-phenylene diamine oligomers linked via the benzene ring in the 1,3 or 1,3,5 manner have been successfully oxidized to high-spin ground-state oligo(cation radical)s that are stable at room temperature (Stickley & Blackstock 1994, 1997; Wienk and Janssen 1996, 1997). Ferromagnetic coupling occurs between p-phenylene diamine cation radicals in regioregular substituted -conjugated chains between neighbors and possibly also between next-nearest neighbors (Meurs & Janssen 2000, 2001). Pendant p-phenylene diamine cation radicals, regioregularly substituted on a conjugated polymer chain, constitute a promising approach toward stable high-spin polymers. Poly(aniline-2-chloroaniline)-p-toluenesulfonic acid salt was obtained by oxidative polymerization of aniline with o-chloroaniline in solutions containing p-toluenesulfonic acid. The copolymer salt was subjected to heat treatment under nitrogen atmosphere at elevated (about 150°C) temperatures. The heat-treated samples acquired electric conductivity 2.7 102 1cm1. According to ESR spectra, the heated poly(aniline-2-chloroaniline)-p-toluenesulfonic salt exists as the poly(semiquinone cation radical), in which unpaired electrons are localized on or near the nitrogen atoms (Palaniappan 1997). The polaronic strategy has also been applied to polymers, incorporating m-phenylene units as coupling links in -conjugated polymer chains of poly(thienylene ketone) (Colle Dal et al. 1999) as well as of polyarylamine, polyacetylene, polythiophene (Kaisaki et al. 1991; Murray et al. 1994; Bushby et al. 1997a). Scheme 1-66 presents these polaronic polymers. Oxidation of the polymers from Scheme 1-66 leads to the polymeric cation radicals with ferromagnetic coupling of spins. Surprisingly, however, the spin concentration in these polymer networks was extremely low. Only a few percent of the cation radical units actually carried an unpaired electron. Several explanations were proposed for this disappointing result. According to Bushby et al. (1997b), for the cross-linked polymers, a more likely interpretation is the steric difficulty of incorporating counterions into a relatively rigid polymer network. The authors consider the PF 6 anion as the counterion to each cation radical part of the crosslinked polymer. If this is so, it is unlikely to be insurmountable. A bulky anion can be replaced by an anion of small size. As to the tailor-made polymers, a more serious, perhaps intrinsic, cause is discussed (Haare et al. 1998). These polymeric cation radicals form -macrodimers according to . usual equilibria 2M ⇔ (M)2 2 . The -dimerization inhibits the formation of high-spin states. Maybe other -conjugated ion radicals with a much lower tendency to form dimers, such as polymeric cation radicals based on p-phenylenediamine, are a promising way to construct intrachain ferromagnetic coupling in polymers. Of course, charge transport in charge-injected polymers is extensively studied because of its importance in practice. Sometimes it is sufficient to introduce a donor or acceptor in a polymer. This leads to the formation of a charge-transfer complex, which transforms into an ion radical salt upon physical excitation. Thus, S. A. Lee and co-authors (1997) studied the photoconductivity of pyromellitic dianhydride-diaminodicyclohexylmethane polyimide and of pyromellitic dianhydride-oxydianiline polyimide in the presence of electron donors, e.g., N,N,N ,N -tetramethyl-p-phenylendiamine. The addition of the electron donor increased the photocurrent generation of polyimide films by about three orders of magnitude. The polyimide from an alicyclic diamine doped with an electron donor showed a larger enhancement of photocurrent. This occurs due to an intermolecular chargetransfer complex. In the case of the aromatic polyimide, both inter- and intramolecular charge-transfer complexes can be formed. By and large, any enhancement of the photocur-
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SCHEME 1-66
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rent by the electron donor is attributed to photoabsorption by the molecular charge-transfer complex formed between the added electron donor and the pyromellitic imide unit of the polymer backbone. The photoexcitation provokes the rapid electron transfer from the donor to the pyrromellitic imide. This produces the anion radical of polymer and the cation radical of the donor, resulting in the photoconductivity in the bulk polyimide films. Localization or delocalization of an excess electron (in anion radicals) and a hole (in cation radicals) along the polymer chain determines the electronic conductivity of the polymeric ion radicals. From this point of view, the cation radical of vinylene-bridged-octithiophene oligomer should be mentioned. Quantum chemical calculation (Casado et al. 2000) indicates that the hole localizes in the middle of the molecule. The hole extends over the central part built from four thyenylene units. Another representatives of such polymers, polysilanes, also attract significant technical interest. Polysilane is a one-dimensional -conjugated polymer. In general, the silyl group can support electron transfer through low-lying-d- and *-orbitals (see, e.g., Gherghel et al. 1995). Several scientific groups investigated the localization length (LL) of the electron/hole conduction in polysilanes. Irie and Irie (1997) and then Seki et al. (2001) recorded the absorption spectra of radical ions of polyalkylsilanes. The spectra had two peaks in the UV and near-infrared regions, and the maxima of the peaks shifted to longer wavelengths with increasing chain length until about 16 monomer units and clearly showed saturation above this length (i.e., LL 16). Kumagai et al. (1996; Ichikawa et al. 1999) measured electron spin resonance and the electronic absorption spectra of anion and cation radicals derived from oligosilanes (up to six monomeric units) or polysilanes (approximately 500 units). The spectra of the polymer ion radicals were similar to those of the oligosilane ion radicals. This suggested that both an excess electron and a hole were not delocalized all over the Si–Si main chain of the polymeric ion radical. The excess electron and hole in the polysilane ion radical are confined to only a part of the polymer chain composed of six silicon atoms (LL 6). Thus, the studies reported that excess electrons and holes are localized on a few silicon atoms, although the localization length reported in each work is slightly different. Tachikawa (1999) also analyzed mobilities of carriers along the silicon chain, and his results should be mentioned separately. As it turned out, the mobility obtained for a positive charge (hole) was several times larger than that for an excess electron. This result suggests that the localization mechanism of a hole and that of an electron are different from each other. Probably, an excess electron is trapped in the defect of the main chain, whereas a hole is not trapped. The defects are mainly structural ones, such as branching points and oxidized sites (Seki et al. 1999). This can lead to a different electron conductivity. Continuation of the polysilane ion radical studies will hopefully result in some important technical applications. 1.7 INORGANIC ION RADICALS IN REACTIONS WITH ORGANIC SUBSTRATES Practically every organic reaction proceeds with the indispensable participation of inorganic substances involved in redox processes. Among these species are oxygen, halogens, compounds of sulfur and nitrogen with oxygen, inorganic ions of the hydroxyl type, and metallic ions. As a result of electron transfer, the inorganic ions are oxidized, reduced, or somewhat changed with respect to the central atom’s valence. Many of the electron-transfer reactions start from an attack of an organic substrate by an inorganic reactant. The lat-
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.
ter often reacts in its ion radical state (the SO 4 anion radical is a typical example). Some ion radical transformations of organic compounds take place with the assistance of active components in a gaseous medium (O2, CO2). These components can give rise to ion radi. . cals (O 2 , CO 2 ). Therefore, some concise review of the conditions of formation and the physical and chemical properties of the most important inorganic ion radicals is in order here. Special attention ought to be given to their reactions with organic compounds, both in an initial uncharged form and in a post-electron-transfer form. It is also very important to understand the difference between redox reactions of the same ions acting under different conditions of solvation—or why ions with close redox potentials are quite different in their activity in liquid-phase electron transfer. For example, the cupric ion is a weak oxidant in aqueous solutions. In acetonitrile (AN) solutions, the same ion is a rather strong oxidant (Ahrland et al. 1983). Ions NO and NO 2 are close in their E0 values (1.51 and 1.56 V, respectively—Bontempelli et al. 1974; Cauquis & Serve 1968). There is a marked solvent dependence of the standard potential E0 for the NO 2 /NO2 couple: 1.56 V in AN, 2.05 V in sulfolane, and 2.32 V in nitromethane (Boughriet & Wartel 1989a, 1989b). In other words, the oxidative properties of NO 2 depend on solvating capability. On account of the weak solvation properties of nitromethane, the NO 2 species is found to be a more powerful oxidizing agent in nitromethane than in sulfolane and particularly in AN. 1.7.1 Superoxide Ion .
The standard potential of the O2 /O 2 pair is equal to 0.15 V in water and 0.60 V in DMF. Usually, dioxygen easily captures one and then another electron, Scheme 1-67: .
O2 e ⇔ O 2 . 2 O e ⇔ O 2 2 In DMSO, dioxygen reductions into the superoxide ion and then into the dioxygen I I dianion are characterized by E 1/2 0.5 and E 1/2 1.5 V (regarding the saturated calomel electrode, SCE) (Sawyer & Gibian 1979). The superoxide ion has an intermediate position in the redox triad of Scheme 1-68: .
2 O2 ⇔ O 2 ⇔ O2
In accordance with such a position in the triad, the superoxide ion possesses a dual reactivity: Depending on the substrate character, the ion can act as an acceptor or as a donor of one electron. For instance, the superoxide ion transfers one electron to quinones, nitro compounds, and transforms them into anion radicals (Sawyer & Gibian 1979). Scheme 1-69 below illustrates the reversible transformation of quinone (Q) into semiquinone (SQ). .
O 2 Q ⇔ SQ O2 As was found out, the direct and reverse processes are fast but have different rates: k1 1.1 . 109, and k1 107 Lmole 1sec1 (Morkovnik & Okhlobystin 1979). According to a widespread opinion, the superoxide ion possesses expressed oxidative properties. Thermodynamically, however, it must be a moderately strong reductant and a very weak oxidant (Sawyer & Gibian 1979). This statement refers, of course, to aprotic mediums, in which the superoxide ion is stable. In the presence of proton donors, the su-
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peroxide ion undergoes disproportionation according to Scheme 1-70: .
.
O 2 H ⇔ HOO . . O 2 HOO → O2 HOO
In all the aprotic solvents (with no proton donor admixtures), the superoxide ion generally cannot act as an oxidant, taking into account a wide range of functionally substituted . compounds. For example, in dry pyridine, O 2 does not oxidize 1,2-dimethoxybenzene (Sawyer & Gibian 1979). This ion, however, reacts with 1,2-dihydroxybenzene. For the OH form, the first step consists of proton transfer from the hydroxyl group to the superox. ide ion. Next, reactions proceed with the participation of HOO . The latter is formed according to the disproportionation of Scheme 1-70. The literature on the oxidation of hydrazines, thiols, and alcohols with the superoxide ion (Sawyer & Gibian 1979) refers to reactions that begin with the formation of the . . HOO radical. It is significant in this respect that there are no reactions between the O 2 ion and dialkyl sulfides, dialkyl ethers (esters), or alcoholates. All of these substrates do not . . contain a labile proton. The free hydroperoxide radical (HOO , which is formed from O 2 after the proton addition) is a conjugate acid of a base, i.e., of the superoxide ion. It is a rather strong acid, with pKa 4.88 (Sawyer & Gibian 1979). Hence, dissociation of the hydroperoxide radical with the generation of a proton and a peroxide ion is possible even . . in acid media. However, the irreversible reaction between HOO and O 2 suppresses the . dissociation. Studies in the chemistry of the O 2 anion radical have achieved noticeable success during the past two decades only. Convenient preparative methods of electrochemical reduction of molecular oxygen were developed to obtain solutions of tetraalkyl ammonium superoxides in aprotic solvents. Usually, oxygen (in a gaseous stream through a phone electrolyte, Alk4NX, solution) is electrolyzed at the controlled potential of 1 V. As solvents, DMF, DMSO, AN, and pyridine are employed. Mercury, platinum, and gold are used as cathodes. The colored solutions formed give an ESR signal at a low temperature. The signal corresponds to the superoxide ion. A full chemical method has been proposed to obtain the superoxide ion (Miyazawa et al. 1985). In this case, oxygen is bubbled through DMF, THF, DMSO, or CCl4 solutions of the potassium or tetrabutyl ammonium salt of 4-benzyloxy-1-hydoxy-2,2,6,6-tetramethyl piperidine. The resulting products are 4-benzyloxy-2,2,6,6-tetramethyl piperidine. 1-oxyl, KO2, or Bu4NO 2 in 90% yields. Alkali superoxides are slightly soluble in organic solvents. Their solubility can be described as moderate only in DMSO. However, in the presence of crown ether, the solubility of superoxides in DMSO becomes excellent. The same applies to benzene as a solvent (Morkovnik & Okhlobystin 1979). Another chemical reaction of oxygen giving the superoxide ion consists of a single-electron transfer to oxygen from Fe1 sandwich cyclopentadienyl (Cp)–arene complexes, Scheme 1-71: .
[CpFeIArene]° O2 → [CpFeIIArene], O 2 . . II [CpFeIIArene], O 2 M , X ⇔ [CpFe Arene] , X M , O 2 In Scheme 1-71, the reversible reaction is shifted to the right when the anion, X, is larger and the cation, M, is smaller. For example, this shift to the right is 100% in the pres ence of Na, PF 6 and only 35% in the presence of Na , F (Hamon & Astruc 1988). The equilibrium takes place as an exchange reaction between the two ion pairs. Reactions of this type are based on the symbiotic-effect premise: The interaction between a hard cation and a hard anion or between two soft ions is stronger than that between two ions of different types.
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One method of chemical generation of the superoxide ion deserves a short description because of the simplicity of the reagents employed. The ion is generated from oxygen in alkaline aqueous solutions of sodium dithionite, Na2S2O4. The reaction follows Scheme . 1-72 below; the concentration of O 2 ion in the oxygenated solutions depends on that of . SO 2 at any moment of the reaction (Wang & Chen 1997). .
S2O 2 4 ⇔ 2SO 2 . . 2SO 2 O2 ⇔ 2SO2 O 2
The reactions of 1-hydroxy- and 1-aminonapthoquinones with O2 open one significant feature of superoxide ion formation. This ion forms a van der Waals complex with an. other product, a semiquinone. Namely, hydrogen bonds are formed between O 2 and the OH and NH2 groups of the corresponding semiquinone. As a result, the reaction equilibrium is shifted to the right (Liwo et al. 1997). 1.7.1.A Reactions of Superoxide Ion with Organic H Acids .
The O 2 ion in a solution promotes proton abstraction from a substrate or a solvent. This results in the formation of organic bases, which are conjugated with the appropriate H acids. All H acids with pKa lower than 23 can take part in such proton transfer (Sawyer & Gibian 1979). For this reason, even organic acids (HB), which are weaker than water, enter the exothermic reaction expressed by Scheme 1-73: .
HB 2O 2 → O2 HOO B
If HB is a weak acid, this reaction is rather slow, but it nevertheless takes place. Sub. stances that are unable to react with O 2 can be involved too when the solvents/reactants contain water or other proton-donor admixtures. Thus, benzaldehyde is absolutely resistant to the action of the superoxide ion. However, benzaldehyde does transform into benzylic . alcohol and benzoic acid upon the action of O 2 in the presence of moisture. Hence, a “superoxide” variant of the Cannizzaro reaction is realized (Sawyer & Gibian 1979); see Scheme 1-74: .
2O 2 H2O → O2 HOO OH 2PhCHO OH H2O → PhCOOH PhCH2OH
Proton-containing admixtures in a solvent or in benzaldehyde can act like water. The superoxide ion abstracts such labile protons and generates a conjugated base. The base in its turn abstracts protons from the solvent, acetonitrile, for example. If benzaldehyde is present, it is converted into cinnamyl nitrile according to Scheme 1-75: PhCHO CH2CN H → PhCH(OH)CH2CN → H2O PhCHBCHCN The treatment of benzaldehyde with potassium hydroxide in acetonitrile results in the formation of the same product, i.e., cynnamyl nitrile (Sawyer & Gibian 1979). Thus, in the . presence of water or other proton sources, the O 2 ion forms strong bases as well as oxygen and the peroxide anion. Therefore, many reactions that are ascribed to the superoxide ion are actually reactions with proton donors. These reactions produce effective oxidants (O2 and HOO) and strong bases (OH or B). This kind of base generation finds applications in synthetic practice. Thus, ethyl nitroacetate is a relatively strong acid (in H2O, pKa 5.75; in DMSO, pKa 9.2). Reacting with . O2, this compound transforms into the corresponding carbanion needed for several kinds of synthesis (Niyazymbetov & Evans 1993).
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1.7.1.B Reactions of Superoxide Ion with Organic Electrophiles .
As already noted, O 2 reacts with effective electron acceptors. It is a one-electron reductant of moderate strength. However, the superoxide ion can act as a nucleophile if a sub. strate has a decreased electron affinity. The reactions in Scheme 1-76 between O 2 and tetranitromethane or alkylhalides are pertinent: .
.
O 2 C(NO2)4 → C(NO2)3 NO 2 O2 . . O 2 RHal → ROO Hal . . ROO O 2 → ROO O2 ROO RHal → ROOR Hal
As seen from the scheme, the initial step of the reaction between the superoxide ion and alkylhalides is a bimolecular nucleophilic substitution. Reactions of KO2 with optically active alkylhalides proceed with the configurational inversion. Moreover, the reactivity changes in the same order as that of the usual SN2 reactions (primary secondary tertiary; I Br Cl) (Sawyer & Gibian 1979). Nevertheless, the SN2 scheme does not exclude a possibility of one-electron transfer according to Scheme 1-77: .
.
.
RHal O 2 → [Hal R O2] → ROO Hal .
This mechanism is rather probable: O 2 is an electron donor, and organyl halogenides can, in principle, accept one electron. With the one-electron mechanism, the inversion of . configuration is determined by O 2 attack on the side, which is opposite to the halogen atom in RHal (Morkovnik & Okhlobystin 1979). The direction of this reaction depends on the nature of the solvent. In pyridine, benzene, and DMF, the main product is alkyl peroxide. In DMSO, an alkyl carbinol is the main product (Sawyer & Gibian 1979). Obviously, the aforementioned intermediary product ROO reacts faster with the solvent Me2SO than with the substrate RHal, Scheme 1-78: ROO Me2SO → MeS(O)(OOR)Me → Me2SO2 RO Alkyl tosylates and mesylates are cleaved upon the action of KO2 in DMSO and give rise to corresponding alcohols. This reaction also proceeds with inversion of configuration at carbon atoms. Such process may be of importance in prostaglandine chemistry (Morkovnik & Okhlobystin 1979). . Reactions of O 2 with esters and halogenides of carbonic acids also pass nucleophilic substitution as an initial step (Sawyer & Gibian 1979). Final products are acyl peroxides or carbonic acids, Scheme 1-79: .
.
1 2 R1COOR2 O 2 → R COOO R O . . 1 R1COOO O 2 → R COOO O2 1 1 2 1 R COOO R COOR → R COOOCOR1 R2O . 1 R1COO-OOCR1 2O 2 → 2R COO 2O2
The reaction of superoxide ion with carbon tetrachloride is important for olefin epoxidations (Yamamoto et al. 1986). The first step of the reaction consists in the formation of the trichloromethyl peroxide radical, Scheme 1-80: .
.
O 2 CCl4 → Cl Cl3COO
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The trichloromethyl peroxide radicals of this scheme oxidize electron-rich olefins. The latter give corresponding epoxides. This peroxide radical is a stronger oxidizing agent than the superoxide ion itself (Yamamoto et al. 1986). Hence, nucleophilic reactions of the superoxide ion are very typical. This ion can be compared with the thiophenoxide and thiocyanate ions with respect to nucleophilicity. The . cause of such high nucleophilicity lies in a so-called effect: in O–O ion, an attacking . site (O ) adjoins directly to a site (O ) with a significant electronegativity. Such an effect usually confers special activity to nucleophiles. The effect can be additionally enhanced by including the O–O group in sulfenate, sulfinate, or sulfonate molecules (Oae et al. 1981). Scheme 1-81 presents typical examples: .
.
RSSR O 2 → RSOO RS .
RSOO O 2 → RSOO O2
Like the trichloromethyl peroxide radical, peroxothio compounds can perform even . nucleophilic oxygenation of substrates that are inert to O 2 in aprotic solvents. For example, stilbene is not changed in dry benzene containing 18-crown-6-ether and KO2. In the presence of diphenylsulfide, however, the interaction does take place and results in the formation of stilbene epoxide. According to Oae and co-authors (1981), stilbene initially . . gives PhCH(OO)CH Ph anion radical adduct. Abstraction of O from the adduct leads to stilbene epoxide with 40% yield (Oae et al 1981). To sum up, the superoxide ion has high basicity and nucleophilicity. It can react as a reducer and cannot serve directly as an oxidant. 1.7.1.C Superoxide Ion in Reactions with Biological Objects .
There are several enzymes that convert O2 into O 2 . The superoxide ion is produced in many biological reactions and especially in respiration (see the review by Tselinski et al. 2001). Also, there are many indications that this anion radical is particularly toxic to cells and can ultimately have deleterious effects on the health and well-being of certain individuals; see reviews by Lang & Wagnerova (1992) and Faraggi & Houee-Levin (1999). More. over, the hydroxyl radical (which can be formed from O 2 by the well-known Haber-Weiss reaction) also has very high toxicity (Sies 1986). The superoxide ion in general does not appear to be intrinsically bactericidal (Hurst . 1997). However, phagocytic cells have the capacity to enzymatically generate NO . Nitric oxide is also not bactericidal at physiologically relevant concentration levels, but reacts . very rapidly with O 2 to form the peroxonitrite ion (ONO 2 ). [The bimolecular rate constant . for this reaction is the largest yet recorded for any reaction of O 2 (Huie 1993)]. Peroxonitrite ion is a powerful, relatively long-lived oxidant with bactericidal capabilities (Hurst 1997). Another route of superoxide anion radical transformation lies in its disproportiona. tion upon superoxide dismutase catalysis (2O 2 2H → O2 H2O2). . Metal cations are present in many biological objects; O 2 is known to coordinate to . the metal ion, yielding the corresponding complex. The binding energies of O 2 with divalent metal ions (including Mg2) have been determined as 0.5–0.7 eV (Fukuzumi & Ohkubo 2000). Being bonded into the complex, the metal cation is reduced easily at more positive potentials. This means that O2, the most important biological oxidant, really can act as a catalyst rather than as an oxidant in the biological reaction of electron transfer (see Fukuzumi et al. 2001).
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1.7.1.D Superoxide Ion–Ozone Anion Radical Relation Ozone is a metastable compound, and its decay requires the activation energy of ca. 105 kJmol to yield atomic and molecular oxygen. However, the ozone anion radical in its turn splits up even more readily. Successful alkene ozonolyses are probably initiated by alkenecatalyzed decomposition of O3. Some reversible single-electron transfer to O3 is possible as a first weakly endothermic step. The ozone anion radical initially formed is the extremely reactive species and should oxygenate alkene cation radicals very quickly and exothermally through mono-oxygen exchange. This process releases the superoxide ion (Schank et al. 2000 and references therein), Scheme 1-82: .
.
.
.
CBC O3 ⇔ CM C O3 → CMC(O ) O2
Depending on the alkene cation radical nature, open-chain oxygenation and epoxidation take place as well as the formation of other trivial ozonolysis products. Alkylaromatic compounds are also oxidized by ozone via the ion radical mechanism. Ethylbenzene, for example, undergoes ozone attack on the ring (80%) and on the alkyl group (20%). According to kinetic studies, the ozone consumption obeys the chain law (Galstyan et al. 2001). The first of the reaction steps in the amine–ozone interaction also consists of oneelectron transfer from the amine to the ozone, with the formation of the corresponding cation and anion radicals. The ozone anion radical has been revealed at low temperatures. Formation of the superoxide ion and the amine nitroxide are the understandable results of the reaction (Razumovskii & Zaikov 1984 and references therein). 1.7.2 Atomic Oxygen Anion Radical There are transformations of organic oxygen-containing ion radicals, which develop through the elimination of an atomic oxygen anion radical. An example of such a transformation is the stilbene epoxidation mentioned earlier. The atomic oxygen anion radical becomes an acting species when hydrogen peroxide or some other source of hydroxy radicals is employed in water media with pH 12 (Vieira et al. 1997). Therefore, properties of the atomic oxygen anion radical should be described, albeit concisely. . . The properties of O , in contrast to these of the conjugated acid HO , are still poorly studied. The review by J. Lee and Grabowsky (1992) describes mass spectrometry of this . radical. In conditions that are usual for synthetic organic chemistry, O displays less reac. . . tivity than its acid OH. Oxidation of acetone by radicals O and OH leads to quite different products (Morkovnik & Orhlobystin 1979); see combined Scheme 1-83: .
.
MeCOCOOH ← ( HO ) M Me2CO M ( O ) → MeCOOH .
Addition of O to double bonds and to aromatic systems was found to be quite slow. . Simic and co-workers (1973) found that O reacts with unsaturated aliphatic alcohols, es. . pecially by H atom abstraction. As compared to O , HO reacts more rapidly (by two to three times) with the same compounds. In the case of 1,4-benzoquinone, the reaction with . O consists of the hydrogen double abstraction and leads to the 2,3-dehydrobenzoquinone anion radical (Davico et al. 1999 and references therein). . Christensen and co-authors (1973) found that O reacts with toluene in aqueous solution to form benzyl radical through an H-atom-transfer process from the methyl group. In . general, the O anion radical is a very strong hydrogen atom abstractor, which can withdraw a proton even from organic dianions (Vieira et al. 1997).
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It is important to emphasize that the atomic oxygen anion radical plays a role in cat. alytic oxidation that occurs on various oxide surfaces. For instance, O reacts with methane at room temperature over various metal oxides (Lee & Grabowsky 1992). Upon solid catal. ysis, O is more reactive toward alkanes and alkenes than are other ionic oxygen species. . Iwamoto and Lunsford (1980) have postulated that O is the active oxygen species that oxidizes benzene to phenol on V2O5/SiO2 at 550°C with 70% selectivity at 10% conversion. The product is formed just on oxide surfaces, before leaving the solid (Azria et al. 1988). 1.7.3 Molecular Oxygen Cation Radical .
Although oxygen has a very high oxidation potential (4.7 V), it can be oxidized to O 2 under special conditions, for example, by means of the ultramicroelectrode technique (Cassidy et al. 1985). This cation radical is of course a very strong oxidant. Its stabilized forms . . are O 2 SbF 6 and O 2 AsF 6 . At present, however, these salts are not considered to contain the oxygen cation radical. Although textbooks (e.g., Cotton & Wilkinson 1988 p. 462) de. scribe O 2 as a firmly established ionic form of oxygen, this is based on the characteristics of O2MF6 (M P, As, Sb, or Pt) and the as. sumption that these compounds are salts [(O 2 )MF 6 ]. Such an ionic formulation is unsupported by their chemical and physical properties, and is inconsistent with the ionization potentials, electron affinities, and electronegativities of the constituent atoms. The more accurate description of O2MF6 molecules is a covalent adduct (OOF)MYF5. (Sawyer 1991) In any event, O2SbF6 or O2AsF6 act as a very effective one-electron oxidant with regard to aromatic amines, nitrogen- and sulfur-containing heterocyclic compounds (Dinnocenzo and Banach 1986, 1988, 1989), as well as perfluorobenzene, benzotrifluoride, and perfluoronaphthalene (Richardson et al. 1986). The oxidation proceeds in freon (CHCIF2) at temperatures from 115 to 145°C, and the formation of the organic cation radicals has been firmly established. 1.7.4 Carbon Dioxide Anion Radical Carbon dioxide, CO2, is a typical component of the gaseous environment for reactions in air or in the presence of air traces. Therefore, both interactions between CO2 and organic . products of an electron transfer as well as reactions of CO 2 with uncharged molecules of organic compounds should be considered. Interaction of CO2 with organic anion radicals leads, as a rule, to carboxylic acids; . . CO 2 anion radicals are not formed. Even such a one-electron reductant as O 2 in aprotic . medium simply adds to CO2. The addition product, in turn, accepts one electron from O 2 (Sawyer & Gibian 1979). The total result consists of the formation of an aprotic equivalent of peroxycarboxyic acid according to Scheme 1-84: .
.
.
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CO2 O 2 → OO-CO 2 O 2 → OO-CO 2
Vacuum UV irradiation of aqueous solutions containing formate is one of the meth. . ods to generate CO 2 . However, this method leads to the excited anion radical CO 2 . Such an excited anion radical acquires the ability to transfer an unpaired electron to nitrobenzene, benzoic acid, or benzaldehyde (Rosso et al. 2001). Radiolysis of water (and water solutions) in air produces carbon dioxide anion radicals (Morkovnik & Okhlobystin 1979). Hydrated electrons, which are generated during the
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radiolysis, are captured with CO2. The carbon dioxide anion radical transforms aliphatic alcohols into -hydroxyl acids according to Scheme 1-85: .
.
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H2O -rays → H OH eeq H; eaq CO2 → CO 2 . . RCH2OH OH → RCH OH H2O . . RCH OH CO 2 → RCH(OH)CO 2 .
Reaction of OH with either CO or HCOOH in water solution also results in the for. mation of CO 2 (Flyuht et al. 2001). Another route to the carbon dioxide anion radical is reduction of hydrogen peroxide with Ti3 in the presence of formate ions (Morkovnik & Okhlobystin 1979), Scheme 1-86: .
Ti3 H2 O2 → Ti4 OH OH; .
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OH HCOO → H2O CO 2 .
The solvated CO 2 anion radical has been observed both in the gas phase and in condensed matter (Holroyd et al. 1997) and has been well characterized by ESR spectroscopy . (Knight et al. 1996). Being solvated, CO 2 anion radicals form complexes that yield quasifree electrons upon photoexcitation. Gas-phase studies (Saeki and others 1999) and ab initio calculations (Tsukuda et al. 1999) indicate that static ion–dipole interactions stabilize . the [(CO2)n–m(ROH)m] type of small clusters. In supercritical carbon dioxide, monomers and dimers of water, acetonitrile, and alcohols also form metastable complexes (Shkrob & Sauer 2001a,b). Such complexation should be taken into account in studies of electron. transfer kinetics in reactions with the participation of CO 2 . . Disproportionation of CO 2 leads to carbon monoxide and a carbonate, while dimerization results in oxalate formation; see the left and right directions of Scheme 1-87: .
2 O2CMCO 2 ← 2CO 2 → CO 3 CO
The carbon dioxide anion radical usually plays a role as one-electron reductant; in . DMF its E0 1.97 V (Amatore & Saveant 1981). In the gas phase, solitary CO 2 loses one electron with an exothermic effect of ca. 45 kJmol1 (Compton et al. 1975). . Interactions of CO 2 with alkyl halogenides or with compounds containing halogen–nitrogen bonds take place as dissociative electron captures. As a result, free radicals are formed according to Scheme 1-88: .
. CO 2
.
CO 2 RHal → CO2 R X . MeCONHBr → CO2 MeCONH Br
The carbon dioxide anion radical was used for one-electron reductions of nitrobenzene diazonium cations, nitrobenzene itself, quinones, aliphatic nitro compounds, acetaldehyde, acetone and other carbonyl compounds, maleimide, riboflavin, and certain dyes (Morkovnik & Okhlobystin 1979). This anion radical reduces organic complexes of CoIII and RuIII into appropriate complexes of the metals in the valence 2 state (Morkovnik & Okhlobystin 1979). In the case of the pentammino-p-nitrobenzoato-cobalt(III) complex, the electron-transfer reaction passes a stage of the formation of the Co(III) complex with the p-nitrophenyl anion radical fragment. This intermediate complex transforms into the final Co(II) complex with the p-nitrobenzoate ligand as a result of an intramolecular electron transfer. Scheme 1-89 illustrates this sequence of transformations: .
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III 2 CO → CO2 [ O2NC6H4COOCoII(NH3)3] → 2 [O2NC6H4COOCo (NH3)3] [O2NC6H4COOCoII(NH3)3]
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Hence, despite the presence of the central atom CoIII, an initial electron transfer is directed not to this center, but to the p-nitrophenyl fragment. Obviously, this fragment, and not the metallic ion, provides an orbital that is symmetrically appropriate and low lying. A similar situation takes place in a reductive cleavage of arylsulfonic salts (see Chapter 3). . A double bond can react with CO 2 (Morkovnik & Okhlobystin 1979); Scheme 1-90 represents one typical example: .
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1 2 1 2 CO 2 R R CBNOH → R R C(COO )NHO
As seen, the transformation into the anion radical is a very effective way to activate carbon dioxide. 1.7.5 Sulfur Dioxide Anion Radical .
0 The couple SO2/SO 2 is characterized by E 0.26 V (in water—Bradic & Wilkins 1984). Such potential falls into the region of redox-protein activity. This suggests that the known damaging action of gaseous SO2 does not correspond to the sulfuric acid formation, but is stipulated by SO2 participation in biological electron transport. Sulfurous acid also causes detrimental effects by irritating the mucous membranes. Interestingly, the resistance of different organisms to sulfur dioxide correlates with the relative significance of redox processes in these organisms. For instance, plants are among the organisms that suffer the most from the damaging actions of SO2. However, different members of the plant kingdom manifest different levels of sensitivity to the gas. Among trees, the most sensitive are fir and pine, while the least sensitive are birch and oak. The highest sensitivity to SO2 among flowering plants is attributed to the rose. In humans, inhalation of air containing more than 0.2% of sulfurous gas causes shortness of breath and confusion. Prolonged and repeated exposure to SO2 leads to loss of appetite, constipation, and inflammatory diseases of the respiratory tract. The levels of sensitivity to the gas differ from individual to individual. Sensitivity usually diminishes with prolonged exposure. It is explained by the induction of the cytochrome defense that allows for a normal electron transport (up to a certain time); see Prousek (1988). It is likely that the reductive properties of dithionites (salts of dithionous acid H2S2O4) are partially ex. plained by S2O 2 4 ⇔ 2SO 2 dithionite reversible dissociation. Both the dithionite ion and the sulfur dioxide anion radical can act as reducers. The former can be oxidized into sulfite or sulfate; the latter can lose an electron, thus transforming into a volatile dioxide, SO2. This is why the use of a stable sodium dithionite upon dye vatting or dye printing is always complicated with sulfurous gas evolution. Dissociation of the dithionite ion can proceed both in aqueous and in nonaqueous media. There is a special equation to determine the average values of the equilibrium constants . 2 (Keq) of the formation of SO (Lough & McDonald 1987), namely, Keq 2 from S2O 4 . 2 2 . [SO 2 ] /{[S2O 4 ]0 0.5 [SO 2 ]}. In DMF, DMSO, and acetonitrile, the Keq values are 42.4, 11.3, and 40 mM at 25°C, respectively. These values exceed Keq in water (1.4 106 mM) by seven orders of magnitude. The salt (Et4N)2S2O4 is very soluble in water and in organic solvents. In nonaqueous media, its dissociation depends strongly on the moisture content. For the DMF/H2O system, Keq values fall from 42.4 to ~104 mM, along with water content increasing from 0 to 20 vol %. Dithionite ions (S2O 2 4 ) are better stabilized in solvents of high polarity and in water than in systems of low polarity (Lough & McDonald 1987).
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Organic solvents are typical for organic electron-transfer reactions, and these sol. vents are usually polar ones. The anion-radical SO 2 is a particularly interesting oneelectron donor. Because this anion radical is formed more easily in less polar solvents, it allows widening the range of solvents acceptable for organic electron-transfer reactions. .
1.7.6 Sulfite Anion Radical, SO 3 .
The sulfite anion radical (SO 3 ) is obtained from sodium sulfite upon the action of titanium trichloride in water in the presence of ethylene diamine tetraacetic acid (as a complexing agent) and hydrogen peroxide (Bradic & Wilkins 1984). Scheme 1-91 below shows this . type of SO 3 generation. .
Ti3 H2O2 → Ti4 OH OH; . . OH SO 2 3 → OH SO 3 .
According to the scheme, the formation and further reactions of anion radical SO 3 take place in an alkaline medium. Therefore, there are some restrictions to the electrontransfer reactions from this anion radical to acceptors. For instance, aliphatic nitro com. pounds react in alkaline media in aci forms. They add SO 3 and form new anion radicals; Scheme 1-92 (Bradic & Wilkins 1984): .
.
RCHBNOO SO 3 → RCH(SO3)NO 2 .
In order to generate SO 3 in acidic media, the reaction between sodium bisulfite and cerium ammonium nitrate should be employed. In (NH4)2Ce(NO3)6, cerium has the oxidation state 4. In acids, valence 4 cerium salts are strong oxidants. With the sulfite ion, the oxidation follows Scheme 1-93: .
3 Ce4 SO 2 SO 3 → Ce 3 .
This method of SO 3 generation is used to perform the addition of the sulfite anion radical to unsaturated compounds, Scheme 1-94: .
.
RCHBCHR SO 3 → RCH(SO 3 )CH R
This addition reaction is also applicable to substrates containing C⬅C and CBS bonds (Ozawa & Kwan 1985). In cases of nonsymmetrically substituted double bonds, the . less shaded carbon atom is attacked by SO 3 anion-radical. For example, acrylic, 3,3. dimethyl acrylic, and crotonic acids react with SO 3 at the expense of the ethylenic carbon, which is remote from the carboxylic group. Because these unsaturated acids are stable in . alkali medium, it was possible to examine their anionic forms in the reaction with SO 3 . It . turned out that the carboxylate group prevents SO 3 from entering the geminal position. In alkaline medium, however, a competition is possible for an unsaturated substrate between . . radicals SO 3 and OH (the last originates from hydrogen peroxide, see earlier). The afore. . mentioned restrictions for SO 3 addition are not relevant to OH addition. In contrast to an. . ion radical SO 3 , radical OH bears no charge. Therefore, there is no electrostatic repulsion between the entering reactant (the hydroxy radical) and the substrate carboxylate group. . For the addition reactions of SO 3 to CBC bonds (which can be performed both in . alkaline and in acidic solution), the following regularity was established: SO 3 anion radical is more active in acidic than in alkaline medium (Ozawa & Kwan 1985). . The reactions of SO 3 anion radicals with organic and inorganic compounds have commanded considerable interest because of the role of these anion radicals in grievous
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health and technological problems. Indeed, biochemical and industrial sources produce 2 . SO 2 3 . SO 3 is often transformed into SO 3 during atmospheric reactions. Both of them are pollutants whose activity in the environment leads to illness (Reed et al. 1986). Although the toxicological role of inorganic SIV and SV is now accepted, recent studies have shown that the outer-sphere reactions of electron transfer are not the sole one. There are many . cases of oxidative addition of SO 3 to unsaturated compounds. Recently, even examples of . the substitutional sulfonation have been described for SO 3 reactivity (Dutta & Ferraudi 2001). .
1.7.7 Sulfate Anion Radical, SO 4
Sulfate anion radical is formed during UV irradiation or heating to 80–100°C of persulfates . M2S2O8 in aqueous solutions. This results in S2O2 8 ⇔ 2SO 4 dissociation (Minisci et al. 1983; Dias & Vieira 1996). Conventional oxidants can, in principle, accept electrons from persulfate ions to form the sulfate anion radical, but the reaction is extremely slow. For instance, M2S2O8 in H2O2 solution gives M2SO4 for two months at O°C (Skogoreva & Ippolitov 1988). The presence of transition metal salts (for example, salts of Ag or TI3) promotes . the reaction (Minisci et al. 1983). In this case, each S2O 2 8 anion generates only one SO 4 according to Scheme 1-95: .
3 S2O 2 → Ti4 SO 2 8 Ti 4 SO 4
The nature of a transition metal is not essential for this redox reaction. However, one . of the reaction products, namely, anion radical SO 4 , can be complexed by a transition . metal in a higher oxidation state. This leads to some stabilization of SO 4 and increases its effective concentration. In other words, further reactions with organic substrates are facilitated (Fristad & Peterson 1984). Cuprous and ferrous salts are preferable. Sometimes, a transition metal salt is deliberately added to a mixture of a substrate and a persulfate salt (Dobson et al. 1986). The free or metal-coordinated sulfate anion radical reacts with an organic substrate, giving rise to a substrate cation radical (Minisci et al. 1983; Itahara et al. 1988; Telo & Vieira 1997). One typical example is the reaction between . toluene and SO 2 in Scheme 1-96: .
.
2 ArCH3 SO 4 → SO 4 (ArCH3)
.
.
Hence, there is a difference between anion radicals SO 4 and CO2 . While the latter is an one-electron reductant (see earlier), the former is a one-electron oxidant. A substrate cation radical is often able to expel a proton and transforms into a radical. The latter regenerates the starting metallic ion. The whole reaction becomes catalytic in respect of the metal; see Scheme 1-97: .
.
(RCH3) → H RCH 2 . RCH 2 Cu2 → Cu RCH 2 RCH 2 H2O → H RCH2OH One-electron transfer from a substrate to the sulfate anion radical mostly follows diffusion rates. For instance, rate constants of one-electron oxidation of benzene and anisole . 9 9 1 1 with SO 4 are equal to 3 10 and 5 10 Lmole sec , respectively (Goldstein & McNelis 1984).
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The sulfate anion radical is not a very strong hydrogen acceptor. It acquires an atomic hydrogen from organic substrates at significantly smaller rates a compared with the rates for one-electron oxidations. For instance, dehydration rate constants are 107, 106 and 105 Lmole1sec1 for methanol, tert-butanol, and acetic acid, respectively (Goldstein & McNelis 1984; Zapol’skikh et al. 2001). Such a peculiarity is very important for the selectiv. ity of ion radical syntheses with the participation of SO 4 . If the sulfate anion radical is bound to the surface of a catalyst (sulfated zirconia), it is capable of generating the cation radicals of benzene and toluene (Timoshok et al. 1996). Conversion of benzene on sulfated zirconia was narrowly studied in a batch reactor under mild conditions (100°C, 30-min contact) (Farcasiu and co-workers 1996; Ghenciu and Farcasiu 1996a, 1996b). The proven mechanism consists of one-electron transfer from benzene to the catalyst, with the formation of the benzene cation radical and the anion radical of sulfate on the catalytic surface. This ion radical pair combines to give a surface combination of sulfite phenyl ester with reduced sulfated zirconia. The ester eventually gives rise to phenol, Scheme 1-98. Coking is not essential for reaction 1-98. To regenerate the catalyst, it is enough just to oxidize the worked-out catalyst. The described reactions of sulfate anion radical characterize it as a radical of strong . oxidation ability (E0 ⬇ 2.60 V, in water—Balej 1984). Like organic anion radicals, SO 4 has some affinity to the hydrogen radical. With cyclic alkynes, the sulfate anion radical acts as an oxygen-transfer reagent, transforming the alkyns into , -epoxy compounds (Wille 2000). In this sense, the sulfate anion radical can also be considered a donor of atomic oxygen in solution. The reaction . . leads to the release of SO 3 , which is significantly less reactive than SO 4 (Muller et al. 1997).
SCHEME 1-98
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.
In contrast to organic anion radicals, SO 4 possesses high electrophilicity and is prone to addition reactions (Davies & Gilbert 1985), Scheme 1-99: .
.
CH2BCHCH2OH SO 4 → O3SOCH2CH CH2OH . . MeCHBCHCOOH SO 4 → MeCH CH(OSO 3 )COOH . . CH2BCHCH2COOH SO 4 → O3SOCH2CH CH2COOH
This reaction leads to the adducts that are observable by means of the ESR method. In some cases (maleic, fumaric acids), ESR spectra can be recorded at high-enough pH val. ues only. Being a strong electrophile, the anion radical SO 4 is more active toward a carboxylate ion than to neutral molecules of unsaturated acids. Anions of saturated acids react with sulfate anion radicals, giving rise to carboradicals (Eberson et al. 1968), Scheme 1-100: .
.
.
2 Me3CCOO SO Me3CCOO 4 → SO 4 . . Me3COO → CO2 Me3C
This reaction resembles the decarboxylation of carboxylates during electrode oneelectron oxidation. The Kolbe electrochemical reaction also consists of one-electron oxidation, decarboxylation, and culminates in dimerization of alkyl radicals formed intermediately. In Reaction 1–100, the sulfate anion radical, and not the anode, acts as a one-electron oxidant. In this case, the caboradical dimerization is hampered. The radicals can be used in preparative procedures. One typical example is alkylation of heterocyclic nitrogen bases (Minisci et al. 1983). The difference between the Kolbe electrochemical reaction and the reaction with the help of a “dissolved” electrode (the sulfate anion radical) deserves some explanation. The concentration of the one-electron oxidation products in the electrode vicinity is significantly higher than that in the bulk of solution. Therefore, in the case of anode-impelled reactions, . the dimerization of radicals produced from carboxylates proceeds easily. Noticeably, SO 4 secures the single-electron nature of oxidation more strictly than an anode. In electrode reactions, radical intermediates can undergo further oxidation if potentials of capture of the first and the second electrons are more or less close. In contrast, the homogeneous reactions . with the participation of SO 4 are characterized by fixed potentials in respect to the reactant and an organic substrate. The route of homogeneous reactions is strongly dependent upon the ratio of reagent concentrations. If the ratio is small, the probability of the secondary reaction is low. The distinction can be illustrated with the oxidation of RCOO: It can give R . and after that, at once, R in an anodic process. In the case of oxidation with the aid of SO 4 , . R becomes the sole product of the reaction (Eberson et al. 1968). . Another advantage of the SO 4 pathway for carboradical generation is its high selectivity. As already noted, this type of oxidation allows one to avoid (practically) the hydrogen abstraction. In the case of alkyl carboxylates, the rate of one-electron oxidation of the carboxyl group exceeds the rate of homolytic disruption of the C–H bond in an alkyl rest by 100 times (Davies et al. 1985). Cleavage of persulfate in the presence of carboxylates is a simple and reliable method of getting alkyl radical for synthesis. 1.7.8 Hydroxide Anion This anion is a typical reagent for organic hydroxylation. Many such reactions have ion radical mechanisms according to which an inorganic anion transfers an electron to an organic
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substrate. It is important to underline the diffuse nature of the excess electron on the OH ion, which allows molecular orbitals of the ion and the substrate to overlap, even with a long distance between them (Takahashi et al. 2001). For example, in the reaction between the hydroxyl anion and o-dinitrobenzene, the OH ion is proved to be an electron donor; . see Section 4.3.4.B. When OH transforms into OH, o-dinitrobenzene gives the anion radical. The o-dinitrobenzene anion radicals are stable. They come out into the reaction volume and meet there with new hydroxyl anions (Abe & Ikegami 1978), Scheme 1-101: .
.
C6H4(NO2)2 OH → OH C6H4(NO2)2 . . C6H4(NO2) 2 OH → C6H4(NO2)(OH)NO2
.
.
C6H4(NO2)(OH)NO 2 C6H4(NO2)2 → C6H4(NO2) 2 C6H4(NO2)(OH)NO2 C6H4(NO2)(OH)NO2 → NO 2 C6H4(OH)NO2
There are several cases of hydroxylation according to the “hidden radical” mechanism, within a solvent cage. As assumed (Fomin & Skuratova 1978), hydroxylation of the . anthraquinone sulfonic acids (AQ–SO3H) proceeds by such a pathway, and OH radicals attack the substrate anion radicals in the solvent cage. Anthraquinone hydroxyl derivatives are the final products of the reaction. In the specific case of dimethylsulfoxide as a solvent, hydroxyl radicals give complexes with the solvent and lose their ability to react with the antraquinone sulfonic acid anion radicals (Bil’kis & Shein 1975). The reaction is stopped just after anion radical formation, Scheme 1-102: .
.
AQ-SO 3 OH → OH (AQ-SO 3 ) . . . Me2SO OH → Me2SO OH → Me HOSOMe
The hidden-radical mechanism (also dubbed as the mechanism of biradical origin) cannot, however, take into account all the peculiarities of the reaction participants. The re. dox potential of the OH/ OH couple is equal to 0.9 V with regard to the saturated calomel electrode in acetonitrile (Tsang et al. 1987). In conventional reference to the normal hydrogen electrode and in water, this potential had been determined as 1.9 V (Berdnikov & Bazhin 1970) or as 1.3 V in acetonitrile (Eberson 1987, Chap. 4, p. 62). These values are too positive to reduce the majority of organic acceptors. Of course, there are some cases of single-electron transfer from the OH ion to an organic acceptor. All these cases, however, refer to substrates with very strong electron affinity, such as tetracyanoethylene and dinitrobenzene (Blumenfel’d et al. 1970). Quinones, ketones, and other substrates have less affinity. In the ground (nonexcited) states, they are unable to capture an electron from the OH ion (Sawyer & Roberts 1988). The described reactions of quinone hydoxylation, as well as electron transfer from hydroxyl anion to quinones, are in general accelerated by light irradiation. Wavelengths are up to 500 nm. The degree of acceleration is higher with greater intensity of irradiation and with an increase in the concentrations of quinone and hydroxide (Blumenfel’d et al. 1970). Direct (outer-sphere) electron transfer takes place from hydroxide to an excited quinone molecule. In the excited molecule, one of the electrons of a double-occupied orbital populates a vacant orbital. The “derelict” orbital is located at the lower level of energy than the vacant orbital. Therefore, the electron transfer from hydroxide to a quinone becomes energetically favorable. This corresponds to an enhanced electron affinity and lifts the prohibition for the electron transfer. . It should be pointed out that the possibility of an OH → OH transformation with one-electron transfer onto a substrate depends on the further fate of the resulting hydroxyl
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radical. The latter is a strong oxidant, and it must be taken out of the reaction sphere in order to shift the electron transfer equilibrium to the right. One opportunity to organize the . shift to the right consists of recombination of two OH radicals, which results in H2O2 formation. As a rule, hydrogen peroxide is formed in reactions between hydroxide and very strong acceptors (Endo et al. 1984). However, an attack of hydroxyl anion upon an unsaturated site of a substrate (an aromatic fragment, the carbonyl group, etc.) is more typical and more interesting chemically. This results in a -complex formation. Such a complex can regroup into a product of hydroxylation. The complex can also decay, with the forma. tion of substrate anion radicals and OH radicals. In short, the electron transfer proceeds in an inner-sphere manner, and the aforementioned restrictions (in potential differences, in . OH oxidation activity) vanish. This inner-sphere mechanism is assumed to be the most frequent. One-electron outer-sphere transfer with radical and ion radical formation (on the one hand) and coupling of an activated substrate with hydroxyl anion with no electron transfer (on the other hand) are two rare extremes. The last but not least peculiarity of the reaction of OH with electron acceptors consists of its dependence on solvation effects. Solvation energy has an influence on the ionization potential of OH and on electron affinity of a substrate as well. For instance, the OH ion is a stronger base and a more effective electron donor in acetonitrile or dimethylsulfoxide than in water. Its solvation energy is lower by 80–100 kJmol1 in these organic solvents than in water (Sawyer & Roberts 1988). This fall in the solvation energy corresponds to a relief of hydroxyl anion oxidation by almost 1 V (Sawyer & Roberts 1988). As for the other participants in the equilibrium reaction of electron transfer (a neutral substrate, . its anion radical, and the OH radical), their solvation susceptibility is not so essential. The main effect should be attributed to OH. Of course, the ability of OH to be an electron donor depends not only on its solvation energy but also on the energy of covalent binding of the hydroxyl group in an adduct with a substrate or in a final product of hydroxylation. 1.7.9 Nitrosonium and Nitronium Ions The nitrosonium ion, NO, is a strong oxidant, with E0 1.51 V vs. the normal hydrogen electrode. It reacts with organic substrates of E0 1.7 V as an outer-sphere acceptor of one electron. In contrast, the nitronium ion, NO 2 , is a weaker oxidant than the nitrosonium ion, although its redox potential (E0 1.56 V vs. the normal hydrogen electrode) is close enough to that of the nitrosonium ion. The cause of such a sharp decrease in acceptor properties is the necessity of energy consumption to change the NO 2 structure upon one-elec. tron transfer. The transformation of NBO into NMO results in an elongation of the nitrogen–oxygen bond by 0.01 nm, which consumes 42 kJmol1 of energy (Eberson & . Radner 1987). Meanwhile, the NO 2 ion is planar, NO 2 is bent (O–N–O angle is equal to 134.3°). The bending consumes considerable energy, estimated to be equal to 218 kJmol1 (Eberson & Radner 1987) or even 587 kJmol1 (Lund & Eberson 1997). It is the factor that decreases the NO 2 electron affinity. Only very oxidizable substrates can participate in the outer-sphere electron transfer to NO 2 . The observed cases of the generation of substrate cation radicals and the nitrogen dioxide radicals are stipulated by the inner-sphere mechanism. Such a mechanism involves the nucleophilic attack of the nitronium ion upon a substrate, the formation of a tetrahedron complex, and the disintegration of the complex. The disintegration results in products, which corresponds formally to one-electron transfer. However, the reorganization energy is lower for the inner-sphere mechanism, since the bending of the nitronium ion is no longer required.
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Thus, there is a great difference in electron-transfer reactivity between the nitronium and the nitrosonium ions, despite the closeness in their redox potentials. Closeness in the electrode potentials supposedly reflects the peculiarities of electrochemical reactions. The liquid-phase reactions discussed here have their own string of distinctions. Electrode processes are heterogeneous. There is a strong electric field in the near-electrode layer. The field can have an effect on the behavior of the NO 2 ion. It is probable, under the electrode . reaction conditions, that both the linear ion NO 2 and the bent NO 2 radical acquire somewhat similar configurations in the tough part of the double electric layer. Energy consumption for such a small bend is revealed in the potential value only vaguely (cf. 1.56 and . . 1.51 V for the couples of NO 2/NO 2 and NO /NO , respectively). Some authors suggest . that the experimental values determined for NO 2/NO 2 might be influenced by significant . orbital overlap between NO 2 and the Pt electrode surface (Lee, K.Y., et al. 1991). This pathway is energetically favored because the outer-sphere reorganization energy is not needed as much. Of course, a solvent can intervene in all of these phenomena. It should be . 0 noted that E0 (NO 2/NO 2 ) in AN has been reported to be over 0.6 V less positive than E . (NO 2/NO2 ) in sulfolane and nitromethane (Boughriet & Wartel 1989a,b, 1993). In homogeneous liquid-phase reactions, there are no strong electric fields, no influence of a tough electrode layer. Geometric factors begin to play their independent roles. These factors determine the equilibrium state, which governs the degree of conversion of the acceptor (NO 2 , NO ) during one-electron transfer. Moreover, the NO 2 ion can form a complex with a solvent (Ciaccio & Marcus 1962; Hunziker et al. 1971). This causes a marked diminution of the NO 2 reactivity and decreases the nitration rates. The NO 2 and NO reactivity problem is also considered with respect to aromatic nitration; see Section 4.5.4. 1.7.10 Tris(aryl)amine Cation Radicals After Walter (1966) developed the preparation and redox properties of tris(aryl)amine cation radicals, they have found wide application in organic synthesis. For example, the salts of the tris(4-bromophenyl)amine and tris(2,4-dibromophenyl)amine cation radicals (occasionally named Magic Blue and Magic Green) with the hexachloroantimonium anion are stable and widely used for the initiation of diverse reactions that begin from cation radical formation. Chapter 6 considers many such reactions. The salts are obtained as a result of the oxidation of tris(bromophenyl)amines with antimonium pentachloride. The nature and reactivity of inorganic counterions control the stability and redox properties of these organic cation radicals. These aspects are analyzed here. Generation of the tris(bromophenyl)amine cation radical follows Scheme 1-103: .
2 (4-BrC6H4)3N 3SbCl5 → SbCl3 2 (4-BrC6H4)3N SbCl 6 The formation of the SbCl 6 anion can be explained by Scheme 1-104: .
.
2SbCl5 ⇔ SbCl 6 SbCl 4 ; SbCl 4 e → SbCl 4; 2SbCl 4 → SbCl5 SbCl3 V III Anion SbCl 6 is a good oxidizing agent because it can be reduced from Sb to Sb . 0 Organic substrates with E up to 1.5 V vs. normal hydrogen electrode transform into cation 0 radicals in the presence of SbCl 6 . Note: tris(4-bromophenyl)amine has E 1.3 V vs. normal hydrogen electrode in acetonitrile (Reynolds et al. 1974). Such acceptor influence of the anion in the salt provides high stability for the salt. In cases of other anions, the cation radical component decays gradually (both as a dry solid and in a solution) (Eberson & Lars-
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son 1986). The main role in such degradation belongs to the reaction between the cation radical and water (from air or as solvent moisture). Cation radicals of tris(aryl)amines can decompose water at room temperature. The reaction resembles the process of water decomposition with oxidized forms of chlorophyll “a” and consists of the formation of hydrogen peroxide and the evolution of oxygen (Pokhodenko et al. 1984); see Scheme 1-105: Ar3N. H2O → Ar3N H 1⁄2H2O2;
H2O2 → H2O 1⁄2O2
The oxygen and hydrogen peroxide evolving in the reaction provoke degradation of tris(4-bromophenyl)amine. Another route of degradation consists of the gradual dimerization of these cation radicals with bromine depletion (Cowell et al. 1970). However, having the hexachloroantimonate counterion, tris(4-bromophenyl)ammoniumyl is stable (Pokhodenko et al. 1984; Cowell et al. 1970). According to IUPAC rules, ammoniumyl is the name for amine cation radicals. Like all cation radicals, ammoniumyls are sensitive to nucleophiles (to reactants or admixtures). At the same time, 1,1,1,3,3,3-hexafluoropropan-2-ol as a solvent drastically curtails nucleophilic reactivity and provides good integrity of the tris(4-bromophenyl)ammoniumyl at ambient temperatures (Eberson, Hartshorn et al. 1996). In cases when tris(4-bromophenyl)ammoniumyl hexachloroantimonate turns out as a weak one-electron oxidant, its 2,4-dibromo- or even 2,3,4,5,6-hexachloro- analogs are employed (Nelsen et al. 1997). After a one-electron oxidation, these ammoniumyls transform into tri(aryl)amines. They are remarkably nonreactive as nucleophiles. After the electron transfer, the hexachloroanytimonate anion begins to serve as a counterion for newly formed cation radicals of a donor molecule. Sometimes, the hexachloroantimonate ion of the ammoniumyl salt is changed with hexabromocarborane, which brings the anion’s nucleophilicity to nil. For example, the fullerene carbocation (C 76) was isolated as the stable salt with the carborane anion after the reaction of C76 with tris(2,4-dibromophenyl)ammoniumyl that had the hexabromocarborane counterion (Bolskar et al. 1996). 1.7.11 Hexachloroantimonate(V) Salts of Trialkyloxonium Cations These salts belong to a widely used class of alkylating agents; their traditional name is Meerwein’s salts (Meerwein et al. 1937). However, they can be involved in reactions other than alkylation (Boettger et al. 1997). One of these reactions is an oxidation of aromatic donors (Rathore and co-authors 1998a), Scheme 1-106: .
2ArH 3(Et3OSbCl 6 ) → 2(ArH SbCl 6 ) 3EtCl 3Et2O SbCl3
So the triethyloxonium cation is an effective oxidant for the production of aromatic cation radicals, but only as the hexachloroantimonate salt. The analogous (Et3OBF 4 ) or (Bu4NSbCl 6 ) cannot be used to prepare any cation radical. This is consistent with the previously described function of SbCl 6 as an oxidant. The slow release of SbCl5 forms the basis for the efficacy of (Et3OSbCl 6) as an aromatic oxidant. Kochi’s group proposed Scheme 1-107 (see below) for this slow release (Rathore and co-authors, 1998a). (Et3OSbCl 6 ) → SbCl5 EtCl Et2O The transformations of SbCl5 caused by one-electron transfer from an aromatic compound have just been described. If the pure Lewis acid SbCl5 is used, its reactivity is very difficult to control, and single-electron oxidation as well as chlorinations of various aro-
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matic donors can occur readily (Mori et al. 1998). Meanwhile, in the case of (Et3O SbCl 6 ), the slow release of the active monomer SbCl5 occurs. In the case of SbCl5 as such, 2SbCl5 ⇔ Cl4Sb–Cl2–SbCl4 dimerization takes place (Cotton & Wilkinson 1988, p. 395). The dimeric form may lead to electrophilic chlorination according to Scheme 1-108: Cl 4Sb-Cl2-SbCl4 2ArH → 2ArCl 2HCl 2SbCl3 Neither Et3O nor SbCl 6 is individually capable of aromatic oxidation. Therefore, the slow-released monomeric SbCl5 is the active oxidant. This is a reason to consider trialkyloxonium hexachloroantimonates together with other inorganic participants of organic ion radical reactions. 1.7.12 Transition Metal Ions In the sense of electron-transfer reactions, transition metal ions can be the one- or two-electron type. The two-electron ions transform into nonstable states upon unit change of the metal oxidation number. In the outer-sphere mechanism, two-electron transfer is a combination of two one-electron steps. The conditions of the outer-sphere mechanism have been analyzed at length in the literature (Eberson 1987, Chap. 7). An organic substrate must fit a metal ion with respect to redox potentials. The substrate cannot have any hydrogen atom that is prone to depart as a proton. The substrate cannot have any positions that are free from any steric stoppage for a metal attack. The metal ion has to have a relatively high redox potential, and this potential has to be away by no less than 0.5 V from potentials of the next redox transformations of the ion. Ligands surrounding the metal ion cannot be substituted with the substrate in electron-transfer reactions. Otherwise, the reaction system is directed to the route of innersphere electron transfer. The same concern has to do with ligand exchange at the expense of medium components. Participation of a transition metal ion in electron transfer leads, of course, to some changes in redox properties. For the outer-sphere mechanism, it is important that the change of redox potential not affect the reorganization energy of the solvate environment of the metallic ion. It is better to restrict the metal ion–substrate interaction sterically at both metal and substrate sides. In all other cases, inner-sphere mechanisms are at work. These mechanisms include addition and subsequent dissociation. For C–H dissociation, so-called metallocomplex activation takes place. This kind of activation consists of inner-sphere electron transfer according to Scheme 1-109: .
.
R CuII(OAc)2 → RCuIII(OAc)2 → R( H ) HOAc CuIOAc One example of outer-sphere electron transfer is the reaction between dipotassium cyclooctatetraene (K2C8H8) and the cobalt complex of bis(salicylidenediamine) (CoIISalen) (Levitin et al. 1971), Scheme 1-110: K2C8H8 2CoIISalen → C8H8 2(CoISalen)K Although C8H8K2 is a two-electron donor, only a one-electron transfer takes place. Only this type of transfer is permitted because of a difference between potentials of the donor and the acceptor. This difference remains a determining factor of electron exchange because the cobalt atom firmly reserves the ligand environment for itself. As for cyclooctatetraene, it is unable to provide the ligand exchange in this case. The results obtained allowed the solving of the reverse problem, namely, the development of an analytical method
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of determination of Co0Salen2. This deeply reduced dianionic form of the complex results in the metallic sodium reduction of CoIISalen. The dianionic form possesses catalytic properties and is actively used in model studies of biological electron transport. After treatment of CoIISalen with excess sodium in THF, the reduction product was allowed to react with cyclooctatetraene; the consumption of cyclooctatetraene was quantitatively determined by means of the GLC method. If an insufficient amount of C8H8 was used, it transferred into C8H8Na2 completely. If the amount of C8H8 employed was 1 mole per each 2 moles of CoII Salen used to obtain [(Co0Salen)2 2Na], then the degree of transformation of C8H8 into C8H8Na2 was 70%. This means that the “deeply reduced” sample contained only 70% of [(Co0Salen)2 2Na]. In that case, outer-sphere electron transfer looks like a “titration” of the double-charged complex with cyclooctatetraene, Scheme 1-111: C8H8 2[(Co0Salen)2 2Na] → C8H8Na2 2[(CoISalen)Na] Titration according to this scheme showed that the treatment of CoIISalen with excess amounts of Na resulted in the nonquantitative formation of [(Co0Salen)2 2Na]. Thus, catalytic and, especially, kinetic investigations of such a complex have to take into account the presence of CoIISalen and/or (CoISalen) in the samples studied. The described convenient method of quantitative electron transfer in solutions is good at determination of other low-valent metallocomplexes. Of note is the solvent role for electron-transfer reactions between organic substrates and metal ions. For instance, a transformation of CuII complex into CuI complex proceeds better in acetonitrile than in water (Ahrland et al. 1983). Coordination of CuI with acetonitrile is much better than that with water. In water solutions, with no significant share of organic solvents, CuI salt disproportionates completely and yields CuII salts and Cu0 metal. In aprotic solvents, the CuI ion is much more stable. In dimethylsulfoxide, the disproportionation is not significant as long as the concentration of CuI is low. In pyridine, it does not proceed at all, irrespective of the concentration. In acetonitrile, CuI salts are absolutely stable, with no signs of disproportionation or air oxidation. Thus, the role of an organic aprotic solvent consists of the enhancement of stability of a partly reduced form of metallocomplex or ion. The cause of such enhancement is the more effective solvation than in the case of water or another proton-donor solvent. In connection with the data just described, one reasonable question arises: Why does the stabilizing role of a nonwater solvent become significant for CuI salts and remain insignificant for CuII salt? As Myagchenko’s group showed (1989), CuII salts (and salts of HgII, PdII as well) exist in nonwater solutions as polynuclear compounds. Thus, CuCl2 forms lamellar lattices with chlorine chains as bridges between copper atoms. By contrast, CuCl does not form such chain structures. CuI ion is easily and effectively solvated by nonwater solvents and forms stable solvates. As far as (CuCl2)n associates are concerned, their interactions with an organic solvent as well as with an organic substrate in general takes place at terminal groups that are energetically active (Myagchenko et al. 1989). Naturally, solvation of the associates, in which the metal ion is encapsulated, cannot be effective. The copper salts belong to the type of one-electron redox systems. Redox couples CoIII/CoII, CoII/CoI, MinIII/MnII, CeIV/CeIII, AgII/AgI, IrIV/IrIII, FeIII/FeII, CrIII/CrII, and WVI/WV also belong to the one-electron string. Two-electron redox systems are represented by couples TIIII/TII, PbIV/PbII, PdII/Pd0, II Mg /Mg0, HgII/Hg0, AuIII/AuI, PtIV/PtII, and PtII/Pt0. While the action of one-electron redox systems is readily understandable in the context of inner- and outer-sphere mechanisms, two-electron redox systems require additional
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considerations. First of all, if a double one-electron transfer is probable from an organic substrate to the same metal ion, does this mean that the same molecule of an organic donor provides these two electrons, or do two molecules of the substrate act as one-electron donors? A study of the kinetics of the oxidation of aromatic hydrocarbons with lead tetraacetate has shown the second order of the reaction. Hence, the rate-determined stage consists of the transformation of PbIV into PbII, with the participation of only one molecule of the hydrocarbon (Dessau et al. 1970). The formed hydrocarbon dication can, of course, react with the uncharged hydrocarbon according to Scheme 1-112: .
.
ArH2 ArH → ArH ArH
The cation radicals depicted in the scheme were really detected, but they originate from the fast reaction of one-electron transfer, which does not affect kinetic constants of the oxidation. The rate constant depends linearly on Brown’s -constants of substituents (Dessau et al. 1970). All these data are in agreement with the formation of the strong polar dication of an aromatic hydrocarbon as an intermediate. Because PbII salts (in particular, the diacetate) are not reductants, the two-electron-transfer reaction proceeds irreversibly. In some cases, two-electron transfer to a metal ion leads to the formation of a reducing form. This gives the possibility of constructing catalytic cycles. Oxidation of ethylene with PdII upon catalysis of CuII as in Scheme 1-113 (see below) is a striking example (Denisov 1978). CH2BCH2 Pd II ⇔ (CH2BCH2)PdII (CH2BCH2)PdII H2O → Pd0 2H CH3CHO Pd0 2CuII → PdII 2CuI . . 2 I II II Cu O2 → Cu O 2 , CuI O 2 → Cu O 2 O 22 2H → H2O2; CuI H2O2 → 2OH 2CuII etc. Metal ion participation in electron transfer reactions can proceed with no changes in its oxidation state, but simply at the expense of complex formations. Scheme 1-114 (below) illustrates such assistance by 3d metal ions (M Mn, Co, Ni, Zn) to electron transfer between a reductant (Red) and an oxidant (Ox). M2 Red Ox ⇔ [Red–M2–Ox] ⇔ Red M2(Ox) The reaction depth correlates to the electron donor ability of Red and the stability degree of M2(Ox) complex. The complexation causes an anodic shift of metal redox potentials, which reaches almost 100 mV for transition metal cations (Maletin et al. 1979, 1980, 1983). Alkali, alkaline and rare earth metal cations also catalyze electron-transfer reactions. Thus, in the pair of CoII–tetraphenylporphyrin complex with benzoquinone, no redox reaction takes place, or it goes too slow to be determined accurately. The metal cations promote this reaction. For example, in the presence of Sc(ClO4)3, the corresponding rate constant of 2.7 105 M2sec1 was observed. Benzoquinone transforms into benzosemiquinone under these conditions (Fukuzumi & Ohkubo 2000). Zinc perchlorate accelerates the reaction between aromatic amine and quinones (Strizhakova et al. 1985). The reaction results in the formation of charge-transfer com-
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plexes [ArNH2Q]. The complexes dissociate in polar solvents, giving ion radicals; see Scheme 1-115: K1
K2
.
.
ArNH2 Q ⇔ [ArNH 2 Q ] ⇔ ArNH 2 Q
Equilibrium-constant values are low enough for reactions in Scheme 1-115. For the interaction of N,N-tetramethylphenylenediamine (ArNH2) with diphenyl benzoquinone (Q) in acetonitrile at 200C, K1 4.5 102, K2 1.1 106 (Levin et al. 1979). However, addition of zinc ions increases sharply the rate of formation of the amine cation radical, leaving the rate of formation of the quinone anion radicals with no change. The influence of zinc ions on the electron transfer from an amine to a quinone can be explained by . the binding of anion radical Q in a complex with ZnII. As a result, the equilibrium may be shifted to the right. However, the stability constant of another possible complex, [Zn(ArNH2)]2, is equal to 105 and is much higher than K1 (Strizhakova et al. 1985). In other words, introduction of ZnII should produce the left, not the right, shift of the equilib. rium. The complexation between Q and ZnII indeed takes place, but it does not preserve these anion radicals from decaying in fast collateral reactions. Moreover, the anion radical does not, prior to its decay, have enough time to leave the inner coordination sphere of the complex with zinc. This causes right-shifting of the equilibrium. Meanwhile, the observed result consists of an increase in the rate of accumulation of cation radicals ArNH 2 , but not . that of anion radicals Q . Complexation between zinc ions and anion radicals of aromatic ketones provides a basis for simple one-electron reduction of these ketones (Handoo & Gadru 1986). In the method, dimethylsulfoxide is employed as a solvent, zinc dust serves as a reductant in the presence of strong alkali (6M, several drops). As widely known, zinc in an alkali medium is a very active reductant. For example, it reduces nitrates under such conditions to ammonia. In the ketone case, electron transfer is stopped at the stage of ketyl formation. Protondonor treatment leads to the respective alcohols: benzophenone gives benzhydrol, fluorenone gives fluorenol, etc. Before protonation, the solution contains anion radicals (ketyls), and these anion radicals give the starting neutral molecules after the addition of pdinitrobenzene. All previously mentioned acceptors have electron affinity values up to 210 kJmol1 (Handoo & Gadru 1986). All of them give anion radicals with practically quantitative yields. A breakthrough feature of such method of anion radical generation is that it does not require anaerobic conditions. The resulting anion radicals (probably because of complexation with zinc ions) become stable enough in air. This is in sharp contrast to the high sensitivity of the same anion radicals to air when they are obtained with the help of alkali metals in ether solvents. Hence, the method based on reduction by a Zn/OH system in dimethylsulfoxide gives a tempting alternative to common methods of anion radical preparation. Furthermore, this method allows one to perform the transfer of only one electron and, in a sense, has advantages over preparative electrolysis at controlled potentials. At the same time, the method remains a purely chemical one, which requires no complicated and expensive instrumentation. Samarium diiodide (SmI2, the Kagan reagent) represents another principal example of concerted reduction and complexation. This salt is very easily prepared in tetrahydrofurane from samarium finely ground as a powder and diiodomethane, diiodoethane, or iodine (Krief & Laval 1999). Samarium iodide can be stored as 101 M THF solution under atmosphere of nitrogen or, better, argon (Kagan et al. 1981). It is well known to have exceptional qualities as a single-electron-transfer reductant. For species soluble in organic medium, it submits a negative-enough reduction potential (E1/2 1.62 V in acetonitrile vs. saturated
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i, NOPF6, acetonitrile-ether;
ii, SmI2, tetrahydrofurane
SCHEME 1-116
calomel electrode—Kolthoff & Coetzee 1957). Like other lanthanide derivatives, SmI2 presents a high coordination number (usually 7). Therefore, it provides a unique opportunity to include molecules of reactants, solvents, and electron-transfer products in the framework of one subunit (Kagan et al. 1981; Yacovan et al. 1996). This facilitates electron transfer and . changes some of the usual reactions of an intermediate radical R . In one example, no coupling product R–R is obtained as a result of the reaction between alkyl halide RX and SmI2 . in THF. The intermediate radical R abstracts hydrogen from THF as the only proceeding reaction (Kagan et al. 1981). In the second example, SmI2 allows one to perform the reduction of the triselenonium cation radical. Reduction with SmI2 proceeds smoothly and in strictly one-electron manner (Ogawa et al. 1994). Scheme 1-116 illustrates the reaction. It would be interesting to examine thulium diiodide in one-electron reduction reactions. On the basis of the work of Evans and Allen (2000), TmI2 has the potential to be an effective replacement for SmI2 when the latter is too weak as a reductant, when subambient reaction temperatures are desirable, etc. Perhaps, TmI2 activity in tetrahydrofurane can be controlled by the addition of hexamethyl phosphorustriamide in the same manner as it regulates the power and reactivity of SmBr2 (Knettle & Flowers 2001). 1.8 CONCLUSION The material discussed in this chapter shows that the organic chemistry of ion radicals has specific features. The transformation of organic compounds into ion radicals changes orbital interactions. This causes changes in electron effects of substituents and enhances the ability of bridge groups to participate in electron delocalization. Acid–base properties of organic compounds are also changed fundamentally. This opens new ways to widen the modern methodology of organic synthesis. Based on numerous examples, the chapter points out ways to enrich the reactivity of metallocomplexes. Comprising an extensive body of work studying organic poly(ion radical)s of the monomeric nature, the chapter examines cases of the separate existence of several unpaired electrons and their ferromagnetic coupling. The absence of spin leaking from one ion radical site to another in the framework of the same molecular skeleton is a very important feature in the design of organic materials with magnetic properties. On the other hand, the section devoted to polymeric ion radicals describes the approaches for the design of conductors that the electronics industry relies on. Every organic reaction proceeds with the participation of inorganic reagents. Ion radical organic reactions also have inorganic participants. The chapter discusses inorganic ion radicals in their reaction with organic substrates. The main aim of this chapter is to provide the foundation for all subsequent chapters.
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2 Formation of Organic Ion Radicals
2.1 INTRODUCTION This chapter describes the preparation of organic ion radicals as well as their properties, i.e., electronic structure, reactivity, and interaction with counterions. In the synthetic chemistry of organic compounds, liquid-phase reactions are most typical. Under these conditions, ion radical salts appear to be surrounded with solvate shells. Either the solvent can solvate the individual ions (having cation and anion individually surrounded by solvent molecules) or an ion pair may be solvated without solvent molecules between the ions. Further transformations of ion radicals (their disintegration, interaction with reagents directly or after the disintegration) also take place in solvents. Formation of transient states and stabilization of final products occur in solvents, too. Medium effects on the generation and structure of ion radicals are of great experimental and theoretical interest (see, for example, Orlov et al. 2001). The nature of the solvent determines the efficiency of the chosen method for ion radical generation. That is why this chapter examines the peculiarities of organic compounds as ion radical precursors under the conditions of liquid-phase electron transfer. Ion radicals are widely discussed as intermediate products of electrode reactions. Regularities of the electrode reactions are, undoubtedly, important for organic synthesis. At the same time, there are features that are distinctive of the liquid-phase reactions only. The electrode reactions are heterogeneous. There is a strong electric field (106–107 Vcm1) in the pre-electrode (double-layer) region. Obviously, this may sometimes act on the reactivity of a depolarizer or on an ion radical that originated from it. The solution in the electrode/solution interface region has special (double-layer) properties that are significantly different from those in the bulk solution. The electrode practically does not change its properties during a reaction with a depolarizer. At the given potential, the influence upon species adsorbed with the electrode surface or located in the framework of the double electric layer remains constant.
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In liquid-phase homogeneous reactions, ion radicals form and react without any influence of either electrode. Some discrepancy also exists in ion-radical reactions between a donor and an acceptor in their ground states, on the one hand, and reactions stimulated by photoirradiation or pulse radiolysis, on the other hand. The energy of the incoming electron is much larger in photochemistry or pulse radiolysis than in homogeneous electron transfer or electrochemistry. Of course, the problem of organic ion radical formation has both chemical and physical aspects. This book concentrates on chemistry, but includes some physical aspects when necessary for better understanding. Nevertheless, organic photochemistry and radiolysis of organic compounds are very specific branches, with their own, very special book repertoire. Therefore, this chapter compresses data on the preparation of organic ion radicals as independent particles that can be free or bound with counterions in ion pairs. It considers liquid-phase equilibria in electron-transfer reactions and compares electrode and liquidphase processes for the same organic compounds. Because isotope-containing molecules have specific features as ion radical precursors, the generation of the corresponding ion radicals is considered in a separate section of the chapter. The chapter also pays some attention to the peculiarities of ion radical formation in living organisms. 2.2 CHEMICAL METHODS OF ORGANIC ION RADICAL PREPARATION 2.2.1 Anion Radicals Electrochemical methods to generate anion radicals consist of potential-controlled electrolysis. The control of potential allows one to detain reduction practically just after a oneelectron transfer to a depolarizer. The one-electron nature of the electron transfer is coincidentally studied by means of coulombometry. One molecule must consume one electron. If less than one electron is consumed in the framework of the one-electron reduction, then the yield of an anion radical is not quantitative. The electrolysis often occurs in a special ampoule placed into a resonator of the electron spin resonance (ESR) spectrometer. This permits one to identify many unstable anion radicals. Electrochemical methods of anion radical generation employ an electrode as an electron donor. In the case of an alkali metal (M), one electron of an organic molecule follows the ensuing scheme (here, ArH is an aromatic hydrocarbon): .
M ArH → ArH , M Solvation of a resulting M takes place in a liquid phase. The solvation may play a crucial role in the generation of anion radicals. For example, benzene, which is a poor electron acceptor, forms its anion radical by reacting with a potassium mirror in dimethoxyethane (DME). DME solvates the potassium cation; the potassium–chelate complex is formed with the participation of two oxygen atoms from the solvent. Solvation of the lithium cation proceeds even more effectively because of the smaller ion radius of lithium than that of potassium (0.078 nm for Li, 0.133 nm for K). With lithium in DME, the equilibrium concentration of the benzene anion radical turns out to be two orders of magnitude higher than that with potassium or sodium (Yakovleva et al. 1960). Without going into detail, the following regularity can be stated: The greater the solvating ability of a medium, the easier is electron transfer from the same metal to the same organic acceptor. In alkali metal reductions, the reducing power follows the order K Na Li.
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Diethoxymethane (DEM, CH3CH2OCH2CH2CH3) has recently become available in commercial quantities. Although DEM contains a reactive methylene group, it is stable to organolithium reagents and, supposedly, can be used for preparation of anion radicals. Two important characteristics of DEM that differentiate it from many other ethers are its immiscibility in water and its nonhygroscopic nature. For example, DEM from a container that had been open for over a year performed superbly as a solvent for several water-sensitive reactions. With no special drying, it is competitive with “anhydrous-grade” THF, which is usually stored under argon (Boaz & Venepalli 2001). Another common solvent where oxygen is easily available for coordination with metal cations is tetrahydrofurane (THF). The ability of anion radicals to remove a proton from the 2-position of THF is sometimes a problem. Dimethyl ether is more stable as a solvent; its oxygen atom is also exposed and can coordinate a metal cation with no steric hindrance from the framing alkyl groups. An added advantage of dimethyl ether is that, because of its low boiling point (22°C), it can be readily removed after reductive metallation and replaced by the desired solvent. The use of aromatic anion radicals in dimethyl ether (instead of THF) is well documented (Cohen and others 2001 and references therein). Commonly used alkali metals have lower ionization potentials than reduction potentials, which are needed for the transformation of an organic molecule into an anion radical. Once a molecule starts getting reduced, it will take up as many electrons as the ionization potential of the metal allows. For alkali metals, the scope of application is limited with substances for which the first electron transfer corresponds to enough negative values of cathodic potentials and is highly separated at this scale from potentials of the second or, generally, next electron steps. If not, metal reduction proceeds deeper. No delay at anion radical formation takes place. There are reductive reactions that show a clear dependence on the homogeneity of the metal reducer. For example, during the alkali metal reduction of bis(gem-dihalocyclopropyl)ethanes in the mixture of liquid ammonia with tetrahydrofuran (THF), the route of any further transformation of primary ion radicals depends on being an alkali metal in the homogeneous versus heterogeneous phase (Oku et al. 1983). Lithium is dissolved in this mixture, and sodium forms a liquid suspension in it. This determines the general result of the reduction. The presence of sodium metal in the heterogeneous phase leads to adsorption of the substrate on the metal surface. The adsorption creates a high concentration of the reducing metal in the reaction zone. This assists the deeper reduction with the full loss of a halogen. In a homogeneous medium (lithium as a reducer), the rate of the primary electron transfer is comparable with the rate of posterior protonation at carbon atoms, and the reaction stops at the stage of elimination of two of the four halogen atoms that are present. Colloidal potassium has recently been proved as a more active reducer than the same metal that has been conventionally powdered by means of shaking it in hot octane (Luche et al. 1984; Chou & You 1987; Wang et al. 1994). In order to prepare colloidal potassium, a piece of this metal in dry toluene or xylene under an argon atmosphere is submitted to ultrasonic irradiation at ca. 10°C. A silvery blue color is rapidly developed, and in a few minutes the metal disappears. A common cleaning bath (e.g., Sonoclean, 35 kHz) filled with water and crushed ice can be used. A very fine suspension of potassium is thus obtained that settles very slowly on standing. In THF, the same method did not work. Attempts to disperse lithium in THF or toluene or xylene were unsuccessful, whereas sodium was dispersed in xylene but not in THF or toluene (Luche et al. 1984). Ultrasonic waves interact with the metal via their cavitational effects (see Section 5.2.4). These effects are closely related to the physical constants of the medium, such as vapor pressure, viscosity, and sur-
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face tension (Sehgal et al. 1982). All of these factors have to be taken into account when one chooses a metal to be ultrasonically dispersed in a given solvent. Several methods have been reported for the preparation of highly reactive metal powders by means of the reduction of metal salts in ethereal or hydrocarbon solvents. These methods used alkali metals as reducing agents in conjunction with or without an electron carrier of the naphthalene type. The reduction of metal salts in this manner produces finely divided metal slurries. These metal slurries are highly reactive and air sensitive and are usually pyrophoric in the absence of solvent. Thus, active magnesium, zinc, indium, copper, and other metals were prepared (Yus 1996; Rieke & Hanson 1997). The metals were used for the hydrogenation of alkynes to alkanes (Alonso & Yus 1997) or for the addition of highly reactive zinc to organic bromides (Guijarro et al. 1999). These active metallic forms can, obviously, be utilized in special cases of anion radical preparation, although the cited works contain no mentions of this. By and large, a finely divided precipitate of a metal is a very effective one-electron reducer. For example, a finely divided precipitate of Zr(0) was obtained on mixing of the naphthalene sodium derivative in THF with ZrCl4. The Zr(0) precipitate was dissolved on the addition of anthracene or benzophenone to form the corresponding zirconium salts of the anion radicals (Terekhova et al. 1996). Neutral organic molecules can also be one-electron donors. For example, tetracyanoquinodimethane gives rise to the anion radical on reduction with 10-vinylphenothiazine or N,N,N ,N -tetramethyl-p-phenylene diamine. Sometimes, alkoxide or phenoxide anions find application as one-electron donors. There is certain dependence between carbanion basicity and their ability to be one-electron donors (Bordwell & Clemens 1981). Sometimes, anion radicals are formed indirectly, by means of special chemical reactions. Photoionization of hydrazine in a mixture of liquid ammonia with THF in the presence of potassium t-butoxide leads to the formation of the diazene anion radical by the sequence of the following reactions (Brand et al. 1985): H2NMNH2 t-BuO → t-BuOH H2NMNH .
H2NMNH h → esolv. H2NMNH . . H2NMNH t-BuO → t-BuOH HNMNH . . HNMNH ↔ H2N BN One unusual case of anion radical preparation is 4- and 3-nitrocatechol or nitrohydroquinone. Reduction of these nitrocompounds in aqueous solutions with sodium borohydride in air leads to the corresponding anion radicals, which are stable at pH 9–12 despite the presence of water (Grenier et al. 1995). As already stated, the interaction between the metal cation and the anion radical that is generated is usually a simple ion association moderated by the coulombic attraction between the two species and the competitive ion-solvating nature of the given solvent. However, interaction between organic one-electron acceptors and some metals results in the formation of a special class of organometallic complexes. The metals reduce the acceptor molecules. The anion radical and the cation form a coordination anion radical complex. Thus, when Cd (d10s2) metal reduces benzoquinone in THF, a solvated organometallic compound is formed (Ch. D. Stevenson, Reiter, Burton, & Halvorsen 1995). According to ESR studies, the benzosemiquinone anion radical is ion associated with Cd2 ion or coordinated to it (Scheme 2-1). Exposure of naphthalene dissolved in liquid ammonia to europium metal immediately results in the characteristic green color of the naphthalene anion radical. ESR analy-
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SCHEME 2-1
sis reveals a signal that comes from an unpaired electron interacting with the 151Eu and 153 Eu nuclei. No hyperfine coupling with the naphthalene protons is observed, although treatment with water leads to 1,4-dihydronaphthalene (Ch.D. Stevenson, Schertz, & Reiter 1999). This means that naphthalene has indeed been reduced to its anion radical and undergone a normal Birch reaction. These results are consistent with the initial donation of the two s electrons to two acceptor molecules. Tight coordination then takes place between the two naphthalene anion radicals and the central atom of Eu2, presumably through the europium s orbital. Since the two electrons, donated by the Eu, are in the same molecular orbital, they are spin paired. Hence, only the remaining seven unpaired f electrons can contribute to the ESR signal. Therefore, the resulting ESR spectrum reflects the interaction of the europium nucleus with a single electron spin. The structure of the complex is depicted in Scheme 2-2. The structure “is tentative, but it is known that: (1) europium ions prefer an octahedral geometry, (2) they complex well with ammonia, and (3) (in this case) two naphthalene units are involved” (Ch.D. Stevenson et al. 1999). . Another rare earth metal cation, Sc3, catalyzes electron transfer from C 60 to phenacyl bromide; the effect is supposedly close to the phenomenon just discussed (Fukuzumi et al. 2000). Benzene and toluene anion radical salts are intriguing examples. Benzene and toluene anion radicals prepared as Cs salts were shown to dimerize readily in THF at 70°C. 3,3 -Bis(cyclohexa-1,4-diene) and the corresponding toluene 1,1 -dimethyl analog
SCHEME 2-2
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were isolated on hydrolysis (Grovenstein et al. 1977). In the presence of 18-crown-6-ether, however, toluene gives, with potassium in THF, the anion radical salt as a tight ion pair. The crown ether encapsulates the potassium cation, which is bound in an approximate 6 fashion to the planar six-membered ring in CH3C6H5. Treatment of potassium with 18crown-6 in a benzene–THF mixture initially yields the very similar benzene-(potassium18-crown-6) anion radical salt in 6 fashion. This salt is barely soluble in both benzene and THF. The solid salt transforms into a dimer. The 3,3 -bis(cyclohexa-1,4-dienyl) ligand is just the dianion of the crystalline salt, being 5:5 bound in a transoid fashion to each of the two crown-encapsulated potassium cations (Hitchcock et al. 2001). It unclear which factor plays the main role in these different reaction pathways: the low solubility of the crystalline cyclohexadienyl dimer, or the electronic and steric effects of the methyl group making dimerization of the toluene-(potassium-18-crown-6) anion radical salt unfavorable. Sometimes, anion radical salts formed in a solution crystallize as aggregates. For instance, if anthracene is reduced at an alkali–metal surface, a slight modification in the nature of the solvent leads to various reduction products that were structurally characterized in single crystals. In the system of Na triglyme, the bare anthracene anion radical crystallizes out as a crystal solvate containing the sodium cation surrounded by two triglyme molecules. In the system of K tetrahydrofurane, the ion quadruple of two anthracene anion radicals connected by a [(K)2(THF)3] bridge crystallizes (Bock, Gharagozloo-Hubmann et al. 2000b). Additionally, one specific topic should be underscored: the use of silicon grease for the fitting of glass in Schlenk vessels, which are usually employed for the preparation of anion radicals. This is merely an experimental detail, at first glance. However, it has significant and general importance. Biphenylene reduction by sodium metal mirror in THF solution containing [2.2.1]cryptand yields as a structurally characterized product the 9,9dimethylsilafluorene anion radical salt containing a Me2-Si-expanded cyclobutadiene ring. When the use of any silicon grease is avoided for all glass fittings, the solvent-separated ion pair crystallizes (Bock, Sievert, et al. 1999). Whether or not this reaction is specific for the initially generated biphenylene anion radical only, the transformation of silicone (Me2SiO)n into the Me2Si-containing reactive species is indisputable. Such species really originate from trace silicon concentration and the residual oxygen, which is disposed on the sodium mirror surface. If the use of silicon grease is avoided, this difficulty in anion radical preparation is circumvented. The unwanted reaction described took place in THF. The specificity of this solvent remains an open question. 2.2.2 Cation Radicals In order to understand features of oxidative one-electron transfer, it is reasonable to compare average energies of formation between cation-radicals and anion-radicals. One-electron addition to a molecule is usually accompanied by energy decrease. The amount of energy reduced corresponds to molecule’s electron affinity. For instance, one-electron reduction of aromatic hydrocarbons can result in the energy revenue from 10 to 100 kJmol1 (Baizer & Lund 1983). If a molecule detaches one electron, energy absorption mostly takes place. The needed amount of energy consumed is determined by molecule’s ionization potential. In particular, ionization potentials of aromatic hydrocarbons vary from 700 to 1,000 kJmol1 (Baizer & Lund 1983). Because many electron-deficient cation-radicals are less stable and have much more reactivity, cation-radical studies were harder to do with slower accumulation of the relevant literature. Thus, the Landolt-Boernstein reference book devotes 149 pages to hydrocarbon anion radicals (1980, vol. 9d1), whereas only 15 pages are reserved for hydrocarbon cation
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radicals (1980, vol. 9d2). The ratio is 10:1. The recent development of physical methods drastically changes this ratio. According to Chemical Abstracts Subject Indices, the ratio was 1:3 during 1984–1993 and reached 1:5 in 1994–1998. Studies of cation radical chemistry began with the work of Weiss (1941) and Michaelis and his coworkers (1941). However, intense investigation in this field began after instrumental methods emerged (ESR and optical spectroscopy in particular). Aromatic hydrocarbons gave products of protonation on dissolving in hydrofluoric acid. Oxidation in aromatic cation radicals did not take place (Kon & Blois 1958). Trifluoroacetic acid is an effective one-electron oxidant (Eberson & Radner 1991). Meanwhile, sulfuric acid caused not only dissolution and protonation, but also one-electron oxidation of aromatic hydrocarbons. Sulfonation, naturally, proceeded too (Weissmann et al. 1957). Anodic oxidation in inert solvents is the most widespread method for cation radical preparation, with the aim of investigating their stability and electron structure. However, saturated hydrocarbons cannot be oxidized in an accessible potential region. There is one exception for molecules with the weakened CMH bond, but this does not pertain to the cation radical problem. Anodic oxidation of unsaturated hydrocarbons proceeds more easily. As usual, this oxidation is assumed to be a process including one-electron detachment from the -system with cation radical formation. This is the very first step of the oxidation. Certainly, the cation radical formed is not inevitably stable. Under anodic reaction conditions, it can expel the second electron and give rise to a dication or lose a proton and form a neutral (free) radical. The latter can either be stable or complete its life at the expense of dimerization, fragmentation, etc. Nevertheless, electrochemical oxidation of aromatic hydrocarbons leads to cation radicals, the nature of which is reliably established (Mann & Barnes 1970, Chap. 3). In the majority of cases, the electrochemical generation of organic cation radicals takes place in an ampoule lowered into an ESR cavity. Sometimes, however, exhaustive external generation and use of a flow system allow one to obtain an ESR spectrum that is far better resolved (see, for example, Seo et al. 1966). Electrochemical methods are very useful in structural studies but are barely applicable for preparative aims. The cause is the limited stability of cation radicals. It is difficult to do low-temperature preparative electrolysis, and the main problem is to dispose of the large amount of heat generated during the electrode work. That is, not much current can be passed through an ordinary-sized electrode without generating too much heat. When potential and temperature control are necessary, only small quantities of a material can be obtained in a reasonable period of time. When potential and temperature control are not necessary, as in Kolbe electrolysis, anodic oxidation is indeed useful as a preparative method. As to chemical routes to cation radical generation, the following oxidants deserve to be mentioned: concentrated (98%) sulfuric acid (Carrington et al. 1959; Hyde & Brown 1962; Carter 1971), persufate (Minisci et al. 1983), iodosobenzene bis(trifluoroacetate) (Alberti et al. 1999), chlorine dioxide (Handoo et al. 1985; Sokolov et al. 1999). The following metal ions found application: Tl(III) (Elson & Kochi 1973; McKillop & Taylor 1973), Mn(III) (Andrulis et al. 1966), Co(III) (Kochi et al. 1973), Ce(IV) (Norman et al. 1973), Ag(I) and Ag(II) (Nyberg & Wistrand 1978), Pd(II) (Eberson & Wistrand 1980). Methods of cation radical generation have also been developed in which oxidants are SbCl5 (Ishizu and co-workers 1973), AlCl3 (Forbes & Sullivan 1966), I2 (Stamirez & Turkevich 1963), Pb(OAc)4 (Elbl-Weiser et al. 1989), and XeF2 (Shaw et al. 1970). Potassium 12tungstocobalt(III)ate was also described as a strictly one-electron oxidant for organic substances (Eberson 1983; Baciocchi et al. 1993).
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This list of oxidants should be accompanied by a consideration of the regularities of their actions; see Chapters 1, 4, and 6. Here, the following three oxidants ought to be em. phasized: Ar3N SbCl 6 , NO , and Ag . Tris(bromophenyl)amine is now available commercially; oxidized with SbCl5 it forms the reagent (see Section 1.7.10). The NOPF 6 and NOBF 4 salts are also available; they are especially convenient when the aim is the isola. tion of oxidized products. This is because the reduction product, NO , is a gas that is evolved (see Section 1.7.9). The Ag ion is used whenever the compounds are easy enough to oxidize: the reduced form is separable solid silver (often as a mirror on the reaction vessel walls). Just like NO, Ag is characterized by a high sensitivity of its standard redox potential to the nature of the solvent. (For selecting chemical oxidants and reductants, the 1996 review by Connelly and Geiger can be recommended.) In some cases, cation radicals are formed from neutral organic molecules upon the action of neutral organic acceptors, such as tetracyanoethylene, tetranitrofluorenone, quinones, and free radicals—aroxyls, nitroxyls, and hydrazyls. As shown, methods for cation radical generation are very various. Each case needs its own appropriate method. The set of such methods is continuously being supplemented. New methods not claiming to be general ones permit one to prepare nondescribed cation radicals, to strive for more resolved spectra, and so on. For one example, it was very difficult to prepare the cation radicals of benzene derivatives with the strong acceptor groups. Recently, some progress has been achieved, thanks to the employment of fluorosulfonic acid, sometimes with the addition of antimony pentafluoride, and lead dioxide (Rudenko 1994). As known, superacids stabilize cationic intermediates (including cation radicals) and activate inorganic oxidants. The method mentioned is effective at 78°C. However, 78°C is the boundary low temperature, because solution viscosity increases abruptly. This leads to anisotropy of a sample and sharply deteriorates an ESR spectrum. Another example involves the preparation of cation radicals from aromatic hydrocarbons. In concentrated sulfuric acid or in a mixture of antimony pentachloride with methylene chloride, polyacenes give cation radicals with ill-resolved ESR spectra (Lewis & Singer 1965). Employment of molten antimony trichloride at 80°C results in improved resolution, diminution of line widths up to the ideal value of 0.005–0.008 mT (Buchanan et al. 1980). Under these conditions, even naphthalene gives the cation radical, with an ESR spectrum of very good quality. The mentioned spectrum surprisingly describes the naphthalene cation radical in its monomeric (not dimeric) form. An essential peculiarity of the method is its special opportunity to regulate the oxidation capacity of SbCl3: It is sufficient to add several mole percents of an acceptor of chloride ions, e.g., AlCl3, or a donor of these ions, e.g., Me4NCl. Because the oxidation process is reversible, a measure of its reversibility is the yield of antimony in the zero-valent state: .
SbCl3 3ArH ⇔ 3(ArH) 3Cl Sb0 Bringing Cl out of the reaction sphere (by the addition of AlCl3) leads to an increase . in (ArH) concentration. Getting Cl into the reaction sphere (by the addition of Me4NCl) . leads to a decrease in (ArH) concentration. As seen, the equilibrium between starting substances and electron-transfer products can be very important, especially for homogeneous reactions. The depicted equilibrium may be more complex in reality and include polyantimonium polyhalides, but the most important peculiarity is the possibility of regulating the oxidative capacity, as it is shown here.
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SCHEME 2-3
A relevant example is a method of indirect generation of cation radicals. Mesoionic 5,5 -azinobis(1,3-diphenyltetrazole) and related mesoionic compounds give cation radicals on oxidation with lead tetraacetate. The reaction proceeds in the presence of sodium tetrafluoroborate. The tetrafluoroborate cation radical salt is stable and can even be purified via column chromatography on silica gel and stored under air for several months without appreciable decomposition (Araki et al. 1999). The cation radicals of N,N -dimethyldiazines are usually produced either by reduction of the corresponding diquarternary salts (Schulz et al. 1988) or by oxidation of the parent diazine (Soos et al. 1977). Lucarini and co-authors (1994) have elaborated a one-pot method for generating such cation radicals from the corresponding aromatic bases via the action of dimethyl sulfate, Me2SO4, as a methylating agent and zinc powder as reductant. The reaction proceeds in benzene according to the Scheme 2-3. To conclude this section, one indirect method of cation radical generation should be mentioned. Organic molecules spontaneously form corresponding cation radicals upon inclusion within activated zeolites (Yoon & Kochi 1988; Yoon 1993; Pitchumani et al. 1997). Zeolites are crystalline alumosilicate minerals that are widely used as sorbents, ion exchangers, catalysts, and catalyst supports. The cation radicals are formed along with carbocations, but the carbocation generation can be diminished by selection of the appropriate zeolite and the manner of activation. For instance, Ca(Y) zeolite activated at above 400°C on a high-vacuum line was found to be ideal for generation of the cation radicals from diarylethenes in cyclohexane solution, even in the absence of light, with minimum interference from carbocations. Oxygen (not necessarily the free oxygen) plays a crucial role in the generation of the cation radicals within the zeolites (Pitchumani et al. 1997). 2.3 EQUILIBRIA IN LIQUID-PHASE ELECTRON-TRANSFER REACTIONS Let us consider a redox reaction of the following type: .
.
A B ⇔ A B
In this reaction, A. acts as a one-electron donor and transforms into A. At the same moment, B gives rise to B.. If the reaction is reversible, the equilibrium is the sum of the following two processes: A ⇔ A e
and
.
B e ⇔ B
Electrochemically, both processes are expressed by their standard potentials E0(A) and E (B). In order to consider some polarographic studies, let us mention that the polaro0
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graphic half-wave potentials, E1/2(A) and E1/2(B), can, under certain conditions, be used instead of E0(A) and E0(B). If both processes are one-electron in nature and they are fast and reversible, their E1/2 values are equal to those of standard redox potentials (referring to the conditions used in a given electrochemical experiment). Values of E1/2 show how easily A or B accepts an electron, i.e., what the A or B acceptor/donor ability is. There is a special equation to count equilibrium constants (K) using Faraday’s constant (F) and E1/2 values of one-electron reversible polarographic waves for A and B (Gennaro et al. 1988): F[E1/2(A) E1/2(B)] ln K RT The equation can give the E1/2(A) value if E1/2(B) and K are known. This estimation has been made for a reaction between the anion radical of tetracyanoquinodimethide (TCNQ) and 2,3-dichloro-5,6-dicyanobenzoquinone (DDQ) (Iida & Akamatu 1967): .
.
(TCNQ) DDQ ⇔ TCNQ (DDQ)
The equilibrium constant of the reaction was determined spectrometrically in acetonitrile. Having K and E1/2(DCNQ) as the known values, E1/2(TCNQ) was determined to be 38 mV. The experimental (polarographic) value was 40 mV. As seen, the calculated and experimental values turned out to be close. Hence, equilibrium constants of homogeneous electron-transfer reactions between . (A) and B are evidently connected to the differences in reduction potentials of A and B. This connection reflects a definite physical phenomenon. Namely, if two redox systems are in the same solution, they react with each other until a unitary electric potential is reached. For the transfer of only one electron at room temperature, the following simplified equation can be employed: ln K 2.3[E1/2(A) E1/2(B)] 0.059 The applicability of this simple equation has been checked for one particularly important case of the electron exchange between species belonging to quite different classes of chemical compounds. Cyclo-octatetraene dipotassium as a donor and -ferrocenylacrylonitrile as an acceptor react in THF. The reaction is reversible, and the presence of all four of the components depicted in Scheme 2-4 has been proved (Todres 1987).
SCHEME 2-4
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According to the yield of C8H8 and the concentrations of the four components, the equilibrium constant of the reaction was determined as 4 102. The value calculated from the potential difference was 4.4 102. Consequently, there is coincidence between the calculated equilibrium constant based on the electrode potentials and the equilibrium constant determined from the liquid-phase experiment. The liquid-phase and electrode processes are similar. In all of the examples considered, E1/2 of the acceptor was much more negative than that of the donor. However, in liquid phase one-electron transfer from a donor to an acceptor can proceed even with an unfavorable difference in the potentials if the system contains a third component, the so-called mediator. The mediator is a substance capable of accepting an electron from a donor and sending it instantly to an acceptor. Julliard and Chanon (1983), Chanon, Rajzmann, and Chanon (1990), and Saveant (1980, 1993) developed redox catalysis largely for use in electrochemistry. As an example, the reaction of terachloromethane with N,N,N ,N -tetramethyl-p-phenylenediamine (TMPDA) can be discussed. The presence of p-benzoquinone (Q) in the system provokes electron transfer (Sosonkin et al. 1983). Because benzoquinone itself and tetrametyl-p-phenylenediamine interact faintly, the effect is evidently a result of redox catalysis. The following schemes reflect this kind of catalysis: .
.
CCl4 TMPDA → (TNPDA) Cl CCl3 . . TMPDA Q → (TMPDA) Q . . Q Q ⇔ Q Q2 . . Q2 CCl4 → CCl3 Cl Q . . Q CCl3 → Q CCl3 CCl3 H → CHCl3 The observed catalytic effect can be explained by comparing the redox potentials of the . reacting species. The potentials of the reversible electron transfers from Q to Q and then from . . Q to Q are right in the middle of the gap between the potentials of CCl4 and TMPDA. This leads to dividing the gap into three components so that each one can be gotten over readily. One curious case of the effect of light on electron-transfer equilibrium involves the reduction of , -di(t-butyl)stilbene with potassium in DME. The reaction leads directly to a diamagnetic dianion; a solution of this dianion remains ESR silent unless subjected to ultraviolet irradiation by a Hg/Xe lamp. The anion radical of , -di(t-butyl)stilbene then formed from the dianion by the loss of an electron. The electron reverted within 5–10 min after ultraviolet irradiation was turned off, transforming the anion radical into the dianion (Gerson et al. 1996). This case deserves to be clarified. Maybe the light effect consists simply in singlet–triplet transformation of the dianion, with the formation of some more or less stable biradical state of the dianion, which possesses two unpaired electrons and can even be a paramagnetic one. 2.4 ELECTROCHEMICAL METHODS VERSUS CHEMICAL METHODS To date, the results of chemical and electrochemical studies generally have been considered in isolation. The purpose of this section is to present a digest of data obtained via the two methods in an effort to demonstrate that a great deal may be gained by both chemical and electrochemical means rather than solely by one method or the other, as is the currently accepted practice.
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In the electrochemical technique, the electrode provides the source (reduction) or sink (oxidation) for electrons. Variation of the applied potential provides the driving force that enables the redox reaction to occur. Organic depolarizers diffuse toward the electrode surface until they are sufficiently close to the electrode (the double-layer region) to accept electrons from the electrode and then to diffuse away from the electrode surface and to become involved in the bulk solution. In chemical redox reactions, organic ion radicals are formed at the final stage of donor–acceptor (D-A) interactions: .
.
.
.
D A ⇔ D , A ⇔ D , A ⇔ D A
All bimolecular reactions include encounter complexes, which involve as much electron transfer as needed to stabilize the system, the amount varying widely depending on what the two molecules are. When there is a lot of charge transfer, it is certainly reasonable to call the complex a charge-transfer complex. The equilibrium constants for complex formation is supposed to vary smoothly from small to large, depending on what the molecules reacting are. It is worthwhile to consider the products of donor–acceptor interaction, namely, charge-transfer complexes and ion radicals. Regarding electrochemical reactions, ion radicals are well-known primary products. Charge-transfer complexes are more usual for chemical processes, but they have their own analogy in electrochemistry: The formation of a charge-transfer complex is in a certain sense similar to the formation of an electrode–substrate complex. As an example, aromatic, especially polycyclic, hydrocarbons are adsorbed on platinum and other metal anodes with adsorption enthalpies in the range from 25 to 42 kJmol1. As calculated, this adsorption must occur with the -electron contour parallel to the surface. The type of bonding involved would be of the same kind as that in a -complex. The distance between the hydrocarbon molecule and the surface would probably be of the same magnitude as that in -donor– -acceptor complexes, or around 0.35 nm (Baizer & Lund 1983). The adsorption of supporting electrolyte or other additives can create a special microenvironment in the reaction layer near the electrode that can favor distinct pathways. Thus the use of tetra-alkylammonium ions as supporting electrolyte depletes the electrode surface of a solvent and sometimes encourages the products of electron transfer to react, with no solvent participation. On the other hand, the intermediates generated at the surface diffuse into the bulk solution. Owing to their high reactivity, the intermediates react in a fairly thin reaction layer that adheres to the electrode, and thus their concentration is higher than in homogeneous reactions, where they are spread uniformly over the medium. This can influence the nature of the product and its distribution. The fundamental electrochemical event, i.e., electron transfer, occurs at the electrode surface. Peculiarities of electrochemical reactions include an electrical field that in a special way complicates the phenomena of adsorption and desorption at the surface. The first layer of the solution, which is in contact with the electrode, possesses a specific structure. It is important for charged particles that the orientation of medium molecules in the vicinity of the electrode produces a decrease in dielectric permeability in the compact part of the double layer (Damaskin & Kryshtalik 1984). Substrate and/or intermediate species adsorb on an electrode surface and orient themselves so that their least hindered sides face the electrode, unless there is another effect, such as polar one. An electrode interface has a layered structure in which a nonuniform electric field (some slope of potential) is generated by polarization of the electrode. An extremely strong electric field of around 108 V/cm in the innermost layer might cause a vari-
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ety of polar effects. Since not only the electron transfer step but also adsorption and some of the chemical steps involved in an electrode reaction take place in the layer, the whole process should be strongly influenced by polar factors. The orientation of polar adsorbed species, such as ion radicals in particular, is electrostatically influenced, and consequently the stereochemistry of their reactions is also controlled by this kind of electrostatic factor. All these phenomena were summarized in several monographs. The collective volume edited by Baizer and Lund (1983) is devoted to organic electrochemistry. This edition is closer to the scope of our consideration than its latest version of 2001 (edited by H. Lund and Hammerich; these editors have changed the authors of the chapters included). As for chemical reactions, the oxidation-reduction (redox) reactions in homogenous medium (i.e., in the bulk of the solution) have been experimentally studied with proper intensity only in the last two decades. There has been some development of the bulk reactions. However, as before, a comparison of one and the same compound in chemical and electrochemical electron-charge-transfer reactions is still of current interest. Such a comparison is made in this section. The examples offered are intended to invoke novel interpretations or to discover new colors in pictures that have already been drawn. 2.4.1 Charge-Transfer Phenomena It is worth noting that there is a significant difference between a conversion of one and the same substrate under one-electron (electrode) transfer and in charge-transfer complexes (homogeneous medium). Anodic oxidation of tetraphenylethylene at a platinum electrode leads to the product of cyclization, namely, to 9,10-diphenylphenanthrene (Stuart & Ohnesorge 1971). The intramolecular coupling reaction does not occur when diphenylethylenes, i.e., stilbene and its methyl derivatives, are electrolyzed under the same conditions (Stuart & Ohnesorge 1971). This difference in the anodic behavior of these substances was attributed to the low stability of the cation radicals of stilbene and its methyl derivatives in comparison to the cation radicals of tetraphenylethylene. The participation of the cation radicals in cyclization of tetraphenylethylene has been unequivocally proved (Svanholm et al. 1974; Steckhan 1977). In homogeneous conditions, when p-chloranil plays the role of electron acceptor, 4-methylstilbene ( -phenyl- -tolylethylene) can, however, be cyclized and converted into 3-methylphenanthrene. The reaction takes place via formation of a charge-transfer complex at a very moderate temperature (36°C) and does not require light radiation (Todres, Dyusengaliev, Buzlanova et al. 1990). Minus details, Scheme 2-5 depicts the transformation. The behavior of 4,4 -dimethoxystilbene on the electrode and with p-chloranil in solutions was compared (Todres & Ionina 1992). It was found that the cis and trans isomers of the compound form different charge-transfer complexes with chloranil that are unlike in color and can be converted;
