摘要:

Chemical Society Reviews Vol 4 No 3 1975 Page FARADAY LECTURE The Electron as a Chemical Entity By F. S. Dainton 323 Bile Pigments By M. F. Hudson and K. M.Smith 363 TATE AND LYLE LECTURE Spin-Lattice Relaxation: A Fourth Dimension for Proton N.M.R. Spectroscopy By L. D. Hall 401 Transition Metal Complexes of Synthetic Macrocyclic Ligands By L. F. Lindoy 421 Co-ordination Chemistry of Aryldiazonium Cations : Aryldiazenato (Arylazo) Complexes of Transition Metals, and the Aryldiazenato-Nitrosyl Analogy By D. Sutton 443 The Application of Electrochemical Techniques to the Study of Homogeneous Chemical Reactions By D. Pletcher 471 The Chemical Society London Chemical Society Reviews Chemical Society Reviews appears quarterly and comprises approximately 25 articles (ca.600 pp) per annum. It is intended that each review article shall be of interest to chemists in general, and not merely to those with a specialist interest in the subject under review. The articles range over the whole of chemistry and its interfaces with other disciplines. Although the majority of articles are intended to be specially commissioned, the Society is always prepared to consider offers of articles for publication. In such cases a short synopsis, rather than the completed article, should be sub- mitted to The Editor, Reports and Reviews Section, The Chemical Society, Burlington House, Piccadilly, London, W1V OBN. Members of the Chemical Society may subscribe to Chemical Society Reviews at 23.00 per annum; they should place their orders on their Annual Subscription renewal forms in the usual way. Non-members may order Chemical Society Reviews for El 0.00 per annum (remittance with order) from: The Publications Sales Officer, The Chemical Society, Blackhorse Road, Letchworth, Herts., SG6 lHN, England. 0Copyright reserved by The Chemical Society 1975 Published by The Chemical Society, Burlington House, London, W1V OBN Printed in England by Eyre 8c SpottiswoodeLtd, Thanet Press, Margate

ISSN:0306-0012    

DOI:10.1039/CS97504FP003

出版商:RSC    

年代:1975

数据来源: RSC

ISSN:0306-0012    

DOI:10.1039/CS97504FX005

出版商:RSC    

年代:1975

数据来源: RSC

摘要:

Chemical Society Reviews Vol 4 No 3 1975 Page FARADAY LECTURE The Electron as a Chemical Entity By F. S. Dainton 323 Bile Pigments By M. F. Hudson and K. M. Smith 363 TATE AND LYLE LECTURE Spin-Lattice Relaxation :A Fourth Dimension for Proton N.M.R. Spectroscopy By L. D. Hall 401 Transition Metal Complexes of Synthetic Macrocyclic Ligands By L. F. Lindoy 421 Co-ordination Chemistry of Aryldiazonium Cations : Aryldiazenato (Arylazo) Complexes of Transition Metals, and the Aryldiazenato-Nitrosyl Analogy By D. Sutton 443 The Application of Electrochemical Techniques to the Study of Homogeneous Chemical Reactions By D. Pletcher 471 The Chemical Society London

ISSN:0306-0012    

DOI:10.1039/CS97504BX007

出版商:RSC    

年代:1975

数据来源: RSC

摘要:

FARADAY LECTURE* The Electron as a Chemical Entity By F. S. Dainton UNIVERSITY GRANTS COMMITTEE, 14 PARK CRESCENTy LONDON WlN4DH 1 Introduction When I heard that Council wished me to give the Faraday Lecture my reaction was compounded of delight and apprehension. The former was a stimulus to accept the invitation whilst the latter gave cause for refusal. My pleasure sprang from the very great honour that the Society had done me and from the fact that Faraday has been a life-long personal hero. I well remember preparing myself to go with a school party to view the solar eclipse of 29th June 1927 from the top of Ingleborough by studying the Ordnance Survey map of that district and seeing on it the words ‘Faraday Ghyll’. I wondered if there were any connection between this topographical feature and the Laws of Electrolysis which an over-ambitious and enthusiastic master had brought to my tender and uncomprehending attention.Indeed there is; for although Michael was born in London in 1791, his father was born in Clapham Wood Hall about 1 mile west-south-west of Clapham Village. When I last visited it about a dozen years ago it was still standing and had a fine outlook over Burn Moor Fell. Somewhat fancifully -some would say extravagantly -I like to think that Michael’s passion for long country walks (20 miles were a mere stroll to him) and some of his personal qualities-his respect for truth, his emphasis on self-education, his steadfastness -were derived from his West Riding ancestors.At all events I became interested in the man. I read his ‘Chemical History of a Candle’ and marvelled at the clarity of his exposition. Later I came to know and admire his scientific work and his ardent campaigning to ensure that natural philosophy should be an element in the education of all who lay claim to the title educated, a cause which, regrettably, still needs its champions today. My countervailing feeling of apprehension stemmed from the hopelessness of trying to commemorate, by emulating, such a brilliant lecturer. I shall try to follow his precepts for lecturers, but I shall not reveal those to you until the end when you can judge as to whether his advice has been properly heeded by me. Having accepted the invitation I had no hesitation about an appropriate topic.It just had to be concerned with electricity -Faraday’s consuming passion, the etymology of which is based on the Greek for amber, ~AEKT~V.In his biography,l Pearce Williams draws attention to von Helmholtz’ suggestion in *Delivered to the Chemical Society Annual Congress, London, 1974. l L. Pearce Williams, ‘Michael Faraday’, Chapman and Hall, London, 1965. The Electron as a Chemical Entity his 1881 Faraday Memorial Lecture that Faraday’s second law of electrolysis combined with Dalton’s Atomic Theory implied that ‘at the atomic level each element received the same quantum of electricity. This quantum is what we would today call the electron’. Faraday missed this implication. Since his last work was concerned with electrical discharges in rarefied gases and since he also observed that these discharges were deflected by magnetic fields it is tempting to think that, with a little luck, Faraday might have identified the electron, in the gas phase at least, almost half a century before elm was measured.The idea that an electron in a condensed phase might have an importance in chemistry was not taken seriously for another half century. Of course the notion of redox potentials led to the writing of equations like Feaq2+-Feaq3++ e- and Iaq--+I2 + e-(1) and it was suggested in the twenties by Franck and Scheibe2 that such reactions could be stimulated by light, but the electron was thought to be unstable in water. I have a vivid memory of giving a lecture on the photo-ionization of aqueous bivalent ions of the first row of the transition elements at the Argonne National Laboratory in October 1952 and, after writing the first part of equation (l), being sternly and publicly rebuked by the late Professors James Franck and Robert Platzman who were confident that the aquated electron, eaq-, would decompose according to equation (2) in a period comparable with the dielectric relaxation time of water, i.e.10-11 s. eaq--+H-+ OHaq-Such Olympian contumely was not to be disregarded and it was not until 1959 that I recovered sufficient self-confidence to carry out projected experiments to see if eaq- could exist long enough to take part in chemical reactions. I was encouraged to resurrect this experimental programme because in 1958 Weisss and his collaborators and in 1959 Barr and Allen4 (see Figure 1) had shown that the reducing entity in the radiolysis of water, then thought to be a hydrogen atom formed in a process written as H20WH-+ *OH (3) reacted in a different manner and at different rates with the same solute at pH < 2 and at pH 2 3.This finding opened the possibility that at higher pH the reducing entity was eaq-, convertible into Ha at low pH via reaction (4). (The alternative explanation that He + Haq+ -Hzaq+was the relevant reaction was thought less likely because H2aq+ was expected to be an oxidizing agent.) Since the solute eaq-+ Haq++H* (4) concentrations used in these experiments were ca. 10-3 mol I-1 these results also a J.Franck and H. Scheibe, Z. phys. Chem. (A), 1928,139’22. J. Weiss and E. Hayon,J. Chem.Soc., 1960,5091;E. Hayon and A. 0.Allen,J. Phys. Chem., * N. 1961,65,2181. F. Barr and A. 0.Allen,J. Phys. Chem., 1959,63,928. Dainton 5 a a (C 1-I-HZ) 4 'a m 3 0 CI sd 2 1 0 0 1 2 3 4 5 6 7 8 PH Figure 1 Yields of hydrogen and chloride ion from the y-radiolysis of 0.1M deaerated aqueous solutions of chloroacetic acid in the pH range 0-6 (Reproduced by permission from J. Phys. Chern., 1961,65,2181) implied that, unless the rates were faster than diffusion controlled reactions (k 1010 1 mol-1 s-I), then the lifetime of eaq- was at least 104 times greaterN than the dielectric relaxation time. The next step was obvious.It was to find whether the reducing entity at pH 2 3 had zero or unit negative charge. This merely required a measurement of the effect of changing the ionic strength, p, on the rate of the reaction of this entity with an ion Am. However, there were no known fast reaction techniques 325 The Electron as a Chemical Entity available and we had therefore to have recourse to a competition method in which AZa and Bo competed as shown in equations (5) and (6) and measuring the relative rates of these reactions at 25 "C.The Brsnsted-Bjerrum Laws then eaq-+ Ah -+ Ah-' (5) eaq-+ BO -+ B- (6) indicate drAa-11 ks[Aza] d[B-]/dt -k6[Bo] ___I-- where loglo [k5/k6] = constant + za (-l)f(p), andf(p) is a function [under appropriate conditions approximating to Jy/(l + Jp)] of the ionic strength of the medium.Quite independently of this work at Leeds,5 Czapski and Schwarze were simultaneously doing similar experiments at the Brookhaven National Laboratories but using different A's and with B = HzOz. The results of both groups (see Figure 2) were concordant and conclusively showed that the reducing entity had unit negative charge and was probably the hydrated electron. 0.2 0.1 -1.9 0-1 0.2 0.1 0.2 0.3 Figure 2 TIte efect of ionic strength, p, on the relative rate constants of reactions (5) and (6) in water at 25 "C.Plotted from data in refs. 5 and 6 AZa NOa-Oa H+ Ag+ KFe(CN), a -B" Ha02 HBOa Ha03 CIHEON NaO (Reproduced by permission from Pure Appl.Chem., 1967, 15, 21) 6 E. Collinson, F. S. Dainton, D. R. Smith, and S. The, Proc. Chem. SOC.,1962, 140; F. S. Dainton and W. S. Watt, Proc. Roy. SOC.,1963, A275,447. 6 G. Czapski and H. A. Schwarz,J. Phys. Chem., 1962,66,471. $!2 Dainton In the past decade knowledge of this entity has grown apace and its properties are now more closely delineated than those of many chemicals. In pondering what to put in this lecture I considered there were two alternative approaches. On the one hand I could assume that much about the electron in condensed media was common knowledge and concentrate on some very recent develop- ments of special interest to the cognescenti or, on the other hand, I could try to paint a broad-brush picture. I chose the latter course as being more appro- priate to the occasion and as being more likely to interest a general audience.This decision inevitably means some sacrifice of rigour and precision of which Michael Faraday would certainly not have approved; but then he would probably have been rather shocked that the Faraday Lecture was given by a mere adminis- strator who in the past few years has had to mix science and politics, because on 9 July 1849 Faraday wrote to Auguste de la Rive in these terms:’ ‘For me, who never meddled with politics and who thinks very little of them as one of the games of life, it seems sad that scientific men should be so disturbed by them and so the progress of pure philosophy be much and so often disturbed by the passions of men’.What I shall do then is to describe in simple terms and in succession the theory of an electron in a condensed system, how this species may be prepared, and how its major physical and chemical properties have been elucidated, and conclude with some illustrations of its utility in chemistry. 2 On ‘Dropping’ an Electron into a Liquid or an Amorphous Solids An electron placed in an assembly of mono- or poly-atomic molecules of zero electron affinity will experience short-range repulsions and induce in them very rapid atomic polarizations which fluctuate as it moves through the medium. It is known as a ‘quasi-free’ electron and although its energy, VO,will certainly be different from its energy in the gas phase there is no known experimental means of measuring it.This energy will change with changing medium density and is likely to decrease with increasing polarizability of the atomic constituents. As the density is increased a point will be reached at which the migrating electron will find itself in a spherical cavity, escape from which may be temporarily prevented by the symmetrical and close array of the molecules constituting the cavity wall. If it is confined long enough for translational and rotational dis- placements of the molecules to occur then, as described below, its energy may fall below VOwhen it is no longer ‘quasi-free’ but ‘localized’. Potential electron-trapping qavities of varying sizes will exist in amorphous solids and in fluids (including gases of more than a critical density) but not in perfect crystals or too attenuated gases.The energy of the confined electron relative to its quasi-free condition will then have two parts viz: electronic and cavity reorganizational. The latter has the following components: (i) that due to See ref. 1 p. 357.* There are many accounts of this theory. The latest, which contains references to most of the previous work, is to be found in ‘Electrons in Fluids’, ed. J. Jortner and N. R. Kestner, Springer, New York, 1973, pp. 1-25. The Electron as a Chemical Entity mechanical work involved in dilation or contraction of the cavity (this is not significant except at high pressure); (ii) that due to the change in surface free energy of the ‘bubble’ as dilation or contraction occurs; (iii) that due to rotation of the molecules of the cavity wall so that the positive ends point to the centre; (iv) the destabilization due to electrostatic repulsions of the permanent and induced dipoles aligned as in (iii) ;(v) inter-hydrogen-atom repulsions due to ‘crowding’ of these in the centre of the cavity: a destabilizing effect which increases in the order ROH, HzO, NH3; (vi) that due to polarization of molecules outside the cavity wall; these are usually treated as a continuum.The electronic energy is evaluated from the usual equation assuming a particular binding potential V(r),which will depend on the number, arrangement, and chemical nature of the molecules of the cavity wall, which of course also influence the reorganizational energy, especially through (ii), (iv), and (v).The optimum cavity size is then determined by minimizing the total energy for the 1s state of the electron. Calculations of this kind can be carriedout for other states and the variations of energy with radius for 2s and 2p (i.e. optically combining) states are shown in typical cases in the configuration diagrams. Such calculations enable the following predictions to be made: (i) whether a stable cavity exists, and, if so, the most probable number of molecules in the cavity wall; (ii) the optimum cavity size and therefore whether ‘dropping’ electrons into the medium will cause dilation or contraction; (iii) estimates of the heat changes accompanying trapping; (iv) because dilation involves structure loosening, it will be accompanied by a diminution of viscosity; (v) possible optical transitions, their energies and oscillator strengths, which in turn allow the form of the optical spectrum, the photoconductivity and photoelectric thresholds to be estimated.A typical configuration diagram is shown in Figure 3. For stable cavities, the optical spectrum is the most important property and it is here that the first important discrepancy with experiment is found. The predicted 2p t 1s lineshape is always narrower and more symmetrical (see Figure 4) than the observed spectrum which has a marked high-energy tail. Figure 5 shows the spectra of electrons solvated in different liquids which illust- rate this point.9 It also shows that small changes in chemical structure can cause significant shifts of Amax which Dorfman has shown cannot be correlated with changes of the dielectric constant.As yet the theory cannot account quantitatively for these shifts but there are qualitative arguments to suggest that within the same homologous series component (v), i.e. lateral repulsions, is a dominant influence. Thus Amax is greater for branched than straight-chain alcohols and for ethers than for alcohols, a result which may be interpreted by enforced greater cavity size due to repulsions of the increasing number of H atoms forming the cavity wall. a L. M. Dorfman, F. Y. Jon,and R. Wageman, Ber. Bunsengesellschufr.phys. Chern., 1971, 75,681. Dainton I 0.8 0.6 0.4 0.2 0 -0.2 -0.4 % 1 P -0.5 $ -0.6 -0.8 -1 -I .2 r 1 1 t 1 r r 1.5 2 2.5 3 3.5 4 R/i Figure 3 Configuration diagrams for electron in liquid ammonia showing energy as a function of cavity radius for 1s and 2p states for four nearest neighbour NH, molecules (Reproduced by permission from ‘Electronsin Fluids’, ed.J. Jortner and N. R. Kestner,Springer, New York, 1973) 329 The Electron as a Chemical Entity 203 KN=4 I Yo =-0.5 eV F Total 1.1 1.2 Energy/eV Figure 4 Predicted spectrum for es-in liquid ammonia at 203 K.Left-hand curves corres- pond to four and six nearest neighbours. Right-hand curve is an appropriately weighted composite. Taken from Figure 4, p. 15 of ref.8 where underlying assumptions are explained (Reproduced by permission from ‘Electrons in Fluids’.ed. J. Jortner and N. R. Kestner,Springer, New York, 1973) 3 The Preparation of Solvated and Trapped Electrons in Condensed Media A. Definitions.1t is useful to distinguish between a solvated electron, designated es-, and a trapped electron, designated et-. The former is an electron in a liquid which is either non-ionized or only very weakly so, such as water, ammonia, amines, ethers, alcohols, or hydrocarbons. It can move through the liquid by processes to be described and can be ascribed a numerical diffusion constant. The latter on the other hand, is an electron in an amorphous, non-ionic solid (i.e. glasses) and cannot move without gross disruption of, or detachment from, its cavity, becoming a mobiZe electron, ern-, which has higher energy than et-and moves rapidly through the medium until it is retrapped at a suitable site which is an actual or potential cavity.As initially produced by any of the methods described below, an electron is, of course, unsolvated or untrapped and is progressively changed into the final fully solvated or trapped state. In the initial state it is sometimes alluded to as dry and in the intermediate state as damp or moist (a reference to water as the medium). These descriptions, though imprecise, are none-the-less useful because they describe realities. B. Preparation.-The obvious method of preparation would be to bombard r animonia 5--n I H0 water ethylenediaminemethanol I7 I 1 I I 1 ?' I I I + I' I I I I 4'4 44+ '4 +I??..400 600 800 1000 1200 1400 1600 1800 2000 2200 Wavelength /nm Figure 5 Absorption spectrum of ea-in various solvents (Reproduced by permission from Ber. Bunsengesellschaftphys. Chem., 1971,75,681) The Electron as a Chemical Entity liquids or glasses with a stream of electrons from some external source. Although it is effective, as may easily be demonstrated by placing blocks of colourless plastics, such as poly(methy1 methacrylate), in the beam of an electron accelerator when the characteristic colour of et- becomes apparent, it is of little use because the whole material develops a negative potential. To avoid this it is usual to extract the electrons from molecules of the medium or from solutes by means of heat, light, or ionizing radiation (a-, p-, and y-rays or fast charged particles delivered by an accelerator) or from solids in contact with the medium which are charged to a negative potential.The most useful methods are described below: (a) Spontaneous. There are two kinds of thermal reaction. The first is the spon- taneous loss of an electron from a substance, M, of low thermionic work function, $M, (or if the vapour is used, of low ionization potential, IM)when it is brought into contact with an appropriate solvent.loa Equation (8) summarizes this process, where S denotes the enthalpy decrease on solvation. It is clear why the alkali metals, having low values of #M, so easily form es- in contact with water, ammonia, or amines.They are less effective in less polar aprotic solvents such as the ethers and hydrocarbons where the solvation energies are much lower. However, it is possible to increase SM+by the use of small amounts of the cyclic ethers (‘crowns’ and ‘cryptates’) which co-ordinate and sequester in soluble form the cation and thus drive reaction (8) further to the rightlob This may cause a further complication in that Ms-has a stability comparable to that of es- and a spectrum which overlaps it, and so complex equilibria involving M, es-, Ms+, and Ms-are established which are difficult to disentangle. For these reasons, and also because M is used as a metal mirror or in mercury amalgam and the concentration distribution of es- initially produced is non-uniform, this method is not of high utility.When used in the rotating cryostat,loc in which jets of M atoms and solvent molecules are sprayed alternately onto a rotating glass cylinder at 77 K, reaction (8) does provide a means of obtaining good e.s.r. spectra of et-. The second method relies on the reversal of reactions of type (2) in protic solvents which dissolve freely the alkali-metal salt of the solvent. Thus if the solvent designated as H2A is capable of ionizing into H+ and AH- it is possible to establish the reaction (9) by passing hydrogen atoms into or generating them within H2A containing high concentrations (> mol 1-l) of NaAH. This was first established for water by Jortner and Rabani.lla The hydrogen atoms He + AH-4es-(9) lo (a) See ‘Solutions Metal-Ammonia’, ed.G. Lepoutre and M. J. Sienko, Benjamin, New York, 1964, and ‘Metal-Ammonia Solutions’, ed. J. J. Lagowski and M. J. Sienko, Butter- worths, London 1970; (b)J. L. Dye et al., J. Amer. Chem. SOC.,1970,92,5226; Ber. Bunsen-gesellschaft. phys. Chem., 1971, 75, 659; J. Phys. Chem., 1972, 76, 2975; (c) J. E. Bennett, B. Mile, and A. Thomas, Nature, 1964, 201, 919. l1 (a)J. Jortner and J. Rabani, J. Amer. Chem. Soc., 1961, 17, 388; (b) W. L. Jolly in ref. 8, p. 169. Duinton need not necessarily be free; it suffices if they are loosely bound to a molecule as in the radical CH&HOH when reaction (10) can take place: CH3cHOH + C2H50s--CH3CHO + es-In passing it may be noted that if the radical anion A*-can be formed then reaction (1 1) is in principle possible and in the case of ammonia (A.-= NH-) H2 + A*--+es-; (1 1) this is known as the Jolly reaction,llb though whether this proceeds as written or is two-stage, i.e. H2 + A*---t He + AH---f es-, remains a moot point. (b) Photochemical. Electrons can be photo-detached from plates,12 films,13 cathodes,l4 and reducing solutes. The disadvantage of the first three methods is that the electrons are formed close to the enitting surface but the advantage is that the method can be used with any solvent towards which the surface is inert. If the surface is negatively polarized the drift velocity of the electrons released can be measured, so that it becomes possible to determine whether they are ‘localized’ or ‘quasi-free’.By this means it has been shown15 that in some cases neither of these descriptions is exclusive, i.e. the electrons show ‘quasi-free’ conduction between cavities in which they are localized for only a short time, r, and from which they may be thermally ejected so that r = ro exp (&RT) where I? is the average escape activation energy. If pqfis the mobility of the quasi-free electron between cavities in which the electron is immobilized then the observed mobility, p, will be given by the equation in which n is the trapping frequency. In unsymmetrical hydrocarbons nrS 1, so that p 21 (pqf/nro)exp(-4-8the range inare ,!?ofand the values E/RT) kJ, indicating very low binding of energy of e- to the trap.The most useful photochemical method, because it ensures a uniform dis- tribution of es- throughout the solvent, is that of detachment of an electron from a solute ion or molecule, Ds, according to equation (12). n may be unity, Ds + nhv -+ Ds+ + les-(12) i.e. the process is monophotonic. This is true for many ions and molecules in aqueous solution16 as is illustrated in Table 1. The biphotonic mode, i.e. n = 2, requires for effectiveness with light of moderate intensity, say, 1014 quantum cm-2 s-1, that an intermediate state of some longevity such as a triplet state is I2 R. A. Holroyd and M. Allen, J. Chem. Phys., 1971,54,5014. l3 F. W. Froben and J. E. Willard, J. Phys. Chem., 1971,75,35. l4 G.C. Barker, Ber. Bunsengesellschaft.phys. Chem., 1971,75,728. Is H. T. Davis, L. D. Schmidt, and R. G. Brown, ref. 8, p. 393. Reviews of this method are to be found by G. Stein, ‘Actions Chimiques et Biologiques des Radiations’, Masson et Cie, Paris, 1969, 13, 119; and H. I. Joschek and L. I. Grossweiner, J. Amer. Chem. SOC.,1966,88, 3261. TheElectron as a Chemical Entity Table 1 Quantum yields of eaq-formation from DBq Daq Br- C1- OH- S042- I- Fe(CN)s4-Fe2+ PhNHz PhOH AJnm 185 185 185 185 254 254 254 254 254 @ 0.34 0.43 0.11 0.71 0.23 0.67 0.06 0.16 0.025 readily accessible17 and from which the ionization level may be reached by an allowed optical transition. It has the virtue that light of longer wavelength than is usually effective for monophotonic processes can be used but, clearly, high intensities such as those provided by laser sources are advantageous.Typical solutes capable of biphotonic photoionization are anthracene and TMPD (tetramethyl-p-phenylenediamine)and their derivatives in hydrocarbon solvents. The advantage of the photochemical method is that it permits the use of flash photolysis to obtain both the spectra and rates of reaction of es-. This is some-what offset by the facts :(i) that the monophotonic quantum yields are not large (see Table 1); (ii) that Ds+ is often an oxidizing agent and the reverse of reaction (12) may be rather fast; and (iii) that the absorption spectra of either Ds or Ds+ or both may overlap that of es-; e.g.when D = TMPD Ds+ is the Wiirster's salt cation which absorbs strongly in the red end of the spectrum. Disadvantages (i) and (iii) cannot be eliminated but it is sometimes possible to convert Ds+ by reaction with a molecule to form a non-absorbing product which does not react with es- [see also (c) below]. (c) Radiolysis of the Solvent. When a fast charged particle is decelerated by passage through an assembly of molecules it loses energy by inelastic collisions in some of which the molecules are ionized. On average the energy lost from all causes per ionization event is ca. 30 eV and the indications are that the average energy of each ejected electron is ca. 10 eV, though this is the average of a wide distribution of actual values. This secondary electron is therefore decelerated in events which involve ionization only rarely and it is eventually thermalized some distance from its co-product, the ion derived from the parent molecule.The thermalized electron may then be solvated or trapped and the ion-molecule may decompose or react with another molecule(s) of solvent. For example in the case of water we may depict the process as: H204AA+HzO' + *e-(1 3)$ deceleration1 f-unsolvated + OH eaq-= hydrated or aquated H+~~ electron (14) in which the rates of formation of Haq+and eaq- must both be limited by the dielectric relaxation rate. The advantages of this method of generating es- are numerous. In the first place, no solute is necessary and the complications inherent in (b) above are absent.I'G. E. Johnson and A. C. Albrecht, J. Chem. Phys., 1966, 44, 3162, 3179; G. Beck and J. K. Thomas, ibid., 1972,57,3649. Dainton Secondly, since the physics of the inelastic collision processes are not strongly influenced by the state of aggregation the primary act is insensitive to whether the molecular ensembles are solid, liquid, or gaseous. Thirdly, since the energy transferred to a molecule is primarily dependent on its proximity to the track of the fast charged particle, solutes in concentration c 10-1 mol 1-1 and are virtually immune from direct excitation by the electron. Finally, accelerator beams are readily pulsed to s and therefore the analogue of flash photolysis, namely pulse radiolysis, is a valuable and now the major tool for elucidating the spectrum and reaction kinetics of eS-.l8 Slightly offsetting these advantages is the fact that, as in the photochemical case, the co-product of the primary act or its breakdown product are free radicals which are therefore electron acceptors, i.e. oxidants, and also that both these species are distributed in ‘spurs’ and in isolated pairs in the close vicinity of the track of the fast charged particle. The first of these countervailing influences can often be removed by the addition of a substance from which the radical co- product can extract a hydrogen atom.Thus in the case of water, saturation with hydrogen gas or ethylene or the presence of M ethanol is sufficient to capture all the hydroxyl radicals formed in reactions (13) and (14), provided the dose-rates and doses are not too large to use up most of the added solute.In each case the OH radical is converted into a reducing radical, respectively He, HOCH2CH20, and CHdHOH, which does not react with es- and which in the cases of H- and CHdHOH may be converted into es- in highly alkaline solutions via reactions exemplified in equations (9) and (lo), thereby augmenting the rate of formation of es-. The non-uniform, track-like distribution of the primary species is not so serious a handicap as might at first sight be presumed. This is not the occasion for a discussion of spur and track kinetics; suffice it to say that at dose-rates generally used in pulse radiolysis with >1 MeV electrons or using kilocurie 7-ray sources and with radical scavenging solutes in the concentration range ca.10-3 to 10-2 moll-1 the kinetics of the subsequent reactions are adequately accounted for by assuming a uniform distribution of primary species. For all these reasons the radiolysis method is the most convenient, controllable, and flexible means currently available for producing es- or et- in solids, liquids, and compressed gases. (d) FieEd Emission Methods. Field emission from cathodes,lg whilst leading in suitable circumstances to formation of es-, is of limited utility because it needs high electric gradients, because it is restricted to media which are electrically insulating, and because it generates electrons in a highly non-uniform distri- bution. See ‘Pulse Radiolysis’, ed.M. Ebert, J. P. Keene, A. J. Swallow, and J. H. Baxendale, Academic Press, London, 1965; A. K. Pikaev, ‘Pulse Radiolysis of Water and Aqueous Solutions’, Indiana University Press, London, 1967; M. Anbar, M. Bambenek, and A. B. Ross, Nut. Stand. Data Ref. Service (US); 1973, 43, 1-67; E. Watson and S. Ray, ibid., 1972, 42, 1-22. W. F. Schmidt and W. Schnabel, Ber. Bunsengesellschaftphys. Chem., 1971,75,654. The Electron as a Chemical Entity 4 The Physical Properties of Solvated or Trapped Electrons A. The E.S.R. Spectrum and Cavity Structure.-By definition, the free electron is the simplest free radical and paramagnetic species. Its e.s.r. spectrum would therefore be expected to be a singlet modified only by the number and type of atoms having non-zero nuclear spins with which it interacts, the precise form of which should therefore allow inferences to be drawn about the conformation of host molecules in the cavity walls.This expectation is fully realized. eaq-, eamm-,20 and et- in aqueous, ethereal, alcoholic, and amine glasses21 all show strong singlet e.s.r. signals with no fine structure which, in the case of et-, are markedly subject to power saturation. The only case for which it has been claimed that hyperfine structure is observable is when et- is produced in the rotating cryostat,22 and Bennett, Mile, and Thomas have used the weak satellite lines as a basis to argue about the number and arrangement of hydrogen atoms in the cavity wall.However, it must be remembered that in this technique alternate layers of solvent molecules and alkali-metal atoms are sprayed on to a rotating surface at 77 K and it is not certain how far the electrons are from the layer of alkali metal or to what extent the solvent layer is microcrystalline. Much better evidence can be obtained in the case of hydrogen-containing glassy media by isotopic substitution with deuterium because the hyperfine splitting constant of the proton is 6.5 times that of the deuteron. This substitution would therefore be expected to cause line-narrowing, provided the molecules lining the cavity wall are oriented with the hydrogen atoms pointing towards the centre as would be expected from the permanent dipole moments of the OH bond in aqueous and alcoholic media.The linewidths for et- in these glasses are given here23a and amply verify this prediction; this is further confirmed by the fact that the mean relaxation time TIT^ depends on the hydrogen isotopic com- position. Host molecules H2O D2O CH30H CH30D CD30H Linewidth in glasses of et- 14 6 11f2 6 14 In passing it is important to recognize that these signals, like the optical absorption spectra of et-, are light sensitive. Care must be taken to avoid this photo-bleaching which, when visible light is used, may be due to photo-decom- position of the traps. In methanolic glasses, for instance, this reaction is et-+ hv (visible) -CHzO-+ H2 (1 5) so that on exposure of the irradiated glass to visible light the singlet e.s.r.spectrum of etalc-, is replaced by the triplet of CHZO-.~~~ E. C. Avery, J. R. Remke, and B. Smaller,J. Chem. Phys., 1968, 49, 951 ; C.A. Hutchison and A. C. Paston, Rev. Mod. Phys., 1953,25,255; J. Chem. Phys., 1953,21,7959. dl P. B. Ayscough, R. G. Collins, and F. S. Dainton, Nature, 1965, 205, 965; F. S. Dainton and G. A. Salmon, Proc. Roy. SOC.,1965, A285, 319; F. S. Dainton, G. A. Salmon, and J. Teply, ibid., 1965, A286, 27; F. S. Dainton, G. A. Salmon, and C. von Sonntag, ibid., 1969, A313, 31. See ref. 1Oc. 83 (a) H. Zeldes and R. Livingston, J. Chem. Phys., 1959, 30, 40; (b) F. S. Dainton, G. A. Salmon, and P. Wardman, Proc. Roy. SOC.,1969, A313, 1. Dainton B. The Equivalent Conductance, Mobility, Diffusion Constant, and Stokes Radius of es-.-These quantities are related by the equations u (mobility/cm s-1 dyn-1) = 1.6 x 10-12 D (diffusion constant/cm2 s-I)/kT = 1.04 x 10-5 h (equivalent conductancelcma 0hm-l equiv-l) (16) and rs (Stokes radius/& = 0.82/(viscosity/poise)x (17) The measurement of Xe-s, is therefore important as the source of values of the diffusion constant, De-s, and the Stokes radius, rs, e-s.This may be achieved by measurement of the changes in conductance during pulse radiolysis of pure liquids. Thus in the case of water the increment in equivalent conductance, AAa, is seen from equation (14) to be due to the formation of two additional conducting species, namely, Haq+ and eaq-and therefore A& = AH^^+ + he,,-Schmidt and Buck2* were the first to carry out measurements of this kind and these were repeated with greater precision by Barker and Fowles24 who also applied the method to eaic-.It is unnecessary to use pulse radiolysis conductivity when stable solutions of es- are available and is known for the counter ion as in the case of ammoniacal solutions of alkali metals. The results for the solvated electron in water and liquid ammonia at room temperature are shown here and illustrate the very high mobility of es-, which u/cml s-1 dyn-1 D/cm2 s-1 rs/A-aq 1.8 x 10-3 4.8 x 10-5 0.5 eNH,-10-2 2 x 10-4 0.35 in water is comparable with that of the hydroxide ion. Since rSis much smaller than the radius of the cavity enclosing and defining es- as calculated from existing models to account for xmsx for the 2p t 1s optical transition (see Section 2), it seems unlikely that the whole solvation shell moves through the medium as an integral unit. More probably the motion of the charge is accom- panied by a (partial?) disintegration of the cavity; perhaps a rotation of one or more molecules of the cavity wall enables the electron to jump to a downfield site where it is temporarily retrapped.This jump distance cannot be large in polar liquids but in non-polar liquids it is large, and in this phase of its migration the electron is to be regarded as probably q~asi-free.1~~~~ C. The Ionic Atmosphere Relaxation Time, .ria.-The solvation of an electron by a solvent containing an electrolyte is followed by the build-up of the ion-atmos- phere of es- in which cations predominate.The sequence of these two processes in water is depicted in equation (19). If a solute Sz reacts with the hydrated K. H. Schmidt and W. L. Buck, Science, 1966, 151, 70;G. C. Barker P. Fowles, D. C. Sammon, and B. Stringer, Trans. Furaduy SOC.,1970,66, 1498. ac See ref. 12, ref. 8 p. 393, and W. F. Schmidt and A. 0.Allen,J. Chem. Phys., 1970,52,4788 The Electron as a Chemical Entity 71678 ionef--eaq--eaq-+ atmosphere electron the half-life of the hydrated electron, t+, will be 0.7/k20[SO] and if this eaq-+ S2-t SZ-l (20) reaction is diffusion controlled, the value of D in the preceding table indicates that by varying [S] from 10-4 to 1 mol 1-1, t+ can be reduced from 10-6 to IO-lO s.Since Debye and Falkenhagen have shown that, at 25 "C 'Ti8 = 10-l' D//A/Ia (20) it is evident that by adjusting [Sz] the hydrated electron can react either after establishing its ion-atmosphere i.e. t+ > TZ~,when k20 will depend upon ionic strength according to the Bransted-Bjerrum equation (this is the basis of the proof of the existence of eaq-; see Section 1) or before, i.e. t+ < Tia. In the latter case the activity coefficient of eaq-, i.e. fe, will be unity and independent of p, whilst the transition state [Sz . . . e-]* with a charge z -1 will have an ion-atmosphere appropriate to an ion of charge zprovided it changes to product before it has time to establish the ion-atmosphere appropriate to charge z -1.This problem has been fully treated by Logan26 and in the case of z = -1 a more simple treatment involving putting loglo f, = -0.51~(z -l)f(,u), leads to the result that, at 25 "Cin water, when t+ > Ti&,loglo k20 = constant + l.O2f(p) (21) when t+ < ria, logio k2o = constant + 0.5lf(/~) (22) Hence if the ionic strength is kept constant and the concentration of a reactive univalent anion is increased the apparent rate constant should diminish by an amount of 0.51 f(p). This was observed by Logan and Dainton26 for the case S-= 13-and is illustrated in Figure 6 using as the reference competitive substrate the uncharged molecule N2O. It is obviously desirable to test this idea more rigorously by using ultra-fast pulse techniques and single solutes of different charge.D. The ThermodynamicProperties of %-.-If the rate constants for the decom- position of es- into H and HA-and for the reverse reaction, i.e. k23 and k-23, es-+ H + HAs-(23) can be measured and if the ionic product [H+s] [HAs-] and the free energy of solvation of the proton are known, then AGO24 and hence the redox potential on the hydrogen scale can be calculated. Combining these data with values of --+ Hs++ es-(24) 26 F. S. Dainton and S. R. Logan,Proc. Roy. Soc., 1965, A287,281; S. R. Logan, Trans. Faraday Soc., 1967,63,3004,3009. 338 Dainton -1.0 -0.8 -0.6 I I I Figure 6 To illustrate the smaller rate constant of the reaction eaq-4-Is-when the electron reacts with Is-before forming its ion-atmosphere, i.e.t+ < rig(Reproduced by permission from Proc. Roy. SOC.,1965, -7,281) Table 2 Signijicant thermodynamic quantitiesa for es-Reaction Ammonia Water AHo' ASoO AGoo Eho AH# ASoO AGOQ Eho &H2-fHs+ + es-157 -101 187 1.58 260 -57 277 2.87 eg--es--97 48 -111 --160 -8 -157.5 -QH and AG in kT mol-l; ASin J mol-l deg-l; Eho in volts. the standard entropy of solution of H2, He, and es-, the last of which can be obtained by semi-empirical relations connecting AS0 of univalent anions and their radii, the thermodynamic quantities in Table 2 are obtained.27a These have a number of interesting features. Firstly, metals with a potential >2.87 V would be expected to form hydrated electrons on contact with water whilst those having values less than this would be expected to be unaffected or to react only slowly or on warming.This is exactly consonant with what is known of the alkali and alkaline earth metals' behaviour. Secondly, the enthalpy and free energy increments on desolvation of es- in both solvents are only slightly greater than ch/hmaxof the optical absorption spectra of es-, which suggests that optical excitation of es- brings the electron to a level which is close to and may *' (a) G. Lepoutre and J. Jortner, J. Phys. Chem., 1972, 76, 683 where references to earlier work are also given; (b)R. R. Hentz and D. W. Brazier, J. Chem. Phys., 1971,54,2777. 339 The Electron as a Chemical Entity overlap the conduction band, a fact which has implications for the photo- conductivity of et- in aqueous and ammoniacal glasses (see Section 5).Thirdly, ASo for the solvation of e- in ammonia is positive which suggests that in this solvent the solvation process is also structure breaking. In principle, estimates of the partial molal volume of es- can be made by studying the effects of pressure on k23 and k-23. In practice, indirect methods have to be used27b and these lead to a value for eaq- between -1.7 and 2.7 cm3, which, provided allowance is made for electrostriction, is not inconsistent with a cavity radius of ca. 1 8, for eaq-. 5 Spectroscopic and other Evidence concerning the Cavities A. The Effwts of Pressure and Temperature on the Spectrum.-The configuration diagrams show that the smaller the cavity radius the larger the energy gap, AE, between the 1s and 2p levels. The distribution of 1s states of different radii can easily be calculated from the form of the configuration diagram and the assump- tion that the cavities are in thermal equilibrium with their surroundings i.e.Nri = Nr, exp[(Er, -Er,)/kT]. Combining these two pieces of information with the transition probabilities leads to the expectation that the absorption spectrum will be slightly unsymmetrical about ch/hmaxand show a Zow-energy tail. In fact a high-energy tail is always observed. This indicates that either one or both of the 1s and 2p energy-radius theoretical relationships is incorrect. If this conclusion is rejected then a non-Boltzmann distribution of 1s states must be presumed and can be found empirically so that the observed spectrum is re- produced.This has been carried out for earn-. Any change of an external variable which causes a diminution in the average size of the cavities is, because of the relationship between &?Z and r, likely to cause hmaxto move to shorter wavelengths. This ‘blue shift’ is observed in both liquid ammonia28 and in liquid aqueousz9 and alcoholic systems30 when the temperature is reduced or the system compressed, as illustrated in Figures 7 and 8. If the configuration diagrams are accepted the data allow crude estimates to be made of the isothermal compressibility and isobaric thermal expansibility of the cavities, and the indications are that the cavities are ‘softer’, i.e. more compressible and expansible, than the bulk liquids.B. Direct Evidence for a Range of Trap Sizes for et-.-The point has been made that if the electron trapped in a glass is not decomposed by light [cf. reaction (15)] the absorption of a quantum may drive an electron from its trap. If therefore the incident light has a wavelength span covering only a portion of the absorp- tion spectrum due to the whole distribution of traps present in the system, the population of those traps which absorb the light will be depleted selectively. The s8 R. Vogelsgesang and U. Schindewolf,Ber. Bunsengesellschuftphys. Chem., 1971,75,651. Ls F. S. Dainton, Ber. Bunsengesellschaft phys. Chem., 1971,75, 608.*O R. R. Hentz, Farhataziz, and E. M. Hassan, J. Chem. Phys., 1971,55,4974; M. G.Robinson, K. N. Jha, and G. R. Freeman, ibid., 1971, 55, 4933. Dainton 63'C t-1 I I I I I J -80 -40 0 40 80 120 160 0 400 800 1200 1600 TIT platm Figure 7 The efect of temperature andpressure on Amax of the absorption spectrum of es-in ammonia (Reproduced by permission from Ber. Bunsengesellschaft phys. Chem., 1971,75,651) electrons then liberated may then fall into holes forming et- species which do not absorb the incident light, in which case they will remain and augment the light absorption in the wavelength range outside that of the incident light. The result should therefore be a photodistortion of the initial spectrum and, in principle, this process should be reversed by re-illumination with light different in wavelength from that of the initial incident light. This photo-shuttle process is represented in equation (25).Furthermore, if there is no possible fate for a photoexcited electron other than retrapping, the total number of trapped electrons will remain constant and throughout all photo-shuttle processes which distort trap distributions the absorption spectrum should satisfy the condition SOD dv = constant. This photo-shuttle under this fully compensatory con- dition has been realized3l for an assembly of trapped electrons in alkaline aqueous glasses at 77 K and is displayed in Figure 9. The view that photo-shuttling perturbs the trap depth and size distribution is confirmed by the observation31 that the e.s.r.signal field strength saturation characteristics change when the optical absorption spectrum is displaced. Moreover, when the spectrum is blue-shifted, with concomitant increase of the 31 G. V. Buxton,F. S. Dainton, T. Lantz, and P. Sargent, Trans. Faraday Soc., 1970,66,2962. The Electron as a Chemical Entity 342 Dainton 2.1 -0 2.0 -1.9 1.9 -U 0 10 M-OH' present work 60 0 Co y-Irradiation 10M 4H- 1.8 -0 H20 Taub and Eiben 60 m Co y-Irradiation Taub and Eiben 0 Shubin etal. 1.7 -OGottschall and Hart I I100 150 200 250 300 Temperature/K Figure 8(b) The eflect of temperature on chlX,ax of eaq-(Reproduced by permission from Ber. Bunsengesellschaftphys.Chem., 1971,75,608) average trap depth and therefore decrease in the trap radius, the mean relaxation time 47172 decreases; a result which is consistent with shortened spin-lattice relaxation time to be expected from the greater proximity to the electron of the nuclei in the cavity wall. Whilst experiments of this kind demonstrate the reality of the distribution of traps of different depths they do not answer the question as to whether the light absorption process detraps the electrons, making them capable of moving to other sites where they can be retrapped or whether the electrons remain in their cavities and the effect of light absorption is merely to change the shape of the cavities and thereby alter both AE for 2p t 1s transition and the observed spectrum.In the case of et-in alkaline aqueous and MTHF (2-methyltetra-hydrofuran) glasses this question has been conclusively answered by showing that when electrodes are inserted in the glass photo-conductivity can be estab-lished.32 It must therefore be presumed that photo-excitationmakes the electron mobile, at least to the degree of bringing it to any conduction band, and equation (26) in which em-represents a mobile electron is a better representation of actuality than equation (25)*. An interesting feature of the photo-conductivity *There is no apriori reason why the only effect of light absorbed in the wavelength range of the et-absorption spectrum should be to mobilize the electron. In alcohols for instance, at wavelengths less than Amax, reaction (15) occurs and it has been shown by Bernas et aZ.3athat in hydrocarbon glasses the detrapping excitation is shifted 0.2 eV relative to the absorption spectrum.sB I. Eisele and L. Kevan, J. Chem. Phys., 1970,53,1867; K. F. Baverstock and P. J. Dyne, Canad. J. Chem., 1970, 48, 2182; A. Bernas, D. Grassel, and J. B. Trueny, J.C.S. Chem. Cornrn., 1972, 759. The Electron as a Chemical Entity I\I I \ I \ Gain . Loss I I r 25 000 20 000 15 000 Wavenumber/cm-' Figure 9 The 'Photoshuttle', i.e. changes in optical density at diferent wavelengths caused by selective illuminationof et-in 10M hydroxideglasses at 77 K. Full line after illumination with h > 640 nm,broken line after subsequent illumination with h < 500 nm of et-formed by y-radiolysis.Lines with dots refer to the corresponding experiments with et-formed by photolysis of Fe(CN)64-(Reproduced from Trans. Faraday Soc., 1970, 66, 2962) is that, as shown in Figure 10, it may be super-ohmic. When this is the case it must be concluded that the higher the energy gained by em-from the higher Dainton RUN SLOPE 50 1 1.6 2 1.5 3 1.4 N2 20 4 1.5 X 5 1.0 3 5 10 20 Field /kV cm-' Figure 10 Photoconductivity of et-in 10M hydroxide glass at 77 K. Current-voltagecurves are plotted bilogarithmically and are for successive doses of 0.6 Mrad (Reproduced by permission from J. Chem. Phys., 1970,53,1867) electrical field, the less likely it is to be retrapped.This is interesting in that it is the only direct evidence supporting the theoretical argument that the higher the energy of an electron ejected from a molecule in the primary radiolytic act the further it will travel before solvation in a liquid or trapping in a glass. When et- is produced in a system by a method which leaves the glass with no net charge, there are balancing cations or radicals or both. These latter can The Electron as a Chemical Entity capture and destroy electrons mobilized by light and whose trajectories before retrapping intersect the site of the cations or radicals. Photobleaching therefore occurs, but since the electrons first removed will be those bound to traps closest to the cations and radicals the number of ‘jumps’ and retrapping and detrapping events before an electron is finally destroyed continually increases as photo- bleaching proceeds.Hence, even though the quantum yield for photomobiliza- tion remains constant the apparent quantum yield for photobleaching progres- sively dim in is he^.^^ This is shown in Figure 11 for alkaline aqueous glasses at OD 83K (after annealing at 140K) Quanta absorbed (arbitrary units) Figure 11 The photobleaching of et-in 10M hydroxide glass at various temperatures in the presence or absence of chromate ionsas scavenger3I (Reproduced by permission from ‘Reactions of Solvated Electrons’, MTP International Review of Science, Physical Chemistry, Series One, Volume 9, Chapter 9, Butterworths,London, 1972) 77 K. There are three interesting consequences of this mechanism.First, if the chance of retrapping increases with the ease of reorientation of molecules in the walls of potential cavities, then it will increase with increasing temperature and the quantum for photobleaching should decrease with increasing temperature. Figure 11 shows that this prediction is fulfilled. Secondly, if efficient electron capturing solutes are added the photobleaching rate should increase and this too is observed. Finally, it will be obvious that the compensatory conditions can never be absolutely attained but are more nearly realized after those electrons Dainton trapped very close to the sites of their formation have undergone geminate recombination either by preliminary thermal- or photo-bleaching.6 The Trapping and Solvation Mechanism All liquids and glasses contain ‘holes’; in liquids they fluctuate in size and position but in rigid glasses these fluctuations are so slow as to be negligible on the time-scale of most experiments. It is highly unlikely that the number and distribution of these ‘holes’ will, before occupancy, accord exactly with those necessary to accommodate an unspecified number of electrons in a thermal equilibrium distribution. The observed final distribution of configurations of cavities each containing an electron will be determined by the balance achieved between the cavity stabilization forces described in Section 2 and the disruptive forces of thermal agitation.An adjustment of the shapes of holes after they have captured electrons will therefore take place and the rate of this adjustment will be regulated by the ability of the solvent molecules to rotate and translate, which, in turn, will depend on temperature. In addition the solvation process is itself exothermic and the extent of relevant molecular reorientation will be influenced by this variable amount of heat released per hole and the rate at which it can be conducted away. Consequently two distinct experimental situations can be envisaged. Firstly, in liquids where mobility is unrestricted or in ‘soft’ glasses, the spectroscopic changes accompanying the adjustment to the final equilibrium distribution might be detectable and, secondly, in gZasses at a sufficiently low temperature the rigidity and higher local temperature gradients favouring heat loss may be sufficient to prevent the attainment of the true equilibrium distribution and a metastable distribution will be permanently preserved.A. Isothermal Optical Changes associated with Solvation.-The change from the initial distribution to the ultimately stable one implies a cavity deepening and narrowing and therefore a net enhancement of absorption at shorter wavelengths and a diminution at longer wavelengths taking place on a time-scale comparable with the dielectric relaxation time of the medium. To observe this phenomenon conveniently it is therefore usual to cool the liquid in order to lengthen the relaxation time.The phenomenon has now been observed in several systems but one of the clearest is that observed when alcohols are pulse-radiolysed. Figure 12 displays some of the data obtained by Baxendale and Wardman33 for n-propanol at 152 K using 5 ns pulses. Evidently at 10 ns only about half the electrons are not in holes of their final configuration and the remainder, which are in shallow ‘red’ traps, adjust in a half-life of ca. 60 ns. The authors point out that this time is substantially shorter than the known dielectric relaxation time for this temperature and it must therefore be presumed that the local heating can contribute to the acceleration of local molecular movement. 33 J. H. Baxendale and P. A. Wardman, Nature, 1971,230,448; J.C.S.Faraday Trans., 1973, 69,586. The Electron as a ChemicalEntity 6 a 4 400 600 800 1,000 1,200 1,400Alnm Figure 12 Spectra observedfollowing a 5 ns, 600 had pulse absorbed by n-propanol at 152 K, (a) immediately ajler the pulse, (b) 200 ns later, (c) 1 ps later (Reproduced by permission from Nature, 1971,230,448) Attempts to detect solvation of es- at room temperature have not been very successful. Even using 10 ps time resolution no changes have been observed in the spectrum of es- in water, glycol, or C1 to C4 alcohols.34 B. Metastable Distributions of et-and their Thermal Relaxation.-When alkaline aqueous glasses which have been y-irradiated at 77 K are subsequently warmed, then, as shown in Figure 13, the absorption band narrows; the width at half height shrinking from 0.99 to 0.92 eV and a blue displacement of hmaxof ca.120 A is observed which is accompanied by an increase in absorption around Amax. This enhancement of absorption is mainly at the expense of the absorption at the red and, to a lesser extent, at the blue end of the original spectrum. Whatever pretreatment a y-irradiated sample receives, warming to 150 K results in an ultimate spectrum which is always the same31 and therefore is likely to be that corresponding to the equilibrium trap distribution at this temperature. Since a rise in temperature normally produces a red shift in the equilibrium spectrum at temperatures above 140 K the inference must be that metastable distributions are possible and can persist if the system is sufficiently refrigerated.Although some of the trapped electrons are destroyed in this warming process, most are not and since some of these changes occur below 135 K, where it has been shown that bulk diffusion cannot occur, it must be presumed that most of 84 M.J. Bronskill, R.K. Wolff, and J. W. Hunt,J. Chem. Phys., 1970,53,4201, Dainton OD 'i\-I!4'/ \! ' '!!4 //&* 0 0 \\ I f I I 30000 25000 20000 15000 Wavenumberjcm'' -0Figure 13 The spectrum of et-in 10M hydroxide glass: -0 -after irradiation at 77 K and ---annealing to 150 K and re-cooling (Reproduced from Trans. FaraCiay Soc., 1970, 66, 2962) the traps which change shape do so by rotation without translation of cavity molecules. A few electrons, particularly those close to a spur or positively charged centres and therefore in a high electrostatic field do migrate, and it must be presumed that this is by a tunnelling mechanism facilitated by rotation of one or more cavity wall molecules which destabilize the original site whilst predisposing an adjacent site to receive the electron.The most striking evidence for metastable trap distribution has been obtained by Higashimura and his colleagues35~3~ for a variety of alcoholic, ethereal, and water-glycol glasses irradiated at 4 K. A typical example is shown in Figure 14 86 H. Hase, M. Noda, and T. Higashimura,J. Chem.Phys., 1971,54,2975; 1972,57,1029. H. Yoshida and T. Higashimura, Canad.J.Chem., 1970,48,504. The Electron as a Chemical Entity Wavelengthlnm 2000 1200 1000 800 600 400 OD I I I I I I -I0 80 4 0 05 00 0’ 0 I/ I I I 5 10 15 20 25 Wavenumber / cm” x lo2 Figure 14 The spectrum of et-in equivolume water-glycolglass y-irradiated at 4 K (1) and after warming to 20 K (2); 30 K (3); 40K (4);and 77 K (5)(Reproduced by permission from J. Chem. Phys., 1971,54,2976) for which the long-wave i.r. absorption disappears irreversibly on warming to 77 K with concomitant intensification of absorption in the visible region. Changes of this kind are usually accompanied by a broadening of the e.s.r. signal attributed to et- which, if the argument given in Section 5B is correct, indicates that the cavities are both deepening and narrowing.7 The Chemical Reactions of es-The reality of the solvated and trapped electron as a discrete physical entity cannot be doubted. We must now inquire into the kinds of chemical reactions in which it may participate and the factors which determine the rates at which these reactions can proceed. A. Reaction Types.-The electron is the simplest free radical and the archetypal reducing agent and nucleophile and all the reactions of es- and et- except their decomposition derive from these characteristics. Thus it will combine with other free radicals such as -OH and H-and with itself (see Table 3). The last reaction is particularly interesting in that in aqueous systems the gross chemistry suggests that the ultimate products are H2+ 20H-but leave open the interesting question Dainton Table 3 Types of reactions of eaq-and their room temperature rate constants Type Equation Rate constant at room temperature Decomposition eaq-4H* + OHaq-14 s-1 2 eaq-+-H2 + 2 OHaq-5 x loglmol-1s-1 With acids eaq-+ Haq+4He 2 x 10101mol-1 s-1 eaq-+ HCOzH+H* + HC02-8 x lO71mol-1s-1 eaq-+ HOAC--+ H* + OAC-8 x 107 1 mol-1 s-1 eaq-+ H2P04- +Ha + HP042--lo61mol-1 s-1 Capture eaq-+ 02 +02-2 x 10101 mol-1 s-1 (non-dissociative) eaq-+ CO2 -+ C02-8 x loglmol-1s-1 eaq-+ I2 -12-5 x 10101mol-1s-1 eaq-+ Cd2+-+ Cd+ 5 x 10101mol-1 s-1 Capture eaq-+ PhCl +Ph* + C1-5 x 10*lmol-ls-l (dissociative) eaq-+ PhBr -Ph- + Br-4 x lo91mol-1 s-1 eaq-+ PhI -+ Ph* + I-1.2 x 10101moI-1s-1 eaq-+ N20 +N2 + 0.-9 x 10glmol-ls-l eaq-+ C(NO2)4 -NO2 + C(N02)3-5 x 10101mol-1 s-1 eaq-+ BrO3--+ BrO2 + 02-4 x loglmol-1s-1 as to whether an intermediate state is the solvated, spin-paired, and therefore diamagnetic, dielectron shown in equation (27).Some interesting pointers may 2 eaq--+ e2,aq2-3H2 + 20H-(27) be obtained from subjecting alkaline aqueous glasses at 77 K to prolonged 7-irradiation. It has long been known that initially this causes an intensification of the characteristic purple colour but that doses in excess of 6 Mrad bring about a bleaching. Kevan and Zimbri~k~~ and Pikaev et aZ.,38 by studying the changes of e.s.r. signals caused by such massive doses have shown (see Figure 15) that whilst the OH + 0 -radical concentration grows throughout the irradiation, the singlet signal ascribed to et- passes through a maximum and declines in a manner similar to the colour.This is consonant with increasingly complete occupancy of all potential et- sites after which any additionally generated electrons can only combine with -OH (which clearly does not happen) or spin- pair with existing et-forming dielectrons. Kevan has further shown thate&- gives rise to a broad absorption band with Amax 21 10 OOO A and decomposes when the glass is either warmed to 135-145 K or illuminated by light of h > 7000A. Non-dissociative capture of electrons by a wide range of molecules, e.g. O2 -+02-, C02 -C02-, R2CO -RzCO-, quinones -Q-, polynuclear aromatic compounds, and by ions, e.g.Cd2+, Fe(CN)e3-, NO2-, Cr042-, is 87 J. Zimbrick and L. Kevan, J. Amer. Chem. Soc., 1967,89,2483; L.Kevan in ref. 16. a8 0.F.Khodzhaw, B. G. Ershov, and A. K. Pikaev, Izvest. Akad. Nauk S.S.S.R., Ser. khim., 1967, 1882, 2253. The Electron as a Chemical Entity I I I I I I 16 16 14 14 12 12 10 10 b 8 6 4 2 I I 1 I I I 0 2 4 6 8 10 12 14 60 CO 'y -Dose, Mrad Figure 15 Relative et-and 0-yields as measured by e.s.r. spectroscopy for y-irradiated10M hydroxide glass at 77 K (Reproduced by permission from J. Amer. Chem. Soc., 1967, 89, 2483) often a very facile process. Since many of the product species are coloured the pulse radiolysis method can be used to study not only the decay of es- but the synchronous build-up of the product and any subsequent reactions which the product may undergo.Much new chemistry has been revealed in this way. Many examples are known of dissociative electron capture. Indeed the high rate of reaction of eaq-with some species [equations (28), (29), and (30)] was responsible for the earliest evidence for the existence of the hydrated electrons.3J9 sg F. S. Dainton and D. B. Peterson,Proc. Roy. SOC.,1962, A267,443. Dainton eaq-+ H202 -+ HO* + OH-(28)-eaq-+ RCI R*+ CI-(29) eaq-+ N2O -+ N2 + 0-(4*OH + OH-) (30) Reaction (30) is especially important because nitrogen is inert and therefore a permanent product and the accompanying radical is the simplest and most powerful oxidizing radical. Provided it is present in concentrations greater than ca.2 x moll-1, nitrous oxide is therefore an excellent titrant for electrons however these are produced,39 and it is used for this purpose and for providing, in conjunction with the radiolysis of water, a reliable and controllablesource of hydroxyl radicals. Of course other electron scavengers are used as titrants in other media. For example tetranitromethane is useful in alcoholic media es-+ C(N02)4 -+*NO2 + C(N02)3-(31) (equation 31) because the nitroform anion is stable and highly coloured. Reactions (28)--(3 1) have been described as dissociative processes but presumably the capture and dissociation are not necessarily simultaneous and the primary electron adduct may have a finite life. This certainly is the case for N2O- in non-polar solvents40 and is probably true for RCl-. The acid of which es- is the conjugate base is the hydrogen atom (with a pK = 9.6) and it may be formed by reaction with the solvated proton or with un- dissociated acids, e.g.es-+ Hs+ (or H3Af) -+ Hs* (or + H2A) (32) es-+ H2P04---f Hs-+ HP0g2-(33) Here again, whether the mechanism is an instantaneous dissociative capture or involves an adduct of hite life as an intermediate is often uncertain. There is, for example, some evidence for the existence of H30. as the first product of reaction (32) in water. But what is certain is that the rate constants of these reactions cover a wide span. Reaction (32) is fast in water, alcohols, and ethers but slow in liquid ammonia.It is also the case that, although the hydrogen atom is a reducing agent, it is less effective than es- as an electron donor and can also act as an abstractor of hydrogen atoms from C-H bonds and is much less reactive to nitrous oxide than is es-. A striking consequence of this is that, as shown in Figures 2 and 16, varying [Hs+] has a dramatic effect on the distri- bution of produ~ts.~l In general, if es- reacts with substrate A to give product Ea (rate constant ka) and the hydrogen atom reacts only with A or some ot&r substrate to form Prb, then product yields vs. pH curves strongly resembling pH titration curves are obtained, e.g. Figure 16. In these circumstances the relation- ship (34) is obeyed, and before pulse radiolysis was available this was the principal (yield of Pra)-l = constant (1 + k22[Hs+]/ka[A]) (34) F. S.Dainton, P. O’Neill, and G. A. Salmon, J.C.S. Chem. Comm., 1972, 1001. 41 P. L. Airey and F. S. Dainton, Proc. Roy. Suc., 1966, A291,478. 353 The Electron as a Chemical Entity 2 4 8 12 PH Figure 16 The eflect of changing pH on the quantum yield of nitrogen produced by the photo-detachment of electrons from Fe(CN),4-ions dissolved in 0.005Maqueous solutions of nitrous oxide and illuminated by 254 nm light(Reproduced by permission from Proc. Roy. SOC., 1966, A291, 479) method of obtaining relative rate constants of reaction of es- with different solutes.42 Little is known about the mechanism of decomposition of es- as represented by equation (23).In the case of eaq- the first-order rate constant depends on temperature in a way which suggests a considerable negative entropy of acti- ~ation~~implying a structure-making process as rate-determining, which is presumably that of full hydration of the hydroxide ion product. It is also associated with a negative volume of activation of 16 ml mol-1 which is also consistent with structure-making and cavity collapse. The solvated electron in ammonia is at least lo8 times longer-lived than eaq- and here there is clear evidence that the negative volume of activation is even larger, a fact which is in accord with the theoretical view that the cavity in eamm- is much larger than that of eaq-. B. Factors Regulating Bimolecular Reaction Rates.-Bimolecular reactions in solution are distinguished from their gaseous counterparts in two ways. Firstly, the solvent impedes the movement of the dissolved reagents and, secondly, when the reagents encounter one another they are ‘caged’ by solvent molecules thus ensuring that they collide several times before they are able to separate.If the 42 See for example, p. 451 of F. S. Dainton and W. S. Watt, Proc. Roy. SOC.,1963, A275’447. 4s R.Olinger and U. Schindewolf, Ber. Bunsengesellschaft phys. Chem., 1971,75,693. Dainton probability of reaction at any collision is high the reactants are always destroyed when both are within the same cage and the rate-determining step of the reaction is the rate of encounter. Such processes are called difision-controlled.If, on the other hand, the reaction rate within the cage is low, then only very rarely are the reactants consumed. Most of them escape, their concentration remains uniform and the impedance of the solvent to the diffusion together of the reactants is no longer the limiting factor. Such processes are often referred to as activation-controZZedbecause it is so frequently necessary for the collision pair to power a minimum activation energy as a pre-condition for reaction. For bi- molecular reactions of es- which come into the latter category treatment by the collision or transition-state theory is as adequate (or inadequate?) as it is for any bimolecular reactions and similar types of correlation are observed. For example the rate constants of reactions of eaq- with substituted phenoxide ions follow Hammett-type relations44 and the slow reactions of eaq- with acids [equation (35)] obey the Brransted relationship,45 k35 = ~[KHB~‘~~.However, the chance of reaction on collision can also be slight because the entropy of reaction is sufficiently negative.This is expected to be the case for reactions of es- in which, probably because of collapse of the cavity, the volume of activation is negative. This is true for the reaction of eaq- with, for instance, saturated amides, where measurements of the effects of high pressures on the rate of reaction46 have led to values of -A V* in the range 7-10 ml mol-l. As illustrated in Table 3, there are many reactions of eaq- which have rate constants 2 1010 I mol-1 s-1. In these cases there is a prima facie case for considering these reactions to be diffusion controlled.The Einstein-Schmolu- chowski treatment leads to expression (36) for the diffusion-controlled rate ks, = 47rr D 8/(eb -1) where 8 = -za/E r kT constant kss (for definition see next section) between es-and an ion of charge za. In this expression r is the interaction distance and will usually be a few A, D is the sum of the diffusion constants of es- and of the other reactant (for Nreagents other than Haq+ which has a specially high mobility, D 5 x cm2 s-l), E is the static dielectric constant of the solvent so that for aqueous solutions and when za = 0, kssis expected to be ca. 2 x 1O1O 1mol-l s-l.The effects of pressure and temperature on these reactions will depend on how sensitive r, D and E are to changes in these variables. In general, D is far more sensitive than either r or E and therefore the magnitudes of the activation energy and activation volume for diffusion, which are properties of the solvent should be the major factors controlling the temperature and pressure dependence of 44 M. Anbar and E. J. Hart, J. Amer. Chem. SOC.,1964,86,5633. 46 J. Rabani, ‘Solvated Electron’, American Chemical Society, Washington, p. 242. 46 R. R. Hentz, Farhataziz, and E. M. Hassan, J. Chem. Phys., 1972,57,2959. The Electron as a Chemical Entity kss. It has been established that, for many reactions in water and methanol,47 dln kss/dT equals dln q/dT, where 7 is the viscosity of the solvent. Similarly for the reaction of es- with nitrobenzene, acetone, and naphthalene in ethanol RT dlnkss/dP corresponds to activation volumes in the range 5.6-7.5 ml mol-l, which is about the value of -RTdln r)/dP for this solvent.48 8 Some Applications of our Knowledge of ee-Faraday is reported to have justified his work on electromagnetic induction to an enquiring well-born lady with the riposte ‘Of what use is a new-born baby, Madame? Modern paymasters of scientific research require less rhetorical answers and it is perhaps appropriate to conclude with some brief mention of the usefulness of knowledge of the chemistry of the electron.Disregarding the importance of knowing about a particle of such intrinsic chemical significance as the electron, its utility rests on its intense colour and high reactivity.These make it possible to produce rapidly and in a completely controlled and measur- able way a wide variety of new chemical species and if, as is often the case the latter are also coloured, their reactions may be investigated. Typical examples of these, in addition to those already mentioned, are the following. 1. The hydrodimerization of vinyl compounds effected either cathodically or by using metal amalgams is important industrially. By pulse radiolysis of aqueous solutions of vinyl monomers, denoted by ml, it is possible to measure directly not only the rate constant of additions of eaq- to the monomer but also that of the dimerization of the protonated add~ct.4~a When ml = acrylonitrile and (m1H)z is succinodinitrile 2 mlH --+ (m1H)z (37) k37 = 1.2 x lo9 1 mol-1 s-1.It is also possible to obtain the rate constants of the first propagation step in the free-radical addition polymerization step of such monomers and of inhibition reactions caused by, for example, 02 and metal ions. 2. Mechanistically allied to this is the industrially important coupling of pyridine to form dipyridyls the mechanism of which can be shown to follow the co~rse~~b 40, es-+ C5H5N -+ CgH5N-f +(CE,H~N)~~--+(C5H4N)a + OH-3. The electron is such a powerful reducing agent and so easily converted into -OH in NzO-containing solutions that it can be used to form a whole range of new species including 02-, 03-, and their acids HO2 and HOs; unstable halogen oxides such as BrO and BrO2; metal ions in unusual valency states e.g.Cd+ and In2+. Moreover, when these reactions are diffusion-controlled, tests can be 47 M. Anbar, Z. B. Alfassi, and H. Bregman-Reisler, J. Amer. Chem. Soc., 1967, 89, 1263; J. K. Thomas, S. Gordon, and E. J. Hart, J. Phys. Chem., 1964,68,1524; and unpublished results of I. Janowsky and F. C. Cattell. 48 K. N. Jha and G. R. Freeman, J. Chem. Phys., 1972,57,1408. 49 (a) K. Chambers, E. Collinson, and F. S. Dainton, Trans. Faraday SOC.,1970, 66, 142; (b) Unpublished work cited by J. D. Rose, Proc. Roy. Soc., 1969, A312,314. Dainton made of the reactivity of these new labile compounds.In this way it was shown that the reducing power of some univalent ions of Group I1 diminishes in the order Zn+ > Cd+ =-Pb+,50 and whether the species MO2+ formed by reactions (38) and (39) are identical M+ + Oz+MOz+ (38) MZ+ + 02---+ MO2+ (39) 4. Most living cells have a high water content and pH > 5, and therefore eaq- is likely to be a primary product of radiolysis. Depending on the degree of anoxia, the eaq- will persist or be converted into 02-. Consequently it is now possible to investigate in vitro some of the reactions likely to occur in vivo when living substances are irradiated and to lay the foundations of a molecular radiobiology. The list is far from exhaustive, but I will not complete the catalogue, because I am mindful of one of Faraday’s famous aphorisms: ‘[The audience] wish for something they can comprehend.This may be deep and elaborate for the learned but for those who are as yet tyros and unacquainted with the subject must be simple and plain.’ I hope I have met the needs of the tyros. Let me conclude with two matters for those more learned in physical chemistry. A. Glass-forming Liquids and Time-Dependent Rate Constants.-In liquids which on cooling increase in viscosity and pass smoothly into a rigid amorphous solid there must be a finite temperature at which transport processes become infinitely slow. It follows that r]-1 for the liquid and the diffusion constant of solutes in it cannot follow the Arrhenius Law which prescribes a dependence on temperature of these two quantities which is incorporated in a term exp( -Et/RT), where Et is the activation energy for transport.It is difficult to test the correctness of this statement because D, 7, and the related quantity, the reciprocal of the dielectric relaxation time, can only be measured over a few orders of magnitude variation. No such limitation applies to the determination of the rate constants of diffusion- controlled reactions of es- in which time resolution down to nanoseconds can be achieved and the concentration of the substrate can be varied, so that variations of k over eight orders of magnitude can be investigated. For reactions of es- in aqueous glass-forming liquids or in MTHF with various solutes, the Arrhenius plot is non-linear and the apparent activation energy, i.e.RT2 dlnkss/dT, increases with decreasing ternperat~re.~~ The inadequacy of the Arrhenius equation is shown by the fact that the frequency factor and apparent activation energy can reach absurd values at low temperature^;^^ thus, in MTHF, values of 1034 1 mol-1 s-1 for A and 59 kcal for Eaare obtained, and in alkaline aqueous G. E. Adams, J. H. Baxendale, and J. W. Boag, Proc. Chem. SOC.,1963,241. 61 (a) G. V. Buxton, F. C. Cattell, and F. S. Dainton, Trans. Faraday SOC.,1971, 67, 687; (b) K. Kawabata, S. Okabe, S. Fujita, H. Horii, and S. Taniguchi, Ann. Reports Rad. Centre Osaka Prefecture, 1968, 8, 70. 61 See F. S. Dainton and A. K. Pikaev in ‘Progress and Problems in Contemporary Radiation Chemistry’, Czech.Acad. Science, Prague, 1971, vol. 2. 357 The Electron as a Chemical Entity solutions the corresponding values at 150 K are 5 x 1021and 15. Almost 50 years ago Tamman and He~se~~ found that the fluidity of glasses above their softening points fitted a different equation: kss = Ass expl-B/(T-To)J (40) Figure 17 shows that this equation applies to the rate of reaction of eaq- with 8-6-4d 2-01 I 1 I I I 5 10 15 20 25 30 10' (7'-135)-'/K Figure 17 The dependence of k on temperature for the reaction of esq-with in 10M aqueous hydroxide solution. The straight line corresponds to equation (40), i.e. log,& = 12.02 -890/(T -135)(Reproduced from Trans. Faraday SOC.,1971,67,697) an oxyanion in aqueous glass-forming liquids.Experiments with other solutes which react with eaq- show that Ass, B, and TOare characteristic of the solvent and largely independent of the nature of the solute. The fact that the diffusion constant of oxygen and the ionic mobilities of Na+ and OH- in the same medium fit a similar relationship with similar values of B and TOsupports the view that the temperature dependence of D far outweighs those of r and E [see equation (36)].51a This point can be tested by using the theory of diffusion-controlled bimolecular 63 G.Tamman and W.Hesse, 2.anorg. Chem., 1962,156,245. Dainton reactions as elaborated by N0yes.~4 He pointed out that if, to a solution of one reagent, the second reagent could be instantaneously introduced in uniform concentration, the true bimolecular rate constant, kr, would only be observed at t = 0.As the reagent concentrations in the immediate vicinity of each solute were depleted by reaction, so average concentration gradients would be set up and kobs would fall from k, to the diffusionantrolled value kss according to equation (41), where x = (I&)* (1 + kr/kss)/r.For x > 4, ex2 erfx is close to JFx. By pulse radiolysis the instantaneous introduction of the second reagent, eaq-, is possible and it can be shown that, for glass-forming aqueous solutions, the second condition is also readily satisfied. Consequently the optical density of esq-, OD, should decline with time according to equation (42), where [S] is the solute concentration.This has been confirmed by experiment the results of which enable kss and D to be independently derived at any temperature, and hence from equation (36) r8/(eb -1) may also be obtained.54 The results show that D does indeed obey the Tamman-Hesse equation with values of B and TOindependent of the nature of the solute and that, as shown in Table 4, Ass is close to 1OI2 1 mol-l s-l but that r, the interaction distance, is highly dependent on the nature of the solute. Evidently the larger and more oxygenated the anion the larger r, and it may be that the very large value of r for S = CrOk2- is, at least in part, due to tunnelling of the electron through its cavity and a small number of hydrogen-bonded water molecules to the hydration shell of the anion.Table 4 Values of Ass, B, TO,and r for reactions of eaq-in aqueous glass-forming liquids Medium Reagent log10Ass B rlnma Acetone 10.7 613 0.6 9.SM-LiCl H+ 10.5 557 0.7 To = 129K NO2- 11.0 666 0.4 Nos- 11.3 716 0.8 Cr042- 11.3 655 1.7 1OM-OH- NO2- 11.8 905 0.8 To = 135K Nos- 12.1 956 1.o Cr042- 11.9 824 2.0 =The quantity in this column contains a factor dependent on temperature and electrostatic energy of the interacting solute and %a-which is however not thought to differ much from case to case. 64 R. M. Noyes, Progr. Reaction Kinetics, 1971, 1, 129. The Electron as a Chemical Entity B. Selective Solvation in Two-component Liquids.-Figure 5 illustrates the marked dependence of the optical spectrum of es- on the nature of the solvent. Subject only to the effects of the ‘blue-shift’ (see Section 6) these differences are preserved at low temperatures.Moreover, when the liquids form glasses the e.s.r. spectrum is generally narrower the greater Amax. Therefore an examination of these spectra of es- in binary mixtures which are sufficiently different in chemical nature should reveal whether the mixtures are truly homogeneous in the liquid and glassy states. Such a pair of miscible substances are ethanol and MTHF. In 1965, Shields55 found from e.s.r. evidence on y-irradiated glasses that, as shown in Figure 18, 1 80 60 % OD % OD 560 nm 1200 40 20 Figure 18 The nature of the immediate environment of et-generated by radiolysis at 77 K of glasses containing ethanol and MTHF in varyingproportions.66 The ordinates correspond to the optical densities ascribable either to eaic-or eeth-expressed as a percentage of the optical density produced by the same radiation dose absorbed by the pure alcohol or pure ether the cavities containing et- consisted, surprisingly, largely of alcohol molecules down to a mole fraction of 0.35 alcohol, and that there appeared to be little evidence of the presence of et- in cavities the walls of which contained both alcohol and ether molecules.This at once raised the question as to whether the 66 L. Shields, J. Phys. Chem., 1965,69,3186. G. V. Buxton, F. C. Cattell, and F. S. Dainton,J.C.S. Faraday Trans., 1975,71,115. Dainton mixture was inhomogeneous, comprising a sea of MTHF containing islands of ethanol to which the electrons preferentially migrated and were solvated or whether, on the other hand, the mixture was homogeneous and ethanol molecules were drawn in to solvate e- preferentially.In view of the rigidity of the glass at the experimental temperature the latter hypothesis seemed less likely than the former. This question was conclusively settled by Dainton and Whewe1157 who studied the spectroscopic changes following 0.2 ps pulse radiolysis of low-temperature, liquid mixtures of these two substances. As may be seen from Figure 19, even in Xlnm 1400 1200 1000 800 0.15 8 6 0.1 GE~10-4 o.d./V 4 0.05 2 Figure 19 Spectroscopic evidence for the progressive conversion of eeth- into eaic-in 0.28 mole fraction ethanol in MTHF at 107 K.The decay of absorption at 1200 nm and the growth of absorption at 560 nm are synchronous and increase with temperature pro- portionately to exp[302/(T -SO)]. Curve A immediately after 0.2 pspulse; B, after 3 ps;C, after 10 ps; D, after 20 ps; E, after 30 ps; F, after 100 ps; and G after 180 ps(Reproduced from J.C.S. Chem. Comm., 1974, 494) a mixture containing only 0.3 mole fraction of ethanol, the species initially produced is predominantly eeth-, but this decays and is synchronously replaced by eaic-. Moreover, the temperature dependence of this reaction follows the Tamman-Hesse equation which also holds (with the same value of TO= 80 K) 57 F.S. Dainton and R.J. Whewell, J.C.S. Chem. Comm., 1974, 493. 361 The Electron as a Chemical Entity for the diffusion-controlled combination of benzyl radicals in this mixture. These results point unmistakeably to the conclusion that, before irradiation, the mixture is homogeneous and the environment of e-when first solvated is predominantly that of the molecules in numerical excess, but that these are progressively displaced by ethanol molecules in a diffusion-controlled process leading to the thermodynamically more stable eaic-. This latter species can decompose spontaneously to CzH50-and H* and it is significant that this unimolecular reaction, which does not involve translation of any molecules through the liquid, shows a normal Arrhenius dependence on temperature.It would be unwise to generalize this result for all binary liquid mixtures and it is notable that, when less compatible compounds such as an alkane and an alcohol are used, this behaviour is not observed. Thus Hentz and Kenney- Wallace58 have found that, in the pulse radiolysis at room temperature of such a mixture containing only 0.1M alcohol, only eaic- species can be detected in as short a space of time as 2~s.This finding suggests that clusters of alcohol molecules pre-exist in the unirradiated liquid. In my opening remarks I foolishly gave hostages to fortune by inviting you to judge whether I had followed the precepts laid down by Faraday. I have already mentioned one of his aphorisms and must now give you the others.They are: ‘1. One hour is enough for anyone. 2. Listeners expect reason and sense whilst gazers only require a succession of words. 3. The most prominent requisite of a lecturer, though perhaps not the most important, is a good delivery. 4. The lecturer should give the audience full reason to believe that all his powers have been exerted for their pleasure and instruction.’ I would like to express my indebtedness to many un-named former colleagues and students without whom much of this work bearing my name and referred to in this lecture could not have been accomplished. R. R. Hentz and G. Kenney-Wallace, J. Phys. Chem., 1972,76,2931.

ISSN:0306-0012    

DOI:10.1039/CS9750400323

出版商:RSC    

年代:1975

数据来源: RSC

摘要:

Bile Pigments By M. F. Hudson and K. M. Smith* THE ROBERT ROBINSON LABORATORIESy UNIVERSITY OF LIVERPOOL, P.O. BOX 147, LIVERPOOL L69 3BX 1 Introduction Bile pigments are open-chain tetrapyrrolic compounds which are derived in Nature by oxidative degradation and ring-opening of the prosthetic groups of haemoproteins. In mammals, bile pigments are simply waste materials which are excreted, and they have no biological function. However, lower animals and plants possess bile pigments which serve important, and often vital, purposes. The protective colorations of some insects are bile pigments, as is the purple defensive secretion of the sea hare. Often, the bile pigments are protein-bound, especially in algae and in green plants; these chromoproteins possess photo- synthetic activity, and one of the most important, phytochrome, is the photo- receptor for the photoregulation of growth in plants.Several books1 dealing with bile pigments and one substantial review2 have been published. All of these are now somewhat dated and it is our intention to give a contemporary view of the position of research in this field, laying parti- cular emphasis on the interaction between the chemical and biochemical disciplines. Literature received in England after 1st January 1975 has not been considered. A. Nomenclature.-Trivial names abound in the natural series of bile pigments. Two systems for numeration are at present in use; the first, which will be used throughout this review because it is consistent with all of the early, and most of the current, literature, is shown in (1).This system also emphasizes the relation- ship between bile pigments and porphyrins, for which the Fischer system of nomenclature is still widely used. The second type of numeration is shown in (2); this system represents an attempt to accommodate all pyrrole-derived pigments within the corrin system of nomenclature. We have rejected this at the present time because few, if any, bile-pigment publications have used it. The completely saturated tetrapyrrole (3) is termed ‘bilinogen’. Pigments with one and two unsaturated interpyrrole links are termed ‘bilienes’ [e.g.(4)Jand ‘bilidienes’ [e.g. (3,(6)] respectively. a,c-Bilidienes (6) are usually termed ‘rubins’. The completely unsaturated pigment a,b,c-bilitriene (7) is called a *Reprint requests should be addressed to this author.A (a) H. Fischer and H. Orth, ‘Die Chemie des Pyrrols’, Vol. IIi, Akademische Verlag., Leipzig, 1937; (h) R. Lemberg and J. W. Legge, ‘Haematin Compounds and Bile Pigments’, Interscience, New York, 1949; (c) C. H. Gray, ‘The Bile Pigments’, Methuen, London, 1953; (d)T.K. With, ‘Bile Pigments’, Academic Press, New York, 1968. a R. Lemberg, Rev. Pure Appl. Chem. (Aitsrraliu),1956, 6, 1. Bile Pigments ‘verdin’ on account of the blue-green colour exhibited by this chromophore. The analogous compounds with terminal pyrrole rings in place of the pyrrolinone rings are called bilane, biladienes, and bilatriene; these compounds do not occur naturally but they do figure prominently as intermediates in porphyrin syn- theses.3 9R 1 18 Ntl HN NH HN b (3) (5) a,b -Bilidienc Ntl HN NH HN C NH HN , (6) (7) a,c-Bilidiene Some bile pigments can be named on the basis of the porphyrin from which they are nominally derived.Thus, biliverdin-IXa (9) is the bile pigment obtained by rupture of protoporphyrin-IX (8) with oxidative removal of the ameso-carbon atom. Ring opening at the P,y, or 6 meso-carbon would give biliverdins-IXP a For a review see A. H. Jackson and K. M. Smith, in ‘The Total Synthesis of Natural Products’, Vol. 1, ed. J. W. ApSimon, Wiley, New York 1973. Hudson and Smith (lo), 4x7 (ll), and -1X8 (12) respectively. As it happens, these pigments (9)-(12) have been obtained by oxidation of the iron complex of protopor-phyrin-IX, but the scheme can also be used for any bile pigment, no matter what its origin, provided that the array of substituents can be related to that in a trivially named porphyrin.With only one exception (see later), all natural bile pigments have been shown to have the -1Xa orientation of substituents. Therefore, except where otherwise stated, the -1Xa isomer will be assumed in the ensuing discussion. V a Me V = -CH=CH, PR= CHzCH2COzR PH r PH PR PR (9) R = H (32) R = Me V Me V Me V Me Me NH HN Me (10) R = H (11) R = H (12) R = H (33) R = Me (34) R = Me (35) R = Me B. Occurrence.-In Animals. Bile pigments have long been recognized as the waste products of haemoglobin catabolism in higher animals.The principal sites of haemoglobin breakdown are the reticuloendothelial cells of spleen, bone marrow, liver, and to a smaller extent, the kidneys. Rupture of the blood pigment, with loss of the QI meso-carbon atom (as carbon monoxide4) and the iron atom (which is re-used), affords biliverdin-IXa (9). In man and mammals, biliverdin is reduced, giving bilirubin (13); the reduction is catalysed by biliverdin reductase, and the sites of highest activity of this enzyme are the spleen and liver, two of the principal sites of haem [iron(ri) protoporphyrin-IX] breakdown. T. Sjoestrand, Acta physiol. Scand., 1952, 26, 328, 334, 338; Ann. New York Acad. Sci. 1970,174, 5; G.D. Ludwig, W. S. Blakemore, and D. L. Drabkin, Biuchem. J., 1957, 66,38P. Bile Pigments Et Me PH -P" PH P" MO calculations5 predict that biliverdin will have pronounced electron- acceptor properties because the lowest vacant molecular orbital is a bonding molecular orbital. The central (b) methine bridge is predicted to be particularly electron-deficient, and this accounts for reduction at this site; furthermore, the loss in resonance energy is only ca. 1 kcal mol-l compared with ca. 8 kcal mol-l if reduction takes place at the a or c positions. Bilirubin is converted ('conjugated') into its water-soluble diglucuronide salt which is excreted with the bile into the duodenum, but hydrolysis in the intestinal tract regenerates free bilirubin, which is reduced by intestinal bacteria to give urobilinoid chromogens as the final products.6 These colourless compounds (bilinogens) produce orange-yellow urobilinoids such as stercobilin (14) and urobilin (I 5).Both stercobilin and d-urobilin are optically active, but i-urobilin is inactive. d-Urobilin is often found7 after administration of broad-spectrum antibiotics such as tetracycline. Under normal conditions a complex mixture of bile pigments is excreted in the faeces; the characteristic yellow skin pigmentation found in jaundice is caused by retention of bilirubin. The most common source of bilirubin is, however, ox gallstones, in which the bile pigment is found as the pure calcium salt. B. Pullman and A.-M. Perault, Proc.Nut. Acad. Sci. U.S.A., 1959, 45, 1476. C. J. Watson, Ann. Internal Medicine, 1969,70, 839. 'C. J. Watson and P. T. Lowry, J. Bid. Chem., 1956,218,633. Hudson and Smith Biliverdin. This green pigment is found in dog placenta,8 in the integumental cells and haemolymphs of insects such as9 the praying mantis, locust, and grasshopper, in the egg shells of some birds,lO in the bones, fins, and skin of certain fish,ll in the blue coral Heliopora coerulea,l-2 and in the root nodules of some ~1ants.l~ In addition, biliverdin-IXy (1l), the only natural bile pigment so far to be shown14 not to have the -1Xa substituent orientation, occurs in the integumental cells of the caterpillar of the Cabbage White butterfly. Aplysioviolin. This is the principal pigment in the purple defensive secretion of the sea hare Aplysia.The pigment has been characterized15 as (16) and because the sea hare contains no significant amounts of porphyrin pigment it appears likely that the bile pigment comes from the biliprotein of red algae which the hare ingests with its food. Phycoerythrobilin. This pigment is obtained from the chromoproteins of red and blue-green algae. It is cleaved from the protein under relatively mild conditions16 and has been assignedl’ structure (17). PMe PH PH PH (16) (17) Phycocyanobilin. Phycocyanin is the photosynthetic biliprotein found in blue- green algae and the prosthetic pigment is called phycocyanobilin. Its structure has been defined18 as (18). * R. Lemberg and J.Barcroft, Biochem. J., 1934,28,978. (a) M. Passama-Vuillaume, Bull. SOC.Zool. France, 1965,90,485; (b) W. Rudiger, AngewChem. Internat. Edn., 1970, 9,473. lo R. Lemberg, Annalen, 1931,488,74. l1 D. L. Fox and N. Millott, Experientia, 1954, 10, 185; W. Rudiger and L. Abolins, ibid., 1966,22,298; 1969’25,574. l2 W. Rudiger, W. Klose, B. Tursch, N. Houvenaghel-Crevecour, and H. Budzikiewicz. Annalen, 1968,713,209; W. Rudiger, W. Klose, M. Vuillaume, and M. Barbier, Experientia, 1968, 24, 1000. l3 A. I. Virtanen and J. K. Mietten, Actu Chem. Scnnd., 1949,3,17. W. Rudiger, in ‘Porphyrins and Related Compounds’, Biochemical Society Symposium No. 28, ed. T. W. Goodwin, Academic Press, London, 1968, p. 121. l5 W.Rudiger, 7.physiol. Chem., 1967,348,1554; see also ref.14. l6 C. O’hEocha, P. O’Carra, and D. M.Carroll, Biochemistry, 1964, 3, 1343; Y.Fujita and A. Hattori, J. Biochem. (Japan), 1962, 51, 89; J. Gen. Appl. Microbiol., 1963, 9, 253; C. O’hEocha and P. O’Carra, Phytochemisrry, 1966, 5,993. l7 (a) D. J. Chapman, H. W. Siegelman, and W. J. Cole, in ref. 14, p. 107; (b) W. Rudiger,P. O’Carra, and C. O’hEocha, Nature, 1967, 215, 1477; (c) H. L. Crespi. L. J. Boucher, G. Norman, J. J. Katz, and R. C.Dougherty, J. Amer. Chem. SOC.,1967, 89,3642. D. J. Chapman, H. W. Siegelman, and W. J. Cole, J. Amer. Chem. SOL.,1967, 89, 3643; see also ref, 17 (6). Bile Pigments Phytochrorne.Phytochrome exists in two forms, ‘P2(724 nm) and ‘Prr’ (665 nm), which are photochemically interconvertible.It exists in all higher plants, and in red and green algae, and acts as a photoreceptor for the regulation of growth and development in plants. Structural investigations have been hampered by the difficulties involved in obtaining large enough quantities of the pigment; a structure (19), which is not yet generally accepted, has been proposed for it.17@ 2 Structure Elucidation A. Intuitive Methods.-Much of the early structural work carried out in Germany relied heavily on intuition and application of experience already gained in the porphyrin field. The colours of bile pigments were particularly helpful in obtaining a reasonable idea of gross structure; bilinogens (3) are, of course, colourless, bilienes (4) are yellow, a,b-bilidienes (5) are violet, a,c-bili- dienes (6) are yellow-red, and bilitrienes (7) are blue or green.Fischer considered the bile pigments to be tetrapyrrylethylenes for many years, but in 1931 he appliedlg a logically pointless but intuitively superb reaction (Section 2B) to bilirubin, and from this he reasoned the structure of bilirubin.la At first, a symmetrical arrangement of the side-chains was assumed, but further work soon revealed20 that the substituent pattern is the same as in protoporphyrin-IX (8); the ramifications of this latter discovery must have been tremendously exciting and led to the initiation of a massive synthetic effort by the Munich school under the direction of Fischer and Siedel, in which a host of bile pigments of both natural and purely synthetic interest were prepared and identified.Once a certain amount of ground work had been done it became possible to identify new bile pigments on the basis of interconversion of these with known, fully characterized, bile pigments. This was not always completely successful ; e.g., Lemberg and Bader21 obtained a pigment (phycoerythrobilin ?) from phycoerythrin by alkaline hydrolysis, and identified it to be mesobiliverdin (20). At that time it was not realized that alkali induces irreversible prototropic shifts in phycoerythrobilin (17), the end product of which is mesobiliverdin (20). On l9 H. Fischer and R. Hess, Z. phvsiul. Chem., 1931, 194, 193. 2o W. Siedel and H. Fischer, Z.physioZ. Chem., 1933,214, 145. a1 R.Lemberg and G. Bader, Annalen, 1933,505, 151. Hudron and Smith (20) (2)) the other hand, Rudiger et aZ.17b were able to challenge the structure (21) proposed for phycocyanobilin by Crespi et aZ.17C on the basis that it was not compatible with the ready isomerization of phycocyanobilin (18) into meso- biliverdin (20). B. Degradative Methods.-No-n-oxidative Degradation. Before the full structure for bilirubin had been elucidated it was realized that the molecule contained vinyl substituents. Thus, Fischer applied19 to bilirubin the resorcinol fusion method22 by which vinyl groups are normally removed from porphyrins, in the hope of preparing a novel de-vinyl (deuter0)bilirubin. Only inconclusive results were obtained but, unaccountably, Fischer was moved to apply the resorcinol fusion to mesobilirubin (22), in which the vinyl groups are replaced by ethyl functions, which could not be expected to be removed in the melt.This gave a mixture of isoneoxanthobilirubic acid (23) and neoxanthobilirubic acid (24) ;20 this told Fischer that breakdown of bilirubin had occurred either side of a methane carbon linking two oxopyrromethene m0ieties.t Once the oxopyrromethene had been identified as a mixture of the two isomers (23) and (24), the structure for bilirubin followed logically. (22) (23) R = H (24) R = H (53) R =CHO (28) R = Me (27) R = Me (54) R =CHBOH tTerminal rings of bile pigments exist as the Iactam ta~tomer;'~ almost all of the early literature used the lactim form for these rings and the corresponding rings in the dipyrrolicdegradation products.la 0.Schumm, Zphysiol. Chem., 1928,178,l. 2s C. H. Gray, D. C. Nicholson, and R. A. Nicolaus, Nature, 1958,181,183. Bile Pigments Mild reduction of bilirubin with hydriodic acid in acetic acid yields24 chiefly bilirubic acid (25) with some of its isomer (26), and oxidation of these with alkaline permanganate affords the orange-yellow oxopyrromethenes xantho- bilirubic acid (27) and isoxanthobilirubic acid (28).25 NH (29) PHPH (26) R = Me(25) (60) R = CHO Oxidative Degradation. (For details of photo-oxidation, see Section 5D).With chromic acid, bilirubin yields haematinic acid (29), a compound similarly obtained from haematin, and which allowed the very early workers to deduce that the porphyrins and bile pigments were in some way related.Nicolaus26 has described a more controlled oxidative procedure, using alkaline permanganate, which can be applied to bile pigments on a micro-scale. Pyrrole-2,5-dicarboxylic acids are produced from the two internal rings of the bile pigment, and these can be identified by paper chromatography in quantities below 5 pg. The most well-known application of this procedure to bile-pigment chemistry was the demonstration2’ that the biliverdin obtained from chemical (‘coupled’) oxidation of haemin [iron(m) protoporphyrin-IX chloride] is a mixture of all four possible biliverdin isomers [(9)--(12)] and not just biliverdin- IXa as had been claimed by Lemberg.2 Clearly, using the Nicolaus procedure, biliverdin-IXa can only yield one pyrrole-2,5-dicarboxylicacid [i.e.(30)]. However, Gray et aL2’ showed that significant quantities of the pyrrole (31) were obtained from the degradation, and reached the inescapable conclusion that some of the bile pigment from the coupled oxidation of haemin must have methyl-vinyl substituents sited on the two internal rings. All four biliverdins have since been ~eparated~~s~9 as the dimethyl esters (32)--(35) (see p. 387). 24 H. Fischer and H. Rose, 2.physiol. Chem., 1912, 82, 391 ;Ber., 1912,45, 1579; 0.Pilotyand S. J. Thannhauser, Annulen, 1912,390, 191 ;Ber., 1912,45,2393; 0.Piloty, ibid., 1913, 46,1000. 25 Ref. 1 (a), p. 121. 2e R. A. Nicolaus, Rass.Med. Sper., 1960,7, Suppl. 2. 27 Z. Petryka, D. C. Nicholson, and C. H. Gray, Nature, 1962,194, 1047. 28 R. Bonnett and A. F. McDonagh, J.C.S. Perkin 1,1973,881. P. O’Carra and E. Colleran, J. Chromafog., 1970,50,458. 370 Hudson and Smith Me PH Me C0,H The permanganate method has the disadvantage that it yields only information on the nature of the two internal rings, and this was clearly not enough for structural elucidation of the algal bile pigments. Most pyrrolic compounds are transformed into maleimides on treatment with chromic acid; its use as a degradative reagent has been adapted by RUdiger.l4 Table 1 shows the oxidizing potential of the reagent to be pH-dependent. The comparative rates of oxidation of bile pigments and protoporphyrin-IX are shown in Table 2; biliverdin is oxidized approximately 100 times faster than protoporphyrin-IX, and bile pigments with unsaturated side-chains degrade much faster than the corre-sponding ones with saturated substituents.Table 1 Oxidation of biliverdin at diferent pH's14 PH t* 1.7 4.5 min 3.1 52 min 4.1 7h 5.1 24 h 5.9 29 h Table 2 Oxidation of tetrapyrroles with chromic acid at pH 1.714 Pigment t+ Aplysioviolin (1 6) 0.5 min Mesobiliverdin(20) 16 min Urobilin (1 5) 20 min Protoporphyrin-IX (8) mmin The chromic acid degradation is usually carried out under two sets of con-ditions: (i) chromium trioxide in 1 M-H2S04-'chromic acid degradation', or (ii) sodium dichromate in 1 % sodium bisulphate (PH 1.7)-'chrornate degrada-tion'.If the degradation is carried out at 20 "Cany ester linkage present remains intact in the products; at 100 "C the esters are hydrolysed. The significant advantage of this procedure is that vinyl substituents survive to yield maleimides bearing vinyl groups. Chromate degradation of biliverdin (Scheme 1) yields 3-methyl-4-vinyl- maleimide (36) in good yield from the terminal rings (A and D), but the middle rings (B and c) do not give maleimides in good yield. Instead, these rings give a 2,5-diformylpyrrole (37), which is oxidized in more acidic media (such as that found in 'chromic acid degradation') to the corresponding maleimide. The Bile Pigments oo Me (36) Me PH Me \" \ \ B\ Me PH P" H (9) (3 7) Scheme 1 Me Et 0 Mc Et Me PH OH&Ao MI2 Me Me PH 11 (37) Scheme 2 2,5-diformylpyrrole can only be obtained from the middle rings (B and c) since they bear the appropriate methine carbons.RUdigerl4 has shown also that with bilitrienes (7) the diformylpyrroles are produced from both of the inner rings, but with a,b-bilidienes (5) [e.g.mesobiliviolin (38) (Scheme 2)] the 2,5-diformyl- pyrrole (37)is derived almost exclusively from ring B. Thus, by degradation and application of chemical intuition, it is possible to establish the sequence of substituents in a given bile pigment. The degradation products can be separated by thin-layer chromatography on silica gel, and the maleimides can be detected by the chlorine-benzidine reagent,30 which is sen- sitive but non-specific.Complete identification can be achieved by comparison with authentic maleimides or by mass spectrometry. The 2,5-diformylpyrroles can be detected by the 2,4-dinitrophenylhydrazinereagent. *O (a) W. Rudiger, Z. physiol. Chem., 1969,350, 1291 ; (b) F.Reindel and W. Hoppe, Chem. Ber., 1954, 87, 1103. Hudson and Smith Me Me PH OHC @JCHON H Me PMe (36) Scheme 3 The technique can be carried out actually on a silica gel thin-layer plate. After spotting the bile pigment onto the base line, the chromic acid reagent is spotted on top of it; subsequent development of the plate in the normal way, followed by detection with the chlorinebenzidine reagent, allows identification of the constituent rings of the bile pigment.Chromate oxidation14 of aplysioviolin [(16); Scheme 31 yielded two imides, one 2,5-diformylpyrrole, and methyl ethylidene succinimide (39). Since the propionic acid side-chain was not esterified in the dialdehyde, the methyl ester (identified earlier) must have been on the middle ring, having a saturated link to one of the terminal pyrrolinone rings (see above); assumption of the -1Xa substituent orientation led14 to the proposal of structure (16) for aplysioviolin. Similarly, the products from oxidation of the pigment from the integumental cells of the caterpillar of the Cabbage White butterfly, showed14 that this pig- ment could not be a -1Xa type, and led to its identification as biliverdin-IXy. Another great advantage of the chromic acid degradation is that it can be carried out on the purified biliprotein itself.l4 R-Phycoerythrin, C-phycoery- thrin (the letters refer to the source and properties of the biliprotein, see ref.14), Bile Pigments and the pigments derived from these yield haematinic acid, methyl ethylidene succinimide (39) and methyl vinylmaleimide (36) upon chromic acid oxidation ; since the phycoerythrobilins contain two carboxylic acid groups, the haematinic acid must be derived from two rings of the pigment. Chromic acid degradation of Gphycocyanin and the pigment from this chromoprotein gavel'b methyl ethylidene succinimide, methyl ethylmaleimide, and haematinic acid. Since C-phycocyanin on hydrolysis with alkali yields mesobiliverdin (20) (and caused Lemberg and Bader21 to assume that phyco- cyanobilin and mesobiliverdin were one and the same compound) it follows that phycocyanobilin has the -1Xa orientation pattern; this led to the pr0posall7~ of structure (18) for phycocyanobilin.It is worth mentioning that if a degradative procedure which can be applied to both the chromoprotein and the isolated pigment had been available in Lemberg's day then progress would have been much more rapid in the field of algal bile pigments and he would not have been satisfied with the mesobiliverdin proposal for phycocyanobilin. Having established the nature of the chromophores of the algal bile pigments, the chromic acid degradation method was also able to offer some information on the bonding of these to their apoproteins.Depending upon the conditions, chromic acid degradation of the native chromoproteins gave different results. For instance, C-phycocyanin at 100 "C is hydrolysed to give peptides which pass into solution, allowing isolation of methyl ethylmaleimide, haematinic acid, and methyl ethylidene succinimide; these same imides are obtained from free phyco- cyanobilin. When the oxidationiscarried out at 20 "C the protein is not hydrolysed and remains intact. Whereas the same quantity of methyl ethylmaleimide can be extracted, only half of the expected amount of haematinic acid and no methyl ethylidene succinimide are obtained. The missing imide (39) and the other half of the haematinic acid can be liberated by hydrolysis without further quantities of oxidant being added.14 The same result is also obtained with R-phycocyanin and, in principle, with R-and C-phycoerythrin. These experi- ments serve to illustrate that two of the rings of the bile pigment are linked to the protein.It is not known which functional groups in the protein are involved in these linkages, and this information can only come after the amino-acid sequences of these biliproteins have been determined. Which of the two inner pyrrole rings is bound to protein has been established for phycoerythrin.gbsl4 On chromate oxidation at 20 "C the 2,5-diformylpyrrole (37) is liberated, while haematinic acid (29) is liberated only upon hydrolysis of the protein residue.Since the 2,5-diformylpyrrole can only be formed from ring B of a,b-bilidienes, this ring must be free and ring c must be bound to the protein in phycoerythrin. 3 Synthesis of Bile Pigments Bile pigments can be prepared either from pyrrolic compounds directly or else by ring-opening of porphyrins. In this section we shall deal only with the former approach; the latter method is discussed in Section 4. Synthetic methods have been discussed in full elsewhere3 and so only a brief Hudson and Smith discussion will be given. However, the importance of total synthesis in the bile- pigment area should not be underestimated because it was on the basis of this, and comparison with authentic compounds, that the early workers in Munich were able to confirm structural proposals and thereafter to piece together the complex relationships between the various pyrrole-derived pigments.la All direct approaches to bile pigments proceed by condensation of two oxo-pyrromethenes, oxopyrromethanes, or of an oxopyrromethene and an oxo- pyrromethane.In a few cases, 5-bromopyrromethenes have been self-condensed, but the terminal bromo-functions are hydrolysed either during or shortly after the condensation, and the bile pigments are isolated with the terminal lactam rings present. Synthesis of Oxopyrromethenes and Oxopyrromethanes-A typical oxopyrro- methene is neoxanthobilirubic acid (24), and two slightly different ap- proaches31as are presented in Scheme 4. These employ classical pyrrole chemistry Me Et Me Me Et Br QCH, ii) Me@Qp"NH HN' ...(11' Br QCOzH H Br +H H Br' i vfit Me Et Me Et HO2C HQ C02Et QCHO NH HNH H Mc PH Mc PH Reagents : i, Br, ; ii ,HBr , 0; iii, HBr, oHcg; iv, NnOMc-H,O,lSO°C: H H v, hydrolysis Scheme 4 which should need no discussion. It is not essential to introduce the lactam function by hydrolysis of a 5-bromopyrromethene [e.g. (40)]. If required, the oxygen function can be carried through from a pyrrolinone; such an approach s1 (a)W. Siedel and H.Fischer, 2.physiol. Chew., 1933, 214, 146, 163; (b) W.Siedel, Z. physiol. Chew., 1935,231,181, 197. Bile Pigments is outlined in Scheme 5, which shows the synthesis of xanthobilirubic acid (27).The pyrrolinone is obtained by hydrogen peroxide oxidation of a 2-unsubstituted pyrrole. Bromination and condensation with the resulting (not isolated) bromo- methylpyrrolinone of a 2-unsubstituted pyrrole affords the required product. Me Et Me Et r Me Et 1 HQMe- a HMe H O H Me PMe Reagents: i,H,02,py;ii, Br2; iii, H Scheme 5 Yet another variation is possible. Scheme 6 shows the synthesis32 of neoxan- thobilirubic acid (24)and isoneoxanthobilirubic acid (23) by condensation of 4-ethyl-3-methyl-(41)or 3-ethyl-4-methyl-pyrrolin-2-one(42)with 2-formyl-3- methylpyrrole-4-propionicacid (43). Gossauer33 has recently reported an ingenious approach to an oxopyrro- methene (44) required for his synthesis of phycocyanobilin. The two rings (Scheme 7) were joined by way of a sulphur-extrusion reaction. Only one direct approach to oxopyrromethanes has been described; Johnson and co-workers found that peroxide oxidation of 5-unsubstituted pyrromethanes affords34 oxopyrromethanes (45)in moderate yield.The most widely used method for synthesis of oxopyrromethanes is reduction of the corresponding oxopyrro- methene, either with sodium amalgam or by hydrogenation over a de-activated palladium catalyst. Many bile pigments possess chiral centres and are thus optically active. A typical example is stercobilin (14). The building blocks for such bile pigments 3* H. Plieninger and U. Lerch, Annalen, 1966, 698, 196. 33 A. Gossauer and W. Hirsch, Tetrahedron Letters, 1973, 1451 ;Annalen, 1974, 1496.34 J. H. Atkinson, R. S. Atkinson, and A. W. Johnson, J. Chem. SOC.,1964,5999. Hudson and Smith EL Me Me Me Scheme 6 Me OHN p. S Br I CHC02CHsPh pMe (44) Scheme 7 must, of necessity, contain the chiral carbon atoms assembled in the correct stereochemical array. Plieninger and his co-workers have made critically important contributions in this area. Hydrogenation of an oxopyrromethane (45) at high temperature and pressure in the presence of Raney nickel furnishes the cis-reduced compound (46) (Scheme 8).32 On the other hand, if the reduction is carried out with sodium in liquid ammonia, the trans-isomer (47) is obtained. The trans-compound can also be obtained by epimerization of the cis-isomer with strong base, such as t-butoxide in t-butyl alcohol.Using these intermediates, Plieninger and co-workers35 have carried out a whole series of sophisticated synthesesof stercobilin and other optically active urobilinoids (see also ref. 3). .N H. Plieninger and J. Ruppert, Annalen, 1970,736,43; H. Plieninger, K. Ehl, and A. Tapia,ibid.,p. 62. Bile Pigments Hz -Ni Na-NH,J I Scheme 8 Pyrro1inones.-As mentioned earlier, pyrrolinones can be ob tained34 by oxidation of 2-unsubstituted pyrromethanes with hydrogen peroxide in pyridine (cf. ref. 36). Contemporary synthetic approaches to bile pigments have tended to rely more upon ring synthesis of pyrrolinones. Perhaps the most widely used method is that of Plieninger.37 Catalytic hydrogenation (Raney nickel) of cyano-hydrins (48) obtained from p-keto-esters gives a diastereoisomeric mixture of hydroxypyrrolidones (49), which can be dehydrated to give the required pyrro- linones (50).It is interesting that reduction of (50) with Raney nickel or sodium in liquid ammonia gives the cis-and trans-pyrrolidones (51) and (52) respectively (cf. oxopyrromethanes) and that (51) can be converted into (52) by treatment with t-butoxide.32 Pyrrolinones can also be obtained via an internal Emmons reaction.38 Condensations to give Bile Pigments.-Treatment of formyl neoxanthobilirubic acid (53) [obtained by the Gatterman reaction from neoxanthobilirubic acid (24)]with isoneoxanthobilirubic acid (23) gives39 mesobiliverdin (20). Reduction of the aldehyde group in (53) to give the corresponding hydroxymethyl derivative (54) followed by coupling with isoneoxanthobilirubic acid (23) similarly gives40 mesobilirubin;these facts firmly established the constitutions of biliverdin and bi1irubin.l" 36 H.Fischer, T. Joshioka, and P. Hartmann, Z. physiol. Chem., 1932,212, 146. 37 H. Plieninger and M. Decker, Annalen, 1956,598, 198; H. Plieninger, and J. Kurze, ibid., 1964, 690, 60. s8 G. Stork and R. Matthew, Chem. Comm., 1970,445. W. Siedel, 2.physiol. Chem., 1935,237, 8. 40 W. Siedel, Z. physiol. Chem., 1937,245,257. 378 Hudson and Smith OH II I HO RZ HO R2-P--C--C-R~ R1,--,4-H CN COtEt COzEt H R' Ra HH HH CH,NHCO,EtI mg + OHC \ Me'0Me', H PH (55) Na2S20,Bilirubin (13) f---Scheme 9 Bile Pigments Total syntheses of bilirubin (13) and biliverdin (8) were complicated by the presence of the vinyl groups in these molecules.Although vinyl-substituted pyrrolinones are now ac~essible,~1~42 Fischer's pioneering work made use of urethane protecting groups (obtained by Schmidt degradation of propionic side-chains). Thus (Scheme 9)43 condensation of the formyl-oxopyrromethene (55) with the Sunsubstituted oxopyrromethene (56) afforded the bilitriene (57). Hydrolysis and exhaustive methylation generated the vinyl groups in biliverdin (8). Reduction with hydrosulphite then gave bilirubin. Slight modifications of this concept have been used in the synthesis33 of phycocyano bilin dimethyl ester (58) using an ethylidene-subst itu ted oxopyrr o-methene, in the synthesis44 of mesobilirhodin dimethyl ester (59) from an (58) (59) oxopyrromethene and an oxopyrromethane, and in the synthesis of optically activea6 urobilin (15).In the last case, a derivative of neobilirubic acid (25) was resolved, and the (+ )-and (-)-enantiomers were condensed separately with racemic formyl isoneobilirubic acid (60). The two diastereomeric mixtures were crystallized three times to give two urobilin-IXa hydrochlorides ([a]D2* + 4500" and -4800" }, presumably the RR-and SS-enantiomers of (1 5). Et 2 Me Me PH (+) -Urobilin-XI11 01 salt Scheme 10 41 H. Plieninger and R. Steinstrasser, Annulen, 1969,723, 149. 48 H. Plieninger, K.-H.Hentschel, and R.-D. Kohler, Annulen, 1974, 1522. * H. Fischer and H. Plieninger, Z. physiol. Chem., 1942,274,23 1. 44 A. Gossauer and D. Miehe, Annalen, 1974,352. Hudson and Smith Treatment of oxopyrromethenes or oxopyrromethanes (Scheme 10) with a suitable one-carbon unit produces symmetrically substituted bile pigments. A useful route to relatively large quantities of a model bilitriene, aetiobiliverdin- Ivy, has been described.45 Heating of the readily available 5-bromo-5’-bromo- methylpyrromethene hydrobromide (61) in wet formic acid causes ‘tail to tail’ self-condensation and hydrolysis of the bromo-function to give aetiobiliverdin- Ivy (62). The product is contaminated with aetioporphyrin-I, produced by ‘head to tail’ self-condensation, and this can be removed by column chromato- graphy.Me 2 HCOOH A * Me Et Et Aetiobiliverdin -IVY (62) 4 Formation of Bile Pigments A. Mammalian Bile Pigments.-The formation of bilirubin from haemoglobin has been known for well over a century; e.g. in 1847 Virchow46 described the formation of ‘haematoidin’ crystals in blood extravasates, but it was not until 1923 that Fischer and Rei~~del~~ proved haematoidin to be identical with bili- rubin (13). After a considerable amount of further work, the final elegant proof that bilirubin is formed by breakdown of the blood pigments was provided, in 1950, by London et aZ.,48who administered [15N]glycine intravenously and orally to adults, and found that up to 90% of faecal bile pigments could be accounted for by destruction of erythrocytes.Most of the labelled erythrocytes were destroyed after about 100-140 days, indicating that mature circulating erythrocytes have an average life span of about 120 days. The classical theory of bile-pigment formation49 summarized in Scheme 11 proposes separation of haem (63) from the protein matrix, followed by loss of iron from ‘haematin’ (64)to give free protoporphyrin-IX (8), which then under- goes oxidative rupture and reduction to give bilirubin (13). Despite poor experi- mental support, this view was accepted for very many years, even though the 45 K. M. Smith, J.C.S. Perkin I, 1972, 1471. 46 R. Virchow, Arch. Path. Anat. u. Physiol. (Virchow’s), 1847, 1,379.H. Fischer and F. Reindel, 2.physiol. Chem., 1923,127,299. 48 I. M. London, R. West, D. Shemin, and D. Rittenberg, J. Biof. Chem., 1950,184,351. 49 M. Nencki and N. Sieber, Arch. Exp. Pharmakol., 1884, 18, 401; ibid., 1888, 24, 430; M. Nencki and J. Zaleski, 2.physiol. Chem,, 1900,30, 384; Ber., 1901,34,997. 381 Bile Pigments 0xidn: -globin -iron redn Haemoglobin -Haematin Protoporphyrin-1X Bilirubin (64) ( 8) (1 3) Scheme 11 direct conversion of protoporphyrin-IX into bilirubin had never been demon- strated. In 1951, London et al. did report50 that 14C-labelled protoporphyrin-IX gave rise to activity in faecal bile pigments several days after injection but these authors suggested that iron may have been introduced into the porphyrin before breakdown took place, and they did not claim this experiment to be proof of the classical theory.The intermediacy of haematin (64) in bile-pigment formation has been a subject of much debate; in cases of haematin jaundice (a condition where haematin accumulates in the plasma without an increase in the serum bilirubin level), S~hottrnuller~~ and Schumm52 concluded that haematin was not a normal precursor of bilirubin, and this view was reinforced53 by the observation that intravenously injected haematin does not cause an increase in the excretion of bilirubin. Haematin was therefore thought of as a side-product in haemoglobin breakdown. Lemberg54 injected mesohaematin (65) into rabbits and found that (63) M = Fe" ; K=V R Me (64)+ M = Fe1I1-0H ; R = V (65)+ M = Fe"'-OH ; R = Et (66) M = py -+ Fell +py ; R = V r\ (75) M = Fe"'-Cl ; R = V (76) M = Fel'l-CI ; R = H PH P" (77) M = Fell ; R=Et (78) M= Fe" ; R=H it caused bilirubinuria and urobilinuria but also that the bile did not contain pigments with ethyl side-chains.The high toxicity of mesohaematin was demon- strated and it was argued that haematin would be unlikely to be a normal biological precursor of bilirubin. 54 However, in more recent times the conversion of haematin into bilirubin has I. 'Haematins' exist as the dehydrated p-oxo-dimer 6o I. M. London, M. Yamasaki, and A. G. Sabella, Fed. Proc., 1951,10,217. I1H. Schottmuller, Miinch. Med. Wochschr., 1914, 61, 230.0. Schumm, Z.physio1. Chem., 1916,97,32. I3 K. Bingold, 2. Klin. Med., 1923, 97, 257; Folia. Haematol. (Leipzig), 1930, 42, 192; R. Duesberg, Arch. Exp. Pathol. Pharrnakol., 1934, 174, 305; Klin. Wochschr., 1938,17, 1353. I4Ref. 1 (b),p. 575. 3 82 Hudson and Smith been demonstrated both in vivo55-58and in ~itro.~~It was observed55 that 15N- labelled haematin is rapidly converted into faecal stercobilin (14) in dogs, and that when injected into rats with external bile fistulae, 14C-labelled haematin is converted into bilirubin with an efficiency between 50 and 70%,56 values which are comparable with those for haemoglobin itself.57 There are two modern theories of bile-pigment formation; one holds that the process involves non-enzymic coupled oxidation and hydrolysis whereas the second proposes an enzymic transformation. Coupled Oxidation of Iron Porphyrins.In 1930, Warburg and Negelein60 des- cribed the formation of a green compound on coupled oxidation of pyridine haemochrome (66) in presence of hydrazine, and the compound was called ‘green haemin’. When treated with methanolic hydrogen chloride, ‘green haemin’ gave an ester which was subsequently shown6l by Lemberg in 1935 to be biliverdin dimethyl ester ferrichloride (67), and only one isomer was isolated. Lemberg62 also obtained crystalline biliverdin dimethyl ester (32) by coupled oxidation of haemoglobin in presence of ascorbic acid; because only one isomer was obtained it was widely believed that this pathway must be operative in vivo.‘Green haemin’ was renamed pyridine verdohaemochrome by Lemberg, and a structure (68) was proposed for it.63 A compound with an absorption maximum at 639 nm was shown to be formed before verdohaemochrome,63 especially when the reaction was carried out with hydrogen peroxide and ascorbic acid. Removal of iron from this intermediate gave an oxyporphyrin; Lemberg favoured the 0 Fe PH 65 I. M. London, J. Biol. Chem., 1950, 184, 373. 6B A. L. Snyder and R. Schmid, J. Lab. and Clin. Medicine, 1965,65, 817. 6’ J. E. Ostrow, J. H. Handl, and R. Schmid, J. Clin. Investigation, 1962, 41, 1628. 68 H. BCnard, A. Gajas, M. Polonovski, and M. Tessier, Compt. rend. SOC. Biol.,1946, 140, 51 ;I. J. Pass, S. Schwarz, and C.J. Watson, J. Clin.Investigation, 1954,24,283. 69 J. E. Kench, C. Gardikas, and J. F. Wilkinson, Biochem. J., 1950, 47, 129; J. E. Kench, ibid., 1954, 56, 669. 60 0.Warburg and E. Negelein, Ber., 1930,63, 1816. 61 R.Lemberg, Biochem. J., 1935,29, 1322. 62 R. Lemberg, J. W. Legge, and W. H. Lockwood, Nature, 1938,142,148. 63 R. Lemberg, B. Cortis-Jones, and M. Norrie, Biochem.J., 1938,32, 171. 383 Bile Pigments keto-tautomeric form (69) for oxyporphyrins and this has since been confirmed,6* with the macrocycle being named ‘oxophlorin’. The 639 nm compound was therefore characterized as the iron(m) complex (70) of the oxophlorin, and was found to be transformed into the corresponding verdohaemochrome by molecular oxygen. Fischer and Lib0witzky6~ also obtained an oxophlorin (71) from the oxidation of iron(m) coproporphyrin-I.Treatment with oxygen in pyridine gave the corresponding verdohaemochrome, which was transformed into the biliverdin bv treatment with alkali and then acid in 73% yield from verdohaemin. When haemoglobin and ascorbic acid were exposed to air for 48 h, a com- pound with an absorption maximum at ca. 674 nm was obtained.62 Treatment with sodium dithionite shifted the maximum to 629 nm. The chromophore was liberated from the 619 nm substance by treatment with ammonium sulphate followed by ethyl acetate in acetic acid. The freed material contained some haem, but biliverdin hydrochloride was formed with 1 % hydrochloric acid. The substance with absorption maximum at 619 nm was termed ‘choleglobin’66 and was thought to be a bile pigment-ironcomplex; the674nm material appeared to be ‘oxycholeglobin’ because on evacuation and dialysis choleglobin was obtained.Choleglobin has never been obtained free from haemoglobin,66 the reaction being prevented from going to completion by denaturation which occurs during its course. Nevertheless, choleglobin has been postulated2 as an intermediate in bile-pigment formation as outlined in Scheme 12. The only difference between Schemes 11 and 12 is whether the protein and iron are removed before or after the oxidation, since it is now accepted that cholehaem is a mixture of compounds containing oxidized porphyrin rings. Lemberg et aL6’ have postulated a pathway eA A.H. Jackson, G. W. Kenner, and K. M. Smith, J. Amer. Chem. Sac., 1966, 88, 4539; J. Chem. Sac. (C), 1968, 302. 6s H. Fischer and H. Libowitzky, 2.physiol. Chem., 1938, 255, 209; H. Libowitzky, ihid., 1940,265,191. 66 R. Lemberg, J. W. Legge, and W. H. Lockwood,Biochem. J., 1941, 35, 328, 339, 363; R. Lemberg and J. W. Legge, ibid., 1941,35,353. 13’ E. C. Foulkes, R. Lemberg, and P. Purdom, Proc. Roy. Suc., 1951, B138,386. Hudson and Smith Oxidn. -Fe redn. Haemoglobin -Choleglobin -Biliverdin -Bilirubin Scheme 12 for the oxidative cleavage of the porphyrin nucleus, and this is shown in Scheme 13. The carbon atom which is lost during the course of the transformation is liberated in the form of carbon monoxide (vide~upra).~ 0 ;o OH -$YO03Fe -W Scheme 13 Protoporphyrin-IX does not undergo ring-opening upon coupled oxidation with ascorbic acid,59 and this seems to suggest some special role for the iron atom, apart from that already recognized for the transport of oxygen in haemo- globin. Bonnett and DimsdaleO* allowed a series of metallo-octaethylporphyrins to react with hydrogen peroxide in pyridine ;metallo-oxophlorins were only ob- tained in those cases (where the central metal ion was FeII, CoII, MnII, and MnlI1) in which the metal atoms possessed an easily accessible higher oxidation state.68 R.Bomett and M.J. Dimsdale,J.C.S. Perkin I, 1972,2540. 3 385 Bile Pigments For example, in the iron(I1) case, the product is an iron(Ir1) oxophlorin, and since the meso-position and the metal ions are oxidized with peroxide, a mechanism (Scheme 14) similar to that occurring in aromatic hydroxylation with Fenton's J reagent was proposed,s8 the great novelty being in the fact that the iron@) species is situated within the porphyrin nucleus.This mechanism also accounts for the fact that a reducing agent is required to be present during the oxidation (i.e. coupled oxidation) of iron(rn) porphyrins because the reaction as outlined (Scheme 14) is carried out on the iron@) porphyrin. Treatment of iron(m) oxophlorins with oxygen in pyridine affords high yields of verdohaemochrome, as mentioned earlier. However, treatment of other metallo-oxophlorins (central metal, e.g. MgII or %I1) with oxygen affords the corresponding metallo-dioxoporphodimethenes (72) via the anion (73) and 0-0' 0 (74) M = Mg,Zn Hudson and Smith radical (74)69 The tendency of iron(II1) oxophlorins to react with oxygen by addition across the carbonyl group (see later) rather than at the opposite methine carbon may be associated with the fact that removal of an electron from the iron(rrr) oxophlorin anion [c-(73)+(74)] might lead to the formation of an iron(Iv) oxophlorin rather than an iron(In) oxophlorin radical ;this possibility of electron loss from the metal rather than the ligand is supported by (i) the observation70 that electrochemical one-electron oxidation of iron(n1) porphyrins gives iron(1v) porphyrins and not the rr-cation radical of the metalloporphyrin as in most other cases, and (ii), the implication71 of iron(1v) porphyrins as intermediates in the functioning of haem systems such as peroxidase and catalase.It might be reasonable to expect an iron@) oxophlorin to react some- what differently with oxygen than does a magnesium or zinc oxophlorin radical. It is also fascinating to consider that the facility of haems to be transformed into iron(1rr) oxophlorins is dependent upon the ability of the metal to be readily converted from the +2 to the +3 oxidation stateG8 and that the further reaction of the oxophlorin with oxygen to give verdohaemochrome is linked with the ability of the metal atom to be transformed from the +3 into the +4 state. Combination of these factors with the facility of iron(@ porphyrins to bind oxy- gen reversibly without oxidation, and the co-ordination chemistry of high- and low-spin iron(I1) which enables triggering of the co-operative effect,72 makes iron a very special metal indeed and may serve to explain why Nature depends so greatly on a metal which is in itself so poisonous.Though kmberg claimedlb that oxidative rupture of haemin (75) led only to the -1Xa isomer of biliverdin, it has since been established in several ways (see p. 370) that a random mixture of bile pigments actually results from the coupled oxidation. Table 314 clearly shows that there is little directive influence by the Table 3 Percentages of biliverdin isomers produced by coupled oxidation with ascorbic acid andlor hydrazine of protohaemin or deuterohaemin14 Biliverdin Pro tohaemin Deuterohaemin isomer ascorbicacid % hydrazine % hydrazine % 01 33.0 & 5.1 26.5 & 3.0 14.8 k 2.1 r6 (8) Y 22.5 26.1 22.3 & 4.0 19.6 & 3.0 20.5 +_ 2.0 45.3 & 3.1 29.2 +, 2.0 8 (b) 21.9 2 2.0 23.5 & 6.0 10.7 & 1.0 vinyl groups in haemin (75) and that in deuterohaemin (76) the /3(8)-*isomer predominates.A considerable degree of specificity has been noted in the coupled oxidation 69 G. H. Barnett, B. Evans, and K. M. Smith, Tetrahedron, in press. 70 R. H. Felton, G. S. Owen, D. Dolphin, and J. Faier, J. Amer. Chem. SOC.,1971,93,6332. 71 D. Dolphin and R. H. Felton, Accounts Chem. Res., 1974,7,26. 72 M. F. Perutz, Nature, 1970,228, 726,734; New Scientist, 1971, 676; W.Bolton and M. F. Perutz, Nature, 1970,228, 551. Bile Pigments of haemoprotein~.~3 Myoglobin, when subjected to coupled oxidation, yields only the -1Xa isomer, and haemoglobin affords both the -1Xa and -Utgisomers with no trace of the y-or 8-isomers. Similar specificity was observed when the native haem was replaced with mesohaem (77) or deuterohaem (78).73 This appears to suggest that the specificity of the rupture is a property of the haem binding rather than of the haem itself, and this is further confirmed by the fact that the cleavage becomes completely non-specific if the protein is denatured before coupled oxidation.73 It is unlikely that the cleavage is made specific due to steric hindrance of the other rneso-positions by the large protein molecule because X-ray studies have shown74 that the a-carbon atom is in fact the most hindered, being situated at the bottom of the haem binding crevice.The Enzymic Oxidation of Iron Porphyrins. In 1963, Nakajima et aL7=reported the isolation of an oxygen-sensitive enzyme, ‘haem a-methenyl oxygenase’, from beef liver and guinea-pig liver. It was shown to catalyse the transformation of pyridine haemochrome (66) and haemoglobin-haptoglobin complex (hapto- globin is a haemoglobin-binding plasma protein) to a possible precursor of biliverdin, which was assigned structure (79) on inconclusive evidence.76 The V Me P” (79) sites of highest enzyme activity were the liver and kidney, but curiously, it was totally absent from bone marrow and spleen, the two principal sites of haemo-globin breakdown.Later, Levin77 conhned that liver homogenates contained a substance which was active in cleaving pyridine haemochrome (66) but the active substance was only slightly heat-labile, was stable to extremes of pH, and was of low molecular weight, suggesting that it was not an enzyme at all. The active compound was shown to be a reducing agent and not to have a catalytic role; it was therefore suggested that the rapid disappearance of pyridine haemochrome was due to a non-enzymic coupled oxidation with endogenous reducing agents, and the proposed structure (79) of the product was discounted.77 Murphy et obtained the same results as Levin and they also showed that extracts from red 75 P.O’Carra and E. Colleran, F.E.B.S. Letters, 1969,5,295;P. O’Carra, in ‘Porphyrins and Metalloporphyrins’, ed. K. M. Smith, Elsevier, Amsterdam, 1975,p. 123. 74 J. C. Kendrew, Science, 1963,139,1259;M.F.Perutz, Proc. Roy. Soc., 1969, B173,113. 75 H.Nakajima, T. Takemura, 0. Nakajima, and K. Yamaoka, J. Biol. Chem., 1963, 238, 3748. 76 H.Nakajima, J. Biol. Chem., 1963,238,3797. 77 E.Y. Levin, Biochemistry, 1966,5, 2845;Biochim. Biophys. Actu, 1967, 136, 155. R. F. Murphy, C. O’hEocha, and P. O’Carra, Biochem. J., l967,104,6C. Hudson and Smith algae contained a compound or compounds of low molecular weight which were active in cleaving porphyrins. The effect of the active factor was enhanced by NADPH; this was presumed to be due to regeneration of the active factor by reduction during the coupled oxidation.Wise and Drabkin79 have claimed that the haemophagous organ of dog placenta contains an enzyme system which catalyses the formation of biliverdin and carbon monoxide from haemin and haemoglobin. This system differed from that of Nakajima in two fundamental ways, (i) in substrate specificity, pyridine haemochrome being relatively ineffective, and (ii) in subcellular location, the activity being associated with the light mitochondria1 fraction. Kench et aLso noted that the splenic pulp of a haemolytic anaemia patient caused an enhancement in the production of bile-pigment precursors, normal splenic tissue being ineffective.Such enzymic activity was also observed in foetal liver, bone-marrow, and spleen, and the biliverdin-forming activity was found V Me ''OH V Me Fe, PO dlSO2 rcdn f------Bilirubin Scheme 15 7g C. D. Wise and D. L. Drabkin, Fed. Proc.., 1964,23,223; ibid., 1965,24,222. 8o J. E. Kench and S. N. Varma, S. African Med. J., 1962,36,794; J. E. Kench, F. E. duToit, and M. Green, S.African J. Lab.and Clin. Medicine, 1963,9,273. 389 Bile Pigments to be somewhat heat-labile, with bilepigment yields considerably lower than reported by Nakajima. Indeed, Kench81 and Schmid82 also were unable to substantiate the work of Nakajima. In 1970, Tenhunen et al.83published details of a microsomal enzyme system from liver which converts haem into bilirubin, and appears to be a mixed function oxidase.The system was called ‘microsomal haem oxygenase’ and requires molecular oxygen and NADPH as co-factors. When the enzymic reaction was carried out in an atmosphere of 1802, the resultant bilirubin (Scheme 15) possessed two 180 atoms and an additional l80atom appeared in the expelled carbon monoxide. When carried out in a medium of HPO, no label was in- corporated, suggesting that no hydrolytic stage occurred between haem and bile pigment. The enzyme system is inhibited by carbon monoxide, and cytochrome P450 was implicated. Scheme 15 proposes the intermediacy of an iron(m) oxophlorin (80); this possibility has been investigated by Kondo et al.84Tritiated a-oxymesoferrihaem (81) was injected into rats with bile fistulae and was found to be extensively converted into bile pigments, the expected mesobilirubin being identified.On the other hand, p-oxymesoferrihaem (82) was found to be a very poor precursor of bile pigments; the metal-free oxophlorins were also converted poorly into bile pigments. The authors concluded by analogy that a-oxyprotoferrihaem (80) is an intermediate in haemoglobin breakdown and that it undergoes further oxidation to bile pigment under the catalysis of an enzyme of definite specificity. This oxidation was postulated (Scheme 16) to proceed by addition of molecular oxygen across the oxophlorin carbonyl group followed by cheletropic fragmenta- tion to give carbon monoxide and the biliverdin iron complex.Finally, the iron is removed to give biliverdin; the overall concepts in this scheme are in accord with the findings of Tenhunen et ~1.~3 A NADPH-dependent microsomal haem cleavage system which does not seem to be associated with cytochrome P 450 has been reported by Yoshida et aZ.85 Haem-cleavage activity had previously been reported86 in chicken macro- phages and this was similarly thought to be unassociated with P450. Treatment of rats with CoCl2 causes8‘ stimulation of haem oxidation by hepatic micro- somes; the rate of haem oxidation, however, has been shown to be unrelated to the microsomal content of P 450, and the conclusion has therefore been drawn that the microsomal enzyme system is not the same as that which metabolizes drugs, and therefore that cytochrome P 450 is not essential for haem oxidation.B. Algal Bile Pigments.-By comparison with the decades of experimentation which have been devoted to the investigation of catabolism of the blood pig- 81 J. E. Kench, M. Green, and M. Hines, s. African J. Lab.and Clin. Medicine, 1964, 10, 33. R. Schmid, S. African J. Lab. and Clin. Medicine, 1963, 9, 276. E3 R. Tenhunen, H. Marver, N. R. Pimstone, W. F. Trager, D. Y. Cooper, and R. Schmid. Biochemistry, 1972, 11, 1716. 84 T. Kondo, D. C. Nicholson, A. H. Jackson, and G. W. KeMer, Biochem. J., 1971,121,601. e6 T. Yoshida, S. Takahashi, and G. Kikuchi, J. Biochem. (Japan), 1974,75, 1187. 86 A. W. Nichol, Biochim. Biophys. Acra, 1970,222,28. 87 M. D. Maines and A.Kappas, Proc. Nat. Acad. Sci. U.S.A., 1974,71,4293. Hudson and Smith Et R1 Me PH (81) R'=OH,R2= H (82) R~=H,R~=oH o=oOH V I Me Scheme 16 ments, relatively little work has been carried out on the origin of the algal bile pigments. However, it is now clear that algal pigments are derived by oxidative rupture of porphyrin precursors, as might be expected from the 'type-IX' orientation of their substituents. Bile Pigments Illuminated cells of Cymidium caldarium incorporated88 8-aminolaevulinic acid (83) (ALA) into phycocyanin, but in the dark, incubation of the alga with ALA P IW ALA resulted in excretion of porphobilinogen (84) (PBG), seven porphyrins, and a blue pigment, shown89 to be the chromophore of phycocyanin. Consideration of the specific activities of products isolated after feeding isotopically labelled ALA allowed the conclusion*9 that the bile pigment was formed from ALA via PBG and a porphyrin, as in mammalian haem catabolism. However, the probable intermediates, uroporphyrinogen-111, coproporphyrinogen-111, proto- porphyrin-IX, and possibly haem, have not been isolated89 from C.caldarium cells actively synthesizing phycocyanin. During the synthesis of phycocyanobilin (the prosthetic group of phycocyanin), wild-type cells of C. caZdarium produced equimolar amounts of carbon monoxide and phycocyan~bilin,~~ suggesting that the CO and bile pigments are derived from the carbon skeleton of protopor- phyrin-IX. These results were interpreted such that haem, [the iron(11) complex of protoporphyrin-IX], or a haemoprotein, should be the substrates for ring- opening to give the algal bile pigments. In our view, it is at least as likely that the magnesium complex of protopor- phyrin-IX is the metalloporphyrin intermediate in algal bile-pigment formation. Magnesium@) protoporphyrin-IX (85) undergoes photochemical oxidationg1 to give compounds which have been formulatedg2 as formylbiliverdins in a model series.The argument against this is that adapted cells of the blue-green alga Tolypothrix tenuis synthesize93 phycocyanobilin in the dark, and therefore that, at least in this species, light is not an essential requirement. However, in the next section of this review, the novel ring-opening of metallochlorins to give94 dihydrobiliverdin derivatives is discussed, and in this reaction the oxidative ring-opening occurs at a meso-position adjacent to the reduced ring.Scheme 17 outlines a possible route to phycocyanobilin from magnesium protoporphyrin- IX (85) in which the 4-vinyl is first reduced to ethyl to give (86); after a proto- R. F. Troxler and L. Bogorad, Plant Physiol., 1966, 41, 491 ;also, in ‘Biochemistry of Chloroplasts’ Vol. 2, ed. T. W. Goodwin, Academic Press, London, 1967, p. 421. R. F. Troxler and R. Lester, Biochemistry, 1967,6,3840. R. F. Troxler, Biochemistry, 1972’11,4235. J. Barrett, Nature, 1967, 215, 733. ga J.-H. Fuhrhop and D. Mauzerall, Photochem. and Photobiol., 1971, 13, 453. g3 Y.Fujita and A.Hattori, ‘Studies on Microalgae and Photosynthetic Bacteria’, University of Tokyo Press, Tokyo,1963, pp.43 1440. O4 J. A. S. Cavaleiro and K. M. Smith,J.C.S. Perkin I, 1973,2 149. Hudson and Smith I= Me I= Me Me Me PH PH tropic shift to furnish the ‘chlorin’ (87), ring-opening could give (88), a possible precursor of the phycobilins. In normal chlorophyll-a biosynthesis, the 4-vinyl group present in magnesium protoporphyrin-IX is reduced to ethyl at a point Biie Pigments as yet uncertain, and the prototropic shift to give (87) from (86) is not unknown because bacteriochlorophyll-b has recently been assigned95 the structure (89), the ethylidene group in ring B presumably being derived from an ethyl group via a proton shift.C. Other Approaches to Bile Pigments from Porphyrins.-In 1938, Fischer and Bock showed96 that photo-oxidation of the sodium complex of aetioporphyrin-I afforded, amongst other things, a bile pigment. Some time later, Barrettgl showed that photo-oxidation of magnesium protoporphyrin-IX gave ring- opened compounds which were identified as biiiviolins. The product from the same reaction in the magnesium octaethylporphyrin series was differently formulated as 1’-formylbiliverdin(90) and a mechanism was suggested92 (Scheme Et Et Scheme 18 18). In a similar type of reaction zinc chlorins give a mixture of the two possible formyl-dihydro-octaethylbiliverdins,in which the ring has been ruptured adjacent to the reduced sub-unit.97 sf, H.Scheer, W. A. Svec, B.T.Cope, M. H. Studier, R. G. Scott. and J. J. Katz. J. Amer. Chem. SOC.,1974, 96, 3714. H. Fischer and H. Bock, 2. physiol. Chenz., 1938, 251, 1. O7 P. K. W. Wasser and J.-H. Fuhrhop, Ann. New York Acad. Sci., 1973, 206, 533; J.-H. Fuhrhop, P. K. W. Wasser, J. Subramanian, and U. Schrader, Annalen, 1974, 1450. Hudson and Smith Treatment of zinc(I1) or thallium(rrr) chlorins with thallium(xxx) trifluoroacetate affordsQ4 the corresponding meso-trifluoroacetoxy-derivative (91) (Scheme 19), Et Et Et Et Et EL-TTFA I Et Et Et H M=Zn or T1-X A1203 Et Et Et Et Et f-Et Et (93) (92) (probably reacts as the radical) Scheme 19 which can be hydrolysed to give the oxychlorin (92) (cf.the oxophlorin inter- mediate in haem catabolism); this compound (92) reacts spontaneously with oxygen to give the dihydro-biliverdin (93). The most recent example of ring-rupture to be reportedg8 is the photo- oxidation of zinc oxophlorins (94), which affords biliverdins via zinc oxapor- phyrins (95) (Scheme 20). 5 Common Reactions of Bile Pigments In this section we shall concentrate only upon those reactions which accomplish some kind of modification of the carbon skeleton of the bile pigment, thereby excluding discussion of reactions which might tend to be more characteristic of side-chains than of the pigment itself. On the other hand, we are also omitting S.Besecke and J.-H. Fuhrhop, Angew. Chem., 1974,86,125. Bile Pigments Et + Et Et Et Et Et (94) Et Et Scheme 20 discussion of hydrogenation of the chromophoric group and complexation with metal ions, both of which are properties of the chromophore, but in which no change in the skeletal composition takes place.A. Gmelin Reaction.-The first reaction of bile pigments to be investigated was the treatment of bile with nitric acid contaminated with nitrous acid, carried out by Tiedemann and Gmeli~~,~~ in 1826. This caused a series of colour changes from the orange-yellow of bilirubin, through green, blue, violet, red, and yellow. A great deal of work on this colour test for bile pigments was carried out by Siede1,100 using both natural and more symmetrical model bile pigments. It is generally accepted that the first stage in the reaction is oxidation of bilirubin (yellow-orange) to give biliverdin (green); after this the absorption maximum gradually moves to shorter wavelength owing to interruption of the chromo- phoric pathway, firstly at a terminal inter-pyrrolic position, then at the other, and finally at the central methine carbon.lb Though the compounds presumed to give rise to these colours have in most cases been given names, their structural assignments are still confused.*9 F. Tiedemann and L. Gmelin, ‘Die Verdauung nach Versuchen’, 1st edn., Karl Groos, Heidelberg and Leipzig, 1826. looW. Siedel and W. Frowis, Angew. Chem., 1939, 52, 38; Z. physiol. Chem., 1941, 267, 37; W. Siedel and E. Grams, ibid., 1941,267,49; W. Siedel, Angew.Chertz., 1943,56, 169. Hudson and Smith B. Diazo Reaction.-The diazo reaction was discovered in 1883 by Ehrlich, who treated a solution of bilirubin in chloroform with diazobenzenesulphonic acid in acidified alcohol; a red colour was produced101 which changed through violet to blue when treated with concentrated hydrochloric acid. [For a thorough treatment of this reaction see ref. l(d), p. 59.1 This colour test is positive with all bile pigments of the a,c-bilidiene type (6) which have a b-CH2 group, but is negative, for example, with biliverdins. With unsymmetrically substituted bilirubins, the products are two azo-compounds (96) and (97) (Scheme 21). The 0 Me H+PhN* 4:,t \ N=NN N PH ‘Ph Ph/ Scheme 21 mechanism of the diazo reaction has recently been investigated102 and the central (b) carbon atom was shown to be liberated as formaldehyde.C. Isomerizationof Bilirubins in Acid Solution.-This reaction has a bearing also on the diazo reaction. Bilirubin has been shown103J04 to undergo reversible acid-catalysed cleavage about the central methylene bridge, which leads to isomeric scrambling (Scheme 22) with production of bilirubins-IXa (1 3), -XIIIa (98), and -1IIa (99) in an equilibrium mixture. Under alkaline conditions, bilirubin does not isomerize. These findings have shown that reaction conditions employed using bilirubins must be carefully scrutinized or else one may be dealing with an isomeric mixture of pigments; e.g., dehydrogenation of bilirubin with 1,4-benzoquinone in DMSO-acetic acid, or with ferric chloride in acetic acid, affords103 a mixture of biliverdins.D. Photo-oxidation of Bile Pigments.-Interest in the photo-oxidation of bilirubin has been stimulated by the observation105 that hyperbilirubinemia in prematurely born infants can be treated by phototherapy, using sunlight or artificial visible light.lo6 The process involves a self-sensitized reaction involving Iol P. Ehrlich, Centr. Klin. Med., 1883,4, 721. lo2 D. W. Hutchinson, B. Johnson, and A. J. Knell, Biochem. J., 1972, 127,907. lo3R. Bonnett and A. F. McDonagh,J.C.S. Chem. Comm.,1970,238. Io4 A. F. McDonagh and F. Assisi, J.C.S. Chem. Comm.,1972, 117; Biochem. J., 1972, 129, 797. Io6 R. J. Cremer, P. W.Perryman, and D. H. Richards, Lancet, 1958,1, 1094. looJ. Lucey, M. Ferreiro, and J. Hewitt, Pediatrics, 1968,41, 1047. 397 Bile Pigments 398 Hudson and Smith singlet oxygen.107 The products isolated108 from methanol solutions are methyl vinylmaleimide (36), haematinic acid (29), and various propentdyopent adducts [e.g. (loo)]. Short irradiation times result in the production of biliverdin, but it has been shown109 that this pigment inhibits the sensitized and unsensitized photo-oxidation of bilirubin, and therefore biliverdin is probably not an inter- mediate in the main pathway of bilirubin photo-oxidation in vitro. E. Cyclization of Bile Pigments to give Macrocyc1es.-Treatment of the bis-iminoether (101) from bilirubin diethyl ester with cobalt or nickel salts affords the cyclized tetradehydrocorrin salts (102).110 The zinc biliverdin v OEt Me OEt OEt Me V clod- Me Me (101) (102) M = Co or Ni obtainedg8 by photo-oxidation of zinc oxophlorins (94) can also be re-cyclized to the corresponding zinc oxaporphyrin (95) (Scheme 20).lo' A. F. McDonagh, Biochem. Biophys. Res. Comm., 1971,44, 1306. lo*For reviews, see D. J. Ostrow, Semin. Hematol., 1972,9, 113; D. A. Lightner, Photochem. andPhotobiol., 1974,19,457. lo9 A. F. McDonagh, Biochem. Biophys. Res. Comni., 1972, 48, 408; D. A. Lightner, D. C. Crandall, S. Gertler, and Q. B. Quistad, F.E.B.S. Letters, 1973, 30, 309. ll0 H. H. Inhoffen, H. Maschler, and A. Gossauer, Annalen, 1973,141.

ISSN:0306-0012    

DOI:10.1039/CS9750400363

出版商:RSC    

年代:1975

数据来源: RSC

摘要:

TATE AND LYLE LECTURE* Spin-Lattice Relaxation :A Fourth Dimension for Proton N.M.R. Spectroscopy By L. D. Hall DEPARTMENT OF CHEMISTRY, UNIVERSITY OF BRITISH COLUMBIA, VANCOUVER, BRITISH COLUMBIA, V6T 1W5 CANADA 1 Introduction The conventional procedure using high-resolution n.m.r. spectroscopy as a tool for studying the structure and stereochemistry of organic molecules in solution is well known: three sets of n.m.r. parameters, the chemical shifts, coupling constants, and integrated areas, are derived from the measured spectrum and these parameters are then used, either singly, or in some suitable combination, as the basis for the structural assignment. Powerful as it is, this approach neglects the fact that each spectrum contains, implicitly at least, two further sets of mag- netic resonance parameters; these are the spin-lattice relaxation times (TI-values) and the spin-spin relaxation times.Recent instrumental developments have made possible the routine measurement of spin-lattice relaxation times, and it is now appropriate for chemists to begin to evaluate the diagnostic potential of what constitutes, for the organic chemist at least, a new class of n.m.r. parameters. The phenomenon of spin-lattice relaxation has long attracted the attention of physicists and of some chemical physicists. Largely as a result of their efforts, there is an extensive literature both of methods for measuring TI-values and of the theory that enables these parameters to be related to a wide variety of mole- cular properties and, importantly, molecular motions.Until recently most practising chemists, even those with an active interest in n.m.r. spectroscopy, have largely neglected the phenomenon of spin-lattice relaxation.? There have been many reasons for this neglect and of these the most cogent is that unti1 very recently, the methods developed by physicists for measuring Tl-values have been incompatible with studies of the complex molecular species which are of interest to the practising chemist. However, with the development of Fourier transform (FT)techniques it has become possible to measure, on a rout-* Delivered in Birmingham, April 1974. t The most notable exceptions to this are the nuclear Overhauser enhancement experiments, the uses of which were first demonstrated by the pioneering studies of Anet and Bourne.' For an excellent summary of the literature and of the full theory of this area the reader is referred to the monograph by Noggle and Schinner.' F.A. L. Anet and A. J. R.Bourne, J. Amer. Chem. Soc., 1965,87,5250. J. H. Noggle and R. E. Schirmer, 'The Nuclear Overhauser Effect. Chemical Applica- tions', Academic Press, New York, 1971. Spin-Lattice Relaxation: A Fourth Dimension for Proton N.M.R. Spectroscopy ine basis, the TI-value of any resonance that can be clearly resolved in an n.m.r. spectrum, and hence to study systems which are chemically quite complex. It is interesting to note that the majority of the laboratories first equipped with FT n.m.r.spectrometers appear to have directed most attention towards studies of 13C Tl-values;3 relaxation studies of other magnetic nuclides, particularly of protons, have been somewhat neglected. In practice there are some complications associated with proton TI-measurements which do not pertain to 13C studies; however, these can be overcome. This review is principally related to proton studies that have been performed in the author’s laboratory. Following a brief introduction to the phenomenon of spin-lattice relaxation and to the pulse n.m.r. methods which are used to measure spin-lattice relaxation times, there is a discussion of experiments which delineate their configurational dependences. This is followed by descrip- tions of several methods which allow extensive simplifications to be made to certain complex lH n.m.r.spectra. The review ends with examples which show that it is a simple matter to manipulate proton spin-lattice relaxation times chemically and thereby to extend the range of molecules which can be studied by FT relaxation techniques. 2 Spin-Lattice Relaxation What, then, is the phenomenon of spin-lattice relaxation? Basically, we are con- cerned with the way in which, and the rate at which, magnetic energy is transferred between the magnetic nuclei under study (the ‘spins’) and their surrounding environment (the ‘lattice’). There are a number of distinct mechanisms whereby this energy transfer can be effected, and the names ascribed to these are listed in Table 1.Fortunately we need not be concerned here with details of the important Table 1 Mechanisms of spin-lattice relaxation 1 Dipole-Dipole (a) Intramolecular (b) Intermolecular (c) Paramagnetic 2 Spin Rotation 3 Quadrupolar 4 Scalar Coupling 5 Chemical Shift Anisotropy differences between these mechanisms; all that need be noted is that in each case a rapidly fluctuating magnetic field, generated and located in the lattice, interacts with the rapidly precessing nuclei of interest and thereby mediates in the energy- transfer process. In principle, and often in practice, several of these different G. C. Levy and G. L. Nelson, ‘Carbon-13 Nuclear Magnetic Resonance for Organic Chemists’, Wiley-Interscience, New York, 1972.relaxation mechanisms operate simultaneously, and in these circumstances the spin-lattice relaxation time measured experimentally (7'1)-is a composite value and contains contributions from each of the several mechanisms (A, B, etc.). Thus (;Jexp= (;)A+ (;)E$+ ... . If this were always the case then it is quite probable that this entire phenomenon would have little, or no, relevance to organic chemists, since although it is a trivial matter to obtain an experimental value for a spin-lattice relaxation time, it is often a more difficult task to derive from that value the individual contribu- tions from several different relaxation mechanisms. Fortunately, it is often possible to conduct experiments in which one of these mechanisms, the intra- molecular dipole-dipole mechanism, dominates the relaxation, and under opti- mum conditions is the only operative mechanism.In practice, this requires that one be interested in molecules which are moving more or less isc tropically in the solution. Furthermore, these molecules should be studied as a dilute solution (ca. 0.1 mol 1-1) in a magnetically inert solvent; both of these requirements minimize contributions from intermolecular effects, and both are attainable if one has access to a FT spectrometer. The requirement of a magnetically inert solvent merely means that the solvent should contain neither fluorine nor proton substituents, and in practice it is usual to approximate to this requirement by using a deuteriated substance, which has the added convenience of providing a heteronuclear-lock signal for the spectrometer.The importance of making experimental measurements under these conditions transcends mere convenience, and this becomes apparent when one studies the mathematical form of the intramolecular dipole-dipole mechanism, which is given below for the contribution that one donor nucleide D makes to the spin- lattice relaxation of a second, receptor nucleus R. In this equation, and YR are the magnetogyric ratios of the two nuclei involved; r, is the internuclear distance between D and R;and rc is a motional correlation time (which reflects the memory that the molecule to which the nuclei D and R are attached has of its motion in solution). The first important point is that if one is considering the relaxation of a proton, R,in an organic molecule, the most probable source of relaxation is another proton; this follows from the fact that only protons have significantly high values of y.The second point to note is that the contribution which the donor proton, D, can make to the relaxa- tion of R falls off as the inverse sixth power of the internuclear separation of D and R;hence one can anticipate the possibility that proton spin-lattice relaxation times will show pronounced configurational dependencies. Of course, in most Spin-Lattice Relaxation: A Fourth Dimension for Proton N.M.R. Spectroscopy ‘real’ organic molecules, each individual proton will be relaxed by interactions with several other protons (D-1, 0-2, .. .) in the same molecule, and hence its total relaxation time (a) will have the form given in equation (2) and the magnitude of each of these contributions will depend on the relative magnitudes of the individual internuclear distances. This leads to the third point: if it is possible to measure quantitatively the contributions which any donor proton, D-1, makes to the relaxation of two other receptor protons, R-1 and R-2, then this should provide a direct inter-comparison of the internuclear separations D-1 to R-1 and D-1 to R-2. Thus,(k)D-1-R-1 (rD-1 +R-2)6 (3)(i)D-l+R-2 This potential can be realized by judiciously conducted nuclear Overhauser experiments.4 In principle it should also be possible to place these distances on an absolute basis by substituting into equation (1) a value for the correlation time, TC;however, this presents a number of major problems because it is not easy to obtain values for correlation times which are sufficiently accurate to justify subsequent calculations of the type envisaged here, and alternative ap- proaches may prove to be more useful.In brief then, under suitable conditions proton spin-lattice relaxation involves through-space interactions between individual protons in the same molecule and, since the efficiency of these interactions falls off rapidly with distance, each proton receives most of its relaxation from its nearer-neighbour protons. 3 Measurement of Spin-Lattice Relaxation Times The basic experimental technology involves pulse-n.m.r.methods, and the simp- lest conceptual model upon which a discussion can be based is the ‘rotating- reference frame’ model. Consider an ensemble of magnetically equivalent nuclei (spins) subject to the influence of some external magnetic field (Ha).At thermal equilibrium between the spins and the lattice, this ensemble will have a net, macroscopic magnetic moment which will, in the rotating reference frame, be directed along the z-axis, which is also the direction of HO(see Figure 1). In Figure 1this magnetic moment is represented by a vector ‘arrow’, whose length is intended to indicate the total amount of magnetization present; when this vector lies along the z-axis, no signal is detected by the spectrometer, which is designed to respond only to that com- R.Freeman, H.D. W.Hill, B. L.Tomlinson, and L. D. Hall, J. Gem. Phys., 1974, 61, 4466. Hall Equilibrium magnetization DelayI Figure 1 The rotating reference frame model for the measurement of a spin-lattice relaxa- tion time. ponent of the magnetization which lies in the x,y-plane. Thus, to assay the amount of magnetization along the z-axis at any particular time it is necessary to tip the magnetization vector through 90" into the x,y-plane. This is easily accomp- lished by the simple expedient of applying a suitable amount of radiofrequency power at the appropriate frequency, in the form of a short pulse (a 9Oo-pulse), and, providing this is done sufficiently rapidly compared with the rate of change of magnetization along the z-axis (see below), this introduces no systematic error.It is important to note that if twice that amount of power is applied say by doubling the length of the pulse, then the original magnetization vector will be tipped through 180" and will lie along the -z-axis. In the rotating reference frame, spin-lattice relaxation refers to the rate at which magnetic energy is transferred from spins which are directed along the z-axis. Initially, in Figure 1 the spin system is at thermal equilibrium with its lattice, and with its macroscopic magnetic moment directed along the z-axis. This equilibrium is then destroyed by applying a 18O0-pulse of power which tips the magnetization rapidly into the -z-direction.Spin-lattice relaxation then restores the system to thermal equilibrium; this is accompanied initially by a decay in the magnetization intensity along the -z-direction and subsequently by a recovery in the +z-direction. The amount of magnetization present at any particular instant can be assayed by applying a 90°-pulse and so tipping the Spin-Lattice Relaxation: A Fourth Dimension for Proton N.M.R. Spectroscopy residual component up, or down, into the x,y-plane, where it induces a signal into the receiver of the spectrometer. In practice, then, a 18O"-pulse is used to destroy the thermal equilibrium, a known amount of time is allowed to elapse (the delay time, t) and the residual magnetization after that period is assayed by a 90"-pulse.This gives the first point on the decay-recovery curve. The system is then left for 5 or more TI-periods to recover its equilibrium level of magnetiza- tion, and then a second 180"-90" pulse-sequence is applied, this time with a somewhat longer value for t. This gives the second point on the decay-recovery curve, and so on; in normal practice a minimum of at least ten such points must be determined. The spin-lattice relaxation time is the time constant of the curve drawn through the individual points [Figure 2(a)]; this value is more conven- iently obtained by making a semi-logarithmic plot of loge (peak height) vs. t [Figure 2(b)]. The significance of the fact that the decay-recovery curve has zero intensity for a particular value of t is discussed below.Pulse spectrometers capable of performing the above manipulation have been available for a number of years.5 Their principle disadvantage is that they lack frequency selectivity and hence their use is mainly confined to molecules having a single set of equivalent nuclei and from which an average 7'1-value is obtained.6 A marked improvement in frequency selectivity has followed from the develop- ment of the 'audiofrequency-pulse' technique by Freeman and Wittekoek.' We constructed a modified version of that spectrometer8 and used it in our own first measurements; however, this technique also suffered from a variety of limitations, not the least being the large amount of time needed to measure the 7'1-values of a complex proton n.m.r.spectrum, transition by transition. Fortunately the adventg of FT n.m.r. spectrometers has made possible a full realization of proton TI-measurements for even complex organic molecules. The required 180"- and 90O-pulses are now applied simultaneously to all of the proton resonances and hence the spin-lattice relaxation times of each proton resonance, indeed of each individual transition in the spectrum, are obtained simultaneously. The output of the spectrometer following the application of the WO-pulse is referred to as the 'free-induction decay' signal (F.I.D.) and this is stored, following digitization, in the memory of a small computer. The F.I.D. consists of a series of overlapping sine-waves and contains all the information concerning the n.m.r.spectrum; this is converted to the more-familiar, frequency- domain, n.m.r. spectrum by the mathematical manipulation known as Fourier transformation. T. C. Farrar and E. D. Becker, 'Pulse and Fourier Transform N.M.R.', Academic Press, New York, 1971. (I T. L. Pendred, A. M. Pritchard, and R. E. Richards, J. Chem. SOC.(A), 1966, 1009. R. Freeman and S. Wittekoek, J. Mugn. Resonance 1969, 1, 238; R. Freeman, S. Witte-koek, and R. R. Ernst, J. Chem. Phys., 1970,52, 1529. R. Burton, C. W. M. Grant, and L. D. Hall, Canad. J. Chem., 1972,50,497. R. R. Ernst and W. A. Anderson, Rev. Sci.Znstr., 1966, 37, 93; R. R. Ernst, Adv. Magn.Resonance, 1966,2, 1 ;R. Freeman and H. D. W. Hill, 'Introduction to Fourier Transform N.M.R.', Varian Associates, Palo Alto, 1970.2.80 2.40 2.00 1.60 Log eY I .20 0.80 0.40 0.00 ' -0.40. I I I I I I 1 0.00 0.80 1.60 2.40 3.20 4-00 4-80 5-60 11s Figure 2 Representation of the data obtained for a two-pulse determination of a spin- lattice relaxation time. (a) The recovery of the magnetization from its inverted value (-Mo) back to thermal equilibrium (+Mo) is shown. Note that for short values of t the peak which would be observed in the n.m.r.spectrum would be negative-going, reflecting the fact that the magnetization was still along the -z-axis when the 9W-pulse was applied. (b) A typical plot of loge (peak height) versus t is shown. Spin-Lattice Relaxation: A Feurth Dimension for Proton N.M.R. Spectroscopy 95.4 1 ISBr Br c1 CI c1 ‘c=c /Br >c=c,/H ‘c=c/ \c=c H /\ /\ /H ‘ClH H Br H H H 58.5 58.5 95.4 78.4 78.4 115 28.8 EtO,C /CO,Et Et0,C /H 26.9 H\ /H 86.4 \c=c \c=c/\ /H H H ‘C0,Et H/“=“\OCOMe 12.5 12.5 28.8 35.8 Figure 3 Valuesfsof the proton spin-lattice relaxation times for alkenes.Tle data for the halgeno-ethanes were determined with a selective, audio-pulse spectrometer10 using ca. 0.5 mol I-’ solutions. The vinyl acetate data were obtained by conventional FT methods, on a ca. 0.1 mol 1-1 solution. UnpublishedlO experiments involving measurements of cis- and trans-alkenes with our home-built, audio-pulse spectrometer are summarized in Figure 3. The differential between the TI-values for the individual pairs of isomers is clear- cut.Although the frequency selectivity of that instrument did not allow a com-plete study to be made, it was possible to measure1OJl the TI-values of some of the protons of 3,4,6-t ri- O-ace t yl- 1-O-benzoyl-2-chloro-2-deoxy-~-~-glucopy-ranose (4), and the study was subsequently completed when the FT method became available in our laboratory. 4 Stereospecific Dependence of Spin-Lattice Relaxation Times The first observations of a pronounced stereospecific dependence for the proton TI-values of a sugar derivative were made independently in 1972 by Professor Richards’ group at Oxford and in the author’s own laboratory in Vancouver. In Oxford,12 a study of ~-glucose-6-phosphate revealed a marked dependence for the TI-values of the anomeric protons, and a similar study of a series of hexose and pentose sugars in aqueous solution in Vancouver gave the data13 shown in formulae (1)-(3).The data for D-glucose (1) and D-galactose (2) immediately show that H-4 has little or no influence on the relaxation of the anomeric proton. The difference between the data for D-glucose and D-mannose (3) shows that H-2 makes some contribution to the relaxation of H-1, but forces one to conclude that the dominant source of the TI-differential between the axial and equatorial anomeric protons must be H-3 and H-5, the implication here being that H-laz lo C. W. M. Grant, Ph.D. Thesis, Department of Chemistry, University of British Columbia, March, 1972. l1 C. W. M. Grant, L. D. Hall, and C.M. Preston, J. Amer. Chem. SOC.,1973,95, 1972. I* R. Dwek, ‘Nuclear Magnetic Resonance in Biochemistry-Applications to EnzymeSystems’, Ciarendon Press, Oxford, 1973. l3 L. D. Hall and C. M. Preston, J.C.S. Chem. Comm., 1972, 1319; Carbohydrate Res., 1974, 37, 267. 408 Hall is nearer to H-3u2and H-L than is H-1 eq, and hence H-luzis relaxed faster by those two axial protons than is H-1 eq. HO Unfortunately the complexity of the proton spectra of these sugars precluded any general identification of other proton resonances, and hence further studies of these interesting stereospecific dependencies required selection of derivatives having a greater dispersion in their proton resonance spectra. The proton TI-data obtainedll from a set of 3,4,6-tri-O-acetyl-l-O-benzoyl-2-deoxy-2-halogeno-D-hexopyranose derivatives are shown in formulae (4)-(8).Com-parison of data obtained for (4) and (9,and for (6) and (7) immediately demon- strated the same pronounced differential between axial and equatorial anomeric protons that had been observed previously. Furthermore, the reciprocal effect of the contribution of H-1 to the relaxation of H-3 can also be seen. Proton H-5, 409 Spin-Lattice Relaxation: A Fourth Dimension for Proton N.M.R. Spectroscopy however, appears not to show this reciprocity; this is because its relaxation is now dominated by contributions from the C-6 protons. AcO OBz AcO 4.1 1.7 3.4 OBz OBz * Data have also been obtained for extensive series of hexose,l3 pentose,l3 inositol,13 and oligo~accharidef~ derivatives, all in aqueous solutions, and studies have been completed of peracetylated pentopyranoses,l5 and of some septano- side derivatives;16 in all instances the experimentally determined 7'1-values reflect the anticipated geometry and, especially in the case of the septanosides, provide insight as to the likely conformational symmetry of the derivative.5 Data Manipulation In this section illustrations are given of the many types of data manipulatior which can be performed once one has access to a FT spectrometer and to the concepts of spin-lattice relaxation. Three distinct opportunities exist; (A) mani-pulation of the magnetization of the sample by application of suitable sequences of pulses to achieve selective elimination of certain resonances ;(B) manipulation of the data contained in the free induction decay signal prior to Fourier trans- formation, thereby changing either the resolution or the signal-to-noise ratio of the linal spectrum; and (C) chemical modification of the relaxation pathways open to the nuclei, thereby obtaining further chemical information about the system under study.A. Manipulationof Magnetization.-It will be recalled from an earlier discussion [Figure 2(a)] that the recovery of the magnetization along the z-axis following application of a 180O-pulse is exponential, and that at a particular time the l4 L. D. Hall and C. M. Preston, Carbohydrate Res., 1973, 29, 522. L. Evelyn, L. D.Hall, and J. D. Stevens, unpublished results. Stevens, unpublished results. D.J.Hall and D.L.Berry,J. 1' Hall magnetization decays to zero intensity. Clearly, if the magnetization of the sample is assayed at that time by means of a 9O0-pulse, no resonance signal will be detected. Since protons in chemically distinct environments can have significantly different TI-values, it follows that by choosing a suitable delay time between the 180"-and 90O-pulse the resonances of individual protons can be made to dis- appear effectively from an n.m.r. spectrum. The use of this experiment to elimi- nate the residual water peak from the spectrum of an organic compound dis- solved in D2O was first demonstrated by Sykesl' and subsequently by Feeney,ls and an example from the author's work19 is given in Figure 4.The salient point here is that the residual water peak has a far longer relaxation time than the resonances of the sugar, and therefore the more rapidly relaxing sugar protons Figure 4 An example of the nulling of a residual water peak using the partial relaxation method. IH N.m.r. spectra of a 5% solution of gentiobose in D,O (99.6%); the sample had previously been lyophiiized once with D,O and then degassed by six freeze-pump-thaw cycles. (a) The normal FT spectrum (one scan); (b),the water-nulled spectrum (10 scans).In this mode of operation the magnetization is inverted by a 18O"-pulse, and this is followed, after a delay time of 0.7 times the TI-value of water, by a W"-pulse (Reproduced by permission from Carbohydrate Res., 1973, 29, 522) l7 S.L. Patt and B. D. Sykes, J. Chem. Phys., 1972,56,3182. l* F. W. Benz, J. Feeney, and G. C. K. Roberts, J. Magn. Resonance, 1972,8, 114. l9 L. D. Hall and C. M. Preston, CurbohydrafeRes., 1973, 27, 286. 411 Spin-Lattice Relaxation: A Fourth Dimension for Proton N.M.R. Spectroscopy recover almost all of their intensity in the z-direction during the time required for the water peak to reach its null point. This experiment provides an invaluable aid to studies of aqueous solutions of polysaccharide derivatives where the residual water peak often obscures some of the resonances of the anomeric protons, and we now use it routinely. We were intrigued by the possibility that this procedure might also facilitate assignments of complex proton spectra-where the success of the experiment would now depend on the presence of some intramolecular differential in relaxa- tion times. An example which illustrates the exciting potential of this method for solving the hidden resonance problem20 is given in Figure 5; we had previously I I 1 HI H3 H, HB2 H2 Figure 5 Partially relaxed 100 MHz proton n.m.r.spectra for 3,4,6-tri-O-acetyl-I -0-benzoyl-Zchloro-Zdeoxy-~-~-glucopyranose(4) in C6D6solution (0.1 mol 1-l) ;I (a), the normal spectrum; (b), the result of using a two-pulse sequence (180"-t-90") (c), the spect- rum determined with a three-pulse sequence (18O"-t-9O0 . . . 90"). For details see text. 2o L. D.Hall, Adv. Carbohydrate Chern., 2, 1974, 911. 412 Hall shownll that the H-2 and H-62 resonances of 3,4,6-tri-O-acetyl-l-O-benzoyl-2-deoxy-p-D-glucopyranose (4) have substantially different 7'1-values, and it was interesting to see whether this three-fold differential in TI-values was sufficient to enable the resonances of H-2 and H-62 to be distinguished. The spectrum in Figure 5(a) was obtained in the usual fashion. The resonance of the more rapidly relaxing proton was obtained with the usual two-pulse sequence (18O0-t-90") with t chosen such that the magnetization of the more slowly relaxing proton had decayed to zero intensity at the time when the monitoring pulse was applied; this gives H-62 as a positive-going resonance. That of the more slowly relaxing proton could be obtained by two distinct methods.In the first, the two-pulse sequence could be used, with t chosen such that now the magnetization of the more rapidly relaxing proton (H-62 in this case) has decayed to zero intensity at the time the 90"-pulse is applied; this would give the more slowly relaxing proton as a negative-going resonance. An alternative and somewhat more convenient method, illustrated in Figure 5(c), usesB1 a three- pulse sequence (18O0-t-90" . . . 90"); the second 90"-pulse now provides a refer- ence spectrum which can be automatically subtracted from that obtained by the first 90O-pulse. The final resultant of this procedure is that the resonance signal of every resonance decays to zero intensity for sufficiently long delay times; in the particular example discussed here, the signal of the more rapidly relaxing proton (H-62) reaches zero intensity faster than that of the more slowly relaxing proton (H-2).It is interesting to note that the H-61, and H-5 resonances, which have closely similar TI-values to H-62, also disappear at the same time as H-62. There are two distinct experimental protocols for the above experiments. If the 7''-values of individual protons are already known, then the value of t required for a resonance to disappear can be calculated. In the normal run of events it is likely that the Tl-values will not be known,and in this casea series of spectra are obtained using arbitrarily chosen values for t; it is then a simple matter to select by interpolation an optimal value of t to ensure the disappear- ance of any required resonance.B. Manipulation of F.I.D. Signal.-The second class of data-processing experi- ment involves mathematical manipulation of the free induction decay signal itself. The salient feature of this experiment is that any F.I.D. signal obtained experi- mentally is converted via Fourier transformation into a n.m.r. spectrum having a specific signal-to-noise ratio and a specitic line-width for the individual transi- tions. This intrinsic relationship between the F.I.D. signal and the ha1 spectrum may be altered simply by mathematically weighting some section of the F.I.D. signal immediately prior to Fourier transformation;22 increasing the weighting of the early section of the F.I.D.signal improves the effective signal-to-noise ratio of the final spectrum, whereas increasing the weighting of the later part improves the resolution. Each of these alternatives involves a compromise, however, since a1 R. Freeman and H. D. W. Hill, J. Chem. Phys., 1971,54,3367. Ba R. R. Ernst, R. Freeman, B. Gestblom, and T. R. Lusebrink, Mol. PAYS., 1967, 13, 283. 413 Spin-Lattice Relaxation: A Fourth Dimension for Proton N.M. R. Spectroscopy any increase in signal-to-noise is accompanied by some loss of resolution, and vice-versa. A simple example23 of the way in which this experiment can be used to improve the resolution of an already well dispersed spectrum is given in Figure 6, which H-3 ! Figure 6Partial 100MHz lH n.m.r.spectra of 1,2,3,4-tetra-O-acetyl-/3-~-ribopyranosein [BH,]acetone(0.1moll-l) at 42°C.(a) shows the result of a single continuous-wave scan, with a sweep-width of 250HZ and a total scan time of lo00 s. The spectrum in (b)shows the FTsummation of 100 transients, each with an acquisition time of 3.0s; the total time used to obtain this spectrum was 1500s. The spectrum given in (c) was derived from the same F.I.D. as (b) but a resolution-enhancement weighting factor of 1.0 units was applied immediately prior to the F.T. (Reproduced byapemission from CarbohydrateRes., 1975,41,41) shows spectra of 1,2,3,4-tetra-O-acetyl-~-~-ribopyranose(9). The spectra in Figures 6(b) and (c) were obtained from the same F.I.D.signal, that in (b) by direct Fourier transformation, and that in (c) with prior mathematical weighting L.D.Hall, C. M. Preston. and J, D. Stevens. Curbohphte Res., in the press. Hall AcO OAc of the latter part of the F.I.D.;* thissameresult is more clearly visible in Figure 7, which shows an expanded representation of the H-5 resonances. The excellent resolution of the small (ca. 0.5 Hz)long-range coupling in the H-5e resonance illustrates the power of this approach; indeed, it now seems sensible to reexamine the entire area of long-range coupling in carbohydrates. It is only necessary to caution here that there is an upper limit to the amount of weighting that can be applied to a F.I.D. signal, beyond which the entire information content of the F.I.D. is destroyed.This feature, together with the systematic relationship between the enhancement of resolution and the accompanying degradation of signal-to-noise ratio, is illustrated in Figure 8. C.Chemical Manipulation,-There are many ways whereby spin-lattice relaxa- tion times can be chemically manipulated. The example considered here derives from two sources; the known binding of lanthanide shift reagents, such as the europium derivative (lo), to an organic molecule via a hydroxy-group,24 and the 7 I (11) M =Ga kMe3 J3 equally well documented effect of paramagnetic metal ions on the relaxation times of associated ligands.25 The 7'1-values of the bicycloheptenol derivative (12) are shown26 in Table 2; these accord with the inverse sixth-power dependence; both the geminal protons relax faster than the methine proton, and of these two, that which is nearest to the methine proton relaxes the faster, but barely so.Table 2 shows the effect of adding ca. molar equivalents of tris(dipivalomethanato)gadolinium(m) (1 1). The relaxation of the protons of the bicycloheptanol derivative is then dominated * This experiment demonstrates another aspect of the FT method; namely the improvement of the signal-to-noise ratio of an n.m.r. spectrum. This ratio can be improved by repeated sampling of the magnetization of the system, and it increases as the square root of the num- ber of times the spectrum is scanned. The improved signal-to-noise ratio of the spectrum in Figure 6(b) was obtained by averaging 100 transients prior to Fourier transformation.p4 I. M. Armitage, G. Dunsmore, L. D. Hall, and A. G. Marshall, Cad. J. Chem., 1972, 50,2110; A. Arduini, I. M. Armitage, L.D. Hall, and A. G. Marshall, Carbohydrate Res., 1973, 26, 255. *E.C. D. Barry, A. C.T. North, J. A. Glasel, R.J. P. Williams, and A. V. Xavier, Nature, 1971, 232, 236; G. N. La Mar and J. W.Faller, J. Amer. Ckem.Soc., 1973,95, 3817. *6 V. M. Gibb and L. D. Hall, unpublished results. Spin-Lattice Relaxation: A Fourth Dimension for Proton N.M.R. Spectroscopy c1 ci by the dipole-dipole relaxation arising from the gadolinium species. This contri-bution would be expected to fall off with the increase in separation between the gadolinium atom and the individual proton.As expected, the methine proton, which is nearer to the gadolinium ion in the gadolinium-bicycloheptenolcomplex, now relaxes faster than either of the geminal protons. Furthermore, of these two, it is now the geminal proton which is cis to the hydroxy-group which relaxes the faster. It is important to note that not only has the entire relative ordering of the Table 2 Relaxation ratesls-1 for 1,2,3,4,7,7-hexachloro-5-hydroxynorborn-2-ene-asa 0.1 mol 1-1 solution in cDcl3. Column A shows the data for a normal solution, without &gassing. The data in column B were obtained afer the addition of 1.92 x 10-4 moll-1 of the gadolinium reagent (11). Column Cgives the numerical diference between columns B and C,and represents the relaxation contributions from (II). Proton studied A B C OH 0.211 4.35 4.14 H-5 0.220 1.61 1.39 H-6czo 0.341 1.47 1.13 H-6cnao 0.391 1.04 0.65 proton TI-values been inverted by the gadolinium complex but also that the differential between the geminal protons has been enhanced.This type of manipulation should significantly extend the scope of spin-lattice relaxation studies of more complex substances, at least for those which have a dominant locus of association with a paramagnetic metal. 6 Conclusions Ever-increasing access to FT n.m.r. spectrometers is clearly going to open up a number of exciting new dimensions to proton n.m.r. studies of carbohydrates and other organic derivatives.It is now possible to obtain routinely spectra of satisfactory quality from the amount of material commonly isolable from a single thin-layer chromatography plate. Providing that a suitable mathematical weight- ing program is available it is also possible to improve the effective resolution of n.m.r. spectrum at least twofold; this means that it is now practical to measure coupling constants of less than 0.1 Hz with high accuracy, and as a result it is Hall Figure 7 Expansions of the H-5eg and H-5ax resonances of 1,2,3,4-tetra-O-acetyl-fi-~-ribopyranose given in Figure 6. (a) corresponds to Figure 6(a) and (b) corresponds to Figure 6@). The half-height width of the H-5eq transitions in (b) is less than that of the H-Saz transitions, implying the presence of a small, but unresolved, coupling into H-5aX.(Reproduced by permission from Carbohydrate Res., 1975,41, 41). now appropriate to study anew the phenomenon of long-range coupling of carbohydrates. Furthermore, the fact that proton spin-lattice relaxation times can, and do, vary substantially from one site to another in a monosaccharide makes possible the routine simplification of spectra via the measurement of partially relaxed spectra. One of the more useful facets of this phenomenon is that it may be possible routinely to eliminate the resonance of methylene protons without the need for specific deuteriation of that site; this has very important implications for studies of nucleosides and nucleotides. Finally, the rapid rate 417 Spin-Lattice Relaxation: A Fourth Dimension for Proton N.M.R. Spectroscopy (e) Figure 8 The H-1 resonance of 1,2,3,4,-tetra-O-acetyl-fl-~-ribopyranose in EaHBlacetone solution (0.1 moll-l) at 42°C.All the spectra were based on the F.I.D.signal resulting from 100 transients with an aquisition time of 3.0 s and a pulse-delay time of 12 s (total time 1500 s) (a) is the result of direct F.T. with no resolution enhancement; (b)was obtained by applying a resolution enhancementfactor of 1.5;(c) a factor of 1.0;(d) a factor of 0.8;and (e) a factor 0.7. In (d) the H-1 resonance is clearly visible as a doublettedquartetwith long-range couplings of ca. 0.5 Hz. The spectrum shown in (e) illustrates the efect of applying too sharp a weighting function.(Reproduced by permission from Carbohydrate Res., 1975, 41,41). at which spectra can be acquired in the form of a F.I.D.signal opens up important areas for kinetic studies. The implications of determining the conformations of carbohydrate derivatives by measuring the relative distances between protons are many and varied, Hall Importantly though, this approach should throw further light on some of the more subtle aspects of carbohydrate conformations, especially of biologically relevant systems such as nucleosides and carbohydrate antibiotics. Beyond that, it is obvious that the methods discussed here have considerable potential for studies of any organic or organometallic system which has a well-dispersed n.m.r. spectrum.In that sense it is to be hoped that the present studies will provide a similar stimulus to that of Lemieux’s pioneering investigation27 of vicinal proton-proton coupling constants. The author is deeply indebted to his students and co-workers for their en- thusiastic and stimulating support and to Professors A. B. Foster, L. Hough, and C. A. McDowell for their steadfast advice and continuing encouragement. The experiments used to illustrate this presentation were performed by Christo- pher W. M. Grant and Caroline M. Preston, and these studies were supported by operating grants from the National Research Council of Canada and by a Fellowship from the Alfred P. Sloan Foundation. Notes added in proof: Since this lecture was presented we have developed4v28s29 a selective-pulse relaxation experiment which provides direct experimental evidence that the proton spin-lattice relaxation of a monosaccharide generally occurs exclusively viathe dipoledipole mechanism.The same type of experiment gives a quantita- tive measure of specific interproton relaxation contributions and, thereby, of interproton distances. Intercomparisons of the relaxation rates of a specifically deuteriated sugar with its ‘normal’ counterpart also al10w13s28 a simple evaluation of individual intramolecular relaxation contributions. These methods have been applied to several series of carbohydrates :(a)deriva-tives of the 1,2 :5,6-di-O-isopropylidene-a-~-glucofuranoseand the allofuranose systems;28 (b) methyl 4,6-O-benzylidene-2-deoxy-a-~-ribo-hexopyranosideand arabino-hexopyranoside;28 fluoride,(c) 2,3,4,6-tetra-O-acety1-a-~-g1ucopyranosy1 its fi-anomer, and the corresponding D-xylopyranose derivatives;30 (d) deriva-tives of ~idine;~O (e) all eight isomers of the 2,3,4-tri-O-acetyl-l,6-anhydro-/h D-hexopyranose system;30 (f) derivatives of 1,2,3,4-di-O-isopropylidene-a-~-galactopyranose;31 (g) derivatives of methyl 2,3-O-isopropylidene-/3-~-rhamno-~yranoside.~~ In every instance the proton Tpvalues appear to reflect sensibly the solution geometry of the derivatives involved.For the 1,6-anhydro-derivatives it has been proven30 that the effects on the proton 7‘1-values of changes in solvent and of solute concentration are associated with changes in rc.Preliminary studies32 R.U. Lemieux, R. K. Kullnig, H. J. Bernstein, and W. G. Schneider,J. Amer. Chem. Soc., 1958, 80,6089. z8 C. M. Preston, Ph.D. Thesis, Department of Chemistry, University of British Columbia, 1975. aB K. Bock, L. D. Hall, T. Marcus, and J. Sallos, unpublished results. so K. Bock and L. D. Hall, unpublished results. 31 L. D. Hall and K. F. Wong, unpublished results. s* K. Bock and L. D. Hall, Carbohodrate Res., 1975, 40,C3. 419 Spin-Lattice Relaxation: A Fourth Dimension for Proton N.M.R. Spectroscopy of the carbon-13 relaxation times of some sugars in aqueous solutions imply that the derivatives are tumbling isotropically. A study has also been made33 of the effects of gadolinium ions on the relaxation times of several carbohydrates in aqueous solutions. 33 C. M.Preston and L. D. Hall, Carbohydrare Res., 1975, 41 53.

ISSN:0306-0012    

DOI:10.1039/CS9750400401

出版商:RSC    

年代:1975

数据来源: RSC

摘要:

Transition Metal Complexes of Synthetic Macrocyclic Ligands By L. F. Lindoy DEPARTMENT OF CHEMISTRY A N D BIOCHEMISTRY, JAMES COOK UNIVERSITY OF NORTH QUEENSLAND, QUEENSLAND, 481 1 , AUSTRALIA 1 Introduction Naturally occurring macrocyclic-ligand transition metal complexes such as complexes of the porphyrin or corrin ring systems and the industrially important metal phthalocyanine complexes have been studied for many years. More recently, a large number of other macrocyclic ligands have been synthesized and their metal complexes have been extensively studied. The present review gives an outline of the transition-metal chemistry of this latter group of cyclic ligands with an emphasis being placed on the more recent work. The chemistry of synthetic macrocyclic ligands can be divided into two broad divisions.Firstly there are the cyclic polyethers of the ‘crown’ type of which (1) is a typical examp1e.l Ligands of this general category have received much recent attention because of their unusual behaviour towards a range of non- transition metal ions2 Few studies involving transition metal ions have been reported3 and it is evident that the majority of such polyether ligands show a limited tendency to form stable complexes with these ions.4 The second category of macrocyclic ligands incorporates the synthetic ring systems containing donor atoms other than oxygen. The majority of such ligands contain nitrogen donor atoms although ligands incorporating sulphur C. J. Pederst-n, J. Amer. Chem.SOC., 1967, 89,7017; C. J. Pedersen and H. K. FrensdorE, Angew. Chem. Znternat. Edn., 1972,11,16. * Structure and Bonding, ed. J. D. Dunitz, P. Hemmerich, J. A. Ibers, C. K. Jorgensen, J. B. Neilands, D. Reinen, and R. J. P. Williams, Vol. 16, 1973. J. J. Christensen, D. J. Eatough, and R. M. Izatt, Chem. Rev., 1974, 74, 351. A. C. L. Su and J. F. Weiher, Znorg. Chem., 1968,7,176. 421 Transition Metal Complexes of Synthetic Macrocyclic Ligands donors as well as phosphorus donors are also known. Typically such ligands form strong complexes with transition metal ions; structures (2),5 (3),6 and (4)’ illustrate early examples of this category. Ligand (2)5~8 was reported by Curtis in 1960 in its nickel complex and was isolated as one product from the reaction of tris(ethylenediamine)nickel(rI) ion with acetone; the corresponding isomer containing cis-imine groups was also formed in the reaction.A large number of other synthetic macrocycles have been investigated sub- sequently and aspects of the transition-metal chemistry of such macrocyclic systems have been discussed.3+11 2 General Considerations The majority of all nitrogen-donor macrocycles that have been studied are quadridentate, e.g. ligands (2) and (3). To fully encircle a first-row transition metal ion a macrocyclic-ring size of between 13 and 16 members is required provided that the nitrogen donors are spaced such that five-, six-, or seven- membered chelate rings are produced on co-~rdination.~JO Smaller great rings can be accommodated if the macrocycle folds and does not completely encircle the metal ion on co-ordination. For ligands incorporating sulphur donors, or for metals of larger ionic radii, there is evidence that ring sizes greater than 13 members are required to surround the metal ion.ll A number of larger ring macrocycles containing more than four donor atoms 6 B.Bosnich, C. K. Poon, and M. L. Tobe, Inorg. Chem., 1965, 4, 1102; C. K. Poon and N. F. Curtis, J. Chem. SOC., 1960,4409. M. L. Tobe, J. Chem. SOC. (A), 1967,2069; 1968, 1549. ’ M. C. Thompson and D. H. Busch, J . Amer. Chem. SOC., 1964,86, 3651. N. F. Curtis, Co-ordination Chem. Rev., 1968,3,3; J. C. S. Dalton, 1974, 347 and references therein. 9 D. H. Busch, Helv. Chim. Acta, Fasciculus Extraordinarius Arfred Werner, 1967, 1974; L.Y. Martin, L. J. DeHayes, L. J. Zompa, and D. H. Busch, J . Amer. Chem. SOC., 1974,96,4046. lo L. F. Lindoy and D. H. Busch, Prep. Znorg. Reactions, 1971,6, I . l1 D. H. Busch, K. Farmery, V. Goedken, V. Katovic, A. C. Melnyk, C. R. Sperati, and N. Tokel, in ‘Bioinorganic Chemistry’, Advances in Chemistry Series No. 100, American Chemical Society, New York, 1971, p. 44. 422 Lindoy have also been prepared;3 (312 and (a13 illustrate ligands of this type. For co-ordination to a transition-metal ion, these large ring macrocycles generally (but not always14) must be capable of twisting so that favourable co-ordination NO, (6) geometries are achieved about the central ion. Isolated examples of related large-ring ligands incorporating more than one transition metal ion are also known; (7) is one such example.'5 However complexes of this type have been observed more frequently with polyether macrocycles and non-transition metal ions.1-3 Me I 2+ I s D.St. C. Black and I. A. McLean, Inorg. Nuclear Chem. Letters, 1970,6,675. l 3 L. F. Lindoy and D. H. Busch, Chem. Comm., 1968, 1589; J. Amer. Chem. SOC., 1969,91, 4690. l4 L. F. Lindoy, D. H. Busch, and V. Goedken, J.C.S. Chem. Comm., 1972,683 ; L. F. Lindoy and D. H. Busch, Inorg. Chem., 1974,13,2495; N. W. AIcock, D. C. Liles, M. McPartlin, and P. A. Tasker, J.C.S. Chem. Comm., 1974,727. l5 N. H. Pilkington and R. Robson, Austral. J. Chem., 1970 23,2225. 423 Transition Metal Complexes of Synthetic Macrocyclic Ligands Another category of polydentate macrocycles includes the polycyclic 'cages' which completely encapsulate the metal ion on co-ordination; these are discussed in a later section.A number of such complexes having unusual co-ordination geometries have now been synthesized. It is pertinent, at this point, to consider why macrocyclic ligands often yield complexes which show unusual properties (compared with similar complexes of related open-chain ligands). There are three main considerations. Firstly, on complex formation, it is apparent that geometrical factors arising from the cyclic nature of the ligands often impose additional constraints on the positions of the donor atoms. As a reflection of these constraints, macrocyclic ligand complexes containing unusual metal-donor atom bond distances, unusual bond angles, or grossly strained chelate-ring conformations are all known.Indeed, a few examples of constraints of the latter type leading to complexes which have unusual co-ordination geometries have also been reported.lOJ4 Secondly, if the cyclic ligand is fully conjugated and incorporates (4n+2) 7~ electrons then enhanced electron delocalization and ligand stability are charac- teristic of the resulting Hiickel aromatic system. Thirdly, cyclic-ligand complexes are almost always found to be considerably more stable thermodynamically and kinetically (with respect to dissociation of the ligand from the metal ion) than their corresponding open-chain analogues.Such properties seem to be an intrinsic feature related to the cyclic nature of the ligands and have been collectively referred to as the macrocyclic effect.16 Studies are still underway concerning the nature of this effect.leJ7 However, it seems clear that the enhanced thermodynamic stability does not reflect solely an increased 'chelate' effect owing to the presence of an additional chelate ring in the cyclic-ligand complex (cf. the analogous open-chain complex). Recent studies17 on a series of nickel open-chain and macrocyclic amine complexes suggest that the enhanced thermodynamic stability associated with the macro- cyclic effect results largely from a more favourable AH" term. This is so even though the donor-metal bond strengths do not appear to differ greatly between similar open-chain and macrocyclic complexes.Rat her the enthalpic differences are attributed to decreased ligand solvation in the macrocycle which thus has less associated (hydrogen-bonded) water to be displaced on complex formation. Hence AH" for complexation of the macrocycle will be more negative than that for complexation of the linear-chain ligand which, because of its open structure, is postulated to be more highly solvated in its uncomplexed form. On the other hand although there will be a greater loss of configurational entropy on co- ordination of the linear-chain ligand, this will tend to be compensated by the more favourable AS" component arising from the release of a greater number of solvent molecules in this case.16 D. K. Cabbiness and D. W. Margerum, J . Amer. Chem. SOC., 1969, 91, 6540; P. Paoletti, L. Fabbrizzi, and R. Barbucci, Inorg. Chem., 1973,12, 1961. l7 F. P. Hinz and D. W. Margerum, J. Amer. Chem. SOC., 1974,96,4993; Znorg. Chem., 1974, 13,2941. 424 Lindoy 3 Typical Synthesis Cyclic ligands have been prepared both directly, by conventional organic synthesis, and by in situ procedures involving cyclization in the presence of a metal ion. The ‘crown’ polyethers are examples of macrocycles which have been prepared mainly by direct syntheses.1 Typically, in these preparations, the cycli- zation reaction is performed under conditions of moderate to high dilution in order to minimize competing linear polymerization reactions.Similarly the quadridentate amine (3 ; ‘cyclam’) can be obtained in small yield by condensation of 1,3-bis-(2’-aminoethyIamino)propane and 1,3-dibromopropane in alcoholic potassium hydroxide under high dilution conditions ;6 however a recent report gives an improved procedure involving an in situ condensation for obtaining this cyclic amine.l*Jg New direct synthetic procedures for obtaining related cyclic amines (and amine-ethers) in excellent yields have been developed recently.20 A typical reaction involved the condensation of the tosylated reactants (8; Ts = tosyl) and (9) by heating in DMF at 100 “C for 1 h. The tetratosyl derivative (10) was obtained in 80% yield; the tosyl groups of (10) were readily removed by heating this product in concentrated sulphuric acid.Ts Ts Ts T S Ts Although a number of macrocyclic ligands can be synthesized both by direct as well as by in situ procedures, it is nevertheless true that a large proportion of other cyclic ligands can only be obtained via in situ techniques. Such cyclization reactions, in some instances, may be aided by metal-ion induced reactivity of one of the condensing functional groups as well as by steric factors involving the appropriate positioning around the co-ordination sphere of the reacting moieties in the ring-closing step. Some representative examples of such metal-ion template reactions will now be discussed. The Schiff-base condensation between a carbonyl compound and an organic amine to yield an imine linkage has formed the basis of many successful macro- cyclic-ligand syntheses;21922 e.g., condensation of 2,6-diacetylpyridine with the l 8 E.K. Barefield, Znorg. Chem., 1972, 11, 2273. l8 E. K. Barefield and F. Wagner, Inorg. Chem., 1973,12,2435. a o J. E. Richman and T. J. Atkins, J , Amer. Chem. SOC., 1974,96,2268. 21 L. F. Lindoy, Quart. Rev., 1971,25, 379. 2a D. St. C. Black and A. J. Hartshorn, Co-ordination Chem. Rev., 1972,9,219. 425 R. W. Alder, R. Baker, and J. Brown, ‘Mechanism in Organic Chemistry’, Wiley-Inter- Transition Metal Complexes of Synthetic Macrocyclic Ligands appropriate triamine in the presence of several transition-metal ions yields corresponding complexes of the 14membered macrocycle (1 1) in high ~ield.~3 This approach has been adopted subsequently to the synthesis of a number of related Schiff base macrocycles based on 2,6-diacetylpyridine;24-28 several of these latter ligands incorporate more than four donor atoms.The in situ condensation of a di-(or poly-)amine with a dicarbonyl moiety has been a common procedure for the preparation of complexes of new macro- cyclic ligands. Such condensations are usually quite facile when, as in the reactions just discussed, the dicarbonyl moiety incorporates one or more donor atoms between the two carbonyl groups. Lewis acids are known to catalyse nucleophilic reactions at carbonyl-carbon atoms2Q and it has been suggested10 that the presence of additional, appropriately spaced, donor atoms between the carbonyl groups will result in the carbonyl-containing moiety bonding more strongly to the metal ion (Lewis acid).With respect to this enhanced bonding, the ability of the dicarbonyl moiety to form a number of chelate rings is un- doubtedly important. In such a situation, the activation of a carbonyl group bound to the metal ion has been ascribed to increased polarization of the carbonyl group so that attack by the amine nucleophile on the carbonyl-carbon is facilitated. Nevertheless the situation can be more complex than the simple polarization argument suggests and caution is needed when rationalizing an observed rate enhancement solely in terms of polarization effects.30 Indeed, for a number of related reactions involving nucleophilic attack at bound carbonyl centres, the observed rate enhancements result largely, in fact, from favourable 23 J . D.Curry and D. H. Busch, J . Amer. Chem. SOC., 1964, 86, 592; L. F. Lindoy, N. E. Tokel, L. B. Anderson, and D. H. Busch, J. Co-ordination Chem., 1971,1,7 and references therein. 24 S. M. Nelson, P. Bryan, and D. H. Busch, Chem. Comm., 1966,641. 26 E. Fleischer and S. W. Hawkinson, J . Amer. Chem. SOC., 1967,89,720. R. H. Prince, D. A. Stotter, and P. R. Woolley, Inorg. Chim. Acfa, 1974,9,51. R. W. Stotz and R. C. Stoufer, Chem. Comm., 1970, 1682. J. Rikernap and D. W. Meek, J.C.S. Chem. Cornm., 1974,442. science, London, 1971, p. 314. 30 C. K. Ingold, ‘Structure and Mechanism in Organic Chemistry’, Cornell University Press, Ithaca, 1953, pp.49-52. 426 Lindoy entropies of activation and not from favourable enthalpic terms as would be expected to occur as a consequence of polarization effects.31 Another related series of in situ cyclizations involving a moiety which contains two terminal aldehyde functions as well as other suitably-spaced donor atoms in the backbone is illustrated by the reaction of simple diamines with dialdehydes of type (12).3a The reaction can be performed in the presence of the metal acetate so that the corresponding neutral complex, of which (13) is a typical example, is produced. Alternatively, for some of the cyclic ligands derived from (12), the free macrocycle can be prepared directly. This is in contrast with the four-nitrogen macrocycle (11) which has not been isolated free of its metal ion.Complexes of the macrocyclic ligands (14) containing two sulphur33 or two n Chem., 1975,14, in the press. x s,o oxygena4 donors have also been prepared; in both cases the in situ cyclizations proceed smoothly, and some of the oxygen ligands have also been isolated in the absence of a metal ion. These latter macrocycles are of interest since they are intermediate in structure between the ‘crown’ polyethers and the macrocycles which incorporate four nitrogen donor atoms. A limited number of other related intermediate macrocycle types have been reported recently3,ZO and there and M. Wein, J. Amer. Chem. Soc., 1972,94,4032. 3* M. Green and P. A. Tasker, Chem. Comm., 1968, 518; M. Green, J. Smith, and P.A. Tasker, Inorg. Chim. Actu, 1971,5, 17; D. St. C. Black and P. W. Kortt; Austral. J. Chem., 31 D. A. Buckingham, J. MacB. Harrowfield, and A. M. Sargeson, J. Amer. Chem. Soc., 1974, 96, 1726 and references therein; D. A. Buckingham, J. Dekkers, A. M. Sargeson, 1972,25,281; E. N. Maslen, L. M. Engelhardt, and A. H. White, J.C.S. Dalton, 1974,1799. 33 L. F. Lindoy and D. H. Busch, Inorg. Nuclear Chem. Letters, 1969,5, 525. 34 L. G. Armstrong and L. F. Lindoy, Inorg. Nuclear Chem. Letters, 1974, 10, 349; Inorg. 427 Transition Metal Complexes of Synthetic Macrocyclic Ligands is evidence that such ligands may show an enhanced specificity towards a number of metal-ions, relative to related ligands containing only oxygen or only nitrogen donor atoms.3~3~ Salts of the diamagnetic complex ion (15) have been isolated from the reaction, in THF, of the dihydrazine (16) with formaldehyde in the presence of nickel ion.35 Inherent in the synthetic strategy for this template reaction is the thought that the use of a dihydrazine moiety such as (16) may well overcome the masking of the nucleophilic character of amines which is usually observed when amines are co-ordinated to metal ions.21 Unlike a co-ordinated amine, a co-ordinated hydrazine still contains a nitrogen with an unshared electron-pair which may participate in nucleophilic condensation reactions.For co-ordinated amines [e.g. (17)36] such reactions are usually inhibited unless the amine first dissociates I ! NH, NH8 n n+ from the metal ion.21 The preparation of (15) is typical of a number of other recent in situ syntheses in which hydrazine (or hydrazone) precursors are in- volved .37 -39 6,6’-Dihydrazino-2,2’-bipyridyl has recently been demonstrated39 to be a highly reactive precursor for the formation of macrocyclic ligands.On com- plexation this dihydrazine chelates strongly through the a-di-imine nitrogens of the bipyridyl residue but steric influences allow only weak interaction between the metal and the terminal amines of the hydrazine groups. Such groups are 35 N. W. Alcock and P. A. Tasker, J.C.S. Chem. Comm., 1972, 1239. 3E L. F. Lindoy, unpublished work. 37 V. L. Goedken and S. Peng, J.C.S. Chem. Comm., 1973,62. 3 8 C . M. Kerwin and G. A. Melson, Znorg.Chem., 1973, 12, 2410; 1972, 11, 726; G. A. Melson, Inorg. Chem., 1974, 13, 994; V. L. Goedken, Y. Park, S. Peng, and J. M. Norris, J . Amer. Chem. SOC., 1974,96,7693. s 9 J. Lewis and K. P. Wainwright, J.C.S. Chem. Comm., 1974, 169. 428 Lindoy thus in a state of unusualIy high lability and nucleophilicity. Indeed the nickel(r1) complex reacts with refluxing aqueous acetone in a few minutes with the forma- tion, in near-quantitative yield, of the nickel complex of the Curtis type macro- cycle (1 8). These conditions are the mildest yet observed for ring-closing reactions of the Curtis type. The synthesis of a quadridentate macrocycle derived from another ‘classical’ chelating ligand (namely, 1,lO-phenanthroline) has also been achieved recently.40 From the reaction of 2,9-dichloro-l , 10-phenanthroline with 2,9-diamino-1 ,lo- phenanthroline in nitrobenzene in the presence of potassium carbonate, the macrocycle (19) can be isolated in high yield.This reaction provides an example of another category of macrocycle direct synthesis. Unlike the previously discussed examples, the present cyclization involves sites which are remote from the donor atoms. Reaction of (19) with copper chloride in nitrobenzene leads to the loss of two protons with the formation of the neutral metal complex. 429 4 Reactions Involving Co-ordinated Macrocyclic Ligands A number of macrocyclic-ligand complexes have been chemically modified to produce new cyclic-ligand complexes. The relative kinetic inertness of most macrocyclic-ligand complexes (even when they contain normally labile metal ions) makes such systems ideal for the study of metal-ion-influenced ligand reactions since the prospect of ligand dissociation during the course of the reaction is much reduced.Redox reactions involving chemical, electrochemical, and catalytic procedures are collectively the most studied group of cyclic ligand reactions. Oxidative dehydrogenations of a range of cyclic-ligand complexes have been reported using a variety of chemical oxidizing agents. In most instances the reaction has involved the conversion of co-ordinated secondary amines to imine functions. The reverse process, viz. hydrogenation of an imine to an amine, is also facile for many cyclic ligand complexes. Hydrogenations of this type have frequently been achieved catalytically (Hz in the presence of Raney nickel or a precious metal catalyst) or chemically, using reagents such as sodium borohydride.Both reactions have also been performed electr~chemically.~~ 4 0 S. Ogawa, T. Yamaguchi, and N. Gotoh, J.C.S. Chem. Comm., 1972, 577. I1 F. V. Lovecchio, E. S. Gore, and D. H. Busch, J. Amer. Chem. SOC., 1974, 96, 3109 and references therein. Dansition Metal Complexes of Synthetic Macrocyclic Ligands Starting from the nickel complex of (1 l), it has been possible, by a combination of selective hydrogenations and chemical oxidations, to produce a range of derivatives containing from zero to three imine linkages, respecti~ely.4~9~3 7 he nickel complexes of a range of Curtis macrocycle derivatives containing from zero to four imine groups have also been isolated as the products of a similar series of redox reactions.8 There is evidence that redox reactions of the type just described are metal-ion dependent both in the ease with which a reaction takes place as well as in the nature of the product formed.These points are well illustrated by comparison of the oxidation of the Curtis iron(@ complex (20) with that of its nickel(ii) analogue. In acetonitrile, the former undergoes a stepwise reaction with molecular oxygen to yield (21) as the final product;44 whereas, for the nickel complex, oxidation leads to the tetra-imine species (22).S As in this case, it is generally true that iron@) complexes undergo more facile oxidative dehydrogenations than do their nickel(ii) analogues.44 Furthermore if the diene complex (23) is used as the starting complex then on reaction with oxygen in acetonitrile, an intermediate of type (24) initially forms but this spontaneously tautomerizes to (25).Compared with the behaviour of the corresponding nickel system,8 the above * % E. K. Barefield, F. V. Lovecchio, N. E. Tokel, E. Ochiai, and D. H. Busch, Znorg. Chem. 1972,11,283. 4 3 J. L. Karn and D. H. Busch, Inorg. Chem., 1969,8,1149. 4 4 J. C. Dabrowiak, F. V. Lovecchio, V. L. Goedken, and D. H. Busch, J. Arner. Chem. SOC., 1972,94,5502; V. L. Goedken and D. H. Busch, ibid, 1972,94,7355. 430 Lindoy reactions clearly indicate the well-known preference45 of iron(n) for ligands containing an a-di-imine linkage.The ligand oxidative dehydrogenations so far discussed have all involved the introduction of unsaturation adjacent to nitrogen donor atoms. Examples of unsaturation being introduced between adjacent carbons in the backbone of a chelate ring are also known although such reactions have been observed less frequently. An example of this class of reaction involves the nickel complex (26) which incorporates a negatively charged electron-delocalized chelate ring derived from acetylacetone.46 On reaction with bromine in acetonitrile, this complex yields (27) in which a localized C=C linkage has been introduced into the six-membered chelate ring; conjugation has also been extended into an adjacent chelate ring. Oxidations in which unsaturation is introduced between carbon atoms in five-membered chelate rings have also been reported recentl~.~7 Such reactions involve nickel@) and copper(rr) complexes of type (28).Treatment of these YY 2+ w CL") 4 5 L. F. Lindoy and S. E. Livingstone, Co-ordination Chem. Rev., 1967,2, 173. 4 6 C. J. Hipp, L. F. Lindoy, and D. H. Busch, Znorg. Chem., 1972,11, 1988. 4 7 T. J. Truex and R. H. Holm, J. Amer. Chem. SOC., 1972,94,4529. fi complexes with three equivalents of trityl tetrafluoroborate in acetonitrile leads to oxidative dehydrogenation such that 15welectron species of type (29) are produced. Reduction of (29) with sodium borohydride in ethanol yields the uncharged 16n-electron species (30).Voltammetric studies indicate that (29), (30), and (31) are interconvertible by means of a series of one-electron (reversible) redox steps; the terminal member of this series (31) contains a stable (4n+ 2) 7r- electron ring system. 431 Transition Metal Complexes of Synthetic Macrocyclic Ligands There is considerable evidence that higher oxidation states of the central metal are involved in the mechanisms of many ligand oxidative dehydrogena- tions of the type so far des~ribed;~~s4~$4~-~* e.g., oxidation of nickel(I1) and iron@) complexes may proceed via generation of the respective tervalent states followed by intramolecular redox processes in which the ligand is oxidized. Indeed, examples have been reported of iron(rr1) macrocyclic-ligand complexes spontaneously generating new imine linkages via oxidative dehydrogenations under certain condition^.^^ In addition, it has been possible to isolate a number of macrocyclic-ligand nickel(Ir1) complexes as products from oxidation reactions on selected corresponding nickel@) specie^.^^^^^ Typically, chemical oxidizing agents such as concentrated nitric acid, ammonium persulphate, or nitrosyl tetrafluoroborate in acetonitrile have been used in these syntheses.Electrochemical studies have amply demonstrated the ability of certain macrocyclic ligands to stabilize a wide range of oxidation states of co-ordinated metal ion~.~l*5~-53 The relative stabilities of the bi- and tervalent oxidation states of iron depend on the degree and nature of the unsaturation in the attached cyclic ligand.44 Cyclic voltammetry indicates that each of the iron(@ complexes (20), (21), (23), and (25) can be reversibly oxidized in acetonitrile to its iron(m) analogue.Such complexes of tervalent iron have also been prepared directly and shown to reduce at the expected half-wave potentials. Whereas the ,??+ value for the FelI/FelI1 couple becomes slightly more positive for each additional isolated double bond introduced into the ligand, the presence of an a-di-imine linkage has a much more dramatic effect and leads to considerably greater positive shifts of Q. Nickel(I1) complexes of a range of neutral (e.g., cyclam) and dianionic macro- cyclic ligands have also been demonstrated to undergo one-electron electro- chemical oxidations to yield stable six-co-ordinate and square-planar nickel(II1) species,41 respectively. The authenticity of the nickel(II1) state (d7 configuration) in these compounds cannot be established unambiguously from magnetic moment measurements alone since the technique is unable to rule out the alternative possibility that the compounds are nickel(I1) species containing stabilized cation radical ligands.However e.p.r. studies clearly indicate that the unpaired electron is localized on the nickel ion in each case; the compounds are therefore correctly formulated as nickel(1rr) species. The complementary one-electron reduction products of the above complexes have also been studied.41 In the case of the uncharged-ligand complexes the e.p.r.spectra fall into two categories. For the complexes of saturated ligands or for those containing isolated imine groups, axial symmetric spectra are observed with gL < g , , . Such spectra are consistent with the presence of nickel@) (d9) species. 4 8 D. F. Mahoney and J. K. Beattie, Znorg. Chem., 1973,12, 2561 and references therein. 4 9 E. K. Barefield and D. H. Busch, Chem. Comm., 1970, 522. 5 0 E. Gore and D. H. Busch, Znorg. Chem., 1973,12, 1. 51 D. C. Olson and J. Vasilevskis, Znorg. Chem., 1969, 8, 161 1. 5 2 D. C. Olson and J. Vasilevskis, Znorg. Chem., 1971,10,463. 63 D. P. Rillema, J. F. Endicott, and E. Papaconstantinou, Znorg. Chem., 1971,10, 1739. 432 Lindoy However, for complexes containing a-di-imine groupings, an isotropic e.p.r. spectrum (containing no hyperfine splitting) is observed with a g value close to that of the free electron. Clearly the unpaired electron in these systems is de- localized on the ligand and such complexes are best formulated as containing nickel(@ co-ordinated to a stabilized anion-radical ligand. Hence, as with the iron complexes discussed above, the presence of an a-di-imine group stabilizes the bivalent state of nickel relative to the tervalent state. These results form part of a recent extensive study41 of the influence of structural and electronic parameters on the electrochemical redox behaviour of a range of four-nitrogen macrocyclic complexes of nickel(r1). The redox behaviour of such complexes in a given solvent is a composite function of a range of factors such as ring size, ring substituents, degree and type of unsaturation, charge type, and co-ordination number.It has been possible to document, in a semi-quantitative manner, the effects of such factors and a concept of additivity of structural contributions to redox behaviour has been proposed. A second category of the reactions of co-ordinated macrocycles involves the acid-base properties of certain ligands of this type. The reversible protonation of the charge-delocalized six-membered chelate ring in nickel(1r) and copper(I1) complexes of type (32) to yield species of type (33) has recently been investi- A H+ 7 OH- (32) (33) gated.54~55 The PKa for the dissociation is markedly metal-ion dependent ; e.g., for 14membered macrocycles the PKa of the nickel complex is 6.45 while for the copper analogue a PKa of 9.3 was observed.The values obtained for the 13- membered macrocyclic complexes are of the same order of magnitude as above but there is a considerable decrease in acidity for the 15-membered macrocyclic complexes. Although the reasons for these differences are somewhat obscure, an explanation based on different degrees of steric crowding has been proposed. Other examples of related reversible protonations have also been docu- mented.11g56157 Ligand-substitution reactions provide a further category of reactions in- volving co-ordinated macrocyclic ligands, and a number of such substitution reactions have been studied. An example is the alkylation of co-ordinated 5 4 J.G. Martin and S. C. Cummings, Inorg. Chem., 1973, 12, 1477. 6 5 J. G. Martin, R. M C. Wei, and S. C. Cummings, Inorg. Chem., 1972,11,475. 5 6 C. J. Hipp and D. H. Busch, J.C.S. Chem. Comm., 1972, 737. 6 7 E. JBger, 2. Chem., 1968, 8, 30. 433 Transition Metal Complexes of Synthetic Macrocyclic Ligandr (secondary) amines of bound macrocyclic ligands. Such co-ordinated amines normally do not show nucleophilic properties; however, an amide nucleophile is generated on deprotonation. The conditions for performing such N-depro- tonations and N-alkylations have been well established previously58 and the reaction has been extended to a number of fully saturated cyclic-ligand com- plexes such as the nickel complex of cyclam (3).59 In this case the nickel complex of tetra-methylated cyclam is produced and this is a typical product of such react ions .59 Tetra-methylated cyclam can also be prepared by directlg,6O organic synthesis and interestingly, when it is treated with nickel ion, a kinetically labile complex is produced.19 In contrast, the similar complex resulting from ligand alkylation in situ is remarkably kinetically i11ert.5~ The different kinetic stabilities of these two products reflect different configurations of the four N-methyl groups.Apparently the product from the synthesis involving pre-formed ligand has the four methyl groups arranged on the same side of the co-ordination plane defined by the four donor nitrogen atoms.19 In contrast, a crystal structure determination on the product from the in situ reaction indicates that this complex contains two of the methyl groups above and two below the co-ordination plane of the four nitrogen donors.59 The reactions of various reagents (particularly electrophiles) towards a range of macrocyclic complexes containing acetyl-su bstit uted charge-delocalized chelate rings has recently been reported.56~6~ The neutral complex (34) is typical of the compounds used in these studies and in several instances the reactivity patterns observed resemble those of transition-metal corrole complexes.The acetyl groups of (34) are displaced by protonation of the carbons to which they are attached. Subsequent studies have indicated that a range of nitrogen electro- philes also effect electrophilic displacement reactions so that the corresponding dinitro-species, e.g.( 3 3 , are produced. Sodium nitrite and trifluoroacetic acid in acetonitrile, nitric oxide in the presence of air, nitrosyl tetrafluoroborate in dichloromethaiie (under oxidizing conditions) and concentrated nitric acid have all been used for reactions of this type. Hence the dinitro-species can be produced either by nitrosation under oxidizing conditions or by direct nitration. The (34) (3 5 ) G. W. Watt and P. W. Alexander, Znorg. Chem., 1968,7,537 and references therein. Kg F. Wagner, M. T. Mocella, M. J. D'Aniello, A. H. J. Wang, and E. K. Barefield, J. Amer. Chem. SOC., 1974, 96, 2625.6o R. Buxtorf and T. A. Kaden, Helv. Chim. Actu (A), 1974,57,1035. C. J. Hipp and D. H. Busch, Inorg. Chem., 1973,12,894. 434 Lindoy success of these reactions thus confirms the considerable nucleophilic character of the reaction sites of the precursor complexes. 5 The Cage Macrocycles Over the past few years complexes of a number of novel three-dimensional ‘cage-like’ macrocyclic ligands (clathrochelates) have been synthesized. These ligands encapsulate the central metal ions and the resulting complexes often exhibit unusual properties. For complexes of this type investigations pertaining to stereochemistry, metal-ion redox behaviour, and metal-ion transport across membranes are of special interest. Examples of cage-ligand complexes are (36)s2 and (37).63964 The synthesis of (36; X = F) was based on the previously reported procedure65 involving reaction of boron trifluoride etherate with bis(dimethylg1yoximato)- nickel(@.The product of this reaction is the very stable complex (38; M = NP) Chem., 1971, 10, 2472. + in which -BFF linkages have bridged adjacent oxime groups. In a related fashion, the reaction of the tris(dimethylglyoximato)cobalt(nI) anion and boron trifluoride yields the cage (36; M = CoIII, X = F, n = 1).G2 Apart from the boron halides, a number of other Lewis acids such as SnC14, SiC14, and H3B03 6 2 D. R. Boston and N. J. Rose, J. Arner. Chern. SOC., 1968,90,6859. 63 J. E. Parks, B. E. Wagner, and R. H. Holm, J . Arner. Chern. SOC., 1970,92, 3500; Znorg.64 E. Larsen, G. N. La Mar, B. E. Wagner, J. E. Parks, and R. H. Holm, Znorg. Chem., 1972,11,2652. 6 5 G. N. Schrauzer, Chem. Ber., 1962,95, 1438; D. Thierig and F. Umland, Angew. Chem., 1962, 74, 1438. 435 (40) E * D. R. Boston and M. J. Rose, J. Amer. Chem. SOL, 1973.95,4163. E7 S. C. Jackels and N. J. Rose, Inorg. Chem., 1973, 12, 1232; S. C. Jackels, D. S. Dierdorf, Transition Metal Complexes of Synthetic Macrocyclic Ligands have been used for similar capping reactionss6 More recently it has been shown that there is no need to isolate the intermediate tris-dioxime complex but that the required clathrochelate can be isolated directly from a mixture of all re- actants.67 This latter in situ procedure has been used to obtain iron(I1) clathro- chelates of type (36) with X = F, OH, OMe, OEt, OPri, or OBun and n = 0.A similar capping procedure has also been used to prepare clathrochelates of type (37; M = FeTI, CoII, NiJ1, or ZnII) from the complexes of the corresponding trigonally symmetric sexadentate ligand containing a phosphorus bridgehead and three terminal oxime gr0ups.~3964 Both boron trifluoride and the tetrafluoro- borate ion were used as capping reagents in this case. It is clear from a number of X-ray structural st~dies6*.~~ that both the cage types above show a tendency to promote a trigonal prismatic co-ordination geometry about the central metal ion; examples of such complexes which are almost exactly trigonal prismatic as well as structures showing various degrees of distortion from this basic geometry are known.The in situ preparations of metal complexes of a new cage derived from three molecules of the bis-hydrazone of biacetyl (and capped by reaction with form- aldehyde) have been reported.37 Although the synthetic procedure falls into the category previously discussed involving condensat ion at terminal hydrazine or hydrazone moieties little can be said, in the absence of further studies, about the generality of such a capping procedure for the preparation of other cages. A number of polyether-amine cages related to the ‘crown’ macrocycles have now been prepared and their behaviour towards an extensive range of non- transition ions has been studied in considerable detai1.2s3 Such ligands are usually synthesized directly; (39) and (40) are two typical examples having different cavity sizes.A feature of such ligands is that they are flexible enough to allow entry of suitably sized cations without ligand rupture. The complexation of (40) with cobalt(I1) has recently been reported.’O The structure of the resulting cation indicates that the cobalt is completely enclosed within the cavity and is bonded to each of the seven donor atoms in a distorted (39) N. J. Rose, and J. Zektzer, J.C.S. Chem. Comm., 1972, 1291. E8 G. A. Zakrzewski, C. Chilardi, and E. C. Lingafelter, J. Amer. Chem. SOC., 1971, 93,4411. Es M. R. Churchill and A. H. Reis, Chem. Comm., 1970,879; 1971,1307; Inorg. Chem., 1972, 11,1811,2299; 1973,12,2280; J.C.S. Dalton, 1973, 1570.‘O F. Mathieu and R. Weiss, J.C.S. Chem. Comm., 1973, 816. 436 Lindoy pentagonal bipyramid co-ordination geometry. One of the factors that appears to influence the formation of this unusual geometry is the small size of the cobalt(I1) ion with respect to the ligand cavity. The situation appears to parallel that which obtains in a number of metallo-enzymes in which constraints in the immediate vicinity of the metal ion are also known to result in unusual co- ordinat ion ge~metries.~~ 6 Macrocyclic Ligand Complexes Related to the Biological Macrocyclic Systems Synthetic ring complexes which copy aspects of the behaviour of the more complicated natural macrocyclic-ring systems (such as those containing por- phyrin or corrin rings) are known and at present the study of such compounds is receiving much attention.Although the results obtained so far do not always closely parallel those in nature, a knowledge of the chemistry is being built up and the biochemical role of metal ions in the natural systems is beginning to be better understood. In this section selected examples of the use of synthetic macrocyclic complexes for model studies are briefly discussed. Since the discovery that the natural product, coenzyme B12, is a cobalt(n1) corrinoid complex containing a o-bonded alkyl ligand (occupying an axial position),72 there has been much interest in the preparation and study of synthetic analogues of vitamin B12 and its coenzymes. A large number of model compounds containing stable cobalt-alkyl moieties have now been reported, and much interest has centred on the factors influencing stabilization of the cobalt-carbon bond.11e73-75 The majority of the model compounds so far prepared have contained unsaturated nitrogen-donor ligands (not necessarily macrocyclic) in the equatorial plane.For example, bis(dimethylg1yoximato)cobalt derivatives73 (the cobaltoximes) show many of the reactions of the BIZ-systems; so do the corresponding cyclic derivatives of type (38). Other cobalt-containing macro- cycles such as (1 1) have also been used for similar studies.11 For such unsaturated ligands, synthesis of the respective cobalt-alkyl species has very often involved the reaction of a corresponding cobalt(1) or cobalt(I1) precursor with alkyl halide although other synthetic procedures have also been successfully empl0yed.7~ It has generally been assumed that unsaturated ligands in the equatorial plane are a prerequisite for the formation of stable cobalt(Ir+-alkyl bonds.However a recent photochemical procedure has enabled the isolation of stable cobalt(xxx)- alkyl complexes of saturated macrocycles such as ~ y c l a m . ~ ~ Thus it seems clear 71 D. D. Ulmer and B. L. Vallee, in ref. 11, p. 187. 7 s P. G. Lenhert and D. C. Hodgkin, Nature, 1961,192,937. G. Schrauzer, Accounts Chem. Res., 1968, 1, 97; H. A. 0. Hill, J. M. Pratt, and R. J. P. Williams, Discuss. Faradav SOC., 1969, 47, 165; J. Lewis, R. H. Prince, and D. A. Stotter, J. Inorg. Nuclear Chem., 1973,35, 341 ; D.G. Brown, Progr. Inorg. Chem., 1973, 18, 177. 74 J. M. Pratt, ‘Inorganic Chemistry of Vitamin BIZ’, Academic Press, New York, 1972; J. M. Pratt and P. J. Craig, Adv. Organometallic Chem., 1973, 11, 331. 7 6 M. Green, J. Smith, and P. A. Tasker, Discuss. Faraday SOC., 1969,47, 172; L. M. Engel- hardt and M. Green, J.C.S. Dalton, 1972, 724. V. L. Goedken, S. M. Peng, and Y. Park, J. Amer. Chem. SOC.. 1974,96,284 and references therein. T. S. Rocheand J. F. Endicott, Inorg. Chem., 1974,13,1575. 437 Transition Metal Complexes of Synthetic Macrocyclic Ligads from the model studies that neither highly unsaturated nor macrocyclic ligands are essential for the stabilization of cobalt-alkyl bonds of the type found in the B12 series. Nevertheless the special properties (e.g.unusual redox properties78) that result from incorporation of the highly unsaturated corrin ring in vitamin B12 and related coenzymes are no doubt of major importance to other aspects of the biological function of this important group of compounds. A common feature of the cobalt in vitamin B12 as well as the iron in such species as cytochrome c, haemoglobin, and myoglobin is that one of the axial co-ordination positions is occupied by a heterocyclic base (viz. an imidazole derivative). Synthetic planar quadridentate macrocycles containing flexible side- chains incorporating appropriately positioned nitrogen donor groups (mainly heterocyclic) have been reported recently.79s80 In these the side-chains are of sufficient length to permit the attached nitrogen donors to occupy a position axial with respect to the plane of the macrocycle.The iron(I1) complex (41) of one such ligand has been used as a model for the myoglobin active site.8o This complex binds oxygen reversibly both in solution (in methylene chloride at -45 "C) and in the solid state and, like myoglobin and haemoglobin, binds carbon monoxide more strongly than it binds oxygen. It has been known for a considerable time that a variety of cobalt(I1) com- plexes of both cyclic and non-cyclic ligands will reversibly bind molecular oxygen at room temperature.8l However, previous solution studies involving corresponding iron(@ complexes have characteristically resulted in oxygen- uptake followed by autoxidation during which the iron(@ is converted into the tervalent state.11v82 Oxygen-uptake is thus irreversible in such cases.Recent s t u d i e ~ , ~ ~ - 8 ~ all at low temperature, have shown that the covalent attachment of an imidazole derivative to the haem group such as occurs in (41) is not essential for reversible oxygenation but that, under controlled conditions in solution, oxygenation will also occur in the presence of non-attached amines. Nevertheless, although not essential for reversible oxygenation, covalent attachment of the imidazole to the haem group does still appear to aid the reaction.84 It has been suggested87 that one mechanism for autoxidation of iron(I1) complexes involves the initial 1 :1 binding of molecular oxygen by the iron@) '* N.S. Hush and I. S. Woolsey, J . Amer. Chem. SOC., 1972,94,4107. 7 9 K. B. Sharpless and H. P. Jensen, Znorg. Chem., 1974, 13, 2617; V. Katovic, L. T. Taylor, C. K. Chang and T. G. Traylor, Proc. Nut. Acad. Sci. U.S.A., 1973, 70, 2647; J . Amer. and D. H. Busch, J . Amer. Chem. SOC., 1969,91,2122. Chem. SOC., 1973,95, 5810, 8475, 8477. R. G. Wilkins, in ref. 11, p. 11 1 ; M. J. Carter, D. P. Rillema, and F. Basolo, J . Amer. Chem. SOC., 1974, 96, 392 and references therein. 88 J. H. Wang, A. Nakahara, and E. B. Fleischer, J . Amer. Chem. SOC., 1958,80, 1109; J. P. Collman and C. A. Reed, ibid., 1973,95,2048. 83 G. C. Wagner and R. J. Kassner, J. Amer. Chem. SOC., 1974,96,5593. 81 W. S. Brinigar and C. K. Chang, J .Amer. Chem. SOC., 1974,96,5595. 8 5 D. L. Anderson, C. J. Weschler, and F. Basolo, J. Amer. Chem. SOC., 1974,96,5599. 86 J. Almog, J. E. Baldwin, R. L. Dyer, J. Huff, and C. J. Wilkerson, J . Amer. Chem. SOC., 1974, 96, 5600. J. E. Baldwin and J. Huff, J. Amer. Chem. SOC., 1973,95,5757 and references therein. 438 FeII + 0 2 + FeIQ + FeII + Fe1IO--0Fe1I --t FeIII Lindoy followed by a rapid bimolecular redox process which eventually leads irreversibly to an iron(rr1) complex : If such is the case then autoxidation should be suppressed by constructing a bulky organic ligand which would effectively sterically prevent the formation of the intermediate iron@) oxygen-bridged ~pecies;8~ the situation might then be similar to that in haemoglobin in which the haem groups occupy molecular cavities in the globin portion of the molecule.88 The syntheses of iron complexes of ligands which are able to enclose a bound 0 2 molecule in a cavity of organic material have been achieved re~ently.~~@J’o Following the initial synthesis of (42)S7, which reversibly binds oxygen at low temperatures and which contains a 0 cavity (with two open sides), the iron(I1) complex of the ‘picket fence’ ligand (43) was prepared.89 This synthetic porphyrin derivative undergoes reversible oxygenation at room temperature in the presence of various bases; moreover the molecular oxygen adduct (with N-methylimidazole as the axial ligand) has been characterized by an X-ray structural analysis.g1 Like oxyhaemoglobin this complex is diamagnetic.The iron(@ lies in the plane of the porphyrin ring and the iron-oxygen bond is bent in a manner previously predicted for oxyhaemo- g l ~ b i n . ~ ~ There also appears to be substantial double-bond character in the iron-oxygen b0nd.~l~~3 Based in part on the results of the studies discussed above, a second iron(@ M. F. Perutz, Nature, 1970,228,726; J. C . Kendrew, Science, 1963,139,1259. 89 J. P. Collman, R. R. Gagne, T. R. Halbert, J. Marchon, and C. A. Reed, J. Amer. Chem. Sac., 1973,95,7868; J. P. Collman, R. R. Gagne, and C. A. Reed, ibid, 1974,96,2629. 90 J. Almog, J. E. Baldwin, R. L. Dyer, and M. Peters, J . Amer. Chem. SOC., 1975, 97, 226; J. Almog, J. E. Baldwin, and J. Huff, ibid., 1975,97,227. 1)1 J. P. Collman, R. R. Gagne, C. A. Reed, W. T. Robinson, and G. A. Rodley, Proc. Nar. Awd. Sci. U.S.A., 1974, 71, 1326. O 2 L. Pauling, Stanford Med. Bull., 1948, 6, 215; Nature, 1964,203, 182. 93 J. P. Collman and R. R. Gagne, J . Amer. Chem. SOC., 1974, 96, 6524. 439 Transition Metal Complexes o f Synthetic Macrocyclic Llgarufs CH3 I H3c \ c / CH3 CH3 H3C I CH3 WC’ I co complex has been synthesizedgO which also binds molecular oxygen at room temperature. This new complex contains the synthetic ‘capped’ porphyrin (44) P 0-c 440 Lidby and is so designed that the organic cavity above the porphyrin plane is large enough to admit molecular oxygen but will not allow entry of other small molecules such as those of the solvent. In the presence of heterocyclic base, an oxygen adduct is readily formed. The stability towards oxidation of the oxy- complex depends markedly on the type and concentration of the heterocyclic base present. The rate of autoxidation is reduced under conditions which favour co-ordination of base in the axial site opposite the ‘capped’ side of the molecule. It appears that when this axial site is unoccupied, attack by molecular oxygen can also occur at this then unprotected position. Indeed, if no base is present in the reaction solution then autoxidation occurs rapidly on exposure to oxygen. Clearly, if the axial base could be held in position by a covalent linkage of the type found in (41) then the resulting complex should be a very efficient oxygen carrier and one that, in essence, contains all the basic structural attributes of the haemoglobin active site. 441

ISSN:0306-0012    

DOI:10.1039/CS9750400421

出版商:RSC    

年代:1975

数据来源: RSC

摘要:

Co-ordination Chemistry of Aryldiazonium Cations : Aryldiazenato (Arylazo) Complexes of Transition Metals, and the Aryldiazenato-Nitrosyl Analogy By D. Sutton DEPARTMENT OF CHEMISTRY, SIMON FRASER UNIVERSITY, BURNABY, B . C . V5A 1S6 CANADA 1 Introduction During the past ten years there has been a growing interest in the synthesis and structure elucidation of transition-metal compounds containing the aryldiazenato- ligand (Arm). The first reported compound, by King and Bisnette in 1964,l was [Mo(CgHg) (CO)2(N2C6H40Me-p)]. Its synthesis utilized displacement of a carbonyl group in [Mo(CgH5) (CO)3]- by the aryldiazonium ion. Subsequently, reactions involving aryldiazonium ions have become an important general route to aryldiazenato-compounds. An immediately obvious analogy exists between this compound and the corresponding nitrosyl IMo(CgH5) (CO)2NO].In each compound an 1 8-electron molybdenum configuration requires nitrosyl and aryldiazenato-ligands to be three-electron donors, and the pursuit of this analogy between the formally isoelectronic nitrosyl and aryldiazenato-ligands has underscored much of the subsequent development. Fortuitously, this has also coincided with certain renewed interest in nitrosyl complexes themselves, when structural evidence showed in 1968 that a formal distinction could be drawn between the familiar three-electron donor nitrosyl ligands with near-linear MNO geometry and a second group of nitrosyl complexes which are characterized by a bent (ca. 120") MNO geometry. The latter correspond to a description of the nitrosyl group acting as a one-electron donor, as in organic nitroso-compounds.2 Similar geometrical differences in aryldiazenato-compounds would be expected and have been found.3 This convergence of structural results on aryldiazenato and nitrosyl compounds, together with the considerable number of new aryldia- zenato-compounds reported in the past ten years, makes this an opportune time for a comparative review of these complexes.A further point of interest and potential utility is the close relationship of the R. B. King and M. B. Bisnette, (a) J . Amer. Chem. SOC., 1964, 86, 5694; (b) Inorg. Chem., 1966, 5, 300. a D. J. Hodgson, N. C. Payne, J. A. McGinnety, R. G. Pearson, and J. A. Ibers, J.Amer. Chem. SOC., 1968, 90, 4486; D. J. Hodgson and J. A. Ibers, Inorg. Chem., 1968, 7, 2345; B. A. Frenz and J. A. Ibers, in 'Chemical Crystallography', ed. J. M. Robertson, MTP International Review of Science, Physical Chemistry, Series One, Vol. 11, Butterworths, London, 1972, p.33. 3 (a) A. P. Gaughan, B. L. Haymore, J. A. Ibers, W. H. Myers, T. E. Nappier, and D. W. Meek, J. Amer. Chem. Soc., 1973, 95, 6859; (b) A. P. Gaughan and J. A. Ibers, Inorg. Chem., 1975,14, 352. 443 (1) A second reaction, and clearly one of considerable consequence in regard to nitrogen fixation, is protonation to yield a co-ordinated N2H2 moiety5 [equation Co-ordination Chemistry of Aryldiazonium Cations aryldiazenato-ligand with dinitrogen. It presently seems that a continued study of the synthesis and properties of aryldiazenato-complexes may prove worthy in understanding the activation of dinitrogen by transition metals and the role of iron and molybdenum in biological nitrogen fixation.The following examples will illustrate. Generally, the co-ordinated dinitrogen ligand has been found to be hardly at all more reactive than molecular dinitrogen itself. One of the few reactions which it has been observed to undergo, in appropriate complexes with Mo, W, or Re, is attack by electrophilic alkyl, aroyl, or acyl groups to form diazenato-complexes as in equation (l).4 IReCl(N2) (PMezPh)4) RCOCl + IReCh(N2COR) (PMe2Ph)aI PMe,Ph (R = Ph or Me) HX trans-[M(N&(diphos)z] -+ [MX2(N2H2) (diphos)z] (M = Mo or W; X = CI or Br) (2) A recent X-ray structure determinations indicates that in [MX(NzHz)(diphos)z] BPh4 the ligand has the hydrazido(2-) structure M-N-NH2, in which case an obvious and close comparison exists between the stereochemistry and electronic configuration of this arrangement and that of an aryldiazenato-ligand protonated on NZ.This seems to be confirmed by the X-ray structure7 of [ReCIz (NN(H)Ph 1 (PMe2Ph)2(NH3)]+, formed by the protonation of the aryldiazenato-ligand in the complex [ReC12(N2Ph) (PMe2Ph)2(NH3)], where the observed dimensions indicate the presence of the RCN-N(H)Ph group. Other compounds formed in the protonation of co-ordinated dinitrogen appear to be complexes of diazene (HN=NH), which has long been postulated as a possible intermediate in nitrogen fixation but never identified.Possibilities for the binding of this ligand include a a-complex (M-NH=NH) and a 7-complex (1). Consequently the binding of diazene itself and substituted diazenes to transition metals is receiving attention,* and an important route to the (a) J. Chatt, J. R. Dilworth, G. J. Leigh, and V. D. Gupta, J. Chem. SOC. (A), 1971, 2631 ; (6) J. Chatt, G. A. Heath, N. E. Hooper, and G. J. Leigh, J. Organometallic Chem., 1973, 57, C67; (c) A. A. Diamantis, J. Chatt, G. J. Leigh, and G. A. Heath, J. Orgunometallic Chem., 1975,84, C11. S J . Chatt, G. A. Heath, and R. L. Richards, J.C.S. Chem. Comm., 1972, 1010; J.C.S. m1. Dalton, 1974, 2074. 6 G. A. Heath, R.Mason, and K. M. Thomas, J. Amer. Chem. SOC., I974,96, 259. 7 R. Mason, K. M. Thomas, J. A. Zubieta, P. G. Douglas, A. R. Galbraith, and B. L. Shaw, J. Amer. Chem. SOC., 1974, 96, 260. 8 S. D. Ittel and J. A. Ibers, Znorg. Chem., 1973, 12, 2290; A. Nakamura, M. Aotake, and S. Otsuka, J. Amer. Chem. SOC., 1974, 96, 3456; P. W. Schneider, D. C. Bravard, J. W. McDonald, and W. E. Newton, ibid., 1972, 94, 8640. Sutton NH 2 Nomenclature NH latter complexes can be through reaction of aryldiazenato-complexes with elect rop hiles . The purpose of this article, then, is to review the synthetic and structural chemistry of aryldiazenato-complexes and their present distribution among the elements, so as to bring out parallels with the corresponding nitrosyls and their relationship with other complexes of nitrogen ligands.The most notable feature regarding the nomenclature currently in use in relation to the compounds described in this Review is its variety and inconsistency. This extends to the diazonium ions themselves, which are properly named by adding the suffix ‘diazonium’ to the name of the parent hydrocarbon, rather than that of the aryl gr0up.9 Nevertheless, the latter terminology is quite common, especially in journals of inorganic or organometallic chemistry. Products containing the M-NzAr group have been most often called arylazo complexes, though recently3 aryldiazo was introduced on the basis of its deriva- tion from aryldiazonium. However, neither name conforms with established rules of nomenclature.A survey of the names currently adopted for related species reveals the following: the molecule HNNH is still frequently called di-itnine6 or di-imide,lO even though general adoption of the name diazene has been urged;llJ2 similarly, PhNNH is variously called phenyldi-imide,lOJ3 phenyldi-imine,7J4 and phenyldiazene,3-11 and an identical name should be, and has been, used for PhNNH as a neutral ligand. As mentioned in the introduction, protonation of M-NN and M-NNPh may give rise to M-NNH2 and M-NN(H)Ph. These ligands have been described as the hydrazido dianion6 and phenylhyd- razido dianion7 respectively. First it must be decided whether different names should be adopted to dis- tinguish the two limiting structural possibilities represented by the linear M-N-N and the bent M-N-N systems which will be discussed in the following Section.It is a principle of chemical nomenclature that a compound must be named without recourse to a determination of its dominant electronic 9 I.U.P.A.C. Nomenclature of Organic Chemistry, Section C, Butterworths, London, 1965, p. 216. lo E. K. Jackson, G. W. Parshall, and R. W. F. Hardy, J. Biol. Chem., 1968, 243, 4952. 11 E. M. Kosower, Accounts Chem. Res., 1971,4, 193. 1% J. H. Fletcher, 0. C. Dermer, and R. B. Fox, ‘Nomenclature of Organic Compounds’, Adv. Chem. Ser., 1974, NO. 126, p. 247. 13 G. W. Parshall, J. Amer. Chem. SOC., 1967, 89, 1822. 14 K. R. Laing, S. D. Robinson, and M. F. Uttley, J.C.S. Dalton, 1973, 2713.445 Co-ordination Chemistry of Aryldiuzonium Cations structure; these systems are essentially different electronic structures and so need not be distinguished by name. Secondly we must decide the systematic name for the ArNN ligand. The names diazene (HN=NH) and aryZdiazene (ArN= NH), recommended12 by the American Chemical Society, will eventually be universally endorsed. According to proposed I.U.P.A.C. rules, the ArNN ligand then logically becomes aryldiazenato (not aryZdiazenidol6) and this name will be used throughout this Review. Other related groups are usually named according to I.U.P.A.C. rules. The nitrogen atoms will be numbered as Ar-W=N1-M. N M 3 Electronic Structure and Geometry A. Three-electron Donor (N&+) Terminal Ligand.-By exact analogy with terminal, so-called NO+, nitrosyl complexes there is a group of aryldiazenato- complexes in which the aryldiazenato-ligand may be viewed formally as a three- electron donor terminal ligand or, equivalently, as the aryldiazonium ion (N2Ar+) co-ordinated through the d o n e pair on the terminal nitrogen together with strong back donation of electrons from the metal.A conventional, simplified picture of the electronic structure is shown in Scheme l(a). It should be noted that, in its usual interpretation, this picture does not imply any change of hybridization at oxygen. However, the simple valence bond model of back-bonding [Scheme l(b)] generates a second lone pair on Scheme 1 oxygen and implies a hybridization change at oxygen from sp to sp2.Unfortu- nately, there are no possible observable geometric changes at oxygen capable of reflecting the adequacy of either picture. The situation in aryldiazenato-complexes is different in this respect. Now, the mere population of n*-orbitals on N2Ar leaves the angle at N2 close to 180°, [Scheme 2(a)] whereas the simple valence-bond picture in Scheme 2(b) predicts this angle to be about 120". In the majority of presently available X-ray structure determinations on formally N2Ar+ complexes the angle at N2 is indeed near to 120" (see the Table), attesting to the importance of back-bonding in these complexes and to the S. D. Ittel and J. A. Ibers, J. Amer. Chem. SOC., 1974, 96, 4804. 446 Ns-C/II(. M-N1/%I N1-N8/II(.M-N1-Ns/" N1-N*-C/" Ref. Table Dimensions of aryldiazenato-complexes from X-ray diflraction studies Compound Group I Terminal 3e--donor @I&+) 1.432(7) 16 [ {HB(PZ)3 )Mo(C0)2(N2Ph) 1 [ReCN2Ph)C12(PMe2Ph)3] 1 21 . 1 (2) 115 17 7 18, 19 [FeNPh) (COhPPh3)2 1 + [Ru(p-NaCsH4Me)C13(3)2] 1.21 l(6) 1.23 1.23(2) 1.201 (7) 1.14(1) 174.2(1) 172 173(2) 179.2(5) 171.2(9) 171.9(5) 1.825(4) 1 .so 1.77(2) 1.702(6) 1.796(9) 1.784(5) 1.74(2) [ReCh(N2COPh) (PMe2Ph)sI 1.43(2) 1.404( 8) 1.40(1) 1.376(6) 1.42(3) 170(2) 1 19(2) 1 %.2(6) 136(1) 137.1(5) 124(2) 19 7 20 36 1.1 5 8( 6) 1.22(3) 1.172(9) 1.961(7) 1.445( 1 1) Group 2 Terminal le--donor (NzAr-) [Rh(NzPh)CI(PPP)] +* 118.9(8) 125.1(6) Group 3 Bridging 3e--donor [MnCNzPh) (COh 12 1.234(3) 1.445(3) 119.7(2) 35 134.4(2) 119.6(2) *PPP = PhP(CH2CH2CH2PPh2)z 2.03 l(2) 2.021(2) * PPP = PhP(CHzCH2CHzPPh,)2 la G.Avitabile, P. Ganis, and M. Nemiroff, Acta Cr-vst., 1971, B27, 725. '7 V. F. Duckworth. P. G. Douglas, R. Mason, and 9. L. Shaw, Chem. Comm., 1970, 1083. W. E. Carroll and F. J. Lalor, J.C.S. Dalton, 1973, 1754. J. A. Ibers and R. L. Haymore, Inorg. Chem., 1975, 14, 1369. IoJ. V. McArdle, A. J. Schultz, B. J. Corden, and R. Ejsenberg, Inorg. Chem., 1973, 12, 1676. + N -Ar --+ M-N=N' \ Ar Scheme 2 qualitative usefulness of the simple VB description. Understandably then, it is this nitrogen atom which is the site for protonation, as indicated by the X-ray structure determination' of [ReCh {NN(H)Ph } (PMe2Ph)2(NH3)]+.The extent to which other N2Ar+ complexes may undergo protonation or reaction with other electrophiles at N2 has at the time of writing received little attention, but this is expected to become an increasingly important aspect of their chemistry. Of course, examples of NNC angles intermediate between 180" and 120" may be expected depending upon the relative importance of the 'linear' and 'singly bent' formalisms (2) and (3) to the electronic structure. In this connection, the + - af + M + : N s N - A r + M-NENAr 136" angle in the listed ruthenium complex is not unexpected. It may be possible eventually to correlate vibrational frequencies with the intra-ligand angles, but at present there are insufficient structural as well as spectral data to allow this in detail.At this point it is worthwhile examining other ligands isoelectronic with N2Ar+ to see whether these display similar geometric changes upon co-ordina- t ion. Examples are cyanides R-CZN,~~ isocyanides R-NEC,~~ and acetylides Co-ordination Chemistry of Aryldiazonium Cations N (b) - R-C-C. There appear to be no recorded instances of a corresponding geometry change of this magnitude for these ligands. Especially interesting is isocyanide, where the nitrogen carries a formal positive charge as in diazonium. The 81 M. Kilner, Adv. Organometallic Chem., 1972, 10, 115. 19 P. M. Treichel, Adv. Organometallic Chem., 1973, 11, 21. 448 Sutton chemistry of this ligand points to it being a much better o-donor and a weaker w-acceptor than CO.On the basis of available X-ray structure determinations, of which some recent examples are cited,23-26 terminal isocyanide complexes exhibit angles close to 180" at nitrogen, presumably reflecting the importance of a-donor rather than ?r-acceptor capacity in its binding to metals in the presently determined examples. B. Ose-electron Donor (NzAr-) Terminal Ligand.-The conventional occurrence of the nitrosyl (nitroso) group or the aryldiazenyl (arylazo) group in carbon chemistry is where these contribute one electron to a two-electron, two-centre o-bond in R-NO or R-N2Ar compounds. These ligands can form analogous bonds to transition metals, and in such complexes they can be viewed formally as one-electron donors and formulated as NO- and N2Ar- ligands.An important difference from the previously described NO+ and N2Ar+ complexes is the expected change in geometry. As a one-electron ligand the nitrosyl group should give rise to an angle of about 120" at nitrogen (4); the corresponding aryldi- azenato-complexes should possess 120" angles at both nitrogen atoms, i.e. be 'doubly bent' and capable of trans- and cis-geometric isomers ( 5 ) and (6). .. .. The ions NO+ and N2Ar+ do, of course, exist and may be utilized directly in the synthesis of many of their complexes, lending at least some degree of reality to their formulation as complexes of the positive ions. This is not true, to any extent, of the ions NO- and NzAr-, and, as yet, NO- and N2Ar- complexes have not been synthesized by direct co-ordination of these anions, nor by reactions involving reagents likely to give rise to these ions in situ.Instead, present methods yield these compounds indirectly, an important route being the formal transfer of two electrons from the metal to the ligand, either as a redistri- bution of electrons in the formal NO+ or N2Ar+ complexes or by electrophilic attack of NO+ or N2Ar+ at an electron-rich metal atom. As an example of the former, nitrosyl complexes of the type [Co(NO)(PR&C12] are believede7 to exhibit structural isomerism between a trigonal-bipyramidal form having a linear equatorial nitrosyl group [CoI(NO+)], and a square-pyramidal form having 23 C.J. Gilmore, S. F. Watkins, and P. Woodward, J. Chem. SOC. (A), 1969, 2833. a4 B. JovanoviC, L. Manojlovit-Muir, and K. W. Muir, J.C.S. Dalton, 1972, 1178. s b D. Baumann, H. Endres, H. J. Keller, and J. Weiss, J.C.S. Chem. Comm., 1973, 853. *c. R. D. Adams, F. A. Cotton, and J. M. Troup, Znorg. Chem., 1974,13,257. $' C. P. Brock, J. P. Collman, G. Dolcetti, P. H. Famham, J. A. Ibers, J. E. Lester, and C. A. Reed, Znorg. Chem., 1973, 12, 1304. 5 449 Co-ordination Chemistry of A ryldiazonium Cations a bent apical nitrosyl group, [Co1Ir(NO-)]. Other examples have been cited.28 The relationship between the limiting forms in this electron redistribution is simply indicated in Scheme 3, and an analogous representation for aryldiazenato- complexes is shown in Scheme 4.Scheme 3 2+ 2+ M z N - N ' M z N - N ' - O M " O M " + n + n - \ Ar Ar \\N \\N - Ar Ar Scheme 4 The synthesis of NO- and N2Ar- complexes by reactions of nitrosonium or aryldiazonium salts can similarly be represented, albeit in a somewhat arbitrary and cumbersome manner, in terms of Schemes 3 and 4, in which it is envisaged that Lewis-base co-ordination by the cation occurs first, followed by an extreme back-bonding process, going beyond (using the aryldiazenato example) the singly bent species to a complete transfer of two electrons from the metal to give the doubly bent ligand. A simpler, and chemically more appealing, viewpoint is that the reaction proceeds directly to the doubly bent ligand by virtue of the N2Ar+ group acting as an electrophile (i.e.Lewis acid co-ordination) as depicted in Scheme 5. In support of this, the metals which undergo such reactions Scheme 5 are typically those for which oxidative addition and other two-electron oxidation processes are important, for example d8 complexes of CoI, RhI, and IrI, and d1O complexes of Pto. A particularly well documented example, which also illustrates the nitrosyl-aryldiazenato analogy very well, is the co-ordination of the dis- sociation-stabilized rhodium(1) phosphine complex [Rh(PPP)Cl] (PPP = ph2P(CH2)3]2PPh} by NO+ and N2Ph+ to give structurally similar square- pyramidal complexes containing apical bent nitrosyl (4) and doubly bent a* J. H. Enemark and R. D. Feltham, Proc.Nar. Acad. Sci. U.S.A., 1972, 69, 3534; Co- ordination Chem. Rev., 1974, 13, 339. 450 Sutton aryldiazenato (5) ligands respectively.3J?9 As can be seen from the Table, the M-N distance in the aryldiazenato-complex (12) (see p. 456) is considerably longer than that of any singly bent aryldiazenato-complex, in keeping with the oneelectron donor formalism and with observed lengths in linear and bent nitrosyl complexes.2 Of the two possible sites for protonation of the doubly bent aryldiazenato- ligand, the presently available evidence, primarily from the magnitude of 15N-’H coupling constant^,^^ is that protonation occurs at N1, the nitrogen atom adjacent to the metal, as for example equation (3).14 No example is presently known of protonation at N2 or at both sites.Protonation therefore yields complexes containing co-ordinated aryldiazenes (NH=NAr) such as phenyldiazene, molecules that are quite unstable in the uncomplexed state.11 HCl [RhC12(N2Ar) (PPh3)el f=lc [RhCb(NH=NAr) (PPh3)21 Et,N (3) Here again, the analogy with nitrosyl complexes is illustrated by equation (4), in which protonation of co-ordinated nitrosyl occurs at the site of co-ordination to give a complex containing co-ordinated HN0.30 HCI [OSCKCO) (NO) (PPh3)2] # [OSC~~(CO) (NH----O) (PPh3)2] - HCI [Co(das)@O)Br]+ --+ [Co(das)2(NH=O)Brl2+ of co. (k) (4b) The stereochemistry of [OsCl(CO) (NO) (PPh3)2] has not been determined by X-rays, but it may be expected that the reaction involves an isomer having the bent Os-N-0 form, rather than the linear.The closely related isoelectronic complex [Os(OH)(N0)2(PPh&] + possesses both linear and bent nitrosyl groups,31 where the linear NO+ group may be considered to be the counterpart As a final comment in this section, the question is raised as to what extent it is realistic, in view of their positive charge, to consider that NO+ and N2Ar+ are actually ever capable of attacking metal complexes as Lewis bases, for example in the displacement of CO and the formation of NO+ and N2Ar+ complexes in equations (5) and (6).1*932133 (a) T. E. Nappier, D. W. Meek, R. M. Kirchner, and J. A. Ibers, J. Amer. Chem. SOC., 1973, 95,4194; (b) T. E. Nappier and D. W. Meek, ibid., 1972, 94, 306. 8o (a) K.R. Grundy, C. A. Reed, and W. R. Roper, Chem. Comm., 1970, 1501; (b) J. H. Enemark, R. D. Feltham, J. Ricker-Nappier, and K. F. Bizot, Inorg. Chem., 1975,14,624. a1 J. M. Waters and K. R. Whittle, Chem. Comm., 1971, 518; G. R. Clark, J. M. Walters, and K. R. Whittle, J.C.S. Dalton, 1975,463. *a (a) B. F. G. Johnson and J. A. Segal, J.C.S. Dalton, 1972, 1268; (b) S. Cenini, F. Porta, and M. Pizzotti, Inorg. Nuclear Chem. Letters, 1974, 10, 983. 9s D. R. Fisher and D. Sutton, Canad. J. Chem., 1973, 51, 1697. 451 Cu-ordination Chemistry of Aryldiazonium Cations An alternative view which merits greater attention is that attack by NO+, NzAr+, and similar cationic species may always be electrophilic at the electron- rich low-valent metal atom, and that displacement of CO, PPh3, etc.occurs by virtue of the competition for the metal's d-electrons between the incumbent ligands and the electrophile. This is illustrated, for a metal carbonyl residue, in simple valencebond terms in Scheme 6. The entering NO+ attacks an electron +O N -C C I II M - M +O +O 0 Ar ni C M C II 0 II pair in a non-bonding d-orbital to give (7), causing no formal change in the electron count for M. By virtue of the increased positive charge on M, it will now be less capable of back-bonding to CO, leading to an electronic arrangement more in favour of (8). Competition then arises between the capability of NO and CO to form multiple-bonds with M. Where this lies in favour of CO, the NO+ may either be expelled or be retained in the bent geometry of the NO- formalism, as occursz in the reaction of FCl(C0) (PPh3)2] with NO+.Where this lies in favour of NO, building of the N=M multiple bond from the nitrogen lone pair will weaken and expedite cleavage of the M-C bond. At the same time this process replaces the electrons being lost with the departure of the CO ligand, main- taining the electron count of the metal. The M(NO+) formalism for the complex results and the geometry is linear. C. Bridging NzAr.-The aryldiazenato-ligand may bridge two metal atoms in structures such as (9) and (10) and in both cases it supplies three electrons to the " . /Ar N I1 111 M ' Ill / \ IN=N M2 452 c 0 Sutton system. Bridging nitrosyls bound in a manner comparable to (9) are well-known,w and this is the geometry of the single well established example35 of a bridging aryldiazenato-ligand in the complex [Mn(N2Ph)(CO)& (24) (see p.463). The bonds from N1 to the two manganese atoms are almost identical (Table) indicating that the three electrons are shared equally by them. As expected, they are also somewhat longer than in the terminal aryldiazenato-complexes. (ref. 69) (8) 4 Synthesis A major route to nitrosyl complexes is by reaction with NO itself, and here an immediate difference in the case of aryldiazenato-complexes becomes apparent : presently ArN2. is not a readily accessible reagent on account of its instability toward dissociation to Ar + N2. As a consequence, aryldiazenato-complexes have been obtained most frequently from diazonium ions, paralleling a number of recent applications of NO+ to the synthesis of nitrosyl complexes.~~36 However, new reagents such as MesSiNzPh promise to be important alternatives for the introduction of ArN2 ligands.35 A.Substitution by Diazonium Ions.-A metal complex frequently reacts directly with an aryldiazonium salt, the product being formed by the displacement of another ligand. Generally, an 18-electron configuration is maintained by the replacement of a two-electron donor ligand such as CO or PR3 by ArN2+. Such a process is favoured by attack at an electron-rich metal site such as a substituted metal carbonyl anion, as in equations (7) and (8), but can also occur with a neutral substituted carbonyl.C 0 (RB(pz)3 )MO(CO)QI- + [(RB(pz)3 )Mo(CO)~(N~A~)I + CO &N,+ DMF For example, [Fe(C0)3(PPh3)2] readily undergoes replacement of one carbonyl group by either NO+ or N2Ar+ to form analogous cationic complexes [Fe(NO)- (CO)z(PPh3)2]+ and [Few&) (CO)z(PPh3)2]+ which are trigonal bipyramids with apical phosphines [equation (9)]. These are formally complexes of iron(0) with NO+ or N2Ar+, on the basis of i.r. spectroscopy [v(NN) z 1730 cm-11, and a structure determination shows a singly bent FeNzAr grOuplg (Table). Both 34 N. G. Connelly, Inorg. Chim. Acta Rev., 1972, 6, 47. 36 E. W. Abel, C. A. Burton, M. R. Churchill, and K.-K.G. Lin, J.C.S. Chem. Comm., 1974, 268. s6 N. G. Connelly and J.D. Davies, J . Organometallic Chem., 1972, 38, 385. 453 Co-ordination Chemistry of Aryldiazonium Cations compounds may be converted into the neutral five-co-ordinate complexes [Fe(NO)X(CO)(PPh3)2] and [F~(NZA~)X(CO)(PP~S)Z] by IphSP=N=PPhs]+ X-QC = CI or Br), but differ in the inertness of the aryldiazenato-complexes towards substitution of CO by PPh3, compared with the nitrosyl complexes.37 It is also possible for the diazonium ion to react with, and abstract, the co- ordinated phosphine group from [Fe(CO)3(PPh3)2] with the formation of highly coloured [RsP-N=NAr]+ cations, which have varying stability depend- ing upon the substituents on phosphorus or the aromatic nucleus.3* Where these cations are unstable and decompose to dinitrogen and a quaternary phosphonium salt, the erroneous impression may be obtained of the formation of transient transition-metal aryldiazenato-complexes, whereas we have found that addition of further diazonium salt may frequently yield isolable derivatives.89 It is appropriate at this point to examine how successful have been the attempted syntheses of other aryldiazenato-complexes of the 3d elements.The binary carbonyls [Fe(CO)5], [CO~(CO)S], and mi(CO)4] react with diazonium salts with vigorous evolution of N2 and CO, whereas [Cr(CO)s] {and [Mo(CO)s]) show little reaction.40~41 No derivatives could be obtained from [CpFe(CO)2]-, ph3SnFe(C0)4]-, [Co {P(OMe)3}4]-, or [CpNiCO]-.1bs42-44 The zwitterion (1 1) C 0 (11) M=Cr or Mo 87 W. E. Carroll, F.A. Deeney, and F. J. Lalor, J.C.S. Dalron, 1974, 1430. 38 J. A. Carroll, D. R. Fisher, G. W. Rayner Canham, and D. Sutton, Cunad. J. Chem., 1974,52, 1914. N. Farrell and D. Sutton, unpublished results. 40 G. N. Schrauzer, Chem. Ber., 1961,94, 1891. 41 J. C. Clark and R. C. Cookson, J . Chem. Soc., 1962, 686. W. E. Carroll and F. J. Lalor, J. Organometallic Chem., 1973, 54, C37. A. N. Nesmeyanov, Y. A. Chapovskii, and L. G. Makarova, Izvest. Akad. Nauk S.S.S.R., Ser. khim., 1965, 1310. a E. L. Muetterties and F. 1. Hirsekorn, J.C.S. Chem. Comm., 1973, 683. 454 i, EtOH-ArCON, [IrCl(NzAr) (PPh3)z]+PF6- pCl(C0) (PPh3)2] ii, ArN,+PFI- Sutton yields an aryldiazenato-derivative by replacement of one CO group in the case of Mo but not Cr.45 The first aryldiazenato-complexes of chromium have been syn- thesized by reaction of the compounds [(CSH~)C~(CO)SI- or [(CaMe~)Cr(C0)3] with a variety of aryldiazonium salts.46 The latter yields [(CaMes)Cr(NzAr)- (CO)z]+ salts, which undergo nucleophilic attack by H- at the arene ring to yield neutral complexes [(CsMesH)Cr(N2Ar>(co)o)z] in apparent preference to attack at the aryldiazenato-ligand.Surprisingly, in view of the iron results already described, the p hosp hine-subs t i t u t ed complex [( CS Mes)Cr( CO)2(PPh)] is not substituted by ArNz+, but is instead oxidized to the monocation.46a Thus far, attempts to secure isolable aryldiazenato-derivatives of unsubstituted metal carbonyls or metal carbonyl anions have been notably unsuccessful.King1 failed to obtain derivatives from v(Co)S]-, [Mn(C0)5]-, and [Co(CO)4]-, and attributed this to labilization of metal-carbon monoxide bonds by the presence of the good electron-withdrawing aryldiazenato-group. More recently, Carroll and Lalor18 confirmed that at - 70 "C reaction does occur with thew anions to give highly coloured species which may be aryldiazenato-derivatives, but that these do not survive warming to room temperature. In the case of [Co(CO)4]-, but not the others, they were able to isolate an unstable triphenylphosphine derivative, believed to be [Co(NzPh)(CO)zPPhs] from i.r. spectroscopy. The stable complex [Fe(N2Ph)(NO)(CO)PPh3] was obtained similarly from [Fe(NO)(CO)3]-. The manganese complex [Mn(N2Ph)(CO)&, containing bridging NzPh ligands,s presently stands as the sole example of an unsubstituted metal carbonyl aryldiazenato-complex, and its synthesis is described in Section 4F.As a matter of contrast, the CO group in the 16electron Vaska's complex is tightly held, and its reactions with diazonium ions (to be described in Section 5C) give a variety of products, in none of which has the CO group been displaced. However, by allowing the carbonyl to react with an aroyl azide followed im- mediately by the addition of an aryldiazonium salt,47 the CO group is effectively substituted by ArN2+, almost certainly owing to the formation of the more labile dinitrogen complex as an intermediate [equation (lo)]. (10) The stoichiometry, synthesis, and high v(") value of 1868 cm-l of this green salt all point to a square-planar IrI complex possessing a singly bent N2Ar+ ligand.Like its carbonyl and nitrosyl analogues, [IrCl(N2Ar) (PPh3)zI + displays numerous reactions. Addition of neutral ligands L, such as COY PR3, AsR3, SbRs, and RNC, occurs to form five-co-ordinate cationic species. These com- pounds may alternatively be synthesized from IrCI(PPh3)zL and ArNz+. From experience with related compounds, the geometries of these compounds are 46 D. Cashman and F. J. Lalor, (a) J. OrgunometaZZic Chem., 1970, 24, C29; (b) ibid., 1971, 32, 351. 46 (a) N. G. Connelly and Z. Demidowicz, J. OrgunometaZZic Chem., 1974, 73, C31; (b) '' B. L. Haymore and J. A. Ibers, J. Amer. Chem. Soc., 1973, 95, 3052.M. Heberhold and W. Bernhagen, 2. Naturforsch., 1974, 29b, 801. 455 Co-ordination Chemistry of Aryldiazonium Cations postulated to be trigonal bipyramids (with singly bent N2A.r) for the tris-(ERs) complexes (E = P, As, or Sb) and square pyramids (with apical doubly bent N2Ar) for the bisphosphine carbonyl or isocyanide. An example14 of the displacement of phosphine, accompanied by co-ordination of halide to give the preferred six-co-ordinate complex,20,48 is given in equation PPh3 EtOH-acetone (1 1). (1211. t ArN.+CI The X-ray structure20of the product indicates it to be a complex of ruthenium@) (T.able) involving no change in ruthenium oxidation state. The reaction therefore belongs to the substitution type, and not to the class of oxidative-addition reac- tions now to be described.B. Oxidative Addition of Diazonium Ions.-There does not appear, as yet, to be any established instance of the simple addition of the diazonium ion (i.e. as a Lewis base) to a transition-metal complex without either ligand displacement or an increase in the oxidation state of the metal. An inspection of all known in- stances of simple addition indicates to this author that they fall into the general category of Lewis-acid addition, i.e. a reaction of the type illustrated in Schemes 4 or 5 is implied and the product is of the N2Ar- complex type.* In only one instance is this presently established by X-ray diffraction, but the chemical, spectroscopic, and structural properties of the remainder, especially where nitrosyl analogues of known structure exist, point strongly to this conclusion.The structurally established example3 is (12), the product of diazonium ion addition to the dissociation-stabilized phosphine analogue (1 3) of Wilkinson’s compound [RhCl(PPh3)3]. Although no analogous compound has been isolated from [RhCl(PPh3)3] itself, probably owing in part to dissociation of a phosphine ligand, simultaneous addition of C1- yields [Rh(NzAr)Ch(PPh3)2], which must be expected to be a complex of rhodium(II1) also.l4 These are examples of the oxi- dative-addition conversion of d8 complexes into d6; the reaction is also exhibited in the conversion of d10 complexes into square-planar d8 species. Cenini et aZ.,49 have described the reaction of a series of p-substituted phenyldiazonium salts with [ Pt (PPh3)3] to yield the plat hum(@ complexes [ Pt (N2CsH4X)(PPh3)3]Y (X = p-NO2, F, H, OMe, Me, NEt2, or NMe2; Y = BF4 or BPh) [equation *Note added in proof: The crystal structure of [IrCl(NPh)(PPh,Me),]+ synthesized from [IrCl(PPh,Me),] and PhN,+ indicates a Ir1(N8Ph+) structure.(B. J. Haymore and J. A. Ibers, personal communication.) Q8 (a) J. A. McCleverty and R. N. Whiteley, Chem. Comm., 1971, 1159; (b) J. A. McCleverty, D. Seddon, and R. N. Whiteley, J.C.S. Dalton, 1975, 839. 4 s S. Cenini, R. Ugo, and G. LaMonica, J. Chem. SOC. (A), 1971, 3441. 456 Sutton Ph2P C1 [ptOpPh3)31 (12) - EtaO-E ArN,+ tOH [PtWzAr) (PPh3)3]+ (1 2) Partial replacement of PPh3 from these complexes can be accomplished to yield some examples of the neutral complexes [Pt(N2C&X) (PPh3)2Z] (Z = N3 or I).These complexes are presumably square planar, with a doubly bent aryldia- zenato-ligand. Similar derivatives could not be isolated with [Pt(PPh3)2L], where L is CH2=CH2 or CH2=CHCN. The cations react with H2 at 1 atm and room temperature in the absence of catalyst to evolve N2 and give [PtH(PPh&]+ BF4-, and no intermediate aryldiazene or hydrazine complexes were observed. They can, however, be reversibly protonated by non-co-ordinating acids to yield the aryldiazene complexes (Pt(NH=NCe&x) (PPh3)3]2+ (BF4)2. These compounds are analogues of the platinum triethylphosphine complexes synthesized by Parshallso and by Garner and using insertion of ArN2+ into [ptHL(PEt&} complexes as described in Section 4C.It is therefore surprising that these neutral and cationic aryldiazenato-complexes are all considered to have v(NN) at about 1580 cm-1, whereas it has been verified by 15N substitution13 to occur at 1463 cm-1 in the neutral complex [PtCI(N2C6&F) (PEt3)2] and appears to occw at about 1580 cm-1 in the corresponding cationic triethylphosphine c0mplexes,5~ though this has not been confirmed by 15N substitution. A band at ca. 1150 cm-1 has also been assigned to v(”) in the triphenylphosphine aryldia- zene c0mplexes,~9 but this appears to be unusually low and the assignment must await additional confirmation. Cook and Jauhals2 have used potentially chelating ligands to displace co-ordinated ethylene from [Pt(PPh3)2(CzH4)]; for example benzenediazonium-1 ,Zcarboxylate (14) reacts to form the PtII complex (1 5).Several other examples of oxidative addition of diazonium ions to noble metal compounds will be found in Section 5. 60 G. W. Parshall, J . Amer. Chem. SOC., 1965, 87, 2133. 61 A. W. B. Gamer and M. J. Mays, J. Organometallic Chem., 1974,67, 153. C. D. Cook and G. S. Jauhal, J . Amer. Chem. SOC., 1968, 90, 1464; T. L. Gilchrist, F. J. Graveling, and C. W. Rees, Chem. Comm., 1968, 821. 457 0 PEt3 (1 3) Co-ordination Chemistry of Aryldiazonium Cations C. Apparent Insertion by Diazonium Ions.-This reaction was first recorded by Parshall,50 who showed that diazonium ion would react with trans- ptHCl(PEt3)2] to yield initially the product of 'insertion' of the diazonium ion into the Pt-H bond namely, the aryldiazene complex (10, which could be deprotonated to the aryldiazenato-complex (17) [equation (1 3)].The overall reaction is equivalent to PEt3 base ArNs+ __+ Cl-Pt-N2Ar Cl-Pt+-NH=NAr fPtHCl(F'Et3)3]+ EtOH 1 I PEt3 I I PEt 3 (1 6) (17) replacement of H- by ArN2- and involves no change in oxidation state of the metal. This reaction was subjected to detailed comparison with the enzymatic reduction of dinitrogen by nitrogenase, and led to important postulates as to the nature of the binding and reduction of dinitrogen in the enzyme system,10J3 The square-planar geometry of (16) has been established by X-rays,15 confirming the formal PtII oxidation state and, by implication, the N2Ar- nature of (17).Additionally the hydrogen atom is found to be located on N1, confirming the conclusion arrived at by Parshall on the basis of 1H n.m.r. spectroscopy using 15N at this position. By replacing C1 with a neutral ligand L such as NH3, pyridine (py), PEt3, and EtNC these compounds have been converted51 into the cationic aryldiazenato- complexes (18) and the dicationic aryldiazene complexes (19). The instability of (1 8) in favour of decomposition to [Pt(Ar)L(PEt3)2]+ by extruding dinitrogen appears to follow a sequence NH3 sz py < PEt3 z EtNC < CO, and in the case of CO only the a-aryl complex was obtained. The mechanism of this reaction is not known, but it clearly appears to be enhanced by the presence of - H+ L-Pt-NH PEt, I I %-Ar 1 2+ PEt, 458 Sutton an-acceptor ligand trans to the aryldiazenato-ligand. The X-ray results on the diazene complex (16) show little trans-influence on the part of either C1 or NH=NCe&F, and since A r m and ArN2- appear at present to be primarily a-bonding ligands, the absence of a trans-ligand high in the trans-influence series readily can be appreciated as a criterion for the stability of these platinurn complexes.The 19F n.m.r. of the p- and m-fluorophenyldiazenato-complexes have been! compared in efforts to estimate the importance of a and w effects in the bonding of the aryldiazenato-ligand to platinum, and the interpretation of the results has (ap - 8,) in the 19F chemical shifts of the neutral p- and m-fluoro-complexes been the subject of some controversy.ParshalP initially interpreted the difference @%cl(&cBfiF) (PEt3)2] to indicate a substantial contribution of resonance structures such as (20) to the electronic structure of the complexes. Garner and PEt, F O N - N = P P t + - C I I 1 PEt, Mays,S1 however, consider that this criterion overestimates resonance effects. They interpret the results for the neutral complexes (and additional new results for the cationic complexes) in terms of the Taft m and OR parameters. They conclude that in neutral [PtCl(N=NAr) (PEt3)2] there is weakn-donor and weak u-acceptor behaviour on the part of the (PEt3)2CIPtN2 fragment towards the FCsfi ring, and that in the cationic complexes [Pt(N=NAr)L(PEt&]+ the (PEt&LPtN2 group acts as a distinctly better a-acceptor (as expected) with the n-donor ability suffering minimal change.It is regrettable that these studies have not included the aryldiazene complexes for better comparison with the crystallo- graphic dimensions. Nevertheless, this interpretation seems to correlate sensibly with the expected properties of the doubly bent ArN2- ligand. In this connexion, Lalor and co-workers53 have measured the 19F n.m.r. chemical shifts of the m- and p-fluorophenyldiazenato-group for a wide selection of complexes of iron, molybdenum, tungsten, and platinum and have attempted to correlate the (8, - 8,) values with v(NN) determined from 15N-substitution studies* and v-effects.A re-examination of these results in terms of the Taft OR and a1 para- meters would clearly also be of some interest. *Note that v(”) values were not corrected (see ref. 806). 6a W. E. Carroll, M. E. Deane, and F. J. Lalor, J.C.S. DaZton, 1974, 1837. 459 (1 5 ) E. Reaction with Ary1hydrazines.-There are several instances where aryl- hydrazines have reacted with a transition-metal complex to form an aryldiazenato- complex by an in situ oxidation of the hydrazine as in equations (16) and (17).17161 (16) Co-ordination Chemistry of Aryldiazoniurn Cations i~m,~4954 ArNs+Cl- -PPha example for fMHC12(PPh3)3] insertion rhodium,l4 of ArN2+ - and into iridium,14,55,56 lM(NH=NAr)C13(PPh3)2 M-H bonds has also equation been J - I_, base HCl observed (14).14 for ruthen- [M(N=NAr)Ch(PPh&] (M = Rh or Ir) (14) I t should be pointed out, however, that sometimes 'insertion' occurs to give an identifiable aryldiazene complex but the corresponding aryldiazenato-derivative is not easily isolated after attempted deprotonation.14~56 D.Reaction of Co-ordinated NO with ArNH2.-Nitrosyl complexes which cor- respond to the NO+ formalism are potential electrophiles, to an extent which may be loosely related to the direct magnitude of v(NO).5' Attack by RNH2 may yield aryldiazenato-complexes according to equation (15), and in this way v(") M-NO + RNH2 3 M-N2R + H20 the compounds [Ru(bipy)z(NzR)Cl] (PF6)2 (R = p-MeOC6H4 or p-MeC6&) have been synthesized from [Ru(bipy)z(NO)CI] ( P F + L ~ ~ These possess remark- ably high values of 2095 and 2080 cm-1 respectively, which at least attest to their formulation as N2Ar+ complexes of ruthenium@) and possibly may indicate a molecular geometry in which the NNAr angle is considerably greater than 120", i.e.tending towards structure (2). This synthesis is probably restricted to nitrosyl complexes having v(N0) greater than about 1900 cm-1 and may be of less general usefulness than even this, as in the author's laboratory and elsewhere the NO group in a number of such nitrosyl complexes has been found not to react similarly with PhNHNH, rner-[ReCh(PPhMe2)3] ___+ [ReCh(NzPh) (PPhMe&] PhNHNH, IMo(C~H~)H(CO)BI d CMO(C~H~) (N2Ph) (C0)zI reflux The linearity of the Re-N-N (17) skeleton and the 118" angle at N2 (Table) suggest a4 D.F, Gill, B. E. Mann, and B. L. Shaw, J.C.S. Dalton, 1973, 311. 55 L. Toni010 and R. Eisenberg, Chem. Comm., 1971, 455. 66 G. Caglio and M. Angoletta, Guzzettu, 1972, 102, 462. 67 F. Bottomley, W. V. F. Brooks, S. G. Clarkson, and S. B. Tong, J.C.S. Chem. Comm., 1973, 919; F. Bottomley, personal communication. ( 8 W. L. Bowden, W. F. Little, and T. J. Meyer, J . Amer. Chem. Soc., 1973,95, 5084. s9 L. Busetto, A. Palazzi, D. Pietropaolo, and G. Dolcetti, J. Organometullic Chem., 1974, 66,453. .O J. A. Carroll and D. Sutton, unpublished results. s1 M. L. H. Green, J. R. Sanders, and R. N. Whiteley, 2. Nuturforsch., 1968, 23b, 106.460 Sutton formulation of [ReCh(NzPh) (PPhMe2)3] as an N2Ph+ compound of rhenium(1) rather than an N2Ph- complex as originally proposed.17 The molybdenum compound may be formulated as a N2Ph+ complex by analogy with the dimen- sions in the Table of the corresponding HB(pz)3 derivative.l6 An interesting example of the ligand tautomerism which may occur is the reaction of benzoylhydrazine with (ReOC13(PPh3)2].4a Spectroscopy favours the presence of a benzoylhydrazido(3-) ligand as in (21) rather than the chelating benzoyldiazenato-ligand illustrated in the resonance structure (22). However, PPh :1 N=N ('I PPh, CI, I ,N-N 3cr, I PPh, nicr - [ ReCl, ( N = NCO I'h) ( PMe PI1 z ) ; ~ ] PPh:, reaction with chlorine or with nucleophiles opens the ring to yield benzoyldia- zenato-complexes, as in Scheme 7.These were described4a as ReIV and ReT1' PPh, PMePh2 Scheme 7 complexes, implying the presence of (N2COPh-), but a Re1(N2COPh+) formula- tion seems to be more in keeping with the X-ray-determined linear Re-N-N skeleton reported4a97 for [ReCh(N2COPh) (PMezPh)a] (Table); it is also con- sistent with the 18-electron rule, though it is by no means certain that these compounds need necessarily conform to this rule. Apparently rhenium acyldia- zenato-complexes could not be synthesized in a comparable manner from acylhydrazines such as MeCONHNH2. However, the rhenium(1) dinitrogen compound [ReCI(N2)(PMe2Ph)4] may be converted4b into the benzoyldia- zenato- or acetyldiazenato-complex using benzoyl or acetyl chloride [equation (l)].The disphosphine complex [ReCl(N2) (diphos)z], does not undergo this reaction, which requires the loss of a labile monophosphine ligand; it may be compared with the aroylation or acylation of the molybdenum and tungsten compounds in equation (18) which proceeds accompanied by the loss of the 461 Co-ordination Chemistry of Aryldiazonium Cations second labile dinitrogen ligand.5~62 The original rhenium(1) dinitrogen complex is RCOX HX IM(Nz)z(PR3)41__+ [ M O R ) (PR&Xl+ [M(N2(H)COR) (PR3)4X] (18) (M = Mo or W) regenerated in the presence of excess phosphine. The reaction illustrated in equation (18) is not confined to acid halides, and importantly it occurs with HX and MeX (X = Cl or Br).The former5 is the first example of protonation of a co-ordinated dinitrogen ligand in a stable complex; it has not proved possible4b to protonate the [ReCI(N2) (PR&] complexes or to protonate, acylate, or aroylate the dinitrogen in [OsC12(N2) (PEtzPh)s]. The sec0nd~~1~3 is important in provid- ing rare examples of alkyldiazenat o-complexes. The latter are generally inacces- sibIe by the more convenient routes used for aryldiazenato-complexes by virtue of the instability of alkyldiazonium ions, but a methyldiazenato-molybdenum complex, [Mo(N2Me)(S2CNR&], has been obtained by the reaction of methyl- hydr azine with [MOO 2( S 2CN R2)2]64 and a t rimet h y 1 sil ylmet hyldiazenat o- complex has been synthesized by the 1,3-insertion reaction [equation (19)].65 [M(C5H5)H(C0)3] + Me3SiCHN= N 3 [M(GH5) (N2CH2SiMe3) (CO)a] (19) This compound may be compared with the structure of (23)66 established by - + X-rays67 where 1,3-addition of Et02CCHN=N across a molybdenum carbonyl centre in [Mo(CgH5) (CO)3]- has taken place, followed by protonation and rearrangement to the carbene.- + (M = MoorW) H 0. '0 F. Reaction with 0rganodiazenes.-This reaction is in principle the replacement of R2 in an organic diazene by a metal [equation (ZO)]. It has received little J. Chatt, G. A. Heath, and G. J. Leigh, J.C.S. Chem. Comm., 1972, 444. *a (a) J. Chatt, J. R. Dilworth, G. J. Leigh, and R. L. Richards, Abstracts of the International Symposium on Nitrogen Fixation, Pullman, Washington, June 1974, p.2; (6) T. A. George '' M. W. Bishop, J. Chatt, and J. R. Dilworth, J. Organometallic Chem., 1974, 73, C59. and S. D. A. Iske, ibid., p. 4. 66 M. F. Lappert and J. S. Poland, Chem. Comm., 1969, 1061. 68 M. L. H. Green and J. R. Sanders, Chem. Comm., 1967, 956. 67 J. R. Knox and C. K. Prout, Acra Cryst., 1969, B25, 1952. 462 R1-N=N-R2 + M 3 R1-N-N-M Sutton attention so far, but recent results for silyl compounds indicate that this may + R2 (20) well turn out to be a generally useful method. Trimethylsilyl(phenyl)diazene, MesSiN=NPh, removes the bromide ligand from [Mn(CO)sBr] with the formations5 of the novel dimer [Mn(N2Ph)(CO)& [equation (21)]. [Mn(CO)sBr] + MesSiN-NPh + [ (Mn(N2Ph) (C0)4 121 + MesSiBr (21) - co The molecular structure (24) of this first established example of a bridging aryldiazenato-ligand has been discussed in Section 3C.N-Methyl-N-nitrosotoluene-p-sulphonamide is a convenient sowce of nitrosyl ligands in the single-step synthesis of numerous nitrosyl complexes of the platinum metals, but the analogous aryldiazenes, p-MeCs&SOzN(Me)N2AT do not appear to afford clean syntheses of corresponding aryldiazenato-com- ~1exes.l~ 1 ,3-DiaryltriazenesY ArHN-N=N Ar or Ar(Me)N-N=NAr, cleave readily in acid solution to produce the diazonium ion in situ, and these reagents have been found to give rise to aryldiazenato-complexes in some instances.l' Ph 'N N ' Ph 5 Aryldiazenato-complexes of 4d and 5d Metals At present, aryldiazenato-complexes have been synthesized containing chro- mium, manganese, iron, cobalt, molybdenum, tungsten, rhenium, ruthenium, osmium, rhodium, iridium, and platinum. The main features of the compounds of the 3d metals and of rhenium and platinum have been dealt with in the preceding sections.The most extensive studies have been concerned with the pairs of elements molybdenum and tungsten, ruthenium and osmium, and rhodium and iridium. An abbreviated review of the aryldiazenato-complexes of these metals is presented here in order to indicate the nature and scope of the present situation. A. Molybdenum and Tungsten.-Kingl prepared the prototype aryldiazenato- complexes (25) for molybdenum [equation (7)] but was unable to obtain a 463 Co-ordination Chemistry of Aryldiazonium Cations tungsten analogue, though this was soon synthesized by Nesmeyanove* via a similar reaction at - 60 "C.The a-aryl complex [W(C5H5)Ph(C0)2] was obtained also, but it is not clear whether this arises from competitive nucleophilic attack at carbon or as a result of nitrogen extrusion from the aryldiazenato-complex the observation that v(") Q oc ( 2 5 ) M = Mo or W (cf. the platinum compo~nds5~~51). The aryldiazenato-complexes may be obtained alternatively from the cyclopentadienylcarbonylhydrides and phenylhydrazinesl [equation (17)]. An extensive series of analogous complexes in which T-cSH6 is replaced by a polypyrazolylborate ligand [RB(pz)3-] has been synthesized by Trofimenko.69 The similarities in the chemistries of the analogous NO and N2Ar complexes However, interesting differences in the HB(pz)3 vs.(C5H5) and N2Ar vs. NO relationships become evident upon closer examination, and the former may in part be accounted for by the greater bulk of the HB(pz)3 ligand. For example, the products in equations (27) and (28) might appear to be of similar structure, but is so much lower than v(N0) in the i.r. spectrum reA. N. Nesmeyanov, Y. A. Chapovskii, N. A. Ustyniuk, and L. G. Makarova, Izvest. Akad. Nauk S.S.S.R., Ser. khim., 1968, 449. O 8 S. Trofimenko, Inorg. Chem., 1969, 8, 2675; J . Amer. Chem. SOC., 1970, 92, 5118. 7 0 M. E. Deane and F. J. Lalor, J . Organometallic Chem., 1973, 57, C61. 71 M. E. Deane and F.J. Lalor, J. Organometallic Chem., 1974, 67, C19. T. A. James and J. A. McCleverty, J. Chem. SOC. (A), 1971, 1068. 464 Sutton points to an entirely different structure for the aryldiazenato-complex con- taining (N2Ar-) rather than (N2Ar +) ; furthermore, considerable differences occur in the behaviour of analogous molybdenum and tungsten aryldiazenato- complexes upon halogenation.71 The synthesis of other molybdenum and tungsten aryldiazenato-complexes has been reported.48, 73 B. Ruthenium and Osmium.-The obvious question to be asked when comparison is made of ArN2+ with isoelectronic molecules is the possibility of synthesizing [Ru(NH3)5(NzAr)]3+, analogous to the known [Ru(NH&Xln+ ions, where X is NO+, N2, or C0.74 The answer is presently negative: attempted direct dis- placement of H2O from [Ru(NH3)5(H20)]2+ by N2Ar+ yielded dinuclear dinitrogen complexes,75 and various other attempts involving phenylhydrazine etc.have led to the same product.60 There is no obvious reason why this ion should not exist, but no method has yet achieved the correct experimental conditions, i.e. to avoid a cation-cation reaction or the in situ liberation of Nz. The reaction of ArN2+ with [RuC12(PPh&] does not seem to be simple, and leads to compounds such as [RLI~(N~A~)~(PP~~)~~C~~~] [BF& where probably n = 2.48 However, ruthenium and osmium aryldiazenato-complexes of the type [MX3(N2Ar) (PPh3)2] result from addition of LiX to the r e a c t i ~ n . ~ ~ ~ ~ These exhibit v(”) frequencies in the range 1850-1900 cm-l, strongly in- dicative of an NzAr+ formulation, and this is confirmed by two independent crystal-structure determinations of solvates of mer-[RuCl3(N2C6HgMe-p)- (PPh3)iJ in which a singly bent N2Ar ligand is obser~ed~g~~~(Tab1e).There is a parallel series of nitrosyl complexes [MX3(NO) (PPh3)2]76 having v(N0) values closely similar to these v(NN) values, indicative of analogous linear NO+ formulations. The aryldiazenato-complexes are resistant to protonation under mild ~0nditions.l~ [RuC13(N2Ar) (PPh3)2]BF4 is also reported to be the product of chlorination of the interesting compound described48 as the bis(ary1diazenato) complex [RuCl(NzAr)2(PPh&]BF4. A complex of this composition may be analogous to the known bis(nitrosy1) [RuCl(NO)2(PPh3)2]PF6, which has been shown by Pierpont and Eisenberg77 to be formally a RuII complex with both linear (NO+) and bent (NO-) nitrosyl groups.A full X-ray structure determination would establish this analogy; unfortunately the attempted pre- paration of single crystals of this compound led instead to isolation of a solvate of [RuC~~(N~A~)(PP~~)~].~O This interesting compound deserves a renewed effort at characterization and structure elucidation. Indeed, its synthesis from [RuHCl(PPh&] and ArN2+ alone deserves further work to elucidate the nature and stability of possible intermediates such as [Ru(NHNAr)CI(PPh3)3]+ and their ease of deprotonation. 73 F. J. Lalor and P. L. Pauson, J. Organometallic Chem., 1970, 25, C51.’* J. N. Armor, Znorg. Chem., 1973, 12, 1959; J. Chatt, G. J. Leigh, and N. Thankarajan, J. Chem. SOC. (A), 1971, 3168. 76 V. B. Shur, I. A. Tikhonova, and M. E. Vol’pin, Zzvest. Akad. Nauk S.S.S.R., Ser. khim., 1972, 212. 76 S . D. Robinson and M. F. Uttley, J.C.S. Dalton, 1972, 1 . 77 C. G. Pierpoint and R. Eisenberg, Znorg. Chem., 1972, 11, 1088. 6 465 Co-ordination Chemistry of Aryldiazonium Cations The reactions of ruthenium and osmium carbonyls and carbonyl hydrides have received some attention. The reaction of nitrosonium ions with [Ru(C0)3- (PPh3)2] differs from the case of the corresponding iron compounds in that no nitrogen-containing complexes result,32a bu trecently32b complexes [Ru(NzAr)- (CO)z(PPh3)2]+ have been obtained with diazonium ions, paralleling the iron case.The monohydrido-complexes muHCl(CO)(PPh3)3] and [OsHCl(CO)- (PPh&] react with diazonium tetrafluoroborates to yield complexes of stoichei- ometry MCI(BF4)(NH=NAr)CO(PPh3)2, believed to contain co-ordinated BF4. These compounds in solution probably have BF4 replaced by s0lvent4~~ and readily absorb CO to yield the cis-dicarbonyl complexes [MCl(NH=NAr)(C0)2- (PPh3)2] BF4 or react with LiCl to yield [MCh(NH=NAr)CO(PPh3)2]. The dicarbonyl salts react with LiCl to yield mixtures of [MC12(NH=NAr)(CO)- (PPh3)2] and cis-[MCh(C0)2(PPh3)2]; the fate of the liberated aryldiazene in the formation of the latter product is intriguing, but was not ascertained. The aryl- diazene ligand is also removed from [MCh(NH=NAr)CO(PPh3)2] by CO to give [MCIz(CO)z(PPh3)2], and is not hydrogenated under mild ~0nditions.l~ Shaw et al.54 observed the insertion of PhN2+ into the Ru-H bond of [T(NH) = - 3.571 but no attempt at deprotonation was reported.The dihydrido-complexes [MH2(CO) (PPh3)3] react with p-tolyldiazonium fluoroborates to give [MH(NH=NAr)CO(PPh3)3] BF4; reaction of the second M-H bond is not reported to occur, despite the fact that an almost 2 : 1 [ArN2+] : [complex] ratio was employed.l4 The aryldiazene complexes react with LiCl to yield [MCh(NH=NAr)CO(PPh3)2] and, for ruthenium, also [RuHCI(NH= NAr)CO(PPh3)2]. Deprotonation of several of these aryldiazene complexes described above probably occurs in solution (as judged by colour changes) but no aryldiazenato-complex analogues of the corresponding known nitrosyls were isolated.The T(NH) values for all these complexes were in the range -3.8 to - 1.6, indicating acidic protons, and the J(16NH) values were ca. 69 Hz. Reactions of diazonium salts with [RuH2(PPh3)4] and [OsHq(PPh3)3] are immediate, but the products decompose spontaneously with evolution of N2.14 [RuHC1(CO)2(PPrn2Bun)2] to produce [RuCI(NH=NPh)(CO)2(PPrn2Bun)2]BF4 C. Rhodium and Iridium.-Baird and Wilkinson78 were the first to report the reaction of diazonium ions with rhodium(1) complexes. [RhCl(PPh&] did not yield an isolable aryldiazenato-complex, but more recently Meek et al. have demonstrated that when a chelating triphosphine ligand {PPP = PhP[(CH&- PPh212) replaces PPh3 the resulting complex (13) easily adds to NO+ and N2Ph+ to yield the expected singly bent nitrosyl complex [RhCI(NO) (PPP)]+ and doubly bent aryldiazenato-complex fRhCI(N2Ph) (PPP)]+ (12) of RhlI1 respecti~ely.~~~~ These compounds have been fully elucidated structurally by X-rays (Table).It may be noted parenthetically that when a slurry of [RhCI- (N2Ph) (PPP)]+BF4 in methanol is warmed to 50 "C [RhHCI(PPP)]BF4 may be isolated in good yield,29b accompanied by anisole, PhOMe. Perhaps MeOH is sufficiently acidic for H+ to displace the aryldiazenato-group as PhN2+ from the I* M. C. Baird and G. Wilkinson, J. Chem. SOC. (A), 1967, 865. 466 Sutton require correc- co-ordination sphere, possibly via a phenyldiazene complex. The RhI cation [Rh(CNBut)4]+ reacts with NO+ and ArN2+ to give [Rh(NO) (CNBut)4I2+ and [Rh(N&) (CNBut)4]2+, presumably by oxidative addition.79 When [RhX(PPh&] (X = C1 or Br) reacts with diazonium ions in the presence of LiX the rhodium(m) complexes [RhX2(N2Ar) (PPh3)2] are obtained.14 These are notable for exhibiting two i.r.bands in the 1550-1615 cm-1 region which are sensitive to 15N substitution, but each shifts by less than the calculated amount for v(ISN=NAr). This effect is not restricted to these compounds, but seems to be due to vibrational coupling53~sO rather than cis-trans isomerism of the aryldiazenato-ligand or the complex. Measured values of v(") tion if this effect is observed.80b Aryldiazonium fluoroborates are believed to react with [RhHCIz(PPh3)3] (free from excess HCI) with the formation of [RhCIz(NH=NAr) (PPh3)3]BF4 by apparent insertion into the Rh-H bond.However, attempts to purify and characterize these salts have been unsuccessful.~4~7~ When the reaction is con- ducted in the presence of excess HCI or LiCl, the products are now believed to be14 the neutral trichlorobis(phosphine) complexes [RhC13(NH=NAr) (PPh3)2] rather than the (less expected) aryldiazenato-complexes of the type [RhC12(N= NAr) (PPh3)2(solvent)] claimed at fist78 and subsequently examined as such by Raman and i.r. spectroscopy.8l This reformulation follows naturally from the syn- thetic conditions employed, and from the fundamental differences in the nature of the products compared with the [R~CIZ(NZA~)(PP~S)Z] complexes synthesized separatelyl4 and described above.Thus the compounds show a weak broad resonance at ca. T -2.0 which is replaced by a sharp doublet [J(15NH)65-70 €321 in the compounds synthesized from 15N=NAr+, indicating protonation of N1. The aryldiazenato and aryldiazene complexes may be interconverted according to Scheme 8. [ llhCl3( CO) ( PPh,),] (A): v ( C 0 ) = ? I O 8 ~ r n - ~ [ Rh 2CIG(NH2NHAr)2( I'Ph,)2] A -b [RhCI(CO)(PPh,),] v ( C 0 ) = 1960 cm-' Scheme 8 'O J. W. Dart, M. K. Lloyd, R. Mason, and J. A. McCleverty, J.C.S. Dalton, 1973, 2039. (a) D. Sutton, Canad. J . Chem., 1974, 52,2634; (b) B. L. Haymore, J. A. Ibers, and D. W. Meek, Inorg. Chem., 1975, 14, 541. 81 G. W. Rayner Canham and D.Sutton, Cunud. J. Chem., 1971,49, 3994. 467 PPh, oc Co-ordination Chemistry of Aryldiazonium Cations Insertion into Ir-H bonds has been a common source of iridium aryldiazene and aryldiazenato-complexes also. [IrHC12(PPh3)3] may be converted into the complexes [IrCh(NH=NAr) (PPh&]BF4 and [IrCh(NH=NAr) (PPh3)2] by reactions analogous to those described above.14 Like the rhodium analogues, [IrC13(NHNAr) (PPh3)2] may be converted by means of HC1 into an unstable brown complex. This reaction may be reversed with NEt3, and the brown complex is therefore tentatively identified as the aryldiazenato-complex [IrClz(NzAr) (PPh3)2]. Toniolo and Ei~enberg~~ and Caglio and A n g ~ l e t t a ~ ~ obtained [IrHz(NH=NAr) (PPh3)3]BF4 from the insertion of diazonium ions into mer- [IrH3(PPh3)3].The results of a preliminary X-ray investigation55 indicated structure (26) for the cation. No further hydrogenation of the aryldia- H zene ligand occurred at 1-3 atm H2 at 50-80 "C, and no indication was made whether insertion of further ArN2+ into the remaining Ir-H bonds could be accomplished. On attempted deprotonation with base no aryldiazenato-complex was formed ; instead the original complex [IrH3(PPh3)3] was regenerated. The compounds are unstable in chloroform or chloroform-ethanol, depositing orange crystals for which v(IrH) is absent. The product from chloroform has been shown by X-rayss2 to be the orthometallated diazene (27), in which the initial hydride ligands have been replaced by chloride derived from the solvent.Vaska's complex, [IrCl(CO) (PPh3)2], occupies a central place in IrI chemistry, (2 7) - PI1 3 P - - + -I (28) *a P. L. Bellon, G. Caglio, M. Manassero, and M. Sansoni, J.C.S. Dalton, 1974, 897. 468 Sutton and several gr0ups~~,~3-*~ have examined its aryldiazenato and aryldiazene derivatives. The synthesis of the Ir1(N2Ar+) complex &Cl(NAr)(PPh3)2]+, isoelectronic with the nitrosyl complexs6 fIr(NO)Cl(PPhs)s] +, has been described in Section 4A. This complex has a considerable chemistry, being a precursor to numerous five- and six-co-ordinate aryldiazenato-complexes which may be expected to show interesting redox properties reflecting NzAr+ and N2Ar- formalisms, as do the corresponding n i t r ~ s y l s . ~ ~ Thus, the reaction of [IrCl- (N2Ar) (PPh3)2]+ with CO, or the reaction of [IrCl(CO) (PPh3)2] with diazonium ions, yields the five-co-ordinate cation [IrCl(NzAr)CO(PPhs)z]+, believed to contain doubly bent ArN2- and to have structure (28) analogous to the struc- turally determined3 NO- bent-nitrosyl complex of iridium(rI1). From this, an extensive series of six-co-ordinate iridium(II1) aryldiazenato-complexes may be obtained by further addition of neutral or negative ligands.47 Studies in the author's laboratory have indicated that the reactions of [IrCl(CO) (PPh3)2] with diazonium ions in the presence of ethanol or propan-2-01 may be considerably more complex than the simple additions just described. The principal products then obtained are the cationic orthometallated diazene85 (29) and the iridium(Ir1) tetrazene complex cation (30).s7 A tentative mechanism for + + I PPh, C,H4F C,H, F the synthesis of the tetrazene has been proposed.8s The diazene is closely analo- gous to the orthometallated complex (27) described by Bellon et aZ.,sz and it might be concluded that it arises as a result of an orthometallation isomerization of the five-co-ordinate simple addition complex (28) with which it is isomeric. Such a mechanism is attractive in view of the vacant co-ordination site in (28), yet there remain doubts about the validity of this mechanism as we have never 83 A. J. Deeming and B. L. Shaw, J. Chem. SOC. (A), 1969, 1128. 84 B. L. Shaw and R. E. Stainbank, J.C.S. Dalton, 1972, 223. 85 (a) F. W. B. Einstein, A. B. Gilchrist, G. W. Rayner Canham, and D. Sutton, J. Amer. Chem. SOC., 1972, 94, 645; (b) F. W. B. Einstein and D. Sutton, J.C.S. Dalton, 1973, 436. 86 C. A. Reed and W. R. Roper, Chem. Comm., 1969, 1459. *' (a) F. W. B. Einstein, A. B. Gilchrist, G. W. Rayner Canham, and D. Sutton, J. Amer. Chem. SOC., 1971, 93, 1826; (b) F. W. B. Einstein and D. Sutton, Inorg. Chem., 1972, 11, 2827. A. B. Gilchrist and D. Sutton, Canad. J. Chem., 1974, 52, 3387. 469 Co-ordination Chemistry of Aryldiazonium Cations been able to induce the orthometallation of (28) to occur other than in the pres- ence of ethanol, propan-2-01, or similar a-hydridic alcohols.*Q From this it seems to us that an intermediate hydride abstraction from the solvent is a necessary step. 6 Conclusion Whilst the chemistry of aryldiazenato-complexes is mostly of very recent origin, it can be said already that they display much of the expected structural and electronic similarity with the isoelectronic nitrosyl complexes, underlining the usefulness of classifying them pragmatically in terms of N2Ar+ and N2Ar- ligands. Despite many aryldiazenato-complexes presently appearing to be less ‘stable’ than their nitrosyl counterparts, the number of successful syntheses to date, the vast number of known nitrosyl compounds, and the already evident greater versatility of the aryldiazenato-ligand toward electrophiles promise much for the future development of their chemistry. N. Farrell, A. B. Gilchrist, and D. Sutton, unpublished results. 470

ISSN:0306-0012    

DOI:10.1039/CS9750400443

出版商:RSC    

年代:1975

数据来源: RSC

摘要:

The Application of Electrochemical Techniques to the Study of Homogeneous Chemical Reactions By D. Pletcher DEPARTMENT OF CHEMISTRY, THE UNIVERSITY, SOUTHAMPTON SO9 5NH 1 Introduction All instrumental techniques require a number of years development before commercial apparatus is available and the theoretical background is thoroughly understood. Happily a number of electrochemical techniques have reached this state of maturity and are now well suited to the investigation of the homogeneous chemical react ions of inorganic, organic, organometallic, and biochemical species. Indeed, polarography, cyclic voltammetry, potential step methods, rotating- disc and ring-disc electrode techniques, coulometry, product isolation tech- niques, and spectroelectrochemical methods may together be adapted to allow the study of reaction mechanisms, the identification of intermediates, and the measurement of rate constants for their chemical processes.These intermediates may be either relatively stable or rather shortlived, and in general electrochemical techniques can cover a range of half lives between 0.1 ms and 104 s. Electro-chemical techniques are particularly well suited to the study of intermediates which may be produced by fast one-electron transfer processes, and hence it is not surprising that electrochemistry has made a significant contribution to the understanding of radical ion chemistry. The study of homogeneous chemical reactions by electrochemical techniques is dependent on one of two concepts. Either a system may be designed where the chemical reaction of interest controls the rate of production of an electroactive species, or the species whose chemistry is to be studied may be produced by an electrode reaction. In the latter case the overall electrochemical experiment effectively consists of two parts; in the first the species is generated and in the second the electrode system is used to analyse for the species remaining in solution after a time lapse.This concept has very general application since a very wide range of species may be produced at electrodes (e.g. carbanions, radicals, carbonium ions, dications, ion radicals, metal ions in unusual oxi- dation states, halogens, organometallic species). This review sets out to describe the most important electrochemical methods used in the study of homogeneous chemical reactions and to illustrate their use with examples taken from the whole of chemistry.The first section, however, discusses the basic principles of electrode processes essential to the understanding of these techniques. Application of Electrochemical Techniques to Study of Chemical Reactions 2 Fundamental Concepts of Electrode Processes The electrochemical experiments to be discussed in this review always consist of an investigation of the relationship between the current flowing through the system and the electrode potential (versus a reference electrode) and the de- pendence of this i-E relationship on experimental parameters such as the time- scale on which the observation is made, the nature of the solution environment, and the effect of convection.Even the simplest electrode process (adsorbed intermediates and surface reactions will not be discussed) is, however, a sequence of several events; for example, the conversion of the solution-free species A into the solution-free product B is a three-step process Asoln -mass transport Asurface electron transfer Asurface k ne Bsurface masstransport Bsurface -Bsoln (2) (3) The current is determined by the overall rate of this sequence and therefore by the slowest of the three steps. Thus an understanding of any electrochemical experiment requires a consideration of both the heterogeneous electron-transfer process and mass transport.The Heterogeneous Electron-transfer Process.-It is the electrode potential which determines whether sufficient energy is being supplied for the electron transfer to occur. The standard electrode potential, E*, for the reaction A+ne+B (4) versus a reference electrode, for example a hydrogen electrode, is related to the standard free energy, AG*, of the overall cell reaction e.g. A k *H+ + B k 2 H~ by the equation AG* = -nFE* (6) Certain electron-transfer processes are fast at the standard potential (k* > cm s-l, see later) and in most experiments with such electrode processes it appears that the electron-transfer process at the electrode surface remains in equilibrium. Under these circumstances the shape of the rising portion of the i-E curve can be derived by calculating the surface concentrations of the reactant and product, CI and C;, respectively, from the Nernst equation a purely thermodynamic equation.Such systems are known as reversible. Pletcher Conversely, irreversible electrode processes are those where some considera- tion must be given to the kinetics, as well as the thermodynamics of the reaction in order to deduce the relationship between current and potential. Such electrode reactions occur only at a low rate at the reversible potential, and in order to obtain a measurable current it is necessary to apply an overpotential. The equation relating current and potential for irreversible reactions has been derived elsewherel-4 -a)nF -E*)] -C;exp [(l RT where k* is the standard rate constant, a the transfer coefficient, and A is the surface area.The former is a measure of the rate at the standard electrode potential. Under normal experimental conditions, one of the exponential terms in equation (8) is usually negligibly small and the equation becomes of the type RTlog i = log nFACIk* --(E -E*) (9)2.3anF Irreversible electrode processes give rise to i-E curves which are more drawn out and displaced along the potential axis when compared with those for a reversible system. Mass Transport.-In general three types of mass transfer must be considered: migration, diffusion, and convection. They arise from the movement of a species under the influence of, respectively, a potential field, a concentration gradient, and an external mechanical force.In kinetic experiments, the contribution of migration is usually suppressed by addition of an excess of an inert electrolyte, and it is normal to carry out experiments under conditions where either it is necessary to consider only diffusion, i.e. the system is unstirred, protected from temperature gradients, etc., or the solution is stirred in a defined manner. It is then necessary to consider both diffusion and convection. The rate of mass transport is dependent on the geometry of the electrode system, but most electrochemical experiments on a timescale of seconds or less and in unstirred solution are well approximated by one-dimensional equations, i.e. by linear diffusion to a plane electrode.Diffusion is then described by Fick’s laws in the form dCzfluxr = -Dz -dx and l J. O’M. Bockris and A. K. N. Reddy, ‘Modern Electrochemistry-Volume II’, Plenum Press, New York, 1970. a B. E. Conway, ‘Theory and Principles of Electrode Processes’, Ronald Press, New York, 1965. K. J. Vetter, ‘Electrochemical Kinetics’, Academic Press, New York, 1967. J. Koryta, J. Dvorak, and V. BohBEkovB, ‘Electrochemistry’, Methuen, London, 1970. Application of Electrochemical Techniques to Study of Chemical Reactions The first law describes the flux of a species through a plane normal to the electrode in terms of the concentration gradient at that plane in space and a diffusion coefficient, D.The second relates to non-steady state conditions and describes the change in the concentration with time of a species due to diffusion alone.In stirred solutions where both diffusion and convection must be considered, the mass-transport equation will be of the type where Vz is the velocity of solution movement in the x direction. Convective diffusion is, of course, dependent on the electrode geometry and the mode of stirring and it is usually necessary to write the equation in three-dimensional co-ordinates. Generally The breakdown of an electrode process into the sequence (1)--(3) readily explains the familiar S-shape of a steady-state i-E curve. At low potentials, it is the electron-transfer process which is the slowest step, and the rate of this pro- cess increases exponentially with potential [see equations (7) and (9)].Eventually step (2) becomes fast compared with (1) and mass transport becomes the rate- determining step; hence a plateau is observed on the i-E plot since the rate of mass transport does not depend on potential. In the potential region of this mass-transport limited current, the electroactive species will undergo the electron-transfer reaction as soon as it reaches the electrode surface so that at the surface its concentration will be zero. In experiments designed to study homogeneous chemical reactions, wherever possible the potential is chosen so that the electrode reaction is mass-transport controlled. Under these circumstances it is not necessary to consider the electron transfer process in any detail and no information about the kinetics of this process is required.3 Formulation of Theory In order to obtain quantitative kinetic data for a homogeneous chemical reaction from an electrochemical experiment it is necessary to formulate a theoretical description of the experiment and to solve the resulting partial differential equations. The solution required will describe the dependence of the i, E, t behaviour on the rate constant for the chemical reaction. The experiment may be described completely by a set of partial differential equations together with their initial and boundary conditions. For each species in the system, it is possible to write down a partial differential equation which describes how its concentration within the diffusion layer close to the electrode surface varies with time and distance from the electrode surface.In general this Pletcher equation will have terms which describe the change in concentration with time due to diffusion, convection, and chemical reactions, e.g. for a one-dimensional mass-transport situation where k is the rate constant for the chemical reaction and v and p are the order of reaction with respect to the species i and j, respectively. Integration of any such set of equations will require (i) information about the concentration of all species at one particular time, normally at the start of the experiment, t = 0, where the concentrations of all species are uniform at all points in space and determined only by the preparation of the solutions for the experiment, and (ii) information concerning the concentrations at two points in space.These are normally at the electrode surface, x = 0, and at an infinite distance from the electrode; the former are determined by the nature of the electrochemical experi- ment and in some cases will depend on time, while the latter again depend only on the solutions prepared for the experiment. In such general terms the situation sounds somewhat complex and is perhaps best simplified by some particular examples. Figure 1 shows the differential equations, initial and boundary conditions for two chemical systems and three electrochemical experiments carried out at stationary electrodes in unstirred solution (i.e. convection need not be considered).The solution of such sets of equations is seldom trivial and usually requires the use of the Laplace Transformation, conversion to integral equations followed by numerical integration or, more recently, simulation techniques. Such tech- niques are beyond the scope of this review and the interested reader is referred el~ewhere.~~~Fortunately a complete solution is not always required since the current is given by i = -nFAD (g)x=o where CA is the concentration of the electroactive species. Hence only the acAevaluation of -at x = 0 is necessary in order to determine the current. Of ax acA course, (z) is commonly a function of time. x-0 The complete solutions of the differential equations describe how the con- centration of each species in the system varies as a function of time and distance from the electrode surface.These are known as concentration profiles and a qualitative consideration of their time evolution during an experiment is often W. H. Reinmuth, Analyt. Chem., 1962, 34, 1446. S. Feldberg, Efectroanalyt. Chem., 1969, 3, 199. Application of Electrochemical Techniques to Study of Chemical Reactions extremely instructive (see, for example, Figure 2) and allows qualitative predictions concerning the results of the experiment. Systems (a) A + e +B (b) A + e +B k k B4P 2B+A + P Diferential equations acA a2cA(a> at = DA -ax2 Initial Conditions t = 0, x 3 0 CA = Cz and CB = 0 Boundary Conditions t > 0, x = co CA = C: and CB= 0 t > 0, x = 0 (i) Potential step experiment to a potential on the limiting current plateau : aCBC;= 0 and acA = -DB(~)DA (z) x =ox =o (ii) Double potential step.First pulse, r s, as (i). Second pulse at a potential where B -e 4 A is diffusion con trolled. t < r C;=Oand DA (!!) x=o = -DB(!!) x=o t > r C;=Oand DA = -DB (iii) Linear potential sweep, scan rate v V s-1 E = Ei + vt acB(..).=. acB DA (")ax z=o = -DB (z)x=o Figure 1 Diferential equations, initial and boundary conditions for potential step and linear potential sweep experiments designed to study the kinetics of two types of homogeneouschemical reaction (a)first-order decay and (b) disproportionation, where the reactive species isproduced in an electrode reaction Pletcher Increasing t X X m X t<Z 2 <z Increasing t Figure 2 The evolution with time of the concentration profiles for species A and B duringa double potential step experiment as Figure 1, a@), for a particular value of k.Lengthof Apulse r s. 4 Polarography In classical polarography,’-ll an i-E curve is recorded by varying the potential of a dropping mercury electrode slowly in a linear manner (ca. 0.1 V min-l). Polarography was the first instrumental electrochemical method to be developed and its popularity over many years may be attributed to the simple equipment required and the high degree of reproducibility which may be attained. It has, however, recently become less fashionable largely because the timescale on which chemical reactions are observed, essentially the drop time of the mercury electrode, is much more restricted than with the techniques described later.Systems for study are, of course, limited to those with an electroactive species within the potential range of mercury. The polarogram for each electron-transfer process has a characteristic S-shape (see Figure 3) and the experimental measurables are: (i) the current on the limiting current plateau; (ii) the half-wave potential, E+, i.e. the potential where the current is half the limiting value; (iii) the slope of the i-E curve in the region where the current changes with potential -this is commonly determined by measuring Ep -Ea.Although the limiting current for many electrode reactions is determined by the rate of diffusion of the electroactive species, it is not uncommon for it to be determined by the rate of a homogeneous chemical reaction.7-13 Kinetic waves L. Meites, ‘Polarographic Techniques’, Interscience, New York, 1965. P. Zuman, ‘The Elucidation of Organic Electrode Processes’, Academic Press, New York, 1969. J. Heyrovsky and J. Kuta, ‘Principles of Polarography’, Publishing House of the Czech. Academy of Science, Prague, 1965. O H. Strehlow, in ‘Techniques of Organic Chemistry-Volume VIII’, ed. A. Weissberger,Interscience, New York, 1963. l1 ‘Physical Methods of Chemistry, Part IIa, Electrochemical Methods’, ed. A. Weissbergerand B. W. Rossiter, Wiley-Interscience, New York, 1971.la R. Guidelli, Electroanalyt. Chem., 1971,5, 149. l3 P. Zuman, Progr. Phys. Org. Chem., 1967, 5, 81. 477 Application of Electrochemical Techniques to Study of Chemical Reactions Current I I I I I ' +Potential. -0.1 -0.2 -0.3 -0.4 -0.5 -0.6 vs. Hg pool Figure 3 D.c. polarograms for (A) O.lM-HzSO1; (B) O.lM--H,SOr + le4M sodium molybdate; (C) 0.1M-H2S04+ 10-4M sodium molybdate + 0.4M sodium perchlorate. The limiting current on the plateau of curve C would allow the estimation of the rate con- stant for the MoIII + C10,-reaction arise when the rate at which the electroactive species is produced at the electrode surface is determined by a chemical reaction, i.e. k +ne A+O+R Where the equilibrium in the bulk solution away from the electrode lies to the left, A is not electroactive and the rate of conversion is slow; the limiting current for the electrode reaction will be small compared with the expected diffusion current if A were reduced directly, and the kinetics of the chemical reaction may be determined directly from a measurement of the limiting current.This method has been widely employed, for example to study the dehydration of diols to aldehydes and slow protonation reactions. 8s 9 l3 When the product of the primary electron-transfer reaction is unstable in the solution environment and the product of the chemical reaction is electroactive, Pletcher the limiting current will again contain a kinetic c~mponent.~-l~ There are two principal types of system where such behaviour occurs +ne k +me O-+R--+O'-+R' ne k 0-+R -0 + product (18)+ reactant In the former case, the polarographic wave will in the limit change from an ne wave when the chemical reaction is slow to an (n + m)e wave when it is fast.In the intermediate situation the wave height will be between these extreme values and will be a function of the rate constant, k. An example of this situation is the reduction of aromatic hydrocarbons in the presence of low concentrations of a proton donor14 Ph + e + Ph-(19) k +C Ph*-+ BH --t B-+ PhH. ---+ PhH-(20) In the latter case, the polarographic wave, known as a catalytic wave, can be many times the height of the diffusion controlled current for a simple reduction of 0 and this catalytic current may be shown to be dependent on the rate constant for the chemical step.The system MoV + 2e -+ MoIII (21) MoIII + c104---+ MoV + C1-(22) is representative of such systems15 and Figure 3 shows the polarograms for reduction of molybdate in the presence and absence of perchlorate ion. In the sulphate medium, molybdate reduces in the two steps as can be seen by the waves E+ = -0.14 and -0.32 V, respectively, and the fact that the second is twice the height of the first. Both are diffusion controlled. The simple test for diffusion control is to examine the limiting current as a function of the head of mercury above the capillary, h; id is proportional to ,/xIn the presence of perchlorate ion the first wave is unchanged but the second at E+ = -0.32V is greatly enhanced in height since it is now a catalytic wave.The shift in half-wave potential with change in the solution environment has been widely employed to study solution equilibria. In particular the variation of E&with pH and the concentration of complexing agents has been used to study the protonation equilibria of organic molecules and the complexation of metal ions respectively. In both cases it is possible to ascertain both the stoicheiometry of the equilibrium and the equilibrium constant. Indeed polarographic studies l4 M. Peover, Electroanalyt. Chem., 1967, 2, 1. lo I. M. Kolthoff and W. Hodara, J. Electroanalyt. Chem., 1963, 5, 2. Application of Electrochemical Techniques to Study of Chemical Reactions have contributed substantially to our knowledge of PKA values and stability constants.7-10.13 Reversible one-electron transfer processes give rise to a wave with a slope [Ea -Ei] = 56 mV.Hence although further corroborative evidence is essential, the observation of a polarographic wave with such a slope is strong evidence that the electrode process involves the transfer of one electron and produces an intermediate stable on the timescale of the droptime of the electrode. Indeed such polarographic results provided the first electrochemical evidence for stable anion radicals14 and for transition metal ions in unusual oxidation states,l6 e.g. sfn3+ =Sm‘+ MeCN Polarography is an excellent technique for analysis and hence may readily be used to follow slow chemical reactions (T+ > mins) provided one of the reactants or products is electroactive. For example the reaction phosphate MeCOCOgH + NADH +MeCHOHCOzH + NAD+ (26)buffer which is catalysed by the enzyme lactate dehydrogenase, may be followed by monitoring the limiting current for the polarographic reduction of the nicotinamide-adenine dinucleotide (NAD+).l7 5 Linear Sweep and Cyclic Voltammetry In linear sweep ~oltammetry~~J*J~ the potential of a stationary electrode (a hanging or sitting drop mercury or micro solid electrode) is varied rather rapidly (0.1-lo00 V s-l) and linearly with time between two fixed potentials EZand E’ and the resulting current response is recorded.In such experiments the rapid change in potential requires that non-steady state diffusion must be taken into account, and as a result the i-E curve for a simple electrode process shows a marked current peak instead of the smooth S-shape for a steady-state technique. In addition the current density for the electrode reaction will be considerably higher and time dependent. A more widely used but closely related technique is cyclic voltammetry,11J8-20 where a saw-tooth potential-time profile is applied to the electrode; in other l6 I. M. Kolthoff and J. F. Coetzee, J. Amer. Chem. SOC.,1957,79, 1852. l7 D. Pletcher, unpublished work. R. S. Nicholson and I. Shain, Andyt. Chem., 1964,36, 706. l9 R. N. Adams, ‘Electrochemistry at Solid Electrodes’, Marcel Dekker, New York, 1969.*O G. Cauquis and V. D. Parker, in ‘Organic Electrochemistry’, ed. M. Baizer, Marcel Dekker, New York, 1973. Pletcher words, the potential is scanned linearly from Ei to Ef but the direction of the potential scan is then immediately reversed and the potential scanned in the reverse direction from E’to Ez. Cyclic voltammetry is an excellent technique for the qualitative study of the stability and homogeneous reactions of species which may be produced in an electrode reaction. For any particular chemical system, the principal experi- mental variables are the potential range, Et-Ef, and the potential scan rate, v. This latter variable is effectively used to vary the timescale on which the chemistry of the reactive species is observed.In addition it will be seen that it is important to note differences between the first potential sweep and later scans when the potential-time profile is applied repetitively. Both cyclic voltammetry and linear sweep voltammetry may be used to obtain the quantitative kinetic data once the mechanism of the reaction is understood. The simplest electrode reaction is one where a single electron is transferred in a fast process to give a product which is stable, i.e. O+e+R (27) Such electrode processes give cyclic voltammograms such as that shown in Figure 4,1* provided the initial solution contains 0 but not R, and its form may be understood in the following way. The portion UV of the curve corresponds V 100mV i=O bE Figure 4 Cyclic voltammogram for a reversible one-electron transfer reaction 0 + e +R. Symbols used in the text are defined on this clcrve.Application of Electrochemical Techniques to Study of Chemical Reactions to the potential region where the ratio of the surface concentration of R to the surface concentration of 0 is increasing rapidly [see equation (7)] and hence the current is increasing. At V, the surface concentration of 0 is tending to zero and the electrode is purely diffusion controlled; the decay in current over the part of the cyclic voltammogram VWX reflects the decrease in the rate of diffusion of 0to the surface as its flux falls off with time towards a steady-state value. The portion XYis the potential region where the ratio of R to 0 at the surface again changes, this time 0 becoming more favourable.Finally Yis the potential where the flux of R to the surface, and hence the current for the reverse process, are at a maximum and beyond this peak both decrease towards zero, when there is again no electron-transfer reaction taking place. It has been shown18 that the cyclic voltammogram will have the following properties at room temperature (seeFigure 4 for the definition of the symbols): (i) the separation of the peaks for the forward and reverse processes is given by dEp = 60 mV; (ii) the peak potentials, Ep, are independent of the potential scan rate, v, as are the functions ii/v* and iF/v+; (iii) the half peak widths for the forward and reverse processes are defined by [Ep -Epp] = 57 mV; (iv) the ratio of the peak currents, ii/ii, is unity; (v) the current beyond the peaks decays according to i cc t-*; (vi) the peak potentials are related to the polarographic half-wave potential by E; = E+ -28.5 mV.Conversely, the observation of a pair of peaks which obey these criteria shows conclusively that the electrode reaction involves the transfer of a single electron and that the product of the electrode reaction is stable. Indeed such cyclic voltammograms form the basis of many electrochemical investigations of the chemistry of anion radicals and cation radicals since they show that under the experimental conditions these species may be prepared and are stable. In addition cyclic voltammetry offers a convenient experimental method for seeking condi- tions where normally reactive species might be stable, at least on the timescale of the experiment.Cyclic voltammetry may readily be carried out over a wide range of solution conditions and of temperature and such studies have been used to show, for example, the stability of a series of square-planar cobalt com- plexes,21 the chromium hexacarbonyl cation,22 the cation radicals of an-thracene,23 triphen~Iamine,~* and hexafluorobenzene,25 the dications of rubrene26 and thianthrene,27 and the dianion of diphenylketone.28 DMF \ !dlG. Costa, G. Mestroni, A. Puxeddu, and E. Reisenhofer, J. Chem. SOC.,1970, 2870. 8a C. J. Pickett and D. Pletcher, J.C.S. Chem. Comm., 1974, 660.a3 0. Hammerich and V. D. Parker, J.C.S. Chem. Comm., 1974, 245. !d4 H. Lloyd Jones, L. G. Boxall, and R. A. Osteryoung,J. Electroanalyt. Chem., 1972,38,476. 86 J. P. Coleman, M. Fleischmann, and D. Pletcher, Electrochim.Acta, 1973, 18, 331. L. Byrd, L. L. Miller, and D. Pletcher, Tetrahedron Letters, 1972, 2419. 87 0.Hammerich and V. D. Parker, Electrochim. Acta, 1973, 18, 537. A. Demortier and A. J. Bard,J. Amer. Chem. SOC.,1973,95,3495. 482 Pletcher C r (CO),i -e THF 253K -e MeCN+l"/oTFAn \ NaAIClo 423K --e \ FSOBH 233K -e FGFF F F Typical results for such experiments are those for the oxidation of rubrene in methylene chloride at 203 K shown in Figure 5; the cyclic voltammogram shows two reversible one-electron transfers and hence both the cation radical and the dication are stable under these conditions.L r 1 T I bE + 0.4 + 0.6 + 0.8 1 VS. Ag/sat reference Figure 5 Cyclic voltammogramfor rubrene at a Pt wire electrode in methylene chloride-0.2M tetrabutylammoniumtetrafluoroborateat 203 K. Thepotential scan rate is 0.01V s-l If the product of the one-electron transfer reaction is not stable on the time-scale of the cyclicvoltametric experiment there will be less of the product of the electron-transfer reaction remaining in the vicinity of the electrode surface (i.e. its concentration profile is affected) and hence the reverse peak on the cyclicvoltam-mogram will be smaller; i,"/i,"will be less than unity.As has already been stated above, however, the timescale of the experiment may be changed by varying the potential scan rate. If the potential scan rate is increased, the timescale of the Application of Electrochemical Techniques to Study of Chemical Reactions experiment is reduced and the chemical reaction will not occur to the same extent, i.e. the ratio i,”/i,”will then be higher. In the limit at very fast scan rates the system will appear reversible and obey all the properties (i) to (vi) listed above. Con- versely by slowing down the potential scan rate, the timescale for observation of the chemical reaction is lengthened and the ratio iE/i: will be less. At this limit all the intermediates will undergo the chemical reaction and no reverse peak will be observed.In practice, if the intermediate produced in the electrode reaction is either highly reactive or very stable only one of the limiting situations may be observable. Thus in the simplest terms, a cyclic voltammetric experiment may be thought of as two parts. During the part of the sweep UVWX of Figure 4 a reactant for the chemical reaction is generated at the electrode surface. The later part of the sweep analyses the amount of this reactant unchanged during the potential scan. The ratio ig/i: varies with change in the potential scan rate when the product of the electron-transfer process is unstable; this variation is accompanied by other changes in the cyclic voltammogram. Thus the peak potentials become dependent on the potential scan rate and because in many cases the product of the chemical reaction is itself electroactive, new peaks may appear.The peak for the oxidation or reduction of this product may appear before or beyond the peak for the primary electron-transfer process depending on whether or not it is easier to oxidize or reduce than the initial electroactive species. In the former case, the product of the chemical reaction will be electroactive at the potential of the primary electron transfer process and the function iE/v+is then dependent on the potential scan rate.29@ The qualitative observation of these changes only tell us that the product of the primary electron-transfer process is unstable, and in order to build up a picture of the detailed mechanism and kinetics of the chemical reaction it is necessary to make quantitative studies of the variation of i:, and E: as a function of potential scan rate.For example it has been shown that the quantity dE;/d log v is 29.1 mV if the chemical reaction is first order with respect to the species produced in the electrode reaction and 19 mV if this reaction is second order;31 the variation in the concentration of the reactants will also help to establish reaction orders. Furthermore, where the product of the chemical reaction is electroactive the investigation of the mechanisms may be helped by the process of ‘peak matching’ i.e. the cyclic voltammogram of likely inter- mediates may be run and the peak potentials compared.Once the qualitative mechanism is known the rate constants can also be found from the variation in either Eg or i,”/ii with log ~18.3~and from the use of analytical expressions or working curves derived theoretically. Fortunately, for most mechanisms these are available in the literature.llJ8 The use of cyclic voltammetry to study homogeneous chemical reactions is best illustrated by some examples. es R. S. Nicholson and I. Shain, Analyt. Chem., 1965, 37, 178, 190. 30 M. Mastragostini and J. M. Saveant, Electrochim. Acta, 1968, 13, 751. 31 J. M. Saveant and E. Vianello, Compt. rend., 1963,256,2597. 33 M. Mastragostini, L. Nadjo, and J. M. Saveant, Electrochim. Acta, 1968, 13, 721. 484 Pletcher The CoI(sa1en)-+EtBr Reaction.-Figure 6 shows cyclic voltammograms for i, .i=O ,S.C.E. iA Figure 6 Cyclic voltammograms run at a vitreous carbon electrode in DMF4.1M sodium tetrafluoroborate. Potential scan rate 0.1 V s-’ (a) ----5 x 10d3M CoII(sa1en);(b) 5 x 10-3M CoII(sa1en) +0.1M EtBr~ the CoII(sa1en) complex [salen is the quadridentate ligand NN-ethylenebis(salicy1i-deneimine)] in DMF in the presence and absence of ethyl bromide.17933 Curve (a), run between -0.5 and -2.0 V in the absence of the halide, shows a single electron-transfer reaction at EpF = -1.16 V which has all the characteristics of a reversible le process CoII(sa1en) +e +CoI(salen)-(35) In the presence of ethyl bromide the ratio i,”/i,”becomes dependent on v, showing that the anion reacts with the alkyl halide.Under the conditions where ipB/ipF < 1, a new reduction peak appears at Ei = -1.76 V. In the limiting situation at slow scan rates shown in curve (b), the peak for the reoxidation of CoI(sa1en)- is not observed and the two reduction peaks are the same height showing that the process at -1.76 V is also a le process. It is, however, irreversible even at very fast scan rates. It was postulated that the chemistry of the cobalt anion is CoI(sa1en)-+ EtBr --f Et-CoIlI(salen) + Br-(36) and that the peak at -1.76 V was due to Et-Co*II(salen) + e -Et-CoII(salen)-(37) where the CoII organometallic is highly unstable. A cyclic voltammogram on 88 G. Costa, A. Puxeddu, and E. Reisenhofer, J.C.S. Dalton, 1973,2034.Application of Electrochemical Techniques to Study of Chemical Reactions Et-CoIII(salen) under the same conditions confirmed this proposal since it gave an irreversible le reduction at =-1.76 V. Costa et aZ.33 have found a rate constant for reaction (36) from the dependence of i,”/i,”on v: it is 7.6 x 102 1mol-1 s-1. The Isomerization of the Anion Radical of Diethyl Maleate.-Curve (a) of Figure 7 shows a cyclic voltammogram for the trans-isomer, diethyl fumarate (DEF) in I=0 vs.Ag wire electrodes Figure 7 Cyclic voltammograms run at a hanging mercury drop electrode in DMF4.1M tetrabutylarnmonium tetrafluoroborate. Potential scan rate 1.O V s-l. (a) 4x~ 10-3M diethyl fumarate; (b) ----4 X lOP3MdiethyZ maleate, first sweep; (c) as (b)but second cycle.DMF. At fairly rapid potential scan rates the i-E response has the properties of a reversible le system DEF +e +DEF-(38) although at slow potential scan rates the ratio i:/ii becomes less than unity due to a reaction which has been sho~n~~l3~ to be dimerization with a rate constant 34 mol 1-1 s-l. The cyclic voltammetry of the cis-isomer, diethyl maleate (DEM), is more complex. The reduction is generally irreversible and only at very fast scan rates is there any sign of reversibility. However, at intermediate sweep rates curve (b), 8q A. J. Bard, V. J. Puglisi, J. V. Kenkal, and A. Lomax, Discuss. Faraday Sol., 1973,56,353. V. J. Puglisi and A. J. Bard, J. Electrochem. SOC.,1972,119, 829, 833; 1973, 120, 748 Pletcher the reverse sweep, shows a reoxidation peak which corresponds in potential to the oxidation of DEF-.The second sweep also shows a new reduction peak which corresponds to the potential for reaction (38). Thus the cyclic voltammetry indicates the scheme DEM + e DEM-(39) DEMO---f DEF*-(40) The rate constant for reaction (40)wasfound to be ca. 10 s-l from a study of the magnitude of the peaks for system (38) as a function of v. Other experiments (see below) show that DEM-can also dimerize and that it does so more rapidly than DEF-. The Oxidation of Tertiary Aromatic Amines.-Adams and co-workers36J7 have used cyclic voltammetry to study the oxidation of a wide range of tertiary aromatic amines, paying particular attention to the stability and reactions of the cation radicals.A typical example is tri-o-methoxyphenyJamine(TOMPA) in acetonitrile. At fast potential scan rates the cyclic voltammogram is simple, showing a single couple with the properties of a reversible le process. -e TOMPA + TOMPA-+ (41) At slower potential scan rates the curves are more complex (see Figure 8). Firstly the reverse peak for reaction (41) diminishes and simultaneously it may be noted that i,F/v+is no longer independent of v and, indeed, increases as the scan rate is slowed down. At the same time two new reduction peaks appear at less positive potentials and the second cycle becomes different from the first. These observa- tions indicate that the cation radical undergoes a chemical reaction and that the OMe OMe \-6 N/t>-ZH+ 2TOMPA + HMH OMe OMe 36 E. T.Seo, R. J. Nelson, J. M. Fritsch, L. S. Marcoux, D. W. Leedy, and R. N. Adams, J. Amer. Chem. SOC.,1966, 88, 3498. 37 R. F. Nelson and R. N. Adams, J. Amer. Chem. SOC.,1968,90, 3925. Application of Electrochemical Techniques to Study of Chemical Reactions 'A i= 0 I I I ------+E +Om6 + 0.8 +1.0 vs.aqu. S.C.E. Figure 8 Cyclic voltammogram at a Pt electrode in acetonitrile-0.1 M tetraethylammoniumperchlorate. Potential scan rate 0.05 V s-' (a) 10-3M tri-o-methoxyphenylamine,first cycle; (b) as (a) but second cycle product is oxidized at a lower potential than the parent amine. The reaction scheme (43) was postulated, and it was found that the cyclic voltammogram for HMB had two reversible systems whose potentials coincided with the peaks of curve (a).The dependence of i:/v+ on v arises because when the potential scan through the peak is fast compared with the half life of TOMPA-+, the overall process involves the transfer of only one electron. At the other limit when the scan is slow compared with this half life, three electrons are transferred. Under these latter conditions such processes are widely known as ecee processes (e denotes electron transfer, c chemical reaction). For an ece process, i,"/i,"changes very rapidly with v in the situation between the limiting le and 3e cases. The rate constant for the dimerization of TOMPA*+ may be obtained from such a study and suitable working curves: it is 6 1 mol-l s-1.In all the three examples discussed above, interpretation is assisted by the observation of a reversible le process under certain conditions. Systems involving only slow electron transfers (for such couple dEp > 60/n mV and both dE, and EpF are dependent on v), multiple electron transfers (for a reversible ne process dEp = 59/n mV), and totally irreversible electrode reactions can also be investigated, but the interpretation must be viewed more cautiously and supplementary evidence will be essential. Cyclic voltammetry has also been Pletcher widely used for the study of more complex systems, e.g. the chemistry of vitamin B1 2,38 catecholamines,39 and dopaq~inone.~~ Thus cyclic voltammetry is probably the best electrochemical technique for the qualitative investigation of the chemical reactions of species which may be formed by electron transfer and whose half-life lies in the range 1 ms to 100 s.6 Potential Step Techniques In such experiments the potential of a stationary micro-electrode is changed instantaneously and the resulting i-t transient as the system relaxes to the steady state is monitored. In chronoamperometrylOJ1 the i-t transient is recorded directly and analysed in that form; it is, however, sometimes advantageous to integrate the current with respect to time (electronically) and to present the transient as a plot of charge (4)vs. time. The technique is then known as chrono- coulornetry.ll Potential step techniques are very poor for investigating the mechanism of homogeneous chemical reactions coupled to the electrode reaction ;all mechan-isms give transients which look qualitatively similar and such an approach would require theoretical analysis to obtain the expected i-t response for all possible mechanisms, followed by curve fitting.On the other hand, if the mechanism is known, they offer a simple method for obtaining quantitative rate data. The method has the advantages over cyclic voltammetry in such studies that it is conceptually and commonly experimentally simpler. Certainly the mathematical analysis is easier since the differential equations describing the experiment have time-independent boundary conditions and the electrode potentials may be selected so that the electrode surface conditions are simple.The i-t transients may be monitored over a period of a few seconds or on a much shorter timescale of a few ps. Since the portion of the i-t transient analysed corresponds to the timescale on which the chemical reaction is observed, it is generally the case that the study of fast reactions requires measurements at short times and some sophisticated innovations have been introduced to overcome the non-Faradaic current observed at times < 0.5 ms.4l Furthermore it should be noted that at t = 0, the rate of diffusion of a species to the electrode is infinite (see Figure 2) and the technique for obtaining the rate constant for very fast reactions usually involves extrapolation of the i-t data to t = 0.Single potential step experiments may be used to obtain kinetic constants for systems where the chemical reaction precedes the electron-transfer process [e.g. equation (16)] and systems where the product of the primary electron-transfer process is unstable and leads to a species which is further oxidized or reduced at the same potential [e.g. equations (17) and (18)]. Generally the potential will be stepped from a value where the electrode reaction does not occur at a measur- able rate to a value where all the electroactive species reaching the surface under- R. L. Birke, G. A. Bryden, and M. F. Boyle, J. Electroanalyt. Chem., 1974,52, 233. s9 M. D. Hawley, S. V. Tatawawadi, S. Piekarski, and R. N. Adams, J. Amer. Chem. SOC., 1967, 89, 447.40 T. E. Young, J. R. Griswold, and M. H. Hulbert, J. Org. Chem., 1974,39, 1981. 41 H. W. Nurnberg, in ‘Polarography, 1964‘, ed. G. J. Hills, MacMillan, London, 1966. Application of Electrochemical Techniques to Stdy of Chemical Reactions goes electron transfer. Thus, for example, chronoamperometry was the method used to obtain the numerical values for the rate constants for the dimerization of the tertiary amine cation radicals [equation (42)];37 other systems which have been studied include the protonation of carboxylate anions,42 complex formation by metal ions,43 catalytic and the dehydration of p-hydroxyphenyl- hydro~ylamine.~~ When the species to be studied is formed by the electrode reaction and the product of the chemical reaction is not electroactive, the kinetics of the chemical k R-P (45) reaction must be obtained by a double potential step experiment. The electrode potential is first stepped from where no electrode reaction occurs to a value where 0 -+ R occurs at a diffusion-controlled rate.After a time interval, T s, the electrode potential is again stepped, this time to a value where the reverse process R 30 is diffusion controlled (note that the electrode reaction need not be electrochemically reversible). The first pulse may be thought of as a period of generation of the reactant, during which a defined concentration profile for R (sensitive to k) will be set up, and the second pulse as an analysis of this profile. Figure 9 shows a typical i-t response for a double potential step experiment.Clearly the current at any particular time t during the second pulse is a reflection of the amount of R remaining in solution and hence the rate constant, k and its value, should be obtainable from a dimensionless plot of say, &+,/it vs. log t/q. The double potential step method has been used to study, for example, the benzidine rearrangement,46 the hydrolysis of 1,4-benzoquinone 4-imine,47 the disproportionation of the technetate ion,48 and the hydration of ascorbic acid.49 7 The Rotating-disc Electrode A rotating-disc eIe~tr~de~~J~~~~~~~ consists of a smooth disc of the electrode material surrounded by a cylindrical sheath of an insulator; the face of the sheath should be flush with the surface of the disc and the radius of the sheath should be large compared with that of the disc.The complete structure is rotated about an axis perpendicular to the surface of the disc. In experiments with such rotating electrodes, the principal mode of mass transport of species to and from the electrode surface is convection rather than diffusion as is the case with experiments at stationary electrodes in unstirred ra P. Delahay and S. Oka, J. Amer. Chem. SOC., 1960,82, 329. 43 H. Gerischer, 2.Elektrochem., 1960,64,29. I4P. Delahay and G. L. Stiehl, J. Amer. Chem. SOC., 1952,74, 3500. 46 G. S. Alberts and I. Shain, Analyt. Chem., 1963, 35, 1859. 48 W. M. Schwarz and I. Shain, J.Phys. Chem., 1965, 69, 30. 47 A. J. Bard and S.V. Tatawawadi, J.Phys. Chem., 1964, 68, 2676. 40 G. Kissel and S. Feldberg, J. Phys. Chem., 1969, 73, 3082. IsS. P. Perone and W. J. Kretlow, Analyt. Chem., 1966, 38, 1760. so A. C. Riddiford, Adv. Electrochem. and Electrochem. Eng., 1966, 4, 47. s1 V. G. Levich, ‘Physiochemical Hydrodynamics’, Prentice Hall, New Jersey, 1962. Pletcher iF is Figure 9Generalized i-t response to a double potential step experiment. The detailed form of the i-t response will depend on both the mechanism and kinetics of the homogeneous chemical reactions for the species produced during the A pulse. solution. The rate of convection is dependent on the rotation rate of the disc, W, and hence the rate of rotation can be employed to control the supply of the electroactive species to the surface and the removal of the product and, in con- sequence, the residence time of species at the electrode surface.In general the complete steady-state i-E curve is obtained as a function of w,although in some experiments it is only necessary to investigate the steady-state current as a function of w at one particular potential. The i-E curve at one rotation rate is very much the equivalent of a polarogram at one droptime and the Ei, Ef -Ei, and iL values may be used in the study of Application of Electrochemical Techniques to Study of Chemical Reactions homogeneous chemical reactions in an analogous way. In most instances, however, use is also made of w as a variable. In the case of chemical reactions which precede the electron-transfer process [e.g.equation (16)], the i-w dependence is investigated at a potential on the limiting current plateau for the electrode process. If the chemical reaction is very slow, the current will be independent of w while, conversely, if it is fast the current will be determined only by the rate of convection and it will be found that i E d.In the intermediate case, the current is determined partly by con- vection and partly by the kinetics of the chemical step (i vs. 0%will be non- linear) and the pure kinetic information is obtained by extrapolating the i-w data to conditions where the rate of convection is infinite, i.e. w -+co.This extrapolation may be carried out employing a plot of l/i vs. l/w* when theory shows that the intercept on the l/i axis is simply related to the rate constant for the chemical process.Systems studied in this way include the dehydration of PhCH(OH)252 and the dissociation of weak a~ids~~p~~ using rotating lead electrodes. The rotating-disc electrode has also been used in the study of ece systems [e.g.equations (16) and (17)].In such studies the i-w data are usually presented as a plot of i/wh vs. o.At slow rotation rates the residence time of any inter- mediate produced at the electrode is relatively long and it has the maximum opportunity to react chemically to form the species which can be oxidized or reduced further. Conversely, at fast rotation rates the intermediate has a greater opportunity to escape the electrode system before the chemical reaction can occur.Thus, i/d, which is a direct reflection of the number of electrons trans- ferred, will decrease with increasing rotation rate. Ideally the i/w%vs. w plot will be S-shaped with limiting situations at low and highw, and the rate constant for the chemical reaction may be obtained from the transition region at inter- mediate w. Examples of reactions studied in this way include the reaction of the diphenylanthracene cation radical with a set of pyridine~~~ and the dehydration of p-hydro~yphenylhydroxylamine.5~ 8 The Rotating Ring-Disc Electrode In order to investigate the chemistry of species which are produced in an electrode process, we have already seen that it is generally necessary to make effectively two separate measurements, e.g.double potential step, cyclic voltammetry. In experiments with the rotating ring-disc electrodes’ the measurements are made at different electrodes. The electrode structure consists of a central disc electrode surrounded by a concentric insulating annulus, a concentric ring electrode and finally an insulating sheath. The ring and disc electrodes may be controlled 62 R. G. Barradas, 0.Kutowy, and D. W. Shoesmith, Canad.J. Chem., 1974,52, 1635. 83 D. Jahn and W. Vielstich, 2.Elektrochem., 1960, 64,43. 64 W. J. Albery and R. P. Bell, Proc. Chem. SOC.,1963, 169. O6 G. Manning, V. D. Parker, and R. N. Adams, J. Amer. Chem. SOC.,1969,91,4584. 88 P. A. Malachesky, L. S.Marcoux, and R. N. Adams, J. Phys. Chem., 1966,70,4068.67 W. J. Albery and M. L. Hitchmann, ‘Ring Disc Electrodes’, Clarendon Press, London, 1971. 492 Pletcher separately and, generally, a reactant is generated at the disc and the amount of it arriving at the ring is monitored. It may be noted that the time interval for transport of the species from the disc to the ring is a function of w and the width of the insulating annulus. Commonly there are three simple controlled potential experiments which are carried out with a rotating ring-disc electrode as a function of rotation rate. (i) The potential of the disc is held constant at a potential where a reactive species is formed and a voltammogram is run at the ring, i.e. an i-E curve is obtained for the reactive species as a function of the rotation rate.(ii) The potential of the disc is varied linearly and the idisc-E curve is recorded, while the potential of the ring is held constant at a value where a particular species is electroactive. By this experiment it is possible to delineate the potential regions where the species is produced at the disc. (iii) The potentials of both the disc and ring are held constant and the current at both is recorded as a function of the rotation rate. The first two types of experiment are useful in the qualitative investigation of reaction mechanisms and jointly allow a similar range of information to cyclic voltammetry. Indeed additional information is sometimes available. The latter experiment is that normally used to obtain the quantitative kinetic data.Even if the species formed at the disc is completely stable, not all of it will arrive at the ring and hence it is necessary to define the collection efficiency, the fraction of the species formed at the disc which reaches the ring; the collection efficiency, NO,is given by the ratio iring/idisc. Clearly, if the species formed at the disc undergoes a chemical reaction the collection efficiency will be less. In both cases it depends both on the gap between the disc and ring electrodes and the rotation rate of the electrode assembly. It is possible to construct ring-disc electrodes suitable for the study of species with a half life in the range 1 ms to 10 s. The rotating ring-disc electrode has been used in the study of the anion radicals of dimethyl fumarate (DEF) and dimethyl maleate (DEM).34Experiments of types (i) and (ii) above were used to confirm the conclusions drawn from the cyclic voltammetry of the two compounds (see earlier section).Further, an experiment of type (iii) could be used to obtain the information that DEM- had a second fast reaction in addition to isomerization to DEF-. In solutions of DEM or DEF, the potential of the disc was held at a value where the anion radicals were produced at a diffusion controlled rate and the potential of the ring was such that both DEF*- and DEM- would be reoxidized. For the solution of DEF the collection efficiency was 0.55, essentially the value for a completely stable intermediate, while for the solution of DEM it was only 0.254.35.Hence in the latter solution, the DEM- disappeared by a reaction whose product was other than DEF.-.This reaction is probably a coupling and it appears that the equivalent reaction of DEF-is much slower. Other reactions which have been studied by this technique include those of Application of Electrochemical Techniques to Study of Chemical Reactions bromine with arsenic(n~),~* ally1 aIcohol,59 and anisole,60 the reactions of the anion radicals of activated ole fin^,^^,^^ and the chemistry of intermediates pro- duced during reduction of p-nitroaniline62 and nitrobenzene63 and oxidation of aniline.64 9 Controlled Potential Coulometry For controlled potential coulometry11~65 a macro working electrode, or rather a cell with a large electrode area to solution volume ratio is employed and the working electrode is placed in a separate compartment to prevent interference from the reactions occurring at the other electrodes.A controlled potential is applied to the working electrode and the i-t response recorded until the current drops to zero, i.e. all the electroactive species are removed from the bulk solution. Controlled potential coulometry is principally used to confirm the overall number of electrons transferred in the electrode reaction; the i-t data may be integrated to give the total charge, q, associated with the electrode process and this is related to the n value by Faraday’s Law. Such electrolysis is also used to confirm that the fhal products of the electrode process, as determined by the normal analytical procedures of organic and inorganic chemistry, are consistent with the mechanistic information found with the other techniques.It should, however, be remembered that the timescale of a controlled potential electrolysis, generally 10-300 min, is much longer than all the experiments discussed above and it is therefore possible for systems with slow homogeneous chemical reactions to give apparently conflicting results. In addition, account must be taken of the increased concentrations of reactants typically used in experiments designed to study products. Controlled potential electrolysis offers a technique for the study of slow chemical reactions.65The i-t data are replotted in the form of log i-t or i-q graphs; such plots are linear for uncomplicated electrode reactions whereas for systems where the product of the electron-transfer step can undergo a slow chemical reaction the plots are more complex.It is common for the concentration of such intermediates to be followed spectroscopically by carrying out the electrolysis in a U.V. or i.r. spectrometer or by sampling the solution and determining the concentration by analysis. 10 Concluding Remarks The aim in this review has been to outline the principles of the most important electrochemical techniques employed for the study of homogeneous chemical reactions. Because of the restriction on space the theoretical equations for each W. J. Albery, S. Bruckenstein, and D.C. Johnson, Trans. Faraday SOC.,1966,62, 1938. W. J. Albeq, M. L. Hitchman, and J. Ulstrup, Trans. Furuday Suc., 1969, 65, 1101. aa W. J. Albery, M. L. Hitchman, and J. Ulstrup, Trans. Faruday SOC.,1968, 64, 2831. a1 W. V. Childs, J. T. Maloy, C. P. Keszthelyi, and A. J. Bard, J. Electrochem. SOC.,1971, 118,874. Oa L. N. Nekrasov, I. P. Ryvkina, and B. G. Podlibner, Soviet Electrochem., 1972,8,1372. 65 W. J. Albery, unpublished results. 64 M. Breitenbach and K.-H. Heckner, J. Electroanalyr. Chem., 1971,33,45. A. J. Bard and K. S.V. Samthanam, Electrmnalyt. Chem., 1970,4, 215. Pletcher of the methods, necessary for quantitative work, have not been presented and the interested reader is referred to the original papers; it must be remembered that the equations will also depend on the detailed mechanism of the homo- geneous chemical reaction.Likewise, details of the instrumentation and glassware and the experimental difficulties and pitfalls are best sought in these references. The reader will note that the review concentrates on controlled potential techniques. The controlled current techniques, chron~potentiometry,ll*~~*~~ reverse current chronopotentiometry,68 and controlled current co~lometry,~~ will often yield the same results as the methods discussed but they never have a marked advantage. Of the other techniques which might have been discussed, thin layer cells,7 0-7g a.c. p0larography~73.74 and spectroelectrochemicl tech- nique~,~~-~~ are perhaps the most important.In the including in situ e.~.r.,~*~7~ spectroelectrochemical methods the intermediate to be studied is formed electro- chemically and monitored simultaneously by an electrochemical and a spectro-scopic technique. They are particularly useful since the optical readout is sensi- tive to the whole of the concentration profile and its variation with time, whereas the current is a reflection only of the gradient of the profile at the surface. Finally I would plead with prospective workers in the field that they attempt to use a number of electrochemical methods and where possible, other tech- niques, on each system. The techniques are complementary and the mechanistic and kinetic conclusions should be uniform! 6~ R. G. Davis, Electroanalyt. Chem., 1966, 1, 157.67 W. H. Reinmuth, Analyt. Chem., 1960, 32, 1514; 1961, 33, 322, 485. e8 H. B. Herman and A. J. Bard, Analyt. Chem., 1964,36, 510,971. II@ J. Janata and H. B. Marks, Electroanalyt. Chem., 1969, 3, 1. 70 C. N. Reilley, Rev. Pure Appl. Chem., 1968, 18, 137. 71 A. T. Hubbard, Crit. Rev. Analyt. Chem., 1973, 201. 72 A. T. Hubbard and F. C. Anson, Electroanalyt. Chem., 1970,4, 129. 73 D. E. Smith, Electroanalyt. Chem., 1966, 1, 1. M. Sluyters-Rehbach and J. H. Sluyters, Electroanalyt. Chem., 1970, 4, 1. 75 T. Kuwana and N. Winograd, Electroanalyt. Chem., 1974, 7, 1. 76 R. F. Broman, W. R. Heineman, and T. Kuwana, Discuss. Faraday SOC.,1973,56, 16. 77 A. Aylmer-Kelly, A. Bewick, P. Cantrill, and A. Tuxford, Discuss. Faraday Soc., 1973, 56, 96. 78 I. B. Goldberg and A. J. Bard, J. Phys. Chem., 1971,75, 3281. 7@ B. Kastening, B. GostiSa-MihelEiE, and J. DiviSek, Discuss. Faruday Soc., 1973, 56, 341. 495

ISSN:0306-0012    

DOI:10.1039/CS9750400471

出版商:RSC    

年代:1975

数据来源: RSC