摘要:

Quarterly Reviews No I Vol 21 1967 Organic Chemistry of Periodates By B. Sklarz The Mechanism of Rusting By U. R. Evans F.R.S. The Chemistry and Biological Function of Isoalloxazines (Flavines) By G. R. Penzer and G. K. Radda Page 3 29 43 Vibration-Vibration Energy Transfer in Gaseous Collisions By J. D. Lambert 67 The Hydrated Electron By D. C. Walker 79 Molecular Polyhedra of High Co-ordination Number By E. L. Muetterties and C. M. Wright 109 The Chemical Society London Quarterly Reviews contains articles by recognised authorities on selected topics from general physical inorganic and organic chemistry. The Journal and Annual Reports interest primarily the research worker Quarterly Reviews is designed for a wider range of readers. It is intended that each review article shall be of interest to chemists generally and not only to workers in the particular field being reviewed. The submission of reviews for publication is welcomed but intending authors are advised to write in the first place to the Editor The Chemical Society Burlington House Piccadilly London W. 1. Such preliminary communications should be accompanied by an outline of the ground to be covered rather than by the completed manuscript Price to non-fellows E4 10s. Od. per annum @ Copyright reserved by The Chemical Society 1967 Published by The Chemical Society Burlington House London. Printed in England by The Thanet Press Margate.

ISSN:0009-2681    

DOI:10.1039/QR96721FP001

出版商:RSC    

年代:1967

数据来源: RSC

摘要:

Organic Chemistry of Periodates By B. SMarz UNIVERSITY OF BIRMINGHAM 1 Introduction Malaprade’s observation1 that mannitol is destroyed by periodic acid and Fleury’s closer definition of the reaction as a specific oxidative fission of 1,2-diols,* launched periodic acid and its salts on a distinguished career among the oxidising agents of organic chemistry. Closely related fissions of a-ketols a-diketones and a-amino-alcohols having been discovered in the 1930’s the reagent became an indispensable tool for the structure determination of the most varied natmal products particularly the carbohydrates and more recently the nucleic acids. The reviews by Jackson3 and by Bobbitt4 fully describe these developments while more recent applications in the carbohydrate and nucleic acid fields will be found in standard w~rks.~n~ Malangeau’s review’ is limited in scope and is in- accessible.The well-established theory and practice of oxidation with periodates particularly in carbohydrate chemistry is described in practical manuals.8 Only a brief outline including recent innovations will be attempted here. In the last decade the organic chemistry of periodate has grown mainly in two directions. The classical glycol cleavage has been studied in detail and a consistent picture of its mechanism is emerging. This has shed light on the oxidation of simple sugars where detailed re-examination of the reactions has revealed some of the intermediates. On the other hand several new well-defined reactions have come to light which extend the use of the reagent in organic chemistry. This Review is devoted principally to these two aspects.The term ‘oxidation’ will refer to reaction with periodic acid or its salts unless otherwise qualified. L. Malaprade Bull. SOC. chim. France 1928 683. P . Fleury and J. Lange Compt. rend. 1932,195 1395. E . L. Jackson Org. Reactions 1944 2 341. J. M. Bobbitt Adv. Carbohydrate Chem. 1956 11 1. W. Pigman (ed.) ‘The Carbohydrates’ Academic Press Ohio 1956; E. Percival ‘Structural Carbohydrate Chemistry’ J. Garnett-Miller London 1962. A. M. Michelson ‘Chemistry of the Nucleosides and Nucleotides’ Academic Press London 1963; see also P. R. Whitfield Biochim. Biophys. Acta 1965 108 202 and references therein. P. Malangeau Miscs au Pointde la Chimie Analytique pure ct appliqued et de l’dnalyse bromato- Zogiquc 1961 9 81. (a) G. F. Smith ‘Analytical Applications of Periodic Acid and Iodic Acid’ G.F. Smith Chemical Co. Ohio 1950; (b) J. R. Dyer Methods Biochem. Analysis 1956,3 11 1 ; (c) R. D. Guthrie ‘Methods in Carbohydrate Chemistry’ ed. R. L. Whistler and M. L. Wolfrom Academic Press London 1962 p. 432 et seq. 3 1. Quarterly Reviews Methods.-Periodate oxidation lends itself particularly to analytical experiments which are often followed by preparative oxidations. Periodic acid and its sodium and potassium salts are the usual reagents. In analytical experiments mixtures containing excess of 0.1-0.01 M-periodate are allowed to react at room temperature or below along with a blank and for long reaction times should be kept in the dark. The photochemistry of periodate solutions which slowly form ozone in daylight has been Organic solvents can be added but may retard some reaction^.^ Tetraethylam- monium periodate which is highly soluble in water and considerably soluble also in several organic solvents including chloroform may be of value in over- coming solubility problems.1° The choice of pH depends on the r e d ~ c t a n t .~ ~ ~ ~ The reaction is usually followed by estimation of the remaining periodate in aliquot portions withdrawn at intervals from the mixture and the blank. The various iodometric titrations are described in the above reviews.8 Methods will also be found there for the estimation of small-molecular oxidation products such as formaldehyde formic acid carbon dioxide and ammonia. The elegant methods of polarographyll have been extended to periodate reactions and in favourable caseslla permit simultaneous estimation of organic products.Periodate can also be estimated by its absorption at 223 mp8c or in more concentrated solution at 260 mp.12 When oxidations are carried out on a preparative scale,86 it may be necessary to remove periodate and iodate ions by precipitation or by ion-exchange if the organic products cannot be extracted. Excess of reagent is sometimes destroyed by adding ethylene glycol or pinacol etc. Several simple compounds frequently used as solvents or appearing as oxida- tion products are hardly attacked by periodate in the dark though more appre- ciably in sunlight .13 They include methanol ethanol and their derived aldehydes and acids. The limited solubility of the commonly used periodates their high molecular weight and the frequen? appearance of iodine are the principal drawbacks to their preparative use.2 Glycol Fission and Related Reactions Introduction.-The oxidative cleavage of vicinal diols is the classical and most widely used reaction of periodate. Oxidation is fast at room temperature except for heavily substituted diols aldehydes and ketones being formed (a; Scheme 1). Analogous cleavage occurs more slowly with a-hydr~xy-carbonyl~~~~ and a-dicarbonyl compounds,1* in which the carbonyl groups appear as carboxyl R. D. Guthrie Chem. and Ind. 1960 691. lo A. K. Qureshi and B. Sklarz J. Chem. SOC. (C) 1966 412. l1 (a) P. Zuman and J. Krupicka CoZZ. Czech. Chem. Comm. 1958 23 598; (b) P. Zuman ‘Organic Polarography’ Pergamon Oxford 1964. l9 J. S. Dixon and D. Lipkin An&. Chem. 1954 26 1092. l3 F.S. H. Head and G. Hughes J . Chem. Sac. 1952,2046. l4 P. W. Clutterbuck and F. Reuter J. Chenr. SOC. 1935 1467. 4 Sktarz functions in the products (b). a-Hydroxy- and a-keto-acids are oxidised rather sl0wly,1~J~ with some exceptions such as glycollic and glyoxylic acids.15 a-Amino-alcohols and a-diamines are likewise oxidised to carbonyl fragments (c) unless one of the amino-groups is tertiary or acylated.17~la These points are discussed in detail later (p. 13). All these reactions are two-electron oxidations requiring one molecule of periodate in which the iodine atom is reduced from the $7 to the + 5 valency state with the formation of iodate ions. Lead tetra-acetate has long been used as a glycol-cleaving reagent.lg It is used in glacial acetic acid or in other organic solvents and oxidises also a-hydroxy- and a-keto-acids and many other structures.20 It is thus a useful complement to rather than ‘a serious c~mpetitor’l~a of periodate.Iodosobenzene diacetate21 resembles lead tetra-acetate as does sodium bismuthate,22 which can be used also in aqueous acid solution. Chromium trioxide has been used for the cleavage of di-tertiary glyc01~,2~a and the mechanism has been Mechanism-A mechanism for the periodate oxidation of glycols suggested as early as 1933 by Criegee% involved the formation of a cyclic periodate complex in analogy with lead tetra-acetate. Much subsequent work has borne out and amplified this idea. The evidence consists of direct observations on the complexes and more extensively of kinetic evidence. The interaction of periodic acid with glycols in water has been observed spectrometrically by an increase in the absorbance of the periodate ion at l6 D.B. Sprinson and E. Chagraff J . Biol. Chem. 1946 164 433. l6 J. E. Courtois and M. Guernet Ann. pharm. franc. 1958 16 119. B. H. Nicolet and L. A. Shinn J. Amer. Chem. SOC. 1939 61 1615. P. F. Fleury J. E. Courtois and M. Grandchamp Bull. SOC. chim. France 1949 88. 19(a) R. Criegee L. Kraft and B. Rank Annalen 1933 507 159; (6) A. S. Perlin Adv. Carbohydrate Chem. 1959 14 9. *O R. Criegee Angew. Chem. 1958,70 173. *l L. K. Dyall and K. H. Pausacker J . Chem. Soc. 1958 3950 and refs. therein. 22 W. Rigby J . Chem. SOC. 1950 1907. 23 (a) M. Uskovic M. Gut E. N. Trachtenberg W. Klyne and R. I. Dorfmann J. Amer. Chem. SOC. 1960,82,4965; (b) J. Rocek and F. Westheimer J.Amer. Chem. SOC. 1962 84 2241. 24 R. Criegee Sitzungsberichte der Gesellschqft fur die Be0 frderrrng der gesamten Wissen- schaften Marburg 1934,69,25 (Chem. Abs. 1935 29 6820). 5 Quarterly Reviews 222.5 mp,a5 while increased absorbance at 226 mp in isopropyl alcohol also suggests some interaction.28 A fall in pH has been observed on addition of ethylene glyc01,2~ other vicinal diolsF8 and certain 1,3-di0ls~~ to periodate solutions. Malaprade30 had already observed this effect and deduced the formation of an 'addition compound' but as to its r81e in the mechanism of the oxidation he could only speculate. Investigating the periodate oxidation of some cyclic monosaccharide com- pounds and cycloalkane polyols Barker and Shaw31 found that several of these compounds rapidly took up one molecule of the reagent which reappeared slowly after the titrations while considerable amounts of substrate (ribose) were recovered.All these compounds had a 1,2,3-cis-triol system capable of existing in a favoured diaxial-monoequatorial conformation from which a terdentate periodate ester (I) was thought to be formed. The other polyols took up periodate rapidly with oxidation. The terdentate complex from 1,2-O-isopropylidene-a-~-glucofuranose has been studied by nuclear magnetic resonance (n.m.r.).32 Its stability in alkaline solution is such that the monosaccharide is not oxidised. Probably the triesters cannot break down directly to oxidation p r o d u ~ t s ~ ~ ~ ~ ~ and any oxidation involves prior hydrolysis to the diester. Despite this stability the terdentate esters have not been isolated.The cyclic diesters are much less stable and the principal evidence for their structure and function in glycol fission comes from kinetic Early experiments had shown that the oxidation of ethylene glycol and of pinacol have different kinetic forms.35 A two-stage reaction course ( d ) was later s ~ g g e s t e d ~ ~ . ~ ~ in which the glycol (G) reacts reversibly with periodate (P) to give an intermediate (X) which breaks down to the products. The rate constants ka k (d) G + P +X-+ products kb k, kb and k and the equilibrium constant K for formation of X have been evaluated for a series of simple glycols in an elegant set of studies by Bunton and his colleagues. Taking the complex equilibria of ionisation and hydration of periodate3' into account they showed that the complex X undergoing breakdown must 25 G.J. Buist C. A. Bunton and J. H. Miles J. Chem. Soc. 1957 4575. 26 J. Klaning and M. C. R. Symons J . Chem. SOC. 1960,977. 27 G. J. Buist and C. A. Bunton J. Chem SOC. 1954 1406. 28 G. J. Buist C. A. Bunton and J. H. Miles J. Chem. SOC. 1957 4567. 29 J. L. Bose A. B. Foster and R. W. Stephens J. Chem. SOC. 1959 3314. 30 L. Malaprade Bull. SOC. chirn. France 1934 883. 31 G. R. Barker and D. F. Shaw J. Chem. SOC. 1959 584. 32 A. S. Perlin and E. von Rudloff Canad J. Chem. 1965 43 2071. 33T. P. Nevell Chem. and Znd. 1959 567. 34 C. A. Bunton Ann. Reports 1959 56 185. 35 C. C. Price and H. Kroll J. Amer. Chem. SOC. 1938 60 2726; C. C. Price and M. Knell ibid. 1942 64 552. 36 F. R. Duke and V. C. Bulgrin J. Amer. Chem. SOC. 1954 76 3803.37 C. E. Crouthamel A. M. Hayes and D. S. Martin J. Amer. Chem. SOC. 1951 73 82. 6 Sklarz be a mono-anion written as (11) or its dehydrated form (III).27 The complex is a stronger acid than periodic acid (cf. borate and tellurate complexes). The stability of X is subject to electronic and steric influences. K increased in going from ethane- through propane-1,2- to (-)-butane-2,3-diol i.e. with increasing methyl substitution and consequent electron availability at oxygen but the sharp drop in the equilibrium constant for meso-butane-2,3-diol is a steric effect.28 A model (Fig. 1) based on atomic dimensions was suggested in which methyl groups interfere sterically with the octahedral periodate oxygen atoms when placed at the hindered positions H,H’ but not at the free positions F,F’.The rate of formation k was determined spectrometrically for some diols of the series and decreased with increasing substitution for steric reasons. Under certain conditions the initial esterification was relatively slow followed by fast cy~lisation.~~ The rate of collapse of X was also estimated and increased with methyl sub- stitution so that steric crowding is probably the dominant factor although electronic effects (hyperconjugation) may also operate.28 The importance of this work lay in the separate evaluation of the various constants for the two reactions steps. With the help of these the difference in over-all kinetics between ethylene glycol and pinacol could be explained. In the former and other lightly substituted diols K is large i.e. the intermediate accumulates and its collapse (k) is rate-determining.The rate of oxidation is greatest near neutrality when the concentration of mono-anion is maximal but decreases at higher pH where the stable dianion is formed but cannot break down. For hindered diols such as pinacol K is small and formation of the complex appears to be rate-determining (ka) leading to overall second-order kinetics. The influence of pH is also more complex and earlier reports are in c o n f l i ~ t . ~ ~ ~ ~ * According to a detailed study by Bunton and his colleagues the oxidation of pinacol is subject to general acid-base catalysis of the ring-closing step39 ammonia being a good catalyst.40 2-Methylbutane-2,3-diol lies on the mechanistic borderline with kinetics which depart from first order between pH 4 and 5 and can be interpreted on@ in terms of an intermediate X.4l Acyclic threo-diols are very generally oxidised more quickly than the erythro- 38S.Senent and P. Escudero Andes real SOC. espaii. Fis. Quim. 1961 57,B 153; Chem. Abs. 1961 55 21761. 39 G. J. Buist C. A. Bunton and J. Lornas J. Chem. SOC. (B) 1966 1094 1099. 40 C. A. Bunton and M. D. Carr J. Chem. Soc. 1963 5854 (p. 5860). 41 G. J. Buist and C. A. Bunton J. Chem. SOC. 1957,4580. 7 Quarterly Reviews isomer and relative configurations have been assigned on this basis.42 Polyols are likewise oxidised first at a threo-diol group glucitol (V) for instance giving glyceraldehyde and D-erythrose when oxidised with limited amounts of perio- date.qs A more recent study has revealed the detailed sequence of oxidation of all the diol pairs in [14C]glucitol [(3,4) > (2,3) > (43) > (5,6) > (1,2)].44 A neat rationalisation of these results was possible in terms of the cyclic complex (IV).Studies with cyclic 1,Zdiols have further clarified the structure of the com- ~ l e x . 4 ~ ~ cis- and trans-Cyclopentanediols (VI; R = R’ = H or CH,) undergo second-order (‘pinacol-type’) oxidation which is faster for the cis compounds. The trans-diol (VI; R ing in the complex. = R’ = CH,) is not attacked at all owing to severe crowd- CHO __c CH,OH cis-Cyclohexane-l,2-diols are oxidised faster than the trans- (diequatorial) isomers and the trans-diaxial diol group is not oxidised.46 In a detailed analysis of the ‘ethylene glycol’ kinetics of cis- and trans-cyclohexanediols Bunton and P 0 Oxygen 0 Carbon FIG. 1. [Reproduced by permission from ref.281 Model of the intermediate complex (two oxygen atoms omitted). 42 P. Zuman J. Sicher J. Kiupicka and M. Svoboda Coll. Czech. Chem. Comm. 1958 23 1237. 43 J. C. P. Schwartz J. Chem. SOC. 1957,276; J. E. Courtois and M. Guernet Bull. Sot. chfm. France 1957 1388. 44 D. H. Hutson and H. Weigel J. Chem. SOC. 1961 1546. 45 (a) V. C. Bulgrin and G. Dahlgren J . Amer. Chem. SOC. 1958 80 3883; (6) C. A. Bunton and M. D. Carr J . Chem. SOC. 1963 770. 46 J. Honeyman and C. J. G. Shaw J. Chem. SOC. 1959 2454. 8 Sklarz his colleagues47 showed that the trans-fused complex is actually more stable than the cis-fused at least at pH 9 and that the higher overall oxidation rate of the cis-diol is due to the faster breakdown of the cis-fused complex (VIII). The flex5bility of the five-membered ester ring is such that the trans-diequatorial junction with the cyclohexane ring is preferred (VII; Fig.2). At the same time FIG. 2. oxygen atoms ‘a’ to iodine are in the plane of the paper. are out of the plane of the paper ‘a” being away from the observer. [Reproduced by permission from ref. 471 (VII) Intermediate complex of the trans-diol (equational conformation); bonds joining (VIII) Intermediate complex of the cis-diol; bonds joining oxygen atoms ‘a’ ‘a” to iodine In both models two of the oxygen atoms joined to iodine are omitted. there is less non-bonding interaction between the periodate oxygen and ring- carbon atoms. Methyl groups increase these interactions and lead to second- order kineti~s.~~b a-Glucopyranose with a cis-diol function at C( 1)-C(2) is oxidised at this position appreciably more quickly than the ,L?-an~mer.~~ Rigid trans-diaxial diols such as truns-decalin-9,1 O - d i 0 1 ~ ~ G are not oxidised ; a cyclic intermediate is clearly impossible.Ditertiary alcohols of this type are cleaved by lead tetra-a~etate,4~~$~~ an observation which among others prompted the suggestion by Levesley Waters and Wright51 that monoesters undergo breakdown in glycol fission. This is certainly excluded for periodate since the monoester of an erythro-diol or of a trans-diaxial cyclic diol should then favour cleavage by trans elimination which is contrary to ~bservations.~~ The stereo- chemical requirements of lead tetra-acetate differ somewhat from those of periodate,5O as seen in the differing initial action on sucr0se.5~ The oxidative cleavage of a-diketones has received relatively little attention.Shiner and Wasmuth5* have shown that it is base-catalysed and of second-order and have postulated a cyclic intermediate formed by nucleophilic attack of the various periodate anions on the carbonyl groups. The glycol-periodate esters arise from electrophilic attack on the diols. 47 G. J. Buist C. A. Bunton and J. H. Miles J . Chem. SOC. 1959 743. 48 S. J. Angyal and J. E. Klavins Austral J. Chem. 1961 14 577. 49 S. J. Angyal and R. J. Young J . Amer. Chem. SOC. 1959 81 (a) 5251 ; (6) 5467. 5 o R . Criegee E. Hoger G. Huber P. Kruck F. Marktscheffel and H. Schellenberger Annalen 1956 599 81. 51 P. Levesley W. A. Waters and A. N. Wright J . Chem. SOC. 1956 840. 52 K. B. Wiberg and K. A. Saegebarth J . Amer.Chem. SOC. 1957 79 2822. 53 A. K. Mitra and A. S. Perlin Canad. J. Chem. 1959 37 2047. 54 V. J. Shiner and C. R. Wasmuth J . Amer. Chem. SOC. 1959 81 37. 9 Quarterly Reviews Support for these views has come from isotopic labelling experiment^.^^ When pinacol and 2-methylpropane-l,2-diol were oxidised in 180-enriched water no label was found in the acetone. Thus oxygen atoms of periodate (which exchange rapidly with water5s) do not become linked to carbon (e; Scheme 2). Experi- ments with [180]biacetyl and with methylacetoin were more difficult to interpret but showed that here periodate oxygen atoms were being linked to carbon. The w * -* 0 * o 0,11 HO-CMe -0 I O-CMe2 Me,CO 8 I 0-CMaR MQ€OR 0 11 HO-CMeR - \\A 1 ___) * '\ * //I I O* *OH Me2C0 6 H 6 $ * - I y 0 II C*.O-CMe -* II,O-CM% - 0-1 I _Ic C=% II'g-T-OH MeCGo* 9 Me 'OH Me (4 O=f\o..3 Scheme 2 mixed nucleophilic and electrophilic r6les of the reagent were strikingly observed with methylacetoin for the acetic acid produced was labelled while the acetone was not (f).The breakdown of the cyclic intermediates in glycol oxidation with periodate and chromatez3 b have been discussed from a theoretical viewpoint.57 Stable glycol chelates have been obtained with some antimony (v) compounds which on pyrolysis give the glycol fission products.58 Contrary to an earlier suggestion glycol oxidation does not involve free radicals methyl methacrylate not being polymerised in the reaction mixture.59 The graft-polymerisation of acrylonitrile on cellulose which is undergoing periodate oxidation is thus surprising but may be due to secondary free-radical reaction of the many aldehyde groups which do in fact terminate polymerisa- tion.60 A study of glycol fission by electron spin resonance is lacking.Intermediates in the Oxidation of Polyo1s.-Because of the importance of periodate oxidation in the carbohydrate field the intermediates arising in the oxidation of various polyols have been examined in detail. We have already referred to the preferential cleavage of certain bonds for steric reason (p. 8). 210; CH20 2HC02H - 55 C. A. Bunton and V. J. Shiner J. Chem. SOC. 1960 1593. 56 Cf. M. Anbar and S . Guttmann J. Arner. Chem. Soc. 1963,83 781. 57 M. C. R. Symons J. Chem. SOC. 1963 4331. 58 F. Nerdel J. Buddrus and K. Hoher Chem. Ber. 1964 97 124. 59 H. Tanabe Chem. Pharm. Bull. (Tokyo) 1960 8 365 (Chem.Abs. 1961 55 10307). 6o T. Toda J . Polymer Sci. 1962 58 411. 10 Sklarz In the complete degradation of glucose five molecules of the reagent are required and formic acid (five molecules) and formaldehyde (one molecule) are formed but the reaction does not proceed at a uniform speed. At pH 3-5 distinct stages in the uptake of reagent and liberation of fragments were observed,61t62 and p-formylglyceraldehyde (IX) was isolated from the reaction mixture.g3 Glucose is oxidised mainly in the pyranose form but exists to about 15 % in the acyclic and furanose forms as estimated from the initial release of formaldehyde by oxidation of the 5,6-b0nd.~~ The formyl ester (IX) is relatively stable at pH 3.6 but is quickly hydrolysed at pH 7 where the oxidation is smooth and complete.Consideration of these esters can be useful in structural studies. Amylose consists of largely unbranched chains of 1,4-a-linked glucose units (X). Periodate oxidation (Scheme 3) releases one molecule of formic acid from each terminal residue by direct fission and a further molecule by subsequent hydrolysis of the formyl ester at the reducing end. The total amount thus represents one un- branched molecule of amylose whose length can then be calculated given the sample weight. Conflicting results were obtained as to the release of the third molecule of formic acid from a simple model compound maltose (X; n = O).64 To avoid the uncertainty Wolff et al.,S5 working with a corn amylose first estimated the formic acid released immediately and then after destroying excess of periodate hydrolysed the intermediate ester (XI) and estimated the additional formic acid.This procedure gave a separate estimate of the number of reducing residues in the amylose molecule which was thus shown to consist on average of two branches. The chemistry of the polyaldehydes produced in the oxidation of poly- saccharides has been reviewed.ss Carbohydrates tend to reduce more periodate than expected from mere glycol fission.67 The new reaction was shown to be a hydroxylation of CH- groups activated by adjacent carbonyl and is discussed further later (p. 15). An a-alkoxy-malondialdehyde moiety (XII) arises from the 4-O-substituted terminal residue of polysaccharides such as (X). The activated CH-group of the largely enolised dialdehyde (XII) is hydroxylated by periodate to give a hemi- acetal (XIII).The formation and breakdown of such intermediates has been studied with model compounds (XII; R = CH or C6H5.CH2).69 At pH 3.6 where (XIII) is most stable to hydrolysis oxidative cleavage predominates 61 F. S. H. Head Chem. and Ind. 1958 360; S. A. Warsi and W. S. Whelan ibid. p. 71. 62 L. Hough T. J. Taylor G. H. S. Thomas and B. M. Woods J. Chem. SOC. 1958 1212. 64 K. H. Meyer and P. Rathgeb Helv. Chim. Acta 1943 31 1545; A. L. Potter and W. Z. Hassid J . Amer. Chem. SOC. 1948 70 3489. 65 I. A. Wolff B. T. Hofreiter P. R. Watson W. L. Deatherage and M. M. Machfasters J. Amer. Chem. SOC. 1955 77 1654. 66 R. D. Guthrie Adv. Carbohydrate Chem. 1961 16 105. 67 T. G . Halsall E. L. Hirst and J. K. N. Jones J . Chem. SOC. 1947 1427. 68 C. F. Huebner S. R. Ames and E. C.Bubl J . Amer. Chem. Soc. 1946 68 1621. 69 (a) J. C . P. Schwartz and M. MacDougall J . Chem. Soc 1956 3065. (b) M. Cantley L. Hough and A. 0. Pittet Chem. and Ind. 1959 1126 1253. C. Schopf and H. Wild Chem. Ber. 1954 87 1571. 11 QuarterZy Reviews -0. ?> H O n -H CHO I CHO RO*YH (XI I) $H,OH HCO,H H-CO,H + C ~ R 3% etc. I 10; CHO i ,!?$& CHO I 04- (XI I) - RO-?-OH YHO (XIII) CHO CHo ROH + 70 L- atc. Scheme 3 giving formic acid and an ester which is slowly hydrolysed further and oxidised. At other pH's hydrolysis of (XIII) occurs followed by oxidation of mes- ~xalaldehyde.~~ In the later stages of carbohydrate oxidation the above hydrolyses and hy- droxylations varying independently with pH are superimposed to produce a complicated pattern of reactions. This has been resolved for some simple mono- and di-saccharides by kinetic cstimation of the various fragments at different aciditie~.~~*~l Under the acidic conditions used in early studies over-oxidation was found to be associated with the liberation of iodine.g7 A method for the detection and estimation of 1,4-links in polysaccharides was even devised on the basis of this fact.72 It has since been shown that sodium iodate in acid solution oxidises benzyloxymalondialdehyde the iodide ion formed being re-oxidised to Over-oxidation is usually undesirable since it confuses the number of true 70 M.Cantley L. Hough and A. 0. Pittet J . Chem. SOC. 1963 2527. J. Chem. Soc. 1954 603. 73 K. Ahlborg Svensk Kem. Tidskr. 1942 54 205; Chem. Abs. 1944 38 4254. L. Hough and B. M. Woods Chem. and Ind. 1957 1421 ; F.S. H. Head and G. Hughes 12 Sklarz glycol-cleavage steps from which structural deductions are most easily made. Low temperatures (5 and reduced concentrations of periodatesb have been recommended to minimise it. In oxidation mixtures exposed to sunlight less specific hydroxylations involv- ing singly-activated CH-groups can occur and oxidations of polyols should therefore be carried out in the dark.74 Amino-alcohols.-The periodate cleavage of N-primary and N-secondary a-amino-al~ohols~~J* (p. 5) is fastest at pH 7-8 and is suppressed in mildly acid solution. (The oxidation of hydroxy-amino-acids is discussed separately below.) NH p 10,- N% CH2W R ( 3 c H 2 M HrOs- R-C~;CH~LR' - c% - + R % COR' (x I v> Scheme 4 4-Hydroxymethyloxazolines (XIV) are oxidised in a two-stage reaction (Scheme 4).75 Acid-catalysed ring opening produces an a-amino-alcohol which is cleaved in slightly alkaline solution.The tertiary amino-group does not prevent cleavage entirely.ls Oxidation of desosamine (XV; R = H) gave in turn the tetrose (XVI) and acetaldol and the methyl acetal (XV; R = CH,) also took up one molecule of the reagent.76 However erythromycin and erythralosamine parent glycosides of 0 gave" the N-oxide (p. 23). An electron-withdrawing substituent on nitrogen such as the acetyP7 or 2,4-dinitrophenyl retards the cleavage which can however proceed in the latter case if there is a second hydroxyl group vicinal to the first. Thus the arylaminoethanol (XVII) resists oxidation but the 3-arylaminopropane-1,2-diol (XVIII) is cleaved giving 2,4 dinitr~aniline.'~ Studies on the oxidation of 2-amino- and 2-acetylamino-2-deoxyglucose 73 Ref.16 in M. Cantley et al. ref. 70. 74 F. S. H. Head J. Text. Znst. 1953 44 ~ 2 0 9 ; Chem. Abs. 1953 47 8378. 76 H. L. Wehrmeister J . Org. Chem. 1961 26 3821. 76 R. K. Clarke Antiobiotics and Chemotherapy 1953 3 663. 77 E. H. fly^ M. V. Sigal jun. P. F. Wiley and K. Gerzon J. Amer. Chem. SOC. 1954 76 3121. K. Hattori H. Harada and Y . Hirata Bull. Chem. SOC. Japan 1962,35 312. 13 Quarterly Reviews (XIX; R = H or Ac) (Scheme 5 ) and the corresponding methyl glucosides confirmed that N-acetylation prevents the aminoalcohol c l e a ~ a g e . ~ ~ ~ ~ ~ The subsequent reactions involving hydrolysis of (XX) and the hydroxylation of 2,4-(02N)2C,H3*NH.CH,*CH20H 2,4-(0,N),C,H3.NH*CHe*CH.CH,0H I (XVII) OH (XVIII) aminomalondialdehyde derivatives (XXI; R = H or Ac) have been elucidated.61 CHO CH,OH CHO CHO -% H t N H R 3.H O t H (xx 0 NHR NHR (x 1x1 (xx) Scheme 5 There have been few studies on the kinetics and mechanism of the fission. It is accderated by increasing pH and probably involves the unprotonated amine.s2~83 With rare exceptions only second-order kinetics have been observeds4 which permit of no deduction as to an intermediate. The kinetic form may be due to the very dilute solutions used by the Czech workers.28p36 Evidence for a cyclic intermediate is again stereochemical. In a series of acyclic a-amino-alcohols the threo-isomers were consistently oxidised more quickly than the erythuo-compounds. This difference was enhanced by N-methylation which when followed by periodate oxidation at pH 6.5-7.0 was recommended as a method for establishing the relative configurations of diasteroisomeric cu-amino-alcohols.84 N-Benzylation retarded the oxidation of the threo-compounds to such an extent that the usual order could be reversed.These effects are interestingly rationalised by the authors in terms of hydrogen bonding.84 cis-2-Aminocyclopentanol is oxidised somewhat more quickly than the trans isomer but the difference is much smaller than for the corresponding diols,82 and the order is reversed in the aminocyclohexanols.84 Differences in hydrogen bonding presumably contribute to these effects. H PhCH-NHEt 10,- P hCH= N Et I - I- (11 Ph CH*NHEt PhCHO H (xx I I) @XI 11) 70 R. J. Jeanloz and E. Forchielli J. Biol. Chem. 1951 188 361. 8o L. Hough and M. I. Taba J .Chem. SOC. 1956,2042. 81 M. Cantley and L. Hough J . Chem. SOC. 1963 2711. ** G. E. McCasland and D. A. Smith J . Amer. Chem. SOC. 1951 73 5164. 83 G. Dahlgren and J. M. Hodson J . Phys. Chem. 1964 68,416. 86 J. Kovar J. Jary and K. Blaha Coll. Czech. Chem. Comm. 1963 28,2199. 14 Sklarz a-Diamines.-Periodate oxidation is similar to that of a-amino-alcohols with regard to rate and pH.ls Piperazine (XXII) gave besides ammonia and formalde- hyde a little formic acid,85 while benzaldehyde and benzylidene-ethylamine were formed from the diamine (XXIII).s6 The second product is probably the precursor of the aldehyde. 3 Oxidation of Enolic Compounds Activated CH-Groups and Eno1s.-The hydroxylation of malondialdehyde intermediates in the oxidation of sugars has already been mentioned.A studys8 with a series of acyclic model compounds (XXIV) showed that one of the activating groups must be an aldehyde or carboxyl function (Rf = H or OH) but that there is no oxidation with a keto-group (as in acetylacetone) or a cyano- group (as in cyanoacetic acid). The intermediate (XXV) undergoes normal cleavage giving ultimately formic acid (R1 = H) or carbon dioxide (R1 = OH) and other expected fragments depending on R2 and R3 (alkyl alkoxyl or hydro- gen). Further hydroxylation of (XXV) competes with the slower cleavage when R1 = OH malonic acid giving some oxalic acid and for alkoxymalondialde- hydes hydrolysis of the hemiacetal (XXV; R2 = OAlk) can occur (p. 12). Cyclic p-diketones are oxidised smoothly.s7 The reductone (XXVII) (p. 17) and the triketone (XXVIII) were oxidised (Scheme 6) just as readily as 1,3- cyclohexanedione (XXVI; R = H) and are the probable intermediates.Predict- ably 2-alkyldiketones gave a carbolcylic acid (XXIX) instead of carbon dioxide and 2-dialkylketones were not oxidised. One of the activating groups may be an aromatics8 or heteroarornati~~~ ring as in benzyl ketones (slow oxidation)88 and in riboflavin,89b and rapid hydroxyla- tion probably occurs also in the p-dinitrone (LXXV p. 25).86 The mechanism of hydroxylation is unknown. Enolisation cannot be the sole factor as pointed out by Bose et aZ.,90 since some weakly enolised compounds such as malonic acid are readily oxidised while certain strongly enolised compounds (e.g. acetyl- acetone) are hardly attacked. The cyclic mechanism (XXX) suggested for malon- dialdehyde and its largely enolised derivativesg0 cannot be extended readily to malonic acid or for steric reasons to cyclohexanedione.A. Wickstrom and A. Valseth Ann. pharm. franc. 1954 12 777. 88 V. M. Clark B. Sklarz and Sir A. R. Todd J. Chem. SOC. 1959 2123. 87 M. L. Wolfrom and J. M. Bobbitt J . Amer. Chem. SOC. 1956 78 2849. 8B(u) C. F. Huebner R. Lohmar R. J. Dimler S. Moore and K. P. Link J. B i d . Chern. 1945,159,503; L. J. Haynes N. A. Hughes G. W. Kenner and Sir A. R. Todd J. Chem. SOC. 1957 3727; (b) H. S. Forrest and A. R. Todd J. Chem. SOC. 1950,3295. H. Felkin Bull. SOC. chim. France 1951 915. J. L. Bose A. B. Foster and R. W. Stephens J . Chem. SOC. 1959 3314. 15 Quarter& Reviews (XXVI) 2 10; R*C02H (XXIX) 2x0,- t + *10; &R &; '0;- k0 C C H *2H (XXVI I) (XXVI I I) Scheme 6 A transient malonic-periodic acid anhydride can be envisaged in which enolisation is promoted but there is no precedent or evidence for this at present.As shown experimentally for catechol (p. 19) a labile enol-periodate complex (XXX) (XXXI) J u may be formed from the enolised cyclic p-diketone and written hypothetically as (XXXI). The contrasting resistance of the acyclic diketones is perhaps related Me Me Me (xxx I I) (XXXIII) to the anomalous properties of their e ~ o l s . ~ l However the above discussion must remain speculative without further experimental evidence. G. S. Hammond in 'Steric Effects in Organic Chemistry' ed. M. S. Newman Wiley New York 1956 p. 452. 16 Sklarz Free radicals were not detected in the oxidation of malonic acid.D2 However a study of the hydroxylation by electron spin resonance would be valuable.Methylcyclopentane-2,3-dione which exists largely as the enol -I) in an exothermic oxidation afforded acidic products one of which was probably (xxxIII) presumably via hydroxylation as above and cleavage.93 The iodo- compound (XXXIV) was isolated when limited amounts of periodate were used. Reductones.-The mductone structure -C( :O)C(OH) = C(0H)- is oxidised rapidly and the triketone subsequently cleaved as in triose r e d u c t ~ n e ~ ~ ~ ~ * and in (XXVII p. 15)J5J7 The dimethyl ether (Xxxvd) of reductic acid was not oxidised but the parent reductone (XXXVa) and the monomethyl ethers (XXXVb) and (XXXVc) gave a-oxoglutaric acid.D3 Iodine was formed under the acid conditions. This recalls the behaviour of the enolic a-alkoxymalondialdehydes already mentioned to which the monoethers (XXXVb and c) are in fact related structurally.As in the oxidation of catechol the reductone oxidation probably proceeds by removal of the second enolic proton rather than by attack of water on carbon (g). Other mechanisms are suggested by Hesse and Mixg3 Phenols.-Several research groups concerned with lignin chemistry have studied the oxidation of phenolic compounds by peri~date.~~ They confined themselves to recording the uptake of the reagent and did not isolate products from the often coloured solutions. In a recent set of elegant studies Adler and his col- leagues have described the complex products which can arise and the essentially simple reactions leading to them.96-100 Adler’s oxidations were carried out with sodium periodate in aqueous or 80 % acetic acid solution.The phenols studied were of two types the dihydric phenols with their mono-ethers and some methylphenols. The study was prompted by the observation that methanol is formed rapidly and quantitatively in the periodate oxidation of guaiacol (XXXVI) and an estimation of guaiacol residues in lignin was based on this.96 The red solution 9a M. C. R. Symons J. Chem. SOC. 1955,2794. gg G. Hesse and K. Mix Chem. Ber. 1959 92 2427. B4 J. C. P. Schwartz Chem. and Ind. 1955 1588. *= J. P. Feifer M. A. Smith and B. R. Willeford J . Org. Chem. 1959,24,90 and refs. therein. s6 E. Adler and S. Hernestam Acta Chem. Scand. 1955 9 319. g7 E. Adler and R. Magnusson Acta Chem. Scand. 1959 13 505. (a) E. Adler R. Magnusson B.Berggren and H. Thomelius Acfa CIiem. Scand. 1960 14 515; (b) E. Adler and B. Berggren Acfa. Chem. Scand. 1960 14 p. 529; (c) E. Adler R. Magnusson and B. Berggren Acta. Chem. Scand. 1960 14 p. 539. O9 E. Adler I. Falkenberg and B. Smith Acta Chem. Scand. 1962 16 529. loo (a) E. Adler L. Junghahn U. Lindberg B. Berggren and G. Westin Acta Chem. Scand. 1960 14 1261 ; (6) E. Adler J. Dahlen and G. Westin Acta. Chem Scand. 160 14 p. 1580. 17 Quarterly Reviews t MuOH (XXXVI) (XXXVI I) (XXXVI 11) from the oxidation of (XXXVI) gave o-benzoquinone (XXXWI) and cis-cis- muconic acid (XXXVIII) obviously derived by further cleavage of the q u i n ~ n e . ~ ~ The reaction is general for catechol and quinol ethers the latter yielding p - benzoquinone but resorcinol monomethyl ether is attacked only very slowly giving methoxy-p-quinone.The quinones are also formed in the very fast oxidation of catechol and quinol themselves. The relative facility of these various reactions is apparent from the major products in the following cases when two pathways are possible. Simple de- hydrogenation is preferred to deinethylation (h i) and the latter reaction is faster for para- than for ortfzo- placed groups (j).97 U OMe It Sodium bismuthate also demethylated guaiacol but lead tetra-acetate and Fremy's salt gave 2-methoxy-p-benzoquinone. para-Hydroxylation by periodate was observed in the by-products (XL) and (XLI) (coerulignone) of the oxidation of 2,6-dimethoxyphenol (XXXIX),9sa but the mechanism is probably ionic. W (XXXIX) (XLV) Besides the expected reaction a + OMe 0 rearrangement occurred (Scheme 7) in the oxidation of the phenyl ether (XLII) with the formation of a biphenyl (XLIII) formed also by the normal oxidation of 2-hydroxydibenzofuran (XLIV).Q7 Dimeric products were obtained in concentrated solution and at room tem- perature from the oxidation of (XXXIX).3,8-Dimethoxy-1 ,Znaphthaquinone 18 Sklarz 0 ' 0.6 / "R,. (XLI v) 0 Scheme 7 (XLVII)98c was formed together with small amounts of an isomer. Adler showed that the Diels-Alder dimer (XLVI) of 3-methoxy-o-quinone (XLV) was oxidised by periodate to (XLVII) in a pathway involving diketone fission and decarb- oxylati~n.~~ b 2 (XLV) - __c (XLVI) (XLVI I) The details of the mechanism by which the dihydric phenols and their ethers are oxidised are not yet understood. Kinetic studies by the stopped-flow method have revealed for catechol the second-order formation of an intermediate which breaks down by a first-order reaction but no intermediate was detected for guaiacol.lol The oxidation of quinol is of the second order.lo2 Using 180-labelled water Adler Falkenberg and Smithg9 demonstrated a difference between the oxidation mechanism of the dihydric phenols and their monoethers.The water-soluble guaiacol derivative (XLVIII) was oxidised with sodium periodate in l80-enriched water but the liberated methanol contained no label. The o- and p-benzoquinones isolated from similar oxidation of the monomethyl ethers were 50% labelled. Since the quinones did not exchange l80 with water at the experimental pH (3-4) the label must have entered during the oxidation envisaged as the attack of water on a periodate ester (XLIX) or a transient cation (L).Attack by water at the methyl carbon would have given labelled methanol . When catechol and quinol were oxidised in H,180 the quinones obtained were not labelled. Here the water molecule removes the hydroxylic proton rather than attacking at carbon (k). lol E. T. Kaiser and S. W. Weidman Tetrahedron Letters 1965 497. lo2E. T. Kaiser and S. W. Weidman J . Amer. Chem. SOC. 1964 86 4354. 19 Quarterly Revfews EtCH-S03Na OObk HO (XLVII I) (XLIX) ( L) The periodate oxidation of alkylphenols is surprisingly fast leading to various products of ortho- and para-hydroxylation which have been elucidated by Adler and his group.1oo Phenol itself is only very slowly atta~ked.~' The oxidation of 2,4-dimethyl- phenol illustrates the various reactions that have been found to occur.Dimeric products (LV and LVI) arise from the o-quinol (LI) itself and from its Diels- Me OH 2(LI) - - O w e (LI) + (LIII) - Me Me (LV) (LVI) Adler addition to the o-quinone (LIII). The latter arises from the oxidation of (LII) the alternative ortho-hydroxylation product.loOa The p-quinols e.g. (LIV) which are minor products do not dirnerise.lo0* The reaction of oestrogens with periodate reported without details may involve hydroxylation of the steroid at C(1Q)).103 The mechanism is probably analogous to that of the first hydroxylation step in the oxidation of guaiacol [cf. (XLIX or L) 1. Other hydroxylations of phenols have been reviewed.lo4 Flavanols.-The flavanols (LVII; R = H or OCH,) were oxidised by periodic acid in aqueous dimethylformamide to give a tautomeric mixture of the 2- lo3 R.A. Harkness and K. Fotherby Experentia 1961 5 253. lo* J. D. Loudon Progr. Org. Chem. 1961 5 46; W. A. Waters ibid. p. 35. 20 hydroxyflavandiones (LVLII ; analogue of (LVIII; R = H) flR 0 (LVI I) Sklarz R = H or OCH3).lo5 In methanol the methoxy- was obtained as its methyl hemiacetal (LX).lOB Q P R 0 (LVIII) O It (LIN Smith has pointed out the analogy of this oxidation with that of phenols and reductones. In those reactions and in the oxidation of simple enols one may consider the attack of water on a transient cation or enol-periodate ester (I). If R is a hydroxyl group a diketone is formed by loss of a proton; otherwise C-hydroxylation occurs. In either case further reactions may ensue as illustrated throughout this section.R HO 6 4 Other Functional Groups Alcohols Olefins and Epoxides-Methanol and ethanol are attacked only in the light.13 A selective oxidation of the 11p-OH group in steroids to the 11- ketone has been reported and at higher concentrations of periodic acid 3a- and 17#l-16p-hydroxyl groups are also attacked.lo3 W) HO (UI I I) HO (W HO Scheme 8 lo5 M. A. Smith J . Org. Chem. 1963,28 933. lo6 M. A. Smith R. A. Webb and L. J. Cline J. Org. Chem. 1965,30,995. 21 Quarterly Reviews Olefins are inert to periodate at least at room temperature and steroids again provide the only well-established exception.lo7 With a three-fold excess of periodic acid cholesterol (Scheme 8) gave 3P7Sa,6P-cholestanetriol (LXI) while with ten-fold excess the 6-ketone (LXTI) was also formed.The trio1 is not an intermediate being unaffected under the conditions (trans-diaxial vic-glycol !). The epoxide (LXIII) was presumed to be involved. With the high concentration of periodic acid acid-catalysed ring-opening by periodate ion cannot be excluded the periodate ester then collapsing to the ketone. A simple hydrolysis of epoxides catalysed by periodic acid was reported earlier by Fieser and Rajagopalan.lo8 Styrene and stilbene oxides were reported (without details) to take up periodate slowly.88 Sulphur Compounds.-The oxidation of sulphur compounds with periodate was studied by Sykes and Todd in connection with the penicillin problem.loQ Thiols are oxidised via disulphides to sulphonic acids though the second stage is effected also by iodic acid. Thus spontaneous hydrolysis of the thiazolidine (LXIV) gave a thiol oxidation of which gave penicillaminic acid (LXV).A periodate estimation of penicillin has been described.l1° Thioethers (sulphides) are oxidised to sulphoxides. Thus excellent yields of (LXVI) were obtained by oxidation of the sulphide at O" although other oxidis- ing agents failed.lll At higher temperatures sulphones are forrned.l1lJl2 Although there was no cleavage of 2-aminoethanethiol it has been observed for thioethers. The ethanol derivatives (LXVII) and (LXVIII) were both oxidised (LX I v> I CO,H / CO H ( U V ) Et S AcO (E t S) &H C H *CH ,CH *CH; NHAc (LXVIII) (LXVII) lo' R. P. Graber C. S. Snoddy H. B. Arnold and N. L. Wendler J. Org. Chcm. 1956 21 1517. lo8 L. Feiser and S. Rajagopalan J. Amer. Chem. Soc. 1949,71 3938.lob P. Sykes and A. R. Todd in 'The Chemistry of Penicillin' ed. H. T. Clarke J. R. Johnson and R. Robinson Princeton 1949 p. 927. lloL. Mazor and M. K. Papay Acta Chim. Acad. Sci. Hung. 1961 26 473; Chem. Abs. 1961 55,20330. 111 N. J. Leonard and C. R. Johnson J. Org. Chem. 1962 27 282. lla W. A. Bonner and R. W. Drisko J. Amer. Chem. SOC. 1951,73 3699. 22 Sklarz by sodium periodate the former alone yielding formaldehyde. Detailed study of the oxidation of thio-derivatives of sugars and amino-sugars showed that cleavage of the group R-S-C-C-X can occur when X is OH NH, or NHAc.I13 In all the examples adduced the carbon atom adjacent to C-X actually carries two thioether groups or one together with a second electronegative function as in (LXVIII). The importance of this is seen in the remarkable cleavage of the acetaldehyde derivative (LXIX) to give methanol and of 2-deoxyglucose diethyl- (but not dibenzyl-)dithioacetal (LXX) to give as an intermediate glycolalde- hyde.l13 The fate of the sulphur-bearing fragment was not reported and the reaction merits further study.Amines.-Formation of an N-oxide from a tertiary amino-group was observed with erythromycin (p. 13),77 and has been found for N-(2-hydroxyethyl)- and N-propyl-piperidine.ll* Triphenylphosphine readily gave the 0 ~ i d e . l ~ ~ Complexes apparently formed between potassium periodate and some primary and second- ary aliphatic arninesllj merit further study in connection with the mechanism of these reactions. A new pattern of periodate oxidation thus emerges summarised by the expression R,Z -t R,Z-0 where 2 is N P or S.Its scope and utility par- ticularly for the first two classes remains to be studied. Mono- and di-alkylanilines toluidines and more slowly halogenoanilines are o ~ i d i s e d . ~ ~ J l ~ The rate of the extensive oxidation of phenylenediamines decreases in the order meta > ortho > para,17 in contrast to the phenols (p. 18). Electron-withdrawing groups such as -CHO -COCH,,117 and -N0,78$116 retard or suppress the oxidation particularly when placed ortho or para to the amino- group- Reaction products have been studied only for aniline and vary considerably with pH. Emeraldine is formed at pH 1,118 various amino- and anilino-quino- neimines and anils at pH 4*5,119 and unidentified products at pH 9. Free radicals are involved.120 Hydrazine Derivatives.-Hydrazobenzene is rapidly oxidised by periodate to a~0benzene.l~~ Monoalkylhydrazines give nitrogen and alkanes in good yield121a via an alkyldi-imine (m) the mechanism and stereochemistry of whose break- down depends on the amount of base present.The reaction has been used in a synthesis of 3-deoxyglucose derivatives,121 and an improved preparation of ll9 L. Hough and M. I. Taha J. Chem. SOC. 1957 3994. 114 B. Sklarz and A. K. Qureshi unpublished observations. 116 K. L. Jaura K. K. Tewari and R. L. Kaushik J. Indian Chem. SOC. 1963 40 1008. 118 J. Kawashiro J. Pharm. SOC. Japan 1953 73,943 (Chem. Abs. 1954,43 10630). 117 H. Tanabe J . Pharm. SOC. Japan 1956 76 1023 (Chem. A h . 1957 51,2598). 118 H. Tanabe J. Pharm. SOC. Japan 1958 78 410 (Chem. Abs. 1958 52 14562). lL0 H.Tanabe Chem. Pharm. Bull. (Tokyo) 1958 6 645 (Chem. Abs. 1960 54 16417) and refs. therein. lao H. Tanabe Chem. Pharm. Bull. (Tokyo) 1959,7 177 316; (Chem. Abs. 1960,54,22425). 121 (a) D. J. Cram and J. S. Bradshaw J. Amer. Chem. SOC. 1963,85 1108; (b) D. M. Brown and G. H. Jones Chem. Comm. 1965 561. 23 Quarterly Reviews nicotinaldehyde by the periodate oxidation of nicotinic acid hydrazide at 0" is a related reaction.122 Strongly acid solutions of various hydrazine derivatives have been titrated with potassium ~eri0date.l~~ Hydroxylamine Derivatives.-Hydroxylamine is instantly oxidised by sodium periodate with formation of nitrous oxide and iodine.lM Analogous oxidation of phenylhydr~xylaminell~ and rnethylhydroxylamine126 gave the nitroso- compound in the latter instance as the cis-dimer without tautomerisation to the oxime.Such oxidation proceeds also in chloroform with tetraethylammonium periodate.lO However a-hydroxyamino-acids undergo instantaneous oxidative decarboxyla tion (n) .126 Primary hydroxamic acids are oxidised to nitrous oxide and the parent acid ( o ) . ~ ~ ~ The formation in fair yields of amides in the presence of primary a m i n e ~ l ~ ~ Th 10,- Th (m) Me-q-Et - Me-<-Et -+ MeCHPhEt + N NH~NH N=NH IO+' (4 RCH (NH. OH) co2 H RCHO + co2 + (HNO) - H ~ N ~ O ~ - N~O points to the existence of an acylating intermediate,124 assumed to be (LXXI).127 N-Alkylhydroxamic acids e.g. (LXXII)86 similarly give the acid or amide,127 together with a nitroso-compound,86J25 e.g. (LXXIII). N-Hydroxypyrrolidines with an a-hydrogen atom are rapidly oxidised to the nitrone (p).10v86 This reaction is also seen in the first two stages of the oxidative OH 0- 120 H.N. Wingfield W. R. Harlan and H. R. Hanmer J. Amer. Chem. SOC. 1952,74 5796. 103 B. Singh and S. S. Sahota J . Sci. Ind. Res. India 1958 17,B 386 (Chem. Abs. 1959 53 7863). lap T. F. Emery and J. B. Neilands J. Org. Chem. 1962 27 1075. 1s T. F. Emery and J. B. Neilands J. Amer. Chem. SOC. 1960 82,4903. les G. A. Snow J . Chem. SOC. 1954 2588; J. B. Neilands and P. Azari Acta Chem. Scand. 1963,17 S190. 1%' B. Sklarz and A. F. Al-Sayyab J. Chern. Soc. 1964 1318. 24 Sklarz degradation of the 2,3'-bispyrrolidinyl (LXXIV) to the nitroso-acid (LXXIII) and the keto-nitrone (LXXVQs6 Hydroxylation at the activated C(3') is a probable step in the sequence. Several Al-pyrroline 1-oxides unsubstituted at C(2) (Lxxvn) were cleaved smoothly by sodium periodate at the double bond with formation of a nitroso- and a carboxyl group (LXXX).Io There is indirect evidence for a reaction path (LXXIV) I -O (LXXV) R ' G R ' m 5 " j o o ~ eR'C).lo; R r c O z H " t 0- OH R y R $ I R T 0- OH OH Fi' R (wxx> -0 (LXXVI 1) (LXXVII I) ( U X I x) via the hydrate (LXXVIII) and the hydroxamic acid (LXXIX). In agreement with this the nitrone-acid (LXXVII; R = CH, R' = C0,H) gave lavulalde- hyde presumably via oxidative decarboxylation of the a-hydroxyamino-acid (LXXVIII; R = CH, R' = CO,H)J0 The oxidation of nl-pyrroline to pyrro- lidone may be of similar nature. Amino-acids and Peptides.-Several studies on the periodate oxidation of amino- acids have been re~orded,1~*-~~~ usually as preliminaries to the oxidation of proteins.In a recent and extensive the periodate uptake and formation of carbon dioxide at various pH's were measured. Products have only occasion- ally been identified. Cleavage of the a-amino-acid function is extremely slow (cf. a-hydroxy-acids) but is promoted by higher temperature 129 and pH 132 and by N-alky1ati0n.l~~ Thus proline is oxidised even at pH 2.2 to A l - p y r r ~ l i n e ~ ~ ~ J ~ which is oxidised further at pH 7.2 to give 2-pyrr0lidonel~~ (cf. this page). Expectedly the polyamide chain of peptides is not susceptible to periodate and oxidation occurs only at the side-chains of certain constituent amino-acids. Free serine threonine and hydro~yprolinel~~ undergo normal cleavage permitting their estimation,l= followed by extensive further oxidation.In peptides they are cleaved only when in the N-terminal position when the laa P. Desnuelle S. Antonin and A. Casal Bull. SOC. Chim. biol. 1947 29 694. lS9 K. Arakawa J. Biochem. (Japan) 1957 44 217. 130 P. D. Bragg and L. Hough J . Chem. Soc. 1958,4050. 131 H. Hormann K. Hannig and G. Fries 2. physiol. Chem. 1959 315 109. 132 J. R. Clamp and L. Hough Biochem. J. 1965,94 17. 133 L. Skursky Z . Naturforsch. 1959 14b 473. 13p See ref. 132 for collected references. 25 Quarterly Reviews amino-group is free>% Thus brief treatment of corticotropin with periodate destroyed the N-terminal serine and the borohydride-reduced product had altered biological properties.lS In yet other amino-acids oxidation of the side-chain is independent of the a-amino-group.Hydroxylysine undergoes normal fission,’37 cysteine and cystine are oxidised to the sulphonic acid cysteic acid,log and the methionine residue of peptides gives the s~1phone.l~~ H (LXXX I) Tryptophan tyrosine (a phenol) and histidine are extensively oxidised to unknown coloured products. The slow dissolution of collagen in aqueous peri~date’~~ probably involves a specific cleavage at tyrosine of the type reviewed by W i t k ~ p l ~ ~ and the cleavage of the model amide (LXXXI) to give ethyl glycinate is also of this type.139b Recently substituted indoles (LXXXII; R = H or Me) have been shown to undergo mild rapid and specific oxidations the products depending on the acidity.140 Sodium periodate cleaves the 2,3-double bond of 3-alkylindoles (LXXXII ; R’ = H) giving ortho-acyl-N-acylanilines.Periodic acid leaves this bond intact but oxidises an alkyl methylene group attached at the 2-position to a carbonyl group (Scheme 9). A recent analysis131 of the effect of periodate treatment on the amino-acid composition of procollagen confirms that oxidation occurs largely at the amino- acids of the last type and is in consonance with earlier less refined studies on ovalbumin,12* ribon~clease,’~~~ chymotrypsin,l&b and 1ys0zyrne.l~~ (LXXXI I) HOCH,*CHR*CHO R*CO,H (LXXXIII) Scheme 9 la5 S. Fujii K. Arakawa and N. Aoyagi J. Biochem. (Japan) 1957 47 471. lS6 H. B. F. Dixon Biochem. J. 1962 83 91. 13* H. Zahn and L. Zurn 2. Naturforsch. 1957 12,B 788. 139 B. Witkop Adv. Protein Chem. 1961 16 (a) 221 ; (b) 252. 140 L. J. Dolby and D. L. Booth J . Amer. Chem. SOC.1966 88 1049. 141 (a) W. F. Goebel and G. E. Perlman J. Exp. Med. 1949,89,479; (6) E. F. Janscn A. L. Curl and A. K. Balls J. Biol. Chem. 1951 189 671. D. D. van Slyke A. Hiller and D. A. MacFayden J . Biol. Chem. 1941 141 681. L. Maekawa and M. Kushibe Bull. Chem. SOC. Japan 1954 27 277. 26 Sklarz Miscellaneous Reactions.-The oxidation of the mixed aldols (LXXXIII ; R = C,H or C,H,,) proceeds readily in cold aqueous dioxan containing bicarbonate the acids RC0,H being formed.l&a The reaction is comparable with the hydroxylation of other activated CH-compounds (p. 15). Cameron and his co-workers have suggested the intervention of free radicals arising from impurities in the dioxan. They observed two hydroxylations in the aphid pig- ment series which required the presence of benzoyl peroxide and thus involve free radi~a1s.l~~ However their phenolic-quinonoid substrates are not strictly comparable with aldol (LXXXIII) and /3-diketones (XXIV XXVI) and in their conditions (periodate in refluxing aqueous dioxan) the specificity of the reagent is lost even acetone being 0xidi~ed.l~~~ Glucose polymethyl ethers are appreciably 0 x i d i ~ e d .l ~ ~ ~ Iodine is liberated slowly from primary alkyl iodides. Prior hydrolysis appears to be promoted by the periodate ion (cf. p. 22) which oxidises the liberated iodide ion.lq4 Periodic acid has been used in a potentiometric estimation of uric acid,145 and dilute sodium periodate in the detection of ferrocene derivatives on paper chroma tog ram^.^^^ 5 Periodate as a Co-oxidant A mild specific reagent for the oxidative fission of olefinic double bonds was developed by Lemieux and von Rud10ff.l~' It consists of aqueous sodium periodate and potassium permanganate at pH 7.7 in a molar ratio of about 60 1.Intermediate aldehydes are oxidised further to acids which with ketones are the oxidation products. Permanganate first oxidises the olefin to the two ketols which are cleaved by periodate. The latter also reoxidises MnV to MnV" at this pH so that the process amounts to a permanganate-catalysed oxidation by periodate. The reagent has been particularly valuable in the analysis of oils and fats butanol or pyridine being used as co-s~lvents.~~~ The fragment acids from an unsaturated lipid are esterified in situ and identified by gas-liquid chromato- g r a p h ~ . ~ ~ ~ This method should also be useful in the oxidative degradation of terpenes,150 steroids,151 etc.The action of the reagent on various functional groups has been examined to define its specificity.152 For example oxidation of isolated alcoholic groups 148 (a) A. J. Birch D. W. Cameron Y . Harada and R. W. Rickards J . Chem. Soc. 1959 889; (b) D. W. Cameron R. I. T. Cromartie Y . K. Hamied E. Haslam D. G. I. Kingston Lord Todd and J. C. Watkins ibid. 1965 6923; (c) P. Fleury and R. Boisson Compt. Rend. 1939 208 1509; (d) G. D. Greville and D. H. Northcote J . Chem. Soc. 1952 1945. 144 A. B. Foster M. Stacey and R. W. Stephens J. Chem. SOC. 1959 2681. 145 A. Berka Analyt. Chim. Acta 1961 25,434. 146 A. N. de Belder E. J. Bourne and J. B. Pridham Chem. and Ind. 1959 996. 14' R. U. Lxmieux and E. von Rudloff Canad.J . Chem. 1955,33 1701. 148 E. von Rudloff Canad. J. Chem. 1956 34 1413. lP9 T. C. L. Chang and C. C. Sweeney J . Lipid Res. 1962 3 170. 150 R. U. Lemieux and E. von Rudloff Canad. J. Chem. 1955,33 1710 1714; 1965,43,2660. 151 M. E. Wall and S. Serota J . Org. Chem. 1959 24 741. ls2 E. von Rudloff Canad. J. Chem. 1965 43 1784. 27 Quarterly Reviews is slow but activating groups accelerated the oxidation (banzyl alcohol giving benzoic acid) and may lead to further degradation as with ally1 and tetra- hydrofurfuryl alcohols. Aldehydes are oxidised smoothly to acids some /3- dicarbonyl compounds not attacked by periodate alone are oxidised and acyclic ethers also suffer a slow oxidation. By the eflective use of very small quantities of permanganate greater selectivity is achieved than is normally associated with this reagent.Disubstituted olefins are oxidised cleanly by a mixture (210:l molar) of sodium periodate and osmium tetroxide in aqueous dioxan but with more substitution oxidation is slow.153 Aldehydes can be isolated. The method is thus selective and effects a valuable economy in the expensive and toxic osmium tetroxide. A combination of ruthenium tetroxide and sodium periodate (1 1 5 molar) has been used in the mild oxidation of steroid C(3)- and C(6)- alcohols in neutral conditions.lM I thank Professor V. M. Clark for suggesting my first periodate experiment. To Professor L. N. Owen and to Dr. G. J. Buist I am indebted for valuable comments and to Professor C. A. Bunton and the Chemical Society for per- mission to reproduce diagrams (IV) (VII) and (VIII). The facilities for experi- ment and study at Cambridge and at Imperial College are gratefully acknow- ledged. 153 R. Pappo D. S. Allen R. U. Lemieux and W. S. Johnson J . Org. Chem. 1956 21 478. 150 H. Nakata Tetrahedron 1963 19 1959. 28

ISSN:0009-2681    

DOI:10.1039/QR9672100003

出版商:RSC    

年代:1967

数据来源: RSC

摘要:

The Mechanism of Rusting By U. R. Evans F.R.S. U N I V E R S I T Y OF CAMBRIDGE Introduction Scaling and Rusting.-The rusting of iron (the formation of hydrated oxide in presence of oxygen and water) must be distinguished from scaling (the formation of anhydrous oxide). In scaling the oxidation-rate falls off as the thickness of the oxide-scale increases. At high temperatures the thickening rate may be inversely proportional to the thickness. At room temperature the rate falls off more abruptly; recent measurements1 record 16 8 after a day and 35 8 after a year. The latter thickness is insufficient to produce the interference colours seen on iron after brief heating in air within the range 250-350"~. Films formed on iron by unpolluted dust-free air at room temperatures generally produce no change in appearance; they can diminish reactivity towards certain reagents.In contrast the rate of rusting sometimes remains almost constant over con- siderable periods. The reason for the difference is that in scaling the oxide is formed upon the metallic surface whilst in rusting it may happen that oxygen is reduced at one place iron passes into solution at a second place whilst the iron oxide appears in hydrated form at a third place where it cannot interfere with continued attack. Conditions for Rusting.-A horizontal iron plate fully immersed in any natural water (fresh or saline) with air above the water surface suffers rusting usually slowly. More rapid rusting occurs if the plate (placed vertically or sloping) is only partly immersed oxygen being readily replenished where the plate cuts the water-line; as explained later there is often an unattacked zone at or just below the meniscus.The rust formed on atmospheric exposure is usually more adherent than that produced by immersion. Exposure outdoors (when the surface is alternately wet and dry) soon sets up rusting but in highly polluted atmospheres sheltered surfaces sometimes rust more rapidly than those kept clean by rain. Indoors an unpainted iron surface often remains bright for some time except where salt particles can settle on it or where the air contains appreciable amounts of sulphur compounds derived from fuel. Mechanism.-The corrosion of iron (or zinc) immersed in water or salt solution is undoubtedly electrochemical; the currents have been measured by two in- ID. Gilroy and J. E. 0. Mayne Corrosion Science 1965 5 55.29 Quarterly Reviews dependent methods2p3 and correlated with the corrosion-rate in the sense of Faraday's law; the good agreement between observed and calculated rates pro- vides quantitative evidence for the electrochemical mechanism. For outdoor atmospheric rusting the evidence is less direct although Russian measurements* on plates covered with a shallow layer of water show that electrochemical action is proceeding. For indoor conditions also it is difficult to explain the facts except by assuming an electrochemical mechanism. Immersed Conditions Rusting in Potassium or Sodium Salt Solution.-The contrast between scaling and rusting is seen (Figure 1) on comparing (A) the oxidation of steel plates 225' w . x 175" X -.-e-r-Q- o-grav ime t ric x- electrometric k' - 0 2 4 6 8 1 0 1 2 1 4 Time (hr) 0 0 2 N -K2 SO /Om"- KC' O'OO 25 5 0 75 100 (hr.) Fig.I. 0.0 215 s'o ;5 I Time Chr.) exposed to dry oxygen at 175" and 22505 with (B) the corrosion of such plates partly immersed in salt solutions.2 In (A) the oxidation-rate falls off with time; in (B) after initial irregularities the corrosion velocity attains a constant value. This depends only slightly on the solution used. A twenty-fold reduction in the concentration of potassium chloride produces only a small reduction in the corrosion-rate nor is there much change when sulphate is substituted for chloride. Other measurements have shown that the corrosion-rate on pure electrolytic iron is of the same order of magnitude as the rate obtained on steel. The fact that the corrosion-rate is so little influenced by either metallic or liquid phase is due to the fact that it is largely controlled by happenings at the meniscus the only region where oxygen can readily be replenished from the air.Here oxygen is reduced by a cathodic reaction which occurs in steps but can be summarised 0 + 2H,O + 4e = 40H- U. R. Evans and T. P. Hoar Proc. Roy. SOC. 1932 A 137,343. J. N. Agar quoted by U. R. Evans J. Iron Steel Inst. 1940 l41,219P esp. p. 2 2 1 ~ . L. L. Rosenfel'd First internat. Congress met. Corrosion 1961 Report (Butterworths) 1962 p. 243; N. D. Tomashov Collected papers on 'Corrosion of Metals and Alloys' 1963 (Moscow) English Translation by A. D. Mercer ed. C. J. L. Booker. D. E. Davies U. R. Evans and J. N. Agar Proc. Roy. SOC. 1954 A 225 443. 30 Evans This requires four electrons which are supplied by iron entering the liquid as cations at points lower down the specimen by an anodic reaction which also occurs in steps but can be summarised 2Fe = 2Fe++ + 4e Since the main ions in a potassium chloride solution are K+ and C1- the cathodic and anodic products can be regarded as potassium hydroxide and ferrous chloride; where they meet they will interact to give a precipitate of a yellow- brown rust consisting of hydrated ferric oxide Fe,O,,H,O or FeO,OH it being assumed that plenty of oxygen is present; if a narrow vessel is used restricting the oxygen-supply a ferroso-ferric compound may be formed.All these facts have been established by direct observation. The alkali and ferrous salt can be detected by simple chemical tests; the precipitation and settlement of the rust are obvious to the eye.The electric current representing the upward movement of electrons through the iron from anodic to cathodic zone can be measured directly on a meter if the specimen is cut along the line dividing the expected anodic area from the expected cathodic area [Figure 2(A)]; this is however a less accurate method of measuring the total current flowing than others w. L MENISCUS Fig. 2. The ‘corrosion pattern’ varies with the physical character of the metal surface. On carefully rolled steel the anodic attack may occur only along the cut edges [Figure 2(B)] where the internal stresses left by the shearing probably prevent the maintenance of a protective film. With less carefully prepared steel corrosion starts at defects on the face and spreads downwards and sideways over arch- shaped areas.These are first seen at points low down [Figure 2(C)] and later at points higher up [Figure 2(D)]. Arch-shaped ‘mantles’ of membranous rust grow out roughly at right-angles to the surface representing the surfaces separating regions wherein OH- and Fe++ respectively predominate. In the end [Figure 2(E) and (F)] nearly the whole surface is corroding except the zone just below the water-line where any iron ions escaping from the metal would find them- selves in alkaline liquid so that precipitation must occur in contact with the metal; this area suffers no appreciable corrosion but usually develops a film thick enough to display interference colours when the specimen is taken out washed and dried. 31 Quarterly Reviews The main body of the rust is precipitated well away from the metal.Its com- position and appearance depends on the cross-section of the containing vessel which controls the oxygen supply. It is not impossible that Fe++ OH- and 0 can interact to give hydrated ferric oxide as the first solid phase but it is generally assumed that ferrous hydroxide or a basic salt is first precipitated and then becomes oxidised. Such oxidation has been studied.g Given plenty of oxygen the product can be a- or 7-FeO,OH according as the pH value is high or low. With deficiency of oxygen green ferroso-ferric compounds may appear; these have been studied at Beme,' and appear to be not simple hydroxides but to contain C1- or SO,- groups. With some geometric arrangements black magnetite is formed.Since however a solid formed out of contact with the metallic surface cannot protect the composition and structure of the rust is not of primary importance. The rate of attack will remain almost constant unless the liquid is one capable of forming a sparingly soluble substance as cathodic or anodic product. Such cases will now be considered. Rusting in Magnesium Salt Solution.-The non-protective character of rust formed in a potassium or sodium salt solution is due to its precipitation at a distance from the metal; if formed in contact with the metal surface rust provides a modicum of protection as is shown by comparing behaviour of steel in magnesium and potassium sulphates. Steel partially immersed in magnesium sulphate has been found* to corrode at slightly more than half the rate caused by distilled water (of the quality used in preparing the solution); the rate was remarkably constant over the concentration-range 0.01 - 0 .4 5 ~ . In contrast potassium sulphate corroded the steel much more quickly than the distilled water. Observations of the attack made the reason quite clear. In a potassium sulphate solution ferrous sulphate is formed below the mantles and potassium hydroxide above them. The potassium hydroxide being soluble is soon dis- persed into the liquid and except on the interference-tint zone just below the meniscus there is no protection. With magnesium sulphate there is a similar formation of arch-shaped areas first low down [Figure 2(C)] and then higher up [Figure 2(D)]. But here the magnesium hydroxide formed by cathodic action just above the lowest arch-shaped mantles does not disperse being sparingly soluble.It is deposited on the metal and when the attack becomes directed on higher points the ferrous ions formed by anodic action underneath the adherent magnesium hydroxide convert it into an iron compound. First a green ferroso- ferric layer appears and then a brown rust; these clinging deposits contain magnesium and the green is far brighter and the brown distinctly lighter than corresponding deposits containing no magnesium; probably in the green sub- stance some Mg2+ ions replace Fe2+. In the end the whole area is covered with clinging rust except for a narrow horizontal strip of white magnesium hydroxide 1 to 2 mm. broad at the meniscus. The important feature of the situation is that J. E. 0. Mayne J. Chem.SOC. 1953 129. ' W. Feitknecht and G. Keller 2. anorg. Chern. 1950 262 61. U. R. Evans J. SOC. Chem. Ind. 1928 47 S ~ T esp. p. 581. 32 Evans the solid matter is deposited upon the metallic surface instead of being precipi- tated at a distance so that there is appreciable interference with attack. Rusting in Sea-water.-Magnesium and calcium are invariably found in sea- water. At one time ~ / 1 0 sodium chloride was often used to represent sea-water in laboratory corrosion tests but generally it corrodes more rapidly. At first the difference may not be great but it becomes important after times sufficient for clinging rust to be formed; in experiments lasting 128 days sea-water collected from the English Channel caused only one third of the corrosion of ~ / 1 0 sodium ~hIoride.~ In certain circumstances however sea-water may be more corrosive than sodium chloride solution.Harbour waters containing organic sulphur com- pounds such as cystine can be very dangerous whilst the mud of estuaries may contain sulphate-reducing bacteria which render the oxygen of SO:- available for the cathodic process. Atmospheric rusting near the coast is probably favoured by magnesium chloride in the salt which prevents drying. Rusting in Hard Fresh Waters.-Many waters used for supply purposes contain calcium bicarbonate as main constituent. A pure solution of calcium bicarbonate is unstable and would deposit solid calcium carbonate if an excess of carbonic acid were not present; the amount of carbonic acid needed for stability depends on the Ca(HCO& concentration and the relationship between stabilising carbonic acid and hardness has been established by German work;l0J1 the results accord well with calculations based on the law of mass action if allow- ance is made for activity coefficients.If water carries more carbonic acid than is needed for stability the excess is known as ‘Aggressive Carbonic Acid’. Such water will dissolve solid calcium carbonate the pH rising as a result. The old ‘marble test’ useful for providing a rough forecast of the corrosive character of a water consisted in determining the pH value before and after contact with marble dust the difference between two values being known as the Langelier Index; its significance as a measure of corrosive power has perhaps been overestimated since other constituents of a water affect behaviour but any water in which marble produces an appreciable pH rise is open to suspicion and should be subjected to more reliable tests based on iron or steel specimens.If a water is free from aggressive carbonic acid (i.e. if it would fail to dissolve calcium carbonate) it will quickly deposit on iron or steel a layer of chalky rust (or rusty chalk) possessing some protective character. If for instance such a water is run through a pipe any incipient corrosion will raise the pH by the cathodic reaction and a layer of chalk later converted in part into ferrous carbonate and thereafter to clinging rust will be deposited gradually spreading C. A. J. Taylor quoted by U. R. Evans ‘Corrosion and Oxidation of Metals’ Arnold London 1960 p. 165. lo G. Bodliinder 2. phys. Chem. 1900,35,23.J. Tillmans and 0. Heubleim Gesundheits-Zngenieur 1912 35 669. 33 QuartprZy Reviews over the surface as in the experiments with magnesium sulphate; this will be equally true whether the cathodic points are particles of mill-scale remaining on the surface or cementite particles within the steel or again areas where the oxygen is preferentially renewed. The protection is not complete but in absence of complicating factors the corrosion will be much less severe than that produced by water containing aggressive carbonic acid which can build no continuous ‘chalky film’. A non-aggressive water may become aggressive if softened by base-exchange since the replacement of Ca(HCO,) by NaHCO, even if only partial lessens the requirements for stabilising carbonic acid so that some of the H2C03 present becomes ‘aggressive’.The danger could be removed by passing the water over limestone or (better) calcined dolomite but this would re-introduce some hardness. Much of the excess of carbonic acid can be removed by simply passing the water down a series of cascades or even allowing it to overflow from the top of a vertical pipe. Such treatment has been found useful at water-works for avoiding corrosion complaints not only from softened water but also from water which contains carbonic acid derived from the vegetation of the collecting area especially at certain seasons of the year. Rusting of Totally-immersed Iron under Stagnant Conditions.-Long-period studies on horizontal discs fully immersed at a known distance below the surface have been carried out at Teddington,12 the reaction being followed by measuring the disappearance of oxygen from the air-space above the liquid generally a solution of potassium chloride.Under stagnant vibration-free conditions the rate of attack is usually slow being controlled by the rate at which oxygen reaches the disc. It is greatest when the upper surface of the disc is very close to the water surface. On a disc of fixed diameter the corrosion-rate rises with increasing breadth of the containing vessel which provides a larger area of water surface; if it is assumed that only a small fraction of water molecules striking a water surface pass into the body of the liquid this would be expected. Corrosion is accelerated if the oxygen pressure in the gas space above the liquid is increased but at a certain level (about 25 atm.) protective-film formation begins to retard attack; even at lower pressures (2-5 atm.) corrosion is diminished if the specimen is exposed dry to oxygen before the liquid is intr0d~ced.l~ Under conditions where oxygen-supply is limited ferroso-ferric products appear as well as ferric rust generally black magnetite rather than the green compounds mentioned above.In the Teddington experiments in ~ / 1 0 potassium chloride the corrosion-rate was constant for twenty days and then declined owing to an ‘alteration in the physical character of the anodic film’. For the first 500-600 days the corrosion product consisted of a thin layer of magnetite over- lain by a loose mass of hydrated ferric oxide which could be shaken off without la G. D. Bengough A. R. Lee and F. Wormwell Proc.Roy. SOC. 1931 A 134,308; 1933 A 140 399; Third Report of Corrosion Committee 1935 p. 123 (Iron and Steel Inst.). l3 G. D. Bengough and F. Wormwell quoted by U. R. Evans ‘Metallic Corrosion Passivity and Protection’ Arnold London 1946 p. 297. 34 Evans OH OH t affecting the corrosion rate. Later a change to a more coherent form was accom- panied by a steady decrease in the corrosion-rate. A French study14 of iron in distilled water or soft water also revealed fairly adherent magnetite covered with brown ferric rust. Apparently the magnetite formed the cathode of a corrosion couple and the iron the anode for the cor- rosion of iron was found to be stimulated by contact with magnetite whereas contact with rust had little effect. Rust Prevention by Inhibitors.-Whilst magnesium or calcium salts merely slow rusting certain inhibitors like sodium hydroxide sodium phosphate or potas- sium chromate can prevent it altogether.Instead of loose rust an invisible film is produced; the film formed on iron placed in a solution of sodium hydroxide or phosphate has been identified by electron diffraction15J6 as y-ferric oxide;* that formed in chromate generally contains chromium.17 In sodium hydroxide the film-formation can be ascribed to OH- ions driven up to anodic points by the current or attracted by adsorption forces; the anodic reaction consists in the removal of the H from the oriented hydroxyl leaving oxide on the surface [Figure 3(A)]; the process can be repeated [Figure 3(B)] until - + METAL LIQUID 1 OH I ME METAL @ LlWlD OH OH CI CI - CI C I OH the oxide becomes sufficiently thick for protection (or perhaps until the store of loosely bonded iron atoms becomes exhausted).If the liquid contains C1- in excess of OH- [Figure 3(C)] there will be places where C1- is adsorbed instead of OH- and the anodic reaction will there be the movement of a metal ion into the * The so-called y-ferric oxide often contains hydrogen and is sometimes regarded as magnetite with the ferrous ions replaced by pairs of protons (M. C. Bloom and L. Goldenberg Corrosion Scieuce 1965 5 623); this view is not held universally. l4 E. Herzog Bull. SOC. chim. France 1936 3 1530; 1938,5 187; Corrosion et Anti-corrosion 1964 12 No. 5. l5 J. E. 0. Mayne J. W. Menter and M. J. Pryor J . Chern. SOC. 1950 3229. 1' R. M. Brasher and E. R. Stove Chem. and Ind. 1952 171. J. E.0. Mayne and J. W. Menter J. Chem. SOC. 1954,99 103. 35 Quarterly Reviews liquid not the detachment of H from OH; the two reactions involve the same electron transfer but only if C1- is present will there be corrosion and rusting. If to a chloride solution OH- is added in quantity just insufficient to prevent corrosion at the most susceptible spots (e.g. where the iron atoms are most loosely packed) it will attack iron locally but intensely; the immune area pro- vides a large cathode for oxygen-reduction and since the attack is concentrated on small anodic spots the corrosion per unit area will there be rapid and the rate of penetration serious. A small anodic area surrounded by a large cathodic area generally represents a dangerous combination. The criteria deciding between (1) film-formation and (2) passage of ferrous ions into the liquid with rust formation where they interact with alkali from the cathode have been discussed elsewhere.l* Under nearly reversible conditions the preferred reaction will be that which leads to the greatest drop of free energy.But where a high anodic current density is imposed by happenings at the cathodic area there may be an insufficient supply of ions possessing the necessary activa- tion energy to pass through the positive zone and the iron may become ‘passive’ even when the free-energy criterion would predict continuing attack. A reducible substance capable of stimulating the cathodic reaction may thus favour passivity; under certain circumstances oxidising agents inhibit corrosion although the sit- uation is complicated.Pertechnetates although less strong oxidising agents than chromates as judged by the redox potention are better inhibitors as shown by work at Oak Ridge;lg they possess however only academic interest since the supply coming from certain atomic energy plants is limited. Atmospheric Conditions The R6le of Sulphur Dioxide.-Iron exposed to humid air in absence of dust and sulphur dioxide suffers little or no rusting. The rapid formation of rust in indus- trial and urban districts has long been attributed to acidity but earlier inves- tigators regarded carbon dioxide as responsible. Work at South Kensington and Teddington20 showed that sulphur dioxide was the main cause of atmospheric corrosion at inland places and that carbon dioxide could even retard attack. The opening stage of rusting in moist air containing sulphur dioxide proceeds rapidly and is probably analogous to the opening stage of the ‘fogging’ of nickel which was found in South Kensington work21 to require sulphur dioxide (or suspended sulphate particles) and a relative humidity exceeding a critical value (about 70 x).Sulphur dioxide is adsorbed and then combines at catalytically active points with atmospheric oxygen and water giving sulphuric acid which collects more water; this attacks the metal by well-established reactions today regarded as electrochemical producing iron (or nickel) sulphate. When once ferrous sulphate has appeared on an iron or steel surface the specimen can be l8 U. R. Evans First internat. Cong. met. Corrosion 1961. Report (Butterworths) p. 1. G. H. Cartledge J. Phys.Chem. 1955,59,979; 1956,60,28 1057; 1957,61,973; Corrosion 1965 21 217. 2o W. H. J. Vernon Trans. Faraday SOC. 1924 19 886; 1927 23 159; 1935,31 1678; Chem. and Ind. 1943 p. 318; J . Roy. SOC. Arts 1949 97 589. 21 W. H. J. Vernon J. Inst. Metals 1932 48 121. 36 Evans exposed to an atmosphere containing moisture but no sulphur dioxide and rust- ing can nevertheless continue; in fact an iron specimen carrying lines of ferrous sulphate crystals develops rust along those lines when exposed to a moist atmo- sphere. One atom of sulphur can cause the transformation into rust of many atoms of iron a point brought out by detailed researches at Stuttgart.22 Any theory of the formation of rust by ferrous sulphate must embody a regeneration mechanism; it must also explain the fact that whilst the outer part of the rust is loose (it can be wiped off on filter paper) the inner part is adherent and resists vigorous scrubbing.One suggestion is that the ferrous sulphate is oxidised to the ferric state and that hydrolysis produces ferric rust and sulphuric acid which then attacks a further quantity of iron. This ‘oxidative hydrolysis’ may indeed be responsible for the loose outer portions of the rust (the detection of a trace of ferric sulphate supports the idea) but it can hardly account for the tightly adherent layer. If the acid liberated by hydrolysis were continually attacking the metallic surface to which the rust was attached adhesion would be impossible. It is more likely that there is electrochemical action with the anodic attack on the iron in pits and reduction of oxygen on the face outside (Figure 4).It has been dern~nstrated~~ (A) (e> A l R Al R RUST FC S v 4 HzO RUST METAL METAL Cathodic bu rfacc pit Fig.4. that such action can indeed produce highly adherent rust at a well-arated cathodic zone. Since as many ferrous ions are produced by the anodic attack as are used up in cathodic rust-deposition complete regeneration is provided and the mechanism explains a curious fact that although soluble salts are easily detectable in the opening stages of rusting they seem largely to disappear later; presumably the anions migrate into the pits thus hiding themselves ‘under- ground’. Although in the early stages the pits may be crevices [Fig. 4(A)] they seem to extend laterally and after a time become saucer-shaped [Fig. 4(B)]. At that stage the ferrous sulphate becomes accessible.On steel exposed outdoors for some years it has been identified independently at N~rthwich~~ and Batter~ea~~ as FeS0,,4H20. The appearance of the tetrahydrate is significant and confirms the fact that the ferrous sulphate is an anodic product; towards 22 G. Schikorr Werkstofe u Korrosion 1963 14 69; 1964 15 457. 23 U. R. Evans Nature 1965 206 980. ** R. S. Thornhill personal communication 1952. 25 A. G. Tanner Chem. and Ind. 1964 p. 1027. 37 2 Quarterly Reviews pure water the heptahydrate is the stable phase but anodic action is to develop acidity and the presence of sulphuric acid would favour the for- mation of a lower hydratemZ8 It would be expected that sooner or later tho lateral extension of the anodic area would undermine the rust layer and this indeed occurs.The periodical shredding of atmospheric rust in large flakes has been noted by several investigators. In the early stages the cathodic reaction may be the reduction of 0 to OH-. When once rust has accumulated however the cathodic reaction can be an entry of ferrous ions into the ferric rust to form a ferroso-ferric compound which is then reoxidised to the ferric state by atmospheric oxygen.* This two-stage (catalytic) mechanism may well proceed more smoothly than the direct teduction of oxygen; the matter is under investigation. Whatever the details the electro- chemical mechanism of atmospheric attack is today generally recognised. In the advanced stages of outdoor rusting formation of fresh rust is seen preferentially at points where nests of ferrous sulphate exist in the existing rust; this supports the catalytic mechanism.The effect of ferrous sulphate nests is shown by work at S t ~ t t g a r t ~ ~ and also in connection with paint-breakdown at Cambridge.3o The presence of ferrous sulphate or chloride in rust is the main reason why painting over rust gives poor protection. Owing to their strong adhesion rust traces are difficult to remove; if shut in below a paint-coat they cause local formation of fresh voluminous rust which pushes away the coat and causes it to break. Corrosion Probability.-Common observation shows that when steel is exposed to the atmosphere rusting starts locally and parts may still be bright after long periods. It is necessary therefore to consider the probability of the inception of corrosion. One method of measuring probability31 consists of placing drops of distilled water on a horizontal surface and counting the proportion which developed rust.When the atmosphere is an oxygen-nitrogen mixture the probability declines as the oxygen content is raised although the 'conditional velocity' of the corrosion produced by those drops which corrode at all in- creases. Below the drops which produce no obvious change the iron probably * Omitting combined water we can write the cathodic reaction Fea+ + 4Fe,0 + 2e = 3Fe30,; this destroys the same amount of Fee+ as is liberated by the anodic reaction Fe I Fea+ + 2e. Since the oxidation of 3Fe30 by air will yield 4.5 Fe,O instead of the original 4-0 Fe,O, it is evident that the iron destroyed in the pits is appearing as rust on the face.Two- stage experiments now being carried out by C. A. J. Taylor (with iron exposed first to moist air containing SO and then to moist air without SO,) show that the rusting continues as quickly as it does when SO is present in the second stage; if between the two stages the iron is immersed in water there is practically no rusting in the second stage. This shows that it is the soluble iron sulphate and not the rust which causes the rusting to continue in the absence of SO,. Other experiments at Cambridge have provided evidence of the electro- chemical mechanism of atmospheric corrosion. *' C. Edeleanu and U. R. Evans Trans. Faraday Soc. 1951,47 1121. 28 T. P. Hoar personal communication 1965. e9 H. Schwarz Werkstofe und Korrosion 1965 16 93 208. 80 J. E. 0. Mayne J. Appl. Chem. 1959 9 673.31 R. B. Mears and U. R. Evans Trans. Furaduy SOC. 1935 31 527. T. P. Hoar Trans. Faraday SOC. 1937 33 1152. 38 Evans develops a protective film. It has been founds* that the presence of sulphur in the steel and also its presence in air (generaIly as SO,) increases both corrosion probability and also conditional velocity This largely explains the bad behaviour of unpainted steel in modern times and also the surprisingly good behaviour of ancient iron in unpolluted atmospheres. The influence of the sulphur dioxide content of the air on the rate of atmo- spheric rusting was clearly brought out in the British corrosion tests33 by a com- parison of behaviour at exposure stations representing different degrees of pollution and equally in the Germana and Indianss tests by a comparison of behaviour at different seasons.Oriental Iron.-Iron produced in Eastern countries in early times presumably with charcoal as fuel must have contained very little sulphur; some of the old iron beams pillars or chains are in places still remote from industrialisation where the air is almost free from sulphur dioxide. The Delhi pillar probably erected in the fourth century and moved to the present site in the twelfth century provides an example. Three analyses quoted in an Indian publicationM record only a ‘trace’ of sulphur in the portion above ground and 0.008 % in the under- ground part. The upper portion of the pillar has remained rust-free and most of it is described as bronzy or bluish by different writers; evidently of the two alternative reactions oxide-formation has prevailed and during sixteen cen- tunes exposure at temperatures periodically elevated by the sun the film has reached visible thickness.The Reviewer is inclined to ascribe the immunity to low probability rather than low velocity; in the lowest part (possibly owing to salts from the soil) corrosion has occurred. However numerous other views have been expressed. One author3’ attributes the good behaviour simply to the absence. of atmospheric pollution; others to the absence of manganese36 or the presence of pho~phorus.~~ Mediaeval British Wrought Iron.-The good performance of wrought iron in our mediaval buildings is responsible for the belief that if only the traditional process of making wrought iron could be revived rusting would be avoided. It is however probable that the absence of sulphur compounds from the air in early days was responsible for the comparative absence of corrosion trouble.Exposure tests38 carried out about 1931-1938 showed that wrought iron suffered 32 R. B. Mears Carnegie Schol. Memoirs (Iron and Steel Inst.) 1935,2A 69. 33 J. C. Hudson Sixth Report of Corrosion Committee 1959 p. 9 (Iron and Steel Inst.). 34 G. Schikorr Werkstoffe und Korrosion 1964 15 457. 35 B. Sanyal and P. D. V. Bhadwar J. SOC. Ind. Res. (Kanpur) 1951,18A 69; B. Sanyal G. K. Singhania and V. K. Nigam Lardev J . Scf. Tech. India 1965 3 104. 36 National Metallurgical Laboratory Technical Journal (Jamshedpur) Feb. 1963 (Delhi Iron Pillar number). The paper by W. E. Bardgett and J. F. Stanners on p. 24 appeared also in J. Iron Steel Inst. 1963 201 3. 37 J. C.Hudson Nature 1953 172 499. 38 J. C. Hudson ‘Corrosion of Iron and Steel’ (Chapman and Hall) 1940 p. 82; J. C. Hudson and J. F. Stanners J . Iron Steel Inst. 1953 180 27. 39 Quarterly Reviews corrosion only slightly more slowly than ordinary mild steel and at a rate rather similar to steel containing 0.2% of copper. It is right to note that the wrought iron used for these tests probably contained more sulphur than the iron made in medizeval times when charcoal was used as fuel both in the puddling furnace and the forge. But Swedish iron made with charcoal corroded more rapidly than the British iron. The subject is complicated by the fact that British wrought iron after puddling and piling contains much slag and consists of a series of parallel layers some very susceptible to attack and others resistant;39 as a result the corrosion velocity parallel to the layers is much faster than in the normal direc- tion.It would however not be possible when beating out iron ornaments at the forge to arrange that the resistant layers are always parallel to the surface. There is little doubt that it is the increased sulphur content of the air and not that of the metal which explains the enhanced rusting of modern times. Low-alloy Steels.-It has been mentioned that the addition of 0.2 % of copper to mild steel reduces the corrosion-rate roughly to the level of wrought iron. Small amounts of nickel chromium aluminium and molybdenum (often in combina- tion with copper) provide even better results as shown by extensive tests in the U.S.A.40 and the U.K.38 These ‘low-alloy’ steels are not non-rusting and must be distinguished from the stainless steels containing 13 % of chromium (or in the austenitic type 18 % chromium and 8 % nickel).They do however corrode more slowly than unalloyed steel and are relatively cheap; one commercial product with 0.5 % copper 1.0 % chromium 0.16 % phosphorus and 0.8 % silicon corrodes at about one-third of the rate of ordinary mild That chromium gives protection will cause no surprise but the action of copper is less easy to explain. An American authority,4O who observes that the rust on the more resistant steels is darker more adherent and of finer texture than that on the less resistant believes that the presence of copper or nickel locks up the sulphate in insoluble form as a complex basic salt. The Role of Salt.-It is common knowledge that atmospheric rusting occurs quickly near the sea.At places where both temperature and humidity are high corrosion is particularly rapid ; probably the presence of hygroscopic magnesium chloride along with the sodium chloride here helps corrosion retarding evapora- tion of droplets that strike a steel surface.4l Careful study in West Africa42 has shown how the amount of air-borne chloride and with it the corrosion-rate decline as the distance from the sea increases; later trials conducted in the U.K. have brought out a similar relationship. Even far inland dry salt particles may be present in dust; if these are hygro- scopic they may set up corrosion where they settle on steel. The fact that dust 8s J. P. Chilton and U. R. Evans J . Iron SteeZInst. 1955,181 113; 1957,185,497; 1957,186 98.40 H. R. Copson Proc. Amer. SOC. Testing Materials 1945 45 554; 1948 48 191 ; 1952 52 1005. I1 U. R. Evans and S. C. Britton J . SOC. Chem. I d . 1930,49 173T. 4a H. R. Ambler and A. A. J. Bain J . Appl. Chem. 1955 5 437; 1960 10,213. 40 Evans can cause rusting has long been known and classical work at South Ken~ington~~ has shown that different types of particle behave very differently; silica was found to be harmless but particles of ammonium sulphate initiated rust spots. When a steel surface was exposed to dust-free air for eleven weeks it was found to become relatively immune from rusting even if later dusty air was admitted. It should not of course be assumed that it requires eleven weeks to produce a film over the surface; optical measurements at TrondheimM show that an invisible oxide-film 20 A thick can appear within 10 minutes; probably the need for the long exposure is due to the fact that at first the film keeps cracking and only after internal stresses have been used up in repeated cracking followed by repair can the invisible film contribute effectively to protection.Particles of hygroscopic salts are the most dangerous at least indoors and an analytical study of dust in factories might assist corrosion control. Early work at South Kensington and Teddingt0n~O3~~ has shown the importance of the relative humidity (R.H.) of the air. Interesting results have also been obtained at Saturated sodium bromide solution stands in equilibrium with air of 59 % R.H. at 20”c; thus a solid particle of sodium bromide will become damp if the air humidity exceeds 59%.It was found that steel inoculated with sodium bromide remained unrusted in air of 50 % R.H. and developed rust at 60 % R.H. But the passage from immunity to rusting was not always sharp. Saturated sodium chloride solution is in equilibrium with air of 78% R.H. Steel carrying sodium chloride becomes strongly rusted at 80% R.H. and remains bright at 60%; at 70% the salt particles become brown and there is some attack. Among hygro- scopic salts liable to set up rusting even in unusually dry air may be mentioned the chlorides of zinc lithium magnesium and calcium; their saturated solutions stand in equilibrium with air of lo% 15 % 32 % and 32-3 % R.H. respectively. The rusting set up at a point where a salt-particle rests on iron or steel may spread if humidity conditions are favourable.Uniform spreading gives expanding circles but frequently if through some irregularity advance commences more quickly at certain points on the periphery than others thread-like growths develop. This ‘filiform’ corrosion cannot be discussed here. The reader is referred to an excellent paper from Teddington,46 in which an electrochemical mechanism is favoured. Final Remarks General Conclusions.-The place of formation of rust is all-important. Under immersed conditions a hard water which can throw down a sparingly soluble compound by the cathodic reaction so that the rust is formed in contact with the metal will cause less rapid corrosion than a water causing loose rust. Under atmospheric conditions the closely adherent character of the rust is a less welcome feature making removal difficult; the rust usually contains soluble 43 W.H. J. Vernon Trans. Faraday SOC. 1924 19 886; 1927 23 159. 44 A. B. Winterbottom Trans. Electrochem. SOC. 1939 76 326. 46 A. Bukowiecki Schweizer Archiv. angew. Wiss. 1957 33 97. 46 R. St. J. Preston and B. Sanyal J . Appl. Chem. 1956 6 26. 41 Quarterly Reviews ferrous salts paint coats applied over rust-traces may be pushed away at the situations of salt-nests suffering breakdown. Special paints designed to con- vert ferrous sulphate to harmless compounds are described elsewhere?' Thanks are due to F. Wornwell R. S. Hudson IT. F. Stanners and J. E. 0. Mayne for useful discussions and to C. A. J. Taylor who drew the diagrams. 47 U. R. Evans and C. A. J. Taylor Trans. Insr. Metal Finishing 1962,39 188; 1965,43 169. 42

ISSN:0009-2681    

DOI:10.1039/QR9672100029

出版商:RSC    

年代:1967

数据来源: RSC

摘要:

The Chemistry and Biological Function of Isoalloxazines (Flavines) By G. R. Penzer and G. K. Radda D E P A R T M E N T OF B I O C H E M I S T R Y U N I V E R S I T Y O F O X F O R D 1 Introduction The significance of derivatives of isoalloxazine (1) in biological systems was recognised in the early 1930's. Warburg and Christian discovered an enzyme containing a yellow chromophorel (now known as the 'old yellow enzyme') and this was closely followed by the isolation of a yellow fluorescent vitamin (ribo- flavine) by Kuhn Gyorgy and Wagner-Jauregg.2 The early chemical work on riboflavine has been summarised by Wagner-Jauregg3 and some other excellent reviews are available on a number of different aspects of flavine coenzyme^.^^^ It is now common to refer to 6,7-dimethyl-9-alkylisoalloxazines as flavines and we shall follow this practice.Three biochemically important flavines are known riboflavine (2) flavine mononucleotide (3) and flavine-adenine dinu- cleotide (4). Some others have been reported.6 $%a [CH-OHI p HNYO k/ N O 713(:,qz 6\ 5 N E a N N G N H ( I) lsoalloxazine (2) Ribof lavine ?- ?- $Hi0 - t-0- R-QCH2 P FH20 - f-0- p q o - y 2 P-+C+%O F 2 O b<Nl "@g,QH0 MQ \ ;:@q; HO OH NH2 (3) Flavine mOnOnUClQOtidQ (FMN) (4) Flavine-adenine dinucleotide (FAD) 0. Warburg and W. Christian Naturwiss 1932 20 688. R. Kuhn P. Gyorgy and T. Wagner-Jauregg Ber. 1933 66 317 576. T. Wagner-Jauregg in 'The Vitamins' ed. W. H. Sebrell and R. S. Harris Academic Press H. Beinert in 'The Enzymes' ed. P. D. Boyer H. Lardy and K. Myrback Academic Press P. Hemmerich C. Veeger and H.C. S. Wood Angew. Chern. 1965,77 1 . Ref. 4 p. 397. New York 1954 vol. 111 p. 299. New York 1960 vol. 11 p. 339. 43 Quarterly Reviews The purpose of the present Review is to discuss aspects of the chemistry of flavines which are related or at least potentially related to their biological function. We hope it will become clear that this involves the application of ideas from different disciplines of chemistry. The Biological Functions of F1avines.-Riboflavine is a well-known vitamin and its action is closely associated with the function of flavine nucleotides in biochemical oxidations. FMN and more often FAD function as coenzymes in the oxidation of amino-acids amines carboxylic acids reduced nicotinamide nucleotides thiol groups etc. Flavines appear to play a part in bacterial de- hydrogenation and nitrate assimilation in fungi and plants.They are also implicated in many other systems like that of thyroid iodination.’ Flavines are involved in several photobiological processes. For instance FMN participates in photosynthetic phosphorylation in isolated chloroplast,8 and it enhances bacterial bioluminescence. Crystalline riboflavine is present in the eyes of some fish and mammals.1° It might also be a photoreceptor in phototropism of plants.ll The chemical problems raised by these natural functions are of three general types the properties related to interaction with proteins catalytic function and photobiology . 2 Non-covalent Interactions of Flavines A. Spectra.- Visible and ultraviolet.-The electronic spectrum of riboflavine in water consists of four bands centred around 220,265,375 and 447 mp (Figure 1).The spectrum of FMN is identical while that of FAD differs in several respects (see Section 2b). The precise positions of the absorption maxima and the extinc- tion coefficients depend on the environment of the flavine chromophore and a knowledge of these effects is essential for the understanding of the state of flavines in biochemical systems. Solvents affect the position of the 375 mp band generally shifting it to shorter wavelengths with decreasing solvent polarity. For instance for 3,6,7,9-tetra- methylisoalloxazine (3-methyl-lumiflavine) the absorption maxima in different solvents are 444,369 in water; 448,362 in formamide; 448,344 in NN-dimethyl- formamide and. 450 347 mp in chloroform.12 Good correlations between transition energies and an empirically derived solvent parameter (Z-value) have been observed in mixed aqueous s01vents.l~ We find that a similar correlation holds for different non-aqueous solvents.It is difficult to decide how important Ref. 4 p. 412. F. R. Whatley M. B. Allen and D. I. Arnon Biochim. Biophys. Acta 1959,32,32. W. D. McElroy and H. H. Seliger in ‘Advances in Enzymology’ ed. F. F. Nord Interscience New York 1963 vol. XXV p. 119. lo A. Pirie Nature 1960 186 352. l1 K. V. Thimann and G. M. Curry ‘Comparative Biochemistry* ed. M. Florkin and H. S. Mason Academic Press New York 1960 vol. 1 p. 281. la H. A. Harbury K. F. LaNoue P. A. Loach and R. M. Amick Proc. Nut. Acad. Sci. 1959 45 1708. l3 J. Koziol and E. Knobloch Biochirn. Biophys. Acta 1965 102 289. 44 Penzcr and Radda the hydrogen-bonding contribution is to the spectral shifts because solvent polarity and ability to form hydrogen bonds often run parallel.The position of the visible band is hardly affected by solvent but the band splits into several inflections in a non-polar environment. I " i t I I I I Wavelength (my) fluorescence ' 260 340 420 500 580 FIG. 1. Riboflavine spectra - absorption; - - - - - fluorescence; . . . . polarisation of The effect of substituents on the electronic transitions of the isoalloxazint chromophore has not been systematically studied except for the work of Dudley Ehrenberg Hemmerich and MUller.l4 There are hpwever examples in the literature and some of these are summarised in Table 1. Methyl substitution in the 6-position results in a bathochromic shift in the two long-wavelength absorption bands although it apparently only affects the 375 mp band in the 5- and 7-positions.It has been suggested that the 5-methyl group forces the N(10) = C-C = N(1)- group out-of-plane producing the bathochromic shift 0b~erved.l~ Since a 7-methyl group has a similar effect this explanation seem unlikely. In general as expected extension of conjugation (e.g. compound XIV Table 1) causes a displacement of the absorption bands to longer wave- lengths. Substituents on nitrogen atoms 9 and 3 have only a small influence on the absorption spectra. Loss of the proton from the 3-position causes a small bathochromic shift. Similar effects may be present in some flavoproteins. The protonated form of riboflavine has a spectrum nearly identical with compound (XVI) 'but different from those of (XVII) and (XVIII) (Table 1) indicating that the most basic position in N(l) in the isoalloxazine nucleus.lo K. H. Dudley A. Ehrenberg P. Hemmerich and F. Miiller Helv. Chim. Ada 1944 47 1354. V. M. Berezovskii and E. P. Rodinova J. Gen. Chem. (U.S.S.R.) 1958,28 1016. 45 P Q\ Table 1 The efect of substituents on isoalloxazine spectra Compound number x 11 111 Iv V VI VII VIII Ix X XI XI1 XIII XIV xv XVI Substituents 1 2 - = o - -0 - = o - =o - - 0 - -0 - =o - = o - =o - = s - = o - = NH - = N-CHz-CH2OH - = N-N = CHPh H =o CH =o 1 3 H H H H H H Me H H H H H H H H H 4 5 = O H = O H -0 H = O H = O Me = O Me = O H = O H = O H = O H = S H = O H = O H = O H = O H = O H 6 7 H H Me H Me Me Me Me Me H Me Me Me Me C1 H H C1 Me Me Me Me Me Me Me Me Me Me Me Me Me Me d d XVII - -OEt - -0Et H Me Me Me+ 262,402 - XVIII - -0Et Et = O H Me Me Me+ 262,394 - ,Z These spectra although mentioned in the literature have been rerecorded to be included in this Review.The compounds were given by the original authors Drs. R. M. Acheson P. Hemmerich and G. Tollin. * V. M. Berezovskii and E. P. Radinova J. Gen. Chem. U.S.S.R. 1958,uI 1016. C R. M. Cresswell A. C. Hill and H. C. S. Wood J. Chem. Suc. 1959 698. d K. H. Dudley A. Ehrenberg P. Hemmerich and F. Muller. Helv. Chim. Acta 1964 47 1354. 9 Me ribityl Me ribityl ribityl Me (CHd NEt (CH,) NEt (CH?,) NEt Me Me Me Me Me Me+ CH,+ Near U. V. (Amax m ~ ) 348 356 375 375 382 394 369 338 361 364,400 367 384 366 not recorded 264,3 90 262,390 Visible (Amax mp)Ref. 433 a 447 a 447 a 447 Q 448 b 446 b 4 4 4 a 446 a 431 a 490 a 452 a 443 c 454 a 510 d d d - - Penzer and Radda A plausible interpretation of the spectrum may be as follows.The two longer wavelength transitions are both of T-T* type (the high extinction coefficients are consistent with this) the 375 m,u band having some contribution from an intra- molecular charge-transfer process leading to a polar excited state. This would account for the solvent-dependence of its position. The inflections in the long- wavelength band observed in non-polar solvents correspond to vibrational levels within a single electronic transition.ls This may be shown by a study of the polarisation spectrum of fluorescence (Figure 1). This polarisation is constant across each of the last two absorption bandsls (0.43 and 0.28 for the 447 and 375 mp bands respectively with glycerol as solvent and linearly polarised light for excitation) indicating that in this region there are only two independent electronic transitions.The transition moments are probably at an angle of about 30" although molecular orbital cal~ulations~~ give this angle as 80". It would be interesting to measure these moments from the spectra of single crystals. Fluorescence.-Flavines show an intense fluorescence with maximal emission at 530 m,u (Figure 1). Riboflavine has a quantum yield of fluorescence of 0.25 in aqueous solutions,18 0.52 in 90% dioxan-water and 0.71 in dioxan.12 Other organic solvents (e.g. dimethyl sulphoxide) on the other hand decrease the quantum yield.19 The fluorescence of flavine nucleotides is weaker than that of riboflavine FAD being considerably weaker.2O Only the uncharged forms of flavines are fluorescent and not their anions or cations.The flavine fluorescence is generally quenched when the chromophore is bound to a protein. Studies of the fluorescence of flavines are therefore mostly concerned with the effect of quenchers and quenching mechanisms. In the case of riboflavine quenching is by a variety of substances electrolytes (e.g. KI) metal ions aromatic substances (such as phenols purines and pyrimidines) or thiols (e.g. lipoic acid and thioglycolic acid).20*21 Two types of quenching mechanism have been observed. The first brought about by electrolytes including most metal ions is due to collisions in which the electronic excitation energy is transformed into kinetic energy. The second (dis- cussed in detail in Section 2(b) involves the formation of non-fluorescent com- plexes in the ground state.These two mechanisms can be easily distinguished. It is known that if the quenching process is competitive with the emission the relation F,/F = T ~ / T holds.22 (Fo is the fluorescence efficiency and r o the life- time of the excited state in the absence of quencher and Fand r are the corre- sponding quantities in the presence of quencher.) Thus collisional quenching causes a decrease in the life-time of the excited state but quenching brought about by the formation of a non-fluorescent complex in the ground state causes 113 G. Weber 'Flavines and Flavoproteins' ed. E. C. Slater and C. Veeger E!sevier Amsterdam 1966. 17 B. Grabe personal communication. 18 G. Weber and F. W.J. Teale Trans. Faraday Soc. 1957,53 646. 19 I. M. Gascoigne and G. K. Radda Chem. Com. 1965 534. 2o G. Weber Biochem. J. 1950 47 114. z1 W. J. Rutter Acta. Chem. Scand. 1958 12 438; K. Sakai Nagoya 3. Med. Sci. 1956,18 245. *2 F. Perrin Ann. Physique 1929 12 169. 47 Quarterly Reviews no change in the life-time of the excited state i.e. T ~ / T = 1. In a situation where both processes may contribute to quenching F,/F > T ~ / T . Fa/F can be directly measured while T ~ / T can be calculated from measurements of fluorescence polarisation according to the e q ~ a t i o n ~ ~ ~ ~ 1 1 where Po refers to the value of polarisation at ‘infinite’ rigidity of the chromo- phore during the life-time of the first excited state and Pi and P are the measured polarisations in the absence and presence of quencher respectively.Quenching by collisions has two other characteristics that can be used to identify it. First it is more efficient at higher temperatures owing to the increase in the number of collisions and the decrease in the viscosity of the solvent. Secondly quenching is not linearly related to the concentration of the quencher at high concentrations for quenching efficiency increases steadily.20 Competitive quenching of flavine fluorescence by substances such as KI and thioglycolic acid is observed while with aromatic systems normally a mixed type of quenching with a major contribution from the static mechanism occurs. On the other hand even for aromatic substances quenching is entirely collisional in non-aqueous solvents (e.g. for caffeine in chloroform and naphthalene- sulphonate ions in formamide; see Table 2).Detailed studies on fluorescence quenching may provide further information. For instance from the concentration-dependence of competitive quenching the life-time of the excited state can be calculated according to the equation derived by Wawi10v.~~ For riboflavine this is approximately sec. From this life- time and measurements of polarisation of fluorescence in media of different viscosities the molecular volumes in solution can be calculated from the relation 1 1 (T = viscosity V = molecular volume). The molar volumes for riboflavine and FAD are 550 and 756 ml. respectively. These might be useful in understanding protein specificity for the different flavine nucleotides in terms of size. Infrared-In principle the infrared spectra of flavines should be useful in elucidating their mode of binding to proteins.In practice such measurements have so far been limited because of the difficulties in using aqueous solutions and low concentrations of flavines and proteins. For the unbound flavine it has been shown by infrared spectroscopy that the 4-oxygen is in the keto-form.2s 29 B. Sveshnikoff Acta. Physicochim. U.S.S.R. 1936,4 453. 24 S. I. Wawilov Z . Physik 1929 53 665. a5 J. T. Spence and E. R. Peterson J. Inorg. Nuclear. Chem. 1962 24 601. 48 Perizer and Radda Table 2 The quenching of pa vine fluorescence* Quencher (ca. 3 X 10-3~) Phthalic acid Nicotinamide DL-Phen ylalanine NAD+ Citric acid Allantoin Urea Benzoic acid Phenol-2,4-disulphonic acid Methionine Thioglycolic acid L-Ascorbic acid L-T yrosine Sulphanilic acid p-Hydroxybenzoic acid p-Aminobenzoic acid Cytosine Thiourea Phenol Pyrogallol Thymine Salicylic acid Adenosine-5’ monophosphate Deoxyadenosine-5’ monophosphate Resorcinol Cat echo1 Quinol Phloroglucinol Adenosine Deox yadenosine DL-Lipoic acid Adenine Adenine in formamide Caffeine Caffeine in CHCl p-Aminosalicylic acid DL-Tr yp t ophan Hypoxanthine Naphthalene-p-sulphonic acid Naphthalene-p-sulphonic acid in formamide Quenching (%’) x 100 Ref.0 0 0 0 <1 < 1 1 2 2 4 5 5 6 8 9 9 9 10 1 1 11 11 12 12 15 12 13 14 15 14 19 14,16 17 16 23 17 17 18 23 19 2 21 1 21,22 22 23 29 2 a a b b a a a a a a a b a a a a a e a a c f f C C C a e a e c c e a e c f f b c f f c d d a e b d d C 49 Quarterly Reviews Quenching Quencher (ca. 3 x l O V 3 ~ ) Ser o t onine creat inine sulp hat e p-Hydroxynaphthoic acid FAD compared to FMN 0.5 M-Phenobarbital in FMN 0.5 M-Phenobarbital in FAD Sodium hydrosulphite Sodium iodide Potassium iodide Silver nitrate Silver acetate 36 b 100 80 33 - 25 6 9 11 91 91 * In order to get comparable figures values are corrected to a quencher concentration of ca.3 x 10-'~ on the assumption that I,,/Ioc concentration (see ref.f). Concentrations of ribo- flavine FMN or in ref. d tetra-O-acetyl-3-N-methylriboflavine where quoted range from 6.25 x lo-'' to 2.6 x 10-6~. Temperatures ranged from 15" to 30° the temperature coefficient of fluorescence depending on the nature of the quenching involved. a K. Sakai Nagoya J. Med. Sci. 1956 18 245. b G. K. Radda Biochim. Biophys. Acta 1966 112 448. C G. Weber Biochem. J. 1950 47 114.d G. Weber in ref. 16. C K. Yagi and Y . Matsuoka Biochem. Z . 1956 328 138. f J. C. M. Tsibris D. B. McCormick and L. D. Wright Biochemistry 1965 4 504. 8 A. Giuditta and L. Casola Biochim. Biophys. Acta 1965 110 17. B. Intra- and Inter-molecular Complexes.-The interactions of flavines with a number of organic substances (mostly aromatic) have been known for many years. Many of these substances have been used to increase the solubility of riboflavine in pharmacological preparations. Their biochemical significance lies in two observations that the flavine and adenine parts of FAD are in close intramolecular association and that aromatic amino-acids are among the sub- stances capable of interacting with FMN or riboflavine. The former observation may be relevant to the coenzymic activity of FAD while the latter may partly explain the mode of interaction of flavines with proteins.Evidence for the formation of inter- and intra-molecular complexes will be discussed first. The absorption spectrum of FAD is different from those of ribo- flavine and FMN. The molar extinction coefficient of FAD at 260 mp is lower than the sum of molar extinction coefficients of riboflavine and adenosine at this wavelength.20s26 There is also a decrease in absorption at 375 and 450 mp the latter being accompanied by a broadening of the band.20 Similar spectral changes are caused by amino-acids such as tryptophan and tyrosineZ7 or by purines and pyrimidines (e.g. caffeine adenosine monophosphate).20 These spectral shifts 26 0. Warburg and W. Christian Biochcm. Z. 1948,296,294; 1948,298 150; L.G. Whitby Biochem. J. 1953 54 437. *'I. Isenberg and A. Szent-Gyorgyi Proc. Nut. Acad. Sci. 1958 44 857; M. A. Slifkin Nature 1963 197 277. 50 Penzer and Radda have often been recorded by ‘difference spectroscopy’ between free and inter- action solutions when the broadening at long wavelengths has been mistakenly identified as a new absorption band.” The fluorescence of FAD is only one fifth that of FMN. Similarly FMN fluorescence is quenched by many other aromatic compounds (Table 2). Quench- ing in these instances is probably mainly by the static mechanism. This has been clearly shown for purines.20 The optical rotatory dispersion of FAD differs significantly from that of FMN and from that of an equimolar mixture of FMN and adenosine monophosphate. In particular a pronounced Cotton effect is observed in the visible region for FAD under conditions when fluorescence indicates complex formation.l9 Further corroborative evidence for complex formation has been obtained from measurements of oxidation-reduction potentials of flavines.For instance caffeine shifts the redox potential of FMN by + 14 mv while L-tryptophan shifts it by - 15 mv.ls Although the redox potential of FAD is very nearly the same as that of FMN the lower reactivity of FAD in the reaction with reduced nicotinamide-adenine dinucleotide (NADH)28 and in photochemical reactions2’ has also been attributed to complexing. The nature of the complexes.-It is reasonable to assume that the planar aromatic systems in these complexes will best interact in a sandwich arrangement although there is no direct experimental evidence to support this.Equally uncertain is the nature of the forces responsible for the stability of the complexes. Szent-Gyorgyi and his associates assume that the FMN-indole derivative and FMN-NADH complexes are of the charge-transfer For the former the spectral shifts produced a red colour. In the latter an electron spin resonance (e.s.r.) signal attributable to the flavine semiquinone was observed. This evidence is in- sufficient since no new absorption band was observed and electron transfer lead- ing to unpairing should not be confused with charge-tran~fer.~~~~~ Similar criticisms apply to the work of Slif kin on amino-acid-riboflavine complexes2’ and of Harbury and his co-workers.12 By varying the structure of FAD analogues in the purine nucleotide McCormick and his co-workers concluded that these complexes are not of the classic charge-transfer type but ‘may be characterised as of the donor-acceptor type as broadly defined.Such a classification involves mutual polarisability or partial charge-transfer effects’.32 An additional r61e is attributed to hydrogen bonding which may become particularly important with 6-aminopurines. Molecular orbital calculations suggest the possibility of regions of complementary charge between the isoalloxazine and adenine rings.38 It is however important that in no solvent other than water have these complexes been detected in spite of a careful search.ls In aqueous organic C. H. Suelter and D. E. Metzler Biochim. Biophys. Acta 1960 44’23. ae W. R. Frisell C. W. Chung and C. G. Mackenzie J.Biol. Chem. 1959,234 1297; G. K. Radda and M. Calvin Nature 1963 200,464. I. Isenberg S. L. Baird jun. and A. Szent-Gyorgyi Proc. Nut. Acad. Sci. 1961 47 245. E. M. Kosower ref. 16. 3B J. C. M. Tsibris D. B. McCormick and L. D. Wright Biochemistry 1965,4 504. 3a B. Pullman and A. Pullman ‘Quantum Biochemistry’ Interscience New York 1963. 51 Quarterly Reviews solvents (e.g. methyl carbitol dimethyl sulphoxide) fluorescence of FAD com- pared to FMN indicates unfolding as the water concentration is decreased.lsJ4 Even in formamide which has a dielectric constant higher than that of water complexing is not significant >6,5 These observations strongly underline the unique position of water in stabilising complexes. It is likely that here we are dealing with the same phenomenon of hydrophobic interactions as is observed for proteins and polypeptide^.^^ We believe that in the flavine complexes contribu- tion of charge-transfer at best can only be very small compared to the stabilisa- tion by hydrogen bonding hydrophobic interactions and other forces such as dipole-dipole and induced dipole-dipole attractions.Molecular complexes between protonated flavines and phenols do show new absorptions that can be correlated with the donor ability of the phenol. Large single crystals of 1 1 complexes show an electrical conductivity that is increased on illumination. The photocurrent is composed essentially equally of positive and negative carriers and e.s.r. shows that complete electron-transfer can occur at lattice imperfection sites. Here we are probably dealing with a true charge- transfer phen~menon.~~ Interactions of jlavines with biopo1ymer.s.-In spite of numerous studies on model systems little is known about the nature of binding of flavine prosthetic groups to proteins except that binding is usually non-covalent.In general the two long-wavelength absorption peaks are altered on binding the 375 mp peak being shifted to shorter wavelengths and the 447 mp band to longer. In about one third of the known flavoproteins however the general shape of the absorptions remain the same while in the others shoulders around 430 and 480 mp appear in the long-wavelength peak.37 It is likely that just as in model systems the splitting of the last band is a result of the partial restoration of the vibrational structure in a non-polar environment.It is interesting that in many cases the 375 mp band is shifted to longer wavelengths in contrast to expectations from models. We feel that this apparently anomalous behaviour is due to specific interactions of the isoalloxazine ring with dipolar groups on the protein. This may mean that while overall the chromophore is surrounded by non-polar groups it still interacts at a particular point with the dipole producing an effect 34 S. F. Velick ‘Light and Life’ ed. W. D. McElroy and B. Glass Johns Hopkins Press Baltimore 1961 p. 108. 35 W. Kauzmann ‘Advances in Protein Chemistry’ ed. C. B. Anfinsen jun. M. L. Ansen K. Bailey and J. T. Edsall Academic Press New York 1959 vol. XIV p. 1. 36 D. E. Fleischman and G. Tollin Biochim. Biophys. Acra 1965 94 248; A. Ray A. V. Guzzo and G.Tollin Biochim. Biophys. Acta 1965,94,258. V. Massey and H. Ganther Biochemistry 1965,4 1161. 52 Penzer and Radda similar to that of a substituent. For instance Theorell and his co-workers3* on the basis of fluorescence studies arrived at a picture for FMN-apoprotein interaction (5). Protein -tyrosyl group-flavine interactions have been implicated in other instances39 and the interaction of riboflavine with tryptophan groups in proteins has been used to characterise the availability of this amino-acid residue to the solvent .40 3 Properties Related to the Catalytic Function of Flavines On the basis of potentiometric titrations and magnetic measurements several authors concluded about thirty years ago that the reduction of riboflavine occurs in two distinct one-electron steps through the semiquinoid intermediate (cf.Scheme l).41 The colour changes noted by early workers were spectroscopically investigated by Beinert4 who observed two bands at 570 and 900 mp during the reduction. E.s.r. showed the presence of radicals in solutions of partially reduced FMN.43 The extinction at 570 mp is related to the amplitude of the e.s.r. signal and is due to absorption by the free radical.44 The broad absorption band in the 900 mp region has been attributed on the basis of temperature- and concen- tration-dependence to a bimolecular (or higher molecular) complex of semi- quinones with each other or with other components of the solution.42 Gibson Massey and Atherton,44 on the other hand believe it to be a charge-transfer absorption of an FMN-FMNH complex. They also observed that the radical concentration was maximal at about 70% reduction of FMN (and not at 50% as might be expected from the disproportionation shown in Scheme 1).It is likely that this asymmetry is due to the presence of another radical species a complex between the semiquinone and FMNH,. Other dimers are present in these equilibria and the system is further complicated by additional acid-base dissociations. The various interactions based on acid-base titrations and spectral measurements are summarised in Scheme 1. The rates of several of the processes in Scheme 1 have been measured by fast- reaction techniques flash phot0lysis,4~ and stopped Temperature-jump relaxation47 gave values of k, = 4 x loa M - ~ sec.-l k,l = 2 x lo6 sec.-l k, = 8 x 10-1 sec.-l and k3 = 4 x lo7 M - ~ sec.-l.The interpretations of the equilibria in Scheme 1 are confirmed by e.s.r. studies. The hyperfine structure of the e m . absorption of FMNH- varies with pH in a 38 H. Theorell ‘Proceedings of the Fourth International Congress on Biochemistry’ ed. 0. Hoffmann-Ostenhof Pergamon Press London 1960 vol. VIII p. 167. 3a P. Strittmatter J. Biol. Chem. 1961 236 2329. 40 J. H. Swineheart and G. P. Hess Biochim. Biophys. Acra 1965,104,205. 41 Ref. 4 p. 372. 42 H. Beinert J. Amer. Chem. SOC. 1956 78 5323. 43 H. Beinert and R. H. Sands ‘Free Radicals in Biological Systems’ ed. M. S. Blois jun. H. W. Brown R. M. Lemmon R. 0. Lindblom and M. Weissbluth Academic Press New York 1961 ch. 2. 44 Q. H. Gibson V. Massey and N. M. Atherton Biochem. J. 1962,85,369. B Holmstrom Photochem. Photobiol. 1964 3 97.46 Q. H. Gibson and J. W. Hastings Biochem. J. 1962 83 368. 47 J. H. Swineheart J . Amer. Chem. SOC.. 1965 87 904. 53 Quarterly Reviews SCHEME 1. Redox and acid-base equilibria in fIavines Y\ 6N HCL FIG. 2. E.s.r. spectra of jlavine semiquinones as a function of pH (from ref. 48) 54 Penzer and Radda characteristic manner (Figure 2). Three forms of the free radical can be clearly distinguished the radical anion at pH higher than 8 the neutral radical between pH 6 and 2 and the protonated form.48 The best resolved spectrum [that of the anion radical of lumiflavine (6,7,9- trimethylisoalloxazine)] has been interpreted in detail with the aid of isotopic substitution. The following assignments were possible. (1) By preparing [9-methyL2HJlumiflavine it was shown that the three methyl protons in lumi- flavine participate in the main hyperfine structure interaction and are equivalent.(2) 15N-Substitution in positions 1 and 3 did not alter the main hyperfine structure and therefore the unpaired spin is virtually isolated from the pyrimidinoid ring. (3) 15N-Substitution in the 10 position alters the spectrum. This proves that N(10) has a considerable spin density on it. (4) By substituting deuterium in the benzenoid ring it was shown that H(5) but not H(8) is involved in the hyperfine interaction. Spin densities on the 6- and 7-positions are not clear as only deriva- tives with different substituents were studied and not deuterated forms. But from the ethyl- and chloro-derivatives it appears that spin density is very low on C(6) but is high on C(7).48 Free radicals are produced by partial reduction of flavoproteins by substrate or some other reducing agent.The hyperfine structures of these radicals are very poorly resolved compared with the unbound flavine radicals in solution. This is a consequence of the slow thermal rotational relaxation of the protein molecule. Free radicals in flavoproteins therefore have to be identified by correlating kinetic and spectroscopic results with e.s.r. measurements and the radical signals have to be characterised by properties such as their widths and power saturation characteristics. Oxidatiun-Reduction Potentials.-The positions at which flavines participate in biochemical electron transfer depends on their oxidation-reduction potentials. Free FMN FAD riboflavine and lumiflavine have very nearly identical potentials the experimental range varying from -0.186 v to -0.218 v at pH 7.The potential depends very largely on the nature of the substituents in the benzenoid ring and to a lesser extent on the 9-s~bstituent.4~ The observation that on complex-formation the redox potentials of flavines are slightly displaced (cf. Section 26) is of interest because they can be changed in either direction on binding to proteins. A shift of +0.067 v is observed for the redox potential of FMN on binding to the apoprotein of the old yellow enzyme indicating a firmer binding of the reduced form and in butyrylcoenzyme A dehydrogenase the shift is +0.374 v. A large negative shift of -0.153 v is observed for lipoyl dehydro- genase. This indicates that flavoproteins just like the cytochromes may occupy several positions in the electron-transport chain.Metals in Flavine Catalysis.-Many flavoproteins contain metal ions (e.g. iron molybdenum copper) which are functionally important in the catalytic processes. A. Ehrenberg ‘Electronic Aspects of Biochemistry’ ed. B. Pullman Academic Press New York 1964; A. Ehrenberg L. E. G. Eriksson and F. Miiller in ref. 16. 40 W. M. Clark ‘Oxidation-Reduction Potentials of Organic Systems’ Williams and Wilkins Co. Baltimore 1960 p. 441. 55 Quarrerly Reviews The interaction of metals with free flavines is therefore an important part of flavine chemistry and has recently been reviewed in detail.5 The similarity of flavines to 8-hydroxyquinoline prompted Albert to investigate the stability of transition-metal-flavine complexes,5o and Foye and Lange claimed to have isolated some of them.51 Hemmerich and his co-workers however have clearly shown that the oxidised forms of flavines do not have any marked affinity for metals in water because of the low basicity of the 10-nitrogen and the unfavour- able energetics of 3,4-prototropy although in acetonitrile interactions do occur.52 Silver cuprous and mercuric ions however form strong complexes in water probably stabilised through charge-transfer.Reduced flavines have no metal affinity. The flavine semiquinone on the other hand is readily chelated by metal ions having d electrons which leads to a shift towards the radical in the FMN- FMNH di~proportionation.~ E.s.r. studies of the radical paramagnetic relaxa- tion show that the semiquinone can interact with metals in flavorproteins as Reactions with reducing agents.-Flavines are reduced by sodium dithionite catalytic hydrogenation zinc and hydrochloric acid etc.The reduced flavine is rapidly autoxidised and so reduction has to be carried out in anarobic conditions. The kinetics of the reduction of riboflavine by ditbionite at pH 9 have been studied using a rapid mixing apparatus. Because of the high pH only the 900 mp absorption was observed. The kinetics of its appearance suggested that it was formed from two molecules of the semiquinone (cf. ref. 44)," but it is still doubt- ful whether the semiquinone is an intermediate in this reaction or whether it is formed by a reaction between oxidised and reduced flavine. The mechanisms of the reduction by NADH and by dihydrolipoic acid which may serve as models for flavoprotein systems have been studied.In the case of NADH the reaction is thought to proceed by a hydride-ion transfer. The evidence is derived from the effects of substituents in both flavine and nicotina- mide solvent effects and the deuterium isotope e f f e ~ t . ~ ~ ~ ~ * ~ ~ ~ ~ The alternative a fast reversible electron transfer followed by a rate-determining hydrogen abstraction has not been entirely excluded. The reaction of dihydrolipoic acid (6) is also believed to be a two-electron process on the basis of similar substituent we11.53 effects but a fast dissociation of one of the thiol groups of lipoic acid precedes the A. Albert Biochem. J. 1953 54 646. 61 W. 0. Foye and W. E. Lange J. Amer. Chem. SOC. 1954,76,2199. 52 P. Hemmerich ref. 16. 53 H.Beinert and P. Hemmerich Biochem. Biophys. Res. Comm. 1965,18,212. 54 G. P. Burn and J. R. P. O'Brien Biochim. Biophys. Acta 1959 31 328. 55 T. P. Singer and E. B. Kearney J . Biol. Chem. 1950,183,409. 56 I. M. Gascoigne and G. K. Radda Chem. Com. 1965,211. 57 G. K. Radda and M. Calvin Biochemistry 1964,3 384. 56 Penzer and Radda Reactions with oxyge:i.-The hydrogen atoms in the 7-methyl group of ribo- flavine are active. They are exchanged in D2015* condense with aromatic alde- hydes and undergo oxidative dimerisation5 under non-solvolytic basic condi- tions. The activity of this methyl group in flavines is similar to its activity in p-nitrotoluene. The kinetics of the autoxidation of reduced FMN followed by rapid reaction techniques are consistent with the occurrence of two simultaneous niechani~ms.~~ In the first the semiquinone of the flavine reacts with oxygen by a free-radical path while the competitive process involves a direct two-electron oxidation through a flavine peroxide QH Chemical evidence for the hydroperoxide intermediate in the reoxidation of reduced isoalloxazines and the analogous tetrahydropyrazines is that N(10)- substituted dihydroalloxazines are resistant to autoxidation while N(9)- substituted ones are very sen~itive.5~ Reduced flavines are not easily autoxidised when bound to proteins of the mitochondria1 electron-transport chain but they are reactive in many other flavoproteins (e.g.amino-acid oxidases). There is no satisfactory explanation yet. Hydroxylation by reducedflavines and oxygen.-Aromatic compounds foreign to an organism often undergo non-specific hydroxylation.Two chemical models for this reaction have been investigated. Hydroxyl radicals produced by Fenton’s reagent give a different pattern of aromatic substitution from the radical pro- duced by a mixture of ascorbic acid ferrous ions EDTA and oxygen. In the latter system perhydroxyl radicals may be Recently it was found that reduced FMN and oxygen together also hydroxylate aromatic compounds. There is disagreement about the nature of the hydroxylating radical in this system. The perhydroxyl radical has again been implicated but we feel that a flavine hydroperoxide [cf. (7)] is more likely.60 58 F. J. Bullock and 0. Jardetzky J. Org. Chem. 1965 30 2056. 59 W. Berends J. Posthuma J. S. Sussenbach and M. I. X. Mager ref. 16. 6o H. Staudinger B.Kerekjarto V. Ullrich and Z . Zubrzycki ‘Oxidases and Related Redox Systems’ ed. T. E. King H. S. Mason and M. Morrison John Wiley and Sons Inc. New York 1965 vol. 11; R. 0. C. Norman and G. K. Radda Proc. Chern. SOC. 1962,138. 57 Quarterly Reviews 4 Photochemistry A. Anaerobic Photochemistry.-When considering the photoreduction of flavines in the absence of oxygen a clear distinction must be made between two possible situations reaction in the presence of external electron donors (photo- reduction) and that in their absence (photobleaching). The two situations though closely related show some basic differences. An aqueous anaerobic solution of riboflavine containing an external electron donor is reduced on illumination with ultraviolet or visible light. There is a con- comitant loss of colour as the reduced form of the flavine has no visible absorp- tion in a dilute solution.On admitting oxygen to the reduced solution riboflavine is quantitatively restored.61 When riboflavine is illuminated in the absence of external donors fading is again observed but in this case colour is only partially restored when oxygen is admitted. The products include lumiflavine and lumi- chrome (6,7-dimethylalloxazine) of which only lumiflavine has a spectrum similar to that of the initial reactant.61,82 Photobleaching therefore is not even partially reversible in complete contrast to photoreduction. In addition the quantum yield ($) of photobleaching is much lower (4 = 0-006)63 than those observed for photoreductions by EDTA (4 = 0.06) and NADH ($ = 0.25).57 The spectra of both photobleached and photoreduced riboflavine solutions resemble those obtained by the dark reduction of riboflavine by dithionite.This has led to the suggestion that the mechanism of photobleaching is reduction by ~ater-splitting~"~~ providing a chemical analogue for the water-splitting reaction in photosynthesis hv FMN + H,O -+ FMNH. + OH. The evidence quoted by Nickerson and his co-workers and by Vernon in support of the mechanism is (a) hydrogen peroxide is produced during anarobic photobleaching. (b) Catalase (a specific enzyme which destroys hydrogen peroxide) inhibits the reaction. (c) Methionine sulphoxide is produced when methionine is added to the reaction mixture. ( d ) The riboflavine spectrum is nearly restored on reoxygenation. (e) The reaction does not occur in a non- aqueous solvent like ethanol.They also thought that photoreduction is faster than photobleaching because the added 'activators' (e.g. EDTA NADH Methionine) reduce the energy required to break the 0-H bonds of water by forming some kind of complex. It was suggested that some degradation of ribo- flavine occurred during reoxygenation in the absence of activators which can somehow protect the side-chain. G. Oster J. S. Bellin and B. Holmstrom Experientiu 1962 18,249. 6a W. M. Moore J. T. Spence F. A. Raymond and S. D. Colson J . Amer. Chem. SOC. 1963 85 3367. ssB. Holmstrom and G. Oster J . Amer. Chem. SOC. 1961 83 1867. 64 G. Straws and W. J. Nickerson J . Amer. Chem. SOC. 1961 83 3187. 65 L. P. Vernon Biochim. Biophys. Acra 1959 36 177. 58 Penzer and Radda This evidence is unsatisfactory.(a) The methods used to detect hydrogen peroxide are not unequi~ocal.”~ (b) The effect of catalase is uncertain; some authors have found it to accelerate photobleaching rather than to inhibit it and in any case it is difficult to visualise how decomposition of a reaction product can retard the reaction in which it is produced. ( c ) The production of methionine sulphoxide can be explained by a mechanism in which methionine acts as an electron donor; its oxidation product is subsequently hydrated to the sulphoxide. In fact more recent studies have shown that the primary product of anaerobic photoreduction by methionine is 3-(methylthio)propionaldehyde not methionine sulphoxide.66 ( d ) The near restoration of the riboflavine spectrum only shows that the isoalloxazine chromophore is not completely destroyed.(e) It has now been shown that photobleaching does occur on illumination of ethanolic ribo- flavine solutions.67 A more satisfactory mechanism has been suggested by Holmstrom Oster and The suggestion is that during photobleaching an intramolecular disproportionation occurs in which the ribityl side-chain is oxidised and the isoalloxazine nucleus is reduced. There is a great deal of evidence for side-chain participation. Smith and Metzler have identified the products of photobleaching as the reduced forms of Iumiflavine lumichrome and 6,7-dimethyl-9-formylisoalloxazine so the side- chain has certainly reacted.g7 Moore and his co-workersa2 observed a kinetic isotope effect (k,/k = 2-5) for the photobleaching of 9-(2’-hydroxyethy1)- isoalloxazine and [2’,2’-2H,]9-(2’-hydroxyethyl) isoalloxazine which shows con- clusively that hydrogen abstraction from the side-chain occurs during the rate- determining step.The lack of isotope effect in D20 compared with water suggests that hydrogen is not abstracted from water or the hydroxyl groups of the side- chain. Lumiflavine in the presence of ribitol is photobleached much more slowly than riboflavine indicating that the reaction is intram~lecular.~~ The suggested mechanism for photobleaching is shown in Scheme 2. Although hydrogen at the 2’-position is by far the most favourable for abstraction it has been shown that w-hydroxyl chains of 3-6 carbon atoms will allow some reac- tion in the order 3 > 5 > 6 > 4.68 Radical stabilisation in the side-chain can be affected by nitrogen as well as oxygen.6Q The mechanism of photoreduction is similar to that of photobleaching.It probably involves hydrogen abstraction from the electron donor not from the side-chain. The reactions of EDTA NADH and methionine have been most throughly studied though probably many amines can fulfil their r6Ie. The stoicheiometry of the EDTA reaction is unknown but it has been shown that one mole of glyoxylic acid is produced for each mole of flavine reduced.66 With NADH the major product is NAD+ (ref. 57) and with methionine me thional . 6a 66 K. Enns and W. H. Burgess J . Amer. Chcm. SOC. 1965 87 5766. 67 E. C. Smith and D. E. Metzler J. Amer. Chem. SOC. 1963 85 3285. 68 C. S. Yang and D. B. McCormick J . Amer. Chem. Soc. 1965,87 5763. 69 G. K. Radda Biochim. Biophys.Acta 1966 112,448. 59 Quarterly Reviews PKO . 7 <O 2 FMNHo kt¶ charge 4t FMN + FMNH FMN-FMNH2 k*! transfer k FMNH + FMNH 0 .A (FMNY-FMNH.) k49 SCHEME 2. Mechanism of photobleaching of riboflavine Finer mechanistic details are similar for both photoreduction and photo- bleaching. The most studied of these is the nature of the active flavine excited state. Small amounts (ca. lo-'^) of KI significantly decrease the quantum yield of photoreduction whereas the concentration required for significant fluorescence quenching is ca. 10-4~.83 Collision-theory calculations suggest that the chemically active species has a minimum life-time of 10-3-104 sec.57,63 Paramagnetic metal ions inhibit photobleaching whereas diamagnetic ones do not. It is clear that excited singlet flavine is too short-lived to be the reactive species so an excited triplet probably participates in the rate-determining step of the reaction.It is conceivable though unlikely that flavine radicals participate in the slow step because their life-time is also of the order of lo3 s~c.,~O and they are known to form complexes with rnetak5 In flash experiments it is possible to observe 7 O B. Holmstrom Arkiv Kemi 1964 22 329. Penzer and Radda the spectrum of the flavine radical and to show that it occurs as a reaction inter- mediate of both photobleaching and photored~ction.~~ An e.s.r. signal has also been detected though the time resolution of the method is too low to permit the study of the radical as an intermediate.67 Addition of KI inhibits the rate of radical formation as well as that of the whole reaction.This indicates that the triplet is involved in the rate-determining process the reaction proceeding by two one-electron steps. The rate of photoreduction has a small positive tempera- ture coefficient and depends comparatively little on donor concentration. This is consistent with the slow step's involving an excited flavine-reducing agent complex. Although under conditions of reaction the triplet is too short-lived to be detected it has been observed at low temperatures. It gives rise to phosphor- escence at 630 m ~ ' ~ and an e.s.r. signal at 1560 gauss. This corresponds to a forbidden drn = 2 transition. The half-life of the FMN triplet in N-HCl is 15 r n ~ e c . ~ ~ Measurements of the anistotropy of the spin-spin coupling tensor for several isoalloxazines show that the electron distribution in the triplet states is not very sensitive to the substituent in the 9-position but does depend on the substituents at C(6) and C(7).73 Photochemical reactivity is also much more dependent on 6- and 7- than on g-substit~ents.~~ Riboflavine can sensitise some other photochemical reactions.For example on illuminating an anzrobic mixture of riboflavine and ergosterol with visible light the steroid is dehydrogenated to give quantitatively the 6,6'-dimer biser- g~stadienol.'~ The ease of reaction decreases FMN > riboflavine > lumiflavine = FAD = 0 in contrast to the flavine reactions so far discussed. Triplet involve- ment is indicated by the action of phosphorescence quenchers and rate-independ- ence of ionic strength suggests the participation of flavine radicals.The reactivity order was attributed to more favoured singlet-triplet conversion in the molecules with large side-chains. Perhaps however the difference between this reaction and that with EDTA is that ergosterol can only react with the flavine radical which is only produced when the side-chain is oxidisable. The inactivity of FAD is common and results from very efficient quenching by the adenine moiety. Illuminated riboflavine solutions can also apparently act as reducing agents. Rutter has observed the riboflavine-catalysed photoreduction of FeIII to FeII. A better explanation than klis suggested water-splitting mechanism is that ribo- flavine is bleached in the usual way and the reduced flavines so formed achieve the reduction of FeIII. A similar mechanism may explain the photoreduction of nitrate to nitrite in the presence of FMN and MoV1 (ref.75) and a likely sequence seems to be FMN -+- reduced flavine -+ MoV + FMN + lumichrome MoV + NO,-+ MovI + NO2-. Flavines can photosensitise cis-trans isomerisation of conjugated double hv MoVI 'l C. Dhere and V. Castelli Compt. rend. 1938 206 2003. 7a T. Shiga and L. H. Piette Photochem. Photobiol. 1964 3 213. 73 J. M. Lhoste A. Haug and P. Hemmerich Biochemistry 1966 5 3290. 74 L. R. Tether and J. H. Turnbull Biochent. J. 1962 85 517. 'ti J. T. Spence and J. A. Frank J . Amer. Chem. SOC. 1963 85 116. 61 Quarterly Reviews bonds e.g. stilbenes and pimaricin an olefinic fungicide. When either cis- or trans-stilbene4carboxylic acid is illuminated in the presence of riboflavine an equilibrium mixture of 40% cis- and 60% trans-isomers is rapidly produced.Lumichrome is as effective as riboflavine so neither the side-chain nor the isoalloxazine nucleus is vital here.S9,76 Oxygen and paramagnetic ions quench the reaction so again triplet involvement seems likely. A triplet-triplet energy transfer is suggested for the mechanism isomerisation occurring in the excited olefin. This can be induced by other molecules with triplet energies greater than that of riboflavine. Greater certainty about energy transfer will be possible when the triplet energy of stilbene-4-carboxylic acid is measured. Singlet-singlet energy transfer between p-terphenyl and riboflavine on the other hand has certainly been observed.77 B. Aerobic Photochemistry.-With more intense light than is required anzerobic- ally riboflavine is photobleached in the presence of oxygen.Halwer found that in alkaline solution the main product is lumiflavine whereas below pH 7 it is l~michrome.~~ Arobic photobleaching is subject to general acid-base catalysis and so an ionic mechanism is probable with a contribution from competing radical processes such as the anaerobic one. Flavines sensitise the photo-oxidation of amino-acids. Labelling experiments show the way the acids generally break up and a suggested mechanism with its obvious correlation to the anarobic reaction is shown in Scheme 3.79 ~NH t -a2- * COz H ~NH + ;r,O + ?O* Scheme 3 Lumichrome as well as riboflavine photosensitises the destruction of biological activity of DNA by destroying the guanine bases. Several products are formed but 14C labelling shows that about 50% of the observed degradation can be expressed by Scheme 4.Scheme 4 The reaction which shows an induction period is inhibited by paramagnetic ions. The peroxides produced are not intermediates. Guanine itself is under no condi- 76 W. Berends and J. Posthuma J. Phys. Chem. 1962 66 2547. ’’ C. L. Gemmill Radiation Res. 1956 5 216. 7g M. Halwer J . Amer. Chem. SOC. 1951 73 4870. 7 9 W. R. Frisell C. W. Chung and C. G. MacKenzie J. Biol. Chem. 1959 234 1297. 62 Penzer and Radda tions photosensitive so energy transfer is unlikely. It has been suggested that triplet flavine interacts with oxygen to give ground-state flavine and singlet oxygen a reactive species known to degrade guanine. This mechanism does not readily explain why degradation is specific to guanine of all the heterocyclic bases in DNA.7s Hydroperoxides produced during the oxidation of reduced flavines may participate in this reaction (see Section 3).In addition to the reactions with nucleotides nucleic acids and amino-acids flavines can sensitise the photo-oxidation of some indole derivatives and proteins. Photoinactivation of enzymes like a-amylase urease and trypsin has also been observed although the mechanisms of these processes are not clearly under- stood.s0 The photo-oxidation of indoleacetic acid a plant hormone may be responsible for the phototropic curvature of plants as the growth hormone concentrations on the light and dark sides of the stem are different. Evidence for this is tenuous. The action spectrum of phototropic movement resembles the visible region of a flavine absorption spectrum but could equally well be ex- plained by the photosensitising action of carotenoids.ll C.Chemiluminescence Involving F1avines.-The emission of light by the bacterium Achromobacter fischeri depends on the presence of FMN. Although the spectrum of the emitted radiation corresponds to neither the fluorescence nor the phosphorescence of free flavines it is still possible that the isoalloxazine chromophore is directly involved in the light process through some kind of com~lex.~ A possible model for this is provided by the chemiluminescence of riboflavine caused by addition of hydrogen peroxide. The radiation emitted has the spectrum of riboflavine fluorescence and its intensity is enhanced by the presence of some metals notably FeII and CuII.The delayed light emission can also be enhanced by illumination of the reaction by visible light. This increase is reduced by small amounts of KI NADH benzene and nitrobenzene which suggests that the primary photochemical event is again mediated through a flavine triplet. Added aromatic molecules are hydroxylated during the reactions and this along with the enhancement by metals ions has led to the suggestion that light emission is a consequence of interaction between riboflavine and hydroxyl radicals. The fact that the energy of the riboflavine triplet (47 kcal./mole) is the same as that of the O-O-bond in hydrogen peroxide fits in with this general scheme of scnsitisation.*l Recently we found that the pattern of aromatic hydroxylation does not corre- spond to hydroxyl radicals and flavine hydroperoxide may again be involved.8e D.Inhibition and Enhancement of Photoprocesses.-The herbicides monuron (3-p-chlorophenyl-1,l -dimethylurea) and diuron [3-(2,4-dichlorophenyl)-l 1 - W. Galston Science 1950 111 619; M. G. Ferri Arch. Biochem. Biophys. 1951 31 R. H. Steele Biochemistry 1963 2 529; J. R. Williams and R. H. Steele Biochemistry 127; C. A. Ghiron and J. D. Spiles Photochem. Photobiol. 1965 4 13. 1965 4 814. O2 G. K. Radda and T. N. Young unpublished observations. 63 Quarterly Reviews dimethylurea] are among the most active inhibitors of photosynthesis. The effect of monuron on cyclic photophosphorylation with FMN as cofactor suggests that it blocks photosynthesis close to the site of the FMN electron- transport system. It is interesting therefore to find that monuron affects several flavine photoreactions.The nature of the response varies with the reaction but in general inhibition results. This is the case for the riboflavine-sensitised aerobic photo-oxidations of EDTA and triethylenetetramine as well as the anaerobic photobleaching and photoreduction by EDTA.sg The inhibition by high oxygen concentration of flavine-sensitised photo-oxidation of ascorbic acid and 2,3-dioxogulinic acid is itself reversed by r n o n u r ~ n . ~ ~ The fact that similar effects are not observed when other dyes like methylene blue and the phthalein dyes replace riboflavine has led to the attractive hypothesis that the flavine-monuron interaction is specific to this pair of substances. More generallysg it is found that compounds expected to be good electron donors towards riboflavine and which form ground-state complexes with it (see Table 2) are powerful inhibitors of photoreduction.Good electron acceptors (e.g. methyl viologen) and substances which do not form a complex with ribo- flavine have no effect. The concentrations of ground-state complexes are in- sufficient to explain the extent of inhibition. Interaction between flavine triplet and inhibitor offers an answer. This is supported by the observed temperature effects. These inhibitions are potentially important to the understanding of photo- synthesis though more work is required completely to elucidate their mech- anisms. We do not believe that the specificity of herbicides in model systems is the result of a different mechanism. The kinetic characteristics of all the inhibition reactions are very similar.Chemically there seems no reason why monuron inhibition should involve different types of interactions from that of e.g. tryptophan. The substituted ureas are comparatively weak inhibitors of even flavine reactions and their lack of effect on other dyes could be the result of molecular geometry.83 It is interesting to note that good inhibitors of anaerobic flavine photoreduc- tion including monuron are easily photo-oxidised aerobically using riboflavine as sensitiser. A different kind of interaction is observed in the presence of some micellar macromolecules (polyvinylpyrrolidone sodium decyl sulphate etc.) which enhance the photosensitivity of many dyes. It is thought that the excited &vine interacts with the macromolecule resulting in increased triplet life-th~e.~* In the final analysis the full value of chemical model systems discussed through- out this Review is only realised when they can be related to biological functions.Addendum-Since this review was written several relevant papers have appeared. It is now likely that the flavine-indole and flavine-purine complexes are of different kinds. In the former charge-transfer forces may play a r61e but they 8a P. Homznn and H. Gaffron Photorhem. Photobiol. 1964 3 499. a4 H. B. Kostenbauder P. P. Delucn and C . R. Kowarski J . Pharm. Sci. 1965 54 1243. 64 Penzer and Radda are unlikely to be significant in flavine-purine complexes.85 In contrast FMNHz acts as an electron donor in forming distinct charge-transfer complexes with nicotinamide adenine dinucleotide.86 In some work the optical rotatory dispersion curves for FAD appear to have been distorted by instrumental factors19 and only the 370 mp absorption band of the flavine is optically active,s7 although in circular dichroism a small activ- ity is associated with the 447 mp band too.88 Further evidence shows that the reaction of NADH with FMN does involve an initial two-electron reduction and not a single electron transfer.89 In a new approach it was suggested that the photoreduction of flavines by phenylacetic acid involves the formation of a benzyl flavohydroquinone.The photochemical behaviour and fluorescence properties of flavines in a solid matrix of methyl cellulose have also been examined.g1 J. E. Wilson Biochemistry 1966 5 1351. 86 T. Sakurai and H. Hosoya Biochim. Biophis. Acta 1966 112,459. R. T. Simpson and B. L. Vallee Biochem. Biophys. Res. Cornm. 1966 22 712; D. Wellner Biochemistry 1966 5 1585; I. Listowsky S. Englard J. J. Betheil and S. Seifter Biochemistry 1966 5 2548. 88 T. Takagi K. Aki T. Isemura and T. Yamano Biochern. Biophys. Res. Comm. 1966,24 501. 89 J. L. Fox and G. Tollin Biochemistry 1966 5 3865. P. Hammerich V. Massey and G. Weber Nature in the press. G. R. Penzer and G. K. Radda Nature 1967 213 251. 65

ISSN:0009-2681    

DOI:10.1039/QR9672100043

出版商:RSC    

年代:1967

数据来源: RSC

摘要:

Vibration-Vibration Energy Transfer in Gaseous Collisions By J. D. Lambert P H Y S I C A L CHEMISTRY L A B O R A T O R Y U N I V E R S I T Y OF O X F O R D 1 Introduction Most chemical reactions only occur when sufficient vibrational energy is acquired in a particular interatomic bond of the molecule concerned. The rate at which this energy is taken up or lost and the ease with which it can flow between the different molecular vibrational modes are thus of fundamental importance in chemical kinetics. In particular the various theories of unimolecular gas reac- tions are based on detailed assumptions about rates of energy transfer.l The conversion of translational into vibrational energy in molecular collisions has been extensively investigated by measurements on ultrasonic disperson and absorption.The velocity of propagation of an adiabatic sound wave in a perfect gas is given by and for a monatomic gas where molecular translational energy is freely inter- changed at every collision is independent of the sound oscillation frequency. This is no longer so for polyatomic gases where the quantised internal degrees of freedom do not freely interchange with translation. Supposing Plo is the probability per collision for a change in vibrational quantum number from 1 - 0 (and is less than unity) the average number of collisions required for a molecule to lose one quantum will be Zlo - l/Plo. If 2 is the total number of collisions one molecule suffers per second a vibrational relaxation time p may be defined by the equation Since Zl0 the collision number is a constant and Z is proportional to the gas pressure fi will be inversely proportional to pressure.For convenience relaxa- tion times p are always referred to a pressure of 1 atm. Equation (2) is an approximation as the gas kinetic collision number 2 requires modifying by a factor taking into account the Boltmann distribution of molecules between quantum states. For a simple harmonic oscillator of fundamental frequency v the correct equation is; Z, - Z/3 [I - exp(- hv/kT)] (3) H. M. Frey Ann. Reports 1960 57 28. 67 Quarterly Reviews which approximates closely to (2) for high values of v. At sound frequencies where the period of the adiabatic oscillation becomes comparable with the vibrational relaxation time the vibrational temperature of the gas lags behind the translational temperature throughout the compression-rarefaction cycle and the effective values of C and V2 in eqn.(1) become frequency-dependent. This phenomenon occurs at ultrasonic frequencies and is known as ultrasonic dispersion; it is accompanied by a non-classical absorption of sound. Measure- ment of either may be used to obtain relaxation times. The measurement of vibrational relaxation times by ultrasonic and other (e.g. shock tube) methods has been fully discussed in an earlier Quarterly Review by McCoubrey and McGrath,2 and in other subsequent articles and books.3 Values of Z, are observed varying from a few collisions to several hundred thousand and showing an exponential dependence on the frequency of the relaxing vibrational mode. In addition for heteromolecular collisions the values are highly specific for the nature of the collision partner; for example C0,-H20 collisions are some lo3 times more efficient than C0,-CO collisions.Such in- efficient transfer and such a high degree of specificity are not in accord with observations on the kinetics of unimolecular gas reactions, and it was pointed out thirty years ago by Patat and Bartholom ‘ that vibration-vibration energy transfer between molecules is likely to be at least as important from the point of view of chemical kinetics as vibration-translation transfer.* Conventional ultra- sonic methods give no direct information about vibration-vibration transfer since the energy transferred remains internal and there is no effect on the adiabatic compressibility of the gas. But indirect information has recently been obtained from experiments on ultrasonic dispersion in polyatomic gases possess- ing several active vibrational modes and in their mixtures.This work forms the main subject of the present Review together with discussion of spectroscopic and theoretical sources of information about vibration-vibration transfer. 2 Intramolecular Transfer of Vibrational Energy. Ultrasonic Dispersion in Polyatomic Gases. Intramolecular transfer of energy between different vibrational modes of a polyatomic molecule can only occur in collision. The energy in the different modes is quantised and except in rare cases where there is exact resonance between harmonics the energy discrepancy must be made up by translational energy. For a molecule with two active vibrational modes of frequency v1 and v2 there are three possible vibrational transitions which are illustrated on the energy level diagram in Figure 1 (a) Transfer of translational energy to 0 - 1 excitation of the mode vl with relaxation time &.(b) Transfer of translational energy to 0 - 1 excitation of the mode v2 with relaxation time p2. (c) The a J. C. McCoubrey and W. D. McGrath Quart. Rev. 1957 11 87. K. F. Herzfeld and T. A. Litovitz ‘Absorption and Dispersion of Ultrasonic Waves’ Academic Press New York 1959; T. L. Cottrell and J. C. McCoubrey ‘Molecular Energy Transfer in Gases’ Butterworths London 1961 ; J. D. Lambert ‘Atomic and Molecular Processes’ ed. Bates Academic Press New York 1962 ch. 20. F. Patat and E. BartholomB 2. phys. Chem. 1936 B,32 396. 68 Lambert complex transfer of quantum of vibrational energy from mode vl plus the necessary increment of translational energy to give 0 3 1 excitation of the mode v2 with relaxation time P12.FIG. 1. Energy-level diagram showing possible transitions for a molecule with two active Vibrational modes. For the majority of polyatomic molecules which have been investigated experimentally a single relaxation time is observed corresponding to relaxation of the whole of the molecular vibrational energy. This means that P2 9 PI & Pl2. Vibrational energy thus enters the molecule via process (a) which is rate-control- ling and rapidly flows in complex collisions via process (c) to the second mode (and any other higher modes). Process (b) is too slow to play any r81e. This mechanism is characterised by a single overall relaxation time 18 which can be shown to be related to 18 by the equation 18 = (C,/C,)/!? where C1 is the specific- heat contribution due to mode vl alone and C the total vibrational specific heat.The general picture is that rapid vibration-vibration transfer maintains con- tinuous equilibrium of vibrational energy between the various fundamental modes of the molecule and that the whole of this energy relaxes in a single vibration-translation transfer process via the lowest mode.5 For a few molecules in all of which there is a large difference between v2 and vl the rate of the complex process (c) is much slower and the condition P2 9 P12 > applies. Process (b) is again too slow to play any rsle but process (a) is now faster than process (c). A double relaxation phenomenon results. The vibrational energy of v2 (and any upper modes) relaxes via complex process (c) followed by the faster process (a).Process (c) is thus the rate-determining step and the vibrational energy of the upper modes is transferred with a relaxation time PI2. The vibrational energy of the lowest mode vl relaxes independently by process (a) with the shorter relaxation time PI. This behaviour has been observed for only three gases SO2 CH2C12,' and C2H6.5 The experimental data J. D. Lambert and R. Salter Proc. Roy. SOC. 1959 A 253,277. J. D. Lambert and R. Salter Proc. Roy. SOC. 1957 A 243 78; P. G. Dickens and J. W. D. Sette A. Busala and J. C. Hubbard J. Chem. Phys. 1955,23,787. Linnett ibid. 1957 A 243 84. 69 3 Quarterly Reviews are summarised in Table 1. Two relaxation times are observed p1 and Isl enabling calculation of the collision numbers Z, and ZI2 corresponding to processes (a) and (c) respectively.For all these molecules v > 215 and theore- tical considerations show that the complex step (c) is a transfer of energy between one quantum of mode v2 and two or three quanta of mode ul. (The frequency gap between v and the remaining upper modes is small in all cases and transfer between these is rapid.) Table 1 Experimental collision numbers for intramolecular vibrational energy transfer at 300°K Substance v Va i Au zia 210 (cm.-l) (cm.-l) (cm.-l) so2 519 1151 2 110 2390 390 CHZCl 283 704 2 140 460 30 C2H6 290 820 3 50 74 20 Z, is the collision number for vibration-vibration transfer between i quanta of mode Y and one quantum of mode v,. The general conclusion may be drawn that intramolecular vibration-vibration energy transfer between modes is usually faster than vibration-translation transfer from the lowest mode.In a few cases where there is a large frequency discrepancy between modes so that multiple quantum transfers are involved vibration-vibration transfer is slower. 3 Intermolecular Transfer of Vibrational Energy A. Ultrasonic Dispersion in Mixtures.-Much more information about vibra- tion-vibration transfer can be derived from ultrasonic measurements on binary mixtures. If a relaxing gas A is mixed with a non-relaxing gas B such as helium there are two collision processes by which vibration-translation energy transfer can occur (1) A* + A --t A + A (2) A* + B 3 A + B (vib. + trans.) (vib. -+ trans.). Since (1) and (2) will have different collision efficiencies the result will be a composite relaxation time for A given by 1 1-x x E=7zF4G where x is the mole fraction of B in the mixture &u is characteristic of process (l) and /IAB of process (2)? This can be represented graphically by a linear plot of reciprocal relaxation time PA against mole fraction x.If both A and B are polyatomic relaxing gases there will also be two collision processes corresponding to (1) and (2) for vibration-translation energy transfer from B* in homomolecular and heteromolecular collisions. In addition there 70 Lambert can be a vibration-vibration transfer between A* and B* making five transfer processes in all (1) A* + A -f A + A (2) A* + B - A -I- B (3) A* + B - A + B* (4) B* + B -+ B + B (5) B* + A - B + A (vib. - trans.) (vib. -+ trans.) (vib.- vib. &- trans.) (vib. +trans.) (vib. -+ trans.). If process (3) vibration-vibration transfer does not occur the mixture will show a double relaxation process characterised by two relaxation times PA and @B which will both be related to molar composition by equations of type (4) each giving a linear plot of 1/p against composition. If vibration-vibration transfer does occur the picture is completely altered. Suppose for convenience that pure A relaxes slowly and pure B rapidly so that processes (4) and (5) are both much faster than processes (1) and (2); there are now two alternative possibilities for the overall relaxation process. If vibration-vibration transfer is much faster than all the other processes (3) will maintain the vibrational energy of the whole system (A* + B*) in continuous equilibrium and the total vibrational heat content of both components will relax via the faster of processes (4) or (5).There will thus be a single relaxation time involving the total vibrational energy. The rate-controlling process will be either (4) or (5). If (4) plays the predominant rele as is likely this will give rise to a near quadvotic dependence of overall reciprocal relaxation time on mole fraction of B since the rate of process (4) is proportional to x2 and eqn. (4) no longer applies. This mechanism is closely analogous to the relaxation behaviour shown by pure polyatomic gases giving single dispersion discussed in Section 2. The near-resonant collision process involving transfer of vibrational energy from mode va of molecule A to mode Y of molecule B plays exactly the same r6le as the complex collision process involving transfer of vibrational energy from mode Y to mode v of a single molecular species.Alternatively if process (3) is slower than (4) or (9 but faster than (1) or (2) A will again relax by the route (3) followed by (4) or (3 but now (3) will be rate-determining. This will give a linear dependence of 1 / p ~ on x. B will relax independently and more rapidly via (4) and (9 with linear dependence of l / p ~ on x. There will thus be a double relaxation process with two relaxation times @A involving only the vibrational heat capacity of A and /?B only that of B; both showing linear concentration-dependence. This mechanism is analogous to the relaxation behaviour of those polyatomic gases discussed in Section 2 which show double dispersion because vibration-vibration transfer between modes is slower than vibration-translation transfer from the lowest mode.The nature of the overall relaxation process for mixtures (whether single or double) and the concentration-dependence of the relaxation times are thus determined by the relative rates of processes (1)-(5). Observations of ultrasonic dispersion in the two pure gases A and B and in a series of mixtures extending over the whole concentration range enable a diagnosis to be made of which 71 Quarterly Reviews type of mechanism is followed. The rates of processes (1) and (4) are obtained from the measurements on the pure components and the rates of processes (3) and (5) which give the best fit to the experimental observations can be estimated by trial and error.The results for a series of mixtures are given in Table 2.8 Table 2 Experimental collision numbers for intermolecular vibrational energy transfer at 300"~ A B VA VB i Av ZAB Z u ZBB (cm.-l) (cm.-l) (cm.-l) Singly dispersing mixtures c2H4 C2H6 810 SF CHClF 344 Doubly dispersing mixtures CC12F2 CH3*OCH3 260 CH3Cl CH3.0CH3 732 CHF c2F4 507 SF CH,.O*CHS 344 SFt? C,F* 344 CF4 c2F4 435 Spectroscopic data NO(A2C+) N 2341 NO(X%) CO 1876 NO(X%) N 1876 co CH4 2143 369 1 25 50 1005 122 821.5 1 11.5 40 970 74 250 1 10 5 73 (3 250 3 18 70 421 <3 164 2 16 80 1005 <3 507 1 0 50 1500 5-5 190 2 36 70 1005 5.5 220 2 5 110 2330 5-5 2330 1 11 790 - - 2143 1 267 10,000 - - 2330 1 454 500,000 - - 1534 1 609 33,000 - - ZAB is the collision number for vibration-vibration transfer between one quantum of mode VA of molecule A and i quanta of mode YB of molecule B.(Frequencies VA and VB are for 0-1 vibrational transition.) Mixtures were chosen for which near-resonant vibration frequencies in the two components would be expected to give rise to rapid vibration-vibration transfer (small value of ZAB). It will be seen that whereZABlies below thecollision numbers for vibrational relaxation of the pure components ZAA and ZBB single dispersion is observed; where ZAB lies between 2' and ZBB double dispersion is observed. The actual experimental results for one singly dispersing mixture SF + CHClF, are shown in Figure 2. The lowest vibrational modes of the two molecules lie close enough for rapid vibration-vibration transfer and single dispersion is observed with near-quadratic concentration-dependence of 1 //!I.This indicates that the homomolecular relaxation of CHClF (B) [process (4)] is the rate-controlling step and is faster than the heteromolecular relaxation [process (5)]. The curve calculated for concentration-dependence of 1 /p was obtained by setting up the detailed energy and temperature-relaxation equations developed by Tanczo~,~ and solving over the whole concentration range on an * J. D. Lambert D. G. Parks-Smith and J. L. Stretton Proc. Roy. SOC. 1964 A 282 380. F. I. Tanczos J. G e m . Phys. 1956 25 439. 72 Lambert 1200 800 - -... - -1 I.... I - t u 0 -0 0.5 S c CHUF mole fraction of CHCLF FIG. 2. Reciprocal relaxation times and energy level diagram for SF6 + CHCIF mixtures. 0 observed points; (Reproduced by permission from Proc.Roy. Soc. 1964 A 282,380) curve calculated from theory. electronic computer. The value of ZAB in Table 2 was estimated to give the best fit and may be taken as 50 -& 15. Similar behaviour is shown by the mixture C2H4 + C2H6 investigated experimentally by Valley and Legvold.lo This case is complicated by the double dispersion shown by pure ethane (Section 2) but the torsional (290 crn.-l) mode of ethane relaxes independently of the upper modes in both pure gas and mixtures. A preliminary report has recently been made of a third mixture CO + C,H,O showing single dispersion.ll The mixtures of the second section in Table 2 which were investigated earlier (when erroneous conclusions were drawn),12 all show double dispersion. The details for one mixture SF + C2F4 are shown in Figure 3.There is near- resonance between the lowest (344 cm.?) mode of SF and the first harmonic of the lowest (190 cm.-l) mode of C2F,. Perfluoroethylene shows very efficient homomolecular vibration-translation transfer (ZBB = 5 3 and the estimated vibration-vibration transfer rate (2- = 70) falls between this and the slower vibration-translation transfer rate of sulphur hexafluoride (ZAA = 1005). Double dispersion is observed with the predicted linear concentration-depend- ence of the two relaxation times. The remaining mixtures in this section all of which involve B components whose homomolecular relaxation is very rapid behave similarly. B. Spectroscopic Evidence.-Information about vibration-vibration transfer involving diatomic molecules with comparatively high vibrational frequencies lo L.M. Valley and S. Legvold J. Chem. Phys. 1962 36,481. l1 T. Seshagiri Rao and E. Srinivasachari Nature 1965 206 926. l2 J. D. Lambert A. J. Edwards D. Pemberton and J. L. Stretton Discuss. Faraday Soc. 1962 33 61. 73 Quarterly Reviews "i mole fraction of C,F 1200 800 n - ' E W 400 3 FIG. 3. Reciprocal relaxation times and energy level diagram for SFs 4- C,F4 mixtures. 0 observed points. (Reproduced by permission from Discuss. Faraday Soc. 1962 33 61) has been obtained by flash-photolytic and spectroscopic techniques. Callear showed that vibrationally excited NO in the ground electronic state NO(X2.rr) can be produced by flashing NO mixtures with ultraviolet light and the rate of relaxation in collision with N or CO followed spectro~copically.~~ He also observed the more rapid rate of vibration-vibration transfer between electronic- ally excited NO(A2C+) and N by observing the quenching of resonance fluores- cence.14 The quenching of infrared resonance fluorescence of CO in the presence of CH has been used by Millikan to measure the rate of vibration-vibration transfer for this mixture.16 The values of 2' obtained for all these mixtures are included in Table 2.In the fluorescence quenching experiments on mixtures of NO(A2C+) + N, the former was produced in vibrational levels v = 3 2 and 1 vibration transfers involving all three levels were followed.13 More than 85 % of the transfers involved exchange of a single quantum e.g. which gives a collision number designated by Zti and the values obtained for the three possible transfer were Zi! = 790; 20; = 440; 2:; = 200 which are in the approximate ratio 3 2 1.Thus vibration-vibration transfer usually involves a single quantum but the efficiency of transfer increases almost propor- tionately with increase in vibrational quantum number.'g l3 A. B. Callear Discuss. Faraday SOC. 1962 33 28. l4 A. B. Callear and I. W. M. Smith Trans. Faraday SOC. 1963 59 1735. l5 R. C. Millikan J. Chem. Phys. 1965 42 1439. l6 A. B. Callear J . Appl. Optics Supplement on Chemical Lasers 1964. 74 Lambert 4 Factors Determining the Ef€iciency of Energy Transfer The general significance of the collision numbers recorded in Tables 1 and 2 can be discussed in terms of quantum mechanical theory. For all mixtures listed in Table 2 except the last three the exchanging vibrational frequencies lie close to exact resonance 36 cm.-l being the largest energy discrepancy.An equation derived by Tanczos for resonant-energy exchange should be appli~able.~ This is based on the quantum mechanical theory of Schwartz Slawsky and Herzfeld (S.S.H.)17 and gives the probability P(a b) for resonant energy exchange in a collision between two molecules a and b as Po(a) and Po@) are geometrical orientation factors. m a ) and P(b) are 'vibra- tion factors' for the quantum jumps involved in each molecule; they depend on the detailed physical nature of the process and the intramolecular force constants and contain terms in the intermolecular repulsion parameter and the inverse of v the vibrational frequency of the mode involved. They result in the probability's being enhanced by a steep intermolecular repulsion potential and decreased for modes of high frequency.The vibration factors are also smaller for multiple quantum jumps and this lowers the probability of transfers involving harmonics of fundamental modes. 01 is an intermolecular force constant p is the reduced mass of the collision pair,+o is the minimum value of the intermolecular potential function used. Insufficient data are available on intramolecular force constants and inter- molecular potentials for most of these polyatomic molecules to make quanti- tative a priuri calculations of P(a b) possible. But the general trends produced by various factors are well illustrated in Tables 1 and 2. The first striking con- clusion is that resonant vibration-vibration exchanges do not have unit efficiency as has often been assumed.This is illustrated experimentally by the mixture CHF + C,F, where both components have vibrational modes of identical frequency 507 cm.-l and the estimated value for Z m is 50. The efficiency of near-resonant collisions decreases with rising frequency of the exchanging modes. Were all other factors equal 2' should be proportional to the square of v. This is illustrated by the increase in ZAB from 5 for interchange between CCI,F (VA = 260 cm.-l) + CH,.OCH (vg = 250 cm.-l) to 790 for the mixture NO(A2C) (VA = 2341 cm.?) + N2 (vg = 2330 cm.?) where the vibration frequencies increase by a factor of approximately 10 and Z m by a factor of rather more than 100. 2' values for the other single quantum near-resonant exchanges CHF (507 cm.-l) + C,F (507 cm.-l) SF (344 cm.-') + CHCI,F (369 cm.-l) GH (810 cm.-l) + C2H (821.5 cm-l) l7 R.N. Schwartz Z. Slawsky and K. F. Herzfeld J. Chem. Phys. 1952,20 1591. 75 Quarterly Reviews which have intermediate frequencies all lie in the neighbourhood of 50. The other factor which is important is the relative inefficiency of multiple quantum transfers. This is illustrated by the remaining near-resonant mixtures in Table 2 all of which involve 2- or 3-quantum transitions and show values of Z- ranging from 70 to 110. The intramolecular 3-quantum transfer between the 290 cm.-l and 820 cm.-l modes of ethane (Table l) with a value of Z, = 74 also falls into this class. For all the mixtures of this group there is no apparent correlation between the size of the energy discrepancy dv and 20 which justifies the approximation of treating all these energy exchanges as resonant.The remaining vibration-vibration transfers recorded in Tables 1 and 2 show energy discrepancies ranging from 110 to 609 cm.-l and can no longer be regarded as approximately resonant. They shouId follow a more complicated expression than eqn. (3 involving an exponential term in the energy discrepancy dv.17918 Callear has shown that for the transfers between NO N, and COY recorded in Table 2 where all three components are diatomic molecules with similar molecular weights and would be expected to show similar gas-kinetic collision parameters the value of log ZAB shows a linear plot against dv as predicted.13 The higher transfer efficiency of CH as a collision partner for CO in spite of the large energy discrepancy (dv = 609 cm.-l) is due to its lower molecular weight and to the fact that vibrations which involve hydrugen atoms show a large amplitude owing to the small mass of the hydrogen atom.lg This results in a substantial increase in the vibration factors (cf.eqn. 5) and is the likely explanation for the comparatively high transfer efficiencies shown by hydrogen-containing molecules in vibration-translation,5 as well as in vibration- vibration transfers. Finally the relatively inefficient intramolecular transfers given in Table 1 for SOz and CH,CI, involve both sizeable energy discrepancies and double quantum jumps. For simple polyatomic molecules where reliable molecular parameters are available a priuri quantum mechanical calculations by the S.S.H. method have shown reasonable success with homomolecular relaxation processes.18 Similar calculations have recently been carried out on a number of 0 mixtures for which some experimenial results are available for comparison.20 The calculated ZAP values are shown in Table 3; they must be regarded as only approximate but are in qualitative agreement with the experimental findings.The latter have all been made with very low concentrations of additive (< 2%) and are in- sufficient for exact interpretation. It is clear that vibration-vibration transfer can give a satisfactory explanation for the striking ‘catalytic’ effect on the vibrational relaxation of 0 shown by many polyatomic additives. The trends shown in the values of ZAB in Table 3 are of the same kind as those discussed above for Tables 1 and 2. The striking efficiency of H,O is an example of the ‘hydrogen effect’ combined with a near-resonant collision.[The even more striking value of ZBB for H,O is due to the strong dipole-dipole interaction l* J. L. Stretton Trans. Faraday SOC. 1965 61 1053. l9 D. G. Jones J. D. Lambert and J. L. Stretton J . Chem. Phys. 1965 43,4541. D. G. Jones J. D. Lambert and J. L. Stretton Proc. Phys. SOC. 1965.86. 857. 76 Lambert Table 3 Calculated collision numbers for intermolecular vibrational energy transfer in oxygen mixtures at 3 0 0 " ~ Additive (€3) vg 1534 1092 1444 729 1596 1178 1402 (cm.-l) i Av (cm.-l) 20 462 110 96 42 376 152 170 16,000 490 1800 80 1160 140 ZBB 1 1,600 4960 4050 490 ca. 1 ca. 1 ca. 1 ZAB is the collision number for vibration-vibration transfer between one quantum of the fundamental mode of O e ( v ~ = 1554 an.-') and i quanta of mode YB of molecule B.For pure oxygen ZAA = 8.31 x lo7 (calc.). between molecules which leads to a very deep minimum in the intermolecular potential so that the+o term (cf. eqn. 5 ) takes control. Theoretical calculations giving such high transfer efficiencies are unreliable.] 5 The Rsle of Vibration-Vibration Transfers in Chemical Kinetics Two factors of fundamental importance in unimolecular reaction kinetics are the efficiency of collisional activation and deactivation of molecules and the efficiency of energy transfer between different vibrational modes in the molecule. The various conclusions about efficiency of vibration-vibration transfer which have been discussed above apply equally to homomolecular and heteromolecular collisions.The vibrational deactivation of molecules in homomolecular collisions can occur either by vibration-translation transfer or by resonant vibration- vibration transfer. The efficiencies of both processes depend critically on the frequency of the lowest vibrational mode; the collision number for vibration- translation transfer varies exponentially with frequency; the collision number for resonant vibration-vibration transfer is roughly proportional to the square of the frequency. Comparison of the values for vibration-vibration transfer given above with those for vibration-translation transfer given by Lambert and Salter5 shows that for hydrogen-containing molecules with frequencies below about 500 cm.-l both processes have collision efficiencies ( ~/ZAB) varying between 1/5 and 1/50 and there is little significant difference between them.For molecules containing no hydrogen atom vibration-translation transfer will be substantially less efficient even in this frequency range. As frequency rises above 500 cm.? both efficiencies decrease but vibration-vibration transfer quickly becomes much more efficient than vibration-translation. Thus for 0 (v = 1554 cm.-l) resonant vibration-vibration transfer (ZAB = ca. lo2) is several powers of ten more efficient than vibration-translation (Zlo = ca. lo7) and even a non- resonant collision with an energy discrepancy dv = 462 cm.-l (ZAB = ca. 104) is much more efficient. The temperature-dependence of resonant vibration- vibration transfer is much weaker than that of vibration-translation transfer 77 Quarterly Reviews so that the difference in efficiency will become less marked at higher temperatures.The chemical implication of this is that for moderately complex organic molecules which have torsional frequencies around 250 cm.-l homomolecular vibrational deactivation may be expected to occur by either process and to have a collision efficiency of 1/3 to 1/5 (allowing for the slightly higher efficiency of transfer from upper vibrational levels). Heteromolecular deactivation by an additive with a suitable range of low vibrational frequencies for easy vibration- vibration transfer may well be somewhat more efficient than homomolecular deactivation. This is in accord with recent observations on the kinetics of thermal isomerisation of cyclobutenes.21 Recent ultrasonic work has shown that for the straight-chain hydrocarbons n-pentane and n-hexane which have even lower torsional frequencies (v < 100 cm.-l) vibration-translation transfer occurs at approximately every collision,22 and other flexible organic molecules may be expected to show the same high efficiency.In contrast for all molecules contain- ing no hydrogen atom and for hydrogen-containing molecules with lowest vibrational frequencies above 500 cm.-l deactivation will occur mainly by vibration-vibration transfer with collision efficiencies varying between roughly 1/50 and 1/500. For this group collisions with monatomic additives which can only deactivate by vibration-translation transfer will be much less efficient than self-collisions. Collisions with polyatomic additives can have slightly higher efficiency than self-collisions but only where suitable vibrational frequencies for near-resonant transfer are present ; complex organic molecules possessing a wide spectrum of frequencies are most likely to fulfil this criterion.Otherwise additive efficiencies will be lower. Relative collision efficiencies for deactivation derived from unimolecular reaction kinetics are in accord with these views.2 The different active vibrational modes of polyatomic molecules are usually fairly closely spaced so that vibration-vibration transfer between modes will be near-resonant and will show an efficiency of thesame order as that of vibrational activation and deactivation. Transfer between modes which are widely spaced may be considerably less efficient. 21 H. M. Frey and D. C. Marshall Trans. Faraday SOC. 1965,61,1715; C . S. Elliott and H. M Frey ibid. 1966 62 895. *a R. Holmes G. R. Jones and R. Lawrence Trans. Faraday SOC. 1966 62,46. 78

ISSN:0009-2681    

DOI:10.1039/QR9672100067

出版商:RSC    

年代:1967

数据来源: RSC

摘要:

The Hydrated Electron By David C. Walker CHEMISTRY DEPARTMENT UNIVERSITY OF BRITISH COLUMBIA VANCOUVER CANADA Hydrated electrons are not easily distinguishable chemically from hydrogen atoms. Certainly in radiation chemistry and doubtless in other chemical studies (such as the reaction of sodium with water) they have been mistaken for hydrogen atoms. In fact hydrated electrons and hydrogen atoms constitute a conjugate acid-base pair the former being the basic and more strongly reducing species. The recent discovery of the hydrated electron by radiation chemists is thus important not only in radiation chemistry but also to investigations on electron- transfer processes novel mechanistic and synthetic reductions redox phenomena electrochemistry and photochemistry. Because of its intriguing simplicity it is of some interest to theoreticians and spectroscopists.Hydrated electrons are extra electrons solvated in liquid water. The process of solvation may be envisaged as the polarisation of the solvent molecules owing to electron-dipole interactions the electric field of the extra electron inducing a potential well which then traps and stabilises the electron. It is thus bound in a self-induced stable quantum state of the polarised dielectric. It may be simply described as a polaron and is variously symbolised as (H,O)*- or e-8q but the latter will be employed here. The solvation energy diffusion coefficient and interaction distance indicate a charge distribution even larger than that of an iodide ion probably extending 25-34 A from the centre Figure 1 m’ively 0 Fm. 1.Diagrammatic representation of a solvated electron. 79 Quartedy Reviews depicts a solvated electron; the size of the arrows represents the extent of the electron-dipolar interaction and the degree of orientation of the water dipoles both diminishing with increasing distance from the centre but extending well beyond the first solvation shell. eFSq is intensely coloured and mobile. It is very reactive and hence short- lived some of its reaction rates being among the fastest known. Its importance is manifest by the fact that since its discovery a few years ago the rates of reaction of e-aq with more than 300 molecules ions and radicals have been measured. This Review is concerned with the existence detection identification stability reactivity and structure of e-aq. Almost all this information is derived from radiation chemistry.l Electrons solvated in media other than water are referred to briefly in the Introduction and for purposes of comparison in the section on theories of electron binding.1 Introduction A. Electrons Solvated in Various Media.-Electrons solvated or trapped in various media are characterised by blue colours paramagnetism and electrical conductivity. The optical absorption bands in the red or near infrared regions exhibited by these systems are invariably broad intense and without fine struc- ture. The paramagnetic susceptibility and electron spin resonance absorption (e.s.r.) indicate free electrons having only weak interactions with neighbouring solvent nuclei. Equivalent conductivities of some of these systems are very high being much larger than for ordinary ions in solution.These properties characterise extra electrons stabilised in media as diverse as liquid ammonia crystalline salts low-temperature glassy ethers solutions of metals in fused alkali halides electrolysed tetralkylammonium salts and liquids and solids of great variety subjected to high-energy radiations. On some the information is prolific on others scanty. More than a century ago Wey17 first showed that alkali metals dissolve in liquid ammonia to give blue solutions apparently perfectly stable when pure. These and solutions of metals in amines and ethers are widely used in organic chemical syntheses as reducing agents.8 Indeed it is likely that any liquid having comparatively high dielectric constant and donor properties will dissolve alkali metals.Even an aprotic solvent hexamethylphosphoramide has recently been shown to form a blue reducing For reviews on the radiolysis of water before the discovery of the hydrated electron see (a) A. 0. Allen ‘The Radiation Chemistry of Water and Aqueous Solutions’ Van Nostrand New York 1961; (6)-E. J. Hart and R. L. Platzman in ‘Mechanisms in Radiobiology’ Academic Press New York 1961 vol. I. M. C. R. Symons Quart. Rev. 1959 13 99. (a) J. H. Schulman and W. D. Compton ‘Colour Centres in Solids’ Pergamon London 1963; (6) M. C. R. Symons and W. T. Doyle Quart. Rev. 1960,14,62. M. R. Ronayne J. G. Guarino and W. H. Hamill J. Amer. Chem. SOC. 1962,84,4230. W. C. Fernelius and G. W. Watt Chem. Rev. 1937 20 195. W. L. Jolly J. Amer. Chem. SOC. 1955 77 4958. See C. A. Kraus ‘Solutions Metal-Ammoniac’ Benjamin Paris 1964 p.7. A. J. Birch and H. Smith Quart. Rev. 1958,12 17. 80 Walker solution stable for several hours.g When dissolved in alcohols however the alkali metals produce hydrogen fairly readily; when dissolved in water the reac- tion is vigorous. Any blue colour in aqueous solution if it occurs is extremely transitory. Despite an earlier report1* of an optical absorption from potassium- water solutions some very careful recent work using a fast stop-flow technique showed that any colour was imperceptible on a sensitive spectrophotometer within 2 milliseconds.ll This observation is in accord with the known rate of combination of hydrated electrons. The increasing ease of dissolution of alkali and alkaline earth metals in ammonia and amines and increasing vigour with which these metals react with water to give hydrogen is also the order of increasing electropositivity.From conductivity and paramagnetic susceptibility measurements from e.s.r. and the optical absorption data the species responsible for the single broad absorption band centred at 15,000 8 in dilute (< 10-2~) metal-ammonia solutions has been fairly positively identified as the solvated electron.2 In more concentrated ammonia solutions and in media of lower dielectric constant such as amines and ethers other absorption bands appear which are generally attributable to monomers or quadrupolar ionic aggregates.12 Electrons solvated in ammonia may be similar to the description given for e-aq but from density measurements (a saturated solution of lithium in liquid ammonia being the least dense liquid known at 2Ooc)l3 and the e.s.r.line widths it is known that the electrons create and reside in large vacancies in the solvent on average ~0.10~ A3 a volume otherwise occupied by about three ammonia molecules. There are many features of electrons solvated in polar liquids beyond the scope of this Review that are relevant to discussions on e-aq. For instance the order of relative stability for electrons solvated in ammonia amines alcohols and water (as shown both by the rate of production of hydrogen in alkali-metal solution and the lifetime of the electron absorption spectrum in pulse radiolysis) is exactly converse to the depth of the potential well (if this is directly related to the energy of the absorption maximum or the dielectric constant).Consequently thermodynamic stability is not important. Thus it must be kinetic stability; but it is not governed by the protolytic power of the solvents. Although the con- centration of HS in water is ca.108 higher than in ammonia increases in the pH do not affect the stabilities very much and the rates of the reactions of e- with H+ are known. Perhaps the differences can be expressed by three equilibria (1)-(3) the third of which can be ignored but for the fact that dimeric species exist in amines and ethers. 0 G. Fraenkel S. H. Ellis and D. T. Dix J. Amer. Chem. SOC. 1965,87 1406. lo J. Jortner and G. Stein Nature 1955 175 893. l1 J. L. Dye personal communication. la (a) J. L. Dye and R. R. Dewald,J. Phys. Chem. 1964,68,135; (b) M. Ottolenghi K. Bar-Eli H. Linschitz and T. R. Tuttle J.Chem. Phys. 1964,40 3729. 13 W. L. Jolly ‘Progress in Inorganic Chemistry’ vol. I Interscience New York 1959. 81 QuarterZy Reviews Equilibrium (2) governs the stability of solvated electrons. If as in water and alcohols it is far to the right and (%2-)s readily forms H2 gas which escapes then e-s is highly unstable. In NH the (ef-)s species is only observed at concentra- tions of electrons much greater than 1 0 - 2 ~ and is stable whereas reaction (2) occurs in irradiated water at < 10-6~. For media of low dielectric constant such as amines equilibrium (1) is to the left and (3) to the right which is in accord with the expectation that on going to lower dielectric constant the tendency is for monomer or dimer to be formed in preference to isolated solvated ions. The existence and relative stability of the (e2-)s species are rather intriguing.They also appear as F1 centres in alkali halide crystals formed by the partial thermal or photolytic bleaching of F-centres and have been shown to be dia- magnetic? F-Centres which are excess electrons trapped at anion vacancies in crystals show absorption bands of similar shape to those of e- in polar liquids and in these cases since the cavity size is determined by the lattice parameters the wavelength of maximum absorption (Amax) and cavity size can be correlated through Ivey’s rule.14 Most glassy solids also show colouration when subjected to ionising radiations; Pyrex becomes brown and silica violet. Trapped electrons have been identified by their optical and e.s.r. spectra in low-temperature glasses of concentrated aqueous solutions,’6 organic ethers,16 and alcoh01s.l~ Again (e,2-)s species are in evidence.In liquid systems the absorption band of (e,”-)s appears on the high-energy side of the e-s band whereas in solids they tend to be on the low-energy side a difference which has been attributed14 to the ability of the cavity in the liquids to adjust its size and shape to the constraint imposed by two electrons presumably having paired spins. B. Methods of producing e-aq.-Although the radiolysis of water is unquestion- ably the most direct and convenient method devised for generating e-aq it is valuable to compare the spectrum and reactivity so obtained with e-aq produced in other ways. Some possible method of producing e-sq are discussed. (i) From aZkaZi metals. The reaction of sodium with water is described classically by reaction (4) followed by the combination of two hydrogen atoms to form molecular hydrogen.If by analogy with ammonia the first step involves ionisa- Na + H,O-+Na+ + OH-+ H (4) tion with the formation of e- which survives long enough to become solvated then hydrogen will be formed by reactions (la) and (5). As will be seen the Na + Na+ 3- e-aq (la) 14 M. J. Blandamer R. Catterall L. Shields and M. C. R. Symons J. Chem. Soc. 1964,4357. l6 (a) Ref. 4; (b) F. S . Dainton and G. A. Salmon Proc. Roy. Soc. 1965 A 285,319; (c) D. R. Smith and J. J. Pieroni Canad. J . Chem. 1965,43 876. l7 F. S. Dainton G. A. Salmon and J. Teplj.. Proc. Roy. SOC. 1965 A -27. D. Schulte-Frohlinde and K. Eiben 2. Nuturforsch 1962 17a 445. 82 Walker conversion of e-aq into hydrogen atoms by reaction (6) is very slow whereas e-aq + H20 -+ H + OH-aq (6) reaction (5) is rapid and will be dominant if the local concentration of eeaq close to the metal surface exceeds ~O-?M (since k,/k > lo-').Because of the magnitude of k the spectrophotometric detection of C a p under steady-state conditions is not feasible at present.ll However it should be possible to test for e-aq chemically. Indeed the presence of N,O in water to which dispersed sodium was added (in the form of sodium amalgam) inhibited the formation of hydrogen with the con- current generation of nitrogen by reaction (7). From the known relative rates of e-aq + N2O -t NS + 0-aq (7) reaction of N20 with eeaq and hydrogen atoms and the observed dependence of this inhibition on the pH of the water it was concluded that reactions (la) and (5) rather than (4) occur.18 In addition Bennett Mile and ThomaslQ have reported depositing sodium atoms within an ice matrix at 77"~ and from the line- width of the e.s.r.absorption concluded that electrons were produced and trapped at defects in the ice structure. They were also able to assign fine structure to the interaction of the electron with protons from four water molecules arranged tetrahedrally around the electron centre. Further there is chemical evidence to suggest that e-89 is the precursor of hydrogen gas formed at the cathode when dilute solutions of Na2S04 are electrolysed.eO (ii) Photochemically. An unsuccessful attempt has been made20 to observe the absorption spectrum of e-aq in pure liquid water subjected to flash photoionisa- tion from a windowless helium resonance lamp the predominant light from which was believed to be at 584 A a photon energy well in excess of the ionisa- tion potential of water.However if a solute is added in dilute concentration this may be readily photoionised by light transmitted by the water. Indeed the flash photolysis of aqueous solutions of I- Fe(CN):- phenol cresol etc.21 has been shown to produce a transient species having the same absorption spectrum and reactivity as caq. Photochemical electron-transfer reactions in general (or indeed any electron- transfer process) may involve hydrated electrons as intermediates. In 1938 Franck and Scheibe= supposed that the intense absorption bands in the region 2000-2500 A which are characteristic of many anions in aqueous solution suggested that absorption of light was accompanied by the formation of a free radical and a hydrated electron.Then Dainton and Jamesz3 much more recently la E. A. Shaede and D. C. Walker Chem. SOC. Special Publ. No. 22 1967 p. 277. l8 J. Bennett B. Mile and A. Thomas Nature 1964 201 919. 2o D. C. Walker unpublished work and Canad. J . Chem. 1966 44,2226. (a) L. I. Grossweiner G. W. Swenson and E. F. Zwieker Science 1963 141 805; (b) G. Dobson and L. I. Grossweiner Radiation Res. 1964 23 290; (c) M. S. Matheson W. A. Mulac and J. Rabani J. Phys. Chem. 1963 67,2613. 22 J. Franck and G. Scheibe 2. phys. Chem. 1938 A,139,22. *3 F. S. Dainton and D. G. L. James Trans. Faraduy Soc. 1958,54,649. 83 Quurtedy Reviews showed that the energy of the charge-transfer spectral band of many bivalent transition-metal ions was linearly related to the redox potential of the M2+/ M3+ couple.Now there is chemical evidence for e-rtq as the precursor of H2 when u3+ salts react with and reduce water.20 (iii) Chemically from the hydrogen atom. The acid dissociation constant of the hydrogen atom can be obtained from equilibrium (6). k- has been evaluated24 as 2.2 x lo7 1. moli-l sec.-l while k is the subject of some uncertainty (see later) but the best available value is 16 1. mole-’ sec.-l (ref. 25). Combination of these gives pKa = 9.7 hence e-aq should be fairly readily generated from hydrogen atoms in alkaline solution. was subse- quently confirmed by competitive studies with introduction of externally generated hydrogen atoms into alkaline s0lutions.2~ It has since been shown to be a fast non-rate-controlling step in the radiation-induced chain reaction between H and N20 at pH > 12.28 Spectroscopic verification for this acid-base transfer was provided by the pulse radiolysis of alkaline water containing high pressures of hydrogen Under these conditions the radiation produced OH radicals reacted with H2 to give hydrogen atoms which were then converted into e-aq the optical absorbance of which was measured.(iv) By the radiolysis of water. Despite the suggestion by Stein in 195230 and Platman’s theoretical prediction the following year:1 the importance of hydrated electrons in aqueous radiation chemistry was not revealed and they continued to be mistaken for hydrogen atoms for almost a decade. Although hydrogen atoms and caq are fairly readily interconverted [k8 = 2.3 x lo1* 1.mole-l sec.-l (ref. 32) and k4 = 2.2 x lo7 1. mole-l sec.-l (ref. 24)] providing 3-7 < PH < 14 interconversion cannot occur before their detection by the scavenging technique. This is because for normal scavenging of reactive species R /cRS[s] = ca. lo7 sec.-l where k ~ s is the rate constant for reaction of R with scavenger S. Except in fairly acidic or very strongly alkaline solution the species existing about sec. after the primary events leading to its formation will be scavenged. Radiation chemistry is concerned with the interaction of high-energy radia- tions with chemical systems. Each high-energy electron either from an external ion source or from Compton or photoelectric absorption of X- or y-rays will 24 M. S. Matheson and J.Rabani J. Phys. Chem. 1965 69 1324. 25 E. J. Hart S. Gordon ant r M. Fielden J. Phys. Chem. 1966 70 150. 26 J. H. Baxendale and G. Hughes 2. phys. Chem. (Frankfurt) 1958 14 323. 27 J. Jortner and J. Rabani J . Amer. Chem. SOC. 1961 83 3868. 2* (a) F. S . Dainton and D. C. Walker Discuss Faraday SOC. 1963,36,313; (b) C. H. Cheke and J. W. Swinnerton J. Phys. Chem. 1964 68 1429. 29 M. S. Matheson and J. Rabani Radiation Res. 1963 19 180. 3o G. Stein Discuss. Faraday SOC. 1952 12 227. 31 R. L. Platzman N.A.S.-N.R.C. Reports 1953 No. 305 p. 22. 32S. Gordon E. J. Hart M. S. Matheson J. Rabani and J. K. Thomas Discuss. Faraday Soc. 1963 36 193. e-sq + H20 = H 3- OH-aq (6) Reaction (-6) originally suggested by Baxendale and e-sq + Hfaq 4 ‘H (8) 84 Walker ultimately give rise to an additional ionisation for about every 30 ev of energy.Thus a 1 Mev electron generates about 33,000 other electrons in the irradiated material. Table 1 indicates the succession of events involving electrons that probably take place when a high-energy electron transfers its energy to liquid water. Very quickly the chain ionisation gives rise to a large number of sub- excitation electrons which become thermalised in ca. sec. by vibrational and rotational and finally just collisional energy exchanges with the water molecules. If the electron then survives the dielectric relaxation time of water 10-l1 sec. it can become solvated and acquire the diffusion and interaction properties of chemical entities. This Review purports to show very good reasons Time (sec) a p y X-ray etc.10-l8 3. 10-l8 J 3 . l 10-16 I/ $1 10-13 .1 10-l1 3. ca. 3. ~-H.E. e-sec. e-sub. e-Th. e-aq e-aq 210-3 3. H + OH- TABLE 1 High-energy particle or quantum High-energy electron (104-106 ev) Secondary electrons (10-105 ev) Sub-excitation electrons (< 10 ev) Thermalised electron (ca. 1 kT) Solvated electron Survivor of intra-spur recombination reac- tions Possible decomposition products for believing that e-aq survives for longer than lo-' sec. and is the principal reducing species in the radiolysis of water. (In addition hydrogen atoms are formed in irradiated water but with about one quarter the yield of e-aq.) Platzman3' estimated that reaction (6) the conversion e-aq -f H might occur within sec. and an early attempt to observe e-aq spectroscopically (Had this estimation for the dissociation time been correct e-ap would not yet have been detected.) Samuel and MageeM developed an admirable diffusion model involving hydrogen atoms and hydroxyl radicals formed by the rapid neutralisation of thermalised electrons in a time too short for solvation to occur.Because the diffusion model explained the radical and molecular yields and their dependence on linear energy transfer there was no cause before 1958 to doubt 33 H. Linschitz N.A.S.-N.R.C. Reports 1953 No. 305 p. 39. 34 A. H. Samuel and J. L. Magee J. Chem. Phys. 1953 21 1080. 85 Quarterly Reviews that radiation chemical affects in water resulted simply from the process H,O@H + OH A need for the existence of e-aq in radiation chemistry then arose. 2 Chemical Evidence for e-aq To distinguish chemically between hydrogen atoms and e-aq a reagent is required which reacts with each of these species either (i) at significantly different rates or (ii) to yield different products.(i) All of the reactions of e-a are electron attach- ments and therefore reductions. Some of these are very rapid being just diffusion- controlled and in general e-aq reductions are faster than those of hydrogen atoms. With N,O for instance there is a 104-fold difference with 0 only 2-fold so relative rate determinations can distinguish them as will be seen. When the reactant is a saturated organic compound with an abstractable hydrogen atom then the rate of reaction with hydrogen atoms greatly exceeds that with e-aq since the latter do not appear to participate in abstractions. Thus only hydrogen atoms produce H2 by reaction with methanol.Further because only one of them is charged the presence of inert salts in the solution will affect their rates of reaction with charged solutes differently. (ii) Occasionally an unstable inter- mediate of the electron-transfer reaction dissociates to give different stable products from those resulting from attachment or abstraction of hydrogen atoms. Monochloroacetic acid exemplifies this category of solute. A. Competition Studies.-In 1958 Baxendale and Hughes26 noted an extra- ordinarily large effect shown by Fe3+ and Cu2+ ions on the yield of hydrogen from irradiated acidified aqueous methanol solutions. Hydrogen gas produced in reaction (9) was reduced by H + CHSOH -+ H + CHaOH H + Fes+ -+ Fe2+ 3. H+ but this was understandable if e-89 was the precursor of H in acid solution [reaction (S)] and kll >> kl,.(9) (10) Fe3+ to a much greater extent than competitive reaction (10) woutd predict eaq + Fe3+ -+ Fez+ (1 1) At about the same time Hayon and WeissS showed that an afteration of the pH of an aqueous solution of monochloroacetic acid gave rise to different products so that discrimination was effected here because of condition (ii). In acid soh- tion H was the predominant product whereas C1- was favoured at pH > 4. Their results are reproduced in Figure 2 which shows the competition at pH - 1 between H+ and CICH,.CO,H for e-aq by reactions (8) and (12). At low con- centrations of chloroacetic acid reaction (8) predominates e-aq are converted 35E. Hayon and J. J. Weiss 2nd Proc. Intern. Conf. Peaceful Uses At. Energy Geneva 1959 vol.29 p. 80. 86 Walker ’ ‘ ‘ ’ 0.5 ’ ‘ ‘ ’ ’ 1.0 ’ Concentration (moles litre) FIG. 2. Dependence of radiation yields upon concentration of monochloroacetic acid in aqueous solution at pH = 1. A G(C1-1; A G(H2) ; OG(H202) (Reproduced with permission from E. Hayon and J. Weisss6) eaq + CICH2C0,H -+ C1- + kH2COaH (12) into H-atoms and H is produced by reaction (13). It has subsequently been H + CI*CH2*C02H + H + CICHC02H (13) shown that hydrogen atoms generated by an electrodeless discharge and passed into chloroacetic acid solution produce H via reaction (1 3):’ corroborating the designation of the acidic species as H and implying that at pH > 4 e-aq is pro- duced. Also concurrently Barr and Allen36 observed a discrepancy in the ratio of rate constants (k’) for reaction with O2 and H202 of ‘H atom species produced by irradiation or by reaction (14).k’ had a value of ca. 500 with the product of reaction (14) whereas it was about unity for the radiation-produced species. Further in the irradiation of acidified solutions k’ was again ca. 500. Barr and Allen concluded that either e a q was produced by radiolysis and H-atoms in (14) or H-atoms in radiolysis and an even more acid species such as H2+ in reaction (14). [In fact Hz+ has been postu- lated to be so weakly reducing that it is itself reduced by iodide?’] By discriminative competitive reactions many systems have now shown the principal reducing species in radiolysis to be the more basic forrn. Confirmation and final identification with the hydrated electron comes from the determination of the magnitude and sign of its charge and the spectroscopic evidence.36 N. F. Barr and A. 0. Allen J. Phys. Chem. 1959,63,928. 37 T. Rigg and J. Weiss J. Chem. SOC. 1952 4198. 87 Quarterly Reviews B. Effect of Ionic Strength.-From the Bronsted-Bjerrum theory of ionic reactions and the extended Debye-Huckel theory of electrolytes the rate constant k of reaction between two species of charge ZA and ZB in water at 25"c is given by eqn. (I).38 In this equation p is the ionic strength of the solution a is a parameter numeric- ally close to unity and k is the rate constant at infinite dilution of ions. Thus k will increase decrease or remain constant depending respectively on whether the charges of the reactants are of the same sign opposite signs or one is zero. I ~ Z A for instance is known then a determination of the dependence of k on p will permit an evaluation of the magnitude and sign of ZB.Because the reactions of e a q are very rapid absolute values of k could not be determined in 1962 so it was necessary to compare the effect of p on rate- constant ratios with a charged relative to an uncharged solute. Czapski and S c h ~ a r z ~ ~ established in this way that the principal reducing species in irradiated water had unit negative charge. They observed the rate of reaction of e-89 with H202 compared with H+ 0, and NO2- as a function of ionic strength their results being shown in Figure 3. In this figure K represents the ratio k(e- + log,& = log,,k + 1.02 z ~ . Z ~ p ' / ~ / ( l + ~ p l / ~ ) Q JJ "2 / (I +JJ '12) FIG. 3 . Efect of ionic strength (p) upon the relative rate constants for reactions of e-aq with various solutes.ON02- relative to H202; 00 relative to Hz02; AH+ relative to H20z. K is the relative rate at ionic strength p KO when p = 0. (Reproduced with permission from G. Czapski and H. A. S c h ~ a r z ~ ~ ) k,- + (where x is H+ 02 or NO2-) at ionic strength p and KO is that ratio when p = 0. The lines drawn with slopes of + 1.02 0 and - 1-02 show that 2 for e-aq = - 1. The same conclusion was reached by Collinson Dainton 3i3 G. Scatchard J. Amer. Chem. Soc. 1930 52 52. 39 G. Czapski and H. A. Schwarz J. Phys. Chem. 1962 66,471. 88 Walker Smith and Tazuki4O using Ag+ ions in competition with acrylamide as the charged and uncharged solutes and recently41 measurements on the absolute value of k, have shown it to vary with ionic strength according to eqn.(I). e-aq + Fe(CN),& 3 Fe(CN)t- (1 5 ) Equation (I) is valid only if species A By and the transition complex have at all times the equilibrium distribution of ionic atmospheres. This will be true for the transition complex even if its lifetime is less than the ionic relaxation time if the ion atmospheres of A and B are united when the complex is formed. Coyle Dainton and Logan42 have discussed the implications to the Brarnsted treatment if B in this case the electron is suddenly generated in a solution of A so concentrated that the formation of the transition complex is complete before the electron could acquire an ion atmosphere. Under these circumstances the complex will have a charge of (ZA + ZB) but an ion atmosphere initially equiv- alent to that of A.The 1-02 in eqn. (I) is then replaced by 0-51 which means that an ion having charge -2 without an ionic atmosphere cannot be distinguished from an ion of unit charge in equilibrium with the ions constituting the medium. Coyle Dainton and Logan have shown that the radiation and photochemically produced hydrated electron establishes its ion atmosphere in less than 3 x sec. Thus the e-ag scavenged some lo-' sec. after its formation resembles a univalent negative ion in being solvated and in equilibrium with an appropriate ion atmosphere. There is substantial chemical evidence for the existence of e-aq in irradiated water. In a series of classic experiments Hart and Boa$3 showed by a fast spectro- photographic technique that a strong absorption band in the visible could be identified with e-aq and this permitted direct precise measurements of the reactivities of hydrated electrons.3 Spectroscopic Evidence for e-aq In 1953 Platzman31 remarked 'irradiated water turns blue we just don't see it'. This was a consequence of his postulation that e-aq should exist in irradiated water that it should have a broad absorption maximum at a photon energy of about 2 ev and that it should have a very transient existence. Lin~chitz~~ had earlier failed to detect an absorption not because his time resolution which only permitted observations to be made lo- sec. after a short pulse of X-rays was inadequate (as was thought at the time) but because the sensitivity of detection was too low. In 1960 Keene4* reported an unidentified absorption in the visible region of the spectrum in water irradiated with a 2 psec.pulse of 4 Mev electrons which Matheson suggested may be due to e-aq.4' In 1962 Hart and Boag observed and identified the absorption spectrum of e-aq. 40 E. Collinson F. S. Dainton D. R. Smith and S. Tazuki Proc. Chem. SOC. 1962 140. 1963 85 1375. 42 P. Coyle F. S. Dainton and S. R. Logan Proc. Chem. Soc. 1964 219. 48 E. J. Hart and J. W. Boag J . Amer. Chem. Soc. 1962 84 4090. p4 J. P. Keene see Radiution Res. 1964 22 1 ref. (1). O5 M. S. Matheson Ann. Rev. Phys. Chem. 1962,13,90. S. Gordon E. J. Hart M. S. Matheson J. Rabani and J. K. Thomas J . Amer. Chem. SOC. 89 Quarter Iy Re views electron 1 otcelerotor H lriqger H delay H :t:::e I A. The Technique of Pulse Radio1ysis.-Owing to the availability of accelerators which provide very short pulses (0.01 to 5 psec.) of high-energy electrons (1.5 to 30 MeV) the radiation analogue of the photochemical technique of flash photol- ysis namely pulse radiolysis has been most fruitfully applied in recent years.The electron beams are sufficiently intense to initiate enough reaction during the short pulse either to measure the rate of formation of a product or the decay of a reactive intermediate immediately afterwards or even to study the build-up or steady-state concentration during the pulse. The first applications of absorp- tion spectroscopy to pulse radiolysis in 196P6 quickly led to a perfection of the technique. Figure 4(u) shows the type of apparatus used by Hart and Boae3 rodiotion shield OSCl lorcope eleclron I H delay trigger H oscli'oscope J time base to obtain the spectrum of e-8q in sodium carbonate solution and then in pure water.A uranium-spark spectroflash was triggered simultaneously with or at a measured time delay after an electron pulse from a 1.8 Mev electron accelerator the absorption spectrum being photographed on a spectrograph. In their initial experiments a 0-5~-Na,cO solution showed a well-defined maximum at about 7000 A which decayed completely within 25 psec. In pure water the e-89 absorp- tion was much shorter lived; indeed the carbonate solution was an extremely favourable environment for this first observation not only because of the extended lifetime of e-aq owing to the pH but also owing to the removal of OH radicals by C032- to give C0,- which absorbs slightly in the same region. This synchronised flash absorption spectrophotographic method can give complete spectra at a variety of time intervals after the electron pulse and can be used to obtain kinetic results.It has a number of disadvantages however. First 46 (a) R. L. McCarthy and A. MacLachlan Trans. Faraday SOC. 1960 56 1187; (b) M. S. Matheson and L. M. Dorfman J. Chem. Phys. 1960,32 1870; (c) J. P. Keene Nature 1960 188,843. 90 Walker the time resolution is limited by an integrated time-intensity factor during the period of the spectroflash (usually a few psec.); secondly the sensitivity is governed by the speed and contrast of the photographic plates; and finally optical density measurements have to be calculated from densitometer traces and compared with blanks showing the sensitivity of the emulsion and spectroflash to various spectral regions.The spectrum obtained by Hart and Boag thus showed certain structural features and secondary bands not observed by the use of the rapid spectrophotometric recording technique though both have similar resolving power (ca. 100 A). The kinetic spectrophotometric method is now more generally used and is depicted schematically in Figure 4(6). Light from a stabilised source passes through the irradiation cell is monitored by a photomultiplier on the exit slit of a monochromator and the resulting signal is fed onto the Y-plates of an oscilloscope to which a suitable time base is applied just before the electron pulse. The oscilloscope screen is photographed and thus a time-dependence of the absorption is recorded for each wavelength setting of the monochromator.This optical detection system compares the transmission of the irradiation cell before the pulse with that during and after the pulse. The smallest change in transmission which can be detected is governed only by fluctuations in the lamp output (very small over a few tens of microseconds) resulting in much greater sensitivity than is possible in the dual-cell systems of conventional spectro- photometry Shot noise in the photomultiplier is the principal limitation over very short times4' and since this is inversely proportional to the square-root of the light flux through the cell an efficient optical system and high intensity lamp yield enormous sensitivities. Optical-density measurements of 0.002 are com- mon4' and 0.00002 is reported to be With this sensitivity concentra- tions of 2 x 10-9~ can be measured in cells 1 cm.in depth for species having E = lo4 1. mole-l cm.-'. Multiple-reflection cells further increase the effective sensitivity of the optical detection and have been used successfully with the higher energy electron B. Absorption Spectrum of e-ag.-Figure 5 shows the absorption spectrum of eaq in pure deaerated water obtained by KeeneS0 using a spectrophotometric method. The band has an absorption maximum at 7200 + 100 8 which corre- sponds with an energy of 1.72 ev. It is a very broad band having a half-height width of about 1.0 ev and is asymmetric on the high-energy side when repre- sented on a linear energy scale; but there is no discernible fine structure to the band between 8000 and 3000 8 on instruments capable of 100 8 resolution.Rabani Mulac and Matheson5I have determined the extinction coefficient with- out the need to measure the radiation intensity or assume a chemical yield of 47 J. P. Keene J. Sci. Instn. 1964 41 493. 48 I. A. Taub Proc. Manchester Symposium on Pulse Radiolysis 1965 Academic Press London. 49 L. M. Dorfman Science 1963 141,493. 5o J. P. Keene Discuss. Faraday Soc. 1963 36 304. 51 J. Rabani W. A. Mulac and M. S. Matheson J. Phys. Chem. 1965 69. 53. 91 QuarterZy Reviews k 8000 6 0 0 0 4000 Wavelength ( A ) FIG. 5 . Absorprion spectrum of e-aq. (Reproduced with permission from J. P. KeeneS0) e-aq. They observed the decay of e-aq concurrent with the formation of the tetraform ion produced in reaction (1 6). From the known E of C(NO,),- the molar decadic extinction coefficient of eag at 7200 8 was computed as 15,800 1.mole- cm.-l and from this and Figure 5 an oscillator strength for the optical transition of about 0.8 can be estimated. So far a full spectral analysis in the near infrared has not been reported and it may prove important in discussions on the optical transitions giving rise to the absorption band to know if fine structure on the low-energy side of the peak exists . Studies with deuterium oxide52 show a small isotopic shift in the absorption maximum to higher energies and probably a very significant increase in extinc- tion coefficient and oscillator strength. r h i s assumes that the radiation yield of solvated electrons is not more than 30 % larger in D20 than in H20. In this regard it would be valuable to measure the extinction coefficient of e-aq in D20 by the tetranitromethane method as a direct check on both E and G(eag).] Anbar and Hart53 have shown the absorption band of e-aq in concentrated solutions of inert electrolytes such as MgCI, KF NaOH NaClO, etc.to be shifted to shorter wavelengths accompanied by a small reduction in the absorbance probably owing to a tightening of the solvent shell by the ionic atmosphere. Decreasing 52 D. M. Brown F. S. Dainton J. P. Keene and D. C. Walker Proc. Chem. Soc. 1964 266. 53 M. Anbar and E. J. Hart J. Phys. Chem. 1965,69 1244. 92 Walker temperature also causes a shift in the absorption band to higher energies.54 Although eaq is very short-lived its intense absorption band spanning the readily accessible region of the spectrum enables its reactions to be studied unambiguously accurately and kinetically.C. Identification of the Absorption Band.-Evidence for the assignment of the transient absorption band shown in Figure 5 to eaq is rather convincing. (i) The energy of Amax agrees reasonably closely with the early prediction?l the current theoretical estimates,% and the solvation energy determined from thermo- chemical (ii) The band’s shape proportions and asymmetry are analogous to those of electron bands in metal-ammonia solutions F-centres and other trapped electron species which have been identified by their e.s.r. absorption. (iii) Additives which are known to react with eaq such as H+ O, H,O, N,O CO, Cu2+ etc. suppress the absorption band whereas reagents which do not react rapidly with e-aq such as CH,OH NH4+ Na+ etc. do not affect the absorption.(iv) The rate of disappearance of the e-aq band is exactly comple- mented by the rate of formation of the C(NO,),- absorption band when C(NO,) is the additive.51 (v) An absolute value for the yield G = 2.6 + 0 ~ 2 ~ ~ of species responsible for the absorption equals the best estimates of G(e-ag) by chemical methods.57 (vi) Measurements of transient electrical conductivity by Boa$* on pulse-irradiated water showed that a charged species was produced which decayed at the same rate as the optical absorption in the presence and absence of additives. (vii) Kinetic ionic-strength effects have shown that the band was due to a species having unit negative charge.41 (viii) A transient species having the same absorption spectrum and reactivity was obtained by flash photolysis of certain aqueous solutions which are believed to undergo photochemical charge- transfer-to-solvent processes involving e-aq.21 (ix) A strengthening in the eaq band is observed when H and OH- are in sufficient concentration for the radiation-produced OH radical to be transformed to eaq via reactions (14) and Measurements of the e.s.r.absorption of eaq are not feasible at present on account of its very transient existence. Since the modulated magnetic field can- not be scanned rapidly enough it may only be possible to obtain the e.s.r. spectrum of e-aq by a sampling technique. Further because of the long time- constant required for good signal-to-noise ratios the decay characteristics may have to be studied by an intermittancy technique similar to that used by Fessenden for hydrocarbon radicals.59 Clearly e.s.r.measurements on eaq are vital in elucidating the nature of the electron binding and the size shape and orientation of the solvent sheath and would corroborate the optical studies on the reactivity of eaq. ( -6).29 54 J. H. Baxendale and J. P. Keene unpublished work. 55 J. Jortner Radiation Res. 1964 Suppl. 4 24. 56 J. H. Baxendale Radiation Res. 1964 Suppl. 4 139. 57 M. S. Matheson Radiation Res. 1964 Suppl. 4 1. 58 J. W. Boag Amer. J . Roentgenol. 1963 40 896. 59 R. W. Fessenden J. Phys. Chem. 1964 68 1508. 93 Quarterly Reviews 4 Reactivity of e-aq A. Reactions in Irradiated Water.-The radiolysis of pure dmrated water by high-energy electrons can be described by the following equation H@We-aq(2-5) + OH(2.3) + H(O.6) + H2(0.45) + H,O2(0-85) + H30+(2*5) where the numbers in parentheses are the yields per 100 ev of energy absorbed.e-aq reacts rapidly with all these species except H and possibly albeit slowly with water. By adjusting the electron current of the accelerator the dose per pulse can be varied up to ca. 5000 rads. Typical oscilloscope traces obtained by the kinetic spectrophotometric (0) 99.6 technique are shown in Fig. 6 for samples of neutral 1 l I " I " I I I ' 0 20 40 60 80 100120 140 160 180 200 Time (y sec) ( b ) n .- .r 94 & 'ifll 5 0 2 4 6 8 10 12 14 16 18 20 Time() sec) (d u 1 1 . 1 I 1 I ' I ' 0 2 4 6 8 10 12 14 16 18 20 Time (J sec) FIG 6. Oscilloscope traces showing the variation of percentage light transmission as a function of time for (a) low (6) intermediate and (c) high-intensity electron pulses.water irradiated with 2 psec. pulses of 4 Mev electrons (a) corresponds to the deposition of ca. 50 rads (b) ca. 500 rads and (c) ca. 5000 rads. At the highest dose the concentration of e-aq at the end of the pulse is ca. 1 O W 6 ~ and the rate of decay of eaq is presented approximately by second-order kinetics. e-aq can undergo the following reactions with radiation products 94 Walker Since e-aq OH and H30+ are produced in comparable amounts reactions (17) (18) and (20) should be truly of the second order but with different rate con- stants. Since H and H202 are produced with smaller yields than e-aq reactions (19) and (21) will contribute to a complex and time-variable extent to the decay of e-aq; in fact since H202 is a stable product (21) may be the dominant process in the later stages of the decay.The initial slope of a high-dose second-order plot would thus give a composite rate constant k given by eqn. (10. where [ lo corresponds to the concentrations of the various species at the end of the pulse. If reaction were not occurring during the pulse the concentration ratios would be given by the ratio of radiation yields. Gordon Hart Matheson Rabani and Thomas32 obtained a series of theoretical curves by numerical computation on the basis of these reactions for selected values of rate constants and radiation yields and compared their experimental curves with the closest- fitting calculated curve to obtain the values k17-k, quoted in Table 2. Dorfman and TaubG0 have shown unambiguously that the combination of two hydrated electrons in reaction (17) give rise to molecular hydrogen not hydrogen atoms.This was done by the pulse radiolysis of O.~M-C,H,OD in D,O which gave yields of D and HD G(Da = G, + 3;Ge-aq and G(HD) = G, expected on the basis of reaction (17). B. Lifetime of e-aq in Water.-At doses per pulse less than 50 rads when the concentration of e-aq is less than ~O-,M first-order decay rates are prevalent. At ~O-,M however the rate of reaction (20) with respect to the H+aq present in neutral water resulting from the ionic dissociation of water should follow second-order kinetics because this ionic dissociation is extremely slow. For smaller doses when [H30+] is essentially constant reaction (20) imposes an upper limit of 300 psec. on the lifetime of e-aq in neutral water so that the pH of the water must be adjusted appropriately for studies of reaction (6).Early estimates of k from first-order decay plots at doses of ca. 50 rads sur- prisingly gave values of similar magnitude in different laboratories [ 2 4 X lo4 (ref. 32) < 4.4 x lo4 (ref. 60) 2-7 x lo4 (ref. 61a) < 5.5 x lo4 (ref. 616) 4.3 x lo4 sec.-l (ref. ~Ic)] equivalent to a half-life for e-aq of 20-25 psec. It 6o L. M. Dorfman and I. A. Taub J. Amer. Chem. SOC. 1963 85,2370. (a) J. P. Keene Radiation Res. 1964 22 1; (6) J. Rabani and G. Stein J. Chem. Phys. 1962 37 1865; (c) G. Dobson and L. I. Grossweiner Trans. Faraday SOC. 1965 61 708 (from flash photolysis). 95 Quarterly Reviews was subsequently showns2 that this half-life (4) could be increased to 47 psec. by more rigorous purification particularly complete elimination of CO and by pre-irradiation.An increase in the pH to 9.8 increased t i further to 58 psec. and a dose effect noted whereby t+ increased as the dose per pulse was diminished. Thus values of 100-300 psec. were obtained for t g for 10 rad pulses with chemically purified water.62 Very recently Hart Gordon and Fielden% have examined reaction (6) using hydrogen-saturated solutions (to alleviate the build-up of H202 during the pre-irradiation purification) at pH = 8-3 (high enough to eliminate reaction (20) but not sufficiently alkaline to involve reaction (-6)) and at very low doses where [e-aq] < 10-8~. They obtain an ultimate limit of t g = 800 psec giving k = 16 1. mole-l sec.-l. Dewald Dye Eigen and DeMae~er,~ have made some most interesting measurements on the decay of the electron absorption when water is added to a dilute solution of electrons solvated in ethylenediamine.They find a first-order decay of e- after mixing is complete the reaction rate being proportional to the water concentration used and conclude from their results that k2 = 20 1. mole-l sec.-l. e-s + H20 + H + OH-s (22) It has been pointed that this value is very close to the best value obtained from pulse radiolysis for k, but some ambiguity exists because the nature of e- in reaction (22) is unknown. If (22) is the reaction of H20 with electrons solvated in ethylenediamine as seems likely since the kinetic results come from the rate of decay of this species then it is probably not pertinent to compare k with k22. Because of the large variation in water concentration studied it is unlikely that the conversion e-s 3 e-aq is a non-rate-controlling step in this experiment and hence its relevance to reaction (6) is unknown.Since in neutral water reactions (17) or (20) will predominate over (6) the latter will have no practical significance. Reaction (6) is however of fundamental importance in the estimation of thermodynamic data and because it determines the relative stability of eag and hydrogen atoms in neutral water. Does e-aq decay spontaneously or react with water i.e. does reaction (6) occur? The evidence for its occurrence is based solely on the fact that under conditions of great purity and at very low concentrations the first-order decay plots are linear. This may be unsubstantial evidence because first the concentration-dependence of the initial rates is in general a better test of the order of a reaction than the time- dependence of the rate and indeed the initial rate has been shown to be con- centration-dependent decreasing with decreasing dose per p ~ l s e ? ~ ~ ~ Secondly because of the inhomogeneity of primary radiation absorption in spurs and h a y s at very low doses the rate of reaction will not be uniform throughout the sample so that second-order plots of reactions (17) (18) and (20) would be non- 62 J.H. Baxendale M. Ebert J. P. Keene and A. J. Swallow Proc. Manchester Symposium on Puke Radiolysis 1965 Academic Press London. 63 R. R. Dewald J. L. Dye M. Eigen and L. DeMaeyer J . Chern. Phys. 1963 39 2388 (value of k,2 revised slightly recently; personal communication from J. L. Dye). 96 Walker linear anyway.Further since the number of species per spur is small (2 to 20) many intraspur combinations will show first-order kinetics and their contribu- tion to the overall decay should increase continuously with diminishing doses. At 2 rads per pulse for instance the mean separation of spurs is ca. 1.6 x 104 A so that about 400 psec. must elapse before the electron concentration can be considered uniform at 5 x 10-9~ and hence the actual local concentrations initially are not given simply by the dose and radiation yields. An evaluation of the Arrhenius parameters for the supposed reaction (6) might resolve the dilemma. If the steric factor for all eaq reactions is close to unity then an activation energy of some 12 kcal. mole-l can be expected if k = 16 1. mole-l sec.-l whereas a typical activation energy for diffusion of 2-3 kcal.mole-l might apply if the bimolecular reactions predominate or reac- tion with impurity is the limiting factor. Further if reaction (6) does occur then its effect should be curtailed and the lifetime of e-aq extended by increasing the pH of the water to 12- 14 whereupon equilibrium (6) will be well to the left. C. Reactivity with Solutes.-If a solute S which reacts with e-aq according to the generalised reaction (23) is added in sufficient concentration so that eaq generated by pulse radiolysis react almost exclusively by reaction (23) then eaq + S - [S-] -+ stable products (23) provided [S] >> [e-aq] pseudo-first-order decay kinetics are observed. kZ3 is hence obtained directly from the observed half-life of eaq. Because of the in- tense broad absorption band of e-,q and the sharp time resolution of the pulse characteristics a vast wealth of rate results pertaining to e-aq is readily obtained and now available.Anbar and Netas4 have tabulated absolute rate constants for 280 reactions represented by reaction (23) including 164 with organic solutes. In Table 2 a few of these are presented principally to exemplify the diversity and trends in the reactivity of eaq. Since all reactions of e-aq are either electron-attachment or dissociative electron-transfer processes one might expect that the availability of a suitable low-energy molecular orbital governs the reactivity of a solute. Although a number of reaction rates involving e-aq appear to be limited only by the en- counter rate such as that with Co(CN)Z+ this is not so for kS0 the reaction between eag and H+aq since the diffusion-controlled limit should be ca.5 x 1O1O 1. mole-1 sec.-l for this charge-neutralisation process. If this is an electron- transfer reaction then the product of reaction (20) will be either H30 or H,O (depending on the nature of H+aq)" but not H immediately. There are other reasons for believing H,O (or H,O,) to have a chemical importance.66 There are striking differences in the reactivity of e-aq with various organic 64 M. Anbar and P. Neta Intern. J. Appl. Radiation 1965 16 227. 65 M. Eigen and L. DeMaeyer in 'The Structure of Electrolytic Solutions' (ed. W. J. Hamer) Wiley New York 1959. 66 (a) H. J. Bernstein J. Amer. Chem. Soc. 1963 85 484; (6) E. J. Hart Science 1964 146 19; (c) M. Anbar and D. Meyerstein J .Phys. Chem. 1965,69,698; ( d ) T. J. Sworski J. Amer. Chem. Soc. 1964 86 5034. 97 Quarterly Reviews TABLE 2 Reactivity of various chemical species ( X) with e-89 in dilute aqueous x (1. mole-l sec.-l) X k(l. mole-l sec-l) solution at pH 7 (values taken from ref. 64) e-aq 5 x lo9 c0(cN)6s+ 9.0 x lolo H 2.5 x 1O'O C(NO,)* 4-6 x lolo H202 1.3 x 1O1O Naphthalene 3.1 x 10' H2O < 16 CH3CO2H 1.8 x 10' OH 3 x 1O'O Fe( CN),* 3.0 x 109 H+aa 2.3 x lolo Ethanol < 105 (PH 5.4) (PH 10) 0 2 2.0 x 1010 CH3COZ- < 1od N20 8.7 x lo9 Benzene < 7 x lo* Cu2+ 3.0 x 1O1O Nitrobenzene 3 x lOz0 Cd2+ 5.2 x 1O1O Glycine < 107 Na+ < lo6 Ribose < 107 NH,+ ca. 2.5 x lo6 Purine 1.7 x 1Olo solutes. Saturated hydrocarbons alcohols ethers and amines having no low- lying vacant orbitals react sparingly k < lo7 1.mole-l sec.-l. Compounds containing unsaturated groups or halogens have high reactivities many showing rate constants well in excess of 1O1O 1. mole-l sec.-l. There is often a striking dependence on the extent of ionic dissociation; thus undissociated acetic acid has a comparatively high reactivity whereas the acetate ion is virtually unreactive. Amino-acids and sugars on the whole are unreactive whereas the pyrimidine and purine bases present in nucleic acids and other biologically important media are extremely reactive. Since e-aq is the predominant radiation-produced reducing species in aqueous systems and since the phenomenon of indirect action will apply in most biological material consisting of 80-90% water these rate- constant data are clearly apposite to an elucidation of the biological effects of high-energy radiations.e-aq is unreactive towards benzene phenol and toluene and in general behaves as an efficient and sharply discriminating nucleophilic reagent.67 Indeed its reactivity towards aromatic compounds is related to the electron density of the n orbitals in the ring as expressed by Hammett's a-function. 5 Structure of e-aq Although much information is available pertaining to the kinetic and thermo- dynamic properties of eaq its nature and structure are still the subject of con- jecture. Jortner has drawn comparisons with models of electron binding in other media to make approximate quantum mechanical calculations of the cavity size and of the energy levels and oscillator strength corresponding to the optical transition all with reasonable success.This section deals with the variety of theories on electron binding. 67 M. Anbar and E. J. Hart J. Amer. Chem. SOC. 1964,86 5633. Walker A. Electron Binding in Various Media.-Whereas many gross similarities exist between the properties of electrons trapped in widely differing environments certain detailed features seem to mitigate against an enveloping structural model. Different theories seem preferable for systems intrinsically very similar as for excess electrons in the non-polar liquids helium and argon. In liquid helium short-range repulsions between the electron and helium atoms strongly suggest a large cavity or bubble structure containing the electron,”* whereas strong long-range attractions in liquid argon lead to a quasi-free electron model in which the core-electron interactions are neglected.69 Sometimes the data are conflicting.For instance Anbar and Harts3 have observed a single absorption band centred at 9200 A attributable to the solvated electron in pulse-irradiated ethylenediamine whereas the absorption band corresponding to this species in solutions of metals in ethylenediamine occurs at 12,600 Excess electrons in various liquids may be free or localised. When the electron moves too freely the liquid structure may not be perturbed as seems to be the case for liquid argon. Localisation in cavities however may arise from local fluid dilation as a result of short-range repulsions between the extra electron and electrons of the medium. This may be very short-lived as in liquid helium or long-range polarisation effects may provide some stability.If the polarisation is merely electronic then it can occur in both polar and non-polar media but for orientational polarisation the molecules of the medium must have a dipole moment and the time required for molecular rotation is necessary for solvation. In 1933 Landau71 suggested a model for F-centres in ionic crystals suggesting the electron was trapped in the crystal structure by the polarisation of the dielectric medium induced by the presence of the extra electron. This ‘polaron’ model is now untenable for F-centres because the latter are best described as electrons at anion vacancies where clearly there is no electrostatic discontinuity and hence no extraneously induced polarisation. However Landau’s polaron theory has been applied to explain electron binding by self-trapping in metal- ammonia solutions by Davydov,72 Deigen,73 and Platzman and Fran~k,~* and these theories are discussed by Jortner Rice and Wilson.75 (i) Metal-ammonia solutions.Both the diminution of density and hyperhe splitting of the e.s.r. absorption indicate that the electron is located in a vacancy of some 3 A radius in metal-ammonia solutions. This species is responsible for the absorption band centred at 15,000 A. In estimating the energy levels available to the electron within such a cavity one could consider it to be either an electron- in-a-box problem the electron being confined to a spherical volume of uniform field with infinitely steep potential walls or nearly hydrogen-atom like the field 68 W. T. Sommers Phys. Rev. Letters 1964 12 271.69 H. T. Davis S. A. Rice and L. Meyer J. Chem. Phys. 1962 37 2470. 70 R. R. Dewald and J. L. Dye J. Phys. Chem. 1964 68 121. 71 L. Landau Physik Z . Sowietunion 1933 3 664. ’* A. S. Davydov Z. eksp. theor. Fiz. 1948 18 913. 73 M. F. Deigen Trudy Znst. Fizik. A . N. Ukrain. S.S.R. 1954 5 105. 74 R. L. Platzman and J. Franck 2. Physik 1954 138 411. 75 J. Jortner S. A. Rice and E. G. Wilson in ‘Solutions M’tal-Ammoniac’ (ed. Lepoutre and Sienko) Benjamin Paris 1964. 99 Quarterly Reviews decreasing coulombically from the central point. An electron-in-a-box calcula- tion by Ogg76 led to the result that for minimum energy the cavity would have to be 10 A in radius though subsequent refinements including those of electronic polarisation and surface-tension considerations yielded a cavity radius of ca.4 A.77 An LCAO calculation was attempted,78 the electron being assumed to have 1s and 2p wave functions centred on hydrogen atoms forming the walls of the cavity but this model is invalidated by the observation that the appropriate proton chemical shift in the nuclear magnetic resonance (n.m.r.) spectrum was absent.79 The compromise model depicts the electron as incompletely confined to a cavity of uniform field and thus experiencing a coulombic fall-off outside. Although the energy levels will then be essentially hydrogenic because of the penetration of the electron into the surrounding medium bulk properties such as the polarisation energy of the medium regarded as a continuous dielectric can be used instead of coulombic or exchange interactions.This forms the basis of the continuum treatments based on Landau's polaron concept. Jortner Rice and Wilson75 have pointed out that a priori calculations of charge distribution and electron energy levels are not yet possible but that two approximations can be applied. The first of these the electronic adiabatic approximation assumes that the extra electron is much more loosely bound and therefore of lower mean velocity than the valence and core electrons of the medium; consequently the extra electron is affected by the potential resulting from the average charge dis- tribution of the electrons of the medium. Jortner's quantum mechanical treat- ment of this using a one-electron wave function leads to an expression for the potential which at large interaction distances is proportional to (l/Dop - l/Ds) where Dop and Ds are the high-frequency and static dielectric constants respec- tively.In the limit then this treatment gives the same results obtained in the earlier polaron rnodeI~.'~,~~,~* Whereas in a crystal lattice the binding energy of an extra electron may be only 0.1 ev and the adiabatic approximation reasonably valid in polar liquids this may be 1 to 2 ev and hence the extra electron is not substantially more weakly bound than valence electrons. Consequently electronic polarisation should be considered in estimating the binding energy since the electrons of the medium will be affected by the mean charge distribution of the extra electron. Jortners0 has discussed this second approximation which treats the extra and medium electrons on an equal basis.Since the solution of the potential functions depends on the charge distribution of the electron itself this treatment is a self- consistent field (S.C.F.) approximation. Using wave functions for a 1s ground state and 2p excited state Jortner has applied the S.C.F. method with consider- able success to metal-ammonia systems. He computes a value for the energy of the (2p t 1s) transition of 0.93 ev when a cavity radius of 3-2 8 is taken. This 76 R. A. O g g Phys. Rev. 1946 69 668. 77 R. A. Stairs J. Chem. Phys. 1957 27 1431. 70 J. V. Acrivos and K. S . Pitzer J. Phys. Chem. 1962,66 1963. J. H. Simpson Proc. Roy. SOC. 1949 A,197 269. J. Jortner Mol. Phys. 1962 5 257. 100 Walker compares favourably with the experimental value of 0-80 ev for the energy of A,, in ammonia.81 (ii) F-Centres.For F-centres the e.s.r. data show the electron to be largely con- fined to the vacancy but interacting to some extent with the first shell of cations.82 For these systems too a semi-continuum model termed a point-ion-lattice approximation,82 has been developed in which it is assumed that the dipoles of the ions surrounding the F-electron point toward the centre of the cavity rather than following the detailed motion of the electron. Thus the absorption band is again described approximately by a (2p t 1s) transition. B. Jortner’s model for eeaq.-By analogy with the metal-ammonia solutions Jortners6 has performed a variational calculation using hydrogenic-type wave functions for the ground and excited states based on the continuum model with the approximations applied previously.Since the binding energy of the extra elec- tron in water should be higher than in ammonia because of the greater dielectric constants one may expect the S.C.F. to be the better approximation. Choosing appropriate one-electron wave functions for the ground Is and first excited state 2p Jortner calculates the energy of the optical transition hv = E2p - E18. Difficulties of non-spherical charge distribution in the excited 2p state are dis- regarded (as if the state were triply degenerate) and in addition it is a non- equilibrium state. This arises because the Franck-Condon restriction on elec- tronic transitions allows the electronic polarisation but not the orientational polarisation to adjust itself to the charge distribution of the excited state. The orientational polarisation in the 2p state is consequently given by the potential function of the 1s charge distribution fls.Thus Jortner describes the energy of the transition by the differences in the zero-point orientational and electronic polarisation energies of the two states according to eqn. (III) where the symbols have their usual meaning in quantum mechanics. The wave functions $le and $2p contain variational parameters which are related to the mean charge distribution of the state through a function governed by the cavity radius R. Jortner thus computes hv for various values of R and finds that hv - 0.93 ev when R - 3.3 A (as for electrons solvated in NH,) whereas hv = 1-35 ev when R - 0. Since the experimental value for hv is 1.7 ev the cavity radius of e-aq appears to be vanishingly small.However it should be noted from these M. Gold and W. L. Jolly Znorg. Chem. 1962 1 818. 82 B. S. Gouray and F. J. Adrian Solid State Phys. 1960 10 127. 101 4 Qmterty Reviews calculations that the mean radius specifying the charge distribution in the ground state is 2.5-3.0 A and in the excited state is 4.9 A so that although the size of the cavity is trivial the electron is rather widely spread out over the sur- rounding water molecules. The value of 2.5-3.0 A should thus be appropriate for the evaluation of the diffusion coefficient and collision radius of e-&q. A success of the model is that values for these parameters so obtained show reasonable agreement with those from kinetic data.32 On the basis of the larger surface tension density and dipole moment of water than of ammonia one would certainly have expected the electron to be associated with a much smaller and deeper trap in water.Despite the evident success of Jortner's theoretical treatment there remain a few factors requiring consideration. (i) Whereas one may qualitatively conjecture that the spectral band derives its broadness from variations in the size and shape of the cavities owing to disorder from thermal motions it is not only very broad (half-width ca. 1 ev) but has similar proportions to the band in ammonia despite the exceptional differences in cavity size calculated for the two solvents. In both solvents there is asymmetry on an energy scale on the high-energy side but no discernible second band corre- sponding to ( 3 p c 1s) transitions. Nor does e-aq show signs of a convergence limit at 5400 A.Perhaps this fine structure is masked by the broadness but accounts for the asymmetry. On this basis at lower temperatures some of the thermal disorder should be removed but D ~ r f m a n ~ ~ reports there to be no fine structure or narrowing of the solvated-electron band in irradiated liquid ethanol even at - 78"c. However using the transition dipole moment Jortner calculates the oscillator strength of the (2p t 1s) transition to be f = 1-1 when R = 0. This compares with the experimental value of ca. 0.8. If this can be regarded as a normal one-electron system in which cf= 1.0 for all transitions from any one energy level then the oscillator strength for higher excited transitions might indeed be too small to observe. (ii) In these calculations the excited state is assumed to be a bound state.Conse- quently eaq should show neither photoconductivity nor photobleaching when illuminated with light corresponding to the wavelength of maximum absorption (Amax). Such an experiment is difficult to perform on e-aq and analogies with other electron systems is conflicting. Irradiated glasses at low temperatures readily undergo photobleaching of the electron bandsls whereas metal-ammonia solutions do not show photocond~ctivity.~~ (iii) The shift of the absorption band with temperature should be mainly due to changes in the cavity size rather than in Ds or D0p.55 Consequently Jortner predicts that d(hv)/dT for metal-ammonia solutions should be 2 to 3 times greater than for eaq. This is quite contrary to the experimental findings. d(hv)/dT has values of 10.5 cm.-l deg.-l for solutions of sodium in ammoniaE5 and 1-5 cm.-l 83 L.M. Dorfman 'Advances in Chemistry' Series 50 1965 American Chemical Society. 8p R. L. Potter R. G. Shares and J. L. Dye J. Chem. Phys. 1961,35 1907. 85 (a) H. Blades and J. W. Hodgins Canad. J . Chem. 1955 33 41 1 ; (6) R. C. Douthit and J. L. Dye J. Amer. Chem. SOC. 1960 82 4472. 102 Walker deg.-l for potassium in ethylamineS6 but 25 cm.-l deg.-l for e-aqM and 29.2 cm.-' deg.-l for e- in ethanoLS3 (iv) Jortner's model appears to predict a (2p t 1s) transition energy for electrons solvated in aliphatic alcohols intermediate between water and ammonia. In fact Amax for methanol appears at ca. 2.1 e ~ . * ~ C. Alternative Models.-If the spherically symmetric potential well of the electron solvation sheath contains various energy levels these must be quantised.If the potential function is hydrogenic then the separation between successive levels diminishes towards the convergence limit and the first transition 2p c Is is three-quarters of the well depth. This is demonstrated in Figure 7(a). Baxendaless has calculated the hydration energy ( AHs) of e-aq to be 1.72 ev exactly equal to EAmax. This may or may not be fortuitous. If the polarisation energy of the medium Ep is substantial then the energy required to remove the electron from the well the ionisation potential I p will be different from AHs as indicated in Figure 7(b). If the coincidence Ep M *EAma then A Hs will be given by E,\,,,. Figure 7(6) may represent the Jortner model the asymmetry arising from weak but significant transitions to higher excited levels than 2p.On the other hand if the disparities cited above are substantiated particularly regarding the photobleaching at Amax which occurs for the low-temperature glasses (similarities between solvated and trapped electrons being discussed later) then perhaps Amax corresponds to the release of the electron from the trap. At wavelengths less than Amax the electron may be excited into a conduction band giving rise to a continuum on the high-energy side. This is depicted in Figure 7(c) the absorption band corresponding to an ionisation-efficiency profile represent- ing the transition probability of excitation to levels above the ionisation potential. The ill-defined low-energy cut-off may result from the numerous structural variations of the potential well and perhaps sub-ionisation excitation levels within the trap such as EAp in Figure 7(c).A combination of ionisation con- tinuum and symmetrical charge-transfer absorption band is also a possibility particularly if neighbouring unexpanded cavities exist as Freeman and Fayadh have suggested even for non-polar liquidsa7 If optical absorption corresponds to the release of an electron from a potential well the broadness asymmetry and shape could alternatively result from the occurrence of many orientations of solvent molecules giving rise to wells of differing depths. Figure 7(d) shows this Amax merely corresponding to the most probable but comparatively tenuous trap. If the distribution of traps is not in thermal equilibrium photolysis at Amax would cause the absorption maximum to shift to higher energies since only traps of energy < EAmax would be affected.Photolysis at Amax in the case represented by Figure 7(c) would lead to a diminu- tion of the absorbance without altering the shape of the band. Despite enormous differences in stability electrons caught in the polarisation fields of polar liquids or glassy or crystalline solids all have very similarly shaped 88 L. R. Dalton J. D. Rynbrandt E. M. Hansen and J. L. Dye personal communication. G. R. Freeman and J. M. Fayadh J. Phys. Chem. 1965,43 86. 103 Quarterly Reviews well depth FIG. 7. Energy levels corresponding to possible absorption processes (see text for details). absorption bands. Whether it is metal-ammonia solutions peaking at 15,000 A or F-centres with band maxima in the visible region the shape and proportions are analogous.Are these ‘trapped‘ or ‘solvated’ electrons ? D. Trapped and Solvated Electrons.-An arbitrary distinction between trapped and solvated electrons could be whether or not the potential well existed before the arrival of the electron at the site. This seems to be more meaningful than simply distinguishing solids from liquids or orderly from randomly polarised centres. On this basis Fcentres are clearly trapped electrons. Crystalline solids in general because of the orderliness and immobility particularly at temperatures much below their melting points may only ‘trap’ electrons and then only at crystal defects of which F-centres are a type. Thus Ghormley and Stewarta3 did not observe any colour due to e-aq in irradiated clear ice at 77”~.As Platz- man3’ has pointed out the relaxation time of ice at low temperature could be several years and hence the electron moves rather freely. The structure of ice apparently does not contain potential wells capable of trapping electrons except when the crystal is distorted or defective nor can orientational polarisation occur. Liquids at room temperature cannot normally give rise to stable traps either. Since in water for instance the dipolar relaxation time is about sec. a potmtial trap formed by the random orientation of neighbouring molecules cannot exist much longer than sec. in the absence of a stabilising influence 88 See C. J. Hochanadel N.A.S.-N.R.C. Reports 1953 No. 305 pp. 434. Walker such as the arrival of an electron. It might be pertinent to ask ife-aq is a solvated electron and formed entirely by self-induced dipolar orientations or if the electron is trapped in an existing polarised site resulting from random orienta- tions.A thermalised electron in water at room temperature will pass a molecule in a period too short for orientational polarisation and may therefore only cause self-induced electronic polarisations. Orientational polarisations may then follow in fluid media; but at 7 7 " ~ rotations will be extremely sluggish if possible at all. Despite this however the absorption spectra of electrons stabilised in low- temperature glasses are virtually identical with spectra obtained in the liquid phase at room temperature. Since water does not readily form a glass com- parisons in pure water are not currently possible but two very closely related systems namely a concentrated aqueous solution and methanol are compared in Figure 8.Curve (c) shows the electron absorption band in pulse-irradiated 15~-sodium hydroxide in water at 2 5 " ~ ~ ~ whereas curve ( d ) was obtained by 7-irradiation of 6~-sodium hydroxide aqueous glass at - 1 9 6 " ~ . ~ ~ These curves are identical except for the shift of the whole spectrum owing to the temperature change. Similarly curve (a) shows the electron-absorption band at 6500 A in pulse-irradiated liquid methanol at 2 5 O g o whereas curve (b) was obtained by y-irradiation of methanol glass at -196".90 Again the spectra are identical but for a normal temperature shift. (The value of GE reported was three times larger for the low-temperature irradiation. This could possibly result from an altera- tion of the extinction coefficient but is more probably a yield effect.) To have identical solvent orientations and distribution of orientations in fluid and rigid phases suggests that the electrons are not in fact polarons but are trapped in existing potential wells resulting from the mean instantaneous dis- tribution of molecular orientations this distribution being frozen-in when glassy solids are made.When cooling produces crystallisation on the other hand an ordered pattern of molecular orientations results which seldom permits the occurrence of electron stabilisation except at defect sites. Thus electron yields in glasses are invariably larger than in the corresponding crystalline phase.g1 The mobility of e-aq and indeed electrons in other polar media could con- sequently arise from jumping or tunnelling between existing potential traps.This may account for the facts that the rate of reaction of eaq with many solutes is governed more by the solvation energy of the activated complex than that of the electron53 and that electrons solvated in water ethanol and methanol react with H+ at similar rates.66b On this basis too the absorption spectrum would correspond to a charge-transfer-to-solvent phenomenon. Recently a linear relationship has been noted between the energy of Amax for the charge-transfer spectra of iodide ions and that for electrons solvated in the various media corroborating the prediction made by Platman and F r a n ~ k . ~ ~ F. S. Dainton and N. Gopinathin unpublished work. F. S. Dainton J. P. Kenne T. J. Kemp (3. A.Salmon and J. Teply Proc. Ckcm. SOC. 1964,265. 91 (a) F. S. Dainton and F. T. Jones Trans. Furuday Soc. 1965 61 1681; (b) T. Henriksen Radiation Res. 1964 23 63; (c) R. Livingston and A. J. Weinberger J. Chem. Phys. 1960 33 449. 105 Quarterly Reviews M A ) 2900 3500 5 3 0 0 6 5 0 0 I I 1 I I -J 0 40 3 0 20 10' 01 10-3 v (cm:' FIG. 8 . Comparisons of absorption spectra of solvated and trapped electrons. (a) Methanol at 25°,90 (6) Methanolat - 196°,Q0 (c) I ~ M - N ~ O H in water at 25°,63 (d)6~-NaOHglassaf - 196°,88 6 Conclusions Whereas the chemical distinction of hydrated electrons and hydrogen atoms is rather esoteric they are readily distinguished spectroscopically in both the optical and microwave regions. eaq has been shown to have a broad intense absorption spectrum throughout the visible region whereas hydrogen atoms absorb only in the far vacuum ultraviolet at ordinary temperatures.Also because of the weak interactions with protons forming the walls of the cavity shown by stable solvated electrons one would expect similar very narrow splittings of the e.s.r. absorption in the case of e-aq in marked contrast to the enormous hyperfine splitting shown by hydrogen atoms. This Review is concluded with a summary of the principal structural kinetic and thermodynamic parameters pertaining to eaq. and the knowledge of many of the fundamental properties of e-aq is summarised in Table 3. Combining the value quoted for the diffusion coefficient of e-aq with the rate constant for reaction (17) k, = 5 x one calculates an effective reaction diameter for eaq of ca.5 A which agrees well with the calculated charge distribu- tion of the 1s ground state corroborating the Jortner model. The thermodynamic results were calculated by Ba~endale~~ with subsequent small corrections applied for the more recent value of k,.25 For equilibrium (6) one can calculate the e-aq + H20 $ H + OH- (6) standard free-energy change Go = + 8-4 kcal. mole-l from the values of k = 16 1. mole-l sec.-l and k- = 2.2 x lo7 1. mole-l sec.-l. Combining this with AGO for the reactions H - &H and H+aq + OH,q - H20 one obtains for the reaction e-aq + H+aq -+ +H2 AGO = -61-8 kcal. mole-l. On the conventional 106 Absorption maximum (Amax) Estimated heat of solution ( AHs) Calculated first excitation energy Molar extinction coefficient at Amax Oscillator strength obs.calc. Calculated cavity size Calculated mean radius of charge (Jortner model) distribution In ground state (1s) In first excited state (2P) Diffusion coefficient Ge-aq in neutral water Lifetime with respect to (e-tq + H,O - H + OH-) Lifetime in neutral water (H,O+ ca. 10-7~) Standard electrode potential PKa for H + e-aq + H+aq TABLE 3 7200 A (1.72 ev) 2 1-72 ev 1.35 ev 15,800 1. mole-l sec.-l ca. 0.8 1.1 ca. 0 2.5-3.0 A 4.9 A 4.5 x 2.5 + 0.1 t+ 2 800 psec. (+ 15 %) cm.2 sec.-' ti = 300 psec. < 2.67 v < 9-7 Walker Ref. 50 a 55 51 43 55 55 55 55 b 57 25 a a a Baxondale's resu!ts (ref. 56) adjusted for the recent value of k = 16 1. mole-' sec.-l b Personal communication quoted in ref. 66b. scale therefore the hydrated electron has a standard electrode potential of 2-67 v. Baxendale makes the comparison with the value of 2.1 v for the hydrogen atom from which it is evident that e-aq is the stronger reducing agent.He suggests in addition that since the convention of quoting standard potentials is based on the arbitrary assumption that AGO = 0 for the standard hydrogen electrode [reaction (24)] it is now possible to adopt an absolute scale which would thus H+aq + e- -+ (24) require all potentials to be 2.67 v more negative. However this suggestion neglects differences in free energy between the electron of reaction (24) which is an electron (free) on the surface of an electrode and e-aq for which the above result was evaluated. These will differ by at least the solvation energy of e-aq (1.72 v). The solvation energy (AH,) of e-,q was evaluated by Baxendale from the cycle shown below a value of -2260.5 kcal.mole-l being used for the abso- e-aq + H+aq -+ (tHz)aq AGs \1 3. 2 107 Quarterly Reviews Baxendale compares this with the free energy of hydration of halide ions and notes that for I- ( AG = 57 kcal. mole-l) which is closest to eaq the contribu- tion made by the entropy of hydration can be neglected. Thus the enthalpy and free energy of hydration are equated so that a Hs - -339.7 kcal. mole-l or 1-72 ev. Since the hydration energy is a function of the ionic radius this small hydration energy implies a distribution of charge of e-89 certainly greater than the 2.2A radius of the iodide iohg2 a further endorsement of Jortner’s calculations. AH calculated in this way must however be regarded as approximate since its evaluation requires an absolute value for a thermodynamic function of an individual ion for which there is no strict thermodynamic baskg2 It should also be noted that since k = 16 1. mole-1 sec.-l is an upper limit the values quoted for the standard potential and pKa are maxima and the value of AH is a minimum. If the value for pKa falls much further (to ca. 8) any studies of the lifetime of e-aq with respect to its reaction with water (reaction 6) will be impossible since for any pH reactions (- 6) or (20) will predominate. The hydrogen atom is evidently a stronger acid than is water and consequently e a q is a weaker base than OH-. It is not surprising then that e-aq abstracts a proton from H,O very slowly i.e. that k is small. 92 K. B. Harvey and G. B. Porter ‘Introduction to Physical Inorganic Chemistry’ Addison- Wesley 19 63. 108

ISSN:0009-2681    

DOI:10.1039/QR9672100079

出版商:RSC    

年代:1967

数据来源: RSC

摘要:

Molecular Polyhedra of High Co-ordination Number By E. L. Muetterties and C. M. Wright PONT DE NEMOURS AND COMPANY WILMINGTON DELAWARE 1 9 8 9 8 CENTRAL RESEARCH DEPARTMENT EXPERIMENTAL STATION. E . I . DU U . S . A . A. Introduction This Review is directed to a characterisation of discrete polyhedra in which an atom is within bonding distance of seven or more other atoms. Specifically seven- eight- nine- ten- eleven- and twelve-co-ordinate complexes are described with primary emphasis on structure. Not long ago this area of co-ordination chemistry might have been described as obscure and certainly of limited scope. Advances in synthesis and the classic structural investigations of Hoard and his students1-' have sharply altered the picture in recent years. It is apparent now that the field of higher-co-ordination polyhedra will be quite large in scope touch- ing elements as light as scandium and titanium.Some of the higher-co-ordinate species are chemically and theoretically beguiling; an especially noteworthy one is the nonahydridorhenium d i a n i ~ n ~ ~ ReH:-. From the structural side the general question of geometry particularly relating to the solution state is a subtle difficult yet intriguing problem. Polyhedral isomerisations should be quite easy in these large polyhedra and this complicates definition of the solution or gaseous state. In brief molecular polyhedra of high co-ordination number have scope interesting chemistry and significant research challenge to those interested in synthesis structural and theoretical chemistry. 1 Ligand Classification.-Aside from the obvious classification of the large co- ordination polyhedra into hepta- octa- nona- deca- and dodeca-co-ordinate structures the general class can be differentiated by the nature of the ligands.Within this context there are three major groups (1) complexes with unidentate ligands (2) complexes with multidentate Iigands (3) structures in which there is extensive multicentre bonding generating polyhedra or clusters in which the high- co-ordinate atom (atoms) is a part of a large molecular overstructure. The first class is severely limited in size because the co-ordinating atoms must be small J. L. Hoard and H. H. Nordsieck J . Amer. Chem. SOC. 1939 61 2853. J. L. Hoard J. Amer. Chem. SOC. 1939 61 1252. (a) J. L. Hoard B. Lee and M. D. Lind J . Amer. Chern. SOC. 1965 87 1612; (b) J.M. D. Lind B. Lee and J. L. Hoard J. Amer. Chem. SOC. 1965 87 1611. J. L. Hoard and J. V. Silverton Inorg. Chem. 1963 2,235. ti M. J. Hamor T. A. Hamor M. D. Lind and J. L. Hoard Inorg. Chem. 1964,3 34. S . Richards B. Pedersen J. V. Silverton and J. L. Hoard Inorg. Chem. 1964,3,27. K. Knox and A. P. Ginsberg Inorg. Chem. 1964,3 555. S . C. Abrahams A. P. Ginsberg and K. Knox Inorg. Chem. 1964,3 559. Hoard personal communication. L. 109 Quarterly Re views and electronegative if a relatively stable molecular complex is to result. Primarily this is the ligand chemistry of hydride ion (TcH,2-),1O fluoride ion (TaFs3-),l1 cyanide ion [W(CN)g4-],12 and oxygen complexes INd(OH2),3+].13 There are many ionic lattices of the lanthanide and actinide halides in which halogen atoms bridge metal atoms to yield high-co-ordination spheres about the metal atoms but no simple halogeno-complex of seven- or eight-co-ordination has been described with the exception of those of the fluorides.Class (2) the metal chelates is by far the largest group within high-co-ordina- tion polyhedra. The most common donor atoms for the multidentate ligands are the electronegative oxygen and nitrogen atoms; in fact these are almost the sole donor atoms for higher co-ordinate complexes of d10 metal ions. Ligands of particular note are derived from /3-diketones oxalic acid tropolone NNN’N’- ethylenediaminetetra-acetic acid and nitrilotriacetic acid. Phosphorus and arsenic are excellent donor atoms for the transition elements but not the main- Group elements. A large number14 of hexta- and octa-co-ordinate complexes based on the diarsine ligand (1) and transition metals have been reported.Sulphur should also be a good donor atom in high-co-ordinate chelate structures and some examples have been described e.g. tetrakis(thiatropono)thorium- As a class the chelates dominate the area of higher co-ordination polyhedra in scope numbers and in kinetic and thermodynamic stability. Bidentate ligands can be distinguished by their rigidity coplanarity of ligand atoms or lack of coplanarity and also the ‘bite’ or the preferred separation of the two donor atoms. The more compact the ligand and the smaller the ‘bite’ the more effective is the ligand in generating high-co-ordination structures. The extreme and almost trivial example here is the peroxide ion which has an oxygen-oxygen separation of about 1.41 A in peroxy-complexes.* Another compact ligand is the (Iv) (2).15 * Cotton and BergmadB suggest the structural principle that ‘a polyatomic ligand in which two chemically equivalent atoms are hdd much closer together than such a pair of atoms would be if independent of each other has a tendency to interact through both of the equivalent atoms in such a way that the mean positions of the pairs of atoms lie roughly at the vertices of one of the usual co-ordination polyhedra’.An example is K,Cr(CN)3(Oz)z which has near-trigonal bipyramidal geometry about the chromium.17 lo A. P. Ginsberg Inorg. Chem. 1964 3 567. l1 J. L. Hoard W. J. Martin M. E. Smith and J. F. Whitney J. Amer. Chem. SOC. 1954,76 3820. la 0. 0. Collenberg Z. anorg. Chem. 1922 121 281 298.l3 L. Helmholz J . Amer. Chem. Soc. 1939 61 1544. l4 R. J. H. Clark D. L. Kepert R. S. Nyholm and J. Lewis Nature 1963 199 559. l5 E. L. Muetterties H. Roesky and C. M. Wright J. Amer. Chem. SOC. 1966 87 4856. l6 F. A. Cotton and J. G. Bergman J. Amer. Chem. SOC. 1964,86,2941; J. G. Bergman and F . A. Cotton Inorg. Chem. 1966 5 1208. l7 R. Stomberg Arkiv Kemi 1964 23 401. 110 Muetterties and Wright coplanar nitrate group which has a short oxygen-oxygen separation ca. 2.1 A and forms octa- and dodeca-co-ordinate structures with ease but the stability of the complexes is low; water instantaneously destroys the nitrate co-ordination polyhedron. Intermediate in effectiveness for high-co-ordination structures are the coplanar and compact ions (3) and (4) derived from oxalic acid and tropolone.Chelates derived from these ligands particularly the tropolone ion,18 are much more stable than nitrate complexes. The oxygen-oxygen separation is 2.58 A in bis(tropolono)copper19 and about 2.56 A in oxalato-chelates.20 The ions derived from p-diketone have greater flexibility as far as adjusting to preferred donor atom separations for a given metal atom. However there is a more significant intraligand repulsion factor in the p-diketone chelates than in those obtained from the more compact coplanar ligands such as tropolone or nitrate ion. The more complex multidentate ligands for maximum effectiveness should bridge vertices of preferred co-ordination polyhedra. For example the nitrilo- triacetato-ion easily spans four of the vertices of the octaco-ordinate dodeca- hedral (see Section C.2a) structure.The critical steric point in designing multi- dentate ligands is non-bonding repulsions within the ligand in the conformation required for chelation. Metal clusters or polyhedra are defined as discrete molecules or ions in which the metal atoms describe a polyhedron e.g. octahedron in Mo6C11,2- or a fragment of a polyhedron e.g. triangle in Re,C1123-. Co-ordination numbers of seven (Re,C1123-) eight (Pt&ll,) and nine ~a6Cl12(OH2>~+] have been found for individual metal atoms in these cluster species. Metal clusteis or polyhedra are properly treated as a distinct structural class since these species are explicable only on consideration of the structure as a whole i.e. molecular orbital theory and not in terms of individual atom co-ordination spheres.In the latter context valence-bond representations generally require invoking bent bonds. For more detailed information on metal clusters the reader is directed to a recent review by Schafter.21 Certain classes of ligand will be excluded from this discussion because there is some ambiguity regarding their co-ordinacy and because their chemistry is so extensive. In particular the huge areas of cyclopentadienyl-metal and olefin- metal complexes in which co-ordination numbers of three and two respectively might be invoked for the organic ligand will not be discussed. [For example l8 E. L. Muetterties J. Amer. Chem. SOC. 1966 88 305. l9 W. M. Macintyre J. M. Robertson and R. J. Zahrobsky Proc. Roy. SOC. 1966 A 289 161. 2o G. L. Glenn J. V. Silverton and J. L. Hoard Inorg.Chem. 1963 2 250. 21 H. Schafer Angew. Chem. 1964,76 833. 111 Quarterly Reviews (T-C~H,),T~H,~~ may be considered as a nonaco-ordinate tantalum structure. J Similarly aromatic hydrocarbon and ~-allyl complexes are excluded. Non- bonding electron pairs which are by inference stereochemically active will be considered quasi-ligands. No attempt will be made to treat ionic or metallic structures in detail however reference will be made to structural types within these two classes. Wells’s treati~e~~a on structural inorganic chemistry provides a more exhaustive assay of ionic and metallic lattices. 2 Metal-Ion Requirements.-Unfortunately present theories of chemical bond- ing are not sufficiently advanced to treat definitively molecules its large as hepta (or higher)-co-ordinate compounds.An attempt even to rationalise the existing knowledge of high-co-ordination structures must be mechanistic and rather empirical. Molecular orbital theory in even its extended forms is quite incapable today of predicting geometries for higher-co-ordination polyhedra. With valence bond approximations the best that can be done is qualitatively to assay orbital over-laps from purely symmetry arguments. For example in octaco-ordination the cube is not an attractive co-ordination polyhedron since only seven hybrid orbitals of proper symmetry can be generated unlessforbitals are used and this is consistent with observation. The stability of the cube versus say the square antiprism is however one of the few relatively dear distinctions that can be made on purely symmetry grounds for idealised geometries in higher-co-ordinate structures.Ground-state geometries are most simply rationalised by using the non-bonding repulsion model first introduced by Sidgwick and PoweU,=b later extended by Gillespie and Nyholm.24 There have been a number of attempts principally by Hoard and Silverton to lend some quantitative aspects to this simple approximation and mention will be made of the results of such calculations for each of the higher-co-ordinate systems. For ionic structures (see ‘Pauling’s rules’25) there is a rather good correlation between cation-anion radius-ratio and the co-ordination number of the metal ion. Further the co-ordination polyhedra in these instances are regular or highly symmetrical polyhedra. For example in AB compositions the relations are as shown in Table 1.However there are exceptions to this generalisation particularly among simple ionic structures where packing of twelve atoms of like charge is not possible and octaco-ordination is the largest co-ordination polyhedron for ionic lattices of the AB type. Co-ordination numbers of nine or more are however found in AB and AB salts or polymeric lattices and in intermetallic compounds. Cation-anion ratio is also an important consideration in molecular polyhedra but by itself it is of no predictive value with respect to idealised geometries. 22 M. L. Green J. A. McCleverty L. Pratt and G. Wilkinson J. Chern. SOC. 1961 4854. 23a A. F. Wells ‘Structural Inorganic Chemistry’ 3rd edn. Oxford University Press London 1961. 23b N. v. Sidgwick and H. M. Powell Proc. Roy. SOC.1940 A 176 153. 24 R. J. Gillespie and R. S. Nyholrn Quart. Rev. 1957 11 339. 25 L. Pallling ‘Nature of the Chemical Bond’ 3rd edn. Cornell Univ. Press Ithaca N.Y. 1960 p. 524; J . Amer. Chern. SOC. 1929,51 1010; 1933,55,1895. 112 Muetterties and Wright Table 1 Minimum radius-ratio 0.225 0.414 0.592 0.732 0.732 0.902 1-OOO Co-ordination number 4 6 7 9 12 12 a Co-ordination polyhedron Examples Tetrahedron LiCl Octahedron NaCl Capped octahedron Cube CSCl Tricapped trigonal prism Icosahedron Cube octahedron For molecular compounds or complexes metal-ion size is obviously an important factor in the formation of high-co-ordination structures. A table of ionic radii will quickly identify the important large ions (see Figure 1) such as I I 1 I I I I 1 1111 I I 1 I I 1 I 4 P fl f2 f3 f 4 f5 f 6 f7 re'* fI f2 f3 P f 5 P Metal ion electronic configuration Fig.1 Plot of metal-ion electronic configuration against ionic radius. the lanthanide and actinide ions. However size alone is insufficient since there is no authentic example of a discrete high-co-ordination complex of for exampIe the large alkali-metal ions. Hydration numbers of the large caesium and potassium ions appear to be less than six whereas the smaller but bivalent magnesium cation has a co-ordination number greater than six. Obviously metal-ion charge is also important. Thus more properly the facility with which an ion can form higher-co-ordination structures is some function of metal-ion size and charge. The filling of the d-electron level also has some impact on ease of formation of high-co-ordination structures.Metal ions of do or low d" electronic configura- tion generally form more stable high-co-ordinate structures than do metal ions of high dx cofiguration. The latter point is explicable on a basis of d-eIectron- 113 Quarterly Reviews ligand repulsions; or as alternatively characterised by Clark et all4 there are fewer low-energy orbitals available for bonding in the high dx configurations. Discrete high-co-ordination structures have been observed for metal ions of high dx electronic configurations although only with chelate ligands. Moreover these have been limited to high-spin species such as Mn2+ (d5) and F e s (d6). Heptaco-ordination has never been established for a low-spin species of for example a ds ion. Here the crystal-field stabilisation is significant and should raise the reorganisation energy for conversion of the octahedron into a heptaco- ordinate structure.It is interesting that studies of substitution reactions of low-spin d6 ions do not appear to be s'2 reactions; i.e. a heptaco-ordinate intermediate or transition state is not implied. In metal clusters where there is significant metal-metal bonding high co-ordination numbers have been observed for metal ions of high d" configuration and low spin e.g. Co,(CO), and Pt,Cl,,. Clark et all4 also consider formation of high-co-ordinate specifically octaco- ordinate structures in the early transition series to be more favourable than with post-transition-metal ions and suggest that this also reflects the greater avail- ability of low-energy s p and d orbitals for the low dx configurations.There is of course a significant metal-ion size decrease between do and d10 configurations in the same periodic series e.g. titanium(1v) has an ionic radius of 0.68 A com- pared with gallium(III) 0-62 A and germanium(Iv) 0.54 A. Hepta- and octa- co-ordinate structures are known for titanium but not for gallium or germanium. On the other hand tin(1v) (0-71 A) which is more comparable in size with titanium(Iv) does form hepta- and octa-co-ordinate species. In fact metal ions of do and d10 configuration of the same size and charge display comparable tendencies toward formation of high co-ordination structures although there exists some dichotomy in behaviour with respect to the nature of the ligand a not unexpected result. 3 Geometry.-For a given co-ordination number the differences between idealised geometries are surprisingly small for the high-co-ordination species and for real 'slightly distorted' structures discussion in terms of idealised geo- metries may become a matter of semantics.Establishment of ground-state structure for the crystalline state demands in many cases in the words of Pro- fessor Hoard 'the full power of three-dimensional X-ray data'. Much of the early work based on powder or two-dimensional results does not meet the necessary conditions. Of all the co-ordination numbers seven is in the most primitive state of structural characterisation. Octa- and nona-co-ordinate geometries are rather well established for the solid state. Few of the higher-co-ordinate structures are sufficiently stable to exist in the gaseous state. Gaseous species appear to be limited to IF, ReF, :XeF, and a number of tropolone and /&diketone derivatives e.g.(tropolono),Th and (acetylacetonato),Zr. Of these only the fluoro-derivatives are sufficiently simple for a meaningful structural analysis of the gaseous ground state by electron diffraction and even these present major problems because of the large heavy- 114 Muet terties and Wright atom scattering. No really rigorous study of the gaseous state has been made although preliminary electron-diffraction studies on IF and :XeF have been reported. It is unfortunate that so few high-co-ordinate species can be examined as gases for this is the only state that will provide an unperturbed environment and answer the question of the preferred ground-state structure. Because there is little difference between idealised geometries for any given co- ordination number there is no assurance that the same geometry for a given species will prevail in all the physical states.The energy differences between isomers is small; smaller than or comparable with packing forces in the solid state solvation energies in the solution state and association energies in the liquid state. The latter forces may in some cases stabilise a geometry that is not the ground-state geometry for the gaseous state. As is invariably the case solution- or liquid-state geometry the state of most general interest to chemists is the most difficult to define. Because of the com- plexity of the species there is no spectroscopic or diffraction technique that will establish geometry for this physical state with a high degree of rigour.Further ionisation ligand dissociation and solvation phenomena must be considered in the solution process of any high-co-ordinate species. For example NbF,2- is well established for the crystalline state. However dissolution of an NbF,2- salt in an aqueous medium even those of high fluoride and hydrogen ion con- centration gives a solution state in which NbF,- predominates and spectro- scopic studies have failed to detect the presence of the NbF,2- ion in s o l ~ t i o n . ~ ~ J ~ Dissolution of chelates may be quite complex. Because the reorganisation energy in going from seven- to eight- or eight- to nine- or nine- to ten-co-ordinate structures is very small there is a distinct possibility that an octaco-ordinate chelate for example may solvate to a nona- or deca-co-ordinate species (e.g.see Plate 1). This is especially true for the larger lanthanide tervalent ions and thorium(1v) complexes. Alternatively dissociation may prevail. A tris-chelate of the type (chel),MX may tend to ionise in polar media to give either (chel),Mf or (chel),M(solvent)+ ions. This type of process is known in titanium chemistry.28 Another type of ligand dissociation partial or complete dissociation of a chelate ligand may be common to chelates and especially to anionic complexes. Partial dissociation may or may not be followed by solvation Species like (5) and (6) may be important intermediates or transition states in chelate exchange reactions of the type M' = H+ or a metal ion 26 0. L. Keller jun. Znorg. Chem. 1963 2 783. 28 M. Cox J.Lewis and R. S . Nyholm J. Chem. SOC. 1964 61 13. K. J. Packer and E. L. Muetterties J . Amer. Chem. SOC. 1963 85 3035. 115 Quarterly Reviews Polar and nucleophilic solvents should favour processes like (1) and (2). Alterna- tively ligand dissociation may be complete and may or may not be followed by solvation Intermediates (7) and (8) could also be important in chelate exchange. Complete dissociation of the chelating ligand unaccompanied by solvation is observed with the octaco-ordinate bisdiarsine complex TiCl,,[diars],* on dissolution in benzene to give a hexaco-ordinate titanium complex TiCl,,[diars]. Characterisation of solution-state geometries and the dynamic processes of solution is the major research challenge in high-co-ordination structures. 4 Stereochemistry.-There are insufficient data to comment meaningfully about stereochemistry in the higher-co-ordinate structures.Characterisation of the more complex mixed-ligand species is necessary before any correlations can be made. There are however qualitative predictions to be derived from some of the mechanistic approaches to the question of bonding and structure. For most of the idealised geometries in any given co-ordination number the differences in geometry are comparable with the distortion imposed by vibra- tionally excited states.29 Thus for the long-term observation or description (> 10-l2 sec.) the high-co-ordination structures must be considered as potential stereochemically non-rigid species. Optical activity which has its origin in asym- metry at the metal atom centre may be very short-lived in solution for the higher-co-ordination species (see particularly Section C for a discussion of this stereochemical aspect) and geometrical isomerisation is also a potentially low- energy process in these structures.B. Heptaco-ordination Heptaco-ordinate structures once presumed to be limited to heavy-metal ions are now known for metal ions ranging from titanium to uranium with electronic configurations of doy P dlO andf” and even for boron in a complex cage struc- ture. The geometry problem for heptaco-ordinate species is far from resolved even for the solid state. Generally discussions of heptaco-ordinate geometry have been based on three idealised? structures which are the pentagonal bi- pyramid (D,J capped trigonal prism (C,,) and capped octahedron(C3y). These are presented in conventional perspectives in Plate 2.None of these has been established for a non-chelate heptaco-ordinate species with the precision of modern three-dimensional X-ray analysis and nothing less will suffice because of the subtle differences between these idealised models. There is a fourth basic * diars is ligand (1) on p. 109. t The term idealised is used in this Review loosely to denote limiting symmetry forms and does necessarily imply an energetically favoured geometry. 29 E. L. Muetterties Inorg. Chem. 1965.4 769. 116 Muet terties and Wright geometry the tetragonal base-trigonal base (Cs) rigorously established for complexes metal clusters and polymeric oxides. There are two idealised forms and one of these is depicted in Plate 2. (The second form has the same symmetry but differs in the relative orientation of the tetragonal and trigonal bases.There is an infinite number of forms intermediate in orientation of the base8 between the two boundary models.) It should be noted that this geom- etry for actual molecules is a refkction of constraints placed on the complex by the ligand the cluster or the solid lattice. This geometry has never been observed in a heptaco-ordinate species with unidentate ligands. The tetragonal base-trigonal base model is only very slightly distorted from the C, capped tri- gonal prism. In fact none of these seven-co-ordinate geometries differs signifi- cantly from another as can be seen from the alternative perspectives in Plate 3. Conformational interconversions require relatively slight bending modes and lifetimes of ground-state geometries may be quite short.29 Furthermore the energy difference between these various idealised states should be small with respect to intermolecular forces generated by ordering or solvation phenomena in the solid liquid or solution states.There have been several attempts*O,al to evaluate the relative stabilities of the idealised heptaco-ordinate MX models simply by considering repulsive forces generated from interactions of like ligand atoms. Considering the repulsive force as an inverse power n of distance the pentagonal bipyramid appears most stable for small values of n the C, model for intermediate values and the C, model for large values to the limit of the hard-sphere model. The energy differences are however very small; and because attractive forces are ignored and because calculations are limited to spherical models i.e.equivalent bond distances within each model the analyses have no predictive value. Another treatment32 of heptaco-ordinate geometries similar to the above but based on a different mathematical approach not limited to preconceived geometries con- firmed the earlier comparisons of the Dbk CZv and C, structures. However the force analyses minimised at a model not previously considered (or so far observed in actual systems). This model has only visualised in the spherical representation (9). *- - - - - - e C2 symmetry and is most easily (9 ) Configurationally the C model is easily generated from the other idealised heptaco-ordinate models by minor bending modes. It will be interesting to see 98 R. J. Gillespie Canad.J. Chem. 1960 38 818. 31 D. Britton Canud. J . Chern. 1963,41 1632. 3a T. A. Claxton and G. C. Benson Canad. J. Chem. 1966,44 157. 117 QuarterZy Reviews whether any actual MX species possess this geometry. Professor J. L. Hoard (personal communication) has noted that the iron(m) derivative of 1 ,Zdiamino- cyclohexane-NN’-tetra-acetic acid (Figure 7 vide infra) approximates the geom- etry of the C model. The most significant result of these analyses of repulsive forces in MX species is that the energy differences between idealised models are probably quite small. None of these analyses is applicable to heptaco-ordinate species in which the ligands are not identical. For such mixed species it is intuitively reasonable that energy minimisation is most favourable in D5h for MX,Y species in C, or C, for MX,Y and in the 4-3 trigonal base-square base for MX4Y3 particularly if X and Y differ significantly in steric or electronic properties.1 Pentagonal Bipyramid.-Pentagonal bipyramidal geometry is reported for the U0,F,3- (ref. 33) ion in the crystalline lattice of the potassium salt on the basis of a two-dimensional X-ray analysis. The U-0 bonds are axial with a 1.76 A separation and the equatorial U-F bond distances are 2.24 A. There is an ordered (tetragonal) and a disordered (cubic) form of K,UF,.= The disordered form appears to be isostructural with K3U0,F and the U-F bond distances are 2.26 A. Powder diffraction data35 suggest that in P-UF each uranium atom is within bonding distance of seven fluorine atoms; the positions of the fluorine atoms were not established.It would be extremely valuable to have precision neutron and X-ray diffraction data for a UF:- salt to fix geometry and bond parameters (will the axial bond distances be shorter than the equatorial dis- tances?). Puckering of the pentagonal ring may be observed for solid-state pentagonal bipyramidal molecules or ions. In the solution state pseudorotation analogous to that characterised for cyclopentane may be a low-energy process for the bonding atoms in the pentagonal ring (10). This has been postulated for iodine heptafluoride.% Hampson and P a ~ l i n g ~ ~ have characterised the ZrF,& ion in the ammonium and potassium salts as C, (capped octahedron) ; however their two-dimensional analysis is not definitive. Zachariasen has argued that 2rF:- is in fact a pen- tagonal bipyramid.= The basic crystallographic problem is complicated by thermal disorder in these salts.Hoppe and Rodder38 report TbF,% (Cs3TbF,) as probably isostructural with ZrF,3-. Other possible heptafluoro-anions include 33 W. H. Zachariasen Acta Cryst. 1954 7 783. 34 W. H. Zachariasen Acta Cryst. 1954 7 792. a5 W. H. Zachariasen Acta Cryst. 1949 2 296. 36 R. E. Lavilla and S. H. Bauer J . Chem. Phys. 1960 33 182. 37 G. C. Hampson and L. Pauling J . Amer. Chem. SOC. 1938 60 2702. 38 (a) R. Hoppe and K. M. Rodder 2. anorg. Chem. 1961 313 154; (6) R. D. Peacock H. Selig and I. Sheft Proc. Chem. SOC. 1964 285; Inorg. Chem. 1967 in the press. 118 Muetterties and Wright XeF,-,38b CeF,* PrF?- NdF?- and DyF,3-?941 A vibrational analysis of iodine heptafluoride implicated Dbh symmetry for this volatile An electron-diffraction of gaseous IF suggests neaf symmetry but with a non-coplanar equatorial belt of five fluorine atoms; the average I-F distance is 1.825 A.Single-crystal X-ray data have also been presented for this molecule43 but are not sufficient to differentiate between the D5h and C, Nuclear magnetic resonance (n.m.r.) results for liquid IF and ReF have been used to argue for a ground-state geometry of very short life.45,46 A preliminary report of a two-dimensional X-ray analysis of K2Mo0,F4,0H indicated that the molybdenum atom is heptaco-ordinate with three fluorine atoms and a peroxide group coplanar with the molybdenum and the remaining fluorine and oxygen (water) atoms on the axis normal to the plane.47 A high- pressure form of U03 has puckered pentagonal bipyramidal geometry for the oxygen atoms about each uranium The collinear uranyl bonds are short ranging between 1.80 and 1.85 A.In Cs,(UO&,(S04) the uranyl group is co- ordinated to five sulphate oxygen atoms to give a near-pentagonal bi~yramid.**~ A near-pentagonal bipyramidal co-ordination sphere is present for three of the cobalt atoms in the metal polyhedron of CO~(CO)~~ vide i1zfra.4~ In Zr(OH),(Cr0,)5(H,0)2 there are infinite chains of zirconium atoms bonded to hydroxide and chromate oxygens?Oa The zirconium atoms are co-ordinated to seven oxygen atoms in the shape of an almost regular pentagonal bipyramid. In K,Cr(O,),(CN) there is near-pentagonal bipyramid geometry with two apical and one equatorial cyano-group (Figure 2).17 The co-ordination polyhe- dron of Cr(O&,(NH& is described as a distorted pentagonal bipyramid?Ob 2 Capped Trigonal Prism (C,,).-The potassium salts of NbF,2- and TaF:- possess nearly identical monoclinic cell units.Analysis of two-dimensional X-ray data2 pointed to a slightly distorted C, capped trigonal prism model for NbF,2- with an average Nb-F separation of 1.97 A. A similar analysis of Na,Zr,F, suggested a linking of trigonal prism polyhedra (ZrFt-) through the square faces with a bridging fluorine atom.5l The Zr-F separations fall in the aB R. Hoppe and K. M. Rodder Z. anorg. Chem. 1961,312,277. 4 1 L. Asprey Rare Earth Research 1961 58. 42 R. C. Lord M. A. Lynch jun. W. C. Schumb and E. J. Slowiski jun. J. Amer. Chem. SOC. 1950,72,522. 43 R. D. Burbank and F. N. Bensey jun. J. Chem.Phys. 1957 27 981. 4p J. Donohue Acta Cryst. 1965 18 1018. 45 E. L. Muetterties and K. J. Packer J. Amer. Chem. SOC. 1964 86 293. 46 R. J. Gillespie and J. W. Quail Cunnd. J. Chem. 1964 42 2671. 47 D. Grandjean and R. Weiss Compr. rend. 1965 261 448. (a) S. Siegel H. Hoekstra and E. Sherry Acta Cryst. 1966 20 295; (b) M. Ross and H. T. Evan jun. J. Znorg. Nuclear Chem. 1960 15 338. 49 P. Corradini J. Chem. Phys. 1959 31 1676. 5* (a) G. Lundgren Arkiv Kemi 1958,13,59; (b) E. H. McLaren and L. Helmholz. J. Phys. Chem. 1959 63 1279. 51 R. M. Herak S. S. Malcic and L. M. Manojlovic Acta Cryst. 1965 18 520. R. Hoppe and W. Liebe Z. anorg. Chem. 1961,313,221. 119 Qluarterly Reviews i C R Fig. 2 Structure of Cr(O,)&N)Ss- in K,Cr(O&,(CN) (ref. 17). range 2.00 A to 2.10 A. Salts of PuF7-5s and WF7-63 ions have been isolated but there are no definitive structural data.The K,PaF salt is not isostructural with K,NbF7 but is nonaco-ordinate. The NbF,O- structure has been recently con- firmed by three-dimensional neutron-diffraction analysis.64 Refined values of the Nb-F separations vary between 1.940 and 1.978 A. The slight distortion from C, symmetry and the exceptionally large vibrational distortions of two of the fluorine atoms are explicable in terms of the packing in the crystal. The monoclinic or B form of the rare-earth oxides (Sm,O,-Gd,Od contains MO units of three types two of which are essentially capped trigonal prisms and the third is octahedral with an additional very long M-O distance of 3.12 A.6595s The two heptaco-ordinate capped trigonal prisms differ only in dimensions with each one having one long Sm-0 distance (2.71 and 2.76 A).The hexagonal form ( A modification) of lanthanum cerium praseodymium neodymium and americium(II1) oxides apparently (X-ray and neutron powder diffraction data) has polyhedra of MO units for which the capped octahe- dron (C3J is a fair approximation of geometry for these ~ n i t s . ~ ' ~ ~ There are three metal-oxygen distances of 2.38 A three of 2-72 A and one of 2.45 A. Reportedly isostructural with A-La,O are La,O,S Ce,O$ and Pu202S59 with four oxygen and three sulphur atoms bonded to the metal atom. 3 Capped Octahedron (C3,).-This geometry has been suggesteds7 for ZrF,'- although challenged by Zachariasei~.~~ Williams and Hoard prescribed the cap- ped octahedron for the NbOFs3- ion from analysis of two-dimensional X-ray measurements.60~61 The Nb-0 bond was placed on the threefold axis although the data were not sufficient to establish this.Metal-ligand separations averaged 2.0 A but the Nb-0 bond should be significantly shorter than the Nb-F bond. 52 R. A. Penneman G. D. Sturgeon L. B. Asprey and F. H. Kruse J. Amer. Chem. SOC. 1965,87. 5803. 53 G. B. Hargreaves and R. D. Peacock J . Chenz. Sac. 1958 2170. 54 G. M. Brown and L. A. Walker Acta Cryst. 1966,20 220. 55 W. H. Zachariasen Acta Cryst. 1949 2 60. 56 D. J. Cromer J . Phys. Chem. 1957 61 753. 57 L. Pauling 2. Krist. 1929 69 415. 58 W. C. Koehler and E. 0. Wollan Acta Cryst. 1953 6 741. 59 0. J. Guentert and R. L. Mozzi Acta Cryst. 1958 11 746. 6o M. B. Williams and J. L. Hoard J . Amer. Chem. SOC. 1942 64 1139.61 J. L. Hoard and W. J. Martin J . Amer. G e m . SOC. 1941 63 11. 120 Muetterties and Wright Both K,ZrF and K3NbOF give closely similar X-ray diffraction patterns. A series of salts based on TaOF,& are isomorphous by diffraction criterion with the NbOFt- salts.g2 In the A form of the rareearth oxides the metal atom has a distorted C, arrangement of seven nearest oxygen a t o r n ~ ~ ~ ~ * (see Section B.2). 4 Tetragomi Base-Trigonal Base.-Two configurations of this polyhedron that possess a plane of symmetry are (1 1) and (12). These conformers are inter-related by a rather trivial twisting mode of the trigonal face and are both very similar to the capped trigonal prism (see Plate 3). Conformation (11) is found for the iron atoms in the cyclobutadiene complex (C6H5)&Fe(CO)3 (Figure 3.6' The structures of the parent cyclobutadiene complex H4C4Fe(C0)3,M,65 has not been established.Conformation (12) is reported66 for a cyclobutadienenickel- (III) chloride solvate with benzene (Figure 4). The monoclinic forms of Zr0267968 0 CI CH3 Fig. 3 Structure of (C,H,),C,Fe(CO) (ref. 63). Fig. 4 Structure of (CH,),C4NiC13 in the (CH$,C,NiCI3,C6H crystal (ref. 66). and HfO,Sg show configuration (12) for the oxygen co-ordination sphere about the metal atoms. In ZrO the metal-oxygen bond distances are 2-07 for the trigonal base and 2-21 A for the square base. A similar configuration is reported 62 A. E. Baker and H. M. Haendler Znorg. Chem. 1962 1 127. 63 R. P. Dodge and V. Schomaker Nafure 1960 186 798. 6p G. F. Emerson L. Watts and R. Pettit J. Amer.Chern. SOC. 1965 87 131. J. D. Fitzpatrick L. Watts G. F. Emerson and R. Pettit J . Amer. Chcrn. Soc. 1965 87 3254. 66 J. D. Dunitz H. C. Mez 0. S. Mills and H. M. M. Shearer Helv. Chim. Acta 1962 45 647. 67 J. D. McCullough and K. N. Trueblood Actu Crysf. 1959 12 507. 6* D. K. Smith and H. W. Newkirk Acru Cryst. 1@65,18,983. 69 J. Adam and M. D. Rogers Acfu Cryst. 1959 12 951. 121 Quarterly Reviews for ZrOS.'O There are four basal sulphur atoms at ca. 2-62 A and three oxygen atoms at 2.13 A. Metal atoms in metal polyhedra or clusters sometimes have co- ordination polyhedra similar to (12) vide infra. 5 Chelate Structures.-Certain metal derivatives of ethylenediaminetetra-acetic acid display a tendency to retain a molecule of solvent. Two of these analysed in the solid state by three-dimensional X-ray methods have been shown to have heptaco-ordinate metal atoms with a strongly bound water molecule.(The heptaco-ordinate ion is probably the dominant species in solution.) In the iron- (111) chelate6 the bonding atoms closely approximate a pentagonal bipyramid whereas in the manganese(@ derivative' the co-ordination polyhedron more closely approximates a capped trigonal prism (Figures 5 and 6). Heptaco-ordina- tion has also been rigorously established for a calcium salt of a hydrated iron(ru) derivative of 1,2-diaminocyclohexane-NN'-tetra-acetic acid. The idealised model is shown in Figure 7. Cohen and Hoard7I describe the co-ordination polyhedron as approximating a capped trigonal prism and alternatively* as a close approxi- mation of the C model (9). The calcium ion in this crystalline iron complex has OH2 Fig.5 Structure of the iron(nr) complex with ethylenediaminetetra-acetic acid in LiFe(OH,)- (ethylenediaminetetra-acetute),2H20 (ref. 6). Fig. 6 Structure of the heptuco-ordinate mangunese(I1) ion Mn(OH,)(ethylenedianrinetetra- in Mn3(H30),(ethylenediuminetetru-ucetate),,8H20. There are also hexaco-ordin- ate manganese ions in this crystal lattice (ref. 7). Fig. 7 Structure of the iron(rr1) complex with 1,2-diaminocyclohexune-NN'-tetru-acetic acid in Ca[Fe(OH~(chelate)],,9H2O (ref. 71). * Personal communication from Professor Hoard. 'O J. D. McCullough L. Brewer and L. A. Bromley Acta Cryst. 1948,1 287. 71 a. H. Cohen and J. L. Hoard J. Amer. Chem. SOC. 1964,86,2749; 1966,88 3228. 122 Muetterties and Wright a virtually perfect C, capped trigonal prismatic co-ordination polyhedron.Cohen and Hoard7 note that in heptaco-ordinate complexes based on ethylenediaminetetra-acetic acid either the pentagonal bipyramid or the capped trigonal prism is a favourable geometry but with the diaminocyclohexanetetra- acetic acid ligand the ring constraints favour the capped trigonal prism (or C,) geometry. In heptaco-ordination one of the most common chelate structures has the composition ( ~ h e l ) M X . ~ ~ s ~ ~ ~ ~ Even in those cases where the chelating ligand is symmetrical there is a large number of possible isomers for each of the four idealised geometries with the exception of the pentagonal bipyramid which has only three isomers. If the ‘bite’ of the chelating ligand is relatively rigid then matching of ‘bite’ with polyhedral edges will be a geometry-determining factor.On the other hand if the ‘bite’ is flexible and if X the unidentate ligand differs significantly from the chelating donor atoms in electronic or steric features energy minimisation will be most favourable in such models as the capped octahedron and the capped trigonal prism where the unidentate ligand can occupy a unique site. Using these models two capped octahedral isomers and three capped prisms to define the minimum set of reasonable isomeric structures in (chel),MX and their closely related isomers in the other geometries we have drawn Figure 8 to represent easy polyhedral isomerisation paths. The vertical columns represent geometries closely related in the sense that small distortions suffice to interconvert structures.6 Metal Polyhedra or Clusters.-Heptaco-ordination is quite often present in polynuclear metal carbonyls and their derivatives. For example analysis of two- and three-dimensional X-ray results for CO,(CO), established a tetra- hedron of cobalt atoms (Figure 9) with three basal cobalt atoms that are bonded to three other cobalt atoms two bridging CO groups and two terminal CO g r o ~ p s . ~ ~ ~ ~ The co-ordination polyhedron for the basal heptaco-ordinate cobalt atoms approximates to a pentagonal bipyramid. Presumably the isoelectronic ion F~CO,(CO),~- is isostructural with c0,(co),,.75 A diethylacetylenederiva- t i ~ e ~ ~ of Co,(CO),, Co,(CO),,(C,H,~C z CC,H,) has been precisely analysed with three-dimensional measurements (Figure 10) and the two basal cobalt atoms are heptaco-ordinate.The two other cobalt atoms can be considered as hexa- or heptaco-ordinate depending on whether the C = C interaction with these two cobalt atoms is taken as unidentate or bidentate. A 4 3 type metal- atom co-ordination sphere is found for the apical iron atom in S2Fe3(C0)z7 and Se,Fe,(C0),78a (Figure 11) which both have tetragonal pyramids of iron 72 G. T. Morgan and A. R. Bowen J. Chem. Soc. 1924,125 1252. 7s E. L. Muetterties and C. M. Wright J. Amer. Chem. SOC. 1964 86 5132. 74 C. H. Wei and L. F. Dahl J . Amer. Chem. SOC. 1966 88 1821. 76 P. Chini L. Colli and M. Peraldo Gazzetfu 1960 90 1005. 7g L. F. Dahl and D. L. Smith J. Amer. Chem. Soc. 1962 84,2450. 77 C. H. Wei and L. F. Dahl Znorg. Chem. 1965 4,493. 78 (a) L. F. Dahl and P. W. Sutton Inorg. Chem. 1963 2 1067; (b) U.Anders and W. A. G. Graham Chem. Comm. 1966,291. 123 Quarterly Reviews c3 Cl c2 c2 c2 c2v c2v c2 Fig. 8 Some of the possible geometrical isomers for the various heptaco-ordinate models in a (chel),MX structure (see text). Each column comprises structural isomers configurationally related by small distortions. oc ko/" -0 Fig. 9 Structure of Co,(CO),,. The three basal cobalt atoms are heptaco-ordinate. The slx cobalr-cobalt distances are identical within experimental error (refs. 49 74). Fig. 10 Structure of Co,(CO),,(CBH,C = CC,H,) (ref. 76). and chalcogenide atoms as the basic metal cluster. The heptaco-ordinate poly- hedron approximates that of (12). The exact geometry in Fe3(C0),2 has now been unequivocally e~tablished'~ as depicted in Figure 12 and two of the iron atoms are heptaco-ordinate ; the co-ordination polyhedron approximates the capped octahedron.The new MnFez(CO),2- anion7a is probably isostructural with Fe,(CO)12. Two of the iron atoms in [C,H,C,C,H,]FQ(CO)~ are heptaco-ordinate 124 MPrcttertim and Wright Q 0 0. c' \ /C/O c I / Fig. 11 Structures of S,Fe,(CO)B and Se,Fe,(CO)o. The structures differ only in the orientdon of the (CO) triangle over the Fe,S and Fe,Se base. In the former a CO group is directly over an iron atom and in the latter over a selenium atom (refs. 77 78a). 0 OC\- i /co Fig. 12 Structure of Fe,(C0),2 (ref. 74). (Figure 13); one of the acetylenic carbon atoms is equidistant from all three iron whereas the second acetylenic carbon atom is bonded to only two iron atoms. Heptaco-ordination is also observed in a related diphenylacetylene-iron derivative [ C6H,C2C,HJ,Fe,( CO)880 which exists in two structural modifications.In the violet metastable form two iron atoms are within bonding distance of two other iron atoms four carbon atoms and three terminal carbonyl groups (Figure 14). The unique iron atom is hexa- or octa-co-ordinate depending on whether the olefhic linkage is considered a mono- or bi-functional moiety. In Fe,(CO),,C (Figure 15) the five iron atoms describe a square pyramid and the basal iron atoms are heptaco-ordinate.81 Heptaco-ordination is found in the rhenium(r1r) halide clusters. The prototype is RqCllZS- which has been rigorously characterised as an equilateral triangle of chlorine atoms with rhenium atoms at the mid-edge points; the remaining nine chlorine atoms are terminally attached in groups of three to each rhenium atom (Figure 1 6).82,83 The co-ordination polyhedron about the rhenium atoms approximates to a pentagonal bipyramid.A similar arrangement is found in the 79 J. F. Blount L. F. Dahl C. Hoogzand and W. Hiibel .J. Amer. Chem. SOC. 1966,88,292. 8o R. P. Dodge and V. Schomaker J . Organornetallic Chem. 1965,3,279. E. H. Braye L. F. Dahl W. Hubel and D. L. Wampler J. Amer. Chem. SOC. 1962,84 4633. 82 J. A. Bertrand F. A. Cotton and W. A. Dollase Znorg. Chem. 1963,2 1166. 83 W. T. Robinson J. E. Ferguson and B. R. Penfold Proc. Chern. SOC. 1963 116. 125 Quarterly Reviews TH5 Fig. 13 Structure of [C6H,.C2.C6H,]Fe,(CO) (ref. 79). Fig. 14 Structure of [C,H,.C2.C,H,],Fe,(C0) in the violet metastable crystal form. See Figure 28 for the isomeric structure (ref.80). The terminal carbon atoms represent one of the carbon atoms of the phenyl substituents. 0 Fig. 15 Structure of Fe,(CO)15C (ref. 81). CI CI Cl Fig. 16 Structure of Re3C1,,3- in CF,R~,CI, (ref. 82). Fig. 17 Structure of Re3C1,[P(C2H5),C,H J9 (ref. 85). /I\\ CL 126 Muetterties and Wright parent trichloride and tribromide; Re3Xg groups are in double layers and are joined together by chlorine-atom bridges to create a q~asi-Re,X,~~ ~ t r u c t u r e . ~ ~ ~ A derivative of the trichloride Re3C19[P(C2H,),C6H5]3 is isostructural with Re,CI,,* (Figure 17),85 and there are related complexes of Re,CI and Re3Br with a number of donor molecule^.^^^^^ Additionally there are Re,CI,,Z- Re,Br,,,- and Re3Brll2- anions in which one or two terminal halogens are missing from the Re3X123- ~ t ~ ~ t ~ r e .~ ~ - ~ ~ ~ ~ ~ ~ ~ For a sulphide (and a selenide) derivative of rhenium carbonyl Re3(CO),(SC,H,), it has been suggested that the C,H,S(C,H,Se) group serves as the bridging ligand as chlorine is in Re3CI ,:-. 90a A heptaco-ordinate polyhedron that is significantly different from the four idealised models discussed above exists in a series of tetrameric copper@) iodide complexes with phosphines arsines e t ~ . ~ l [The analogous silver(1) halide com- plexes are also tetrameric and may be isostru~tural.~~] These have tetrahedra of copper atoms with iodine atoms above the faces and donor ligands at the vertices (Figure 18).92 Each copper atom is within bonding distance of three iodine atoms three copper atoms and one donor atom (of the donor ligar~d).~~ Six atoms of the co-ordination sphere are below the central copper atom (13).This geometry is simply a reflection of the entire structure rather than any As \ As Fig. 18 Structure of [(C,H,),AsCuI] (ref. 92). 84 (a) F. A. Cotton and J. T. Mague Inorg. Chem. 1964 3 1402; (b) F. A. Cotton and S. J. Lippard ibid. 1965 4 59; (c) B. R. Penfold and W. T. Robinson ibid. 1966 5 1758; (d) M. Elder and B. R. Penfold ibid. 1966 5 1763. 85 F. A. Cotton and J. T. Mague Inorg. Chem. 1964 3 1094. 86 F. A. Cotton S. J. Lippard and J. T. Mague Znorg. Chem. 1965 4 508. F. A. Cotton N. F. Curtis C. B. Harris B. F. G. Johnson S. J. Lippard J. T. Mague W. R. Robinson and J. S. Wood Science 1964 145 1305. J. E. Ferguson B. R. Penfold and W. T. Robinson Nature 1964,201 181 ; Inorg.Chem. 1966 5 1763. M. Elder and B. R. Penfold Nature 1965 205 276. * O (a) A. G. Osborne and F. G. A. Stone Chem. Comm. 1965 361 ; (6) E. W. Abel B. C. Crosse and G. V. Hutson Chem. and Ind. 1966 238. O1 F. G. Mann D. Purdie and A. F. Wells J. Chem. Suc. 1936 1503. BzA. F. Wells 2. Krist. 1936 94 447. 127 Quarterly Reviews indication of preferred stereochemistry for heptaco-ordinate copper. A closely related configuration is foundgS for one of the iron atoms in Roussin's Black Salt CsFe,S,(NO),,H,O. The apical iron atom (Figure 19) is within bonding Fig. 19 Structure of the iron-sulphur polyhedron in Roussin's Black Salt CsFe,S,(NO),,OH (ref. 93). distance of three irons three sulphurs and a nitrogen of a nitrosyl group. The interaction of the unique iron atom with the three other iron atoms may be accounted for with a single delocalised molecular orbital.There are two isomeric boron hydrides of the composition B,,H, which have heptaco-ordinate boron atom^.^*^^^^^ The centrosymmetric species may be described as two B, units from decaborane (B,,H,J sharing a pair of boron atoms with the latter atoms each within bonding distance of six other boron atoms and a bridge hydrogen atom (Figure 20).Q40 The iso-B,,H, s t r u ~ t u r e ~ ~ ~ is formally analogous in that two B, units from decaborane are joined by sharing two boron atoms but the points of fusion are different from that in the normal structure (Figure 21). The 'shared' boron atoms again are the heptaco-ordinate ones; one is within bonding distance of the five boron and two hydrogen atoms and the other is near seven boron atoms.There are a number of solid-state structures derived from the early transition elements in which metal-metal interactions occur and heptaco-ordination is found for the transition element. Niobium(1v) iodide is a polymeric chain with two cis-bridging Nb-I-N b bondsg5 The niobium-niobium separation is only 3.3 1 A indicating metal-metal bonding and establishing heptaco-ordinate G. Johansson and W. N. Lipscomb Acta Cryst. 1958,11 594. 9g (a) P. G. Simpson and W. N. Lipscomb J. Chem. Phys. 1963 39,26; ( b ) P. 0. Simpson K. Folting R. D. Dobrott and W. N. Lipscomb ibid. p. 2339; (c) P. G. Simpson K. Folting and W. N. Lipscomb J. Amer. Chem. SOC. 1963 85 1879. 95 L. F. Dahl and D. L. Wampler Acta Cryst. 1962 15 903. 128 Muetterties and Wright Fig.20 Structure of n-B,,H,,. The two heptaco-ordinate boron atoms are represented by circles. The black dots represent boron atoms each of which has a terminal hydrogen atom (not depicted) (ref. 94a). Fig. 21 Structure of iso-B,,H,,. The notations are those of Figure 20 (refs. 94b 94c). niobium atoms (Figure 22). Tantalum(1v) iodide is isomorphous95~gs and niobium(1v) chlorides7 is structurally analogous with niobium iodide. Other niobium compounds which have Nb-Nb interactions and are heptaco-ordinate include NbOC1,Q8 and NbO,,QQ and a formally related species is M O C I . ~ ~ ~ In the W,C1,3- ion the two tungsten atoms are within bonding distance and are bridged by three chlorine atoms.lo1 The remaining halogens are terminally bonded three to each tungsten atom. Fig 22 Structure of i L NbI (ref.95). 7 Quasi-heptaco-ordinate Polyhedra.-Zenon hexafluoride unlike all other known neutral metal hexafluorides does not possess 0 symmetry.lo4 The precise geometry is as yet uncharacterised but the distortion from Ok symmetry 96 L. F. Dahl and D. L. Wampler J. Amer. Chem. SOC. 1959 81 3150. 97 H. G. Schnering H. Wohrle and H . Schiifer see ref. 21 p. 842. 98 H. G. Schnering and H . Wohrle see ref. 21 p. 1841. 99 B. 0. Marinder Arkiv Kemi 1962, 19 435. loo H. G. Schnering and H. Wohrle Naturwiss. 1963 50 91. lol W. H. Watson jun. and J. Waser Acta Cryst. 1958.11 689. loe L. S. Bartell R. M. Gavin jun. and H. B. Thompson J. Chem. Phys. 1965 43 2547. Quarterly Reviews is presumed by some103 to reflect the presence of a non-bonding pair of electrons in a directed orbital.Accepting such a non-bonding electron pair as a quasi- ligand we can enlarge the definition of heptaco-ordination to encompass such molecules ions and complexes as XeF6,103 IF6- BrF, IF5,NC,H5 and IF,,0CHN(CH,),.104 It should be noted that in some isoelectronic species such as TeC1,2- the non-bonding electrons must be in an s orbital because regular octahedral geometry pre~ai1s.l~~ Thus structural data are required to define precisely this class of quasi-heptaco-ordinate structures. 8 Solution Data.-There is no definitive structural datum for heptaco-ordinate species in solution. Nuclear magnetic resonance results for a number of these species in solution have been equivocal in a structural sense because invariably there was spectroscopic equivalence of ligand atoms. This has been true for ReH,p(C,H,),] (‘H resonance of ReH atoms is a triplet due to HP coupling),los as well as ReF7,4 IF,- BrF6- and IF5 complexes (19F spectra).lo4 The spectroscopic ligand-atom equivalence in IF and ReF has been ascribed to fast intramolecular is0merism.4~,~~ A similar explanation might be suggested for ReH,p(C,H,)a,* but the remaining species are labile and fast intermolecular ligand exchange may provide the fast averaging process.More recently Professor L. S . Bartell (personal communication) has suggested for XeF, which is dis- torted from Oh symmetry and is a quasi-heptaco-ordinate structure with the non-bonding electron pairs in a directed orbital a rapid intramolecular re- arrangement formally analogous to the ammonia inversion. 7 The polyhedral rearrangement is estimated to be very fast possibly faster than that for ammonia.There are a number of unimolecular tris-chelates of the type (chel),MX that are most certainly heptaco-ordinate but the geometries are as yet undefined. Examples from the tropolone and substituted tropolone ions are (tropolono),- SnC6H5,73 (tropol0no),TiC1,~~~ and (y-isopropyltropolono),Nb015 which are all non-ionic and monomeric in organic solvents. Analogous acetylacetone derivates of Zr and HfN e.g. (acac),HfCI are monomeric and give no evidence of ionisation to the tris-chelate metal cation which is established for titanium (a~ac),Ti+.,~ However the hafnium and zirconium derivatives based on the bulky anion derived from dibenzoylmethane although non-ionic themselves react with ferric chloride to give the salt (bzbz),M+FeC1,-.28 Catechol forms an apparent hept aco-ordinate s tructurelo8 with t ant alum(v) methoxide [(cat),- TaOCH,I2-.Reactions of niobates and tantalates with catechol in aqueous media give species of the type (cat),Nb03- and [(cat)3TaOTa(cat),]*-.109 pyramidal geometry if the hydrogen atoms are all in the pentagonal plane. Io3 R. J. Gillespie ‘Noble Gas Compounds’ Chicago University Press Chicago 1963 p. 333. lo* E. L. Muetterties ‘Advances in The Chemistry of The Co-ordination Compounds’ The MacMillan Co. New York 1961 p. 509. lo5 E. E. Aynsley and A. C. Hazell Chem. and Ind. 1963,611. lo6 L. Malatesta M. Freni and V. Valenti Gazzetfa 1964 94 1278. lo7 E. L. Muetterties and C. M. Wright J . Amer. Chem. SOC. 1965 87 4706. lo* R. Gut H. Buser and E. Schmid Helv. Chim. Acta 1965 48 878.lo9 A. Rosenheim and E. Roehrich 2. anorg. Chem. 1932,204 342. * All five metal-hydrogen atoms would be environmentally equivalent in pentagonal bi- t See also Trans. Amer. Crysf. ASSOC. 1966,’ 2 149. 130 Muetterttes and Wright Halogen oxidation of halogenopentacarbonylmetallate anions X(CO),M- gives rise to a series of heptaco-ordinate anionic species:l1° The solid products are salts which function as typical 1 1 electrolytes in nitro- benzene. This oxidation route to heptaco-ordinate anions is general to XW(CO)5- and XMo(CO) where X and halogen oxidant is bromine or iodine. This pro- procedure has also been applied to a series of diarsine (0-phenylenebis-di- methylarsine) complexes:111,112 1 2 Mo(CO),(diars) - [Mo(CO),(diars),I]+I- Analogous triarsine [bis-(0-dimethylarsinophenyl)methylarsine] chemistry1lsU pc exists xs M(CO),(triars) - [M(CO),(triars)W+X- (M = Cr Mo W; X = Br I) A [M(CO),(triars)xl+X- -+ M(CO)2(triars)X Molecularity and ionicity were established by molecular-weight and conductivity studies as well as metathetical ion-exchange reactions.Similar chemistry pervades bis-phosphine chelate chemistry of molybdenum and tungsten c a r b o n y l ~ . ~ ~ ~ For example iodination or bromination of the 1,2-bisdiphenyIphosphinoethane derivative Mo(CO),(diphos) yields Mo(CO),(diphos)X which is diamagnetic and a non-electrolyte. 1 ,2-Dithian complexes of molybdenum and tungsten carbonyls are oxidised by iodine to give non-ionic unimolecular species of the type [W(CO),(dith)I,] that are probably heptaco-ordinate with bridging dithian groups.ll5 M.C. Ganorkar and M. H. B. Stiddard J . Chem. SOC. 1965 3494. ll1 W. J. Kirkham A. G. Osborne R. S. Nyholm and M. H. B. Stiddard J. Chem. SOC. 1965 550. llx H. L. Nigam R. S. Nyholm and M. H. B. Stiddard J. Chem. Soc. 1960 1806. l13(a) C. D. Cook R. S. Nyholm and M. L. Tobe J. Chem. SOC. 1965 4194; (6) R. S. Nyholm M. R. Snow and M. H. B. Stiddard ibid. p. 6570; (c) M. R. Snow and M. H. B. Stiddard Chem. Comm. 1965 580. 116 H. C. E. McFarlane and W. McFarlane J. Znorg. Nuclear Chem. 1965 27 1059. J. Lewis and R. Whyman J. Chem. SOC. 1965,5486. 131 Quarter& Reviews An interesting type of heptaco-ordinate complex comprises the metal deriva- tives of v(co) in that there is metal-metal bonding e.g. (C,H,),PAU-V(CO)~ and (triars)Cu-V(CO),.l16 These complexes are monomeric and non-ionic.A related V-V complex is [V(CO)B(diars)], derived from V(CO) and the diarsine.l16 Niobium and tantalum pentahalides react with some donor molecules to give bis-adducts that may be heptaco-~rdinate.ll~-~~* The fluoride adducts are quite insoluble.120 Complexes of the pentachlorides or pentabromides with two molecules of donors such as tetramethylene sulphide pyridine trimethylamine and triphenyl-phosphines -arsines and -stibines have been reported but there is no measurement on molecularity. The sulphide adducts have rather high thermal ~tabi1ity.l~~ Well characterised are the 1 1 diarsine (1) complexes of NbCl and TaCl,; the complexes are monomeric and non-conducting.lZ3 There is a recent review of complexes formed by Group V metal halides.124 Conditions for synthe~is'~~J*~ have been characterised for the isolation of salts of the NbF,2- TaF:- and TaF,& ions from acidic aqueous media but the nature of the solution species has only recently been studied.It appears that the con- centration of ions like NbF72- is very small in aqueous media; the major solution species are NbF,- and NbOF,2-.26,27 On the other hand the tantalum system indicates a much higher stability for TaF,2- in aqueous media.127 This ion and the TaF,- ion are the major ions present in hydrofluoric acid solutions of tan- taluni(v); addition of NH,F to such solutions favours TaF,2- ion formation at the expense of the TaF,- ion. In KF-LiF melts tantalum(v) and niobium(v) are present predominantly as the heptafluorornetallate a n i ~ n s . ~ ~ * J ~ ~ Sodium ion favours formation of the hexafluorometallate ion; both TaF6- and TaF?- ions are present in NaF-LiF melts.Zirconium and hafnium tetrachlorides react with pyridine in benzene to give tris-adducts.130 Molecularity in solution is unknown on warming the adducts lose one molecule of pyridine at ca. 70". Tris-chelates of rare-earth ions particularly those exemplified by the ,8- diketone derivatives separate from protonic solvents with one to three molecules 116 A. S. Kasenally R. S. Nyholm R. J. O'Brien and M. H. B. Stiddard Nature 1964 204 871. 117 G. W. A. Fowles and C. M. Pleass J. Chem. SOC. 1957 2078. 11* P. J. H. Carnell and G. W. A. Fowles J. Chem. SOC. 1959,4113. 119 (a) F. Fairbrother and J. F. Nixon J. Chem. SOC. 1962 150; (6) F. Fairbrother K. H. Grundy and A. Thomson J . Less Common Metals 1966 10 38.120 (a) H. C. Clark and H. J. Emelbus J. Chem. SOC. 1958 190; (b) R. G . Cave11 and H. C. Clark J . Znorg. Nuclear Chem. 1961 17 257. lal J. Desnoyers and R. Rivest Canad. J. Chem. 1965 43 1879. lZ2 K. Lindner and H. Feit 2. anorg. Chern. 1924,132 10. lZ3 R. J. H. Clark D. L. Kepert and R. S. Nyholm J . Chem. SOC. 1965 2877. lZp M. Webster Chcm. Rev. 1966 66 87. 121 A. W. Laubengayer and C. G . Polzer J. Amer. Chem. SOC. 1941 63 3264. lZ6 0. Hahn and K. E. Piitter 2. anorg. Chcm. 1923 127 153. lP7 0. L. Keller jun. and A. Chetham-Strode jun. Inorg. Chem. 1966 5 367. 128 J. S. Fordyce and R. L. Baum J. Chent. Plzys. 1966 44 1 1 59. J. S. Fordyce and R. L. Baum J . Chem. Plzys. 1966,44 1166. I3O T. C. Ray and A. D. Westland Inorg. Chern. 1965 4 1501. 132 Plate 1 Perspectives of tlze idealised octaco-ordinate D4d square antiprism and nonaco-ordinate Dan symmetrically tricapped trigonal prism models illustrating the small distortion required in increasing co-ordination number.Plate 2 Conventional perspectives of the idealised heptaco-ordinate models. From left to right are the C, capped trigonal prism Cs tetragonal base-trigonal base D5h pentagonal bipyra- mid and C, capped octahedron. Plate 3 Perspectives illustrating structural similarities in the idealised heptaco-ordinate models. From left to right are the CaV capped trigonal prism Cs tetragonal base-trigonal base D5h pentagonal bipyramid and C, capped octahedron. Plate 4 Idealised octaco-ordinate geometries. From left to right are the square antiprism (D4d) dodecahedron (D2& undecahedron or bicapped trigonal prism (C2v) s-bicapped trigonal anti- prism (&& and hexagonal bipyramid (D6h).Plate 5 Perspective illustrating similarities in the cube s-bicapped trigonal antiprism and hexagonal bipyramid. Plate 6 Perspective illustrating the similarities in the square antiprismatic dodecahedral and undecahedral models. Muetterties and Wright of s o l ~ e n t . ~ ~ - ~ ~ ~ The solvate molecules cannot (with several exceptions) be removed without solvolysis and decomposition of the tris-chelate strongly suggesting a significant bonding interaction between the metal atoms and the donor atom of the solvent. The monosolvate phases may well contain heptaco- ordinate metal atoms. In solution the co-ordination number of these chelates is not established.Eight or nine may be the average co-ordination number for the large ions such as lanthanum but this probably falls to a limiting value of seven at lutetium. In the tropolone system a dihydrate of tris(tropo1ono)erbium has been isolated but in general the tris(tropo1onates) of the Ianthanides separate from water in anhydrous ~OI-III.~O~ These anhydrous forms are very intractable and are believed to be polymeric through bridging oxygen atoms.lo7 Metal-ion co-ordination numbers in the polymeric structures could be seven or eight.lo7 A tetracyanotrihydroxytechnetium anion Tc(OH),(CN),~ has been described on the basis of isolation and analysis of a thallium salt; the salt was obtained by dissolution of hydrated technetium(rv) oxide in aqueous hydrocyanic acid.140 Protactinyl sulphates PaO(SO& and selenates may be heptaco-ordinate with bidentate sulphate gr0ups.1~~ Arguments142 have been presented for a heptaco-ordinate iridium intermediate IrH,(CO)[P(C,H&l, which is formally related to the well-establishedlo6 ReH,[P(C,H,)& and to ReHs[P(C,H5)J4?43 This class of high-co-ordinate phosphino-metal hydrides may prove to be fairly large.C. Octaco-ordination Octaco-ordination is rather common in complexes of the larger metal ions. The ‘larger metal’ ions begin about scandium so the span of octaco-ordinate structures is really quite large.lo7 (In lattices octaco-ordination has been established for atoms as small as boron e.g. CO,,V,B,.~~) These co-ordination polyhedra seem more pervasive in the chemistry of the lanthanide and actinide ions and the early transition-metal ions particularly those of do dl and d2 configuration although there are apparent octaco-ordinate structures based on metal ions of d10 configuration such as Pb4+ Sn4+ and In%.It remains to be established how extensive octaco-ordination is for metal ions of d10 configuration; only recently 131 G. W. Pope J. F. Steinbach and W. F. Wagner J . Znorg. Nuclear Chem. 1961 20 304. lsa L. C. Thompson and J. A. braas Znorg. Chem. 1963 2 89. lS3 J. G. Stites C. N. McCarty and L. L. Quill J . Amer. Chem. SOC. 1948 70 3142. 130 R. C. Ohlmann and R. G. Charles J . Chem. Phys. 1964,40 3131. lSs L. R. Melby N. J. Rose E. Abramson and J. C. Caris J . Amer. Chem. Soc. 1964 86 5117. 138 (a) C. Brecher H. Samuelson and A. Lempicki J. Chem. Phys. 1965 42 1081; (b) C. Brecher A. Lempicki and H.Samuelson J . Chem. Phys. 1964 41,279. 13’ R. G. Charles and A. Perrotto J . Inorg. Nuclear Chem. 1964,26,373. 138 F. Halverson J. S. Brinen and J. R. Leto J . Chem. Phys. 1964 40 2790. 139 J. R. Ferraro and T. V. Healy J. Inorg. Nuclear Chem. 1962,24 1449. la0 W. Herr and K. Schwochau Angew Chem. 1961,73,492. lol K. W. Bagnall D. Brown and P. J. Jones J. Chem. Soc. 1965 176. la* L. Vaska Inorg. Nuclear Chem. Letters 1965 1 89. 143 M. Freni and V. Valenti Gazzetra 1961 91 1357. H. D. Stadelmaim and J. G. Avery 2. Metall. 1964,56,508. 133 Qwrterly Reviews has this aspect of d1* metal-ion chemistry come under serious investigation A number of idealised geometries have been established or suggested for octaco-ordination. Simplest and most symmetrical of these is the cube which is found in ionic lattices like caesium chloride and calcium fluoride but it has never been observed in a molecular species.Non-bonding repulsions are at a maximum in this regular polyhedron; moreover only seven s p and d orbitals possess proper symmetry for bonding. Although an unfavourable model for most cases the cube is not out of the question for a metal ion offz configuration where there are sufficient orbitals of proper symmetry but the non-bonding repulsions would not of course be relieved in this case. Actually a variant of the cube has been suggested for some uranyl compounds. The variant the symmetrically bicapped trigonal antiprism or puckered hexagonal bipyramid (&) is derived from a cube by a trigonal distortion. One of the uranyl structures 01U03 is purpor- t e d l ~ l ~ ~ only slightly distorted from a cubic array but most uranyl compounds appear to be closer to hexagonal bipyramidal geometry (see Plates 4 and 5).The next most symmetrical of the octaco-ordinate polyhedra is the square antiprism (D4d) (Plate 4). Non-bonding repulsions are significantly lower in the antiprism relative to the cube and there is no symmetry problem for utilisation of eight orbitals with a dx electronic configuration. The square antiprism con- formation is commonly found in the solid state for MX8"- ions and for octaco- ordinate chelate structures. Very closely related to the square antiprism is the D2d dodecahedral structure which is based on two interpenetrating trapezoids. The similarity of the D4d antiprism and geometries is illustrated in Plate 6.In the dodecahedral model there are two types of ligand environment each equally populated. The dode- cahedral structure has been observed in a number of octaco-ordinate chelates and ions; the classic example is the dodecahedral Mo(CN)$- ion.' In one sense the dodecahedral structure bridges the gap between the cube and the square anti- prism; any slight distortion of the cube in the manner outlined in Figure 23 brings it into the Dza point-group. Fig. 23 Distortion of the cube to dodecahedral geometry. Another possible idealised geometry and least symmetrical of the polyhedra l4ti W. H. Zachariasen Actu Cryst. 1948,1,281. 134 Muetterties and Wright is the C, hendecahedron derived from a trigonal prism by the capping of two square faces (Plate 4). This geometry first accorded theoretical status by Kimbal114s from symmetry arguments has not been established for a molecular octaco-ordinate structure but has been defined for some lanthanide and actinide halide lattices.This geometry is a hybrid of the D2d and geometries; the striking similarities of these three geometries are shown in Plate 6. A special case of octaco-ordination is found in metal polyhedra. For example in Pt&2 the octahedral Pt6 group describes the faces of a cube whose edges are defined by the twelve chlorine atoms.147 Each platinum atom is nearly coplanar with four chlorine atoms and is bonded to four other platinum atoms. The geometry for individual metal atoms in these metal polyhedra or clusters is not really relevant to the question of ground-state geometry in molecular species.Moreover Pt,CIl2 can be considered isostructural with Ta,CI,,(OH,)~~ if filled non-bonding platinum orbitals are directed out from the cube faces thus yielding a quasi-nonaco-ordinate platinum atom analogous to the nonaco- ordinate tantalum atoms [Ta(Ta),Cl,OH,] (see Section D). Some rather unusual octaco-ordinate geometries are also found in other metal polyhedra e.g. Fe6(CO),,C,148 but these are defined by the overall metal cluster arrangement. 1 Structural Relationships.-(a) Unidentate Zigands. It is of interest to draw some qualitative energy relationships among the various possible octaco-ordinate structures by considering such factors as steric interactions and promotional energies. Of the octaco-ordinate geometries only three D4d D2d and C2, have proper symmetry to use eight metal orbitals derived solely from s p and d levels.Since these three models are very closely related in geometry (Plate 6) it is not unreasonable to presume that the energy levels for these structures should be comparable. The C, model which is of significantly lower symmetry than the other two may be slightly higher in energy particularly if all eight ligands are identical. These three models differ electronically in that the lowest-energy d orbital is d,a for d for D2d and dxs-,,B for C,,. The cube should be the highest-energy species on the basis of promotional energy and repulsion con- siderations. The repulsion term is relieved to some degree in the D3d and DBd models. Another possible model is the symmetrically bicapped trigonal prism which should be less attractive than the DSd model on purely steric grounds (this DSh structure has never been observed).The lowest-energy d orbitals in the Oh and D systems are a degenerate set d,* and dxa-,, for Oh dxa-ut and d, for D3a d, and d for D6h and dxa-,,a and d, for D3h. A speculative outline of energy relationships for the octaco-ordinate models is depicted in the potential-energy diagram of Figure 24. This diagram is only schematic and no meaning should be read into the finer details of line shapes. The energy barriers for conformational isomerisation should be relatively small. 146 G. E. Kimball J. Chem. Phys. 1940 8 188. 14' K. Brodersen G. Thiele and H. G. Schnering Z. anorg. Chem. 1965 337 120. lC8 E. H. Braye L. F. Dahl W. Hubel and D. L. Wampler J . Amer. Chem. Soc. 1962 84 4633. 135 Quarterly Reviews Beginning at the left only slight distortions are required to isomerise D4& to C2, C2 to D2d to Oh Oh to D3d and D3d to Dan.Unfortunately there is no physical measurement relevant to isomerisation processes and this concept of potential stereochemical non-rigidity cannot be quantitatively delineated at this stage. These comments should not be taken as a generalisation for all octaco- ordinate structures. There should be complexes where a given geometry is signi- ficantly stabilised with respect to the other geometries by.rr bonding or by very low coulombic or non-bonding repulsions. In such cases the barrier to con- formational isomerisation may become quite large and the isomerisation rates may be very low except at high temperatures. E Qi Fig. 24 Possible energy diagram for octaco-ordinate species as a function of configuration.Qi is a shape function relating to interbond angles which describes a continuous path for configuration conversions. Except for the cited uranyl complexes all octaco-ordinate molecular species have either the square antiprismatic or the dodecahedral structure. Since a large number of such structures have been determined one may ask if there are any evident correlations between the configuration adopted and properties asso- ciated with metal ions or the ligands. The answer is no at least with respect to any simple correlation. Clark Kepert Nyholm and Lewis1* note that the more polarisable ligands like cyanide tend to yield a dodecahedral array about a metal atom ; however they emphasise that this correlation is ‘purely tentative’.The only real attempt to elaborate and make the parameters favouring either the prism or the dodecahedron quantitative is that of Hoard and Silvertons and an extension of this treatment by K e ~ e r t . l ~ ~ Hoard and Silvertons consider the direct bonding interaction to be essentially equivalent for the D4d and DZd co-ordination polyhedra and assign stabilisation of either polyhedron to effects arising from (1) non-bonding repulsions in the lQ9 D. L. Kepert J. Chem. SOC. 1965 4736. 136 Muetterties and Wright primary co-ordination sphere (2) coulombic repulsions and (3) geometrical restrictions from ligand stereochemistry. It must be emphasised that the attrac- tion forces are not taken into account and this could severely affect the con- clusions. Hoard and Silvertons conclude that for octaco-ordinate complexes where the ligands are alike and unidentate (1) Calculated angles for the two polyhedra closely approximate those experi- mentally observed.(Hybridisation schemes for metal s p and d orbitals in square antipri~maticl~~~ and d~decahedrall~~~ geometries also yield bond angle values in good agreement with those experimentally observed.) However this is not necessarily a justification of the repulsion model since the basic geometry is assumed for each polyhedron. (2) Energetically the DZd and models are equivalent within the meaning- ful limits of the arguments and assumptions. These two models are preferred to the cube and the hexagonal bipyramid. (3) The intrinsic ligand non-equivalence i.e. environmental as well as metal- ligand bond distance logically provides an operator for energy minimisation of the dodecahedral model relative to a square antiprism for MX4Y4 species in which X and Y are significantly different ligands electronically or sterically.If size is the distinguishing feature the smaller donor atoms should go to the A (Figure 25) sites. This important prediction should be examined experi- Fig. 25 The dodecahedron ( D z d ) with edge and vertex notations following the Hoard and Silverton6 convention. Only two of the eight “g” edges are labelled. mentally. The chelate TiC14,2(diars) a special case of MX4Y4 does in fact have the dodecahedral configuration.151 To speak of stabilisation of a hexagonal bipyramid is conceptually misleading. The hexagonal bipyramid is not a particularly attractive polyhedron for MX or MX,Y species.If however octaco-ordination is to be considered for an MX,Y2 compound where YMY is a rigid collinear array (resonance or n-bond stabilisation) the hexagonal bipyramid the bicapped trigonal antiprism or the cube are the only reasonable models. This is to say that a cylindrical system such as uranyl deformed to adjust to a or D4d polyhedron would represent a very unstable high-energy state. For such systems the dodecahedron and the square antiprism are destabilised. lao(a) G. H. Duffey J. Chem. Phys. 1950 18 746; (b) G. Racah ibid. 1943 11 214. lS1 R. J. H. Clark J. Lewis R. S. Nyholm P. Pauling and G. B. Robertson Nature 1961 192 222. 137 Quarterly Reviews (b) Chelates. The isomer possibilities in dodecahedral or square antiprismatic geometry where the ligands are multidentate is obviously quite large.152 For example consider the possible stereoisomers in a M(chel) structure in which the ligands are identical bidentate and symmetrical? For the square antiprism there are three isomers of plane of symmetry.D, and C symmetry none of which possesses a There are twice as many stereoisomers in the less symmetrical dodecahedron; these have DZd S4 D (two isomers) C, and C symmetries; none except the @@@@ *2d D2 s4 D2 c2 Cl DZd isomer has a plane of symmetry. Generation of these isomers may be more readily visualised from Figure 25 and Table 2 which identify the isomers by the polyhedral edges employed. There are close interrelationships between all nine stereoisomers; that is to say relatively slight distortions can convert a dode- cahedral isomer into a square antiprismatic isomer.Two isomer relationships are outlined for conversion of dodecahedral into square antiprismatic geometry; the fist group is generated by relatively slight distortions and the second group requires much larger distortions (Table 2). Motions required by conversion of the first order are so small that vibronic excitation may suffice for traversing the energy barrier. There are important consequences to this inherent stereochemical non-rigidity namely that two or more geometrical isomers related by one or more first-order conversions should be substantially present in the solution liquid or gaseous state of an octaco-ordinate species if the states of the 152 L. E. Marchi W. C . Fernelius and J. P. McReynolds J. Amer. Chem. SOC.1943,65,329. 138 Muetterties and Wright Table 2 Dodecahedron isomer (dodecahedral edge notation*) Related square antiprism isomers First order Second order D2d (mmmm) D2 D4 c2 s 4 (gggg) c2 D2? 0 4 D,’ (aabb) 0 4 D2 c2 c (mmgg) C and D2 D4 c (abmg) c2 D2 0 4 D2 (gggg) D4and D2 C2 This precisely follows the notations of Hoard and Silverton.’ isomers are comparable. Moreover the racemisation of an optically active form (+) or (-) forms exist for all isomers except D2d) should proceed quite rapidly particularly for the D square antiprism which is the only geometry other than D2d observed for the crystalline state of tetrakis-bidentate ~helates.2~ An estimate for the ground-state lifetime of a (+) or (-) form of the D2 isomer is 10-1 to lo-* second. That is not to say that a single crystal of a (+) or (-) form will not be isolated but that optical activity will not be generally observed for solutions of a D2 square antiprismatic isomer.The longest-lived optical isomers appear to be those with D,’ dodecahedral geometry. This type of argument can easily be extended to unidentate structures of the type MA,B,,. Hoard and Silverton5 have analysed relative stabilities of ground-state geo- metries for chelates and presented several conclusions based on the nature of the chelate and the overall charge on the complex. (1) If the ligand is coplanar and if the complex is charged minimisation of the coulombic interaction occurs with dodecahedral geometry and D2d symmetry. This class is exemplified by Zr(C204)44- and Cr(02)43-. (2) For neutral tetrakis-complexes in which the metal-oxygen bond is long e.g.> 2.30 A for acetylacetonates the shorter polyhedral edges (e.g. a b in Figure 25) are clearly preferred to minimise ring strain. This argument uniquely identifies the DZdDOD and D2S* models as the low-energy species as is observed for the cerium thorium and uranium tetrakis(acety1acetonates). We note an extension of this point. Matching of polyhedral edges to the oxygen-oxygen separation in the highly rigid tropolone ligand (4) should be the critical factor in determining ground-state geometry for tetrakis(tropol~nates).~~~ On this basis the D2 square antiprismatic and the DZd dodecahedral isomers should be the more stable isomers for all tetrakis(tropo1onates) except possibly those based on very small metal ions such as tin(rv).S Any departure from these predictions may reflect a contraction or elongation of the metal-oxygen bonds generated by coulombic repulsions in charged complexes.Since the range of tetrakis- (tropo1ono)-chelates is large Sc3+ to U4+ this series is worth detailed structural investigation. (3) Neutral M(chel) molecules in which the metal-oxygen bond distance is $ Molecular mDdels for (tropolono),Zr support this proposal. Ease of connectivity goes D < C < D in square anti-prismatic geometry and D,’ < C < C f D < S c Datl in dodecohedral geometry. 139 QuarterIy Reviews less than 2-30 A may have D or D symmetry in square antiprismatic geometry or D2d S, or D symmetry in dodecahedral geometry. It is rather difficult to present a more detailed energy differentiation among these five models without considering specific metal ions and specific ligands.So far all determined structures have essentially D square antiprismatic or D2d dodecahedral geometry. More detailed analyses of this class of chelates is in the paper by Hoard and Sil~erton.~ (4) Constraints imposed upon geometry by ter- quadri- quinque-dentate etc. chelates are severe and generalisations for this class are not significant. Any conclusions must be based on a specific multidentate ligand. In the ethylene- diaminetetra-acetate anion the six donor ligands readily conform to six of the eight general positions for the dodecahedron namely the two nitrogen atoms at adjacent A vertices and four oxygen atoms at B vertices (Figure 25). In the nitrilotriacetato-anion with the nitrogen atom at an A vertex the glycinate rings span a rn and g edges and a twofold rotation will generate the other half of a quasi-dodecahedral complex as found for bisnitrilotriacetatozirconate(1v) (-2).2 Structural Data for the Crystalline State.-(a) Dodecahedron (Table 3). The dodecahedron can be visualised as two structurally equivalent mutually per- pendicular trapezoids. Alternatively the structure can be visualised by dividing the eight ligands into two sets of four that form an elongated tetrahedron and a flattened tetrahedron which are interlocked. Figure 23 shows the distortion of the cube leading to the dodecahedral structure. This structure was first estab- lished by Hoard and Nordsieck in 1939 as the co-ordination polyhedron of the eight cyanide groups about molybdenum in K,Mo(CN),,~H,O.~ This analysis based on two-dimensional data has been reexamined with three-dimensional data to obtain more highly refined values for bond angles and distances.Within experimental error there is no difference in bond distance between the A and B type (Figure 25) Mo-C bonds (personal communication from Professor J. L. Hoard). Dodecahedra1 geometry has also been established for ‘chelate’ structures based on the compact peroxy and nitrate ligands. In K3Cr0, the chromium ion is surrounded by eight oxygen atoms as first shown by Stomberg and Brosset from two- and three-dimensional X-ray re~u1ts.l~~ A refinement of this structure by SwaIen and I b e r ~ I ~ ~ with the three-dimensional data established the oxygen- oxygen bond length in the peroxide group to be 1.405 f 0.039 A which is shorter than a normal peroxy-separation and the chromium-oxygen separations to be 1-846 f 0-022 A and 1.944 f 0.024 A.Isomorphous with K,CrO are K3Nb0, K,TaO, and K3V08>55,156 The nitrate ion is bidentate in [As(C,H,),],- [Co(NO,)J with an oxygen-oxygen separation of 2-1 A.16 The co-ordinated lS3 R. Stomberg and C. Brosset Acra Chem. Scand. 1960 14 441. lS4 J. D. Swalen and J. A. Ibers J. Chem. Phys. 1962 37 17. lS5 J. E. Fergusson C. J. Wilkins and J. F. Young J. Chem. Soc. 1962 2136. lS6 G. Boehm 2. Krist. 1926 63 319. 140 Muett ert ies and Wr&h t oxygen atoms form a distorted dodecahedron around cobalt so that the observed symmetry is only C,,,. Four metal-ligand distances are 24-2.11 A and four are 2.36-2.54 A which define an elongated and severely flattened tetrahedron respectively.The nitrogen atoms (centre of the nitrate group) are at vertices of a flattened tetrahedron. The two N-Co-N angles bisected by the molecular C axis are 152" and 144" compared with 109" for a regular tetrahedron. Cotton and Bergman have synthesised an analogous compound [As(C,H,),],- [Co(O,CCF,)J which has virtually the same electronic spectrum as the nitrate complex.l6 A three-dimensional structural determination20 of tetrasodium tetrakis- oxalatozirconate(rv) trihydrate established the co-ordination polyhedron around zirconium as dodecahedral and thus fixed the structure of the isomorphous hafnium complex. The average zirconium-oxygen distances are 2.1 68 8 (form- ing the flattened tetrahedron) and 2.230 8 (forming the elongated tetrahedron) and the average intra-ring oxygen separation is 2.563 A with essentially planar rings.The angle at zirconium in the ring is 71.3" and the average 0-Zr-0 angle between adjacent rings in the same plane is 70.4" and 147". The tetrakisdibenzyol- methane derivative of cerium(rv) also has the dodecahedral geometry with Ce-0 distances of 2.4&2.44 A and the thorium and uranium(1v) analogues are isomorphous by diffraction ~riteri0n.l~' A structural determination by three-dimensional X-ray analysis of the dipo- tassium salt of bis(nitrilotriacetato)zirconate(Iv) monohydrate established the [N(OAc),I3- ligand as quadridentate. The co-ordination polyhedron adheres to the constraints of the dode~ahedrall~~ model although the Zr-N distance is abnormally long (Figures 25 and 26). Two oxygens and two nitrogens lie in *-- 0 Zr -0 =2.13A Zr -0 = 2.25 Zr-N =2.75 Fig.26 The dodecahedral co-ordination of the nitrilotriacetato-ligands about W,. in K,Zr- (nitrilotriacetate)2,0H2. The broken lines represent the bridging groups of the chelate ligand (ref. 158). one tetrahedron with bond lengths of 2-251 and 2.439 A respectively and the zirconium-oxygen bonds in the other tetrahedron are 2.124 and 2.136 A which is in agreement with those found for unidentate ligands. Isomorphous with Zr[N(OAc),],2- is H~[N(OAC),],~-.~~~ 15' L. Wolf and H. Barnighausen Acta Cryst. 1960 13 778. 158 J. L. Hoard E. Willstadter and J. V. Silverton J . Amer. Chem. SOC. 1965 87 1610. lS9 J. L. Hoard J. V. Silverton G. L. Glen and E. Willstadter Proc. Seventh ICCC Stock- holm Uppsala (June 1962). 141 Quarterly Reviews A number of diarsine (1) complexes of the type MCI4,2(diars) have been pre- pared by Clark Lewis Kepert and Nyholm.14Jso The crystal structure has been established for TiC14,2(diars) by P.Pauling et aZ.151 and is octaco-ordinate with the diarsine rings bridging edge a and a (Figure 25) of the two interpenetrating trapezoids. The titanium-arsenic distance is 2-71 f 0.02 A the titanium- chlorine distance is 2-46 & 0.02 & and the separation between adjacent arsenic atoms is 3.21 A. Isomorphous with TiCI4,2(diars) are MX4,2(diars) where M = Zr Ti Hf Nb or V and X = Br or CI.l5l The dodecahedron is also found in polymeric octaco-ordinate units. The zirconium and hafnium atoms in K2ZrF6 and K2HfF6 respectively are sur- rounded by eight fluorine atoms four of which are shared. The polyhedra are connected in a chain-like fashion by sharing two opposite edges of the MF polyhedron.The crystal-structure determination of potassium hexafluorozir- conium(1v) gave 2-1 2-2-26 A as the zirconium-fluorine separation.161 The zirconium atom in Li,BeF,ZrF is also surrounded by eight fluorine atoms; four at 2-05 A (forming the flattened tetrahedron) and four at 2.16 A from the zirconium.162 An X-ray powder study of thorium tetrachloride and uranium tetrachloride suggests that the structures are isomorph~us.~~~ Each metal appears to be surrounded by eight chlorine atoms. In ThCI four chlorines are apparently at a distance of 2.46 A and four at 3.11 8 from the metal. The corresponding distances in UCI are 2.41 A and 3.09 A. Isomorphous with UC14 are NpC14 and PaC14.1aa The lattice of gallium dichloride is based on an array of Ga+ and GaC14- ions.ls5 The tetrachlorogallate anion is tetrahedral with gallium-chlorine distances of 2.19 A.Each Ga+ ion is surrounded by eight chlorine atoms from six different tetrahedra. The chlorine atoms occupy positions at comers of an irregular dodecahedron with four at 3-18 A and four at 3.27 A from the Ga+ ion. The dodecahedron is also found in zircon (ZrSi04)166 in which zirconium is bound by four oxygen atoms at 2-15 A and four at 2.29 A. The same oxygen polyhedron is found in yttrium phosphate (YPo4)186~u7 In addition the oxygen dodecahedra1 co-ordination polyhedron is found in the garnet structure. The garnets are a group of orthosilicates of which grossularite Ca3AI,(Si04), uvarovite Ca,Cr,(SiO,), and andradite Ca,Fe,(SiO,), are exampIes.168,16B The general formula is R311 RZII1 (SO,) with packing of the SiO tetrahedron such that the RII ions are eight- and the RII1 ions are six-co-ordinate.Structurally related to these silicates are the fluoroaluminates e.g. Na3A12(LiF4)3.170 The lanthanum cerium and neodymium phosphates exist in the monoclinic 1 6 0 R. J. H. Clark D. L. Kepert and R. S. Nyholm Nature 1963 199 559. 161 H. Bode and G. Teufer Acta Cryst. 1956 9 929. lci2 D. R. Sears and J. H. Burns J. Clzem. Phys. 1961 41 3478. 163 R. C. L. Mooney Acta Cryst. 1949 2 189. 164 W. H. Zachariasen Acfu Cryst. 1949 2 388. 165 G. Garton and H. M. Powell f. Inorg. Nuclear Chem. 1957 4 84. 166 I. R. Kirstanovic Actu Cryst. 1958 11 896 167 M. K. Carron M. E. Mrose and K. J. Murata Amer. Min. 1958 43,985.168 A. L. Gentile and R. Roy Amer. Min. 1960 45 701. 1-59 A. Durif Electron Telecomm. Internat. Conf. Brussels vol. 3 part 1 500 (1958). 1'0 G. Menzer Z. Krist. 1930 75 265. 142 Muetterties and Wright or hexagonal structure. The monoclink form is isomorphous with the mineral monazite. In the hexagonal structure the metal is octaco-ordinate.171 In CeP04 the cerium is surrounded by four oxygens at 2.34 8 and four at 2.66 A and the co-ordination polyhedron approximates dodecahedra1 geometry. In a sulphate of zirconium Zr,(OH),(S04),(H20)4 the zirconium is octaco-ordinate and the oxygen co-ordination polyhedron approximates the dodecahedron.172 Each Table 3 Dodecahedra1 octaco-ordinate structures NdPO4 Y PO Ce(DBM)4j Ga+[GaCI,l- ThC14 UCl NPC4 PaCI /3-ThBr4 Hard-sphere model Hoard-Silverton model d4sp3 Hybrid orbitals M-Aa(av.) Bond distance (A) 6~"(av.)Q 2.71 2.202 2-22 2.29 2-16 2.21 2-25(0) 2*44(N) 2.230 1.846 2.15 2.11 2.03 2.34 2-42 3.18 3.1 1 3.09 2-86 1.00 1 -03 36 43 35.3 35.2 43.4 36.0 47 36 36-85 35.2 34.55 M-Ba(av.) Bond distance (A) SBo(av.)a Ref.* 2.46 2.185 2.22 2.15 2.05 2.16 2.13 2.1 68 1 -944 2.1 5 2.36 2.54 2.66k 2.42 3.27k 2.46 2.41 2-61 1.00 1 -00 73 65.5 71-8 73.5 86-8 71.8 81 75 69.46 73.5 72.78 151 160 172 173 166 162 161 161 158 159 20 20 153 154 155 156 155 156 155 156 192 1 1 16 171 171 171 166 167 157 165 163 163 164 164 1 149 5 150b aSee Figure 25 for bond distance edge and angle notation; bdiars = o-phenylenebisdi- methylarsine; clsomorphous with TiCI4,2diars are M = Zr Hf V Nb; X = C1 and M - Ti Nb Zr Hf; X = Br.Probably also isomorphous with TiCI4,2diars are M = W Tc Re; X = Ci; M = Re; X = Br and M = Nb X = I; da = 2.78 b = 3.00-3.42 g = 2.62- 2.82 m = 2-34-2.63 A; ea = 2.42 g = 2.85 m = 2-54 A; fa = 2.52 g = 2-68 m = 2.47 A; QN(OAC),~- = nitrilotriacetic acid ion a = 2.68 g = 2.785 m = 2.62 A; ha = 2-57 b = 3.19,g = 2-735,m = 2*563A;tK3Cr0,,a = 2-57 b = 2.75,g = 2.74.m = 1*49A;jDBM= dibenzoylmethane ion a = 2.667 m = 2.866 A; kSpecific assignments of distance to the M-A and M-B bonds were not made by the authors; L D.E. Scaife Inorg. Chem. 1966 5 162. * For references see Text. 171 R. C. L. Mooney Acra Cryst. 1950 3 337. 172 D. B. McWhan and G. Lundgren Inorg. Chem. 1966 5,284. 143 Quarterly Reviews zirconium atom has two hydroxide oxygens four sulphate oxygens and two water oxygens at a mean distance of 2.19 A.Dodecahedra1 geometry is also found in Zr(OH),(N03),(H20) with a Zr-0 distance of 2-22 A.173 In other zirconium salts there is a near-square antiprismatic geometry (see Section C.2b). Lanthanum telluride La2Te3 described as a Th3P4 type structure has a lanthanum co-ordination sphere of eight tellurium atoms with La-Te distances of 3-244 and 3-418 A?74 A number of heavy-metal phosphides arsenides anti- monides and tellurides have this structure e.g. U3As4 and U,Te, as well as some rare-earth sulphides selenides and tellurides e.g. Ce2S3.175J76 The poly- hedron might be described as a highly distorted dodecahedron or square anti- prism; however any description in terms of idealised geometries is unrealistic for this particular polyhedron. (6) Square antiprism (Table 4).The square antiprismatic geometry is well established by X-ray analysis for discrete octaco-ordinate unidentate and bidentate compounds and for polymeric or ionic lattices. The only discrete complexes with unidentate ligands are the octafluorometallate ions. Sodium octafluorotantalate exists in the crystalline state with near D4d symmetry for the TaFS3- complex ion at least within the accuracy of the two-dimensional X-ray analyskll There is however a high probability that some distortion from full D4d symmetry arises from packing forces. The Ta-F bond distance ranges from 1.93-2.01 A with an average length of 1.98 A and the average F-F separation is about 2.41 8 and 2.42 A for square edges and triangular edges respectively. A similar structure apparently obtains in the K2ReF8 lattice (two-dimensional X-ray The Re-F distance varies in the range 1437-1.93 A.The struc- tures of such fluro-complexes as R u F ~ - ) ~ ~ WF 8 9 2- 179 M o F ~ - - ~ ~ ~ ~ ~ ~ XeF8”;38b the apparently isostructural UFSs and PaFss,lsl and TeF,2-ls2 have not been determined. Preliminary analysis of X-ray measurements for Na,UF indicates the uranium atom is octaco-ordinate (U-F = 2-29 A).183 The tetrakis(acety1acetonates) of zirconium cerium(Iv) and thorium are reported to use the square antiprism co-ordination polyhedron. The zirconium- (IV) acetylacetonate whose structure was determined by Silverton and HoardlS4 with three-dimensional X-ray data is of nearly D symmetry with an average zirconium-oxygen bond length of 2.198 A and a 0-Zr-0 ring angle of 75’. The intra-ring oxygen-oxygen distance is 2.674 A and the acetylacetonate ring 173 D.B. McWhan and G. Lun‘dgren Acfa Cryst. 1963 16 36. 174W. L. Cox H. Steinfink and W. F. Bradley Znorg. Chem. 1966 5 318. 175 J. Flahaut M. Guittard M. Patrie M. P. Pardo S. M. Golabi and L. Domange Acfa Cryst. 1965 19 14. 176 W. H. Zachariasen Acta Cryst. 1949 2 57. 177 P. A. Koz’min J. Struct. Chem. (U.S.S.R.) 1964 5 60. 178 E. E. Aynsley R. D. Peacock and P. L. Robinson Chem. and Znd. 1952 1002. 179 B. Cox D. W. A. Sharp and A. G. Sharpe J. Chem. SOC. 1956 1242. 180 G. B. Hargreaves and R. D. Peacock J . Chem. SOC. 1958 4390. l e a E. L. Muetterties J . Amer. Chem. SOC. 1957 79 1004. 1e3 J. G. Malm H. Selig and S . Siegel Inorg. Chem. 1966 5 130. 1e4 J. V. Silverton and J. L. Hoard Inorg. Chem. 1963 2,243. D. Brown and J.F. Easey J. Chem. SOC. ( A ) 1966 254. 144 M uetterties and Wright folds out about this edge of the antiprismatic square face by an angle of 22.6f7". The inter-ring square edge distances average 2.590 A. The distortion of the structure from square antiprismatic geometry is toward D dodecahedral. Hafnium(1v) acetylacetonate is isomorphous with the zirconium chelate by diffraction criterion. An isomorphous p-thorium(Iv) acetylacetonate of D point-group symmetry has been reported by Grdenid and MatkovidlS5 although the structure was not sufficiently refined unequivocally to establish stereo- chemistry. The average thorium-oxygen distance is 2.41 A with an 0-Th-0 angle of 70"; the oxygen-oxygen separation within acetylacetonate rings is 2.74 A; between adjacent rings it is 3-05 A and 3.10 A along the square and tri- angular face edges respectively.Tetrakis(acety1acetonato)plutonium is reportedlse to be isomorphous with P-thorium(rv) acetylacetonate. From a two-dimensional X-ray analysis of cerium(1v) tetrakis(acetylacetonate) square antiprismatic geometry has been suggested with an 0-Ce-0 angle of 72" and a cerium-oxygen bond distance of 2.40 Within the ring the oxygen-oxygen separation is 2.81 A and the distance between adjacent rings is 2-97 A and 2.95 8 along square faces and triangular faces respectively. Isomorphous with t e t rakis(acety1acet 0nato)cerium are the a-ur anium(1v) and a-thorium(1v) acetylacetonate~.~~~~ The crystalline 01 form of the thorium(1v) tetrakis(acety1acetonate) is formed spontaneously at room temperature from the /? Grdenid and M a t k o ~ i i l ~ ~ ~ attribute the two modifications of the thorium chelate to differences in packing arrangements but the structures of the two forms are not sufficiently well established to preclude alternative explana- tions such as changes in stereochemistry.A polyhedron of eight oxygens in a distorted square antiprismatic arrange- ment surrounds yttrium in yttrium(rr1) tris(acety1acetonate) trihydrate.ls8 The yttrium is co-ordinated to the six oxygens of the acetylacetonate ligands which bridge square edges and to two water oxygens which are at vicinal positions. The average Y-Oacac distance is 2.367 A the Y-Owater distance is 2-41 A and the average 0-Y-0 angle formed with the acetylacetonate ring is 72.5". Co-ordination polyhedra approximating that of the square antiprism have been suggested for a number of ionic or polymeric lattices.In all cases there is distortion of the polyhedron and the distortion is reportedly high in some of the polyhedra. Unfortunately the rigour of the analysis in most of the X-ray structural determinations of these 'polymeric lattices' falls short of establishing unequivocally the geometry of the co-ordination polyhedron. Therefore the assignments cited below might possibly be changed upon full three-dimensional analysis to dodecahedral or hendecahedral geometry. The lattice structures of a-ZrF,189a (the structure of the /? form has not been lE5 D. Grdenit and B. MatkoviC Nature 1958 182 465. lS6 A. E. Comyns Actu Cryst. 1960 13 278. lS7 (a) B. MatkoviC and D. Grdenic Acta Crysf. 1963 16 456; (6) D. Grdenic and B.Matkovii Acta Cryst. 1959 12 817. lE8 D. E. Sands J. A. Cunningham and W. F. Wagner; D. E. Sands personal communi- cation 1966. (a) R. D. Burbank and F. N. Bensey jun. U.S. Atomic Energy Corn. K-1280 (1956); (b) A. C. Larson R. B. Roof jun. and D. T. Cromer Actu Cryst. 1964,17,555. 145 Quarterly Reviews deterrninedlgoa) and UF,1S4JsQb are essentially identical. The heavy-metal atoms are within bonding distance of eight fluorine atoms which describe a slightly dis- torted square antiprism. Metal-fluorine distances range from 2-23-2.35 A in UF and 2-03-2-18 8 in a-ZrF,. The tetrafluorides of hafnium thorium neptunium plutonium cerium,lU praseodymium,190b terbium,lgoc americi~m,lgO~ and curium1goe are isomorphous with a-ZrF by diffraction criterion. In HfF,,- 3H20 there is a polymeric chain of hafnium atoms bridged by fluorine atom~.l~l The co-ordination polyhedron is described as a distorted square antiprism.For the compositionally related ZrF4,3H201g2 there are dimers instead of polymers and the zirconium co-ordination polyhedron apparently more closely resembles the dodecahedra1 geometry. An X-ray analysis of Eu(H,O),C& established a lattice of EU(H20),C12+ and Cl- The geometry of the europium co-ordination sphere was described as a highly distorted square antiprism with europium-oxygen and -chlorine distances of 2-44 and 2.77 A respe~tive1y.l~~ Acceptable alternative descriptions of the geometry are a distorted dodecahedron and a distorted hendecahedron. There are a number of isomorphous hydrated lanthanide and actinide chlorides (Nd Sm Er Gd and Pu).lg4 Analysis of the gadolinium salt gave a geometry quite comparable with the europium case.lg4 The average bond lengths are 2.77 and 2-40 A for the Gd-CI and Gd-0 bonds.Barium hydroxide octahydrate has based on a two-dimensional X-ray analysis a slightly distorted square antiprismatic array of eight water molecules about the barium atom at a distance of 2.69-2.77 A?95 The same metal-ion co- ordination polyhedron is suggested for the strontiumlg6 analogue and for Ca02,8H20.1g7 These analyses are based on two-dimensional data. Octaco-ordination is found for iodine in a number of iodates e.g. the iso- morphous Ce(I0,)4 and h(Io3)4,1g8 Ce(103)p,H20,199 Zr(10,)4,200 and NaIO,) ;201 and the co-ordination polyhedron is described as a ‘crude antiprism’. Five non- bonded oxygen atoms plus the three associated with the iodate complete the polyhedron.In Ce(IO,) the cerium atom has seven oxygen atoms at 2.18- 2.41 A and one at 2-82 8 to form a ‘much distorted Archimedean antiprism’.lg8 Each of two oxygen atoms per iodate group is shared with the cerium atom. Less lgo (a) V. Amirthalingam and K. V. Muralidharan J. Inorg. Nuclear Chem. 1964 26 2038; (b) J. Soriano M. Givon and J. Shamir Znorg. Nuclear Chem. Letters 1966,2 13; (c) B. B. Cunningham D. C. Feay and M. A. Rollier J. Amer. Chem. SOC. 1954,76 3361; ( d ) L. B. Asprey ibid. p. 2019; (e) L. B. Asprey F. H. Ellinger S. Fried and W. H. Zachariasen ibid. 1957 79 5825. lgl D. Hall C. E. F. Richard and T. N. Waters Nature 1965 207 405. lg2 T. N. Waters Chem. and Ind. 1964 713. lg3 N. K. Bel’skii and Yu. T. Struchkov KristallograJ’iya 1965 10 16.lgP M. Marezio H. A. Plettinger and W. H. Zachariasen Acta Cryst. 1961 14 234. lg5 H. Manohar and S. Ramaseshan Z . Krist. 1964 119 357. lg6 H. G. Smith Acta Cryst. 1953 6 604. lg7 R. S. Shineman and A. J. King Acta Cryst. 1951 4 67. lg8 D. T. Cromer and A. C. Larson Acta Cryst. 1956 9 1015. lg9 J. A. Ibers Acta Cryst. 1956 9 225. A. C. Larson and D. T. Cromer Acta Cryst. 1961 14 128. 201 C. H. MacGillavry and C. L. P. Van Eck Rec. Truv. chim. 1943 62,729. 146 Muetterties and Wright distorted is the antiprismatic arrangement of oxygen atoms about cerium in Ce(I0,)4,H20.199 The average Ce-0 distance in the hydrate is 2.33 A. Three- dimensional X-ray studies of Zr(XOJ4 show a nearly perfect antiprism of oxygen atoms about the zirconium with an average Zr-0 distance of 2.210 A.2oo Octaco-ordination has been suggested for zirconium cerium thorium and uranium in a number of oxygen derivatives and a discussion of this class of basic salt is given by Lundgren.202 Zirconium sulphate tetrahydrate has a layer lattice of Zr(SOk),,4H20 composition.203 Each zirconium atom is within bonding distance of four oxygen atoms from water molecules and four oxygen atoms from individual sulphate ions.The average Zr-0 distance is 2.18 A. This structure is closely related to U(S04)2,4H20.204 The zirconyl halide octahydrates ZrOCl,,- 8H,O and ZrOBr2,8H20 are isomorphous and have a near-antiprismatic co- ordination sphere of eight oxygen a t ~ m ~ . ~ ~ ~ ~ ~ ~ ~ Lundgren has suggested octaco- ordination with a near-square antiprismatic arrangement of oxygen atoms in Zr(OH)2S0,,172 Th(OH)2S04,207 U(OH)2S04,208 Th(OH)2Cr04,H20,209 A recent crystal-structure detemination2l3 of RbLiF and isostructural CsLiF shows that rubidium is surrounded by eight fluorines in a square anti- prismatic arrangement at distances of 2.78-3.16 A (average 2.95 A).In CsLiF, casium is surrounded by six fluorines at 2.96-3.15 A and two at 3.50 and CeOSO4,H2O,2l0 U(SOp),,4H2OY U604(0H),(S0~)6,211 and Ce604(OH)4(S0~)6.212 3.53 A. Table 4 Square antiprismatic octaco-ordinate structuresa Compound NaIOs RbLiP2 CsLiF Ca02,8Hz0 Sr02,8 H20 Ba02,8H,0 SrCI2,2H,O BaC1,,2H20 Sr(OH),,8 Ha0 MOW23 Ha0 HfF, 3 H2O a-ZrF4b UF HfF* C e F I M-X Squar 8" Bond distance (A) edge (A) 2.78-3.1 6 2.96-3-15 3.50-3.53 3.1 3-3.36 2-60 2.74 209 2.12 2.28 3.23 57 202 G. Lundgren Svensk Kem.Tidskr. 1959 71 200. J. Singer and D. T. Cromer Acra Cryst. 1959 12 719. 204 P. Kierkegaard Acra Chem. Scand. 1956 10 599. A. Clearfield and P. A. Vaughan Acra Cryst. 1956 9 555. 206 G. M. Muha and P. A. Vaughan J. Chem. Phys. 1960,33 194. 207 G. Lundgren Arkiv Kemi 1950 2 535. 208 G. Lundgren Arkiv Kemi 1952 4 421. 209 G. Lundgren and L. G. SillCn Arkiv Kemi 1949 1,277. *lo G. Lundgren Arkiv Kemi 1953 6 59. 212 G. Lundgren Arkiv Kemi 1956 10 183. 213 J. H. Burns and W. R. Busing Znorg. Chern. 1965.4 1510. G. Lundgren Arkiv Kemi 1953 5 349. TriangIe edge (4 2oyf.* Y e 21 3 21 3 197 e 197 e g 196 3.50 195 191 189a 189b 164 164 147 Quarterly Reviews Table &continued Compound ThF4 NPF4 PuF4 Zr ( a ~ a c ) ~ Rf( a ~ a c ) ~ flh(ac44 Pu(acac) Ce(acac) a-U(a~ac)~ ~x-Th(acac)~ ZrO C12,8H20 ZrOBr2,8H20 Zr(IOd4 Zr(S04)8,4&0 Th(0H j~Cr04,H20 U(S04)2,4HzO U(OHhSO4 u 60doH)4(so4) 6 Hard-sphere model Hoard-Silverton dSp3 Hybrid orbitals d4sp3 Hybrid orbitals model eo 57.3 58.2 58.5 58 57 59 58 58 54 59-25 57.3 60.9 57.6 M-X Bond distance (A) 2.198 2-41 2.40 2.24 2.21 2.178 2.19 1.98 1.87-1.93 2.41 2.33 2.18-2'41 2.82 2-44(0) 2*77(C1) 2*41(0) 2-77(C1) 3-20 2-4 2.5 2.3 2.3 2-3 1.000 1.000 Square Triangle edge (A) edge (A) 164Ref .* 164 164 2.59,2.676 2.74 184 184 3.05 2-74d 3.10 185 186 2.97 2.81 2.95 187 187 187 205 205 2.65 2.73 200 2531,2623 2719 203 2859.2648- 172 2.41 2.42 11 177 212 210 2.83 2.75 199 198 198 2.91 3.33 2864.19 193 194 194 194 194 194 207 209 204 208 21 1 1.215 1.215 149 1.190 1.258 5 3.85 4.02 223a b 150a 150a aThe groups indicate isomorphous series; 6There are two kinds of zirconium atom one with a metal-fluorine distance of 2.09 the other with a metal-fluorine distance of 2.12 A.All Zr-F distances within a given polyhedron are equal; Cacac = acetylacetonate ion; dThe length of the edge the acac ring spans; * For references see Text except as follows eR. W. G. Wyckoff 'Crystal Structures' Interscience New York 1965 2nd edn. vol. 3 p. 842; PA. T. Jensen Kg1. danske Videnskab Selskab Mat.-fys. Medd. 1943 20 No. 5; gldem ibid. 1945 22 NO. 5. (c) Hexagonal bipyramid (Table 5). Although not yet subjected to the test of a three-dimensional X-ray analysis the existence of hexagonal bipyramidal geometry in uranyl compounds seems fairly well established. First the collinear or very nearly collinear nature of the uranyl (UO,) group appears to be on a moderately sound basis.In RbUO,(NO& the best fit of two-dimensional X-ray and neutron diffraction data is with an arrangement in which the collinear 148 Muetterties and Wright OUO group (U-0 distance 1.78 A) is bisected by a plane of three nitrate groups.214,216 The hexagon of oxygen atoms is slightly puckered; oxygen atoms are alternately 0.09 A above and below the plane. Each nitrate group is effectively bidentate with respect to the uranium atom yielding a non-regular (threefold symmetry) hexagonal co-ordination. A similar arrangement has been found consistent with two-dimensional X-ray measurements for NaU02(0COCH& with three acetato-groups providing the threefold symmetry of hexaco-ordina- tion in the plane perpendicular to the OUO group,2lS and an analogous three- dimensional variant in U02C03 with bidentate carbonato-gr~ups?~~ Powder data for U02(N03)2,6H20218 and U02(N03)2,20P(OC2H,),219 have been interpreted in a formally similar fashion with two nitrate and two 'para' oxygen atoms from the solvate molecules describing the hexagonal co-ordination sphere.Recently the U02(N03),,6H20 structure was reexamined with three- dimensional neutron diffraction The uranyl group (U-0 distance 1.76 A) is surrounded by a near-planar hexagon of four oxygen atoms from two bidentate nitrate groups and two equivalent water oxygen atoms. Table 5 D3d and Dan Octaco-ordinate structures Compound Bonda distance (A) Ref.* Compound Bond" distance (A) Re/.* CaUOaOa 2U-01= 1-92 145 RbUOz(N03)s 2U-01 = 1.78 215 6U-011 = 2-29 8Ca-O = 2-45 [UO,(NO*),.2U-01 = 1.76 220 Bicapped trigonal antiprism Planar hexagonal bipyramid 6U-011 = 2.48 2HaOI.4HaO 2U-011 = 2.397 SrUOaOa 2u-011= 1-91 145 2U-011 = 2.504 2U-011 = 2.547 6U-Or1 = 2.33 8Sr-0 = 2.58 UO~COI 2U-01= 1-67 UO,F 2U-01 = (1-91)b e 4U-011 = 2.52 217 6U-F = 2.50 2U-011 = 2,44 6Am-F = 2.47 ~Pu-OII= 2.55 KAmOaFz 2Am-01 = (1*93)b f KpuOaCOs 2Pu-01 = (1*94)b g RbAmO&O3' g or-U03 2U-01= 2-08 222 6U-011 = 2.39 145 NaUOa(O,CCHa) 2U-01= 1.71 216 6U-011 = 2.49 NaMOa(OsCCH3),C M = Np Pu Am 164 216 uo z(N03)~f(c~H 60)3POI*d 219 "The uranyl oxygens are designated 01 and the other six oxygens 011; m e s e distances are estimated; "This compound is isomorphous with the one directly above; NO parameter is available; * For references see Text except as follows eW.H. Zachariasen Acta Cryst. 1948,1 277; Idem ibid. 1954,7 795; OF. Ellinger and W. H. Zachariasen J. Phys. Chem. 1954 58 405. 214 J. L. Hoard and J. D. Stroupe National Nuclear Energy Series Div. 111 2 13 McGraw Hill Book Co. New York 1949. 215 G. A. Barclay T. M. Sabine and J. C. Taylor Act@ Cryst. 1965 19 205. alb W. H. Zachariasen and H. A. Plettinger Acta Cryst. 1959 12 526. 217 D. T. Cromer and P. E. Harper Acta Cryst. 1955 8 847. 218 J. E. Fleming and H. Lynton Chem. and Ind. 1960 1416. 219 J. E. Fleming and H. Lynton Chem. and Znd. 1959 1409. 2zo (a) J. C . Taylor and M. H. Mueller Acta Cryst. 1965 19 536; (6) B. 0. Loopstra and E. H. P. Cordfunke Rec. Trav. chim. 1966,85 135. 149 Quarterly Reviews Like the cube the hexagonal bipyramid is not an attractive model on purely symmetry grounds for metal ions that do not utilisef orbitals.( d ) Bicapped trigonal antiprism (puckered hexagonal bipyramid) (Table 5). There are insufficient structural data to place the symmetrically bicapped trigonal antiprism (0 d ) among the established octaco-ordinate geometries. Such a geometry has been suggested for the layer lattices found in a-U03 CaUO, SrU04 and UO,F but the X-ray data consisted solely of powder diffraction patterns.145 A recent neutron diffraction study of the hexagonal a-U03 suggests that this form is an imperfectly crystalline form of rhombohedral 01-U03?20b (e) Undecahedron or bicapped trigonal prism (Table 6). Experimental evidence definitive for the C, undecahedron is now available from the single-crystal X-ray analysis of terbium trichloride.221 The chlorine atoms at trigonal vertices are at 2.70 and 2.79 A from the terbium atom; the two chlorine atoms at the square pyramidal vertices are at 2-95 A.A ninth chlorine atom is at a distant approach of about 4.0 A and is non-bonding. This type of structure was first suggested by Zachariasen222 for several lanthanide and actinide halides e.g. PuBr and NdBr, on the basis of X-ray powder diffraction data. The undeca- hedral species may be considered as frustrated nonaco-ordinate structures in that close approach of the ninth halogen is inhibited by the steric hindrance of the other eight halogen atoms. Table 6 Undecahedral octaco-ordinate structures Compound Re$ * Compound YF3a 224 PLIB~,~ SmF 224 TbC136 EuFB 224 NbBr GdFs 224 SmBr TbFs 224 EuBr HoFS 224 AmBrs ErFs 224 mF3 224 PrIs DYF3 224 P-NPBr TmF 224 CeI BiF 224 f UI Th,SUb g NPI ThI,C 2236 PUI m 3 LuF 224 NdIs Ref.* 222 221 222 222 h 222 222 222 h h h 222 222 222 222 W3y-F = 2.3 A; b8ThII-S = 2.94 A (See also Table 10); C8Th-I = 3.20 A (See also Table 4); dSPu-Br = 3.08 A; e2Tb-C1 = 2-74,4Tb-CI = 2-79 2Tb-CI = 2.95 A; * For references see Text except as follows fW. H. Zachariasen U.S. Atomic Energy Commission Argonne National Laboratory Report ANL-4400 (January 1950); O W . H. Zachariasen Acta Cryst. 1949,2,288 ; hF. H. Spedding and A. H. Daane Ames Laboratory Report IS-350 (September 1961). 921 J. D. Forrester A. Zalkin D. H. Templeton and J. C. Wallmann Inurg. Chem. 1964 3 185. 222 W. H. Zachariasen Acfu Cryst. 1948 1,265. 150 Muetterties and Wright Thorium tetraiodide has been described in terms of a distorted square anti- prism.223a Because of the close similarity of the antiprism and the undecahedron the thorium tetraiodide structure can equally well be described as a distorted unde~ahedron.,~,~ The yttrium atom in yttrium t r i f l ~ o r i d e ~ ~ ~ has essentially the same geometry about it as does terbium in terbium trichloride except that the ninth distant halogen atom is closer in YF than in TbCl,.This is essentially a compromise between an octa- and a nonaco-ordinate structure. The yttrium-fluorine dis- tances are 2.3 A with the ‘distant’ ninth fluorine atom at 2.6 A. Isostructural with YF are the orthorhombic lanthanide trifluorides samarium to lutetium inclusive. In a second (hexagonal) lanthanide fluoride structure formed by lanthanum to europium inclusive and also holmium and thulium there is nonaco-ordination by fluorine about the metal atom.cf) Cube. As noted earlier a co-ordination cube has not been observed for any molecular metal complex. It is found extensively in ionic lattices exemplified by the czsium chloride and calcium fluoride structures. (8) Metal polyhedra. Hexadecacarbonyl hexarhodium has a basic octahedral array225 of rhodium atoms each attached to two terminal carbonyl groups with an interpenetrating tetrahedron of bridging carbonyl groups (Figure 27). The co-ordination sphere about each rhodium atom is a hybrid of the square anti- prism and the dodecahedron. Four rhodium atoms describe a square face and four carbonyl groups form a near-trapezoidal array about a single rhodium atom. ? s C Fig. 27 Structure of Rh,(CO)l, (ref.225). Fig. 28 Srructure of[C,H,.CII.C,H JzFe,(CO)8 in the black stable crystal. See Figure 14for the isomeric structure (ref. 80). In the black stable form of [C6H5C,C6H5],Fe3(CO), one iron atom is un- ambiguously octaco-ordinate.80 This unique iron atom is bonded to two terminal carbonyl groups two bridging carbonyl groups two iron and two carbon atoms. The other two iron atoms may be described as hexa- or octaco-ordinate depend- 223 (a) A. Zalkin J. D. Forrester and D. H. Templeton Znorg. Chem. 1965 3 639; (b) D. H. Templeton personal communication. 424 A. Zalkin and D. H. Templeton J . Amer. Chem. SOC. 1953 75 2453. 2p5 E. R. Corey L. F. Dahl and W. Beck J. Amer. Chem. SOC. 1963 85 1202. 151 Quarter Iy Reviews ing upon the definition of co-ordination number for the olefin-iron interaction (Figure 28).The apical iron atom in Fes(CO),,C is octaco-ordinate.sl This iron atom has an unique geometry with three terminal carbonyl groups in half the co-ordina- tion sphere and a square pyramidal Fe4C arrangement in the other half (Figure 15). The remaining four iron atoms are heptaco-ordinate. Crystalline PtCl contains discrete Pt&1,2 units in which there is an octa- hedron of platinum atoms and nearly coplanar with each platinum atom are four chlorine atoms which in sum total describe the twelve edges of a cube.147 Each platinum atom is thus bonded to four chlorine atoms and four platinum atoms (Figure 29). An analogous structure is apparently present226 in the solu- tion state for Bi,(OH)126+. There may be octaco-ordination in [Re(CO),H] with an Re triangle bridged by hydrogen U Fig.29 Structure of Pt,Cl, (ref. 147). Fig. 30 The structure of discrete Zr4(OH)8(OH2)1,8+ in ZrOCI2,8H20. The eight oxygen atoms oj'the hydroxyl groups describe the basic cube. Each cube face is capped with a zirconium atom which is bonded to four oxygen atoms from water molecules as illustrated for one face in the diagram (refs. 205 206). Eight oxygen atoms form a square antiprism about each uranium atom in [U604(OH)J12+6S042-. The six uranium atoms describe an octahedron with a uranium-uranium distance of 3-84 A and the U-0 separation is 2-3 A.212 If the uranium atoms were considered to be within bonding distance each uranium atom would be dodecaco-ordinate. This structure is based on a two-dimensional X-ray analysis and may be subject to significant refinement or change.The zirconyl halide octahydrates mentioned above as having square antiprismatic co-ordination about the zirconium atoms exist in the solid state and apparently also in solution as discrete Zr4(0H),(OH~,,s+,8C1- groups (Figure 30).205*20s The zirconium atoms are in a square configuration with bridging hydroxyl groups and each zirconium atom has four terminal water groups. Apparently bound by coulombic forces the eight halogen atoms occupy specific sites about the polyhedron. Octaco-ordinate metal atoms are found in the eicosahedral metallo-boranes. For example rhenium is octaco-ordinate in B,C2HllRe(CO),-. The rhenium atom completes the carborane eicosahedron and is bonded to three boron and a26 V. A. Maroni and T. G. Spiro J. Amer. Chem.SOC. 1966,88 1410. (a) D. K. Huggins W. Fellmann J. M. Smith and H. D. Kaesz J. Amer. Chem. SOC. 1964 86 4841; (6) A. Zalkin T . E. Hopkins and D. H. Templeton Inorg. Chem. 1966 5 1189. 1 52 Muetterties and Wright two carbon atoms in the adjacent pentagonal plane. Three carbonyl groups are terminally bonded to the rhenium atom.227b 3 Solution State.-There are no measurements that provide any significant information concerning the geometry of octaco-ordinate complexes in solution. With the recent interest in laser materials a number of publications have ap- peared claiming structural definition of tetrakis-chelates of rare-earth ions a very extensive class of octaco-ordinate compounds by analysis of electronic or fluorescence ~ p e ~ t r a . ~ ~ ~ ~ ~ ~ - ~ ~ ~ Although the conclusions may be correct the analyses possess little rigour.Unfortunately there is no physical technique that will provide sufficient data for an unequivocal structural determination for such complex molecules in solution. For special cases a technique such as nuclear magnetic resonance (n.m.r.) may provide information regarding non-equivalence that will unambiguously eliminate some of the possible stereoisomers from con- sideration; however in all cases examined so far only spectroscopic equivalence of magnetic nuclei in the ligands has been observed. These include Mo(CN),4- (13C),29 ReH,[P(C,H&] (lH),loS and a number of chelate struc- t u r e ~ . ~ ~ ~ ~ ~ ~ ~ Intermolecular exchange can be rigorously excluded as the origin of the equivalence for the fist two examples and for a few of the chelate structures.The octacyanometallate ions are a particularly interesting case for solution studies. It has been argued on the basis of Raman and infrared data that in solution Mo(CN):- has D4d square antiprism geometry rather than the D2d dodecahedra1 geometry established for the solid potassium salt.233 The optical spectrum has been interpreted in terms of D4d as well as D 2 d symmetry but these analyses are far from definiti~e.~%-~~' As noted above the cyano-complex in solution shows a single sharp 13C n.m.r. resonance which is consistent withD,d symmetry but attention is called to the potentially short ground-state lifetimes of an octaco-ordinate structure which can disqualify any physical observation of a relatively long time scale for the determination of point-group ~ymmetry.2~ The closely analogous paramagnetic ions Mo(CN),3- and W(CN)t- have been examined by electron spin resonance t e c h n i q ~ e s .~ ~ ~ ~ ~ ~ Analysis of spectra at 25" in aqueous solution and at -196" in frozen glycerol solution indicates a similarity in structure in these two states and the results are consistent only with a D4d ground state (with the reasonable assumption that configurational exchange at -196" is slow relative to the e.s.r. time scale).237 The spectra of 228 J. Blanc and D. L. Ross J. Chem. Phys. 1965 43 1286. 23s H. Bauer J. Blanc and D. L. Ross J. Amer. Chem. SOC. 1964 86 5125. 230 N. J. Rose and E. Abramson J . Chem. Phys. 1965 42 1849. aslT. J. Pinnavaia and R. C. Fay Inorg. Chem. 1966,5 233. 232 A. C. Adams and E. M. Larsen Inorg. Chem.1966 5 228. a33 H. Stammreich and 0. Sala Z. Efektrochem. 1960 64 741 ; 1961 65 149. 234 R. M. Golding and A. Carrington Mol. Phys. 1962,S 37 I . 235 E. Konig Theor. Chim. Acta. 1962 1 23. 236 G. Gliemann Theor. Chim. Acta 1962 1 14. 237 B. R. McGarvey Inorg. Chem. 1966,5,476. 238 R. G. Hayes J . Chem. Phys. 1966,44,2210. 153 Quarterly Reviews dilute concentrations of K,Mo(CN) in &MO(CN) are consistent with D2d symmetry.237 This is the only well-defined example of a high-co-ordinate struc- ture which has different ground-state geometries in the solid and the solution state. The nature of octaco-ordinate species in solution is not well characterised in terms of solvation phenomena dissociation and ligand exchange. These extremely important considerations are probably inter-related and should certainly be more extensively examined.There is good evidence that lanthanide and actinide tetrakis-chelates interact significantly with strong donor molecules such as amides and sulphoxides. For example the electronic spectrum of the tetrakis(acetylacetonato)europium(m) anion in alcohol is grossly altered on addition of dimethylformamide suggestive of complex f0rmati0n.l~~ Variational concentration studies indicate that the complex is highly dissociated in solu- t i ~ n . l ~ ~ An analogous chelate anion derived from dibenzoylmethane reacts with dimethylformamide to give a crystalline 1 1 add~ct.l,~ Tetrakis(tropo1ono)- thorium(1v) forms 1 1 crystalline complexes with donor molecules like dimethyl sulphoxide.l* On the other hand the related neutral tropolone derivatives of zirconium and hafnium(1v) give no evidence of complex formation; and this is true also for cationic species such as tetrakis(tropolono)tantalum(v)+.18 Clearly steric factors must control formation of nonaco-ordinate complexes significantly.Ligand dissociation in octaco-ordinate species is generally a relatively easy process although there is one notable exception the non-labile Mo(CN):- ion. The stability of the cyanides probably reflects the much more effective delocalisa- tion of charge possible with cyano-groups; this is a critical point for these highly charged (-3 to -5) anionic complexes. In contrast to the cyanides the octa- fluorometallate ions are completely or nearly completely dissociated in aqueous media e.g. there is no evidence of the TaF,& ion in HF or NH,F-HF solu- t i o n ~ .~ ~ ~ In chelates many of the acetylacetonate complexes exhibit tendencies to dissociate in solution and this is quite well documented for the anionic tetrakis- /hikctone derivatives of europium;136 one or two molecules of solvent un- doubtedly take up co-ordination sites in the chelate x Solvent + E~(acac)~- + acac- + Eu(acac),(solvent). In the tropolone chelates stability of the tetrakis-derivatives is a function of the formal charge on the complex and the size of the metal ion.lo7 Here the cationic derivatives are more stable than the neutral ones which are in turn more stable than the anionic complexes. The hydrolysis rates of the tetrakis(tropo1onates) of niobium(v) and tantalum(v) are strikingly different although the ionic radii of these metal congeners differ only slightly (Ta5+ = 0.73 A; Nb5+ = 0.70 A).In neutral solutions the tantalum chelate is not significantly hydrolysed but the niobium derivative rapidly hydrolyses with the separation of tris(tropo1ono)- oxyniobium(v).107~239 The theimodynamic stabilities of the niobium and tantalum tropolonates may be more similar than is apparent for the case of neutral solu- E. L. Muetterties and C. M. Wright J . Amer. Chem. Soc. 1965 87 21. 1 54 Muetterties and Wright tions because of the very low solubility of T3Nb0 a species that has no analogue in the tantalum system. Quantitative soholytic data for octaco-ordinate species are lacking except for a series of europium acetylacetonates. The degree of dissociation of the tetrakis- anions in alcohol is 24 37 43 and lOO% respectively for the acetylacetone benzoylacetone dibenzoylmethane and benzoyltrifluoroacetone derivatives.13s A cation effect upon dissociation has been observed for tetrakis(benzoy1tri- fluoroacetonato)europium and is ascribed to ion-pair formation with the benzoyltrifluoroacetonato-anion.240 Lability of ligands in octaco-ordinate structures has been demonstrated in a number of systems.For example dissolution of tetrakis(acety1acetonato)- zirconium(1v) and tetrakis(trifluoroacetylacetonato)zirconium(Iv) in benzene yields a solution in which all possible isomers are The mixed complexes Zr(tfac),(acac) Zr(tfac),(acac), and Zr(tfac)(acac) are favoured at the expense of Zr(tfac) and Z r ( a ~ a c ) ? ~ ~ * ~ ~ The deviation from a purely statis- tical distribution of ligands is ascribed to entropy changes; enthalpy changes are nearly zero.A similar exchange is observed in the analogous Hf Ce and Th systems. Exchange rate is lowest in the hafnium case and highest for thorium. All of the fluorine n.m.r. resonances in the mixed thorium system coalesce at ca. 43"; lifetimes are less than a second at such temperatures. It is interesting that the thorium chelates are the most labile. Thorium unlike Zr Hf and Ce shows a tendency to be nonaco-ordinate i.e. the tetrakis-chelates tend to solvate in strongly basic media. The exchange studies were however done in non-basic media such as benzene. Perhaps nonaco-ordinate binuclear species like (14) are important intermediates for the thorium system in benzene whereas unassisted ligand dissociation may prevail with the smaller metal chelates.Rate of exchange of tropolone with tetrakis(tropolono)tantalum(v)f cation is estimated to be less than lo3 sec.-l from n.m.r. data; tracer studies have shown the exchange is 94 "/ complete in 30 minutes.lo7 Rate of ligand exchange increases in going to neutral and again in going to anionic tetrakis(tropo1ono)metallates. This ligand lability makes the outlook for isolation of optical isomers in octaco- ordinate chelates rather p00r?99107 Nuclear magnetic resonance studies of octaco-ordinate p-diketone or tropolone chelates have as yet failed to discern the environmental non-equival- ence of ligand atoms (e.g. methyl groups in acetylacetonates) expected for D (square antiprism) or D2 (dodecahedral) ~ymmetry.2~~ No evidence of geo- metrical isomers was obtained by n.m.r.studies of the tetrakis(trifluoroacety1ace- tonato) derivatives of Zr Hf Ce and Th even at temperatures as low as - 105". z40 E. P. Riedel and R. G. Charles J. Appl. Phys. 1965 36 3954. 155 Quarterly Reviews These studies point to but do not define the stereochemical non-rigidity of octaco-ordinate structures.231 Some of the chelates particularly those based on tropolone and related oxo- bidentate ligands are readily degraded by hydroxide ion. Interestingly as shown by tracer studies in the case of cationic tetrakis(tropolonates) a significant part of the hydrolysis comprises hydroxide ion attack of the chelate structure at a ligand po~ition.2~~ There have been extensive syntheses and physical studies of actinide com- plexes and this enormous area of chemistry cannot be treated in detail in this Review.A review of the co-ordination chemistry of the actinides by C ~ m p ~ and a spectroscopic review by Rabinovitch and B e l f ~ r d ~ ~ ~ provide considerable detail (584 references). In any case there is a large number of actinide complexes particularly uranyl complexes of the type UO,(chel),- that are undoubtedly Table 7 Possible octaco-ordinate uranyl compounds Compound Ligand Ref. ~ ~ ~ 3 ~ ~ 6 H 7 ~ 3 ~ 3 1 3 Acetylacetonate a *M+UOz(CSH70&- Acetylacetonate b c UOZ(C~H$OZF~)%,~HZO (4,4,4-Trifluoro-1-(2-thienyl)-l,3- d Hfruo,(c,o~lloz)31- 3-Tsopropyltropolonate UO2(CpH6NOSO3)3- 5-Sulphonato-8-quinolinolate * M+UOz(HOC 8H4C00)3- Salicylate i NH4UOZ(C 6H 5N202)3 N-Nitrosophenylhydroxylamine ion j KU~Z(CSH~OSZ)~ Ethyl xanthate k KU02(C4H7OS2)3 Tsopropyl xanthate k K[UOP(C~H~~N~,),~,HZO Diethyl dithiocarbamate k rUOz(Cl4Hlo~2N)Pl5 Benzoylnicotinoylmethane ion I U O Z ( C ~ ~ H I ~ ~ Z N ) ~ ( N ~ ~ ) ~ Benzoylnicotinoylmethane ion I UO,(C,,H,z0,),,2.5HzO Dibenzoylmethane ion m U ~ ~ ( N ~ ~ ) Z ~ C S H ~ N O N-Meth ylacet anilide P P ~ O ~ ( N O ~ ) ~ ~ ( C ~ H S O ) ~ O Tributyl phosphate 4 butanedionate] *M+[UOz(CpH,N0)31- 8-Quinolinolate g U0,~(C7Hl,0),POOl,[(C,H,,o),PooHl~ 2-Ethylhexylphosphoric acid h UO,(NO3),,2CH3CN n UOZ(N~~)Z,~CSHCTN 0 HP02(C9H6NO)3 8-Quinolinolate r *M = K NH4 CHaNH, CsHsNH3 H =A.E. Comyns B. M. Gatehouse and E. Wait J . Chern. SOC. 1958,4655; bA. E. Comyns ref. 241 p. 125 132; eK. Hager 2. unorg. Chem. 1927 162 82; dA. E. Comyns ref. 241 p. 125; eD. Dyrssen Acra Chem. Scund. 1956,10 353;fC.F. Richard R. L. Gustafson and A. E. Martell J . Amer. Chem. SOC. 1959 81 1033; eE. P. Bullwinkel and P. Noble jun. ibid. 1958 80 2955; hC. F. Baes jun. R. A. Zingaro and C. F. Coleman J . Phys. Chem. 1958 62 129; {J. T. Barr and C. A. Horton J . Amer. Chem. SOC. 1952 74 4430; 5W. S. Horton J. Amer. Chem. SOC. 1956,78 897; "R. A. Zingaro ibid. p. 3568; ZL. Sacconi and G. Giannoni J . Chem. SOC. 1954 2368; mL. Sacconi G. Caroti and P. Paoletti ibid. 1958 4257; nG. I. Kobyshev and D. N. Suglobov Proc. Acad. Sci. U.S.S.R. Sec. Phys. Chem. 1958 120 325; oJ. T. Barr and C. A. Horton J . Amer. Chem. SOC. 1952 74 4430; PR. Rascanu Ann. Sci. Univ. Jussy 1931,17,70; Chem. Abs. 1933,27,2647; QG. F. Best H. A. C. McKay and P. R. Woodgate J . Znorg. Nuclear Chem. 1957 4 315; rB. G. Harvey H. G.Heal A. G. Maddock and E. L. Rowley J. Chem. SOC. 1947 1010. 341 A. E. Comyns Chem. Rev. 1960,60 115. 943 E. Rabinovitch and R. L. Belford 'Spectroscopy and Photochemistry of Uranyl Com- pounds' Macmillan New York 1964. 156 Muetterties and Wright octaco-ordinate although the data are not definitive. A representative list of these complexes is in Table 7. Possible octaco-ordinate structures are listed in Tables 8 and 9. Table 8 Possible octaco-ordinate compounds with unidentate ligands Compound * ZrCI4,4RNH2 ZrC14,4NH ZrC1,,8NH3 TaC13[NHR],,3NH2R (R = CH, C2Hd PbCI4,4RNH (R = CH n-GH,) TeF,,2N( CHa La13,8DMF Ce(C104),,8DMA Pr(ClO&,,8DMA Pr13,8DMF Nd(C104),,8DMA Nd13,8DMF Sm13,8DMF (R = CH3 CGHJ La( ClO,&,8DMA Gdl,,8DMF ThCl4,4ROH (R = C2H6 n-C4H,) ThC14,4CH3.COC6H6 ThBr4,2CH3COOC2H6 ThBr4,4C6H,CH0 ThBr4,4CH,CN ThCl4,4RNH (R = CH, C,H& Th14,4CH3CN ThC14,4DMA ThT4,4DMA Th(SCN),,4DMA UC1,,4n-C3H,.NH UC14,6N2H2 UC1,,4NMA U12CI,,5DMA U13Cl,5DMA wC14,2.5DMA] [U(NOLJ~,~-~DMAI~ rITh(N03)4,2.5DMAI2 UId,4DMA &At C C C 118 C d 288 e 288 288 e 288 e e 293 f g h h h d i j k 264 k k I 1 k 264 264 264 289 264 6 157 Quarterly Reviews Table Qontinued Compound* Ref3 U(SCN)4,4DMA rn K4Mo(CN)7,2H20 n K3M0(CN)8 237,238 K3W(CN)8 237,238 KS[M~(OH)~(~N)~I,~H~O 0 K4[M o(OH)4(CN)d,4H@ P K3W(OH)4(CN)J Q KdMo(OH)dCN)5(NO)] 0 KdMo(CN)5(OH),] t MOC14,4(C,H&ASO P 4 r [Mo(CN),(CH,CN),(H20)~,4H20 8 Na,MoF 179 u K3MoF 180 K3WF8 179 u [WCO)&I3 227 v Rb,WF 53 K2Re(CN) 8 W K3Re(CN)8 W Ag,[Re(CN),NOI X ReH,l?(C,H5)3l3 106 K,RuF 178 c1(C0)[(C6H5)3P130sHD2a 142 [(c6H,)4As12[c0(02ccFd~ 16 Pd(N03) Y Am03 z Cs2TeF 182 Cs2XeF8 386 K4Th(NCS)8 aa lTh(OCHdJ bb K,PaF 181 IN(CH,)J,PaC1 cc Rb4UWCS) dd K,U(CNS)JH,O dd K4U(CNS)8 rn Na2UFSb 183 Li,UF 181 K2UCIg ee * NMA = N-Methylacetamide; DMA = A"'-dimethylacetamide; DMF = "'-dimethyl- formamide.QSuggested intermediate. *From an analysis of X-ray data; UF co-ordination (U-F = 2.29 A) was established but the geometry is as yet undefined. t For references see Text except as follows CJ. M. Matthews J . Amer. Chem. SOC. 1898 20 815; dE. L. Muetterties and W. D. Phillips ibid. 1957 79 2975; CT. Moeller and V. Galasyn J . Znorg. Nuclear Chem. l960,12,259;fD. C. Bradley M. A. Saad and W. Wardlaw 1 58 Muetterties and Wright J. Chem. SOC. 1954 109; gG. Jantsch and W. Urbach Helv. Chim. Acta 1919 2 490; hR.C. Young J. Amer. Chem. SOC. 1934 56 29; tK. W. Bagnall D. Brown P. J. Jones and J. G. H. du Preeze J. Chem. SOC. 1965 350; jT. Moeller and D. S. Smith U.S. Report Air Force Office of Scientific Research TN-58-559 (1958); kK. W. Bagnall D. Brown P. J. Jones and P. S. Robinson J. Chem. SOC. 1964 2531; 11. Kalnins and G. Gibson J. Inorg. Nuclear Chem. 1958,7 5 5 ; "K. W. Bagnall D. Brown and R. Colton J. Chem. SOC. 1964 2527; nP. C. H. Mitchell and R. J. P. Williams ibid. 1965 4570; OW. P. Griffith J. Lewis and G. Wilkinson ibid. 1959 872; PW. R. Bucknall and W. Wardlaw ibid. 1927 2981 ; qK. N. Mikhalevich and V. M. Litvinchuk J. Znorg. Chem. (U.S.S.R.) 1964,9,1293; X. N. Mikhalevich and V. M. Litvinchuk ibid. 1959 4 800; 8F. Holzl Monatsh. 1927 48 689; tS. M. Horner and S .Y. Tyree Znorg. Chem. 1962 1 947; S. UKatz Zbid. 1964 3 1598; uJ. M. Smith. W. Fellmann and L. H. Jones ibid. 1965,4,1361; WR. Colton R. D. Peacock and G. Wilkinson Nature 1958,182,393; %R. L. Colton R. D. Peacock and G. Wilkinson J. Chem. SOC. 1960 1374; YC. C. Addison and B. G. Ward Chem. Comm. 1966,155; "P. F. Lindley and P. Woodward J. Chem. SOC. ( A ) 1966 123;aaY. Y. Kharitonov A. K. Molodkin and A. V. Babaeva Bull. Acad. Sci. (U.S.S.R.) Div. Chem. Sci. 1964 No. 4 578; bbD. C. Bradley Nature 1958 182 121 1 ; C C D . Brown Colloques Internationaux du Centre National de la Recherche Scientifique Orsay France No. 154 (1965); d d ( a ) V. P. Markov and E. N. Traggeim Zhur. neorg. Khim. 1961 6 2316; (b) V. I. Belova Y. K. Syrkin and E. N. Traggeim Zhur. neorg. Khim. 1964 9,2673; eeK.W. Bagnall D. Brown and J. G. H. du Preeze J. Chem. SOC. 1964,2603. Table 9 Possible octaco-ordinate compounds with multidentate ligands *ZrQ- *Zr(edta)(PDS)2- *Zr(edta)(L)a- bZr(edta)(acac)2- bZr(edta)(L)2- bZr(L)(L')a- Zr(tftac) Zr(tfa~)~ bZr(aca~)~-,(tfac)~ Hf( t fac) *Hf(aca~),-~(tfac)~ T4Ta+ N~(~C~C)~(OC~HD)B TANb+ Ta(acac)3(OC2H6)!4 Ligand (L) Ref. d e Dimethyldithiocarbamate ion 290 107 Dimethyldithiocarbamate ion 290 f f NN-Dihydroxyethylglycinate f N-Hydroxyethyliminodiacetic f Diethylenetriaminepenta-acetic f acid ion acid ion 3,ddisulphonate ion g Disodium 1,8-dihydroxynaphthaIene- g g Citrate ion g N-Hydroxyethylethylenediamine- g h triacetic acid ion oxalate ion 23 1 231.232 231 231,232 i 107 i 107 Dimethyldithiocarbamate ion 290 Oxalate ion i k 1 Oxalate ion m Oxalate ion n 73 0 159 Quarterly Reviews RE(acac),- Na+RE(T,)- b[RE(L)(L')ls- RE(CH,CO,),(o-phen) RE(N03)3(o-~hen)2 bPm(L),- bPm(L),- bE~(L)4- Eu(acac) ,2 H20 ELI( tfac),- Eu(tftac),- Eu( tftac),(o-phen) 'Eu(L)~- Eu(L),- Eu(L),,2pyridine Gd( t ft ac),- Gd(L),- Tb(tftac),- Tb(tftac),(o phen) Tb(L)4- 'M( EDTA),SHZO "m(L),- bTmf Lh- EW)dPFd3 Eu(L)3,2H,O Th(02CC6H5)r bTh(L)4 Th(tftac) Ligand (L) Ref.P 73 73 15 4 r a-Hydroxyisobutyrate ion S t Glycolate S Lactate S t-Butylpenta-lY3-dione ion t 23 1 107 136 229 240 t u v 107 N-Hydroxyethylethylenediamine- triacetic acid ion iminodiacetic acid ion Glycolate Lactate Glycolate 4-Picoline N-oxide ion Hexafluoroacetylacetonate ion Lactate Dibenzoylmethane ion Dibenzoylmethane ion Dibenzoylmethane ion Dibenzoylmethane ion Glycolate a-Hydroxyisobutyrate ion or-Hydroxyisobutyrate ion Oxalate ion y-Isopropyltropolone ion Thiotropone ion 5 Sulphonato-8-quinolinolate ion Diethylenetriaminepenta-acetic acid ion Diethylenetriaminepenta-acetic acid ion trans- 1,2-Diarninocyclohexane- NNN"'-tetra-acetic acid ion l-(o-Arsonophenylazo)-2-naphthol- 3,6-disulphonic acid ion 132 W W S S S 135 136a 138 229 135 135 135 134 229 229 135 229 229 3 S S S X S 261 YY 284 Z aa 107 bb 18 cc 15 dd ee 3- 8- gg hh ii Muetterties and Wright Table 9-continued Comuounda Ligand (L) Ref.ThCj,(o-phen)p ii Pa( t f t a ~ ) ~ 285 kk kk U(==) U( t f a ~ ) ~ KbU(L)4,5H,O Oxalate ion U(edta),ZH,O Np( oxine)4 11 Pu(oxine) 11 Am(oxine) 11 if) UAbbreviations are as follows 1,lO-Phenanthroline (o-phen) disodium-l,2-dihydroxy- benzene-3,5-disulphonate ion (PDS) ethylenediaminetetra-acetic acid ion (edta) acetyl- acetonate ion (acac) 1,l ,l-trifluoropentane-2,4-dione ion (tfac) o-phenylenebisdimethylarsine (diars) phthalocyanine (pc) 8-quinolinolate ion (oxine) thenoyltrifluoroacetone ion (tftac) Rare Earth metal ion (RE) tropolone ion 0; bThese compounds were not isolated.These formulations are based on spectroscopic or potentiometric titration data; = Tb-Lu; * For references see Text except as follows dJ. D. Miller and R. H. Prince J. Chem. SOC. 1965,3185; eD. Dyrssen J. Inorg. Nuclear Chem. 1958,8,291 ;fB. I. IntorreandA. E. Martell J. Amer. Chem. SOC. 1960 82 358; UIdem ibid. 1961,83 3618; hR. E. Connick and W. H. McVey ibid. 1949 71 3182; ‘P. N. Kapoor and R. C. Mehrotra J. Less Common Metals 1965,8 339; jM. C. Steele Austral.J. Chem. 1957,10 368; kJ. E. Fergusson W. Kirkham and R. S. Nyholm ‘Rhenium’ Elsevier Amsterdam 1962 pp. 36-44; U. E. Fergusson and R. S . Nyholm Chem. and 2nd.. 1958 1555; mR. Charonnat Ann. Chem. 1931 16 186; nM. Platsch 2. anorg. Chem. 1899 20 308; OW. J. Kroenke and M. E. Kenney Inorg. Chem. 1964,3,251; PK. Ramaiah and D. F. Martin Chem. Comm. 1965 130; QL. Pokras and P. M. Bernays J. Amer. Chem. SOC. 1951,73 7; rL. Pokras M. Kilpatrick and P. M. Bemays ibid. 1953 75 1254; 8L. W. Holm G. R. Choppin and D. Moy J. Inorg. Nuclear Chem. 1961 19,251 ; tJ. S. Brinen F. Halverson and J. R. Leto J. Chem. Phys. 1965,42 4213; uM. L. Bhaumik ibid. 1964,40,3711; 1964,41,574; WG. H. Dieke H.M. Crosswhite and B. Dunn J. Opt. SOC. Amer. 1961 51 820; wF. A. Hart and F. P. Laming J. Inorg. Nuclear Chem.1965 27 1605; ZW. W. Wendlandt Analyt. Chem. 1957 29 800; VJ. P. Phillips J. F. Emery and H. P. Price ibid. 1952 24 1033; “(a) G. N. Wyrouboff and A. Verneuil Ann. Chim. phys. 1905,6,441; (6) F. A. Johnson and E. M. Larson Inorg. Chem. 1962 1 159; aaD. C. Madigan J. Appl. Chem. 1959 9 252; bbD. Dyrssen Acta Chem. Scand. 1955 9 1567; CcH. Iinuma J. Chem. SOC. Japan 1943 64 742; ddRef. f Table 7; eeG. H. Carey R. F. Bogucki and A. E. Martell Inorg. Chem. l964,3,1288;ffR. F. Bogucki and A. E. Martell J. Amer. Chem. SOC. 1958,80,4170; Sfl. H. Seu S. Wu and T . Chuang J. Inorg. Nuclear Chem. 1965 27 1655; hhH. Brintzinger H. Thiele and U. Muller Z. anorg. Chem. 1943 251 285; a‘A. R. Palmer Analyt Chim. Acta 1958 19 458; jjB. W. Fitzsimmons P. Gans B. C. Smith and M. A. Wassef Chem.and Ind. 1965 1698; kkH. Gilman R. G. Jones E. Bindschadler D. Blume G. Karmas G. A. Martin jun. J. F. Nobis J. R. Thirtle H. L. Yale and F. A. Yoeman J. Amer. Chem. SOC. 1956 78 2790; 43. H. Eberle and C. Keller Angew. Chern. Internat. Edn. 1965 4 971. D. Nonaco-ordination Nonaco-ordinate structures are fairly common in molecular complexes and in ionic lattices of the larger lanthanide and actinide ions. There are a number of idealised ground-state geometries for hepta- and octaco-ordinate structures but only one polyhedron has been reported for nonaco-ordinate molecular com- plexes. The polyhedron is the symmetrically tricapped trigonal prism (D3 J. (This polyhedron devoid of a central atom may be present inBiSS+.243) Interestingly this geometry is uniquely defmed by hybridisation of one s three p and five d orbitals.The only other reasonable polyhedron and one observed in metal cage compounds and in certain metal tellurides and arsenides is the monocapped a48 A. Hershaft and J. D. Corbett Inorg. Chem. 1963,2,979. 161 Quarterly Reviews square antiprism (C,,) which can be generated from the trigonal prism by rela- tively small distortions (Plates 7 and 8). The monocapped square antiprism should also be observed for some molecular nonaco-ordinate species in the solid state because packing forces may favour this less symmetrical polyhedron in some instances. Reorganisation energy for going from an octaco-ordinate square antiprism or dodecahedron to a nonaco-ordinate trigonal prism or monocapped square antiprism should be very small (see Plate 1).There are data that bear on the ease of distortion of nonaco-ordinate geo- metries. In the polyhedral B,H:- ion the nine boron atoms describe a Dsh tricapped trigonal prism in the solid state. For the solution state of B9H:- the llB n.m.r. data rigorously rule out the C, model and are wholly con- sistent with the D3h tricapped trigonal prism.244 Since the B polyhedron lacking a central atom is generated primarily by B-B attractive forces it is intrinsically a stereochemically more rigid structure than an MX co-ordination structure. Two of the most extraordinary nonaco-ordinate molecular species are TcH,2- and ReH,2-. Three-dimensional X-ray analysis and neutron diffraction studies*~ established the capped trigonal prismatic structure for ReH2- (Re-H = 1-68 A). This ion had previously been characterised as Re* then ReHZ- and then ReH,2-.The technetium ionlo is isostructural with ReH,2-. It will be interesting to see if this heavy-metal hydride chemistry extends to other metals in the transition group. Nonahydration is quite common for the tervalent rare-earth ions. Ketelaar245 first suspected a trigonal prismatic arrangement for M(OH&,3+ after an X-ray examination of the nonahydrated ethyl sulphates of yttrium and several rare earths including lanthanum and dysprosium. Later H e l m h ~ l z ~ ~ ~ in a now classic X-ray study established such a polyhedron for Nd(OH2),3+ in Nd(OH,),(BrO,),. The Nd-0 bond distances were 2.51 8 for the three central bonds and 2.47 for the outer six; the difference in bond distances was within experimental error. The trigonal prism of M(OH&?+ in the erbium yttrium and scandium ethyl sulphates [M(C,H,S0,)3,9H,0] was confirmed by Fitzwater and R ~ n d l e .~ ~ In the erbium compound the six M-O distances to the corners of the prism are 2-37 A and the remaining three are 2.52 A. Not all hydrated rare-earth ions are necessarily nonaco-ordinate; Hoard Lee and Lind3 point out that there is probably an equilibrium between M(OH,),3+ and M(OH2),3+ ions in solution and that the average hydration number is closer to nine for La3+ and probably closer to eight for the last member of the series Lu3+Fb which is significantly smaller (by 0.213 A). Lanthanide hydration numbers have not been rigorously established for the solution state. Some of the large bivalent metal cations are also nonahydrated at least in the crystalline state.Strontium chloride hexahydrate which is typical of the alkaline earth halide hydrates has a columnar structure (two-dimensional X-ray 244 F. Klanberg and E. L. Muetterties Inorg. Chem. 1966 5 1955. 246 J. A. A. Ketelaar Physica 1937 4 619. 246 L. Helmholz J . Amer. Chem. Soc. 1939 61 1544. a47 D. R. Fitzwater and R. E. Rundle Z. Krist. 1959 112 362. 162 Muetterties and Wright with strontium atoms coplanar with three water-oxygen atoms at 2.80 A and the strontium atom are connected through trigonally arranged bridging oxygen atoms of water molecules at 2.62 A. The co-ordination polyhe- dron is the tricapped trigonal prism. Lanthanide trihydroxides also have a trigonal prismatic grouping; in this case the nonaco-ordinate unit is M(OH),6-. The isostructural series includes the trihydroxides of La Pr Nd Sm Gd and Dy.Surprisingly the very small lanthanide ions E?+ and Ybw are also reported to be nonaco-ordinate in the hydroxide lattices.MQ Nonaco-ordination with symmetrically tricapped trigonal prismatic geometry is found in a large number of ionic lattices of the lanthanide and actinide as well as large bivalent metal halides. A partial list of these nonaco- ordinate structures is in Table 10. The rubidium ion in RbCdC1 is nonaco- ordinate.250 The capped trigonal prism is also found in complex metal fluorides. In hexagonal NaNdF, there are two Nd3+ sites and in each site the Ndw ions are co-ordinated to nine fluorine atoms with near D3h symmetry for the poly- h e d r ~ n . ~ ~ ~ One prism is smaller than the other; the average Nd-F distances are 2.377 A and 2.426 A for the twe sites.NaLaF and NaCeF are probably iso- structural.261 X-Ray data also indicate an isostructural relationship of NaCaMF (M = rare earth or yttrium tervalent ion) and NaYF to NaNdF,.251-253 The hexagonal rare-earth trifluoride lattices (Tysonite) have nine fluorine atoms within bonding distance of the metal atoms; the space-group is P’3 ~ 1 . ~ ~ The co- ordination polyhedron is quite irregular. Zachariasen presents arguments (intensity calculations or consideration of interatomic distances) for nonaco- ordination in the lattices of the following actinide halides UC13(U-9CI = 2.96 A) /3,-Na,ThF6(Th-9F = 2-41 A) p2-K2UF6 p2-Na2UF, Pl-K2UF (U-9F = 2.73 A) and U2F,(U-9F = 2.31 A).2229255 In K,PaF each protac- tinium atom is surrounded by nine fluorine atoms which approximate the vertex positions of a tricapped trigonal prism; the Pa-F distances range from 2.19- 2.46 A?5s Nonaco-ordination has also been established3 in a rare-earth chelate KLa(OH&[(OCCH&2NCH2CH2N(CH2CO)J,5H20.(Isomorphous with this complex are the potassium salts of the La Nd and Gd chelates the sodium salts of La Nd Tb Gd and Er and the ammonium salts of Nd and Gd.3 The co-ordination polyhedron in the terbium complex is isostructural with the lan- thanum complex. The distances are Tb-N 2.665 A Tb-0 2.377 A and Tb-Owater 2.481 The co-ordination polyhedron is defined by the six 248 A. T. Jensen Kgl. danske Videnskab. Selskab Mat.-fys. Medd. 1940,17 No. 9 1. 249 K. Schubert and A. Seitz Z. anorg. Chem. 1947 254 116. 193. 251 J. H. Burns Inorg. Chem.1965 4. 881. a52 B. P. Sobolev D. A. Mineev and V. P. Pashutin Kristallografiya 1963 8 545. 253 A. A. Voronkov N. G. Shumyatskaya and Yu. A. Pyatenko J. Struct. Chem. (U.S.S.R.) 1962 3 665. 254 A. Zalkin D. H. Templeton and T. E. Hopkins Inorg. Chem. 1966 5 1466. 255 W. H. Zachariasen Acta Cryst. 1949 2 390. 256 D. Brown and A. J. Smith Chem. Comm. 1965 554. C. H. MacGuillavry H. Nijveld S. Dierdorf and J. Karsten Rec. Trav. chim. 1939,58 163 Table 10 Non-molecular structures with nonaco-ordinate metal atoms Co-ordination Co-ordinationa 5 Class Compound group polyhedron Bond distance (A) UCls M(z)@ TCTP UCIs 9U-C1= 2-96 MC18 MBr M(OH)s MFs (M = La-Gd Ac U Np Pu Am) (M = La Ce Pry Ac U a-Np) (M = La Pry Nd Sm Gd Dy Er) (M = La-Nd Sm Eu Ho Tm Ac ma MOO TCTP Pbcl TCXP U-Am) PbCls 9Pb-CI = 2.8-3.5 Pb(O€l)Cl Pb(OH)4(C1) 5 2Pb-OR = 2.67,2Pb-OH = 2.93 PbBr SPb-Cl = 3.23 PbFa BaXt' EuCl SmCl ThSt 9Th-S = 2.95 ThSe US BiOCl BiOBr 4Bi-0 = 2.31 4Bi-Cl = 3.07 4Bi-0 = 2.32.4Bi-Br = 3.18. Bi-Cl = 3.49 NaCa(RE)Fl* NaNdF NaMF SrC1,,6H20 SrCI,,GH,O TCTP TCTP TCTP P 8 Ref.* 5 222 i3 249 s 222 222 e 224 164 f g h h h h h i k 4 4 i Bi-Br = 3-92 6(RE,Ca-F) = 2.39 3(RE,Ca-F) = 2.30253 6Nd-F = 2.426,3Nd-F = 2.377 25 1 (M = La-Nd Eu-Er Y Sm Tm Dy) 251 6Er-0 = 237 3Er-0 = 2-52 247 6Y-0 = 2.37 3Y-0 = 2.55 247 6Pr-O = 2.47 3Pr-0 = 2.65 247 M = La-Sm Gd Dy 245 Sr-0 = 2.62 Sr-0 = 2.80 248 252 Table 10 Non-molecular structures with nonaco-ordinate metal atoms Co-ordination Co-ordinationa 5 Class Compound group polyhedron Bond distance (A) UCls M(z)@ TCTP UCIs 9U-C1= 2-96 MC18 MBr M(OH)s MFs (M = La-Gd Ac U Np Pu Am) (M = La Ce Pry Ac U a-Np) (M = La Pry Nd Sm Gd Dy Er) (M = La-Nd Sm Eu Ho Tm Ac ma MOO TCTP Pbcl TCXP U-Am) PbCls 9Pb-CI = 2.8-3.5 Pb(O€l)Cl Pb(OH)4(C1) 5 2Pb-OR = 2.67,2Pb-OH = 2.93 PbBr SPb-Cl = 3.23 PbFa BaXt' EuCl SmCl ThSt 9Th-S = 2.95 ThSe US BiOCl BiOBr 4Bi-0 = 2.31 4Bi-Cl = 3.07 4Bi-0 = 2.32.4Bi-Br = 3.18. Bi-Cl = 3.49 NaCa(RE)Fl* NaNdF NaMF SrC1,,6H20 SrCI,,GH,O TCTP TCTP TCTP P 8 Ref.* 5 222 i3 249 s 222 222 e 224 164 f g h h h h h i k 4 4 i Bi-Br = 3-92 6(RE,Ca-F) = 2.39 3(RE,Ca-F) = 2.30253 6Nd-F = 2.426,3Nd-F = 2.377 25 1 (M = La-Nd Eu-Er Y Sm Tm Dy) 251 6Er-0 = 237 3Er-0 = 2-52 247 6Y-0 = 2.37 3Y-0 = 2.55 247 6Pr-O = 2.47 3Pr-0 = 2.65 247 M = La-Sm Gd Dy 245 Sr-0 = 2.62 Sr-0 = 2.80 248 252 Table 10 Non-molecular structures with nonaco-ordinate metal atoms Co-ordination Co-ordinationa 5 Class Compound group polyhedron Bond distance (A) UCls M(z)@ TCTP UCIs 9U-C1= 2-96 MC18 MBr M(OH)s MFs (M = La-Gd Ac U Np Pu Am) (M = La Ce Pry Ac U a-Np) (M = La Pry Nd Sm Gd Dy Er) (M = La-Nd Sm Eu Ho Tm Ac ma MOO TCTP Pbcl TCXP U-Am) PbCls 9Pb-CI = 2.8-3.5 Pb(O€l)Cl Pb(OH)4(C1) 5 2Pb-OR = 2.67,2Pb-OH = 2.93 PbBr SPb-Cl = 3.23 PbFa BaXt' EuCl SmCl ThSt 9Th-S = 2.95 ThSe US BiOCl BiOBr 4Bi-0 = 2.31 4Bi-Cl = 3.07 4Bi-0 = 2.32.4Bi-Br = 3.18. Bi-Cl = 3.49 NaCa(RE)Fl* NaNdF NaMF SrC1,,6H20 SrCI,,GH,O TCTP TCTP TCTP P 8 Ref.* 5 222 i3 249 s 222 222 e 224 164 f g h h h h h i k 4 4 i Bi-Br = 3-92 6(RE,Ca-F) = 2.39 3(RE,Ca-F) = 2.30253 6Nd-F = 2.426,3Nd-F = 2.377 25 1 (M = La-Nd Eu-Er Y Sm Tm Dy) 251 6Er-0 = 237 3Er-0 = 2-52 247 6Y-0 = 2.37 3Y-0 = 2.55 247 6Pr-O = 2.47 3Pr-0 = 2.65 247 M = La-Sm Gd Dy 245 Sr-0 = 2.62 Sr-0 = 2.80 248 252 OI Table 10-continued Class BaCIz,H20 Nd(Br03),9Ha0 PbFCl Compound SrXz b,6H20 CaXsb,6HZO Ba12,6H20 BaC12,Hz0 Nd(Br03),9H20 NdCla99H2O Co-ordination Co-ordinationa polyhedron Bond distance (A) group 7Ba-Cl = 3-25,2Ba-OH = 2-76 Ba(C1),(OH2)2 MOI20)0 TCTP 6Nd-0 = 2-47 3Nd-0 = 2.51 TCTP 9K-F = 2.73 9Th-F = 2.38 9K-F = 2.73 9U-F = 2.36 9U-F = 2.31 9Th-F = 2-40 6ThI-S = 3.07 3ThI-S = 2.82 TCTP m1(s)6(s)3 M(Z)4(Z')5 CSPA Pb(F),(Cl),Cl Th(O),(S)5 4Th-0 = 2.40 5Th-S = 3.00 Ref.* 248 248 248 t 246 246 5 t t 293 U 222 255 255 40 38 222,255 222 V 222,255 c; 2 255 9' -Y 2 182 1 1 m 1 - w 8 Table 10-continued class Compound uos NpOS AcOCl AcOBr PUOCl PuOBr YOCl YOBr LaOCl LaOBr SmOCl HoOCl REOClb REOBrb Co-ordination Co-ordination" group polyhedron Bond distance (A) 4NpO = 232,SNp-S = 2.91 4U-0 = 2.34 5U-S = 2.93 4La-O = 2-39 4La-Cl = 3-18 4La-O = 2.40 4La-Br = 3.28 4Sm-0 = 2.30,4Sm-C1 = 3.11 4Ho-0 = 2.25,4Ho-C1 = 3.05 L a 4 = 3-14 La-Br = 3.49 Sm-Cl = 3.09 Ho-Cl = 3.04 Ref.* ki i 2 i 2 164 s 164 164 164 164 n 0 0 P P 164 n 164 n "Tricapped trigonal prism (TCTP) monocapped square antiprism (CSPA); bX = C1 Br I; CRare Earth metal ion (RE); dPossibly iso- morphous with ,B,-Na,ThF,; * For references see Text except as follows eD. H. Templeton and C. H. Dauben J. Amer. Chem. SOC. 1954 76 5237; fH. Braekken 2. Krist. 1932 83 222; BH. Brasseur Bull. SOC. Roy. Sci. LiZge 1940,9,166; hW. Doll and W. Klemm Z. anorg. Chem. 1939,241,239; tW. H. Zachariasen Actu Cryst.l949,2,291;jR. W. M. D'Eye P. G. Sellman and J. R. Murray J. Chem. SOC. 1952 2555; kE. D. Eastman L. Brewer L. A. Bromley P. W. Gilles and N. L. Lofgren J. Amer. Chem. SOC. 1950 72 4019 ZW. H. Zachariasen Actu Cryst. 1949 2 288; mW. Nieuwenkamp and J. M. Bijvoet 2. Krist. 1932 81 469; "1. Mayer S. Zolotov and F. Kassierer Inorg. Chem. 1965 4 1637; OL. G. Sille'n and A. L. Nylander Svensk kem. Tidskr. 1941,53 367; PD. H. Templeton and C. H. Dauben J. Amer. Chem. SOC. 1953 75 6069; gL. G. Sill6n Svensk kern. Tidskr. 1941,53,39; rB. K. Veinstein and Z. G. Pinsker Zhur. fiz. Khim. 1949 23 1058; SR. Fitzwater and R. E. Rundle U.S. Atomic Energy Commission Report ISC-241 (1952); W. I. Iverovna V. P. Tarasova and M. M. Umanskii Izvest. Akad. Nauk S.S.S.R. Ser. Fiz. 1951 15 164; *H.Lambot Bull. SOC. Roy. Sci. Lie'ge 1943 12 439; VL. A. Harris Actu Cryst. 1960 13 502. Muetterties and Wright donor atoms of the ethylenediaminetetra-acetate ion and three water oxygen atoms. Because of the spatial limitations imposed by the ligand the geometry does not closely approximate the trigonal prism; Figure 3 1 illustrates geometry and relevant bond distances. Hoard Lind and Lee3b suggest that in this anionic lanthanum complex decaco-ordination will prevail in solution and that decaco- ordination will extend to analogous complexes of Ce3+ to Gd3+ inclusive. Bohigian and M~te11~~' suggest that the triethylenetetraminehexa-acetic acid derivative of thorium(rv) is either a nona- or deca-co-ordinate complex. There is a tris- (tetrahydrofuran) solvate of tris(dipyridy1)yttrium that might be a nonaco- ordinate structure.258 Fig.31 Structure of lanthanum(rrr) trihydrate ethylenediaminetetra-acetate complex in KLa(ethylenediaminetetra-acefate),(OH~,,5H,O (ref. 3). It has been suggested that in sodium guanidinium trisulphatothorium hexa- hydrate the thorium is nona~o-ordinate~~~ with three water molecules at the square pyramidal vertices and six sulphato-groups (which bridge co-ordination polyhedra) at the trigonal vertices of a tricapped trigonal prism. In K2Th(S04)4,2H20 the thorium is surrounded by two water oxygen atoms and seven sulphato-oxygen atoms of which two come from the same sulphato- Octaco-ordinate thorium chelates clearly show acceptor character and there are specific reports of a dimethyl sulphoxide adduct of tetrakis(tropo1ono)- thorium1* and an acetic acid adduct of tetrakis(thenoyltrifluoroaceto)thorium.261 There is also a series of thorium and uranium(rv) halide or perchlorate complexes with donor molecules such as amides but there is no definitive datum as to the molecularity of these species.262-264 Thorium(1v) nitrate forms a complex with two tributylphosphine l i g a n d ~ .~ ~ ~ Since the nitrate groups appear to be bidentate this could be another example of decaco-ordinate thorium. Nonaco-ordination is also found in complex intermetallic structures. An group -259 260 257 T. A. Bohigian jun. and A. E. Martell Znorg. Chem. 1965,4 1264. 258 S. Herzog and K. Gustav Z. Naturforsch. 1962 17b 62. 25D A. K. Molodkin G. A. Skotnikova and E. G. Arutyunyan J . Znorg. Chem. (U.S.S.R.) 1964 9 1458.260 E. G. Arutyunyan and M. A Porai-Koshits J. Struct. Chem. (U.S.S.R.) to be published. 262 K. W. Bagnall D. Brown and A. M. Deane J. Chem. SOC. 1962 1655. 263 K. W. Bagnall A. M. Deane T. L. Markin P. S. Robinson and M. A. A. Stewart J. Chem. SOC. 1961 1611. 264 K. W. Bagnall D. Brown P. J. Jones and J. G. H. DuPreez J . Chem. SOC. 1965,3594. 265 J. R. Ferraro A. Walker and C. Cristallini Inorg. Nuclear Chem. Letters 1965 1 25. G. Goldstein 0. Menis and D. L. Manning Analyt. Chem. 1960 32 400. 167 Quarterly Re views example is NiBi which Zhdanov266 claims from analysis of diffraction data has an Ni(Ni,Bi,) co-ordination polyhedron approximating the tricapped trigonal prism. A preliminary report of the structure of NbS,Br and NbS,Cl indicates nonaco-ordination for the niobium atoms (Figure 32).267 There may be nonaco- ordinate thorium atoms in some of the thorium-sulphur (e.g.Th,S,,) and thorium-selenide phases. In LaTe and NdTe, each metal atom is surrounded by nine tellurium atoms which roughly describe a monocapped square anti- prism.2sg The La-Te distances in LaTe range from 3.26-3.38 A. Isostructural with the tellurides is Fe2As.269 Fig. 32 Niobium co-ordination polyhedron in NbS2C1 (ref. 267). Fig. 33 The Mo,,CI,~+ polyhedron in the Mo,CI,,~- anion; each molybdenum atom has a terminal chlorine atom within bonding distance (ref. 270). Two other kinds of nonaco-ordinate geometry are found in metal cages or clusters. These geometries are defined by the complex polyhedron of metal atoms and ligands. One is represented by the Mo6CI,G+ framework in which the metal atoms describe an octahedron and the chlorine atoms which bridge these metals describe a cube (Figure 33).270 To this framework six relatively labile ligands take terminal bonding positions to the metal atoms.These terminal ligands are of the type halide ions (as in Mo6Cl,2- and MO6ClgBr;-) hydroxide ion water and organic donor molecules such as pyridine and triethy- lamine.271-253 In these polyhedra the metal atom is within bonding distance of 3 8 0 Cl Mo Cl cr 288 G. S. Zhdanov Academy Sciences (U.S.S.R.) 1954 10 99; Chem. Abs. 1956 SO 2234d. 287 H. G. Schnering W. Beckmann and H. Schafer see ref 21 p. 842. 268 R. Wang H. Steinfink and W. F. Bradley Inorg. Chem. 1965,5 142. 269 A. F. Wells ‘Structural Inorganic Chemistry’ 3rd edn. Oxford University Press London 1961 p.520. 270 C. Brosset Arkiv Kemi. Min. Geol. 1946 22 A 1. 271 J. C. Sheldon J . Chem. SOC. 1960 3106. 272 J. C. Sheldon J. Chem. SOC. 1960 1007. 279 C. Brosset Arkiv Kemi 1949 1 3 5 3 . 168 Muetterties and Wright four other metal atoms four cage halide ions and one terminal ligand to give a monocapped square antiprismatic co-ordination polyhedron as in (1 5 ) with a ninth site directed up out of the plane of the paper from the central position. This polyhedron devoid of a central atom may prevail in B,CI,H or B,C1,H2-.274 Other metal (e.g. tungsten)-halogen systems may have analogous cage struc- tures. Moreover the curious (n-C6H6)6Mn6(NO)B is isoelectronic with MO,CI,,~- and probably has Mn6 octahedra with a cubic array of bridging 3 NO groups.276* The second class of metal cage structure with nonaco-ordinate metal atoms is typified by the Ta6C1,;+ structure in which the metal atoms again form an octahedron (edges are - 2.9 A) but the halogen atoms are at the mid-edges of a cube (Figure 34).276 Again a labile ligand site is probably tangential to the cube Fig.34 The Ta6Cl12a+ polyhedron. In solution each tantalum atom is probably terminally attached to a water molecule (ref. 276). Four of the chlorine atoms are omitted for clarity. faces as in Ta6C1,,(OH2)~+ and Nb,CI12(OH2)62+. The co-ordination poly- hedron for the metal atoms is a monocapped cube (16). This type of structure is found in (Nb6Ts)I,, (the %,I polyhedron is distorted from Oh to Ci sym- metry through shifts of two apical niobium atoms away from the line normal to the plane of the four other niobium atoms277) (Nb6CI12)C12 ("a6112)12 Ta6Br12(OH2)62+ va6c112)cI, (Ta6Br12)Br2~ fl@r12)Br3 and (Ta6Br12) Recent X-ray and molecular weight data by F.A. Cotton et. al. show a composition 274 J. A. Forstner T. E. Haas and E. L. Muetterties Inorg. Chem. 1964 3 155. 276 R. B. King and M. B. Bisnette Inorg. Chem. 1964 3 791. 276 P. A. Vaughan J. H. Sturdivant and L. Pauling J. Amer. Chem. Soc. 1950 72 5477. L. R. Bateman J. F. Blount and L. F. Dahl J. Amer. Chem. Soc. 1966 88 1082. (n-C,HJSMn,(NO),. 1 69 Quarterly Reviews Br6?1,277-280 The tantalum and niobium polyhedral cations appear to undergo two-electron oxidations.281 The reported Ir6114 may be isostructural with (Ta611& and contain nonaco- ordinate iridium atoms.282 E. Decaco-ordination Decaco-ordination in molecular complexes is found or seriously suspected only for lanthanide or actinide-metal complexes.This may be fortuitous but only for these metal ions can$-orbital overlap be reasonably considered and a valence- bond representation of decaco-ordination would utilise f orbitals. On the other hand a multicentre-bond molecular orbital approximation would obviate invocation of f-orbital participation. Actually f-orbital overlap in decaco- ordinate complexes is necessarily real on symmetry grounds alone; however the magnitude of the overlap integral is at this stage simply a matter of conjecture. Size of the metal ion as well as compactness of ligands is obviously important if a high-co-ordination polyhedron is to be generated. Accordingly the earlier members of the lanthanide and actinide series are the most favourable candidates for decaco-ordination because of the severe ion contractions with increasing numbers off electrons.This is illustrated in Figure 1 in which ionic radii are plotted against atomic number; a few other large metal ions particularly the alkaline earth bivalent ions are included for reference. High oxidation state also favours maximum co-ordination number but this must often be compromised because of the tremendous decrease in ionic radius with ionic charge e.g. compare Ce4+ with Ces+ in Figure 1. Decaco-ordination has been rigorously established for only one complex. A three-dimensional X-ray analysis4 of La[(02CCH2),NCH2~CH2N(CH,C0,H)- CH2C0],4H,0 shows two nitrogen atoms four carboxylic-oxygen atoms and four water-oxygen atoms within bonding distance of the lanthanum atom.Geometry is defined primarily by the constraints of the sexadentate ligand. The lanthanum atom is slightly above the six donor atoms of the chelate leaving adequate room for co-ordination of the four water molecules. The structure is presented in Figure 35 with pertinent bond distances. This lanthanum example provides no information about preferred ground-state geometry for decaco- ordination because of distortion arising from the chelate moiety. A decaco- ordinate complex with all ligands identical is required and none is known. Possible models for such polyhedra are the symmetrically bicapped square antiprism (D4 d) (established for the B,,H1,2- polyhedron which has no central atom) and a bicapped dodecahedron (C,) which are generated from the two common polyhedra for octaco-ordinate structures and are of proper symmetry 278 R.E. McCarley and J. C. Boatman Znorg. Chem. 1965,4 1486. 279 P. J. Kuhn and R. E. McCarley Znorg. Chem. 1965 4 1482. 280 A. Simon H. G. Schnering H. Wohrle and H. Schiifer Z. anorg. Chem. 1965 339 155. R. E. McCarley B. G. Hughes F. A. Cotton and R. Zimmerman Inorg. Chern. 1965 4 1491. 282 S. F. A. Kettle Nature 1965 207 1384. 1 70 Muetterties and Wright Fig. 35 Structure of the decaco-ordinate lanthanum atom in La[(O,CCHz),NCH,CHz*N- (CH&OZH)CH,COe],4Ho0 (ref. 4). for an sp3d5fhybrid model. The disparities and similarities of these two poly- hedra are reflected in Plates 9 and 10. Both polyhedra are easily generated from the tricapped trigonal prism the idealised model for nonaco-ordination.Re- organisation energy for either conversion should be small. Professor Hoard in a personal communication has noted that the D4d bicapped square antiprism is not an unreasonable model on steric grounds but it is definitely less favourable than the D3h nonaco-ordinate and Ih dodecaco-ordinate models. With a bond distance of 2.40 A the eight pyramidal edges and eight zig-zag edges are 2.60 A and the square edges are long 3.09 A. The radius-ratio is 0-848 which is two- thirds of the way from the value of 0-732 for the D3h nonaco-ordinate model and 0.905 for the eicosahedron. In a pentakis-bidentate chelate such as the Th(trop01onate)~- ion described below the preferred stereoisomers for the chelate based on a bicapped square antiprism model and the 2-58 A ‘bite’ of the tropolone ion would have no square edges bridged by a tropolone ion.The tropolone anion (4) forms decaco-ordinate complexes18 with Th4+ and U4+. Stability of the anionic complexes M(0,C7H5)5- is only modest outside of the crystalline state of the alkali-metal salts. The anionic complex dissociates in polar media but not in non-polar solvents. Even in the solid state packing considerations loom significant and crystalline salts are obtained only with the smaller alkali-metal ions. This is most apparent for U(02C,H,),- where only lithium salts have been isolated. Multivalent ions of do configuration e.g. Hf4+ and Taw do not complex tropolone ion beyond the tetrakis stage. Another apparent decaco-ordinate chelate is the thorium derivative of tri- ethylenetetraminehexa-acetic a ~ i d . 2 ~ ~ Some of the octaco-ordinate thorium chelates tend to be complexed with solvent e.g.the bis-acetic acid complex of tetrakis(thenoyltrifluoroaceto)thorium(Iv),261 and these may be decaco-ordinate. Additionally some of the tetrakis-bidentate chelates of thorium separate from solution with an extra molecule of the parent ligand in the protonated form. This is particularly characteristic of the chelates derived from 8-hydroxyquinoline and its derivatives. The additional molecule of 8-hydroxyquinoline can be removed at high temperature in a These complexes may be decaco-ordinate in the solid state; however they appear to dissociate in solution but here the solvent may be displacing 8-hydroxyquinoline to give either a p83 T. Moeller and M. V. Ramaniah I . Amer. Chem. SOC. 1954,76 2022.2a4 W. W . Wendlandt Analyt. Chem. 1956,28,499. 171 Quarterly Reviews nona- or deca-co-ordinate species f.e. Th(~hel)~(solvent) or a. There is also a report of a pentakis(thenoy1trifluoroacetone) derivative of protactinium(~v).~*~ Salts of the pentakis(carbonato)thorium ion Th(C03),g- and of MO(CO,),~- have been isolated and these may have five chelating carbonato-groups arrayed about a central metal atom.286,287 There are many reports of lanthanide and actinide complexes that might be decaco-ordinate at least in the solid state but no definitive structural data are available. For example there is a large number of amide and phosphine oxide complexes of the type La(C104)3,8CH,CON(CH3)2 Er(C104),,7- and Th14,6CH,CON(CH3)2.262-2M~288~289 The actinide complexes have been shown to ionise in polar organic solvents like acetonitrile and nitromethane; degree of ionisation is often a function of concentration and of time.There is a preliminary report of a curious pentakis(dimethyldithiocarbonamato)tantalum(v) compound but no structural or spectral data were presented.290 It has been reported that in the presence of hexamethylenetetramine the salts of many bivalent ions tend to separate from aqueous solution with ten molecules of water. The formation of these complexes e.g. BaC12,[(CH~,N~,,10H20 has been ascribed to stabilisation by hexamethylenetetramine of high-hydration forms of the inorganic salts. An X-ray examination291 of the calcium salt how- ever established that the calcium ion is simply octahedrally co-ordinated to six oxygen (water) atoms. CH,*CON(CHJ, U(C104)4,6CH3CON(CH& U(NOJ4,2.5CH3CON(CH& F.Undecaco-ordination Eleven-atom polyhedra are not to our knowledge very common in inorganic chemistry and this probably reflects the fact that there are no particularly good idealised models for an eleven-atom polyhedron. The lower stabilities of idealised undecaco-ordinate polyhedra relative to the more favourable polyhedra for deca- and dodeca-co-ordination contribute to the dearth of eleven-atom co- ordination complexes. Recently Ueki Zalkin and T e m p l e t ~ n ~ ~ ~ ~ concluded from a three-dimensional X-ray analysis of Th(NO3),,5H2O that the crystals are orthorhombic and that the thorium atom is within bonding distance of eleven oxygen atoms three from water molecules and eight from four nitrate groups which function as bidentate ligands.The light-atom positions were also charac- terised in a neutron diffraction study;292b the average bond distances are 2-46 for thorium to oxygen in the bound water molecules and 2.57 for the thorium- 285 G. Bouissitres and J. Vernois Compt. rend. 1957 244 2508. x86 I. I. Chernyaev V. A. Golovnya and A. K. Molodkin J. Inorg. Chem. (U.S.S.R.) 1958 3 100. 287 M. C. Steele Austral. J. Chem. 1957 10 367. 289 K. W. Bagnall P. S. Robinson and M. A. A. Stewart J. Chem. Suc. 1961,4060. ego S. E. Rasmussen and N. C. Broch Chem. Comm. 1965 289. 291 P. DeSantis A. L. Kovacs A. M. Liquori and L. Mazzarella J. Amer. Chem. SOC. 1965 87 4965. 292 (a) T. Ueki A. Zalkin and D. H. Templeton Actu Cryst. 1966 20 836; (b) J. C. Taylor M. H. Mueller and R. L. Hitterman ibid. p. 842.T. Moeller and G. Vicenti J. Inorg. Nuclear Chem. 1965 27 1477. 1 72 Muetterties and Wright oxygen distances in the nitrate interactions. The polyhedron itself cannot readily be described as any particular polyhedron but Taylor Mueller and H i t t e m ~ a n ~ ~ ~ ~ point out that if the nitrate interactions are considered as single entities rather than bidentate functions the geometry approximates the NbF,2- mono-capped trigonal prism (C2w). The low symmetry (C& of ligand atoms about the thorium atom may reflect distortions due to packing effects or to the very short ‘bites’ of the nitrate group. Multidentate ligands can impose serious distortions on an idealised geometry in co-ordination chemistry Eleven-atom polyhedra are found extensively among boron hydrides. Examples of these are B,C2HI1 and its derivatives wherein nine boron atoms and two carbon atoms comprise the polyhedron with exopolyhedral BH bonds.The crystal structure of this ‘carborane’ series has been determined for the bis-dimethyl (C substitution) derivative (personal communication from Professor Hawthorne) and in Plate 11 the B,C2 polyhedral array is compared with a stick model of the thorium eleven-atom co-ordination complex. The B,,HIl2- ion of Klanberg and Muet- terties may be isostructural with the B9C2 carboranes. As pointed out above when the possible constraints of the bidentate nitrate ligands are considered the distortion of the thorium complex from the C, geometry observed in the BgC carborane is really not very severe. The only other idealised geometry that might be seriously considered in eleven-atom metal co-ordination compounds would be derived directly from the eicosahedron by removing one of the eicosa- hedral ligands leaving an open-faced structure of C, symmetry.This particular geometry has not been rigorously established either in metal co-ordination chemistry or in boron chemistry. However a modified version of it might be observed in a periodic lattice if packing arrangements were such that the eicosa- hedron could not be completed owing to strong non-bonding repulsions from large or bulky ligands and that the twelfth ligand position would be absent or else there might be another ligand at a very long distance on the unique fivefold axis. G. Dodeca- and Higher Co-ordination The eicosahedron (Figure 36) is established for the co-ordination sphere of a lanthanum cation complex; twelve sulphate oxygen atoms are within 2.60- 173 Quarterly Reviews 2.74 A of type I lanthanum ions in the crystal lattice of La,(S0,)3,9H,0.293u The eicosahedron is slightly distorted possibly owing to packing factors.Because the other smaller lanthanide ions failed to crystallise in this form it was sug- gested that lanthanum(1u) is just barely able to support the co-ordination of twelve oxygen atoms. There is a hexasulphatothorium ion Th(S04)6*- isolated in the form of salts that may be analogous to the lanthanum There are reports of dodecaco-ordination in other structures and these also involve bidentate ligands which have very short ‘bites’. Especially noteworthy are nitrate and oxalate complexes. Cerium in Ce2Mg,(NO3),,,24H2O is sur- rounded by twelve oxygen atoms at an average distance of 2.64 A.294 These oxygen atoms belong to six nitrate ions and are at the vertices of a slightly irregular eicosahedron.The oxygen-nitrogen-oxygen bond angle of the bidentate nitrate function is smaller than in free nitrate ion; apparently some distortion of the nitrate is necessary to generate a stable eicosahedral co-ordination poly- hedron about cerium(II1). In a polyhedron such as Ce(NO3),& the bridging nitrate groups are symmetrically arranged as the handles in the Argentine ‘pato ball’ (see Figure 37). This arrangement appears to be the most favourable for Fig. 37 The structure (Th) of Ce(N03)ss- in Ce,Mg3(NO.&,24H,O (ref. 294). any eicosahedron with six bidentate ligands. The idealised geometry of a discrete polyhedron of this type has the unusual Th point-group symmetry.In Mg(H2O),Th(N0,),,2H2O there is an eicosahedral arrangement of six nitrate groups about thorium with near Th symmetry for the Th(NO,),Z- polyhedron.295 The Th-0 separation is 2-63 A. There is an eicosahedral arrangement of six water and six perchlorate oxygen atoms about the barium ion in Ba(C10,),,3HZO with Ba-0 distances ranging from 2.96-3.18 A similar arrangement is found in Ba(N03)2296,297 with six bidentate nitrate ions arranged as in ce(NoJt+; and in BaSiF and BaGeF each barium atom is nearly equidistant from twelve fluorine atoms from eight 193(a) E. B. Hunt jun. R. E. Rundle and A. J. Stosick Acta Cryst. 1954 7 106; (b) A. Rosenheim V. Samter and J. Davidsohn Z . anorg. Chem. 1903,35,424. 294 A. Zalkin J.D. Forester and D. H. Templeton J. Chem. Phys. 1963 39 2881. 295 S. Scavnicar and B. Prodic 1965 Acta Cryst. 18 698. z96 N. V. Mani and S. Ramasesham Z. Krist. 1960 114 200. 297 F. M. Jaeger and F. A. van Melle Proc. Acad. Sci. (Amsterdam) 1928.31 651. 174 Muetterties and Wright neighbouring hexafluorometallate ani0ns,2~* which form a nearly regular eicosahedron. Dodecaco-ordination thus appears to be limited to the almost trivial cases where the 'chelating ligand' has a very short bite as in the nitrate ion. Alternatively such polyhedra as La(NO,),& can be considered as octahedral with a preferred LaW-NO3- orientation in which La-0 interactions are maximal (or in which the La-N separation is minimised). Inorganic oxides which are based on close-packed XO layers tend to have the heavy X atoms surrounded by twelve oxygen atoms as in the BaNi0,,299 perov- skite,30° and hexagonal BaTi0Zo1 structures.In the idealised perovskite the twelve oxygen atoms described a cube octahedron (Figure 38) about the heavy atom. The hydrogen-atom positions in 6-uranium hydride UH, have been deter- mined by neutron diffractionFo2 There are two types of uranium atoms. One is surrounded by twelve hydrogen atoms at the corners of an eicosahedron of Th symmetry (faces composed of isosceles triangles). The second type of uranium atom is also dodeca~o-ordinate~~~ but the hydrogen atoms are in sets of three; the configuration approximates the truncated tetrahedron (Figure 39). Fig. 38 The cube octahedron. Fig. 39 The truncated tetrahedron. Very high co-ordination numbers are commonly found in close-packed structures such as the heavy-metal silicides and in complex alloy structures particularly those derived from the heavy metals.Co-ordination numbers of ten twelve fourteen fifteen and sixteen have been reported. Decaco-ordination is found in three structural classes of heavy metal silicide TiSi,,304 CrSi,,305 and MoSi,Fo6 The isomorphous series ThSi,,307 LaSi2,308 CeSi,,309 a-USi, NpSi, e9* J. L. Hoard and W. B. Vincent J. Amer. Chem. Soc. 1940 62,3126. 299 J. J. Lander Acta Cryst. 1951 4 148. 300 R. S. Roth J. Res. Nut. Bur. Stand. 1957 58 75. 301 R. D. Burbank and H. T. Evans jun. Acta Cryst. 1948 1 330. 302 R. E. Rundle J. Amer. Chem. Soc. 1951 73 4172. 303 R. N. R. Mulford F. H. Ellinger and W. H. Zachariasen J. Amer. Chem. Sue. 1954,76 297. 304 F.Laves and H. J. Wallbaum 2. Krist. 1939 101 78. 305 B. BorCn Arkiv Kemi Min. Geol. 1933 11 A No. 10. 306 W. H. Zachariasen Z. phys. Chem. 1927 128 39. 307 G. Brauer and A. Mitius 2. anorg. Chem. 1942 249 325. 308 F. Bertaut and P. Blum Acta Cryst. 1950,3 319. 309 W. H. Zachariasen Acta Cryst. 1949 2 94. 175 Quarterly Reviews PuSi, and the uranium silicides U3Si and U3Si:O7 contain dodecaco-ordinate metal atoms. Tetradecaco-ordination is found in the metal-rich silicide Mo3Si305 where molybdenum has two molybdenum four silicon and eight molybdenum atoms at distances of 2-44 A 2-73 A and 2.99 A respectively. The data are not sufficiently precise to define the geometry about the metal or silicon atoms. Further even the matter of co-ordination number becomes a bit of a semantic point in these close-packed structures.Frank and Kasper310 define the co- ordination sphere of an atom as encompassing all other atoms whose centres are nearer the atom in question than any other; alternatively co-ordination number is the number of nearest neighbours. The former definition unam- biguously assigns twelve and fourteen to the co-ordination numbers for hexagonal close-packed metals and body-centered cubic metals respectively. In complex intermetallic phases the co-ordination number is fourteen although the nearest-neighbour rule sets the number as one or two. Four idealised co- ordination polyhedra for complex alloy structures are shown in Figures 36 and 40-42. All four have triangular faces a reflection of the fact that there is tetra- hedral grouping of atoms throughout these structures thus requiring triangulated co-ordination polyhedra.In contrast for cubic close packing there are square as well as triangular faces and the co-ordination polyhedron for twelve in this system is the cube octahedron (Figure 38). Fig. 40 Idealised geometry for tetradecaco-ordination. There is an example of eicosaco-ordinate uranium in uranium dib~ride.~ll Each metal atom in a position of D,h symmetry is close to twelve boron atoms at the vertices of a hexagonal prism and to an additional eight others through the faces of the B, prisms at a distance comparable with the metallic diameter. slo F. C . Frank and J. S . Kasper Acra Cryst. 1958,11 184. m1 B. W. Howlett J. Inst. Metals 1959 88 91. 176 A Muetterties and Wright Fig. 41 Idealised geometry for pentadecaco-ordination.Fig. 42 Idealised geometry for hexadecaco-ordination. H. Comments on Lanthanide and Actinide Chemistry There is an apparent complexity in the solid-state chemistry of lanthanide and actinide compounds. Compositions or stoicheiometries which are quite diverse provide few clues to the underlying structural principles. Moreover poly- morphism is very common. Especially noteworthy in this context is the fluoro- complex chemistry of uranium(@ and (IV). The stoicheiometries of the reactions of alkali-metal fluoride with uranium@) and uranium(1v) are quite varied. This complexity in lanthanide and actinide chemistry is explicable in the light of the 177 Quarterly Reviews structural considerations presented in the preceding sections. There is the pos- sibility of co-ordination numbers for the metal atom of six seven eight and nine as well as some intermediate cases that may be described as 6+ 7+ and 8 f (see Sm,O in Section B.2 and YF and TbCI in Section C.2e).If one uses a trigonal prism as the hypothetical base model then as shown in Figure 43 it is quite easy to go from 6+ to nonaco-ordination. All of the geometries illustrated in Figure 43 have been established in rare-earth compounds. The difference in 1 I AH:tok other idea I i zed geometries Fiz. 43 Hypothetical transformation mechanisms for changes in co-ordination number. energy levels for these various co-ordination models should not be gross and will be very sensitive to the radius-ratio of metal atom to ligand atom and to the polarising power and the size of other cations that may be present in the lattice.Moreover the barrier for polyhedral rearrangements to other idealised geo- metries (for any given co-ordination number) will not be very large. The differ- ences in energy levels and the barriers to changes in co-ordination number or geometry should be the same order of magnitude as foices generated by packing and ordering factors in the solid state. Thus the extraordinary range of stoicheio- metries in the MF-UF or UF systems may simply reflect very subtle differences in the actual arrangement of fluorine atoms about the uranium atom or in some cases to changes in co-ordination number. Polymorphism reflecting again changes in the co-ordination number or in the geometry of the polyhedron is also expected in the solid state because of the relatively small energy differences among the co-ordination polyhedra.The rare-earth oxides (see Section B) are an excellent example of the subtle structural changes that can occur. In the cubic C form the metal atoms are hexaco-ordinate with distorted octahedra of two types C, and C,, which share a vertex whereas in the hexagonal A form they are heptaco-ordinate with near C, symmetry (monocapped octahedron). There are three distinct co-ordination polyhedra in the monoclinic B form with two of these having heptaco-ordinate metal atoms and near C, symmetry (mono- capped trigonal prism) and one polyhedron that is essentially an octahedron although there is a seventh oxygen atom on a threefold axis at a non-bonding distance of 3.12 A. The two heptaco-ordinate polyhedra differ only in the dimen- sions of the capped t r i g o d prisms.178 Muetterties and Wright I. Suggested Numbering Rules for Co-ordination Polyhedra 1 Discrete Polyhedra.-To provide a consistent system for the specification of stereochemistry in all co-ordination polyhedra it is important that a universal numbering procedure be adopted. This is particularly critical for the description of metal co-ordination polyhedra that have multidentate ligands. Several types of convention can be visualised that would label identical vertices and identical edges would simply identify vertices or would label identical vertices and identify all vertices with a combination of letters and numbers. We have con- sidered in detail a number of alternative procedures and would like to suggest the following numbering rules.This suggestion conforms to previously adopted rules set forth for octahedral complexes,312 and for borane and carborane :a C4y-SOUARE PYRAMID 6 { 4 1 0' -OCTAHEDRON j !!- -7? 4 3 I uh-cueE 4 @ d - T ~ ~ ~ ~ ~ ~ ~ ANTIPRISM & -TRIGONAL PRISM $ 3 3 7 - 6 ,a SF[ SF[ 4 7 4 7 CS c s 4 MONOCAPPED OCTAHEDRON 3 CLV MONOCAPPED SQUARE BASE-TRIGONAL CAP c3v TRIGONAL PRISM 8fjj6 4 3 7 0,ySOUARE' ANTIPRISM 4 bH-8lCAPPED TRIGONAL PRISM pv -MONOCAPPED SOUARE ANTIPRISM b 6 h - ~ ~ ~ ~ ~ ~ ~ ~ ~ EIPYRAMID I 4 0 56 BICAPPEO TRIGONAL ANTIPRISM 8 Cpv-BICAPPED TRIGONAL PRISM 9 4 h - I-TRICAPPED TRIGONAL PRISM 31p W. C. Fernelius 'Advances in Chemistry' Series No. 8 Amer. Chem. SOC. 1951. 179 Quarterly Reviews 8 SQUARE ANTIPRISM I 7 C~V-MONOCAPPED PENTAGONAL ANTIPRISM 1 I~-ICOSAHEDRON 1 04 -OMNICAPPED CUBE 6 c2 I ( E s C ~ H I ~ STRUCTURE ) rd -TRUNCATED TETRAHEDRON oh-cueE OCTAHEDRON 6 1,@14 1 1 10 Og,j- S-EICAPPED H EXAGO I4 AL ANT I PR ISM 4 9 74 -DISTORTED ICOSAHEDRON; " PATOHEDRON 1 Ih -PENTAGONAL DODECAHEDRON Fig.44 Numbering convention for the more common co-ordination polyhedra. poIyhedra31S (this is not strictly true for the borane and carborane polyhedra in that the directions given by Adams recommend initiation of the numbering system at an apex position). Illustrations of the suggested numbering for poly- hedra are given in Figure 44 and a procedure for numbering any polyhedron is given below. Numbering Rules.-(1) Do not number the central atom in co-ordination com- plexes. (24 Select the proper axis of highest order (or one from a degenerate set of 313 R .Adams Inorg. Chem. 1963 2 1087. 180 Muetterties and Wright highest order). If there are one or two polyhedral vertices on the axis then begin numbering with one of these vertices. (If there are two vertices on the principal axis and if there is no plane of symmetry perpendicular to this principal axis begin the numbering at the vertex which is nearer to the largest number of other poly- hedral vertices This qualification is necessary for some of the less regular polyhedra.) If there are no vertices on the principal axis of symmetry follow directions in (3) below. If there is no symmetry axis follow directions in (4). (2b) Drop down from vertex number (1) until one or more other vertices lie in a plane ‘A’ perpendicular to the highest symmetry axis. (2c) Number the vertices in plane ‘A’ identified in (2b) in a clockwise fashion.(If the vertices are non-equivalent in this plane ‘A’ begin by numbering the first two vertices whose connecting line is parallel to the element generating non- equivalence of symmetry in plane ‘A’. This qualification is necessary for poly- hedra such as the C, monocapped trigonal prism. If there is still an ambiguity with the above qualification then select the element generating the asymmetry whose vertices are nearest plane ‘A’. This problem arises for poly- hedra such as the C, bicapped trigonal prism whose numbering is illustrated in Figure 44.) Continue down to the next plane ‘B’ of vertices and again number in a clockwise fashion. Pick up the numbering at a vertex that is in a plane ‘X’ which includes the vertex with the lowest number in the top plane ‘A’ and the principal symmetry axis identified in (24.If no vertex is in the intersection of planes ‘B’ and ‘X‘ then rotate plane ‘X’ about the principal axis in a clockwise fashion to the first polyhedral vertex then pick up the numbering at this vertex and continue in a clockwise fashion. Continue this procedure until all planes encompassing polyhedral vertices are numbered. The vertex of lowest number in plane ‘A’ remains the reference point throughout. (3) If no vertices lie on the axis of highest symmetry then go to the plane perpendicular to this axis and begin numbering as in (2b) and (24. [For poly- hedra which do not have a plane of symmetry perpendicular to the principal axis of symmetry begin numbering at the end which has the largest number of vertices in the first plane perpendicular to the principal axis.There remains an ambiguity as to the vertex which is identified as (1) in some truncated polyhedra e.g. the axially truncated square pyramid. In these cases start with the polygon of smaller dimensions.] (4) If the polyhedron possesses only a plane of symmetry select if possible a unique vertex and identify this as vertex (1). (It is very difficult to state a general unambiguous procedure for polyhedra of such low symmetry.) Number the remaining vertices in a fashion formally analogous to the procedures above. (See the numbering convention for the two square base-trigonal capped poly- hedra in Figure 44.) With these rules other polyhedra not illustrated in Figure 44 can be numbered in a consistent fashion. As noted above there are alternative systems for number- ing of certain polyhedra which would be specific to these polyhedra and would at the very least be aesthetically more pleasing.A case in point is the problem en- countered with our suggested system for all of the polyhedra based on the tri- 181 Quarterly Reviews gonal prism in particular the various capped trigonal prismatic polyhedra. In these cases the trigonal prismatic vertices have different numbers for all the capped derivative polyhedra. Alternatively the trigonal prismatic vertices might have been given a non-varying set of numbers one to six and the capped apices numbers seven to nine. However we recommend sacrifice of aesthetics in favour of a wholly consistent set of rules. Attention is called to the Hoard and Silverton labelling system6 for numbering the octaco-ordinate dodecahedron which is particularly attractive for this polyhedron in that like vertices and like edges are immediately apparent from the labelling procedure.However this procedure does not uniquely identify edges and vertices and does present some problems in chelate structures for uniquely specifying isomers. Our suggested numbering system provides ready identification of isomers in chelate structures. For example the possible eight-co-ordinate chelate (17) based on the thiatropone chelate ligand would be identified as 1,2 (S 0); 3 4 (S 0); 5 6 (0 S); 7 8 (S 0) square antiprism and the isomer (18) as 1 5 (0 S); 2,6 (S 0); 3 7 (S 0); 4 8 (S 0) square antiprism. 2 Linked and Fused Polyhedra.-Polyhedra which are linked at one or more sites should be numbered as they would be (1.1 above) if not linked.Numbers are employed as subscripts to distinguish polyhedra. If these polyhedra have a plane of symmetry perpendicular to the principal proper axis of symmetry the poly- hedra should be numbered so that the vertices which are linked have the lower numbers i.e. numbering should commence at the end at or nearest the linkage site. Examples for linked s-bicapped square antiprisms are (19) (20) and (21). This system is applicable to metal co-ordination polyhedra and to polyhedra lacking a central atom such as the polyhedral boranes and carboranes. If polyhedra are fused at a vertex edge or face and if these polyhedra lack a central atom (e.g. polyhedral boranes) the whole structure should be con- sidered as a new polyhedron subject to rules (1) to (4) in paragraph (1) above.Fused polyhedra which individually contain central metal atoms are members of the broad class of compounds commonly described as metal clusters. Because the metal atoms may be environmentally non-equivalent and additionally may actually be different nuclei the metal atoms in these condensed polyhedra must be numbered. With this modification of rule (1) in the procedure above the 1 82 M uetterties and Wright 9 6 9 71 suggested numbering procedure is applicable to metal clusters. However it may not be generally necessary to apply a numbering convention to metal clusters (unless an extensive chelate chemistry develops at the terminal or exopolyhedral positions). Words alone should provide a simple and straightforward descrip- tion for metal clusters.It is suggested that resolution of this issue be postponed until metal cluster chemistry is more highly elaborated. Appendix The following information relevant to the subject of high co-ordination structures has come to our attention since the submission of the original manuscript. Ligand.-The adaptability of the non-rigid chelating ligand acetylacetone is quite systematically documented in an article by Lingafelter and Braun3I4 in which interatomic distances and angles in published and some unpublished complexes with acetylacetone chelates have been compared. In complexes rang- ing from pentaco-ordinate to octaco-ordinate species the oxygen-metal-oxygen angle varies from about 70" up to 98". In the tetrakis-derivatives of zirconium(Iv) and cerium(1v) the bond angles are nearly the same 75" and 73" respectively.Metal-oxygen distances in these two specific tetrakis-complexes are about 2-20 A in zirconium and vary in the cerium case from 2.37-2.43 .$; the latter distances are based on two-dimensional data whereas the zirconium was deter- mined with three-dimensional data. 314 E. C. Lingafelter and R. L. Braun. J. Amer. Chern. Suc. 1966 88,2951. 189 Quarterly Re views Heptaco-ordination.-Antimony(m) and selenium(rv) have a potential for quasi- heptaco-ordinate geometry in the hexahalogeno-anions. However analysis of crystals of (NH,)4Sbr~rSbVBr,2 has shown that in this lattice the SbBr63- ions have strict Oh symmetry and are he~aco-ordinate.~~~ An infrared study of salts of TeC1,2- TeBre2- Te162- and the related selenium species was made and it was concluded that although the extra pair of electrons in the tellurium and selenium complexes does not distort the individual octahedra it does affect the position of the stretching and bending modes.These results however even in comparison with the related tin complexes are not definitive because they are based solely on the solid state either as Nujol mulls or as pressed disks.31g Bartlett et al. have presented a preliminary analysis of the crystal structure of [XeF,]+ (PtF6]-. This structure is presented as an ionic lattice with square pyramidal XeF,+ groups. However the xenon atom is almost within bonding distance of one of the fluorine atoms associated with the hexafluoroplatinate atom. If this is accepted as a bonding interaction an alternative description of the species is a xenon- fluorine-platinum bridge structure in which xenon then may have the quasi- heptaco-ordinate structure with a non-bonding pair of electrons residing in a directed orbital .317 Numerous new fluoro-complexes have been reported in recent months by a number of workers.These include such species as nitrosyl salts of transition- metal fluoro-complexes e.g. NOWF, (NO),WF, (NO),ReF, and N00sF,?l8 as well as alkali fluoride complexes of quinquevalent neptunium319 and similar complex fluorides of ter- and quadrivalent uranium.320 No structural information is available for the nitrosyl salts so the actual co-ordination number in the solid- state species has not been established although co-ordination numbers of at least seven are not unrealistic for these particular metals.In the case of the neptunium complex it has been shown that Rb,NpF is monoclinic and iso- structural with K2TaF, which is apparently heptaco-ordinate and presumably has the capped trigonal prismatic geometry. The new work on the ter- quadri- and quinque-valent uranium complexes has not progressed to a stage such that the actual co-ordination number has been definitely established but again co-ordination numbers of at least seven are reasonable for the relatively large U3+ U4+ and U5+ ions. In work by K a t ~ ~ ~ l it has been shown that the hexa- fluorides of uranium tungsten and molybdenum react with sodium fluoride to achieve stoicheio-metric ratios of 1 2 and 1 1 at high rates if the sodium fluoride has been preformed by decomposition of Na2UF8. Equilibrium pres- sures are reported.Definitive structural information is as yet unavailable. Potential heptaco-ordinate fluoride complexes of hafnium(1v) with NH,F 315 S. L. Lawton and R. A. Jacobson Inorg. Chem. 1966,5 743. 316 N. N. Greenwood and B. P. Straughan J. Chem. SOC. (A) 1966 962. 317 N. Bartlett F. Einstein D. F. Stewart and J. Trotter Chem. Comm. 1966 550. 318 N. Bartlett and S. P. Beaton Chem. Comm. 1966 167. 319 L. B. Asprey T. K. Keenan R. A. Penneman and G. D. Sturgeon Inorg. Nuclear Chem. Letters 1966 2 19. 320 R. E. Thoma H. A. Friedman and R. A. Penneman J . Amer. Chem. SOC. 1966,88,2046. 321 S. Katz Inorg. Chem. 1966 5 666. 184 Muet t ert ies and Wright HF and H20 have been reported?, A series of neptunium plutonium ameri- cium and curium complex fluorides have been prepared e.g.Na,Np,F,, Na7fi&1 Na,h,F,, and Na,Cm,F,,. The structures have not been estab- lished and co-ordination numbers may be anything from seven to nine.323 In NaNb6015F there is a periodic lattice which contains a pentagonal bipyramidal NbO,F group as a repeating structural unit which shares edges with a ring of octahedra. The mean metal-oxygen distance is 2-03 A. Isostructural with the fluoride complex is NaNb60,50H?24 Relative to the capped trigonal prism as an idealised model in heptaco- ordination RandiC and Mak~iC,~ have calculated the splittings of d orbitals for heptaco-ordinate complexes of the type NbP,2- using a point-charge model and find the results very sensitive to assumed geometrical angles of the molecule. In a review by J. L. Hoard326 on some aspects of haem stereochemistry structural data for pure iron porphyrins are presented and it is pointed out that a substantial displacement of the iron atom from the plane of the porphine nitrogen atoms is a normal structural property of all high-spin iron porphyrins and that significant alterations in the conformation about the co-ordination group are expected to accompany transition to the low-spin state.It has been suggested that heptaco-ordination exists about the iron atom of the haem and in oxymyoglobin in which the oxygen molecule is symmetrically attached to the iron. If this is the case Hoard has suggested that the capped trigonal prism is the more probable geometry but Hoard favours the standard octahedron as the model for this particular complex. In chelate chemistry Clark Greenfield and Nyh~lrn,~' have shown that terdentate arsines such as tris-l,l,l-(dimethylarsinomethy1)ethane react with the trihalides of titanium and vanadium to give monomeric species of 1 1 stoicheiometry which display no evidence of electrical conductivity in solution.Nuclear magnetic resonance suggests that all three arsine atoms of each ligand are co-ordinated to the metal and therefore it appears that the titanium and vanadium atoms are heptaco-ordinate in these complexes. There have been a number of recent developments in cluster chemistry in which the metal atoms are in fact heptaco-ordinate. Most of these are deriva- tives of the Re,CI,,* ion its analogues or derivatives of dodecacarbonyltri-iron. Cotton and L i ~ p a r d ~ ~ have prepared a rather intriguing cluster retaining the basic geometry of the Re,Br,,% ion by reaction of the bromide with silver arsenate followed by treatment with dimethyl sulphoxide.The composition of the compound is Re,Br3(As0,),[(CH3),SO],. The dimethyl sulphoxide groups probably occupy terminally co-ordinating sites in the plane of the three rhenium atoms with the two terdentate arsenate groups at terminal positions normal to 322 M. B. Gaudreau Compt. rend. 1966 263 67. 323 T. K. Keenan Inorg. Nuclear Chem. Letters 1966 2,211. 325 M. RandiC and Z. MaksiC Theor. Chim. Acta 1966 4 145. 326 J. L. Hoard to be published. 327 R. J. H. Clark M. L. Greenfield and R. S. Nyholm J. Chem. SOC. (A) 1966 1254. 328 F. A. Cotton and S. J. Lippard J. Amer. Chem. Soc. 1966,88 1882. S. Ar,derson Acta Chem. Scand. 1965 19 2285. 185 Quarterly Reviews the planes described by the three rhenium atoms.In two independent studies carbonyl analogues of the Re3Br,,3- structure have been obtained. These species have a composition [M(CO),SR] where M represents manganese or rhenium and R an organic group such as ethyl or phenyl. Structurally these complexes have not been defined but by analogy with the rhenium case it is presumed that the thiol groups serve as bridging groups and that the carbonyl groups are all terminally b0nded.3~~*~~* Stewart and O'Donnell~l suggested that the core of the Mo,C1,3+ cation is similar to that found in Re3C1,23- and Re3CI,. In this particular case unless molybdenum were solvated in solution molybdenum would only be hexaco-ordinate. A phosphine derivative of dodecacarbonyltri-iron has recently been described which has the same structure as the parent compound with the phosphine group at a terminal positio11.3~~ A crystal-structure determinationw of this phosphine derivative has established that there are two isomeric structures in the single crystal one with the phosphine terminally bonded to the unique hexaco-ordinate iron atom and one with the phosphine terminally bonded to a heptaco-ordinate iron atom (see Figure 12).Notably the two bridging carbonyl groups appear to be asymmetric in both isomers. A complete three-dimensional structural analysis of the complex oxide UO3,3Y,O has been completed and shows that the distribution of metal ions in the structure is not fixed by symmetry. Rhombohedra1 uranium atoms are octahedrally co-ordinated with all uranium-oxygen distances equal; however reduction of this rhombohedral phase followed by oxidation leads to a uranium co-ordination number of seven.The existence of a rhombohedral structure in the range of U02+z,Y203 compositions is apparently made possible by the partial substitution of uranium at yttrium positions. With the uranium atoms at the cell origin the co-ordination number increases to seven or eight with the extra U-0 bonds being perpendicular to opposite faces of the octahedron at a distance of 2.46 A. At temperatures above lOOO" the structure rearranges to the normal fluorite structure with octaco-ordinate (cubic polyhedron) uranium.= Octaco-ordination.-A review specific to eight-atom co-ordination polyhedra will appear shortly by S. J. Li~pard.3,~ Chelate chemistry still dominates most of the recent advances.Most significant of these is the discovery of the volatile CS+Y(CF~-CO-CHCOCF,)~- salt which sublimes in vacuo or in air at 180-230". Structural analysis by X-ray diffraction * More recently it has been established that :the manganese derivative is a tetramer not a trimer (A. Wojcicki and J. Lewis personal communication). 3z9 E. W. Abel and B. C. Crosse J . Chem. SOC. (A) 1966 1141. 330 A. G. Osborne and F. G. A. Stone J . Chem. SOC. ( A ) 1966 1143. 331 D. F. Stewart and T. A. O'Donnell Nature 1966 210 836. 332 R. J. Angelici and E. E. Siefert Znarg. Chem. 1966 5 1457. 333 D. J. Dahm and R. A. Jacobson Chem. Comm. 1966,496. 334 S . F. Bartram Inarg. Chem. 1966 5 749. 385 S. J. Lippard 'Progress in Inorganic Chemistry' ed. F. A. Cotton Interscience New York 1966.186 Muetterties and Wright is in progress.ssg The structure of calcium dipicolinate trihydrate has been deter- minedB7 by a three-dimensional X-ray structural analysis. The co-ordination poly- hedron is an irregular dodecahedron consisting of the planar terdentate dipicol- inate ligand (Ca-0 = 2-406 2.524 A; Ca-N = 2.491 A) a bridging ligand oxygen atom (Ca-0 = 2.363 A) and four water molecules (Ca-0 = 2-405-2.567 A). Two calcium dipicolinate units form a dimer by sharing a ligand oxygen. The co-ordination polyhedra are linked through the oxygen atom of a combined water molecule to give a three-dimensional network in the solid state. Sulphito- complexes of thorium(1v) have been described with unit-cell compositions of Na,8Th(S0,)l,,CO(NH2)2,1 8H20 in which the sulphite moiety may be bifunc- tional to give octa(or higher)co-ordinate thorium polyhedra.B8 Further work with the unique diarsine (o-phenylenebisdimethylarsine) has shown that TiF4and TiI do not behave like TiCI,.The fluoride gives a polymeric presumably fluorine bridged structure of the composition (TiF,),,diarsine in contrast to the molecular dodecahedral TiC142diarsine complex.161 Behaviour of the iodide also departs from that of the chloride and Clark et aZ.339 suggest an ionic formulation (TiI,,2diar~ine)~+21-. Qualitative observations of the reaction of diarsine with ZrC1 and HfCI show a significantly higher rate of precipitation of ZrC14,- 2diarsine; conjecture on this observation is tempting; but as the authors note in the absence of detailed kinetic data speculation is ~ n r e a l i s t i c .~ ~ Extension of the work with diarsine to the analogous phosphineaO suggests that high-co- ordinate structures can also be achieved with the diphosphine and the early transition elements. With titanium tetrachloride and the diphosphine a very reactive 1 2 complex is formed which is isomorphous with dodecahedral TiCI, 2diarsine. The authors conclude that stereochemistry rather than intrinsic elec- tronic properties associated with the donor atom is a dominating parameter.=O Anhydrous tetrakisnitrate complexes of tin@) titaiium(Iv) and zirconium(1v) have been reported. Preliminary X-ray analysis of Ti(NOa4 suggests D2d dodecahedral geometry.341,342 Thorium (IV) complexes with pyrocatechol-3,5- disulphonic acid may be eight (or higher) co-ordinate species.343 Differentiation between the lanthanides and actinides (periodic analogues) is apparent in the stability of the complexes of americium to fermium with 1,2-diaminocyclohexane- tetra-acetic acid.This suggests diminution of the chelate functionality and in- creasing hydration with increasing atomic number possibly owing to the sharp decrease in atomic radius going from americium to fermium.344 Lead(1v) acetate is octaco-ordinate with basic dodecahedral geometry although 336 S. J. Lippard J. Amer. Chem. SOC. 1966 88 4300. a37 G. Strahs and R. E. Dickerson Spore Newsletter 1965 2 30. 338 V. A. Golovnya A. K. Molodkin and V. N. Tverdokhlebov Rum. J. Inorg. Chem. 1965 10 1195. 339 R. J. H. Clark W. Errington J. Lewis and R. S. Nyholm J . Chem. Soc. (A) 1966 989. 340 R. J. H. Clark R. H. U. Negrotti and R.S. Nyholm Chem. Comm. 1966,486. 341 C. C. Addison C. D. Garner W. P. Simpson D. Sutton and S. C. Wallwork Proc. Chem. SOC. 1964 367. 342 B. 0. Field and C. J. Hardy Proc. Chem. SOC. 1962 76. 343 Y. Murakami and A. E. Martell Bull. Chem. SOC. Japan 1966,39 1077. 344 R. D. Baybarz J. Inorg. Nuclear Chem. 1966 28 1055. 187 Quarterly Reviews there is some distortion.345 Carbonate complexes of the lanthanides may also be octaco-ordinate ; complexes of the type M(CO,),& have been describeds46 (M = Cd Pr Nd Sm and Eul*I). Tetrakis-oxalates of Cew and PrN are possibly o~taco-ordinate.~~ Further evidence of ligand lability in octaco-ordinate acetylacetonate-type derivatives comes from studies by Adams and L a r ~ e n ~ ~ * and by Cotton Legzdins and L i p ~ a r d . ~ ~ In mixtures of the neutral Zr HfW or anionic YIII tetrakis- complexes with a given metal ion ligand exchange with various substituted acetylacetonates is rapid and non-random.Mixed complexes are preferentially formed. Equilibrium constants for some of these systems have been e v a l ~ a t e d . ~ ~ ~ ~ ~ A number of new lanthanide and actinide periodic lattices and complex salts have been reported all of which may have octaco-ordinate metal atoms. These include PaBr (which is isomorphous with dodecahedra1 ThBr,) Pax,,- salts PaX4,2(and 4)-ligand complexe~,3~~ an isomorphous series of NpF,- PI,#,- AmF6- and CmF,- saltss1 (with the compact fluoride ion the co-ordination number may exceed eight) hexachloro- and hexabromo-complexes of Th U and Nprv,s2 as well as chloro-complexes of the lanthanides e.g.Na,EuCI, Na,EuCl, Na,HoCl, and Na,ErC1,.55S Probable octaco-ordinate thorium(rv) uranyl and neptunyl complexes recently described are UO,(NO,), Th(NO,) (see below for possible decaco- ordination),354 Np(N03),,2H20 and related degradation phosphite uranyl complexes e.g. U0,(N03),[OP(C6H,)3],356 (see octaco-ordination). Preliminary X-ray analysis of U02(02)3 diffraction data indicates near-hexagonal bipyramidal geometry with coplanar 0 groups in the hexagonal plane with a linear UO group [a long uranyl-oxygen distance of 1.88 A (normally about 1.77 A) and a 2.27 A U-0 peroxide ~eparation1.3~~ In UO,(NO,),(H,O) and in UO,(NO,),(H,O),,H,O the uranyl grouping retains the characteristic linear array. The two nitrate groups are trans and bidentate and significantly bent from the equatorial UO plane to give essentially a chain conformation.The uranyl U-0 separations are long (1.9 A) the U-O(H,O) separations are 2-0& and the U-O(NO,) separations are 2.1 and 2.4 A in the dihydrate and trihydrate respec ti vel y .35 8335 345 B. Kamenar Acta Cryv. 1963 16 A 34. 346 H. S. Sherry and J. A. Marinsky Inorg. Chem. 1963,2,957. 347 S . W. Pajahoff Monatsh. 1963 94 526. s4* A. C. Adams and E. M. Larsen Inorg. Chem. 1966,5 814. 349 F. A. Cotton P. Legzdins and S. J. Lippard to be published. 350 D. Brown and P. J. Jones Chem. Comm. 1966,279. 351 T. K. Keenan Inorg. Nuclear Chem. Letters 1966 2 153. 352 D. Brown J. Chem. SOC. (A) 1966 766. 353 B. G. Korshunov and D. V. Drobot Russ. J Inorg. Chem. 1965,10 1256. 354 J. R. Ferraro and A. Walker J. Chem. Phys. 1966 45 550.355 J. B. Laidler J. Chem. SOC. (A) 1966 780. 356 F. A. Hart and J. E. Newberry J. Inorg. Nuclear Chem. 1966 28 1334. 357 N. W. Allcock Chem. Comm. 1966 536. a5* V. M. Vdovenko E. V. Stroganov A. P. Sokolov and G. Lungu Radiokhimiya 1962 4 59. 359 V. M. Vdovenko E. V. Stroganov and A. P. Sokolov Radiokhimiya 1963,5,97. 188 Muetterties and Wright Analysis of the three-dimensional X-ray measurements for the tetrapheny- larsonium salt of Co(0,CCF3),2- established near S4 symmetry. The co-ordina- tion polyhedron more nearly approximates to a tetrahedron than a dodecahedron and the authors prefer a characterisation of this anion as a tetrahedral Co2+ complex.ssO The high-temperature tetragonal form of ZrO in contrast to the heptaco- ordinate zirconium in normal ZrO has octaco-ordinate zirconium atoms with near dodecahedra1 geometry (four Zr-0 separations at 2-07 8 and four at 2.45 A)?a1 In cluster chemistry the niobium atom cluster in the fl-Nb3Br periodic lattice is analogous to the Re,ClIz3- cluster but additionally has a centrally located bromine atom to give effectively octaco-ordinate niobium Specific directions for the preparation of metal-atom cores for the square antiprism and dodecahedron models for Fischer-Taylor-Hirschfelder atomic models have been recently reported by Homeier and Larsen.363 Nonaco-ordination.-Dimethyl sulphoxide complexes of thorium(1v) chlorides of the type ThCl,,S(CH&$O are monomeric in acetonitrile and non-conducting in nitromethane suggestive of a nonaco-ordinate complex provided these basic solvents are not entering the co-ordination sphere.The uranium complex UCl4,3(CH,),SO is on this count heptaco-ordinate. Although uranium(W) is smaller than thorium(Iv) it is somewhat surprising that there appears to be a significant drop in co-ordination numbers in these related systems.364 Perhaps the authors’364 suggestion of solvent participation is the answer and actual co- ordination numbers in the two systems are the same or differ by only one. Complexes of lanthanide nitrates with dimethyl formamide M(DMF),(NO,), have been prepared and nonaco-ordination suggested.365 In nitromethane the complexes do not behave as electrolytes but in dimethyl formamide there is a displacement reaction apparently to give the cation [M(DMF),(NO,),]+ which also may be nonaco-ordinate provided that the nitrate groups are bidentate.Similarly the oxonitrate anionic complexes of niobium and tantalum OM(NO,) have been isolated as tetramethylammonium salts and may be nonaco-ordinate.ss6 The volatile casium salt of Y(CF3COCHCOCF3)4- noted above shows a parent mass ion corresponding to CSY[CF,COCHCOCF,]~+ although this is very weak.33s A strong peak assignable to CsY[CF3CO-CHCOCFJ3+ suggests that the caesium is very tightly bound. The site of binding is unknown and it seems unlikely that there is a close Cs-Y interaction to give nonaco- ordination. Lippard suggests that the caesium ion may either be bound to the 360 J. A. Bergman jun. and F. A. Cotton Inorg. Chem. 1966,5 1420. 361 G. Teufer Acta Cryst. 1962 15 1187. 362 A. Simon and H. G. von Schnering J. Less Common Metals 1966 11 31. 363 E. Homeier and E.M. Larsen J. Chem. Educ. 1966,43 376. 364 K. W. Bagnall D. Brown P. J. Jones and J. G. H. du Preez J. Chem. SOC. (A) 1966,737. 365 S. S. Krishnamurthy and S. Soundarajan J . Inorg. Nuclear Chem. 1966 28 1689. 366 D. Brown and P. J. Jones J. Chern. SOC. (A) 1966,733. 189 Quarterly Reviews 'sheath of negative fluorine atoms which surround the yttrium atom' or may be bound to the methylene (= CH-) carbon atom. Nonaco-ordination has been suggested for the lanthanide(rI1) complexes with 1 -hydroxycyclopentanecarboxylic acid.367 In optical studies of europium and neodymium complexes with nitrilotriacetic ethylenediaminetetra-acetic and diethylenetriaminopenta-acetic acids the species have been analysed in terms of metal co-ordination numbers of six to eight.3s8 These estimates seem abnormally low and octa- to at least nona-co-ordination is more reasonable for these chelate derivatives of these large tervalent ions.The structural investigation of LaF by Zalkin Templeton and HopkinsMs have characterised the lanthanum polyhedron as irregular with a nonaco- ordinate lanthanum atom. Man~mann~~O also reported an X-ray analysis of LaF and gave the same space-group P&l as Zalkin et al. However Mansmann describes the lanthanum co-ordination polyhedron as did Ofteda13'l who had the incorrect space group as undecaco-ordinate with nine LaF separations of 2.42-2.61 8 and two at 3.10 A. These separations agree closely with those found by Zalkin et al. We prefer the characterisation as nonaco-ordinate since a separation of 3-01 8 is greater than the sum of the van der Waals radii and these two long La-F separations may be only weakly bonding.Lanthanide fluoride complexes of the type NaMF where M = YIIl and lanthanide("I) excepting Yb and Lu are isomorphous with NaLaF,372 which has nonaco- ordinate lanthanum atoms.70 The oxyfluoride Pa20F8 is isomorphous with U2F (body-centred cubic) and may have a nonaco-ordinate protactinium at0m.3~~ The Pa20CI analogues may be nonaco-ordinate but steric factors may effect a significant change from the fluoride lattice and in the protactinium co- ordination ~phere.3~~ Analysis of the X-ray diffraction data375 for a single-crystal of NdTe (periodic lattice) shows the neodymium to be surrounded by nine tellurium atoms with the basic geometry of monocapped square antiprism (Nb-Te distances are 3.21 (4 Te) 3.36 (1 Te) and 3.35 A (4 Te).In new cluster chemistry developments B~rbank,~ has analysed the Ta6C1,,,7H20 structure by X-ray diffraction. The crystals are trigonal with a probable space-group P31 rn. Burbank describes the structure as distorted and subject to faulting. The basic tantalum co-ordination polyhedron is the previously described Ta,C1,22+ unit with two chlorine atoms and four water molecules filling the terminal tantalum sites to give nonaco-ordination. The Ta 'octa- hedron' is elongated on the four-fold axis. Other new and presumably costruc- 367 J. E. Powell and D. L. G. Rowlands Inurg. Chem. 1966 5 819. 388 T. V. Ternovaya and N. A. Kostromina Russ. J. Znorg. Chem. 1965 10 1100. 369 A. Zalkin D. H. Templeton and T. E. Hopkins Inorg. Chem. 1966 5 1466. 370 V. M.Mansmann 2. Krist. 1965 122 375. 371 I. Oftedal 2. phys. Chem. 1931 B 13 190. 372 R. E. Thoma H. Insley and G. M. Hebert Inorg. Chem. 1966,5 1222. 373 L. Stein Znorg. Chem. 1964 3 995. 374 D. Brown and P. J. Jones J. Chem. SOC. (A) 1966 874. 375 B. K. Norling and H. Steinfink Znorg. Chem. 1966 5 1488. 376 R. D. Burbank Inorg. Chem. 1966 5 1491. 190 Muetterties and Wright tural clusters of this type are W,Br,* W6Br16,377 Nb,Br14,8H,0 Ta,Br,4,8H20S78 Nb,C116,3C2H,0H379 and Nb,Fl,.380 Deca- and Dodeca-co-ordination.-Tetrakisacetatouranium(w) a fibrous mole- cule has been characteri~ed~~l by X-ray analysis as consisting of linear arrays of bicapped square antiprisms (with co-ordinated oxygen atoms serving as bridging groups). This type of structure may be found in other very large quadrivalent metal acetates or nitrates e.g.Th(N03)4 although the analogous lead(1v) derivative has a molecular lattice with octaco-ordination. Brown and Jones382 have prepared a series of nitrate complexes with protac- tinium. Depending upon the functionality of the nitrate ligands co-ordination may range from ten in CH,CN,Pa(NO&OPa(No,),NCCH to twelve in CsPa(NOJ,. In the ytterbium-antimony phase diagram YbSb is one of the most stable phases present. Preliminary X-ray analysis383 of this phase indicates that the ytterbium atom is decaco-ordinate with a near-square antiprismatic arrangement with eight antimony atoms at distances of 3.19-3-30 A with two more antimony atoms at relatively long distances of 3.57 A but these two are over one of the square faces. This significant departure from idealised symmetrically bicapped square antiprismatic geometry is not disconcerting in that YbSb is a periodic lattice.Line Notation for Co-ordination Compounds.-McDonnell and Pa~ternack,3~~ using a reference-structure approach developed a notation system for represen- tation of inorganic structures in co-ordination compounds based on considera- tion of compatibility with systems formulated for organic structures.385387 Their suggested notation system is based384 on letters and closely follows the H a y ~ a r d ~ ~ rules. The lettering system differs from ours only in the matter of where the numbering starts in successive planes reflecting the minor conflict between I.U.P.A.C. recommendations and the usage in the United States partic- ularly for polyhedral boranes.McDonnell and Pasternack’s system is not compre- hensive for established co-ordination polyhedra and is arbitrary if not mislead- 377 H. Schafer and R. Siepmann J. Less Common Metals 1966 11 76. 378 H. Schafer and B. Spreckelmeyer J. Less Common Metals 1966 11 73. 379 B. Spreckelmeyer and H. Schafer J . Less Common Metals 1966 11 74. 380 H. Schafer H. G. von Schnering A. Simon D. Giegling D. Bauer R. Siepmann and B. Spreckelmeyer J. Less Common Metals 1965 10 154. 381 I. Jelenic D. GrdeniC and A. Bezjak Acra Cryst. 1964 17 758. 382 D. Brown and P. J. Jones J. Chern. SOC. (A) 1966 733. 383 R. Wang R. Bodnar and H. Steinfink Inorg. Chem. 1966 5 1468. 384 P. M. McDonnell and R. F. Pasternack J . Chem. SOC. 1965 56. 385 ‘Rules for I.U.P.A.C. Notation for Organic Compounds’ John Wiley and Sons Inc.New York N.Y. 1961. 386 H. W. Hayward ‘A New Sequential Enumeration and Line Formula Notation System for Organic Compounds’ Patent Office Research and Development Report No. 21 Depart- ment of Commerce Washington D.C. 1961. 387 W. J. Wiswesser ‘A Line-Formula Chemical Notation’ Thomas Y. Crowell Co. New York N.Y. 1954. 191 Quarterly Reviews ing in the capital-letter notation for the specific polyhedron in any given polyatom system ‘The symmetry designators although assigned somewhat arbitrarily reflect frequency of occurrence of individual configuration.’ The latter qualifi- cation was not correctly applied to the seven- and eight-atom polyhedra con- sidered by them. Polyhedral Holes.-Many minerals and also synthetic periodic lattices have a network of holes that may be described as polyhedral.These polyhedra for the most part do not correlate with the polyhedra that dominate in high-co-ordina- tion structures and in polyhedral boranes. The pores or holes in silicate minerals may be generated by the incorporation or inclusion of polyhedral ions during the condensation polymerisation of the silicate chains or nets.38s The types of polyhedral holes include the cube octahedron in faujasite and s0dalite,3~~ hexagonal prisms in certain and very large ones such as the hexacosahedron found in Linde Sieve A,389,392 and the triacontrahedral (30) and dotetracontahedral (42) cavities in synthetic ana1~ites.s~~ New Structural Data.-The crystal structure of Cs+Y(CF,COCHCOCF,),- which was previously is c0mplete.3~~ The yttrium atom is co- ordinated to eight oxygens (Y-0 = 2-33 A) in a dodecahedral co&guration with the chelate groups spanning “g” edges (see Figure 25) rather than the “m” edges resulting in the first known example of idealized D symmetry for the dodecahedral configuration.The D dodecahedral structure could possibly result from a slight distortion of the D square antiprism. A cause of such distortion might be the observed close association of the Cs+ cation with the Y(CF,COCHCOCF,),- anion in the solid state.394 The previously known dodecahedral configurations have D2d symmetIy (m edge span). Based on a crystal structure determination the ligands in NH,[Pr(CH,.SCCHCOCF,),] are reported 395 to also span “g” edges but the symmetry is lowered by the reversed attachment of one of the asymmetric ligands. A recent three dimensional single crystal study of [Fe(C,H,N),]2+[Fe4(CO),,]2- establishes the configuration of the [Fe4(C0),,I2- anion as a tetrahedron of iron atoms396 (Feb-Fea = 2-58 A Feb-Feb = 2.50 A) with the same general structure as Co,(CO), (Figure 9) and the thirteenth CO group on the three-fold axis triply bridging the three basal iron atoms aver.Feb-COb = 2-00 A. The three carbonyl groups in the basal plane of the iron atoms are not symmetrically placed between the two iron atoms (Fe-CO = 2.24 2.28 2-33; Fe-VO = 1.80 388 R. M. Barrer Chem. in Britain 1966 2 380. 389 R. M. Barrer Endeavour 1964 23 122. 390 L. S. Dent and J. V. Smith Nature 1965 181 1794. 391 L. Broussard and D. P. Shoemaker J. Amer. Chem. SOC. 1960,82 1041. 392 W. M. Meier and G. T. Kokotailo 2. Krist. 1965 121 14.393 R. M. Barrer and I. S. Kerr J . Chem. SOC. 1963 434. 394 S. J. Lippard F. A. Cotton and P. Legzdins J. Amer. Chem. SOC. 1966,88,5930. 395 M. Cefola W. Hamilton R. Lalancette and S. Laplaca reference 394 footnote 11. 396 R. J. Doedens and L. F. Dahl J . Amer. Chem. SOC. 1966,88,4847. 192 Muetterties and Wright 1-81 1-85) and do not bridge as in CO,(CO)~~ but a weak interaction exists. Considering the nine carbonyls in the Fe,CO fragments as terminal (aver. Fe-CO = 1-72 A) the basal iron atoms are seven co-ordinate with a distorted version of a 4 3 orientation of ligands. The apical iron atom is six-co-ordinate. The results of the structure determination of [(C,H5)aAs]2Re3C11~97 and Cs2Re3Br11398 are now published. The basic structure is the R%X,23- structure (Figure 16) with one terminal halogen atom missing.A significant structural change is the shortening of the two Re-Re bonds to the halogen deficient rhenium atom to 2.435 A vs. 2.483 A in the chloride and 2.43 A vs. 2-49 A in the bromide. The two terminal halogen rhenium bonds on the halogen-deficient rhenium atom are also significantly shorter than the others. In the Cs2Re3Brll structure the Re-Br bonds on the bromine-deficient rhenium are bent towards each other 133" vs. 159" for the angle subtended by a pair of off-plane bromine atoms at a non-bromine deficient rhenium. on IF is presented using the rotating sector microphotometer method; the previous study employed the visual methodF6 Preliminary analysis indicates that the structure is close to a penta- gonal bi-pyramid and that the axial bonds are shorter than the equatorial bonds.A three dimensional single crystal X-ray study of anhydrous Ti(N03),400 established co-ordination of eight oxygen atoms to the titanium atoms from four bidentate nitrate groups. The basic geometry approximates the dode- cahedron (DZd symmetry). The nitrate groups in K3[Hg(N02)4]N03401 are bidentate by a three-dimensional crystal structure analysis and arranged around the mercury atom in roughly a tetrahedral orientation with the eight oxygen atoms 2.4 A from the mercury. The structure of a-Gd,S3 was solved by a three dimensional single crystal X-ray study402 and is found to contain two nonequivalent gadolinium atoms. One gadolinium site is co-ordinated to seven sulphur atoms at an average distance of 2.81 A and the co-ordination polyhedron is a slightly distorted square-capped trigonal prism.The co-ordination polyhedron of the other gadolinium site is a bicapped trigonal prism with eight sulphur atoms at an average distance of 2.90 A. Within each polyhedron the gadolinium-sulphur distances are approxi- mately equivalent. This rare-earth a-sesquisulphide structure may exist for Ce to Dy?03 A neutron powder diffraction study of HoDtM indicates a structure in which holmium has nine nearest hydrogen neighbours at a distance of 2-10 to 2-29 A and two hydrogen neighbours at 2.48 A. The holmium atom is nine co-ordinate and the co-ordination polyhedron is the tricapped trigonal prism with distortion 897 B. R. Penfold and W. T. Robinson Inorg. Chem. 1966,5,1758. s98 M. Elder and B. R. Penfold Inorg. Chern. 1966,5 1763. s99 H.B. Thompson jun. and L. S. Bartell Trans. Amer. Cryst. Ass. 1966,2 190. 400 C. D. Garner and S. C. Wallwork J. Chem. SOC. (A) 1966,1496. 401 D. Hall and R. V. Holland Proc. Chem. Soc. 1963,204. 402 C. T. Prewitt and A. W. Sleight Amer. Cryst. ASSOC. Jan. 25 (1967) Atlanta Georgia. Cryst. 1965,19 14. 404 D. Mansmann and W. E. Wallace J . Phys. (France) 1964.25,454. A second electron diffraction J. F. Lahaut M. Gutlard M. Patrie M. P. Pardo S. M. Golabi and L. Domange Acta. 193 Quarterly Reviews in the capping so as to lower the point group symnietry about holmium to C,. The other hexagonal rare-earth trihydrides SmH, GdH, TbH, DyH, ErH, TmH, LuH, HoH and YH probably have the same crystal structure as We thank Professor J. L. Hoard for critical comments. HOD^ 194

ISSN:0009-2681    

DOI:10.1039/QR9672100109

出版商:RSC    

年代:1967

数据来源: RSC

摘要:

Volume 21 1967 Cumulative Indexes for Authors and Titles are to be found at the end of Volume 20. Copies may be obtained from The Publications Sales Officer The Chemical Society. Burlington House London W.l price 5/-d. post free. Quarterly Reviews INDEX OF AUTHORS Abraham E. P. 231 Ashby E. C. 259 Bransden B. H. 474 Bruce J. M. 405 Buckingliam A. D. 195 Davidson R. S. 249 Evans U. R. 29 Fensham P. J. 507 Howe A. T. 507 Lambert J. D. 67 Lee J. B. 429 Luz Z. 458 McLennan D. J. 490 Muetterties E. L. 109 Orr B. J. 195 Parker W. 331 Pelletier S. W. 525 Penzer G. R. 43 Radda G. K. 43 Ramage R. 331 Riddell F. G. 364 Roberts J. S. 331 Robinson D. L. 314 Silver B. L. 458 Sklarz B. 3 Theobald D. W. 314 Uff B. C. 429 Waley S . G. 379 Walker D. C. 79 Weatherston J. 287 Wehry E. L. 213 Wright C.M. 109 INDEX OF TITLES Alkaloids the chemistry of the Hydrogen abstraction in the liquid C,,-diterpene 525 phase by free radicals 249 Arthropod defensive substances the Isoalloxazines (Flavines) the chemistry chemistry of 287 and biological function of 43 Biogenesis sesquiterpene 331 Light-induced reactions of quinones Carbanion mechanism of olefin- 405 forming elimination 490 Liquid phase hydrogen abstraction 231 in the by free radicals 249 Cephalosporin C Group stances Conformational analysis heterocyi: Nuclear and electronic spin resonance 364 chemical applications of oxygen-1 7 Co-ordination number molecular 458 109 Olefin-forming elimination carbanion 490 polyhedra of high Cyclohexane chemistry conforma- mechanism Of tional abnormalities in 314 Organic chemistry of Periodates 3 Organic reactions involving electro- the C, 525 philic oxygen 429 Electron the hydrated 79 Oxygen organic reactions involving Electron spin resonance chemical electrophilic 429 and 458 resonance chemical applications of Electronic properties of binary com- 458 metals 507 Photochemical behaviour of transition- Electrophilic oxygen organic reactions metal complexes 21 3 involving 429 Sesquiterpere biogenesis 33 1 Elementary particles 474 Rusting the mechanism of 29 Enzyme action mechanism of 379 Quinones light-induced reactions of Energy transfer vibration-vibration 405 in gaseous collisions 67 Transition-metal complexes photo- Free radicals hydrogen abstraction in &emical behaviour of 21 3 the liquid phase by 249 Transition metals electronic proper- Grignard reagents.Compositions and ties of binary compounds of the mechanisms of reaction 259 first-row 507 Heterocyclic conformational analysis Vibration-vibration energy transfer 364 in gaseous collisions 67 ChemistrY of arthropod defensive sub- Molecular hyperpolarisabilities 195 Diterpene alkaloids the chemistry of applications of oxygen-1 7 nuclear Ox~gen-17 nuclear and electron spin pounds of the fist-row transition Periodates9 organic chemistry of 3 550

ISSN:0009-2681    

DOI:10.1039/QR9672100549

出版商:RSC    

年代:1967

数据来源: RSC